Principles of Environmental Chemistry

Roy M Harrison School of Geography, Earth and Environmental Sciences, University of Birmingham, Birmingham, UK

ISBN-13: 978-0-85404-371-2

A catalogue record for this book is available from the British Library r The Royal Society of Chemistry 2007 All rights reserved Apart from fair dealing for the purposes of research for non-commercial purposes or for private study, criticism or review, as permitted under the Copyright, Designs and Patents Act 1988 and the Copyright and Related Rights Regulations 2003, this publication may not be reproduced, stored or transmitted, in any form or by any means, without the prior permission in writing of The Royal Society of Chemistry, or in the case of reproduction in accordance with the terms of licences issued by the Copyright Licensing Agency in the UK, or in accordance with the terms of the licences issued by the appropriate Reproduction Rights Organization outside the UK. Enquiries concerning reproduction outside the terms stated here should be sent to The Royal Society of Chemistry at the address printed on this page. Published by The Royal Society of Chemistry, Thomas Graham House, Science Park, Milton Road, Cambridge CB4 0WF, UK Registered Charity Number 207890 For further information see our web site at www.rsc.org Typeset by Macmillan India Ltd, Bangalore, India Printed by Henry Ling Ltd, Dorchester, Dorset, UK

Preface While this book is in its ﬁrst edition, it nonetheless has a lengthy pedigree, which derives from a book entitled Understanding Our Environment: An Introduction to Environmental Chemistry and Pollution, which ran to three editions, the last of which was published in 1999. Understanding Our Environment has proved very popular as a student textbook, but changes in the way that the subject is taught had necessitated its splitting into two separate books. When Understanding Our Environment was ﬁrst published, neither environmental chemistry nor pollution was taught in many universities, and most of those courses which existed were relatively rudimentary. In many cases, no clear distinction was drawn between environmental chemistry and pollution and the two were taught largely hand in hand. Nowadays, the subjects are taught in far more institutions and in a far more sophisticated way. There is consequently a need to reﬂect these changes in what would have been the fourth edition of Understanding Our Environment, and after discussion with contributors to the third edition and with the Royal Society of Chemistry, it was decided to divide the former book into two and create new books under the titles respectively of An Introduction to Pollution Science and Principles of Environmental Chemistry. Because of the authoritative status of the authors of Understanding Our Environment and highly positive feedback which we had received on the book, it was decided to retain the existing chapters where possible while updating the new structure to enhance them through the inclusion of further chapters. This division of the earlier book into two new titles is designed to accommodate the needs of what are now two rather separate markets. An Introduction to Pollution Science is designed for courses within degrees in environmental sciences, environmental studies and related areas including taught postgraduate courses, which are not embedded in a speciﬁc physical science or life science discipline such as chemistry, v

Preface

physics or biology. The level of basic scientiﬁc knowledge assumed of the reader is therefore only that of the generalist and the book should be accessible to a very wide readership including those outside of the academic world wishing to acquire a broadly based knowledge of pollution phenomena. The second title, Principles of Environmental Chemistry assumes a significant knowledge of chemistry and is aimed far more at courses on environmental chemistry which are embedded within chemistry degree courses. The book will therefore be suitable for students taking second or third year option courses in environmental chemistry or those taking specialised Masters’ courses, having studied the chemical sciences at ﬁrst-degree level. In this volume I have been fortunate to retain the services of a number of authors from Understanding Our Environment. The approach has been to update chapters from that book where possible, although some of the new authors have decided to take a completely different approach. The book initially deals with the atmosphere, freshwaters, the oceans and the solid earth as separate compartments. There are certain common crosscutting features such as non-ideal solution chemistry, and where possible these are dealt with in detail where they ﬁrst occur, with suitable cross-referencing when they re-appear at later points. Chemicals in the environment do not respect compartmental boundaries, and indeed many important phenomena occur as a result of transfers between compartments. The book therefore contains subsequent chapters on environmental organic chemistry, which emphasises the complex behaviour of persistent organic pollutants, and on biogeochemical cycling of pollutants, including major processes affecting both organic and inorganic chemical species. I am grateful to the authors for making available their great depth and breadth of experience to the production of this book and for tolerating my many editorial quibbles. I believe that their contributions have created a book of widespread appeal, which will ﬁnd many eager readers both on taught courses and in professional practice. Roy M. Harrison Birmingham, UK

Contents Chapter 1

Chapter 2

Introduction R.M. Harrison

1.1 1.2 1.3 1.4

The Environmental Sciences Environmental Chemical Processes Environmental Chemicals Units of Concentration 1.4.1 Atmospheric Chemistry 1.4.2 Soils and Waters 1.5 The Environment as a Whole References

1 3 3 5 5 6 7 7

Chemistry of the Atmosphere P.S. Monks

2.1 2.2 2.3 2.4 2.5

2.6 2.7 2.8 2.9

Introduction Sources of Trace Gases in the Atmosphere Initiation of Photochemistry by Light Tropospheric Chemistry Tropospheric Oxidation Chemistry 2.5.1 Nitrogen Oxides and the Photostationary State 2.5.2 Production and Destruction of Ozone 2.5.3 Role of Hydrocarbons 2.5.4 Urban Chemistry Night-Time Oxidation Chemistry Ozone-Alkene Chemistry Sulfur Chemistry Halogen Chemistry 2.9.1 Tropospheric Halogens and Catalytic Destruction of Ozone vii

8 10 11 17 20 26 28 35 37 40 46 46 51 56

viii

Chapter 3

Contents 2.10 Stratospheric Chemistry 2.10.1 The Antarctic Ozone Hole 2.11 Summary Questions References

58 63 72 72 76

Chemistry of Freshwaters M.C. Graham and J.G. Farmer

3.1 3.2

Introduction Fundamentals of Aquatic Chemistry 3.2.1 Introduction 3.2.2 Dissolution/Precipitation Reactions 3.2.3 Complexation Reactions in Freshwaters 3.2.4 Species Distribution in Freshwaters 3.2.5 Modelling Aquatic Systems 3.3 Case Studies 3.3.1 Acidiﬁcation 3.3.2 Metals and Metalloids in Water 3.3.3 Historical Pollution Records and Perturbatory Processes in Lakes 3.3.4 Nutrients in Water and Sediments 3.3.5 Organic Matter and Organic Chemicals in Water Questions and Problems Further Reading References Chapter 4

80 82 82 91 94 97 121 122 122 130 139 145 150 157 159 159

Chemistry of the Oceans S.J. de Mora

170

4.1

170 170 173 177 179 182 182 184 199 201 204 210

Introduction 4.1.1 The Ocean as a Biogeochemical Environment 4.1.2 Properties of Water and Seawater 4.1.3 Salinity Concepts 4.1.4 Oceanic Circulation 4.2 Seawater Composition and Chemistry 4.2.1 Major Constituents 4.2.2 Dissolved Gases 4.2.3 Nutrients 4.2.4 Trace Elements 4.2.5 Physico-Chemical Speciation 4.3 Suspended Particles and Marine Sediments

Contents 4.3.1

Description of Sediments and Sedimentary Components 4.3.2 Surface Chemistry of Particles 4.3.3 Diagenesis 4.4 Physical and Chemical Processes in Estuaries 4.5 Marine Contamination and Pollution 4.5.1 Oil Slicks 4.5.2 Plastic Debris 4.5.3 Tributyltin Questions References Chapter 5

Chapter 6

210 213 218 219 223 223 226 228 230 231

The Chemistry of the Solid Earth I.D. Pulford

234

5.1 5.2

Introduction Mineral Components of Soil 5.2.1 Inputs 5.2.2 Primary Minerals 5.2.3 Secondary Minerals 5.2.4 Weathering Processes (See also Chapter 3) 5.3 Organic Components of Soil 5.4 Soil pH And Redox Potential 5.4.1 pH and Buﬀering 5.4.2 Soil Acidity 5.4.3 Soil Alkalinity 5.4.4 Inﬂuence of pH on Soils 5.4.5 Redox Potential 5.4.6 Reduction Processes in Soil 5.5 Chemical Reactions in Soil 5.5.1 Reactions in Soil Solution 5.5.2 Ion Exchange (Physisorption) 5.5.3 Ligand Exchange (Chemisorption) 5.5.4 Complexation/Chelation 5.5.5 Precipitation/Dissolution 5.5.6 Soil Processes Questions References

234 238 238 238 240 246 248 254 254 255 257 258 260 261 263 263 267 271 273 273 275 275 278

Environmental Organic Chemistry C.J. Halsall

279

6.1

279

Introduction

Contents 6.2

Chapter 7

The Diversity of Organic Compounds 6.2.1 Identifying Sources of Hydrocarbons 6.3 The Fate of Organic Contaminants 6.4 Chemical Partitioning 6.4.1 Important Partitioning Coeﬃcients 6.4.2 Temperature Dependence 6.4.3 Partition Maps 6.5 Chemical Transformation and Degradation 6.6 Chemical Transformation through Photochemistry 6.6.1 Light Absorption and the Beer-Lambert Law 6.6.2 Photolysis in Aqueous Systems 6.6.3 Photochemistry of Brominated Flame Retardants (BFRs) 6.7 Conclusions 6.8 Questions References

280 282 284 284 286 292 295 299 301 302 303

Biogeochemical Cycling of Chemicals R.M. Harrison

314

7.1

314 316 317

Introduction: Biogeochemical Cycling 7.1.1 Environmental Reservoirs 7.1.2 Lifetimes 7.2 Rates of Transfer between Environmental Compartments 7.2.1 Air–Land Exchange 7.2.2 Air–Sea Exchange 7.3 Transfer in Aquatic Systems 7.4 Biogeochemical Cycles 7.4.1 Case Study 1: The Biogeochemical Cycle of Nitrogen 7.4.2 Case Study 2: Aspects of Biogeochemical Cycle of Lead 7.5 Behaviour of Long-Lived Organic Chemicals in the Environment Questions References Glossary Subject Index

305 309 310 310

321 321 324 330 333 335 335 340 344 345 347 354

CHAPTER 1

Introduction ROY M. HARRISON Division of Environmental Health and Risk Management, School of Geography, Earth and Environmental Sciences, University of Birmingham, Edgbaston, B15 2TT, Birmingham, UK

1.1 THE ENVIRONMENTAL SCIENCES It may surprise the student of today to learn that ‘the environment’ has not always been topical and indeed that environmental issues have become a matter of widespread public concern only over the past 20 years or so. Nonetheless, basic environmental science has existed as a facet of human scientiﬁc endeavour since the earliest days of scientiﬁc investigation. In the physical sciences, disciplines such as geology, geophysics, meteorology, oceanography, and hydrology, and in the life sciences, ecology, have a long and proud scientiﬁc tradition. These fundamental environmental sciences underpin our understanding of the natural world and its current-day counterpart perturbed by human activity, in which we all live. The environmental physical sciences have traditionally been concerned with individual environmental compartments. Thus, geology is centred primarily on the solid earth, meteorology on the atmosphere, oceanography upon the salt-water basins, and hydrology upon the behaviour of freshwaters. In general (but not exclusively) it has been the physical behaviour of these media which has been traditionally perceived as important. Accordingly, dynamic meteorology is concerned primarily with the physical processes responsible for atmospheric motion, and climatology with temporal and spatial patterns in physical properties of the atmosphere (temperature, rainfall, etc.). It is only more recently that chemical behaviour has been perceived as being important in many of these areas. Thus, while atmospheric chemical processes are at least as important as physical processes in many environmental problems such as stratospheric ozone depletion, the lack of chemical knowledge has been 1

Chapter 1

extremely acute as atmospheric chemistry (beyond major component ratios) only became a matter of serious scientiﬁc study in the 1950s. There are two major reasons why environmental chemistry has ﬂourished as a discipline only rather recently. Firstly, it was not previously perceived as important. If environmental chemical composition is relatively invariant in time, as it was believed to be, there is little obvious relevance to continuing research. Once, however, it is perceived that composition is changing (e.g. CO2 in the atmosphere; 137Cs in the Irish Sea) and that such changes may have consequences for humankind, the relevance becomes obvious. The idea that using an aerosol spray in your home might damage the stratosphere, although obvious to us today, would stretch the credulity of someone unaccustomed to the concept. Secondly, the rate of advance has in many instances been limited by the available technology. Thus, for example, it was only in the 1960s that sensitive reliable instrumentation became widely available for measurement of trace concentrations of metals in the environment. This led to a massive expansion in research in this ﬁeld and a substantial downward revision of agreed typical concentration levels due to improved methodology in analysis. It was only as a result of James Lovelock’s invention of the electron capture detector that CFCs were recognised as minor atmospheric constituents and it became possible to monitor increases in their concentrations (see Table 1). The table exempliﬁes the sensitivity of analysis required since concentrations are at the ppt level (1 ppt is one part in 1012 by volume in the atmosphere) as well as the substantial increasing trends in atmospheric halocarbon concentrations, as measured up to 1990. The implementation of the Montreal Protocol, which requires controls on production of CFCs and some other halocarbons, has led to a slowing and even a reversal of annual concentration trends since 1992 (see Table 1). Table 1

Atmospheric halocarbon concentrations and trendsa Concentration (ppt)

Annual change (ppt)

Halocarbon

Pre-industrial

2000

To 1990

1999–2000

Lifetime (years)

CCl3F (CFC-11) CCl2F2 (CFC-12) CClF3 (CFC-113) C2Cl2F4 (CFC-113) C2Cl2F4 (CFC-114) C2ClF5 (CFC-115) CCl4 CH3CCl3

0 0 0 0 0 0 0 0

261 543 3.5 82 16.5 8.1 96.1 45.4

þ9.5 þ16.5

1.1 þ2.3

þ4–5

0.35

þ2.0 þ6.0

þ0.16 0.94 8.7

50 102 400 85 300 1700 42 4.9

a Data from: World Meteorological Organization, Scientiﬁc Assessment of Ozone Depletion: 2002, WHO, Geneva, 2002.

Introduction

1.2 ENVIRONMENTAL CHEMICAL PROCESSES The chemical reactions affecting trace gases in the atmosphere generally have quite significant activation energies and thus occur on a timescale of minutes, days, weeks, or years. Consequently, the change to such chemicals is determined by the rates of their reactions and atmospheric chemistry is intimately concerned with the study of reactions kinetics. On the other hand, some processes in aquatic systems have very low activation energies and reactions occur extremely rapidly. In such circ*mstances, provided there is good mixing, the chemical state of matter may be determined far more by the thermodynamic properties of the system than by the rates of chemical processes and therefore chemical kinetics. The environment contains many trace substances at a wide range of concentrations and under different temperature and pressure conditions. At very high temperatures such as can occur at depth in the solid earth, thermodynamics may also prove important in determining, for example, the release of trace gases from volcanic magma. Thus, the study of environmental chemistry requires a basic knowledge of both chemical thermodynamics and chemical kinetics and an appreciation of why one or other is important under particular circ*mstances. As a broad generalisation it may be seen that much of the chapter on atmospheric chemistry is dependent on knowledge of reaction rates and underpinned by chemical kinetics, whereas the chapters on freshwater and ocean chemistry and the aqueous aspects of the soils are very much concerned with equilibrium processes and hence chemical thermodynamics. It should not however be assumed that these generalisations are universally true. For example, the breakdown of persistent organic pollutants in the aquatic environment is determined largely by chemical kinetics, although the partitioning of such substances between different environmental media (air, water, soil) is determined primarily by their thermodynamic properties and to a lesser degree by their rates of transfer. 1.3 ENVIRONMENTAL CHEMICALS This book is not concerned explicitly with chemicals as pollutants. This is a topic covered by a companion volume on Pollution Science. This book, however, is nonetheless highly relevant to the understanding of chemical pollution phenomena. The major areas of coverage are as follows: (i) The chemistry of freshwaters. Freshwaters comprise three different major components. The ﬁrst is the water itself, which inevitably contains dissolved substances, both inorganic and organic. Its properties are to a very significant degree determined by the

Chapter 1

inorganic solutes, and particularly those which determine its hardness and alkalinity. The second component is suspended sediment, also referred to as suspended solids. These are particles, which are sufﬁciently small to remain suspended with the water column for significant periods of time where they provide a surface onto which dissolved substances may deposit or from which material may dissolve. The third major component of the system is the bottom sediment. This is an accumulation of particles and associated pore water, which has deposited out of the water column onto the bed of the stream, river, or lake. The size of the sediment grains is determined by the speed and turbulence of the water above. A fastﬂowing river will retain small particles in suspension and only large particles (sand or gravel) will remain on the bottom. In relatively stagnant lake water, however, very small particles can sediment out and join the bottom sediment. In waters of this kind, sediment accumulates over time and therefore the surface sediments in contact with the water column contain recently deposited material while the sediment at greater depths contains material deposited tens or hundreds of years previously. In the absence of significant mixing by burrowing organisms, the depth proﬁle of some chemicals within a lake bottom sediment can provide a very valuable historical record of inputs of that substance to the lake. Ingenious ways have been devised for determining the age of speciﬁc bands of sediment. While the waters at the surface of a lake are normally in contact with the atmosphere and therefore well aerated, water at depth and the pore water within the bottom sediment may have a very poor oxygen supply and therefore become oxygen-depleted and are then referred to as anoxic or anaerobic. This can affect the behaviour of redox-active chemicals such as transition elements, and therefore the redox properties of freshwaters and their sediments are an important consideration. (ii) Salt waters. The waters of seas and oceans differ substantially from freshwaters by virtue of their very high content of dissolved inorganic material and their very great depth at some points on the globe. These facets confer properties, which although overlapping with those of freshwaters, can be quite distinct. Some inorganic components will behave quite differently in a very high salinity environment than in a low ionic strength freshwater. Historically, therefore, the properties of seawater have traditionally been studied separately from those of freshwaters and are presented separately, although the important overlaps such as in the area of carbonate equilibria are highlighted.

Introduction

(iii) The chemistry of soils and rocks. There are very significant overlaps with freshwater chemistry but the main differences arise from the very large quantities of solid matter providing very large surfaces and often restricting access of oxygen so that conditions readily become anoxic. However, many of the basic issues such as carbonate equilibria and redox properties overlap very strongly with the ﬁeld of freshwater chemistry. Soils can, however, vary very greatly according to their location and the physical and chemical processes which have affected them during and since their formation. (iv) Environmental organic chemistry. Much of the traditional study of the aquatic and soil environment has been concerned with its inorganic constituents. Increasingly, however, it is recognised that organic matter plays a very important role both in terms of the contribution of natural organic substances to the properties of waters and soils, but also that speciﬁc organic compounds, many of them deriving from human activity, show properties in the environment which are not easily understood from traditional approaches and therefore these have become a rather distinct area of study. (v) Atmospheric chemistry. The atmosphere contains both gas phase and particulate material. The study of both is important and the two interact very substantially. However, as outlined previously, chemical processes in the atmosphere tend to be very strongly inﬂuenced by kinetic factors, and to a large extent are concerned with rather small molecules, which play only a minor part in the chemistry of the aquatic environment or solid earth. Inevitably, there are important processes at the interface between the atmosphere and the land surface or oceans, but these are dealt with more substantially in the companion volume on Pollution Science.

1.4 UNITS OF CONCENTRATION 1.4.1 Atmospheric Chemistry Concentrations of trace gases and particles in the atmosphere can be expressed as mass per unit volume, typically mg m3. The difﬁculty with this unit is that it is not independent of temperature and pressure. Thus, as an airmass becomes warmer or colder, or changes in pressure, so its volume will change, but the mass of the trace gas will not. Therefore, air containing 1 mg m3 of sulfur dioxide in air at 01C will contain less than 1 mg m3 of sulfur dioxide in air if heated to 251C. For gases (but not particles), this difﬁculty is overcome by expressing the concentration of

Chapter 1

the trace gas as a volume mixing ratio. Thus, 1 cm3 of pure sulfur dioxide dispersed in 1 m3 of polluted air would be described as a concentration of 1 ppm. Reference to the gas laws tells us that not only is this one part per 106 by volume, it is also one molecule in 106 molecules and one mole in 106 moles, as well as a partial pressure of 106 atm. Additionally, if the temperature and pressure of the airmass change, this affects the trace gas in the same way as the air in which it is contained and the volume-mixing ratio does not change. Thus, ozone in the stratosphere is present in air at considerably higher mixing ratios than in the lower atmosphere (troposphere), but if the concentrations are expressed in mg m3 they are little different because of the much lower density of air at stratospheric attitudes. Chemical kineticists often express atmospheric concentrations in molecules per cubic centimetre (molec cm3), which has the same problem as the mass per unit volume units.

Worked Example The concentration of nitrogen dioxide in polluted air is 85 ppb. Express this concentration in units of mg m3 and molec cm3 if the air temperature is 201C and the pressure 1005 mb (1.005 105 Pa). Relative molecular mass of NO2 is 46; Avogadro number is 6.022 1023. The concentration of NO2 is 85 mL m3. At 201C and 1005 mb, 85 106 273 1005 22:41 293 1013 ¼ 161 106 g

85 mL NO2 weigh 46

NO2 concentration ¼ 161 mg m3 This is equivalent to 161 pg cm3, and 161 1012 46 molecules

161 pg NO2 contain 6:022 1023 ¼ 2:1 1012

and NO2 concentration ¼ 2.1 1012 molec cm3.

1.4.2 Soils and Waters Concentrations of pollutants in soils are most usually expressed in mass per unit mass, for example, milligrams of lead per kilogram of soil. Similarly, concentrations in vegetation are also expressed in mg kg1 or mg kg1. In the case of vegetation and soils, it is important to distinguish

Introduction

between wet and dry weight concentrations, in other words, whether the kilogram of vegetation or soil is determined before or after drying. Since the moisture content of vegetation can easily exceed 50%, the data can be very sensitive to this correction. In aquatic systems, concentrations can also be expressed as mass per unit mass and in the oceans some trace constituents are present at concentrations of ng kg1 or mg kg1. More often, however, sample sizes are measured by volume and concentrations expressed as ng L1 or mg L1. In the case of freshwaters, especially, concentrations expressed as mass per litre will be almost identical to those expressed as mass per kilogram. As a kind of shorthand, however, water chemists sometimes refer to concentrations as if they were ratios by weight, thus, mg L1 are expressed as ppm, mg L1 as ppb and ng L1 as ppt. This is unfortunate as it leads to confusion with the same units used in atmospheric chemistry with a quite different meaning. 1.5 THE ENVIRONMENT AS A WHOLE A facet of the chemically centred study of the environment is a greater integration of the treatment of environmental media. Traditional boundaries between atmosphere and waters, for example, are not a deterrent to the transfer of chemicals (in either direction), and indeed many important and interesting processes occur at these phase boundaries. In this book, the treatment ﬁrst follows traditional compartments (Chapters 2, 3, 4, and 5) although some exchanges with other compartments are considered. Fundamental aspects of the science of atmosphere, waters, and soils are described, together with current environmental questions, exempliﬁed by case studies. Subsequently, the organic chemistry of the environment is considered in Chapter 6, and quantitative aspects of transfer across phase boundaries are described in Chapter 7, where examples are given of biogeochemical cycles. REFERENCES 1. For readers requiring knowledge of basic chemical principles R.M. Harrison and S.J. de Mora, Introductory Chemistry for the Environmental Sciences, 2nd edn, Cambridge University Press, Cambridge, 1996. 2. For more detailed information upon pollution phenomena Pollution: Causes, Effects and Control, 4th edn, R.M. Harrison (ed), RSC, Cambridge, 2001 or R.M. Harrison (ed), Introduction to Pollution Science, RSC, Cambridge, 2006.

CHAPTER 2

Chemistry of the Atmosphere PAUL S. MONKS Department of Chemistry, University of Leicester, LE1 7RH, Leicester, UK

2.1 INTRODUCTION The thin gaseous envelope that surrounds our planet is integral to the maintenance of life on earth. The composition of the atmosphere is predominately determined by biological processes acting in concert with physical and chemical change. Though the concentrations of the major atmospheric constituents oxygen and nitrogen remain the same, the concentration of trace species, which are key to many atmospheric processes are changing. It is becoming apparent that man’s activities are beginning to change the composition of the atmosphere over a range of scales, leading to, for example, increased acid deposition, local and regional ozone episodes, stratospheric ozone loss and potentially climate change. In this chapter, we will look at the fundamental chemistry of the atmosphere derived from observations and their rationalisation. In order to understand the chemistry of the atmosphere we need to be able to map the different regions of the atmosphere. The atmosphere can be conveniently classiﬁed into a number of different regions which are distinguished by different characteristics of the dynamical motions of the air (see Figure 1). The lowest region, from the earth’s surface to the tropopause at a height of 10–15 km, is termed the troposphere. The troposphere is the region of the active weather systems which determine the climate at the surface of the earth. The part of the troposphere at the earth’s surface, the planetary boundary layer, is that which is inﬂuenced on a daily basis by the underlying surface. Above the troposphere lies the stratosphere, a quiescent region of the atmosphere where vertical transport of material is slow and radiative transfer of energy dominates. In this region lies the ozone layer which 8

Chemistry of the Atmosphere

Figure 1 Vertical structure of the atmosphere. The vertical proﬁle of temperature can be used to deﬁne the different atmospheric layers

Chapter 2

has an important property of absorbing ultraviolet (UV) radiation from the sun, which would otherwise be harmful to life on earth. The stratopause at approximately 50 km altitude marks the boundary between the stratosphere and the mesosphere, which extends upwards to the mesopause at approximately 90 km altitude. The mesosphere is a region of large temperature extremes and strong turbulent motion in the atmosphere over large spatial scales.1 Above the mesopause is a region characterised by a rapid rise in temperature, known as the thermosphere.2 In the thermosphere, the atmospheric gases, N2 and O2, are dissociated to a significant extent into atoms so the mean molecular mass of the atmospheric species falls. The pressure is low and thermal energies are significantly departed from the Boltzmann equilibrium. Above 160 km gravitational separation of the constituents becomes significant and atomic hydrogen atoms, the lightest neutral species, moves to the top of the atmosphere. The other characteristic of the atmosphere from mesosphere upwards is that above 60 km, ionisation is important. This region is called the ionosphere. It is subdivided into three regimes, the D, E and F region, characterised by the types of dominant photo ionisation.3 With respect to atmospheric chemistry, though there is a great deal of interesting chemistry taking place higher up in the atmosphere,1–3 we shall focus in the main on the chemistry of the troposphere and stratosphere. 2.2 SOURCES OF TRACE GASES IN THE ATMOSPHERE As previously described, the troposphere is the lowest region of the atmosphere extending from the earth’s surface to the tropopause at 10–18 km. About 90% of the total atmospheric mass resides in the troposphere and the greater part of the trace gas burden is found there. The troposphere is well mixed and its bulk composition is 78% N2, 21% O2, 1% Ar and 0.036% CO2 with varying amounts of water vapour depending on temperature and altitude. The majority of the trace species found in the atmosphere are emitted into the troposphere from the surface and are subject to a complex series of chemical and physical transformations. Trace species emitted directly into the atmosphere are termed to have primary sources, e.g. trace gases such as SO2, NO and CO. Those trace species formed as a product of chemical and/or physical transformation of primary pollutants in the atmosphere, e.g. ozone, are referred to as having secondary sources or being secondary species. Emissions into the atmosphere are often broken down into broad categories of anthropogenic or ‘‘man-made sources’’ and biogenic or

Chemistry of the Atmosphere

natural sources with some gases also having geogenic sources. Table 1 lists a selection of the trace gases and their major sources.4 For the individual emission of a primary pollutant there are a number of factors that need to be taken into account in order to estimate the emission strength, these include the range and type of sources and the spatio- and temporal-distribution of the sources. Often these factors are compiled into the so-called emission inventories that combine the rate of emission of various sources with the number and type of each source and the time over which the emissions occur. Figure 2 shows the UK emission inventory for a range of primary pollutants ascribed to different source categories (see caption of Figure 2). It is clear from the data in Figure 2 that, for example, SO2 has strong sources from public power generation whereas ammonia has strong sources from agriculture. Figure 3 shows the (2002) 1 1 km emission inventories for SO2 and NO2 for the UK. In essence, the data presented in Figure 2 has been apportioned spatially according to magnitude of each source category (e.g. road transport, combustion in energy production and transformation, solvent use). For example, in Figure 3a, the major road routes are clearly visible, showing NO2 has a major automotive source (cf. Figure 2). It is possible to scale the budgets of many trace gases to a global scale. It is worth noting that there are a number of sources that do not occur within the boundary layer (the decoupled lowest layer of the troposphere, see Figure 1), such as lightning production of nitrogen oxides and a range of pollutants emitted from the combustion-taking place in aircraft engines. The non-surface sources often have a different chemical impact owing to their direct injection into the free troposphere (the part of the troposphere that overlays the boundary layer). In summary, there are a range of trace species present in the atmosphere with a myriad of sources varying both spatially and temporally.5 It is the chemistry of the atmosphere that acts to transform the primary pollutants into simpler chemical species. 2.3 INITIATION OF PHOTOCHEMISTRY BY LIGHT Photodissociation of atmospheric molecules by solar radiation plays a fundamental role in the chemistry of the atmosphere. The photodissociation of trace species such as ozone and formaldehyde contributes to their removal from the atmosphere, but probably the most important role played by these photoprocesses is the generation of highly reactive atoms and radicals. Photodissociation of trace species and the subsequent reaction of the photoproducts with other molecules is the prime initiator and driver for the bulk of atmospheric chemistry.

Table 1

Chapter 2

Natural and anthropogenic sources of a selection of trace gases

Compound

Natural sources

Carbon-containing compounds Carbon dioxide Respiration; oxidation of (CO2) natural CO; destruction of forests Methane (CH4) Enteric fermentation in wild animals; emissions from swamps, bogs, etc., natural wet land areas; oceans Carbon monoxide (CO)

Forest ﬁres; atmospheric oxidation of natural hydrocarbons and methane

Light parafﬁns, C2–C6

Aerobic biological source

Oleﬁns, C2–C6

Photochemical degradation of dissolved oceanic organic material Insignificant

Aromatic hydrocarbons Terpenes (C10H16) CFCs and HFCs

Trees (broadleaf and coniferous); plants None

Nitrogen-containing trace gases Nitric oxide (NO) Forest ﬁres; anaerobic processes in soil; electric storms Nitrogen dioxide Forest ﬁres; electric storms (NO2) Nitrous oxide (N2O) Ammonia (NH3)

Emissions from denitrifying bacteria in soil; oceans Aerobic biological source in soil. Breakdown of amino acids in organic waste material

Sulfur-containing trace gases Dimethyl sulﬁde Phytoplankton (DMS) Sulfur dioxide Oxidation of H2S; volcanic (SO2) activity

Anthropogenic sources Combustion of oil, gas, coal and wood; limestone burning Enteric fermentation in domesticated ruminants; emissions from paddy ﬁelds; natural gas leakage; sewerage gas; colliery gas; combustion sources Incomplete combustion of fossil fuels and wood, in particular motor vehicles, oxidation of hydrocarbons; industrial processes; blast furnaces Natural gas leakage; motor vehicle evaporative emissions; reﬁnery emissions Motor vehicle exhaust; diesel engine exhaust Motor vehicle exhaust; evaporative emissions; paints, gasoline, solvents Refrigerants; blowing agents; propellants Combustion of oil, gas and coal Combustion of oil, gas and coal; atmospheric transformation of NO Combustion of oil and coal Coal and fuel oil combustion; waste treatment

Landﬁll gas Combustion of oil and coal; roasting sulﬁde ores

(Continued )

Chemistry of the Atmosphere Table 1 (Continued ) Compound

Natural sources

Anthropogenic sources

Other minor trace gases Hydrogen Oceans, soils; methane oxidation, isoprene and terpenes via HCHO Ozone In the stratosphere; natural NONO2 conversion Water (H2O)

Motor vehicle exhaust; oxidation of methane via formaldehyde (HCHO) Man-made NONO2 conversion; supersonic aircraft Insignificant

Evaporation from oceans

Source: From ref. 4.

Sources of emission in the UK 10 2

2 7

SO2

NOx

NMVVOC 1

2 8

CH4 79

4 7

CO2

N2O

NH3

Public power, cogeneration and district heating

Road transport

Commercial, institutional and residential combustion

Other mobile sources and machinery

Industrial combustion

Waste treatment and disposal

Production processes

10. Agriculture

Extraction and distribution of fossil fuels

11. Nature

Solvent use

Figure 2 UK emission statistics by UNECE source category (1) Combustion in Energy production and transformation; (2) Combustion in commercial, institutional, residential and agriculture; (3) Combustion in industry; (4) Production processes; (5) Extraction and distribution of fossil fuels; (6) Solvent use; (7) Road transport; (8) Other transport and mobile machinery; (9) Waste treatment and disposal; (10) Agriculture, forestry and land use change; (11) Nature

Chapter 2

Figure 3 UK emission maps (2002) on a 1 1 km grid for (a) NO2 and (b) SO2 and in kg (data from UK NAIE, http://www.naei.org.uk/)

The light source for photochemistry in the atmosphere is the sun. At the top of the atmosphere there is ca. 1370 W m2 of energy over a wide spectral range, from X-rays through the visible to longer wavelength. By the time the incident light reaches the troposphere much of the more energetic, shorter wavelength light has been absorbed by molecules such as oxygen, ozone and water vapour or scattered higher in the atmosphere. Typically, in the surface layers, light of wavelengths longer than 290 nm are available (see Figure 4). In the troposphere, the wavelength at which the intensity of light drops to zero is termed the atmospheric cut-oﬀ. For the troposphere, this wavelength is determined by the overhead stratospheric ozone column (absorbs ca. l r 310 nm) and the aerosol loading. In the mid- to upper-stratosphere, the amount of O3 absorption in the ‘‘window’’ region at 200 nm between the O3 and O2 absorptions controls the availability of short-wavelength radiation that can photodissociate molecules that are stable in the troposphere. In the stratosphere (at 50 km), there is typically no radiation of wavelength shorter than 183 nm. The light capable of causing photochemical reactions is termed the actinic ﬂux, Fl(l) (cm2 s1 nm1), which is also known as the scalar

Chemistry of the Atmosphere 2.5

Solar flux (W m-2 nm-1)

Energy curve for blackbody at 6000K 2.0

Solar irradiance curve outside atmosphere Solar irradiance curve at sea level for a O2 zenith angle of 0° H2O O2

O3 1.5

H2O

1.0

H2O H2O

0.5

H2O

UV O3

H2O CO 2

H2O

Visible Infrared

0 0 200

600

1000

1600

2000

2600

3000

Wavelength (nm) Figure 4

Solar ﬂux outside the atmosphere and at sea level, respectively. The emission of a blackbody at 6000 K is included for comparison. The species responsible for light absorption in the various regions (O2, H2O, etc.) are also shown (after ref. 60)

intensity or spherical radiant ﬂux, viz Z Fl ðlÞ ¼ Ll ðl; W; jÞ do

ð2:1Þ

where Ll(l) (cm2 s1 sr1 nm1) denotes the spectral photon radiance, o is the solid angle and (W, j) are the polar and azimuthal angles of incidence of the radiation interacting with the molecule of interest. In essence, all angles of incident light must be considered when measuring or calculating the actinic ﬂux (i.e. W ¼ 0–1801, o ¼ 0–3601). Photolysis rates are often expressed as a ﬁrst-order loss process, e.g. in the photolysis of NO2 NO2 þ hn - NO þ O(3P)

(2.2)

i.e.

d½NO2 ¼ j2:2 ½NO2 dt

ð2:3Þ

where the photolysis frequency, j, can be expressed as j¼

Zlmax lmin

sðl; TÞfðl; TÞFl ðlÞdl

ð2:4Þ

Chapter 2 9.E-03

J214 (clear) 8.E-03

J224 (cloudy)

π sr j ( NO2)/ s-1 2π

7.E-03 6.E-03 5.E-03 4.E-03 3.E-03 2.E-03 1.E-03 0.E+00 04:00

06:00

08:00

10:00

12:00

14:00

16:00

18:00

20:00

22:00

time (UT)/ h Figure 5 Diurnal proﬁle of j(NO2) (reaction (2.2)) measured on a clear sky and cloudy day61

where s the absorption cross-section (cm2), f the quantum yield of the photoproducts and T the temperature. Figure 5 shows a typical measured photolysis rate for NO2 (reaction (2.2)) in the atmosphere. The photolysis rate reaches a maximum at solar noon concomitant with the maximum in solar radiation.

Example for the calculation of photolysis rates From the following data calculate the photolysis rate of O3 into O(1D) at T ¼ 298 K. Wavelength (nm) F(l) (photon cm2 s1) s (cm2 molec1) 295–300 300–305 305–310 310–315 315–320

2.66 4.20 1.04 1.77 1.89

10 1014 1015 1015 1015

59.95 10 26.50 1020 15.65 1020 7.95 1020 4.15 1020

f 1.0 1.0 0.83 0.26 0.21

The data used represent a gross approximation, owing to the large changes in all the quantities in the region of the atmospheric cut-oﬀ, but

Chemistry of the Atmosphere

are illustrative of the method and controlling quantities for ozone photolysis. The most commonly used form of Equation (2.4) becomes jðs1 Þ ¼

lmax X

sav ðlÞfav ðlÞFav ðlÞ

lmin

The av denotes that average of wavelength intervals. Therefore, j ¼ 3.14 104 s1.

2.4 TROPOSPHERIC CHEMISTRY In general, tropospheric chemistry is analogous to a low-temperature combustion system, the overall reaction is given by CH4 þ 2O2 - CO2 þ 2H2O

(2.5)

Unlike combustion, this is not a thermally initiated process but a process initiated and propagated by photochemistry. The chemistry that takes place in the troposphere, and in particular, the photochemistry is intrinsically linked to the chemistry of ozone. Tropospheric ozone acts as initiator, reactant and product in much of the oxidation chemistry that takes place in the troposphere and stratospheric ozone determines the amount of short wavelength radiation available to initiate photochemistry. Figure 6 shows a typical ozone proﬁle through the atmosphere illustrating a number of interesting points. First, 90% of atmospheric ozone can be found in the stratosphere (see Section 2.10); on average about 10% can be found in the troposphere. Second, the troposphere in the simplest sense consists of two regions. The lowest kilometre or so contains the planetary boundary layer6 and inversion layers which can act as pre-concentrators for atmospheric emissions from the surface and hinder exchange to the so-called free troposphere, the larger part by volume, that sits above the boundary layer. For a long time, transport from the stratosphere to the troposphere was thought to be the dominant source of ozone in the troposphere.7,8 Early in the 1970s, it was ﬁrst suggested9,10 that tropospheric ozone originated mainly from production within the troposphere by photochemical oxidation of CO and hydrocarbons catalysed by HOx and NOx. These sources are balanced by in-situ photochemical destruction of ozone and by dry deposition at the earth’s surface. Many studies, both experimental- and model-based have set about determining the

Chapter 2

35 20

Stratospheric Ozone 15

20 10

15 10 5 0

Figure 6

Ozone Layer

Ozone increases from pollution

Altitude (miles)

Altitude (kilometers)

Tropospheric 5 Ozone

Ozone concentration

A typical atmospheric ozone proﬁle through the atmosphere.58 The concentration is expressed as a volume mixing ratio

contribution of both chemistry and transport to the tropospheric ozone budget on many different spatial and temporal scales. There is growing evidence that the composition of the troposphere is changing.11 For example, analysis of historical ozone records has indicated that tropospheric ozone levels in both hemispheres have increased by a factor of 3–4 over the last century. Methane concentrations have effectively doubled over the past 150 years and N2O levels have risen by 15% since pre-industrial times.12 Measurements of halocarbons have shown that this group of chemically and radiatively important gases to be increasing in concentration until relatively recently.12 One of the difﬁculties about discussing tropospheric chemistry in general terms is that by the very nature of the troposphere being the lowest layer of the atmosphere it has complex multi-phase interactions with the earth’s surface, which can vary considerably between expanses of ocean to deserts (see Figure 7). The fate of any chemical species (Ci) in the atmosphere can be represented as a continuity or mass balance equation such as dCi duCi dvCi dwCi dCi ¼ þ Kz þ Pi Li þ Si dt dx dz dz dy dCi þ dt clouds

ð2:6Þ

Chemistry of the Atmosphere

Figure 7 A schematic representation of the atmosphere’s role in the earth system

Chapter 2

where t is the time, u, v and w are the components of the wind vector in x, y and z accounting for the horizontal and vertical large-scale transport. Small-scale turbulence can be accounted for using Kz, the turbulent diffusion coefﬁcient, Pi and Li are the chemical production and loss terms and Si are the sources owing to emissions. Cloud processes (vertical transport, washout and aqueous phase chemistry) are represented in a cloud processing term. The application of this type of equation is the basis of chemical modelling. In this chapter we will concentrate mainly on the chemical terms in this equation and the processes that control them, but inherently, as the study of tropospheric photochemistry is driven by observations, these must be placed within the framework of Equation (2.6). 2.5 TROPOSPHERIC OXIDATION CHEMISTRY Though atmospheric composition is dominated by both oxygen and nitrogen, it is not the amount of oxygen that deﬁnes the capacity of the troposphere to oxidise a trace gas. The ‘‘oxidising capacity’’ of the troposphere is a somewhat nebulous term probably best described by Thompson.13 The total atmospheric burden of O3, OH and H2O2 determines the ‘‘oxidising capacity’’ of the atmosphere. As a result of the multiple interactions among the three oxidants and the multiphase activity of H2O2, there is no single expression that deﬁnes the earth’s oxidising capacity. Some researchers take the term to mean the total global OH, although even this parameter is not deﬁned unambiguously.

Figure 8 gives a schematic representation of tropospheric chemistry, representing the links between emissions, chemical transformation and sinks for a range of trace gases. Atmospheric photochemistry produces a variety of radicals that exert a substantial inﬂuence on the ultimate composition of the atmosphere.14 Probably the most important of these in terms of its reactivity is the hydroxyl radical, OH. The formation of OH is the initiator of radicalchain oxidation. Photolysis of ozone by UV light in the presence of water vapour is the main source of hydroxyl radicals in the troposphere, viz O3 þ hn(l o 340 nm) - O(1D) þ O2(1Dg)

(2.7)

O(1D) þ H2O - OH þ OH

(2.8)

Chemistry of the Atmosphere

Figure 8 A simpliﬁed scheme of tropospheric chemistry. The ﬁgure illustrates the interconnections in the chemistry, as well as the role of sources, chemical transformation and sinks

The fate of the bulk of the O(1D) atoms produced via reaction (2.7) is collisional quenching back to ground-state oxygen atoms, viz O(1D) þ N2 - O(3P) þ N2

(2.9)

O(1D) þ O2 - O(3P) þ O2

(2.10)

The fraction of O(1D) atoms that form OH is dependent on pressure and the concentration of H2O; typically in the marine boundary layer (MBL) about 10% of the O(1D) generate OH. Reactions (2.7 and 2.8) are the primary source of OH in the troposphere, but there are a number of other reactions and photolysis routes capable of forming OH directly or indirectly. As these compounds are often products of OH radical initiated oxidation they are often termed secondary sources of OH and include the photolysis of HONO, HCHO, H2O2 and acetone and the reaction of O(1D) with methane (see Figure 9). Table 2 illustrates the average contribution of various formation routes with altitude in a standard atmosphere. Two important features of OH chemistry make it critical to the chemistry of the troposphere. The ﬁrst is its inherent reactivity; the second is its

Chapter 2 PRODUCTION j(O1D) j(HONO) j(HCHO) j(H2O2) O3 + CnH2n j(ROOH) j(RCHO) j(H2O)

LOSS

RO2

CH4, HC

NO CO

RADICAL - NOx OH + NO2 HO2 + NO2 RO + NO → Nitrate RCO3 + NO2

HO2

RADICAL - RADICAL HO2 + HO2 HO2 + RO2 RO2 + RO2

NO, O3

Heterogeneous P

Figure 9

The sources, interconversions and sinks for HOx (and ROx) in the troposphere14

Table 2

Calculated fractional contribution of various photolysis rates to radical production with altitude

Altitude Ground Midtroposphere Uppertroposphere Lower stratosphere

j(O(1D)) þ H2O

j(O(1D)) þ j(Acetone) CH4

j(H2O2)

j(HCHO)

0.68 0.52

0.0 Neg.

Neg. 0.03

0.15 0.20

0.17 0.25

0.35

0.02

0.1

0.25

0.28

0.40

0.1

0.25

0.1

0.15

Note: (Neg: Negligible).

relatively high concentration given its high reactivity. The hydroxyl radical is ubiquitous throughout the troposphere owing to the widespread nature of ozone and water. In relatively unpolluted regimes (low NOx) the main fate for the hydroxyl radical is reaction with either carbon monoxide or methane to produce peroxy radicals such as HO2 and CH3O2, viz OH þ CO - H þ CO2

(2.11)

H þ O2 þ M - HO2 þ M

(2.12)

Chemistry of the Atmosphere

and OH þ CH4 - CH3 þ H2O

(2.13)

CH3 þ O2 þ M- CH3O2 þ M

(2.14)

In low-NOx conditions, HO2 can react with ozone leading to further destruction of ozone in a chain sequence involving production of hydroxyl radicals. HO2 þ O3 - OH þ 2O2

(2.15)

OH þ O3 - HO2 þ O2

(2.16)

Alternatively, it can recombine to form hydrogen peroxide (H2O2) HO2 þ HO2 - H2O2 þ O2

(2.17)

or react with organic peroxy radicals such as CH3O2 to form organic hydroperoxides, CH3O2 þ HO2 - CH3O2H þ O2

(2.18)

The formation of peroxides is effectively a chain termination reaction, as under most conditions these peroxides can act as effective sinks for HOx. In more polluted conditions (high-NOx), peroxy radicals catalyse the oxidation of NO to NO2 HO2 þ NO - OH þ NO2

(2.19)

leading to the production of ozone from the subsequent photolysis of nitrogen dioxide and reaction of the photoproducts, i.e. NO2 þ hn(l o 420 nm) - NO þ O(3P)

(2.2)

O þ O2 þ M - O3 þ M

(2.20)

Hydroxyl radicals produced in reaction (2.19) can go on to form more peroxy radicals (e.g. via reactions (2.11) and (2.13)). Similarly to HO2, CH3O2 can also oxidise NO to NO2. CH3O2 þ NO - CH3O þ NO2

(2.21)

Chapter 2 O3 + hν

H2O

OH + CH4

CH3 + O2 CH3O2

HO2

O3 hν + O2

CH3O

CH3OOH

NO2

+ O2

HO2

hν + O2 O2

hν

NO2

HCHO hν CO(+2HO2)

CO2

Figure 10 Simpliﬁed mechanism for the photochemical oxidation of CH4 in the troposphere62

The resulting methoxy radical reacts rapidly with O2 to form formaldehyde and HO2. CH3O þ O2 - HCHO þ HO2

(2.22)

The oxidation of methane is summarised schematically in Figure 10. The OH radical may have another fate, dependant on the concentration of NO2, it can react with NO2 to form nitric acid. OH þ NO2 þ M - HNO3 þ M

(2.23)

The formation of HNO3 represents an effective loss mechanism for both HOx and NOx.

Chemistry of the Atmosphere

Example lifetime calculation For the reaction, CH4 þ OH - CH3 þ H2O that is the key loss process for CH4 (a greenhouse gas), the rate coefﬁcient for the reaction at atmospheric temperatures is given by k ¼ 8.4 1015 cm3 molecule1 s1. Given the mean atmospheric concentration is [OH] ¼ 5 105 molecule cm3, what is the atmospheric lifetime of CH4? Answer: For a reaction of the type A þ B - P (i.e. second-order) the atmospheric lifetime is given by 4 tCH OH ¼

1 1 ¼ ¼ 3:1 108 s ¼ 9:9 years 15 k½OH ð6:3 10 Þð5 105 Þ

The type of calculations are useful as they give an indication of the likely chemical lifetime (i.e. the amount of time it will take before a molecule is reacted away) of a molecule in the atmosphere. Clearly, this form of calculation does not take into account any other chemical loss routes other than reaction with OH or any other physical process that may remove a molecule.

Temperature dependence of a reaction The lifetime of a compound in the atmosphere will depend on how fast it reacts with main atmospheric oxidants. Many reaction rates vary with temperature, therefore in the atmosphere the lifetime will vary with altitude. For the reaction between OH and CH4, the temperature dependence of the reaction is given by k ¼ 1.85 1012exp(1690/T). How does the reaction rate vary between 0.1 and 10 km and therefore effect the lifetime? For the lifetime calculation see previous description. Answer:

Region Boundary layer Lower troposphere Middle troposphere Lower stratosphere

Altitude (km)

T (1C)

0.1 1 6 10

25 9 24 56

tOH k (cm3 molecule1 s1) (years) 6.3 4.6 2.0 7.6

1015 1015 1015 1016

9.9 13.8 30.3 82.7

Chapter 2

2.5.1 Nitrogen Oxides and the Photostationary State From the preceding discussion, it can be seen that the chemistry of nitrogen oxides are an integral part of tropospheric oxidation and photochemical processes. Nitrogen oxides are released into the troposphere from a variety of biogenic and anthropogenic sources including fossil fuel combustion, biomass burning, microbial activity in soils and lightning discharges (see Figure 2). About 30% of the global budget of NOx, i.e. (NO þ NO2) comes from fossil fuel combustion with almost 86% of the NOx emitted in one form or the other into the planetary boundary layer from surface processes.5 Typical NO/NO2 ratios in surface air are 0.2–0.5 during the day tending to zero at night. Over the timescales of hours to days NOx is converted to nitric acid (reaction (2.23)) and nitrates, which are subsequently removed by rain and dry deposition. The photolysis of NO2 to NO and the subsequent regeneration of NO2 via reaction of NO with ozone is sufﬁciently fast, in the moderately polluted environment, for these species to be in dynamic equilibrium, viz NO2 þ hn - NO þ O(3P)

(2.2)

O(3P) þ O2 þ M - O3 þ M

(2.20)

O3 þ NO - NO2 þ O2

(2.24)

Therefore, at suitable concentrations, ambient NO, NO2 and O3 can be said to be in a photochemical steady-state or photostationary state (PSS),15 provided that they are isolated from local sources of NOx and that sunlight intensity is relatively constant, therefore ½O3 ¼

j2:2 ½NO2 k2:24 ½NO

ð2:25Þ

The reactions (2.2, 2.20 and 2.24) constitute a cycle with no net chemistry. The PSS expression is sometimes expressed as a ratio, viz f¼

j2:2 ½NO2 k2:24 ½NO½O3

ð2:26Þ

If ozone is the sole oxidant for NO to NO2 then f ¼ 1. This situation often pertains in urban areas where NOx levels are high and other potential oxidants of NO to NO2 such as peroxy radicals are suppressed. In the presence of peroxy radicals, Equation (2.25) has to be modiﬁed as

Chemistry of the Atmosphere

the NO/NO2 partitioning is shifted to favour NO2, viz ½NO2 ¼ ðk2:24 ½O3 þ k2:19 ½HO2 þ k2:21 ½RO2 Þ=j2:2 ½NO

ð2:27Þ

Though the radical concentrations are typically ca. 1000 times smaller than the [O3], the rate of the radical oxidation of NO to NO2 is ca. 500 times larger than the corresponding oxidation by reaction with O3. We shall return to the significance of the peroxy radical catalysed oxidation of NO to NO2 when considering photochemical ozone production and destruction. From the preceding discussion, it can be seen that the behaviour of NO and NO2 are strongly coupled through both photolytic and chemical equilibria. Because of their rapid interconversion they are often referred to as NOx. NOx, i.e. (NO þ NO2) is also sometimes referred to as ‘‘active nitrogen’’.

Photostationary state At what NO2/NO ratio will the PSS ratio be equal to 1 for midday conditions (j2.2 ¼ 1 103 s1) given that k2.24 ¼ 1.7 1014 cm3 molecule1 s1 and that O3 ¼ 30 ppbv. Using the Equation (2.26) f¼

j2:2 ½NO2 k2:24 ½NO½O3

Converting O3 from ppbv to molecule cm3, as 1 ppbv ¼ 2.46 1010 molecule cm3 (at 251C and 1 atm) 30 ppbv ¼ 7.38 1011 molecule cm3. Given that f ¼ 1 then ½NO2 k2:24 ½O3 ¼ ½NO j2:2 [NO2]/[NO] ¼ 12.55 The extent of the inﬂuence of NOx in any given atmospheric situation depends on its sources, reservoir species and sinks. Therefore, an important atmospheric quantity is the lifetime of NOx. If nitric acid formation is considered to be the main loss process for NOx (i.e. NO2), then the lifetime of NOx (tNOx) can be expressed as the time constant for reaction (2.23), the NO2 to HNO3 conversion. 1 ½NO tNOx ¼ 1þ ð2:28Þ k2:23 ½OH ½NO2

Chapter 2

Therefore, using this simpliﬁcation, the lifetime of NOx is dependent on the [OH] and [NO]/[NO2] ratio. Calculating tNOx under typical upper tropospheric conditions gives lifetimes in the order of 4–7 days and lifetimes in the order of days in the lower free troposphere. In the boundary layer, the situation is more complex as there are other NOx loss and transformation processes other than those considered in Equation (2.29), which can make tNOx as short as 1 h. Integrally linked to the lifetime of NOx and therefore the role of nitrogen oxides in the troposphere is its relation to odd nitrogen reservoir species, i.e. NOy. The sum of total reactive nitrogen or total odd nitrogen is often referred to as NOy and can be deﬁned as NOy ¼ NOx þ NO3 þ 2N2O5 þ HNO3 þ HNO4 þ HONO þ PAN þ nitrate aerosol þ alkyl nitrate, where PAN is peroxyacetlynitrate (see Section 2.5.4). NOy can also be thought of as NOx plus all the compounds that are products of the atmospheric oxidation of NOx. NOy is not a conserved quantity in the atmosphere owing to the potential for some of its constituents (e.g. HNO3) to be efﬁciently removed by deposition processes. Mixing of air masses may also lead to dilution of NOy. The concept of NOy is useful in considering the budget of odd nitrogen and evaluating the partitioning of NOx and its reservoirs in the troposphere.16 In summary, the concentration of NOx in the troposphere

determines the catalytic efﬁciency of ozone production; determines the partitioning of OH and HO2; determines the amount of HNO3 and nitrates produced; and determines the magnitude and sign of net photochemical production or destruction of ozone (see Section 2.5.2).

2.5.2 Production and Destruction of Ozone From the preceding discussion of atmospheric photochemistry and NOx chemistry, it can be seen that the fate of the peroxy radicals can have a marked effect on the ability of the atmosphere either to produce or to destroy ozone. Photolysis of NO2 and the subsequent reaction of the photoproducts with O2 (reactions (2.4) and (2.20)) are the only known way of producing ozone in the troposphere. In the presence of NOx the following cycle for the production of ozone can take place: NO2 þ hn - O(3P) þ NO

(2.2)

O(3P) þ O2 þ M - O3 þ M

(2.20)

OH þ CO - H þ CO2

(2.11)

Chemistry of the Atmosphere

H þ O2 þ M - HO2 þ M

(2.12)

HO2 þ NO - OH þ NO2

(2.19)

NET: CO þ 2O2 þ hn - CO2 þ O3

(2.29)

Similar chain reactions can be written for reactions involving RO2. In contrast, when relatively little NOx is present, as in the remote atmosphere, the following cycle can dominate over ozone production leading to the catalytic destruction of ozone, viz HO2 þ O3 - OH þ 2O2

(2.15)

OH þ CO - H þ CO2

(2.11)

H þ O2 þ M - HO2 þ M

(2.12)

NET: CO þ O3 - CO2 þ O2

(2.30)

Clearly, there is a balance between photochemical ozone production and ozone loss dependent on the concentrations of HOx and NOx. Figure 11 shows the dependence of the production of ozone on NOx taken from a numerical model. There are distinct regions in terms of N(O3) vs. [NOx]

net ozone production (ppbv day-1)

C 6

-2 1

100 NOx (pptv)

1000

Figure 11 Schematic representation of the dependence of the net ozone (N(O3)) production (or destruction) on the concentration of NOx

Chapter 2

on Figure 11. For example, in region A the loss of ozone (L(O3)) is greater than the production of ozone (P(O3)), hence the net product of this process, i.e. N(O3) ¼ P(O3) L(O3)

(2.31)

leads to a net ozone loss. The photochemical loss of ozone can be represented as L(O3) ¼ (f j2.7(O1D) [O3]) þ k2.15[HO2] þ k2.16[OH]

(2.32)

where f is the fraction of O(1D) atoms that react with water vapour (reaction (2.8)) rather than are deactivated to O(3P) (reactions (2.9) and (2.10)). Evaluation of Equation (2.32) is effectively a lower limit for the ozone loss rate as it neglects any other potential chemical loss processes for ozone such as cloud chemistry,17 NO3 chemistry (see Section 2.6) or halogen chemistry (see Section 2.9). The balance point, i.e. where N(O3) ¼ 0 is often referred to, somewhat misleadingly, as the compensation point and occurs at a critical concentration of NOx. Above the compensation point P(O3)4L(O3) and therefore N(O3) is positive and the system is forming ozone. The in-situ formation rate for ozone is approximately given by the rate at which the peroxy radicals (HO2 and RO2) oxidise NO to NO2. This is followed by the rapid photolysis of NO2 (reaction (2.4)) to yield the oxygen atom required to produce O3. P(O3) ¼ [NO] (k2.19[HO2] þ Ski[RO2]i)

(2.33)

It is also worth noting that P(O3) can also be expressed in terms of the concentrations of NOx, j2.2(NO2), O3 and temperature by substitution of Equation (2.27) into Equation (2.33) to give P(O3) ¼ j2.2[NO2] k2.24[NO][O3]

(2.34)

At some concentration of NOx the system reaches a maximum production rate for ozone at dP(O3)/d(NOx) ¼ 0 and even though P(O3) is still significantly larger than L(O3) the net production rate begins to fall oﬀ with increasing NOx. Until this maximum is reached the system is said to be NOx limited with respect to the production of ozone. The turn-over, i.e. dP(O3)/d(NOx) ¼ 0 is caused by the increased competition for NOx by the reaction OH þ NO2 þ M - HNO3 þ M

(2.23)

In reality, the situation is somewhat more complicated owing to the presence at high concentrations of NOx of increased levels of nonmethane hydrocarbons (NMHCs), especially in places such as the urban

Chemistry of the Atmosphere

atmosphere. The oxidation of NMHCs in common with much of tropospheric oxidation chemistry, is initiated by reaction with OH, leading to the rapid sequence of chain reactions. OH þ RH - R þ H2O

(2.35)

R þ O2 þ M - RO2 þ M

(2.36)

RO2 þ NO - RO þ NO2

(2.37)

RO - carbonyl products þ HO2

(2.38)

HO2 þ NO - OH þ NO2

(2.19)

This cycle is similar to the preceding one for the oxidation of CO, in that it is catalytic with respect to OH, R, RO and RO2, with HO2 acting as the chain propagating radical. The mechanism of reaction (2.38) is strongly dependent on the structure of RO.

Photochemical production of ozone Calculate the photochemical production rate of ozone at midday, given that j2.2 ¼ 1 103 s1, O3 ¼ 30 ppbv and k2.24 ¼ 1.7 1014 cm3 molecule1 s1 for a 2:1 ratio of NO2/NO. Using the equation P(O3) ¼ j2.2[NO2] k2.24[NO][O3] As the rate constant is in the units of cm3 molecule1 s1, all the concentrations should be converted into molecule cm3 (as per PSS example). The answer should be converted into ppb h1 by dividing the answer in cm3 molecule1 s1 into ppbv and turning the seconds into hours. For NO2 ¼ 2 ppbv, NO ¼ 1 ppbv, P(O3) ¼ 7.1 ppbv h1. The equation scales as per the absolute values of NO2/NO used. With the involvement of volatile organic compounds (VOCs) in the oxidation chemistry, Figure 11 represents a slice through an n-dimensional surface where there should be a third axis to represent the concentration of VOCs. The peak initial concentrations of ozone generated from various initial concentrations of NOx and VOCs are usually represented as an ‘‘O3 isopleth diagram’’, an example of which is shown

Chapter 2

in Figure 12.18 In an isopleth diagram, initial mixture compositions giving rise to the same peak O3 concentration are connected by the appropriate isopleth. An isopleth plot shows that ozone production is a highly non-linear process in relation to NOx and VOC, but picks out many of the features already highlighted in Figure 11, i.e. when NOx is ‘‘low’’ the rate of ozone formation increases with increasing NOx in a near-linear fashion. On the isopleth, the local maximum in the ozone formation rate with respect to NOx is the same feature as the turn over in N(O3) in Figure 11. The ridgeline along the local maximum separates two different regimes, the so-called NOx-sensitive regime, i.e. N(O3) p (NOx) and the VOC-sensitive (or NOx saturated regime), i.e. N(O3) p (VOC) and increases with increasing NOx. The relationship between NOx, VOCs and ozone embodied in the isopleth diagram indicates one of the problems in the development of air quality policy with respect to ozone. Reductions in VOC are only effective in reducing ozone under VOC-sensitive chemistry (high NOx) and reductions in NOx will only be effective if NOx-sensitive chemistry predominates and may actually increase ozone in VOC-sensitive regions. In general, as an air mass moves away from emission sources, e.g. in an urban region, the chemistry tends to move from VOC-sensitive to NOx-sensitive chemistry.

100 50

NOx (ppbv)

20 10 1 5

Figure 12

5 10

100 VOC (ppbC)

1000

Isopleths giving net rate of ozone production (ppb h1) as a function of VOC (ppbC) and NOx (ppbv) for mean summer daytime meteorology and clear skies18

Chemistry of the Atmosphere

P(O3) (ppbv day-1)

1 3

(a)

2 1

HO2 (pptv)

0 (b) 2

OH (pptv)

0.15

(c)

0.10 0.05 0.00

100 NOx (pptv)

1000

Figure 13 Observed ozone production rates (P(O3)) and concentrations of HO2 and OH during SONEX (8–12 km altitude, 40–601N latitude) plotted a function of the NOx concentration (NOx ¼ observed NO þ modelled NO2). The observed rates and concentrations are averaged over 24 h, using diel factors obtained from a locally constrained box model. The lines on the three panels correspond to model-calculated values for median upper tropospheric conditions as encountered during SONEX19

Case Study I – A practical example of the dependence of ozone production on NOx – An example of the experimental determination of these relationships is shown in Figure 13, a comparison of observed ozone production rates (P(O3)) and concentrations of HO2 and OH from the NASA SONEX mission19 plotted as a function of NOx. The data were taken from a suite of aircraft measurements between 8 and 12 km altitude at latitudes between 40 and 601N. The model data suggest that P(O3) becomes independent of NOx above 70 pptv and the turn-over point into a NOx-saturated regime occurs at about 300 pptv. The bulk of the experimental observation below [NOx] o 300

Chapter 2

pptv observe the P(O3) dependency predicted by the model, but above [NOx] B 300 pptv P(O3), computed from the measured HO2 and NO, continues to increase with NOx, suggesting a NOx-limited regime.

Case Study II – Photochemical control of ozone in the remote marine boundary layer (MBL) – An elegant piece of experimental evidence for the photochemical destruction of ozone comes from studies in the remote MBL over the southern ocean at Cape Grim, Tasmania (411S).20 In the MBL, the photochemical processes are coupled to physical processes that affect the observed ozone concentrations, namely deposition to the available surfaces and entrainment from the free troposphere. The sum of these processes can be represented in the form of an ozone continuity equation (a simpliﬁed version of Equation 2.6), viz d½O3 Ev ð½O3 ft ½O3 Þ vd ½O3 ¼Cþ þ dt H H

ð2:39Þ

where C is a term representative of the photochemistry (the net result of production, P(O3), minus destruction, L(O3) Equation (2.32)), Ev the entrainment velocity (a measure of the rate of ozone transport into the boundary layer), [O3]ft the concentration of free-tropospheric ozone, vd the dry deposition velocity (a measure of the physical loss of ozone to a surface) and H the height of the boundary layer. In general, the MBL is particularly suitable for making photochemical measurements owing to its stable and chemically simple nature. Figure 14 shows the average diurnal cycle of ozone and total peroxide (mainly H2O2) in clean oceanic air as measured at Cape Grim during January 1992. During the sunlit hours an ozone loss of about 1.6 ppbv occurs between mid-morning and late afternoon. This loss of ozone is followed by an overnight replenishment to a similar starting point. In contrast, the peroxide concentration decreases overnight from 900 to 600 pptv and then increases from 600 to 900 pptv between midmorning and late afternoon. It is worth noting that the magnitude of this anti-correlation of ozone and peroxide is dependent on season. The daytime anti-correlation between O3 and peroxide can be interpreted as experimental evidence for the photochemical destruction of ozone, as the ozone is destroyed via reactions (2.7, 2.15 and 2.16) while simultaneously peroxide is formed from chemistry involving the odd-hydrogen radicals OH and HO2 (reactions (2.11, 2.12 and 2.17)).

Chemistry of the Atmosphere

The night-time replenishment of ozone is caused by entrainment of ozone from the free troposphere into the boundary layer. The overnight loss of peroxide is due to deposition over the sea surface (and heterogeneous loss to the aerosol surface), as peroxide has a significant physical loss rate, in contrast to ozone which does not. Therefore, the daytime anti-correlation of ozone and peroxide is indicative of the net photochemical destruction of ozone.

2.5.3 Role of Hydrocarbons The discussion up to this point has focused in the main on the role of CO and CH4 as the fuels for atmospheric oxidation. It is clear that there are many more carbon compounds in the atmosphere than just these two.21 One of the roles of atmospheric photochemistry is to ‘‘cleanse’’ the troposphere of a wide-range of these compounds. Table 3 illustrates the global turnover of a range of trace gases including hydrocarbons and illustrates, for a number of trace gases, the primary role played by OH in their removal. NMHC have a range of both biogenic and anthropogenic sources.5,21 Carbon monoxide chemistry is not independent from NMHC chemistry as 40–60% of surface CO levels over the continents, slightly less over the oceans, and 30–60% of CO levels in the free troposphere, are estimated to come from NMHC oxidation.22 15.0 14.8 1,000

14.6 14.4

900 14.2 14.0

800

13.8 700

13.6 13.4

600

Concentration of ozone (ppbv)

Concentration of peroxides (pptv)

1,100

13.2 500 1

12 Hour of day

13.0 21

Figure 14 Average diurnal cycles for peroxide (open squares) and ozone (ﬁlled squares) in baseline air at Cape Grim, Tasmania (411S) for January 199220

36 Table 3

Chapter 2 Global turnover of tropospheric gases and fraction removed by reaction with OH

Trace gas

Global emission rate(Tg yr1)

Removal by OHa (%)

CO CH4 C2H6 Isoprene Terpenes NO2 SO2 (CH3)2S CFCl3

2800 530 20 570 140 150 300 30 0.3

85 90 90 90 50 50 30 90 0

a Assuming mean global [OH] ¼ 1 106 molecule cm3. Source: After ref. 59.

A major component of the reactive hydrocarbon loading are the biogenic hydrocarbons.21 As previously indicated, the hydrocarbon oxidation chemistry is integral to the production of ozone. Globally, the contribution of NMHC to net photochemical production of ozone is estimated to be about 40%.23 There are a number of inorganic molecules such as NO2 and SO2 (See Table 3) which are also lost via reaction with OH. A number of halocarbons also exist that posses insubstantial tropospheric sinks and have importance in the chemistry of the stratosphere (see Section 2.10). The ultimate products of the oxidation of any hydrocarbon are carbon dioxide and water vapour, but there are many relatively stable partially oxidised organic species such as aldehydes, ketones and carbon monoxide that are produced as intermediate products during this process, with ozone produced as a by-product of the oxidation process. Figure 10 shows a schematic representation of the free radical catalysed oxidation of methane, which is analogous to that of a hydrocarbon. As previously discussed, the oxidation is initiated by reaction of the hydrocarbon with OH and follows a mechanism in with the alkoxy and peroxy radicals are chain propagators and OH is effectively catalytic, viz OH þ RH - R þ H2O

(2.35)

R þ O2 þ M - RO2 þ M

(2.36)

RO2 þ NO - RO þ NO2

(2.37)

RO - carbonyl products þ HO2

(2.38)

Chemistry of the Atmosphere

HO2 þ NO - OH þ NO2

(2.19)

As both reactions (2.37) and (2.19) lead to the oxidation of NO to NO2 the subsequent photolysis leads to the formation of ozone (see reactions (2.2) and (2.20)). The individual reaction mechanism depends on the identity of the organic compounds and the level of complexity of the mechanism. Although OH is the main tropospheric oxidation initiator, reaction with NO3, O3, O(3P) or photolysis may be an important loss route for some NMHCs or the partially oxygenated products produced as intermediates in the oxidation (see reaction (2.38)). In summary, the rate of oxidation of VOCs and therefore by inference the production of ozone is governed by the concentration of the catalytic HOx radicals. There are a large variety of VOCs with a range of reactivities; therefore this remains a complex area. 2.5.4 Urban Chemistry In some respects, the story of atmospheric chemistry and particularly ozone photochemistry begins with urban chemistry and photochemical smog. The term smog arises from a combination of the words smoke and fog. In the 1940s, it became apparent that cities like Los Angeles (LA) were severely afﬂicted with a noxious haze.24 Though, at the time it was thought to be a relatively local phenomenon, with the understanding of its chemistry came the development of a photochemical theory for the whole of the troposphere. The LA smog is often termed photochemical smog and is quite different in origin to the London smogs of the 19th and 20th centuries, which have their origins in abnormally high concentrations of smoke particles and sulfur dioxide. The London smogs have been alleviated with the effective application of legislation that has reduced the burning of coal in the London area. The major features of photochemical smog are high levels of oxidant concentration in particular ozone and peroxidic compounds, produced by photochemical reactions. The principal effects of smog are eye and bronchial irritation as well as plant and material damage. The basic reaction scheme for the formation of photochemical smog is VOC þ hn - VOC þ R

(2.40)

R þ NO - NO2

(2.41)

NO2 þ hn(l o 420 mm) - NO þ O3

(2.2)

38 Table 4

Alkanes Alkenes Aromatics Other

Chapter 2 Percentage of NMHC classes measured in the morning at various locations Urban Los Angeles

Urban Boston

Rural Alabama

42 7 19 33

36 10 30 24

9 43a 2 46b

Source: From ref. 31. a Large contribution from biogenic alkenes. b Mainly oxygen containing.

NO þ O3 - NO2 þ O2

(2.24)

RþR-R

(2.42)

R þ NO2 - NOy

(2.43)

There is a large range of available VOCs in the urban atmosphere, driven by the range of anthropogenic and biogenic sources.5,21,31 Table 4 illustrates the different loadings of the major classes of NMHC in urban and rural locations. These NMHC loadings must be coupled to measures of reactivity and the degradation mechanisms of the NMHC to give a representative picture of urban photochemistry. The oxidation of the VOCs drives, via the formation of peroxy radicals, the oxidation of NO to NO2, where under the sunlit conditions the NO2 can be dissociated to form ozone (reaction (2.2)). Pre-existing ozone can also drive the NO to NO2 conversion (reaction (2.24)). The basic chemistry responsible for urban photochemistry is essentially the same as that takes place in the unpolluted atmosphere (see Section 2.5). It is the range and concentrations of NMHC fuels and the concentrations of NOx coupled to the addition of some photochemical accelerants that can lead to the excesses of urban chemistry. For example, in the Los Angeles basin it is estimated that 3333 ton day1 of organic compounds are emitted as well as 890 ton day1 of NOx. In addition, to the reactions forming OH in the background troposphere, i.e. via the reaction of O(1D) with H2O, viz O3 þ hn(l o 340 nm) - O2 þ O(1D)

(2.7)

O(1D) þ H2O - 2OH

(2.8)

Chemistry of the Atmosphere

under urban conditions OH may be formed from secondary sources such as HONO þ hn(l o 400 nm) - OH þ NO

(2.44)

where the HONO can be emitted in small quantities from automobiles or formed from a number of heterogeneous pathways,25 as well as gasphase routes. OH produced from HONO has been shown to be the dominant OH source in the morning under some urban conditions, where the HONO has built up to significant levels overnight. Another key urban source of OH can come from the photolysis of the aldehydes and ketones produced in the NMHC oxidation chemistry, in particular formaldehyde, viz. HCHO þ hn(l o 334 nm) - H þ HCO

(2.45)

H þ O2 þ M - HO2 þ M

(2.12)

HCO þ O2 - HO2 þ CO

(2.46)

HCHO þ 2O2 þ hn - 2HO2 þ CO

(2.47)

Net

Smog chamber experiments have shown that the addition of aldehyde significantly increases the formation rates of ozone and the conversion rates of NO and NO2 under simulated urban conditions.3 A marked by-product of oxidation in the urban atmosphere, often associated, but not exclusive to, urban air pollution is PAN. PAN is formed by OH þ CH3CHO - CH3CO þ H2O

(2.48)

CH3CO þ O2 - CH2COdO2

(2.49)

addition of NO2 to the peroxyacetyl radical (RCOdO2) leading to the formation of PAN. CH3COdO2 þ NO2 þ M " CH3COdO2dNO2 þ M

(2.50)

PAN is often used as an unambiguous marker for tropospheric chemistry. The lifetime of PAN in the troposphere is very much dependant on the temperature dependence of the equilibrium in reaction (2.50), the lifetime varying from 30 min at T ¼ 298 K to 8 h at T ¼ 273 K. At midtroposphere temperature and pressures PAN has thermal decomposition

Chapter 2

lifetimes in the order of 46 days. It is also worth noting that peroxyacyl radical like peroxy radicals can oxidise NO to NO2 CH3COdO2 þ NO - CH3COdO þ NO2

(2.51)

PAN can be an important component of NOy in the troposphere (see Section 2.5.1) and has the potential to act as a temporary reservoir for NOx and in particular it has the potential to transport NOx from polluted regions into the background/remote atmosphere.26 There are other side-products of urban air pollution that are giving increasing cause for concern. These include the production of both aerosols (both organic and inorganic) and particles. The effects of photochemical smog/urban air pollution remain on the political agenda owing to their potential impact on human health and the economy. In summary, urban photochemistry is not substantially different from tropospheric photochemistry. It is the range and concentrations of the VOCs involved in oxidation coupled to the concentration of NOx and other oxidants that lead to a large photochemical turnover. 2.6 NIGHT-TIME OXIDATION CHEMISTRY Though photochemistry does not take place at night, it is important to note, within the context of tropospheric oxidation chemistry, the potential for oxidation chemistry to continue at night. This chemistry does not lead to the production of ozone, in fact the opposite, but has importance owing to the potential for the production of secondary pollutants. In the troposphere, the main night-time oxidant is thought to be the nitrate radical formed by the relatively slow oxidation of NO2 by O3, viz. NO2 þ O3 - NO3 þ O2

(2.52)

The time constant for reaction (2.52) is of the order of 15 h at an ozone concentration of 30 ppbv and T ¼ 290 K. Other sources include N2O5 þ M - NO3 þ NO2 þ M

(2.53)

but as N2O5 is formed from NO3 þ NO2 þ M - N2O5 þ M

(2.54)

the two species act in a coupled manner. Dinitrogen pentoxide, N2O5, is potentially an important product as it can react heterogeneously with water to yield HNO3. During the daytime the NO3 radical is rapidly photolysed as it strongly absorbs in the visible, viz NO3 þ hn - NO þ O2

(2.55)

Chemistry of the Atmosphere

NO3 þ hn - NO2 þ O(3P)

(2.56)

having a lifetime in the region of 5 s for overhead sun and clear sky conditions. Further, NO3 will react rapidly with NO NO3 þ NO - NO2 þ NO2

(2.57)

which can have significant daytime concentrations in contrast to the night time, where away from strong source regions, the NO concentrations should be near zero. The nitrate radical has a range of reactivity towards VOCs. The nitrate radical is highly reactive towards certain unsaturated hydrocarbons such as isoprene, a variety of butenes and monoterpenes, as well as reduced sulfur compounds such as dimethylsulﬁde (DMS). In the case of DMS, if the NO2 concentration is 60% that of DMS then NO3 is a more important oxidant than OH for DMS in the MBL.27 In general, NO3 abstraction reactions of the type NO3 þ RH - HNO3 þ R

(2.58)

are relatively slow, with the alkyl radical reacting with oxygen under atmospheric conditions to form a peroxy radical. In the case of RH being an aldehyde, acyl products will form acylperoxy radicals (R CO O2), potential sources of peroxyacylnitrates. In contrast, the reaction of NO3 with alkenes occurs by an addition mechanism, initiating a complex chemistry involving nitrooxy substituted organic radicals, which can either regenerate NO2 or produce comparatively stable bifunctional organic nitrate products.28 For example, the products derived from the reaction of NO3 with propene in the presence of O2 and NOx include CH3CHO, HCHO, 1,2-propanedioldinitrate (PDDN), nitroxyperoxypropylnitrate (NPPN) and a-(nitrooxy)acetone (See Figure 15).3 The reaction channel that produces the nitrated acetones also yields peroxy radicals, leading to the potential for a night-time source of OH, either by reaction (HO2 þ O3) or by the direct reaction of the peroxy radical with NO3. For the reaction of NO3 with propene the initial addition can take place at either end of the double bond, viz NO3 þ CH2QCHCH2 þ M - CH2CHCH2(ONO2) þ M

(2.59)

NO3 þ CH2QCHCH2 þ M - CH2CH(ONO2)CH2 þ M

(2.60)

The reaction can then proceed by the mechanism shown schematically in Figure 15. The ratio of ﬁnal products is dependent on the structure of the individual alkenes. In general, for branched alkenes, there is

Chapter 2 O

CH3CCH2ONO2

OO O2

CH3CHCH2ONO2

+ HO2 α-(nitrooxy)acetone CH3CCHCH2ONO2 O

RO2 NO3

NO2

NO3

PDDN

CH3CHO + HCHO + NO2 M

NO3 O2

ONO2

RO2 NO3

CH3CHCH2O2

CH3CHCH2O

NO2 NPPN

Figure 15 A schematic representation of the chain propagation reactions in the NO3 radical initiated oxidation of propene28

Organic nitrates

Aerosol surface reactions

NO 2

Alkenes O3

N 2O5

NO 3 H2O

NO 2

NO, RO2

RH, DMS

HNO 3

Aerosol losses

Surface losses

Figure 16 A simpliﬁed reaction scheme for night-time chemistry involving the nitrate radical27

significant regeneration of NOx and production of unsubstituted carbonyl products while comparatively they are a minor source of HOx.28 For the less alkyl substituted alkenes, there is a greater yield of HOx and bifunctional organic nitrate products but a lesser regeneration of NOx. Therefore, depending on the mix of hydrocarbons, NO3 chemistry can act to either recycle NOx, therefore inhibiting the formation of nitrate

Chemistry of the Atmosphere

aerosol or HNO3 at night and potentially lead to the generation of a night-time source of HOx. Figure 16 provides a simpliﬁed summary of the relevant night-time chemistry involving the nitrate radical. One important difference between NO3 chemistry and daytime OH chemistry is that NO3 can initiate, but not catalyse, the removal of organic compounds. Therefore its concentration can be suppressed by the presence of fast-reacting, with respect to NO3, organic compounds.

Reactive lifetime with respect to different oxidants Different oxidants react at different rates in the atmosphere. Further, the average concentrations of the main oxidants can vary. For the following table of data, calculate the relative lifetime with respect to reaction with OH and NO3, given that the average [OH] E 106 molecule cm3 and the [NO3] E 109 molecule cm3.

Compound

k(OH) (cm3 tOH molecule1 s1) (days)

CH4 C2H6 C2H4 (CH3)2CQC(CH3)2

8.5 2.7 8.5 1.1

1015 1013 1012 1010

1361 43 1.3 0.1

k(NO3) (cm3 molecule1 s1

tNO3 (days)

o1 1019 8 1018 2 1016 4.5 1011

115,740 1446 58 2104

The data nicely illustrate that lifetime is a product of rate and the concentration of the radical species.

Case Study III – Evidence for the role of NO3 in night-time oxidation chemistry from experimental based studies.27,28 Significant NO3 concentrations have been detected over a wide range of atmospheric conditions, indicating a potential role for NO3 over large regions of the atmosphere.28 The atmospheric lifetime of NO3 can be estimated using the steady-state approximation (cf. PSS) to be tðNO3 Þ ¼

½NO3 k52 ½NO2 ½O3

ð2:61Þ

Chapter 2

A useful quantity with which to compare this with is the reciprocal of the lifetime calculated from the sum of the ﬁrst order loss processes involving NO3 and N2O5 X tðNO3 Þ1 k ½HCi þ kðNO3 þDMSÞ ½DMS þ khet ðNO3 Þ i ðNO3 þHCi Þ þ kI ½H2 O þ kII ½H2 O2 þ khet ðN2 O5 Þ :K2:54 ½NO2 ð2:62Þ where the pseudo ﬁrst-order loss rates over the i reactive hydrocarbons are summed and kI and kII are the ﬁrst- and second-order components with respect to H2O of the reaction N2O5 þ H2O - 2HNO3

(2.63)

khet(NO3) and khet(N2O5) are the heterogeneous loss rates for these species and K2.54 is the equilibrium constant for reactions (2.53 and 2.54). Figure 17 shows the observed lifetime of NO3 in the MBL at Mace Head, Ireland29 segregated by arrival wind sector. Mace Head experiences a range of air masses from clean marine air to European continental outﬂow. The measured NO3 lifetime varies from 2 min to 4 h (see Figure 17). An assessment of the parameters controlling NO3 atmospheric lifetime (cf. Equation (2.61)) highlights that, under the conditions encountered the lifetime of NO3 chemistry is very sensitive to DMS and NMHC chemistry in clean marine air. However, in more polluted air the terms involving the indirect loss of N2O5 either in the gas-phase with H2O or through uptake on aerosol tend to dominate.

From the preceding discussion it can be seen that the involvement of NO3 chemistry in gas-phase tropospheric chemistry has potentially six significant consequences.28 (i) The radical can control NOy speciation in the atmosphere at night (via reaction (2.54)). (ii) Nitric acid can be formed, by hydrolysis of N2O5, as a product of a hydrogen abstraction process or indirectly via the NO3-mediated production of OH, which can react with NO2 (reaction (2.23)) to produce nitric acid. (iii) Primary organic pollutants can be oxidised and removed at night. (iv) Radicals (HOx and RO2) produced by NO3 chemistry can act as initiators for chain oxidation chemistry.

Chemistry of the Atmosphere 0 30 [DMS] = 10 pptv τNO3 = 57.8 minutes

90 10

90 00

30 0

90 00

90 10

NO3 /mins.

[DMS] = 44 pptv τNO3 = 14.6 minutes

Figure 17 Polar plots showing the observed lifetime of NO3 vs. arrival wind direction for Mace Head, Ireland (top) and Tenerife (bottom).29 The black circle represents the calculated lifetime of NO3 from reaction with the campaign average [DMS] only

(v) Toxic or otherwise noxious compounds such as peroxyacylnitrates, other nitrates and partially oxidised compounds, may be formed. (vi) Nitrate products or NO3 itself may act as temporary reservoirs in the presence of NOx.

Chapter 2

2.7 OZONE-ALKENE CHEMISTRY Another potential dark source of HOx in the atmosphere, more particularly in the boundary layer, is from the reactions between ozone and alkenes. The ozonolysis of alkenes can lead to the direct production of the OH radical at varying yields (between 7 and 100%) depending on the structure of the alkene, normally accompanied by the co-production of an (organic) peroxy radical. As compared to both the reactions of OH and NO3 with alkenes the initial rate of the reaction of ozone with an alkene is relatively slow, this can be oﬀset under regimes where there are high concentrations of alkenes and/or ozone. For example, under typical rural conditions the atmospheric lifetimes for the reaction of ethene with OH, O3 and NO3 are 20 h, 9.7 days and 5.2 months, respectively in contrast, for the same reactants with 2-methyl-2-butene the atmospheric lifetimes are 2.0 h, 0.9 h and 0.09 h. The mechanism for the reaction of ozone with alkenes was ﬁrst suggested by Criegee30 in the late 1940s and involves the addition of ozone to form a primary ozonide, which rapidly decomposes to form a vibrationally excited carbonyl oxide (Criegee intermediate) and carbonyl products. The Criegee intermediate can then either be collisionally stabilised by a third body (M), or undergo unimolecular decomposition to products (see also Section 2.8). It is now widely believed that alkylsubstituted Criegee intermediates can decompose via a vibrationally hot hydroperoxide intermediate to yield an OH radical, along with another radical species of the general form R1R2CC(O)R3, which is expected to react rapidly with O2 to form a peroxy radical (RO2) in the atmosphere.31 Figure 18 shows a schematic representation of the ozone-alkene reaction mechanism. The OH and peroxy radical yield is dependent on the structure and mechanism of the individual alkene-ozone reaction.32 Table 5 shows typical OH yields for the reaction of a range of anthropogenic and biogenic alkenes with ozone. There is growing experimental evidence33 of the importance of ozone-alkene reactions as significant oxidative sinks for alkenes and a general source of HOx.

2.8 SULFUR CHEMISTRY Sulfur chemistry is an integral part of life, owing to its role in plant and human metabolism. Sulfur compounds have both natural and anthropogenic sources. In modern times, the atmospheric sulfur budget has become dominated by anthropogenic emissions, particularly from fossil fuel burning. It is estimated that 75% of the total sulfur emission budget is dominated by anthropogenic sources with 90% of it occurring in the

Chemistry of the Atmosphere R1

+ O3

O R3 R4

Primary Ozonide O

O R1

O O

(1−α)

O O

R4 O

+ R1

R4 R

Stabilisation

R' R

Decomposition Products

O O

H R

H O O

H H

R O2

O R

Figure 18 A schematic representation of the oxidation of an alkene initiated by reaction with ozone (after ref. 14)

Table 5

Range of OH yields from the reaction of ozone with alkenes

Alkene

OH yield

Ethene Propene Methylpropene D3-carene

0.18 0.06 0.35 0.07 0.72 0.12 1.00

Source: All data taken from ref. 32.

Chapter 2

Northern Hemisphere. The natural sources include volcanoes, plants, soil and biogenic activity in the oceans.5 In terms of photochemistry the major sulfur oxide, sulfur dioxide (SO2) does not photodissociate in the troposphere (cf. NO2), i.e. SO2(X1A1) þ hn(240 olo 330 nm) - SO2(1A2, 1B1)

(2.64)

SO2(X1A1) þ hn(340 olo 400 nm) - SO2(3B1)

(2.65)

The oxidation of sulfur compounds in the atmosphere has implications in a number of different atmospheric problems such as acidiﬁcation, climate balance and the formation of a sulfate layer in the stratosphere, the so-called Junge layer. By far the largest sulfur component emitted into the atmosphere is SO2. Figure 19 shows the spatial distribution of SO2 emissions in 1980 and 2000 from EMEP.34 In Europe, the source regions for SO2 are quite apparent, the so-called black triangle region (southern Poland, eastern Germany and the northern part of the Czech Republic) are the largest sources of anthropogenic sulfur pollution. There are a number of other large emission sources including central UK and the Kola Peninsula, also apparent in Figure 19. The absolute maximum in emissions is in southern Italy around Sicily where the largest single source of both natural, the volcano Mt. Etna, and anthropogenic SO2 is found. SO2 can be detected from space-borne sensors35 as a product of volcanic activity and fossil fuel burning. Figure 19 also illustrates another interesting point in that over much of Europe between 1980 and 2000 there has been a decrease of 1000 ton or more owing to legislative limits. Interestingly, with decreasing land emissions the importance of ship emissions has increased (from 5% to 16% of total European SO2 emissions). The atmospheric oxidation of SO2 can take place by a number of different mechanisms, both hom*ogeneously and heterogeneously in the liquid and gas phases (see Figure 20). The gas-phase oxidation of SO2, viz SO2 þ OH þ M - HSO3 þ M

(2.66)

HSO3 þ O2 - HO2 þ SO3

(2.67)

SO2 þ H2O þ M - H2SO4

(2.68)

can lead to the formation of sulfuric acid, which owing to its relatively low vapour pressure can rapidly attach to the condensed phase such as

Chemistry of the Atmosphere Mg

SO2 100

50000

10000

5000 1000

500 20

100 0

(a)

120 Mg

100

50000 80

10000 5000

1000 40

500 20

100

(b)

Figure 19

100

120

Spatial distributions of sulfur dioxide emissions in (a) 1980 and (b) 2000.34 The units are ton year1 grid sq1

aerosol particles. The bulk of the H2SO4 is lost via wet deposition mechanisms in cloud droplets and precipitation. There is another potential gas-phase loss route for SO2 that can lead to the formation of sulfuric acid in the presence of H2O, i.e. the reaction of SO2 with Criegee intermediates (see Section 2.7). The aqueous-phase oxidation of SO2 is

Chapter 2

more complex, depending on a number of factors such as the nature of the aqueous phase (e.g. clouds and fogs), the availability of oxidants (e.g. O3 and H2O2) and the availability of light. An overview of the mechanism is given in Figure 20. The key steps include the transport of the gas to the surface of a droplet, transfer across the gas–liquid interface, the formation of aqueous-phase equilibria, the transport from the surface into the bulk aqueous phase and subsequent reaction. In brief, the SO2 gas is dissolved in the liquid phase, establishing a set of equilibria for a series of S(IV) species, i.e. SO2 H2O, HSO3 and SO32. SO2(g) þ H2O " SO2dH2O(aq)

(2.69)

1 SO2dH2O(aq) " HSO 3 þ H

(2.70)

2 1 HSO 3 " SO3 þ H

(2.71)

The solubility of SO2 is related to the pH of the aqueous phase, decreasing at lower values of pH. The oxidation of sulfur (IV) to sulfur (VI) is a complex process dependent on many physical and chemical factors. The main oxidants seem to be O2 (catalysed/uncatalysed), O3, H2O2, the oxides of nitrogen and free-radical reactions in clouds and fogs. For example, H2O2 is highly soluble in solution so even at relatively low gas-phase concentrations (typically ca. 1 ppbv) there is a

Figure 20 Summary of emission, oxidation and deposition of S(IV) and S(VI) (after ref. 63)

Chemistry of the Atmosphere

significant concentration of H2O2 present in solution. The oxidation proceeds as O−

HSO3−

+ H2O 2

OOH + H 2 O

ð2:72Þ

O− S OOH + HA

H2SO 4 + A−

ð2:73Þ

where HA is an acid. The ubiquitous occurrence of H2O2, its solubility, its high reactivity and pH independence (under atmospheric conditions) of the rate constant for the reaction with SO2 makes H2O2 one of the most important oxidants for SO2 in the troposphere. A more detailed description of aqueous-phase oxidation of SO2 is given in ref. 36. 2.9 HALOGEN CHEMISTRY In comparison to the atmospheric chemistry taking place in the stratosphere where halogen chemistry is well known and characterised (see Section 2.10), there has been much debate as to the role of halogen species in the oxidative chemistry of the troposphere. There is growing experimental evidence as to the prevalence of halogen chemistry as part of tropospheric photochemistry (see Table 6).14,37 Much of the proposed halogen chemistry is propagated through the reactions of a series of halogen atoms and radicals. Table 6

Sources of reactive halogen species found in various parts of the troposphere

Species, site

Likely source mechanism

ClOx in the polar boundary layer ClOx by salt ﬂats BrOx in the polar boundary layer BrOx in the Dead sea basin BrOx in the free troposphere

By-product of the ‘‘Bromine Explosion’’

BrOx in the MBL IOx in the MBL Source: After ref. 37.

By-product of the ‘‘Bromine Explosion’’ Autocatalytic release from sea-salt on ice, ‘‘Bromine explosion’’ mechanism ‘‘Bromine Explosion’’ (salt pans) (1) Photo-degradation of hydrogen-containing organo-halogen species (e.g. CH3Br), (2) ‘‘Spill-out’’ from the boundary layer, (3) Transport from the stratosphere? ‘‘Bromine Explosion’’ mechanism (1) Photo-degradation of short-lived organo-halogen species (e.g. CH2I2) (2) Photolysis of I2?

Chapter 2

Bromine oxide species can be formed in the polar boundary layer38,39 and areas with high salt levels such as the Dead Sea.40 The major source of gas-phase bromine in the lower troposphere is thought to be the release of species such as IBr, ICl, Br2 and BrCl from sea-salt aerosol, following the uptake from the gas phase and subsequent reactions of hypohalous acids (HOX, where X ¼ Br, Cl, I).41 HOBr þ (Br)aq þ H1 - Br2 þ H2O

(2.74)

The halogen release mechanism is autocatalytic38 and has become known as the ‘‘Bromine explosion’’. The Br2 produced in reaction (2.74) is rapidly photolysed, producing bromine atoms that can be oxidised to BrO by O3 (reaction (2.76)) BrX þ hn - Br þ X

(2.75)

Br þ O3 - BrO þ O2

(2.76)

the resultant BrO reacting with HO2 (reaction (2.77)) to reform HOBr. BrO þ HO2 - HOBr þ O2

(2.77)

Thus, the complete cycle has the form BrO þ O3 þ ðBr Þaq þ ðHþ Þaq

surface; HOx

2BrO þ products

ð2:78Þ

where effectively one BrO molecule is converted to two by oxidation of bromide from a suitable surface. Figure 21 shows a schematic representation of the Bromine-explosion mechanism.37 It is worth noting that the bromine explosion mechanism only occurs from sea salt with a pHo6.5 therefore requiring acidiﬁcation of the aerosol potentially caused by the uptake of strong acids likely to be of anthropogenic origins or naturally occurring acids.42,43 The wide-spread occurrence of these heterogeneous mechanisms are supported by the observed depletion of bromide ions in sea-salt aerosol.42,43

Case Study IV – BrO in the springtime Arctic – In the spring time in both the Arctic and Antarctic large clouds of BrO-enriched air masses are observable from space (see Figure 22).44,45 These clouds cover several thousand square kilometres over the polar sea ice with BrO levels up to 30 pptv. The BrO is always coincident with low levels of ozone in the MBL.38 In order to observe these events, there is a requirement for meteorological conditions that stop mixing between

Chemistry of the Atmosphere

the boundary layer and the free troposphere and sunlight to drive the required photolysis of gaseous bromine (see reaction (2.75)) released heterogeneously through chemical processes on the ice. Because of the prerequisites for strong surface inversions to conﬁne the air and sunlight, episodes of bromine explosion events and boundary layer ozone depletion tend to be conﬁned to spring. It is thought that frost ﬂowers in the Antarctic, which are formed on a liquid layer with high salinity on top of the sea ice surface, are likely to provide a large surface area for the heterogeneous release of reactive bromine but the evidence is not unequivocal. With respect to iodine chemistry, the major sources of iodine are thought to be from macroalgal sources releasing organoiodine compounds.46 Photolysis of the organoiodine compounds releases the iodine. RIx þ hn - R þ I

(2.79)

I þ O3 - IO þ O2

(2.80)

During daylight hours iodine monoxide, IO exists in a fast photochemical equilibrium with I, viz. IO þ hn - I þ O

(2.81)

The Bromine explosion BrO +BrO

+O3 Br

Sea ice surface Aerosol

+HO2 HOBr

HOBr

hν

+Br Br2

HOBr

Heterogeneous release of bromine on sea salt surfaces Catalytical ozone destruction

BrO + HO2 → HOBr + O2

HOBr ( g ) + Br − (aq ) + H + → Br2 ( g ) + H 2O Br , BrO 2O3 → 3O2

Figure 21 A schematic representation of the so-called Bromine explosion mechanism where effectively one BrO molecule is converted to two by oxidation of bromide from a suitable aerosol surface64

Chapter 2

GOME BrO, Mar 2000 BrO VC [x1013 molec/cm2] >8 7 6 5 4 3 2 >1 Gwin DOAS 1.02, Oct.2000 IASB-BIRA

GOME BrO, Oct 1999

BrO VC [x1013 molec/cm2] >8 7 6 5 4 3

DLR ESA-ESRIN BIRA-IASB Contact : [emailprotected]

2 >1

Figure 22 [BrO] in the Antarctic and Artic during spring derived from GOME satellite measurements45

The aerosol ‘‘explosion’’ mechanism, previously described for bromine, acts effectively to recycle the iodine back to the gas phase.

Case Study V – Iodine monoxide chemistry in the coastal margins – Figure 23 shows the measured concentration of iodine monoxide (IO)

Chemistry of the Atmosphere

at Mace Head in Ireland.47 The data in Figure 23 show not only a clear requirement for radiation for photochemical production of IO but also a strong correlation with low tide conditions. The correlation with low tidal conditions is indicative of the likely sources of IO, in that the photolysis of organoiodine compounds (reaction (2.79)) such as CH2I2 and CH2IBr emitted from macroalgae in the intertidal zone are possibly the candidate iodine sources. The photolysis lifetime of a molecule such as diodomethane is only a few minutes at midday. Recent experimental observations have suggested the I2 can be emitted directly from the macroalgae at low tide making this a substantial source of iodine.48 The chemistry can also lead to an enrichment of iodine in the aerosol in the form of iodate providing a route to move biogenic iodine from ocean to land.47

Figure 23 [IO], tidal height and solar radiation measured at Mace Head in Ireland. The grey areas represent the low-tide areas during the day. The dotted line in the upper panel is the IO detection limit46,47

Chapter 2

2.9.1 Tropospheric Halogens and Catalytic Destruction of Ozone Potentially, the most important effect of reactive halogen species maybe that their chemistry may lead to the catalytic destruction of ozone via two distinct cycles Cycle I: XO þ YO - X þ Y þ O2 (or XY þ O2)

(2.82)

X þ O3 - XO þ O2

(2.83)

O3 þ O3 - 3O2

(2.84)

Net:

In Cycle I, the rate-limiting step involves reaction (2.82), the self- or cross-reaction of the halogen monoxide radicals. Cycle I has been identiﬁed to be the prime cause for polar boundary layer ozone destruction.38 The second-ozone destruction cycle, which is more prevalent at low halogen levels, has the form Cycle II: XO þ HO2 - HOX þ O2

(2.85)

HOX þ hn - X þ OH

(2.86)

X þ O3 - XO þ O2

(2.83)

OH þ CO - H þ CO2(þM) - HO2

(2.11, 2.12)

Net: O3 þ CO þ hn - CO2 þ O2

(2.87)

The rate-determining step in this reaction sequence is reaction (reaction (2.85)) making ozone destruction linearly dependent on [XO]. The fraction of HOX that photolyses to give back OH depends critically on the accommodation coefﬁcient of HOX on aerosols. Currently, there is a large uncertainty in this parameter. An important side effect of Cycle II is the potential for the reduction of the [HO2]/[OH] ratio by consumption of HO2. The inorganic halogen chemistry described is summarised in Figure 24.

Chemistry of the Atmosphere

Figure 24 Simpliﬁed scheme of the inorganic halogen reactions in the boundary layer (X ¼ Cl, Br, I). The reactive halogen sources are release from sea salt or photochemical degradation of organohalogen species37

Table 7

Rates of ozone depletion by various catalytic cycles involving XO using measured data typical of MBL just after sunrise

Rate-determining reaction

Cycle (see text)

BrO þ IO IO þ HO2 IO þ IO BrO þ HO2 BrO þ BrO

I II I II I

O3 removal rate (ppbv h1) 0.13 0.08 0.06 0.03 0.01

Source: After ref. 48. Note: All rates calculated at T ¼ 295 K for BrO ¼ 3 pptv, IO ¼ 2 pptv and HO2 ¼ 3 pptv.

Case Study VI – The potential effect of tropospheric halogen chemistry on ozone – An example of the potential effect of the ozone depletion cycles (Cycles I and II) can be assessed using data from the MBL.49 Table 7 lists the rates of ozone depletion (from Cycles I and II) in the ﬁrst hour after sunrise at Mace Head in Ireland. Measurements of

Chapter 2

[BrO] show a pulse in the ﬁrst two hours after dawn ascribed to the photolysis of inorganic bromine compounds produced either by the bromine explosion mechanism41 or the photolysis of mixed bromo/ iodo-organohalogens46 built up overnight. Using measured concentrations of BrO, IO and HO2, the data in Table 7 show the ozone depletion cycle (Cycle type I) involving the BrO and IO cross-reaction is the most important with an O3 depletion rate of 0.3 ppbv h1.

2.10 STRATOSPHERIC CHEMISTRY The stratosphere is characterised by increasing temperatures with increasing height (see Figure 1). The presence of ozone and oxygen that absorb the UV light and emit infrared radiation heats this region. Approximately, 90% of ozone molecules are found in the stratosphere (see Figure 6). The chemistry of the stratosphere, as compared to that of the troposphere, maybe described as being more chemically simple but initiated by shorter wavelength (more energetic) light. A British scientist Sydney Chapman50 suggested the basic ideas of stratospheric ozone in the 1930s, which have become known as the Chapman cycle. Short wavelength UV (hn) can dissociate molecular oxygen and the atomic oxygen fragments produced react with oxygen molecules to make ozone. O2 þ hn - O þ O M

O þ O2 ! O3

(2.88) ð2:20Þ

There are some further reactions that complete the picture, which are involved in the interconversion and removal of ozone and atomic oxygen O3 þ hn - O þ O2

(2.89)

O þ O3 - O2 þ O2

(2.90)

As O and O3 can rapidly interconvert (reactions (2.89 and 2.90)) they are often referred to as odd oxygen. Even from this simple chemistry, it is possible to see why there is an ozone layer (see Figure 6). For the production of ozone, both UV radiation (hn) and molecular oxygen are required (O2). High up in the atmosphere there is a plentiful supply of short-wavelength UV radiation but little oxygen, lower down in the atmosphere the opposite situation

Chemistry of the Atmosphere

pertains. Thus, the point where there is both enough molecular oxygen and UV light is where the maximum in the ozone concentrations are going to be found. In the terrestrial atmosphere this point is at a height of about 15–40 km, as shown in Figure 6. When the Chapman oxygenonly mechanism is applied to real atmospheric measurements, it is found that about ﬁve times as much ozone is predicted as is measured. This therefore follows that there must be some other, faster, way of removing ozone in the stratosphere. One idea put forward by Bates and Nicolet51 is that a catalytic process could be speeding up the destruction of ozone. A catalyst is something that can promote a chemical reaction without itself being consumed. In the stratosphere, a catalytic cycle like X þ O3 - XO þ O2

(2.83)

O þ XO - X þ O2

(2.91)

Net O þ O3 - O2 þ O2

(2.90)

can lead to the net destruction of ozone (reaction (2.90)). In this case the catalyst is the molecule X, which by destroying ozone is converted into XO, but the XO then reacts with an oxygen reforming X. This, the socalled chain process, can proceed many hundreds of thousands of times before termination. As to the identity of X, it is molecules like chlorine atoms (Cl), nitrogen oxide (NO) and the hydroxyl radical (OH) and therefore the corresponding XO is ClO, NO2 and HO2, i.e. Cl þ O3 - ClO þ O2

(2.92)

O þ ClO - Cl þ O2

(2.93)

These cycles are often referred to as the HOx, NOx and ClOx cycles and the groups of species as families. On the surface there is now an apparent contradiction; with only the Chapman cycle we produce too much ozone, whereas the catalytic cycles could destroy all the ozone! There are a number of cycles that can interconvert the catalytic cycles without odd oxygen removal these will be in competition with the catalytic cycles, e.g. NO þ O3 - NO2 þ O2

(2.24)

Chapter 2

NO2 þ hn - NO þ O

(2.2)

Net O3 þ hn - O2 þ O

(2.89)

The products of reaction (2.2) recombining via reaction (2.20) to reform ozone. This form of cycle (2.24, 2.2, 2.89) is often referred to as a null cycle. The fraction of the XOx tied up in null cycles is ineffective as a catalyst. Another key feature with respect to the effectiveness of catalytic cycles is the formation of reservoir species via holding cycles. At any given time about 99% of active Cl is held as reservoir species. Cl þ CH4 - HCl þ CH3

(2.94)

Cl þ H2 - HCl þ H

(2.95)

ClO þ NO2 ! ClONO2

ð2:96Þ

These reservoirs are of great importance to the chemistry of the stratosphere as they act to divert potential catalytic species from active to inactive forms, but they remain available to release the active catalysts again. HCl is the longest-lived and most abundant Cl reservoir species having a lifetime of about one month in the lower stratosphere. It is returned to active Cl largely via reaction with OH OH þ HCl - Cl þ H2O

(2.97)

ClONO2 formed via reaction (2.96) is destroyed mainly by photolysis or via reaction with O atoms, leading to regeneration of active Cl species. ClONO2 þ hn - Cl þ NO3

(2.98a)

ClONO2 þ hn - ClO þ NO2

(2.98b)

O þ ClONO2 - Products

(2.99)

The lifetime of ClONO2 is approximately 6 h in the lower stratosphere (o30 km) decreasing to about an hour at 40 km owing to the increase in UV light. Figure 25 shows the key chemical interconversions of the chlorine chemistry in the stratosphere, delineating the molecules into source, active and reservoir species. Where do these catalytic species come from and what effect does man have in increasing the levels of these catalysts? The catalytic families of

Chemistry of the Atmosphere

Chlorine Chemistry in Stratosphere

CH3Cl hν

CFCl3

ClONO2

OH O3

hν

NO2

NULL CYCLE

ClO HO2

O(3P) Loss cycles

hν

HOCl

HCl ACTIVE SPECIES RESERVOIRS SOURCE GASES O3

Figure 25 Inorganic chemistry involved in the interconversion chlorine species in the stratosphere. The relationship between source gases, active species and reservoir species is illustrated

HOx, NOx and ClOx are present in the natural atmosphere but in the contemporary atmosphere these have been supplemented by anthropogenic sources. For example, with the chlorine species the most abundant natural precursor is methyl chloride, CH3Cl. There is little interhemispheric difference in the concentration of CH3Cl indicative of a strong oceanic source, though there is a rising contribution for biomass burning as well as small contribution from volcanic emission.52 As CH3Cl has a tropospheric lifetime of a year, it can be transported to the stratosphere where on reaction with OH or via photolysis it can release the chlorine. Figure 26 shows the primary sources of chlorine species for the stratosphere separating the anthropogenic and biogenic source categories. There are a number of potential sources of anthropogenic pollution that could ﬁnd there way to the stratosphere. One of the ﬁrst problems that scientists were worried about was Concorde or stratospheric supersonic transport (SST). Aircraft engines produce large amounts of nitrogen oxides (e.g. NO) on combustion of the fuel. Injection of these nitrogen oxides directly into the atmosphere could supplement the relevant catalytic cycle plus they can be very long-lived at high altitudes. Initial estimates suggested that SST would lead to substantial losses of

Chapter 2

Primary Sources of Chlorine and Bromine for the Stratosphere in 1999 Chlorine source gases

4% 5%

3000

2000

Human-made sources

Total chlorine amount (parts per trillion)

Other gases Methyl chloroform (CH3OCl3) HCFC3 (e.g. HCFC22-CHCIF2) CFC-113 (CCl2FCCF2) Carbon tetrachloride (CCl4)

12%

CFC-11 (CCl3-F) 23%

CFC-12 (CCl2F2) 32%

1000

Methyl chloride (CH3Cl) Natural Sources

16%

Figure 26 Stratospheric source gases – A variety of gases transport chlorine into the stratosphere. These gases are emitted from natural sources and human activities. For chlorine, human activities account for most that reaches the stratosphere. The CFCs are the most abundant of the chlorine-containing gases released in human activities. Methyl chloride is the most important natural source of chlorine58

ozone. For two reasons it became apparent that this would not be the case. First, the number of aeroplanes required to observe a loss of ozone would be 1500, ﬂying 7 h a day, 7 days a week. This number of aeroplanes would be able to transport the population of London to New York, or vice versa in a matter of weeks. Second, reﬁnements to the understanding of the fundamental chemistry used to make the initial predictions have shown the impact is likely to be less. A warning note should be interposed at this point because the rise in international air travel and the new generation of aircraft means that this problem is back on the scientiﬁc agenda.53,54 The potential of subsonic aircraft to pollute the upper atmosphere is now being studied extensively. Other areas of potential stratospheric pollution have been identiﬁed such as NASA space shuttle launches and the extensive use of agricultural fertilisers.

Chemistry of the Atmosphere

A more serious culprit with respect to the depletion of the atmospheric ozone layer was identiﬁed from about 1974 onwards. US scientists postulated55 that some industrial chemical starting to reach the stratosphere might interfere with the mechanisms of ozone formation and destruction. The chemicals were chloroﬂuorocarbons, or CFCs, and had been introduced by General Motors in the 1930s mainly as coolants for refrigerators and air conditioning systems, aerosol propellants and blowing agents, etc. The perceived virtue of these compounds was their physical properties, chemical inertness and lack of toxicity. The potential problem becomes apparent at higher altitudes where the CFCs (e.g. CF2Cl2) are not so inert and can be broken down by UV light CF2Cl2 þ hn - Cl þ CF2Cl

(2.100)

releasing a chlorine atom that can enter the previously described catalytic cycles (reactions (2.92 and 2.93)) and destroy ozone. Fortuitously, as described previously, most of the active chlorine, including that provided by CFC degradation is tied up in the reservoir compounds. There are many myths about CFCs and chlorine in the atmosphere, such as if CFCs are heavier than air how can they make it high up into the atmosphere? The truth is that thousands of measurements have found CFCs high in the atmosphere and though heavier than air, the atmosphere is not still and the winds can carry and mix gases high into the stratosphere.52 Figure 26 shows the amount and variety of CFC compounds that can act as source gases for chlorine in the stratosphere. Based on predictions of the effect of CFCs on the ozone layer, in 1987 a previously unprecedented step was taken when many countries signed the UN Montreal protocol specifying the control and phase-out of these ozone-depleting chemicals.52 Since that time the protocol has been modiﬁed in order to speed up the schedule and extend the range of chemical covered to further lessen the effect of these chemicals (see Figure 27). One of the factors that lead to more rapid world action on CFCs was the discovery of the so-called Antarctic ozone hole. 2.10.1 The Antarctic Ozone Hole British scientists from the British Antarctic Survey (BAS) had been making measurements of ozone from their base at Halley Bay (761S) for many years. They detected a decline in the springtime ozone since 1977, and by October 1984, they had detected a 30% decline in the total ozone.56 Today, this value has reached 60%. The ozone hole begins to develop in August, is fully developed by early October and has normally broken up by early December (see Figure 28).

Chapter 2

Effect of the Montreal Protocol 20000

Predicted abundance (parts per million)

Effective stratospheric chlorine Montreal 1987

15000

10000 No protocol London 1990

5000 Copenhagen 1992 Beijing 1999

0 1980

2000

Zero emissions

2020

2040

2060

2080

2100

Year Figure 27 Predictions for the future abundance of effective stratospheric chlorine are shown in (1) the assuming no Protocol regulations, (2) only the regulations in the original 1987 Montreal Protocol and (3) additional regulations from the subsequent Amendments and Adjustments. The city names and years indicate where and when changes to the original 1987 Protocol provisions were agreed upon. Effective stratospheric chlorine as used here accounts for the combined effect of chlorine and bromine gases. The ‘‘zero emissions’’ line shows stratospheric abundances if all emissions were reduced to zero beginning in 200358

Figure 29 shows a recent picture of the ozone hole taken from space. Strictly speaking the use of the word ‘‘hole’’ to describe what happens to ozone in the Antarctic is an exaggeration. There is undoubtedly a massive depletion of ozone, particularly between 12 and 20 km in the Antarctic stratosphere (up to 100%) but the total column of ozone is depleted rather than removed altogether (see Figure 28). The exact location and size of the hole varies with meteorological conditions, but the area covered has increased over the past 10 years or so (see Figure 30). Currently, in the austral spring the hole extends over the entire Antarctic continent, occasionally including the tip of South America, covering an area equivalent to the North American continent (ca. 22 million km2) (see Figure 31).

Chemistry of the Atmosphere Polar Ozone Depletion

Arctic Ozone Sodankyla, Finland (67°N)

Antarctic Ozone South Pole (90°S)

Altitude (km)

Altitude (miles)

20 10

10 5

15 0

Ozone abundance (mPa)

Figure 28

Antarctic and Artic ozone distributions in the SH and NH spring periods.58 The ﬁgure shows the substantial decrease in ozone in the 14–20 km range after polar sunrise. For the Antarctic the data show there was little depletion before the 1980s. For the Arctic, there is substantial interannual variability owing to the temperature of the stratosphere

Earth Probe TOMS Total Ozone, September 16, 2000

<120

Figure 29

200 280 360 440 Ozone (Dobson Units)

520>

The 2000 ozone hole over the Antarctic seen by the satellite EP-TOMS (Image courtesy of NASA)

Chapter 2 Antarctic ozone minimum (60° - 90° S)

209 205 205 9/17

200

10/1510/10 189

1979 - 1992 Nimbus 7 TOMS 1993 - 1994 Meteor 3 TOMS 1995 (no TOMS in orbit) 1996 - 2003 Earth Probe TOMS

9/25

Mnimum ozone (DU)

180

173

169

9/19 10/18

160

159

154 10/7 10/3 146 10/24

140

140 9/9

128 124 120

120

124 10/4 117 10/10 10/7

10/5

111

10/5

100

94 10/6

88 9/28

104 99 9/24 94 90 92 9/26 9/29 9/30 9/5

97 9/28

80 1980

1985

1990

1995

2000

2005

Year

Figure 30

Lowest values of ozone measured from satellite each year in the ozone hole. Global average ozone is about 300 Dobson Units. Before 1980 ozone less than 200 Dobson Units was rarely seen. In recent years ozone near 100 Dobson Units has become normal in the ozone hole. Ozone in the year 2002 ozone hole was higher than we have come to expect because of unusually high temperatures in the Antarctic stratosphere (from NASA data)

What happens in Antarctic stratosphere to make it different from many other places? It is colder than the rest of the earth, ozone depletion occurs in the Austral spring when it is particularly cold (801C). Polar meteorology leads to very strong westerly winds (up to 100 ms1) forming a stable vortex, which prevents mixing of air from lower latitudes (see Figure 32). The vortex develops a core of very cold air and it is these temperatures that allow polar stratospheric clouds (PSCs) to form. When the sun returns in September, the temperatures begin to rise, the winds weaken and typically this drives the vortex to break down by November. In a sense the air over the pole is more or less conﬁned into what can be thought of as a giant chemical reactor. There is a slow downward circulation (see Figure 32), which draws air down through the vortex modifying this to a notional ﬂow reactor. In 1986 and 1987, a huge international expedition was mounted to examine the chemistry and meteorology that leads to the polar ozone loss. Many different measurements were made using satellites, aeroplanes and ground-based instruments.57 The results showed some stark

Chemistry of the Atmosphere Average area of Ozone Hole

Size (million sq km)

area of North America

area of North Antarctica

area of ozone <220 DU average over 30 days vertical line = minimum and maximum area

Year

Figure 31 Growth of the average area of the ozone hole from 1979 to 2003. The ozone hole is deﬁned as the area for which ozone is less than 220 Dobson Units, a value rarely seen under normal conditions. It shows that the ozone in this area hardly occurred at all in 1980, but by year 2000 covered an area larger than North America, 26.5 million km2 (from NASA data)

contrasts between the chemistry inside and outside the vortex before and after sunrise in the Austral spring (see Figure 33). Within the vortex concentrations of e.g. ClO were a factor of 10 larger than outside. These measurements also showed that the stratosphere within the disturbed region was abnormally dry and highly deﬁcient in nitrogen oxides. The reason for the dehydration and denitriﬁcation of the Antarctic stratosphere is the formation of the PSCs, whose chemistry perturbs the composition in the Antarctic stratosphere. Polar stratospheric clouds can be composed of small (o1 mm diameter) particles rich in HNO3 or at lower temperatures (o190 K) larger (10 mm) mainly ice particles. These are often split into two categories, the so-called Type I PSC, which contains the nitric acid either in the form of liquid ternary solutions with water and sulfuric acid or as solid hydrates of nitric acid, or Type II PSCs made of ice particles. The ice crystals on these clouds provide a surface for reactions such as ClONO2 (g) þ HCl (ice) - Cl2 (g) þ HNO3 (ice)

(2.101)

Chapter 2

Figure 32 The Winter vortex over Antarctica. The cold core is almost isolates from the rest of the atmosphere, and acts as a reaction vessel in which chlorine activation may take place in the polar night (after ref. 3)

ClONO2 (g) þ H2O (ice) - HOCl (g) þ HNO3 (ice) (2.102) HOCl (g) þ HCl (ice) - Cl2 (g) þ H2O (ice)

(2.103)

N2O5 (g) þ H2O (ice) - 2HNO3 (ice)

(2.104)

In reactions (2.101 and 2.102), the normally inactive reservoir for Cl and ClO is converted into molecular chlorine. It is worth noting that the HNO3 is left in the ice-phase. These reactions can continue all winter long without light converting inactive chlorine to active chlorine. The two main effects therefore of the heterogeneous chemistry involving

Chemistry of the Atmosphere Observations of stratospheric constituents from ER2 aircraft

3.5 ppmv 10 ppbv

Chemically perturbed region

H2O NOy

~0.75 ppbv 250 ppbv

N2O

CFC11

ClO

~5 pptv

BrO 60°S

65°S

70°S

Polewards Figure 33 Schematic representations of the changes in the concentration of some stratospheric species on entering the chemically perturbed region of the Antarctic vortex (after ref. 57)

PSCs are the release of active chlorine from its usual reservoirs and the removal of NOx from the system by the production of HNO3. When the sun rises in the spring it triggers the depletion chemistry. Cl2 þ hn - Cl þ Cl

(2.105)

Cl þ O3 - ClO þ O2

(2.92)

But there is one further unforeseen modiﬁcation as the concentration of oxygen atoms is too low ﬁrst thing in spring to regenerate Cl via the normal reaction (2.93), so there is instead a further low-temperature reaction that regenerates the chlorine atoms ClO þ ClO þ M - (ClO)2 þ M

(2.106)

(ClO)2 þ hn - Cl þ ClOO

(2.107)

ClOO þ M - Cl þ O2 þ M

(2.108)

that can re-enter the catalytic cycle and destroy large amounts of ozone. The absence of significant concentrations of NOx stops rapid

Chapter 2

reformation of the ClOx reservoir species (e.g. via reaction (2.96)). The ClO self-reaction ozone destruction cycle is supplemented by a coupled ClO/BrO cycle. ClO þ BrO - Cl þ Br þ O2

(2.109)

ClO þ BrO ! BrCl þ O2 BrCl þ hv ! Cl þ Br

ð2:110; 2:111Þ

Cl þ O3 - ClO þ O2

(2.92)

Br þ O3 - BrO þ O2

(2.76)

The stratospheric bromine comes from a range of source gases including CH3Br.58 The contribution of bromine to ozone loss in the polar regions has increased faster than that of chlorine because abundances of bromine continue to increase at a time when those of chlorine are levelling oﬀ (see Figure 27).52 Cycles (2.105, 2.92, 2.106–2.108) and (2.76, 2.92, 2.109–2.111) account for most of the ozone loss observed in the Antarctic stratosphere in the late winter/spring season. At high ClO abundances, the rate of ozone destruction can reach 2–3% day1 in late winter/spring. Figure 34 shows a summary of the photochemistry and dynamics in the polar stratosphere, illustrating the time proﬁles of the key chlorine species coupled to the temperature requirements in the vortex. Figure 35 shows measurements of a range of chemical species and temperature in the 2004 ozone hole. The data reﬂects the main chemical and physical features on the ozone hole. In summary, the key features of Antarctic ozone loss are (i) The circulating winds in the polar regions enable the formation of a stable vortex which provides a gigantic ‘‘reaction vessel’’ for ozone depletion to occur within. (ii) The low temperatures in the vortex encourage the formation of PSCs, which enhance the production of active chlorine species. (iii) Pre-conditioning of the atmosphere takes place during the polar winter releasing chlorine from reservoir molecules. (iv) At sunrise, molecular chlorine is dissociated into free atoms that can destroy the ozone. (v) Chlorine atoms are regenerated by the dimer reactions.

Chemistry of the Atmosphere

Abundance

Inactive chlorine

Surface reactions

Active chlorine

Gas-phase reactions

Inactive chlorine

ClONO2

ClONO2 (Cl2+ClO+(ClO)2) HCl

Autumn

HCl

Early

Winter

Late

Spring

Time

Denitrification and dehydration

Temperature

Surface processing Cl-catalysed O3 destruction Formation, Cooling, descent

Maximum intensity

Inactive chlorine recovery Breakup Polar vortex evolution Time

Figure 34 Photochemistry and dynamics in the polar stratosphere (after ref. 65)

Figure 35 Chemical species and temperature as measured by satellite in the 2004 ozone hole (image courtesy of NASA). The white line on the ozone ﬁgure gives the extent of the vortex

Chapter 2

(vi) The removal of nitric acid (and other oxides of nitrogen), by slow, gravitational settling means that the reservoir species cannot be regenerated. The HNO3 is liberated within the particle. In more recent times, there have been discoveries of ozone depletion in the Arctic that occur by similar mechanisms as the ones described here (see Figure 28). The Arctic equivalent does not tend to be as dramatic owing to the fact the Artic stratosphere does not get as cold as the Antarctic, mainly owing to a less well-formed vortex, largely owing to northern hemisphere topography.52 2.11 SUMMARY The chemistry of the atmosphere is diverse, driven in the main by the interaction of light with a few molecules that drives a complex array of chemistry. The type and impact of atmospheric chemistry varies in concert with the physical and biological change throughout the atmosphere. An integral understanding of atmospheric chemistry within the earth system context underpins many contemporary global environmental problems and is therefore vital to sustainable development. This chapter has outlined the main features of the chemistry of the troposphere and stratosphere. In some senses it must be remembered that our understanding of atmospheric chemistry is still evolving, but the science expressed gives entry into the rich and changing world of atmospheric science. In the coming years we need to understand the interaction of atmospheric chemistry with climate as well as better understand the multi-phase impacts of the chemistry. There is still much to discover, rationalise and understand in atmospheric chemistry. QUESTIONS (i) The earth system is made up of the atmosphere, hydrosphere, lithosphere and biosphere. Describe, with examples, the chemical interactions between these different elements of the earth system. (ii) The amount of any chemical substance in the atmosphere is dependent on the source–transport–sink relationship for that compound. Describe the different types of sources of trace gases in the atmosphere. (iii) Discuss the following statement: ‘‘The atmosphere can be split into different regions according to different physical or chemical properties’’.

Chemistry of the Atmosphere

(iv) Explain the following statement: ‘‘Ozone can be photolysed both in the troposphere and stratosphere leading to quite different chemistry’’. (v) Why is the stratosphere separate from the troposphere? (vi) Discuss the chemistry of ozone in the stratosphere in terms of the relationships between catalytic species and the different types of reservoir chemistry. (vii) Explain the main chemical and physical features of the so-called ozone ‘‘hole’’. (viii) The hydroxyl radical (OH) has different role in the stratosphere and troposphere. Give an example of the reactions of OH in troposphere and stratosphere and describe their significance. (ix) What role does the boundary layer play in the mixing of gases? (x) What are the main reactions that contribute to the formation of the so-called photochemical smog? (xi) What physical, chemical and social effects can aggravate the formation of photochemical smog? (xii) Peroxyacetylnitrate (PAN) is of interest as a characteristic product of tropospheric photochemistry. (i) How is PAN formed in the atmosphere? (ii) The unimolecular decomposition of PAN to peroxy acetyl radicals and NO2 is strongly temperature dependent. The following expression gives the temperature dependence of the unimolecular rate constant: 14; 000 k ¼ 1:0 1017 exp : T If the lifetime is given by 1/k, calculate the atmospheric lifetime of PAN at T¼310, 298, 290 and 280 K. (iii) In the light of your answer to part (b), comment on the potential of PAN as a reservoir for reactive nitrogen in remote regions. (xiii) Explain the following: (i) In the shadow of a thick cloud the concentration of OH falls to nearly zero. (ii) On a global scale the concentration distribution of CO2 is about the same whereas the concentration of isoprene varies from one location to the next. (iii) There is more than one chemical pathway to the initiation of night-time chemistry.

Chapter 2

(iv) Both HONO and HCHO are detected in abundance in areas suffering from photochemical smog and may lead to the acceleration of the production of smog. (v) The effect that NOx from aeroplanes and lightning can have on ozone can be quite different from NOx emitted from the surface of the earth. (xiv) The ﬁgure below shows two sets of ozone proﬁles taken over Halley Bay in Antarctica during Winter/Spring. (i) What measurement platforms are normally used to take an ozone proﬁle? (ii) Why under normal conditions (August 15th) is the ozone maximum at 20 km? (iii) Describe the key processes that have contributed to the loss of ozone by October 13th. (iv) What molecules are thought to be responsible for the ozone loss shown in the ﬁgure?

(xv) The ozone continuity equation below is a reasonable representation of the processes that controls ozone in the boundary layer, viz

Chemistry of the Atmosphere

d½O3 Ev ð½O3 ft ½O3 Þ vd ½O3 ¼Cþ þ dt H H C is a term representative of the photochemistry (production or destruction), Ev the entrainment velocity, [O3]ft the concentration of free tropospheric ozone, vd the dry deposition velocity and H the height of the boundary layer. The ozone budget shown in Table 8 has been calculated both in the summer and winter using the ozone continuity equation for a site in the marine boundary layer. (i) Describe the chemistry involved in calculating the term C in the ozone continuity equation. (ii) Given the magnitude of the photochemical terms in Table 8, comment on the amount of available NOx. (iii) How does the ability to make or destroy ozone change with season? (iv) What factors affect the magnitude of a deposition velocity? (v) What role does entrainment play in the boundary layer? (vi) What are the weaknesses of using this approach for the calculation of chemical budgets? (vii) Predict the likely diurnal cycle of ozone and hydrogen peroxide in the remote marine boundary layer. (xvi) Nitric acid has many potential loss processes in the troposphere. Given that the reaction OH þ HNO3 ! H2 O þ NO3 has a rate constant of 2.0 1013 cm3 molecule1 s1 and the hydroxyl radical concentration is 4 106 molecule cm3. The

Table 8

Calculated average ozone removal and addition rates according to pathway, ppbv day1 (upper part) and fractional contributions to overall production or destruction patyways (lower part) O3 removal

O3 addition

Pathway

Photochemistry (%)

Deposition (%)

Photochemistry (%)

Entrainment (%)

Summer Winter

1.19 0.61

0.18 0.35

0.56 0.29

2.1 0.1

Summer Winter

87 64

13 36

21 74

79 26

76 Table 9

Chapter 2 Photoprocess data for HNO3

Wavelength (nm) 295–305 305–315 315–325 325–335 335–345 345–355

F(l) (photon cm2 s1) 2.66 4.20 1.04 1.77 1.89 2.09

10 1014 1015 1015 1015 1015

s (cm2 molecule1) 0.409 0.146 0.032 0.005 0 0

10 1020 1020 1020

f 1 1 1 1 1 1

deposition velocity for HNO3 is 0.005 cm s1 over a 1 km wellmixed boundary layer and the photoprocess data given in Table 9. What is the atmospheric lifetime of HNO3?

REFERENCES 1. A.K. Smith, J. Atmos. Sol.-Terr. Phy., 2004, 66, 839. 2. J.M.C. Plane, Chem. Rev., 2003, 103, 4963. 3. R.P. Wayne, Chemistry of Atmospheres, 3rd edn, Oxford University Press, UK, 2000. 4. R.A. Cox and R.G. Derwent, Specialist Periodical Reports Chem. Soc., 1981, 4, 189. 5. A.V. Jackson, in Handbook of Atmospheric Science, Blackwell publishing, Oxford, 2003, Chapter 5. 6. R.R. Stull, An Introduction to Boundary Layer Meteorology, Kluwer Academic Publishers, London, UK, 1988. 7. C.E. Junge, Tellus, 1962, 14, 363. 8. P. Fabian and P.G. Pruchniewz, J. Geophys. Res., 1977, 82, 2063. 9. P.J. Crutzen, Pure Appl. Geophys., 1973, 106–108, 1399. 10. W.L. Chameides and J.C.G. Walker, J. Geophys. Res., 1973, 78, 8760. 11. A. Volz and D. Kley, Nature, 1988, 332, 240. 12. IPCC, Climate Change 2001: The Scientiﬁc Basis, Cambridge University Press, Cambridge, UK, 2001. 13. A.M. Thompson, Science, 1992, 256, 1157. 14. P.S. Monks, Chem. Soc. Rev., 2005, 34, 376. 15. P.A. Leighton, The Photochemistry of Air Pollution, Academic Press, New York, 1961. 16. J.M. Roberts, Reactive odd-nitrogen (NOy) in the atmosphere, in Composition and Climate of the Atmosphere, H.B. Singh (ed), Van Nostrand Reinhold, New York, 1995. 17. J. Lelieveld and P.J. Crutzen, Nature, 1990, 343, 227. 18. S. Sillman, Atmos. Environ., 1999, 33, 1821.

Chemistry of the Atmosphere

19. L. Jaegle´, D.J. Jacob, W.H. Brune, I.C. Faloona, D. Tan, Y. Kondo, G.W. Sasche, B. Anderson, G.L. Gregory, S. Vay, H.B. Singh, D.R. Blake and R. Shetter, Geophys. Res. Lett., 1999, 26, 3081. 20. (a) G.P. Ayers, S.A. Penkett, R.W. Gillett, B.J. Bandy, I.E. Galbally, C.P. Meyer, C.M. Elsworth, S.T. Bentley and B.W. Forgan, Nature, 1992, 360, 446–448; (b) P.S. Monks, G. Salisbury, G. Holland, S.A. Penkett and G.P. Ayers, Atmos. Environ., 2000, 34, 2547–2561. 21. C.N. Hewitt (ed), Reactive Hydrocarbons in the Atmosphere, Academic Press, London, 1999. 22. N. Poisson, M. Kanakidou and P.J. Crutzen, J. Atmos. Chem., 2000, 36, 157. 23. S. Houweling, F. Dentener and J. Lelieveld, J. Geophys. Res. (Atmos.), 1998, 103, 10673. 24. A.J. Haagen-Smit, Ind. Eng. Chem., 1952, 44, 1342. 25. G. Lammel and J.N. Cape, Chem. Soc. Rev., 1996, 25, 361. 26. H.B. Singh, D. O’Hara, D. Herlth, J.D. Bradshaw, S. Sandholm, G.L. Gregory, G.W. Sachse, D.R. Blake, P.J. Crutzen and M. Kanakidou, J. Geophys. Res., 1992, 97, 16511. 27. B.J. Allan, N. Carslaw, H. Coe, R.A. Burgess and J.M.C. Plane, J. Atmos. Chem., 1999, 33, 129. 28. R.P. Wayne, I. Barnes, P. Biggs, J.P. Burrows, C.E. Canosa-Mas, J. Hjorth, G. LeBras, G.K. Moortgat, D. Perner, G. Poulet, G. Restelli and H. Sidebottom, Atmos. Environ., 1991, 25, 1. 29. B.J. Allan, G. McFiggans, J.M.C. Plane, H. Coe and G.G. McFadyen, J. Geophys. Res., 2000, 105, 24191. 30. R. Criegee, Agnew. Chem. Internat. Edit., 1975, 14, 745. 31. J.G. Calvert, R. Atkinson, J.A. Kerr, S. Madronich, G.K. Moortgat, T.J. Wallington and G. Yarwood, The Mechanisms of Atmospheric Oxidation of the Alkenes, Oxford University Press, Oxford, 2000. 32. S.E. Paulson, M.Y. Chung and A.S. Hasson, J. Phys. Chem. A, 1999, 103, 8125. 33. G. Salisbury, A.R. Rickard, P.S. Monks, B.J. Allan, S. Bauguitte, S.A. Penkett, N. Carslaw, A.C. Lewis, D.J. Creasey, D.E. Heard, P.J. Jacobs and J.D. Lee, J. Geophys. Res. (Atmos.), 2001, 106, 12669. 34. EMEP, in EMEP Assessment, Part I: European Perspective, G. Lo¨vblad, L. Tarraso´n, K. Tørseth and S. Dutchak (eds), Oslo, October, 2004. 35. M. Eisinger and J.P. Burrows, Geophys. Res. Lett., 1998, 22, 4177.

Chapter 2

36. B.J. Finlayson-Pitts and J.N. Pitts Jr., Chemistry of the Upper and Lower Atmosphere, Academic Press, San Diego, 2000. 37. U. Platt and G. Honniger, Chemosphere, 2003, 52, 325–338. 38. L.A. Barrie, J.W. Bottenheim, R.C. Schnell, P.J. Crutzen and R.A. Rasmussen, Nature, 1988, 334, 138. 39. S.M. Fan and D.J. Jacob, Nature, 1992, 359, 522. 40. K. Heibestreit, J. Stutz, D. Rosen, V. Matveiv, M. Peleg, M. Luria and U. Platt, Science, 1990, 283, 55. 41. R. Vogt, P.J. Crutzen and R. Sander, Nature, 1996, 383, 327. 42. G.P. Ayers, R.W. Gillett, J.M. Cainey and A.L. Dick, J. Atmos. Chem., 1999, 33, 299. 43. R. Sander, W.C. Keene, A.A.P. Pzszenny, R. Arimoto, G.P. Ayers, E. Baboukas, J.M. Cainey, P.J. Crutzen, R.A. Duce, G. Ho¨nninger, B.J. Huebert, W. Maenhaut, N. Mihalopoulos, V.C. Turekian and R. Can Dingenen, Atmos. Chem. Phys., 2003, 3, 1301. 44. (a) T. Wagner and U. Platt, Nature, 1998, 395, 486–490; (b) T. Wagner, C. Leue, M. Wenig, K. Pfeilsticker and U. Platt, J. Geophys. Res., 2001, 106, 24225–24235. 45. M. Van Roozendael, C. Fayt, J.-C. Lambert, I. Pundt, T. Wagner, A. Richter and K. Chance, Development of a bromine oxide product from GOME, Proceedings of ESAMS’ 99-European Symposium on Atmospheric Measurements from Space, ESTEC, Noordwijk, The Netherlands, 18–22 January 1999, WPP-161, 1999, 543–547. 46. L.J. Carpenter, Chem. Rev., 2003, 103, 4953. 47. L.J. Carpenter, K. Hebestreit, U. Platt and P.S. Liss, Atmos. Chem. Phys., 2001, 1, 9–18. 48. A. Saiz-Lopez and J.M.C. Plane, Geophys. Res. Lett., 2004, 31, L04112. 49. A. Saiz-Lopez, J.M.C. Plane and J.A. Shillito, Geophys. Res. Lett., 2004, 31, L03111. 50. S. Chapman, Mem. Roy. Meterol. Soc., 1930, 3, 103. 51. D.R. Bates and M. Nicolet, J. Geophys. Res., 1950, 55, 301. 52. WMO, Scientiﬁc assessment of ozone depletion: 2002, Global Ozone Research and Monitoring Project, Report No. 47, Geneva, 2003. 53. IPCC, Aviation and the Global Atmosphere, Special Report, Intergovernmental Panel on Climate Change, 1999. 54. S.R. Kawa, J.G. Anderson, S.L. Baughcum, C.A. Brock, W.H. Brune, R.C. Cohen, D.E. Kinnison, P.A. Newman, J.M. Rodriguez, R.S. Stolarski, D. Waugh and S.C. Wofsy, Assessment of the effects of high-speed aircraft in the stratosphere: 1998, NASA Tech. Pub. 1999-209236, NASA Goddard Space Flight Center, Greenbelt, MD, 1999.

Chemistry of the Atmosphere

55. M.J. Molina and F.S. Rowland, Nature, 1974, 249, 810. 56. J.C. Farman, B.G. Gardiner and J.D. Shanklin, Nature, 1985, 315, 207. 57. J.G. Anderson, D.W. Toohey and W.H. Brune, Science, 1991, 251, 39. 58. WMO, Twenty questions and answers about the ozone layer, in WMO, Scientiﬁc Assessment of Ozone Depletion: 2002, Global Ozone Research and Monitoring Project, Report No. 47, Geneva, 2003. 59. D. Ehhalt, Phys. Chem. Chem. Phys., 1999, 24, 5401. 60. J.N. Howard, J.I.F. King and P.R. Gast, Thermal radiation, Handbook of Geophysics, Macmillan, New York, 1960. 61. P.S. Monks, A.R. Rickard, S.L. Hall and N.A.D. Richards, J. Geophys. Res., 2004, 109, D17206. 62. P.D. Lightfoot, R.A. Cox, J.N. Crowley, M. Destriau, G.D. Hayman, M.E. Jenkin, G.K. Moortgat and F. Zabel, Atmos. Environ., 1992, 10, 1805. 63. D. Lamb, D.F. Miller, N.F. Robinson and A.W. Gertler, Atmos. Environ., 1987, 21, 2333. 64. U. Frieß, Spectroscopic measurements of atmospheric trace gases at Neumayer station, Antarctica, Ph.D. Thesis, University of Heidelberg, Germany, 2001. 65. WMO, Scientiﬁc assessment of ozone depletion: 1998, Global Ozone Research and Monitoring Project, Report No. 44, Geneva, 1999.

CHAPTER 3

Chemistry of Freshwaters MARGARET C. GRAHAM AND JOHN G. FARMER School of GeoSciences, University of Edinburgh, Crew Building, King’s Buildings, West Mains Road, Edinburgh EH9 3JN

3.1 INTRODUCTION Water is essential for all life on Earth. As the Earth’s population continues to increase rapidly, the growing human need for freshwater (e.g. for drinking, cooking, washing, carrying wastes, cooling machines, irrigating crops, receiving sewage and agricultural runoﬀ, recreation, and industrial purposes) is leading to a global water resources crisis. Rees1 has commented that ‘there is a growing consensus that if current trends continue, water scarcity and deteriorating water quality will become the critical factors limiting future economic development, the expansion of food production, and the provision of basic health and hygiene services to millions of disadvantaged people in developing countries’. In 2003, the UN International Year of Freshwater, 20% of the world’s people lacked good drinking water and 40% lacked adequate sanitation facilities.2 The inventory of water at the Earth’s surface (Table 1)3 shows that the oceans, ice caps, and glaciers contain 98.93% of the total, with groundwater (1.05%) accounting for most of the rest, and lakes and rivers amounting to only 0.009% and 0.0001%, respectively.3 The average annual water withdrawal for use by humans is currently about 4000 km3, of which 69% is used by agriculture, 21% by industry/power, and 10% for domestic purposes.2 Figures vary widely between continents and countries. On a per capita basis, North America withdraws seven times as much freshwater as Africa. The USA (41% agriculture, 46% industry/ power, 13% domestic) withdraws almost 20% of its renewable freshwater resources each year, in comparison with 7% for the UK (3%, 75%, 22%), 22% for China (68%, 26%, 7%), 33% for Poland (8%, 79%, 13%), 51% 80

Chemistry of Freshwaters Table 1

Inventory of water at the Earth’s surface

Reservoir Oceans Mixed layer Thermocline Abyssal Ice caps and glaciers Groundwater Lakes Rivers Soil moisture Atmosphere totala Terrestrial Oceanic Biosphere Approximate total

Volume 106 km3 (1018 kg) 1400 50 460 890 43.4 15.3 0.125 0.0017 0.065 0.0155 0.0045 0.0110 0.002

Per cent of total 95.96

2.97 1.05 0.009 0.0001 0.0045 0.001 0.0001

1459

Source: Global Environment: Water Air Geochemical Cycles by Berner/Berner, r 1996. Adapted by permission of Prentice-Hall, Upper Saddle River, NJ.3 a As liquid volume equivalent of water vapour.

for India (86%, 5%, 8%), and 76% for Bangladesh (96%, 1%, 3%).2 Many developing countries are using up to 40% of their renewable freshwater for irrigation.2 As a consequence of population growth, pollution and climate change, the average supply of water per person is likely to drop by one-third over the next two decades,4 with a global supply crisis projected to occur between 2025 and 2050, although much earlier for some individual countries.1,2,4 The quality of freshwater is also of concern. For example, about 75% of Europe’s drinking water is taken from groundwater,2 a resource described as overexploited in almost 60% of European industrial and urban centres and threatened by pollutants.5 River and lake eutrophication caused by excessive phosphorus and nitrogen from agricultural, domestic, and industrial eﬄuents is a problem across most of Europe. The Eastern European countries in particular fare badly when judged against the criteria of pathogenic agents, organic matter, salinization, nitrate, phosphorus, heavy metals, pesticides, acidiﬁcation, and radioactivity, and even the Nordic countries of Norway, Sweden, and Finland, despite having generally better freshwater conditions than the rest of Europe, suﬀer from acidiﬁcation.5 In Asia, 86% of all urban wastewater and 65% of all wastewater is discharged, untreated, into the aquatic environment.6 There is also the unexpected and insidious chemical contamination of apparently clean drinking water supplies, exempliﬁed most notably by the arsenic contamination of groundwater that aﬀects millions in Bangladesh and other Asian countries.2

Chapter 3

The eﬀects of contaminants or pollutants on freshwater depend upon their chemical, physical, and biological properties, their concentrations and duration of exposure. Aquatic life may be aﬀected indirectly (e.g. through depletion of oxygen caused by biodegradation of organic matter) or directly, through exposure to toxic or carcinogenic chemicals, some of which may accumulate in organisms. Such toxicity may, however, be modiﬁed by the presence of other substances and the characteristics of the particular water body. For example, metal toxicity is aﬀected by pH, which inﬂuences speciation; dissolved organic carbon has been shown to reduce bioavailability by forming metal complexes and by subsequent adsorption to particulate matter in freshwater.7 The toxicity of many heavy metals to ﬁsh is also inversely related to water hardness, with Ca21 competing with free metal ions for binding sites in biological systems. In biologically productive lakes which also develop a thermocline in summer, the bottom water may become depleted in oxygen, leading to a change from oxidizing to reducing conditions. As a result, the potential exists for remobilization of nutrients and metals from bottom sediments to the water column, one more example of how alterations in master variables in the hydrological cycle, in this case the redox status, can aﬀect the fate and inﬂuence of pollutants. A further demonstration of the importance of fundamental properties of both pollutants and water bodies is provided by the behaviour of chemicals upon reaching a groundwater aquifer. Soluble chemicals, such as nitrate, move in the same direction as groundwater ﬂow. A poorly soluble liquid which is less dense than water, such as petrol, spreads out over the surface of the water table and ﬂows in the direction of the groundwater. Poorly soluble liquids which are denser than water, such as various chlorinated solvents, sink below the water table and may ﬂow separately along low permeability layers encountered at depth in the aquifer and not necessarily in the same direction as that of the overlying groundwater.7 The rest of this chapter therefore consists of two major sections, ﬁrst on fundamental aquatic chemistry of relevance to the understanding of pollutant behaviour in the aquatic environment and second on associated case studies and examples drawn from around the world, including reference to water treatment methods where appropriate. 3.2 FUNDAMENTALS OF AQUATIC CHEMISTRY 3.2.1 Introduction 3.2.1.1 Concentration and Activity. All natural waters contain dissolved solutes. In order to understand their chemical behaviour in

Chemistry of Freshwaters

aqueous systems, it is ﬁrst necessary to consider the relationship between solute concentration and solute activity. The activity of a solute is a measure of its observed chemical behaviour in (aqueous) solution. Interactions between the solute and other species in solution lead to deviations between solute activity {i} and concentration [i]. An activity coeﬃcient, gi, is therefore deﬁned as a correction factor, which interrelates solute activity and concentration. {i} ¼ gi [i]

(3.1)

Concentration and activity of a solute are only the same for very dilute solutions, i.e. gi approaches unity as the concentration of all solutes approaches zero. For non-dilute solutions, activity coeﬃcients must be used in chemical expressions involving solute concentrations.w Although freshwaters are suﬃciently dilute to be potable (containing less than about 1000 mg L1 total dissolved solids (TDS)),8 it cannot be assumed that activity coeﬃcients are close to unity. Calculation of activities using the expression in Equation (3.1) would appear to be straightforward, requiring only a value for the activity coeﬃcient and the concentration of the solute, i. The activity coeﬃcients for individual ions, gi, cannot, however, be measured. The presence of both cations and anions in solution means that laboratory experiments quantifying either the direct (e.g. salt solubility) or indirect (e.g. elevation of boiling point, depression of freezing point, etc.) eﬀects of the presence of solutes in aqueous solution generally lead to the mean ion activity coeﬃcients, g. Traditionally, single-ion activity coeﬃcients have been obtained from the mean values, e.g. using the mean-salt method (Example 3.1),9,10 but this approach has been shown to be inadequate because it ignores ion pairing eﬀects at higher solute concentrations.11,12 This is generally more signiﬁcant for anions than for cations because the extent of anion association at higher solute concentration is much greater.13

Example 3.1: Determination of single-ion activity coeﬃcients for Ca21 from mean values for CaCl2 using the mean-salt method. The solubility of calcium chloride in water can be expressed as CaCl2(s) " Ca2þ KSP (aq) þ 2Cl(aq) w

Concentration is expressed on the molar scale in terms of mol L1 of solvent or on the molal scale in terms of mol kg1 of solvent. The molal scale gives concentrations that are independent of temperature and pressure. In this chapter, the molar scale will be used on the basis that molarity and molality are almost identical for the low ionic strengths commonly associated with freshwaters.

Chapter 3

where the solubility product, KSP ¼ Ca2þ fCl g2 ¼ gCa2þ g2Cl ½Ca2þ ½Cl 2 We deﬁne the mean ion activity coeﬃcient as the geometric mean of the single-ion activity coeﬃcients gCaCl2 ¼ ðgCa2þ gCl gCl Þ1=3 The mean ion activity coeﬃcient values can be obtained from experiments where the eﬀect of electrolyte concentration on the KSP value for a salt is determined. The mean values are then compared with those for KCl under the same solution conditions. The single-ion activity coeﬃcient for Ca21 can then be computed if an assumption is made about the individual values for K1 and Cl. These ions have the same magnitude of charge and similar electronic conﬁguration, ionic radii, and ionic mobilities. On the basis of these properties, the MacInnes convention (1919) states that gK1 ¼ gCl ¼ gKCl In general: gcation ¼ ðgcation Clx1 Þx =ðgKCl Þx1 Thus gCaCl2 ¼ ðgCa2þ g2Cl Þ1=3 ¼ ðgCa2þ g2KCl Þ1=3 and gCa2þ ¼ ðgCaCl2 Þ3 =ðgKCl Þ2 For I ¼ 0.1 mol l1: gCa2þ ¼ ð0:518Þ3 =ð0:770Þ2 ¼ 0:234 (Mean salt data from Robinson and Stokes.9)

Although other methods using such experimental data have been documented,14 theoretical and/or empirical considerations have led to various expressions involving a dependence on ionic strength, I, which is a major factor inﬂuencing the activity of solutes in aqueous solution. 3.2.1.2 Ionic Strength. Ionic strength, I, is a measure of the charge density in solution and is deﬁned as X 2 ½ z ð3:2Þ i I ¼ 0:5 i i

Chemistry of Freshwaters

where [i] is the concentration of solute ion, i; zi is the integral charge associated with solute ion, i. Ionic strength, expressed in terms of mol L1, is most accurately determined by carrying out a total water analysis, i.e. quantiﬁcation of the concentrations of all ionic species in solution and calculation of I using Equation (3.2). It may, however, be estimated from measurements of either the TDS in solution or, preferably, the speciﬁc conductance (SpC) of the solution, e.g. I ¼ 2.5 105 TDS (mg L1) or I ¼ 1.7 105 SpC (mS cm1).12 The simplest theoretical relationship between I and gi is expressed in the Debye–Hu¨ckel (DH) Equation (3.3).z The DH model assumes that ions can be represented as point charges, i.e. of inﬁnitely small radius, and that long-range coulombic forces between ions of opposite charge are responsible for the diﬀerences between the observed chemical behaviour, i.e. activity, and the predicted behaviour on the basis of solute concentration. log gi ¼ 0:5 z2i log gi ¼ 0:5 z2i

p p I = 1 þ 0:33 ai I

log gi ¼ 0:5 z2i

DH equation

ð3:3Þ

Extended DH equation ð3:4Þ

p p ¨ I = 1 þ I Guntelberg approximation

ð3:5Þ

An inﬁnitely dilute solution is deﬁned as having an ionic strength o105 mol L1 and activity coeﬃcients calculated by any of the above equations would have a value of unity. For freshwaters where I is in the range 105–102.3 mol L1, activity coeﬃcients can be calculated using the DH Equation (3.3). Where I is in the range 102.3–101 mol L1, however, the values calculated using Equation (3.3) diﬀer from experimental data and this stems from the assumption that the solute ions are point charges. The extended DH expression (3.4) includes a parameter, ai, which is related to ion size. Example 3.2 illustrates the use of both DH and extended DH expressions to calculate single-ion activity coeﬃcients. z

The general forms of the DH and extended DH equations are log gi ¼ Azi2OI and log gi ¼ Azi2(OI/(1 þ BaiOI)), respectively, where A and B are temperature dependent constants. A ¼ 1.824928 106 r00.5 (eT)1.5 B0.5 and B ¼ 50.3 (eT)0.5 B0.33 at 298 K, where r0 is the density and e is the dielectric constant of water at 298 K.

Chapter 3

Example 3.2: Calculation of single-ion activity coeﬃcients using the DH and extended DH equations. If a solution contains 0.01 mol L1 MgSO4, 0.006 mol L1 Na2CO3, and 0.002 mol L1 CaCl2 then the ionic strength of solution is I ¼ 0:5ð0:01 ðþ2Þ2 þ 0:01 ð2Þ2 þ 0:006 2 ðþ1Þ2 þ 0:006 ð2Þ2 þ 0:002 ðþ2Þ2 þ 0:002 2 ð1Þ2 Þ ¼ 0:5ð0:04 þ 0:04 þ 0:012 þ 0:024 þ 0:008 þ 0:004Þ ¼ 0:064molL1 Calculating the single-ion activity coeﬃcient for Na1 using both the DH and the extended DH equations (use a value of 4.5 for the ion size parameter, ai) p DH log gNaþ ¼ 0:5ðþ1Þ2 ð0:064Þ ¼ 0:126 gNaþ ¼ 0:747 p Ex-DH log gNaþ ¼ 0:5ðþ1Þ2 ð0:064Þ p =ð1 þ 0:33 4:5 ð0:064ÞÞ ¼ 0:0919

gNaþ ¼ 0:809

At I ¼ 0.064 mol L1, the single-ion activity coeﬃcient calculated by the DH model is B8% lower than that calculated by the extended DH model. Now consider a solution which has an ionic strength of only 0.001 mol L1. DH

log gNaþ ¼ 0:0158

gNaþ ¼ 0:964

Ex DH

log gNaþ ¼ 0:0151

gNaþ ¼ 0:966

Clearly, at lower ionic strength, the single-ion activity coeﬃcient is much closer to unity. Also, the DH and the extended DH models give almost exactly the same value. This is because the denominator (1þ0.33aiOI) of the extended DH equation approaches a value of unity, i.e. 0.33aiOI approaches zero, as the ionic strength decreases towards zero. In other words, the two equations become identical at very low ionic strength. It should be noted that, for solutions of ionic strength o101 mol L1, either Equation (3.4) or the Gu¨ntelberg approximation (3.5), which incorporates an average value of 3 for ai, can be used.13 The Gu¨ntelberg approximation is particularly useful in calculations where a number of ions are present in solution or when values of the ion size parameter are poorly deﬁned.

Chemistry of Freshwaters

For solutes in higher ionic strength solutions, however, Equations (3.3)–(3.5) tend to underestimate activity coeﬃcients in part because of the assumptions that are inherent in the DH and extended DH Equations, e.g. only long-range coulombic interactions, ion size parameters are independent of ionic strength, ions of like charge do not interact, each ion can construct its own atmosphere of ions of opposite charge.8,12,15,16 Additional factors are that (i) short-range interactions (including those between ions of like charge and also between ions and neutral molecules),12,15–17 rather than the long-range coulombic interactions, between ions become increasingly important and (ii) the number of water molecules involved in ion hydration spheres means that the activity of water decreases to a value signiﬁcantly less than unity.12 In order to compensate particularly for the eﬀect of (ii), various empirical correction factors have subsequently been added to the extended DH expression. log gi ¼ 0:5 z2i

p p I = 1 þ 0:33 ai I bI Truesdell Jones equation ð3:6Þ

log gi ¼ 0:5 z2i

p p I = 1 þ I bI

Davies equation

ð3:7Þ

The Truesdell–Jones18 equation contains an empirical term, bI, as well as the ion size parameter, ai. Values of both ai and b are often obtained by ﬁtting this equation to the individual activity coeﬃcients obtained from mean salt data using the MacInnes convention (see Example 3.1).9,10 This equation is used in some geochemical computer models, e.g. PHREEQE.19 The Davies equation, which also incorporates an empirical term bI but uses an average ion size parameter (ai ¼ 3), is more commonly used because the ion size parameter is not well deﬁned for all ions. The Davies equation can be used to calculate activity coeﬃcients for monovalent and divalent ions in aqueous solutions, where I ¼ 0.1–0.7 mol L1 and I ¼ 0.1–B0.3 mol L1, respectively. Originally, Davies used a value of 0.2 for b but later suggested that the value be increased to 0.3.20,21 The geochemical computer model MINTEQA2 employs the Davies equation to calculate single-ion activity coeﬃcients but uses a value of 0.24 for b.22 Speciﬁc Ion Interaction Theory (SIT) is a more advanced empirical approach that was developed by Guggenheim23 and extended by Scatchard and others.24–26 The SIT equation has an extended DH term that accounts for the non-speciﬁc long-range interactions but also has an ion- and electrolyte-speciﬁc correction factor, eij [j], for each pair of ions, i and j, of opposite charge. Brønsted27 also recognized the importance of short-range interactions in higher ionic strength solutions and Equation (3.8) is often referred to as the Brønsted–Guggenheim equation. log gi ¼ ‘Extended DH’ þ Sjeij [j]

(3.8)

Chapter 3

Originally the extended DH term was similar to that in Equation (3.5), i.e. with an average ion size parameter, ai ¼ 3. Scatchard28 showed that a better ﬁt with experimental data was obtained using an average value of ai ¼ 4.6, i.e. 0.33ai ¼ 1.5, and so the extended DH expression used in the SIT model is often p p ð3:9Þ log gi ¼ 0:510 z2i I = 1 þ 1:5 I Overall, however, the SIT approach is limited in application to intermediate ionic strength solutions (0.1–3.5 mol L1). The Pitzer model can be used to obtain activity coeﬃcients for solutes in low (o0.1 mol L1), intermediate (0.1–3.5 mol L1) and high (43.5 mol L1) ionic strength solutions. The Pitzer equations29 include terms for binary and ternary interactions between solute species as well as a modiﬁed DH expression. The general formula is X X D ½j þ E ½j½k þ . . . ð3:10Þ ln gi ¼ z2i f g þ ij j jk ijk where fg is the modiﬁed DH term fg ¼ 0.392 (OI/(1 þ 1.2OI) þ (2/1.2) ln(1 þ 1.2 OI))

(3.11)

and the Dij terms describe interactions between pairs of ions, i and j, while the Eijk terms describe interactions among three ions, i, j, and k. Higher-order terms can also be added to this general formula. A key diﬀerence between the SIT and Pitzer models is that interactions between ions of like charge are also included. For ionic strength solutions of up to B3.5 mol L1, there is good agreement between the two models but, at 43.5 mol L1, the ternary and higher-order terms in the Pitzer equations become more important and this model gives the more accurate ion activity coeﬃcients.12 Although naturally occurring brines and some high ionic strength contaminated waters may require the more complicated expressions developed in the Davies, SIT, or Pitzer models, the use of Equations (3.3)–(3.5) is justiﬁed for the ionic strengths of many freshwaters. Finally, this section has focused primarily on ions in aqueous solution but it should be remembered that aqueous solutions also contain important dissolved neutral species, e.g. O2, CO2, H2CO3, Si(OH)4. For non-ideal solutions, the Setschenow Equation (1899) has traditionally been used to describe the ionic strength dependence of the activity coeﬃcients for dissolved neutral species and can be written as log gi ¼ ki [i]

(3.12)

Chemistry of Freshwaters

where ki is the temperature dependent salting out coeﬃcient for molecular species, i. At higher ionic strength values, an additional dependence on [i]2 is often required to ﬁt the observed solubility data.17 Alternatively, the Pitzer model can be used with a high degree of accuracy to describe the shortrange binary (neutral–neutral, neutral–cation, neutral–anion) and ternary (neutral–neutral–neutral, neutral–cation–anion) interactions between ions and neutral species in single and mixed electrolyte solutions.17 3.2.1.3 Equilibria and Equilibrium Constants. Many of the important reactions involving solutes in freshwaters can be described by equilibria. This approach means that an equilibrium constant, K, relating the activities of the solutes, can be deﬁned for each stoichiometric equilibrium expression. aAþbB"cCþdD

(3.13)

K ¼ {C}c {D}d / {A}a {B}b

(3.14)

The equilibrium constant can also be deﬁned in terms of the concentrations of the solutes. K ¼ (gc [C])c (gd [D])d/(ga [A])a (gb [B])b

(3.15)

For inﬁnitely dilute solutions, { } ¼ [ ] and the equilibrium constant can be written as K ¼ [C]c [D]d/[A]a [B]b

(3.16)

For solutions of ﬁxed ionic strength, or, for example, where major ions in solution, e.g. conservative cations and anions, are present at concentrations several orders of magnitude greater than the species involved in the chemical equilibrium, e.g. A, B, C, and D in Equation (3.13), it can be assumed that the solute activity coeﬃcients are also constants and can be incorporated into the equilibrium constant. The equilibrium constant for a ﬁxed ionic strength aqueous solution is termed a constant ionic strength equilibrium constant, cK. K ¼ [C]c [D]d/[A]a [B]b where cK ¼ (gaa gbb / gcc gdd) K

(3.17)

This again enables the use of concentrations rather than activities in equilibrium calculations. It is sometimes advantageous to use a mixture of activities and concentrations and a mixed equilibrium constant, K 0 , is deﬁned as K 0 ¼ {C}c [D]d / [A]a [B]b

where K 0 ¼ (gaa gbb / gdd) K

(3.18)

Chapter 3

For example, a mixed acidity constant is frequently used where pH has been measured according to the IUPAC convention as the activity of hydrogen ions but the concentrations of the conjugate acid-base pair are used. The relationship between K and cK or K and K 0 as deﬁned above can be used to calculate the eﬀect of ionic strength of solution on the true equilibrium constant, K. cK and K 0 can be calculated using the Equations (3.3)–(3.5) together with experimentally determined values of K (Example 3.3).

Example 3.3: Calculation of the mixed acidity constant, K 0 , using the Gu¨ntelberg approximation and the Davies equation. For an acid-base equilibrium: HA " H1 þ A

K ¼ {H1}{A}/{HA}

K 0 ¼ ðgHA =gA Þ K Taking logs: logK 0 ¼ log ðgHA =gA Þ þ log K pK 0 ¼ pK þ log gA log gHA Using the Gu¨ntelberg approximation: p p I= 1 þ I pK 0 ¼ pK þ 0:5 z2HA z2A Using the Davies equation: p p I = 1 þ I 0:3I pK 0 ¼ pK þ 0:5 z2HA z2A Now calculate the mixed acidity constant, K 0 , at ionic strength values of (i) 0.05 mol L1 using the Gu¨ntelberg approximation and (ii) 0.4 mol L1 using the Davies equation for a monoprotic acid, HA, with K ¼ 6.8 108: (i)

p p pK 0 ¼ 7:17 þ 0:5 ð0 ð 1Þ2 Þ ð 0:05=ð1 þ 0:05ÞÞ ¼ 7:17 0:0914 ¼ 7:08 K 0 ¼ 8:34 108

Chemistry of Freshwaters

(ii)

p p pK 0 ¼ 7:17 þ 0:5ð0 ð 1Þ2 Þ ð 0:4=ð1 þ 0:4Þ 0:3 0:4Þ ¼ 7:17 0:5 ð0:387 0:12Þ ¼ 7:04 K 0 ¼ 9:19 108

The values for the corrected constants, K 0 , are 22% and 35% greater, respectively, than the inﬁnite dilution constant, K. At higher ionic strengths, the activity coeﬃcients could be calculated using the SIT or Pitzer equations and thus a more accurate correction made to the inﬁnite dilution constant, K. Finally, it should be noted that all values of the equilibrium constants apply only at a speciﬁed temperature. In Sections 3.2.2–3.2.4, the values of the equilibrium constants apply at 298 K unless stated otherwise. 3.2.2 Dissolution/Precipitation Reactions Section 3.2.1 has provided relationships between solute activities and concentrations in solution in order that solute behaviour can be quantiﬁed. This section discusses dissolution and precipitation reactions that impart or remove solutes to/from natural waters, and therefore modify the chemical composition of natural waters. 3.2.2.1 Physical and Chemical Weathering Processes. Initial steps involving physical weathering by thermal expansion and contraction or abrasion lead to the disintegration of rock. Disintegration increases the surface area of the rock, which, in the presence of water, can undergo chemical weathering. Water acts not only as a reactant but also as a transporting agent of dissolved and particulate components, and so weathering processes are extremely important in the hydrogeochemical cycling of elements. Mineral dissolution reactions often involve hydrogen ions from mineral or organic acids, e.g. acid hydrolysis of Nafeldspar in Equation (3.21). Alternatively, the transfer of electrons (sometimes simultaneously with proton transfer) may promote the dissolution of minerals, e.g. reductive dissolution of biotite. Chemical weathering of minerals results not only in the introduction of solutes to the aqueous phase but often in the formation of new solid phases. Dissolution is described as congruent, where aqueous phase solutes are the only products, or incongruent, where new solid phase(s) in addition to aqueous phase solutes are the products. These reactions

Chapter 3

can be represented by equilibria, where the equilibrium constant is related only to the activities of the aqueous solutes, i.e. an assumption is made that the activities of solid phases have the value of unity. Examples of weathering of primary minerals, those formed at the same time as the parent rock, are shown in (3.19)–(3.22): Congruent dissolution of quartz SiO2ðsÞ þ 2 H2 O Ð SiðOHÞ04ðaqÞ quartz silicic acid

ð3:19Þ

K ¼ {Si(OH)40}

(3.20)

At pHo9, the equilibrium between quartz and silicic acid can be represented as in (3.19) and the value of K at 298 K is 1.05 104 mol L1,30 indicating that the solubility of quartz is low. At higher pH values, the dissociation of silicic acid results in the increased solubility of quartz (see Section 3.2.4.1). Incongruent dissolution of Na-feldspar 2 NaAlSi3 O8ðsÞ þ2 Hþ þ 9 H2 O Ð Al2 Si2 O5 ðOHÞ4ðsÞ þ2 Naþ ðaqÞ Na-feldspar kaolinite þ 4 SiðOHÞ04ðaqÞ

ð3:21Þ

silicic acid

K ¼ fSiðOHÞ04 g4 fNaþ g2 =fHþ g2

ð3:22Þ

Alteration of primary minerals such as the feldspars gives rise to secondary minerals such as kaolinite (Al2Si2O5(OH)4), smectites (e.g. Na0.33Al2.33Si3.67O10(OH)2), and gibbsite (Al(OH)3). As the composition of the new mineral phases and the potential for further alteration of secondary minerals are a function of the prevailing geochemical conditions, many equilibrium expressions must be used to fully describe chemical weathering processes in natural systems. For example, reactions leading to the formation of kaolinite from the primary mineral, Na-feldspar, as well as alteration of the secondary mineral, e.g. to gibbsite, must be considered. The composition of the aqueous phase is of major importance in determining the nature of the solid products formed during chemical weathering. In particular, the solution activities of silicic acid, metal ions, and hydrogen ions are key parameters inﬂuencing the formation and alteration processes of new solid phases (Example 3.4).8

Chemistry of Freshwaters

Example 3.4: Investigation of the stability relationship between gibbsite and kaolinite in natural waters at pHo7. Consider the acid hydrolysis weathering reaction undergone by Nafeldspars (pHo7): þ 3þ NaAlSi3O8(s) þ Hþ (aq) þ 7H2O " Na(aq) þ Al(aq) þ 3OH(aq) þ 3Si(OH)4(aq) This equilibrium expression represents congruent dissolution of Nafeldspar but the solution quickly becomes saturated with respect to the solid phase Al(OH)3 (gibbsite KSP ¼ 1033.9) and so the secondary mineral, gibbsite, is formed. þ NaAlSi3O8(s) þ Hþ (aq) þ 7H2O " Al(OH)3(s) þ Na(aq) þ 3Si(OH)4(aq) The concentration of silicic acid in solution is a key parameter inﬂuencing the stability of gibbsite. In the presence of even low concentrations of silicic acid, e.g. 104 mol L1, kaolinite would be formed instead. þ 2NaAlSi3O8(s) þ 2Hþ (aq) þ 9H2O " Al2Si2O5 (OH)4(s) þ 2Na(aq) þ 4Si(OH)4(aq) In natural systems, the role of water as a transporting agent is important. Geochemical conditions promoting removal of silicic acid from the solid-water interface (e.g. water ﬂow) favour the formation of gibbsite over kaolinite. (Adapted from Drever.8) 3.2.2.2 Solubility. Dissolution of minerals during chemical weathering releases species into solution but aqueous phase concentrations are limited by the solubility of solid phases, such as amorphous silica, gibbsite, and metal salts. For example, the solubility of AlIII and CaII may be limited by the formation of gibbsite and gypsum, respectively. Al(OH)3

gibbsite

" Al3þ (aq) þ 3OH(aq)

2 CaSO4.2H2Ogypsum " Ca2þ (aq) þ SO4(aq) þ 2H2O

(3.23) (3.24)

The equilibrium solubilities are deﬁned by the respective solubility products, KSP (values from ref. 30). KSP ¼ {Al31} {OH}3 ¼ 1033.9

(3.25)

Chapter 3 4.58 KSP ¼ {Ca21} {SO2 4 } ¼ 10

(3.26)

A solution is considered to be undersaturated, saturated, or oversaturated with respect to a solid phase, for example gypsum, if KSP 4 2 21 21 {Ca21}{SO2 4 }observed, KSP ¼ {Ca }{SO4 }observed, or KSPo{Ca } 2 {SO4 }observed, respectively. 3.2.2.3 Inﬂuence of Organic Matter. Natural organic matter is present in most natural systems. With respect to weathering processes, its importance in freshwaters and indeed soils and sediments in contact with freshwaters can be attributed to the presence of organic acids, which include low molecular weight compounds, e.g. oxalic acid, and extremely complex, high molecular weight coloured compounds described as humic substances – a heterogeneous mixture of polyfunctional macromolecules ranging in size from a few thousand to several hundred thousand Daltons (Da). In addition to inorganic acids, e.g. H2CO3, these provide hydrogen ions for the acid hydrolysis of minerals. As well as promoting dissolution, natural organic matter can inﬂuence the formation of new mineral phases.31 3.2.3 Complexation Reactions in Freshwaters In this section, as a starting point, it is assumed that all species in solution are in hydrated form, e.g. MIII(H2O)3þ 6 , where the six water molecules form the ﬁrst co-ordination sphere of the metal ion, MIII. The hydrated form is often represented as, for example, M3þ (aq). 3.2.3.1 Outer and Inner Sphere Complexes. Outer sphere complexation involves interactions between metal ions and other solute species in which the co-ordinated water of the metal ion and/or the other solute species are retained. For example, the initial step in the formation of ion pairs, where ions of opposite charge approach within a critical distance and are then held together by coulombic attractive forces, is described as outer sphere complex formation. 2 21 2 Mg2þ (aq) þ SO4(aq) " (Mg (H2O)(H2O)SO4(aq))

(3.27)

The formation of ion pairs is inﬂuenced by the nature of the oppositely charged ions, the ionic strength of solution, and ion charge. Ion pairs are generally formed between hard (low polarizability) metal cations, e.g. 2 Mg21, Fe31, and hard anions, e.g. CO2 3 , SO4 , and ion pair formation is generally most signiﬁcant in high ionic strength aqueous phases. In

Chemistry of Freshwaters

freshwaters, which are frequently of low ionic strength, ion charge is a key parameter. Greater coulombic interactions occur between oppositely charged ions with high charge. Inner sphere complexation involves interactions between metal ions and other species in solution which possess lone pairs of electrons. Inner sphere complexation involves the transfer of at least one lone pair of electrons. Those species which possess electron lone pairs are termed ligands and reactions may involve inorganic or organic ligands. 3.2.3.2 Hydrolysis. In aqueous systems, hydrolysis reactions are an important example of inner sphere complexation for many metal ions. The interaction between hydrated metal ions and water can be written as a series of equilibria for which the equilibrium constants are denoted *Kn.y 2þ þ Al3þ (aq) þ H2O(1) " Al(OH)(aq) þ H(aq)

*K1 ¼ 104.95 (3.28)

þ þ Al(OH)2þ (aq) þH2O(1) " Al(OH)2(aq) þ H(aq)

*K2 ¼ 105.6 (3.29)

þ Al(OH)þ 2(aq) þH2O(1) " Al(OH)3(aq) þ H(aq)

*K3 ¼ 106.7 (3.30)

þ Al(OH)3(aq) þH2O(1) " Al(OH) 4(aq) þ H(aq)

*K4 ¼ 105.6 (3.31)

Hydrolysis and, more generally, complex formation equilibria may be described by cumulative stability constants, bn. Using the Equations (3.28)–(3.31) 2þ þ 4.95 Al3þ (3.32) (aq) þH2O(1) " Al(OH)(aq) þ H(aq) *b1 ¼ *K1 ¼ 10

The hydrolysis of hydrated metal ions can also be written as the interaction between the hydrated metal ion and the hydroxyl ligand. For example Al31(aq) þ OH(aq) " Al(OH)21(aq) K1

where K1 and *K1 are related as follows Al31(aq) þ OH(aq) " Al(OH)21(aq) K1 H2O(l) " H1(aq) þ OH(aq) Kw so Al31(aq) þ H2O(l) " Al(OH)21(aq) þ H1(aq) *K ¼ K1Kw The values for *Kn are those obtained by Wesolowski and Palmer;32 all values are given in unitless form but fundamentally are deﬁned in terms of the products of the activities of ions on the molar or molal scales.

Chapter 3 þ þ Al3þ (aq) þ2H2O(1) " Al(OH)2(aq) þ 2H(aq)

*b2 ¼ *K1 *K2 ¼ 1010.55

(3.33)

þ Al3þ (aq) þ3H2O(1) " Al(OH)3(aq) þ 3H(aq) *b3 ¼ *K1 *K2 *K3 ¼ 1017.25 (3.34) þ Al3þ (aq) þ4H2O(1) " Al(OH)4(aq) þ 4H(aq) *b4 ¼ *K1 *K2 *K3 *K4 ¼ 1022.85 (3.35)

Clearly, these hydrolysis reactions are dependent on hydrogen ion activity. The relationship between speciation and pH, which is inﬂuenced by the characteristics of the metal ion, will be discussed further in Section 3.2.4.1. 3.2.3.3 Inorganic Complexes. The main inorganic ligands in oxygenated freshwaters, in addition to OH, are HCO3, CO32, Cl, SO42, and F, and, under anoxic conditions, HS and S2 (see Section 3.2.4.4). The stability of complexes formed between metal ions and inorganic ligands depends on the nature of the metal ion as well as the properties of the ligand. Table 2 shows some of the major complexed metal species involving inorganic ligands in oxygenated freshwaters at pH 8. 3.2.3.4 Surface Complex Formation. Metal ions form both outer and inner sphere complexes with solid surfaces, e.g. hydrous oxides of iron, manganese, and aluminium. In addition, metal ions, attracted to charged surfaces, may be held in a diﬀuse layer, which, depending upon ionic strength, extends several nanometres from the surface into solution. Table 2

Major species in freshwaters

Metal ion

Major species

Mg CaII AlIII MnIV FeIII NiII CuII ZnII PbII

Mg Ca21 Al(OH)30, Al(OH)21, Al(OH)4 MnO20 Fe(OH)30, Fe(OH)21, Fe(OH)4 Ni21, NiCO30 CuCO30, Cu(OH)20 Zn21, ZnCO30 PbCO30

[Mx1(aq)]/[MTOT

(aq)]

0.94 0.94 1 109 – 2 1011 0.4 0.01 0.4 0.05

Source: Adapted from Sigg and Xue,33 and reproduced with permission from W. Stumm and J.J. Morgan, Aquatic Chemistry 3rd Edn, r Wiley, 1996.

Chemistry of Freshwaters

Diﬀuse layer metal retention and outer sphere complex formation involve electrostatic attractive forces, which are characteristically weaker than co-ordinative interactions leading to inner sphere surface complex formation. A number of factors inﬂuence metal interactions with surfaces, including the chemical composition of the surface, surface charge, and the nature and speciation of the metal ion. The importance of the pH of the aqueous phase in these interactions will be discussed further in Section 3.2.4.1. 3.2.3.5 Organic Complexes. Dissolved organic matter consists of a highly heterogeneous mixture of compounds, including low molecular weight acids and sugars as well as the high molecular weight coloured compounds termed humic substances. Humic substances are secondary synthetic products derived from decaying organic debris. Although they are structurally poorly deﬁned, it is accepted that large numbers of mainly oxygen-containing functional groups are attached to a ﬂexible, predominantly carbon backbone. Individual organic molecules, in particular those of humic substances, can provide more than one functional group for complex formation with a hydrated metal ion. Ligands which provide two and three functional groups are termed bidentate and tridentate, respectively. Complexes where more than one functional group from the same organic molecule is involved are more stable than those where functional groups are from discrete organic molecules. The concentration of dissolved organic matter in freshwaters is generally low (2–6 mg C L1) and humic substances, comprising molecules which possess large numbers of functional groups in numerous diﬀerent chemical environments, are implicated as the component of natural organic matter most important in metal binding.31 3.2.4 Species Distribution in Freshwaters 3.2.4.1 pH as a Master Variable. pH is one of the key parameters which inﬂuences species distribution in aqueous systems. Many equilibrium expressions contain a hydrogen ion activity term, e.g. acidity constants, complexation constants. It is therefore useful to consider the relationship between the activity of the species of interest (e.g. contaminant metal ions, organic pollutants, naturally occurring inorganic and organic solutes, and weakly acidic functional groups on mineral surfaces) and the activity of hydrogen ions. For example, an acid – conjugate base pair can be represented as HA and A, respectively. HA " A þ H1

(3.36)

Chapter 3

K ¼ {A}{H1}/{HA}

(3.37)

The activity of acid initially in solution, C, termed the analytical activity, is equal to the sum of the equilibrium activities of the acid and its conjugate base. (3.38) C ¼ {HA} þ {A} Equations (3.37) and (3.38) can be combined to give an expression for the activity of the conjugate base in terms of the equilibrium constant, the analytical activity, which is also a constant, and the hydrogen ion activity. K ¼ {A}{H1}/(C{A})

(3.39)

{A} ¼ KC/(K þ {H1})

(3.40)

Similarly an expression for the activity of the acid can be written as {HA} ¼ {H1}C/(K þ {H1})

(3.41)

Alternatively, the activities of acid and conjugate base can be represented as a fraction of the analytical activity. a0 ¼ {HA}/C ¼ {H1}/(K þ {H1})

(3.42)

a1 ¼ {A}/C ¼ K/(K þ {H1})

(3.43)

a0 þa1 ¼ 1

(3.44)

a0 and a1 are termed dissociation fractions for the acid and conjugate base, respectively. The expressions for the dissociation fractions are now independent of the analytical activity. A plot of dissociation fraction against pH for an acid-conjugate base pair is shown in Figure 1. More frequently, log{ } against pH is plotted using (3.45)–(3.46). log {HA} ¼ log{H1} þ log C log(K þ {H1})

(3.45)

log {A} ¼ log K þ log C log(K þ {H1})

(3.46)

The equilibrium pH can be obtained directly from the log{ } vs. pH plot by adding lines showing log {H1} and log {OH} and by using an expression for charge balance. {H1} ¼ {A} þ {OH}

(3.47)

Chemistry of Freshwaters 1.0

α0

pH=pK

α1

0.8

0.6

0.4

0.2

0.0

Figure 1 The inﬂuence of pH on a0 and a1 for a monoprotic acid

For a solution containing a weak acid, the charge balance relationship (3.47) is satisﬁed at the equilibrium pH (Example 3.5). The same log{ } vs. pH plot can also be used to determine the equilibrium pH of a solution containing the salt of a weak acid, e.g. NaA, by using the appropriate charge balance expression. {Na1} þ {H1} ¼ {A} þ {OH} 1

(3.48)

Using C ¼ {Na } ¼ {HA} þ {A } gives {HA} þ {H1} ¼ {OH}

(3.49)

Example 3.5: Graphical illustration of the relationship between (i) log{2,4,6-TCP} and pH; (ii) log{2,4,6-TCP} and pH for an organic pollutant, 2,4,6-trichlorophenol (2,4,6-TCP), that has been released into a natural water (assume I ¼ 0 mol L1). Use C ¼ {2,4,6-TCPTOTAL} ¼ 4 104 mol L1 and K ¼ 106.13. Graphical representation is best approached by dividing the pH range into two regions, pHopK and pH4pK. The slope of the lines in each of these two regions can be determined by diﬀerentiating (3.45) and (3.46) with respect to pH. For pHopK, (Kþ{H1})-{H1} and d log {2, 4, 6-TCP}/d pH ¼ d log{H1}/d pH þ d log C/d pH d log{H1}/d pH ¼ 0

100

Chapter 3

d log {2, 4, 6-TCP}/d pH ¼ d log K/d pH þ d log C/d pH d log{H1}/d pH ¼ þ1 For pH4pK, (Kþ{H1})-K and d log {2, 4, 6-TCP}/d pH ¼ d log {H1}/d pH þ d log C/d pH d log K/d pH ¼ 1 d log {2, 4, 6-TCP}/d pH ¼ d log K/d pH þ d log C/d pH d log K/d pH ¼ 0 To construct the straight line sections of the graph, note that log{2,4,6-TCP} ¼ log C for pHopK and log{2,4,6-TCP} ¼ log C for pH4pK. Then using the full expressions for log{2,4,6-TCP} and log{2,4,6-TCP}, calculate several points on lines of slope þ1 and 1 to complete the graph. 3

-log { }

OH-

H+

TCP

TCP-

6 {H +} ~ {TCP -}

The equilibrium pH for a solution containing the weak acid is obtained by approximating the charge balance expression by {H1}B{TCP} (because the solution will be acidic and thus {OH} is very small). Conversely, the equilibrium pH for a solution of the salt of the weak acid is obtained by approximating the appropriate charge balance expression by {TCP}B{OH} (because the solution will be basic and so {H1} is very small). Note also that, as shown in Example 3.4, the equilibrium constant, K, can be corrected for ionic strength of solution. In general, the equilibrium pH value decreases with increasing ionic strength. Carbonate Equilibria. Another example of the importance of acid-base behaviour, not only in aerated freshwaters but also in seawater (see

101

Chemistry of Freshwaters

Chapter 4), is the dissociation of diprotic acids such as carbonic acid, H2CO3. The hydration of dissolved CO2(aq) in natural waters gives rise to carbonic acid. The dissociation of carbonic acid not only inﬂuences the pH of the water but also provides ligands which can complex trace metals. Example 3.6 illustrates the relationship between log { } and pH for a closed aqueous system which contains dissolved CO2 (the presence of mineral phases has been ignored). By assuming that a system is closed to the atmosphere, it is possible to treat carbonic acid as a non-volatile acid and to consider only the hydration reaction which converts CO2(aq) to H2CO3. H2CO3* is used to represent the sum of the activities of CO2(aq) and H2CO3, i.e. to take account of the presence of CO2(aq), which is in equilibrium with H2CO3. The dissociation of carbonic acid can then be described by (3.50)–(3.51) and the equilibrium constants, K1 and K2, are deﬁned by (3.52)–(3.53). 1 H2CO*3 " HCO 3 þ H

(3.50)

2 1 HCO 3 " CO3 þ H

(3.51)

1 * K1 ¼ {HCO 3 } {H }/{H2CO3}

(3.52)

1 K2 ¼ {CO2 3 } {H }/{HCO3 }

(3.53)

The analytical activity of dissolved inorganic carbon species can be written as the sum of undissociated and dissociated acid species (3.54). For a closed system, the value of C is constant over the entire pH range. 2 C ¼ {H2CO*3} þ {HCO 3 } þ {CO3 }

(3.54)

Combining (3.52)–(3.54), an expression for each of these (H2CO3*, HCO3, and CO32) is obtained. {H2CO*3} ¼ C ({H1}2/({H1}2 þ K1{H1} þ K1K2))

(3.55)

1 1 2 1 {HCO 3 } ¼ C (K1{H }/({H } þ K1{H } þ K1K2))

(3.56)

1 2 1 {CO2 3 } ¼ C (K1K2/({H } þ K1{H } þ K1K2))

(3.57)

The dissociation fractions, a0, a1, and a2, are obtained by dividing each of (3.55)–(3.57) by the analytical activity, C (cf. (3.42)–(3.43)). a0 ¼ {H1}2/({H1}2 þ K1{H1} þ K1K2)

(3.58)

102

Chapter 3

a1 ¼ K1{H1}/({H1}2 þ K1{H1} þ K1K2)

(3.59)

a2 ¼ K1K2/({H1}2 þ K1{H1} þ K1K2)

(3.60)

a0 þ a1 þ a2 ¼ 1

(3.61)

where

Graphical representation of log { } vs. pH can again be used to obtain the equilibrium pH (Example 3.6).

Example 3.6: Graphical illustration of the relationship between (i) log{H2CO3*}, (ii) log{H2CO3}, (iii) log{HCO3}, and (iv) log{CO32} and pH for a closed aqueous system. Use C ¼ 2 105 mol L1, K1 ¼ 5.1 107 and K2 ¼ 5.1 1011. The presence of dissolved CO2 is taken into account by the following equilibrium: CO2ðaqÞ þ H2 O ÐH2 CO3ðaqÞ Khydration ¼ fH2 CO3 g=fCO2 g ¼ 1:54 103 We deﬁned {H2CO3*} ¼ {CO2} þ {H2CO3} but {H2CO3*}B{CO2} because the hydration equilibrium lies far to the left. So Khydration ¼ {H2CO3}/{H2CO3*} and {H2CO3} ¼ {H2CO3*} 1.54 103. In this way the presence of dissolved CO2 is taken into account while the true concentration of carbonic acid can still be determined. By assuming a closed system, the total analytical activity of carbonic acid, C, is constant [this is not true for an open system (Example 3.7)]. Thus (3.52)–(3.54) can be combined to give (3.55)–(3.57). For a diprotic acid such as H2CO3*, the pH range can be divided into three regions, pHopK1, pK1opHopK2, and pH4pK2, and the procedure outlined in Example 3.5, i.e. obtaining logarithmic expressions for {H2CO3*}, {HCO3}, and {CO32} and diﬀerentiating each with respect to pH, can then be utilized to construct the plot of log{ } against pH. For {H2CO3*} pHopK1 : logfH2 CO3 g ¼ 2 logfHþ g þ log C 2 logfHþ g ¼ log C d logfH2 CO3 g=d pH

¼0

103

Chemistry of Freshwaters

pK1 opHopK2 :

logfH2 CO3 g ¼ logfHþ g þ log C log K1 d logfH2 CO3 g=d pH ¼ 1

pH4pK2 :

logfH2 CO3 g ¼ 2logfHþ g þ log C log K1 K2

d logfH2 CO3 g=d pH ¼ 2 The slopes for the other lines are obtained in the same way. Approximating the charge balance expression (3.62) by {H1}B{HCO3} gives the equilibrium pH for a closed aqueous system containing carbonic acid. H

-log{ }

HC O

H2CO3*

O H-

{H +} ~ {HC O3- }

H2CO3

Alkalinity. An important deﬁnition arising from the considerations of carbonate equilibria is that of alkalinity, which is derived with reference to the charge balance expression for carbonic acid. 2 (3.62) {H1} ¼ {HCO 3 } þ 2{CO3 } þ {OH } Alkalinity is described as the acid neutralizing capacity of an aqueous system or equivalently as the amount of base possessed by the system. 2 1 Alkalinity ¼ {HCO 3 } þ 2{CO3 } þ {OH } {H }

(3.63)

Species other than carbonate can contribute to alkalinity and an alternative deﬁnition is 2 2 Alkalinity ¼ fHCO 3 g þ 2fCO3 g þ fNH3 g þ fHS g þ 2fS g 2 2 þ fH3 SiO 4 g þ 2fH2 SiO4 g þ fOrg g þ fHPO4 g þ þ 2fPO3 4 g þ fOH g fH g fH3 PO4 g

ð3:64Þ

104

Chapter 3

In addition to the charge balance expression for H2CO3, Equation (3.64) includes appropriate expressions for NH41, H2S, H4SiO4, HOrg, and H2PO4. In each case, the reference point has been selected on the basis of the pK values for these species. For example, phosphoric acid, H3PO4, has pK1 ¼ 2.15 and pK2 ¼ 7.20 and so H2PO4 will be the main ‘P’ species in solution at the reference point, i.e. the equilibrium pH of a solution of carbonic acid and water (pH B5.65). The contribution of ‘P’ species to alkalinity is thus determined by using the charge balance expression for NaH2PO4. Another approach to determine carbonate alkalinity is by using a charge balance expression for the major conservative ions in aqueous solution. g þ 2fCO2 Carbonate Alkalinity ¼ fHCO 3 g X 3 ¼ ðconservative cationsÞ X ðconservative anionsÞ ¼ fNaþ g þ fKþ g þ 2fCa2þ g þ 2fMg2þ g fCl g 2fSO2 4 g fNO3 g

ð3:65Þ

Alkalinity is an important parameter in assessing the eﬀects of environmental change on aqueous systems (see Section 3.3.4.1). It is also important to understand that, by deﬁnition, alkalinity (Equation (3.61)) is independent of addition or removal of CO2 (or H2CO3) from the system (cf. Equation (3.62) – H2CO3 does not appear in the charge balance expression). This can be very useful in the determination of the concentration of dissolved inorganic carbon species in aqueous systems that are in equilibrium with an atmosphere containing CO2(g) (Example 3.7). Example 3.7: Graphical illustration of the relationship between (i) log{H2CO3*}, (ii) log{H2CO3}, (iii) log{HCO3}, and (iv) log{CO32} and pH for an open aqueous system. Use pCO2 ¼ 3.5 104 atm, KH ¼ 3.2 102, K1 ¼ 5.1 107 and K2 ¼ 5.1 1011. CO2(g) þ H2O " H2CO3*

KH ¼ {H2CO3*}/pCO2 ¼ 3.2 102 mol L1 atm1

where KH is the Henry’s Law constant for atmosphere-aqueous phase equilibria involving gases.

105

Chemistry of Freshwaters

Clearly, {H2CO3*} is now a constant for a ﬁxed partial pressure of CO2 and is independent of pH, i.e. {H2CO3*} ¼ KHpCO2. This means that {H2CO3} is also constant over the entire pH range. Combining this new expression for {H2CO3*} with (3.52)–(3.53) gives expressions for {HCO3} and {CO32}: 1 2 1 2 {HCO 3 } ¼ K1KHpCO2/{H } and {CO3 } ¼ K1K2KHpCO2/{H } The analytical activity, C, is not a constant but can be expressed as C ¼ KHpCO2/a0 which increases with increasing pH according to the function (1 þ K1/ {H1} þ K1K2/{H1}2) as shown by the heavy black line on the accompanying graph. Plotting log{ } vs. pH for {HCO3} and {CO32} gives lines of slope þ1 and þ2, respectively. The charge balance expression, which is the same as that for the closed system, can again be used to determine the equilibrium pH. {H +} ~ {HC O3- }

OH -

2 3 -

-log { }

4 H2CO3*

HC O

6 H2CO3

Finally, it is useful to note that the expressions describing the pH dependence of carbonate species concentration for both closed (Example 3.6) and open systems can also be used in mass balance expressions for metal complexation. pH Dependence of Complex Formation. Other equilibria such as metal complexation reactions can be considered as acid–base reactions and plots of log { } against pH also provide information about the dominant species present in solution under diﬀerent geochemical conditions. Hydrolysis of metal cations occurs progressively with increasing pH, e.g.

106

Chapter 3

M31(aq), M(OH)21(aq), M(OH)21(aq), M(OH)30(aq), and M(OH)4(aq) are the hydrolysis products for many MIII cations (see Section 3.2.3.2). M31(aq) is generally found at low pH and only at extremely high pH is M(OH)4 formed. The dependence of CrIII hydrolysis reactions on pH is illustrated in Example 3.8.

Example 3.8: Determination of CrIII speciation in an aqueous solution at pH 1, 6, and 11. Use {CrIIIT} ¼ 106 mol L1 and log*b1 ¼ 4.00, log*b2 ¼ 9.62, log*b3 ¼ 16.75, and log*b4 ¼ 27.77. Assume that no Cr(OH)3 is precipitated, i.e. hom*ogeneous solution, and that inorganic and organic ligands are absent. Cr3þ þ H2 O Ð CrOH2þ þ Hþ 3þ

þ 2H2 O Ð

Cr3þ þ 3H2 O Ð 3þ

þ 4H2 O Ð

þ CrðOHÞþ 2 þ 2H CrðOHÞ03 þ 3Hþ þ CrðOHÞ 4 þ 4H

þ 0 3þ 2þ fCrIII T g ¼ fCr g þ fCrOH g þ fCrðOHÞ2 g þ fCrðOHÞ3 g 6 mol L1 þ fCrðOHÞ 4 g ¼ 10

¼ fCr3þ g ð1 þ b1 =fHþ g þ b2 =fH þ g2 þ b3 =fHþ g3 þ b4 =fHþ g4 Þ This equation can be solved for {Cr31} at various pH values by using the equilibrium constants for each of the four equilibria above, e.g. at pH 2, {Cr31} ¼ 9.90 107 mol L1. The concentrations of all other species can be calculated using {Cr31} at each selected pH value, e.g. at pH 2, {CrOH21} ¼ (*b1/{H1}) {Cr31} ¼ 0.01 9.90 107 ¼ 9.90 109 mol L1. Plotting log{ } against pH illustrates that the dominant species of CrIII are Cr31, Cr(OH)21, and Cr(OH)4 at pH 1, 6, and 11, respectively. Alternatively, the same information may be presented in a plot of mole fraction, w, against pH. The mole fractions can be obtained by dividing species activity by the total activity of CrIII, e.g. wCr31 ¼ {Cr31}/{CrIIIT}.

107

Chemistry of Freshwaters 4 Cr(OH) 30

CrOH 2+ Cr(OH) 2+

Cr 3+

Cr(OH) 4-

-log ( }

8 10 12 14 16

1.0

0.8

Cr 3+

CrOH 2+ Cr(OH) 2+

Cr(OH) 30

Cr(OH) 4-

0.6

0.4

0.2

0.0

It should be noted that, for higher values of {CrIIIT}, equilibria for polymeric species should be included and that Cr(OH)3 may precipitate. In the presence of iron, the mixed precipitate FexCr1x(OH)3 may be formed.

Example 3.9 shows, however, that, for AlIII, ligands such as F and SO42 may compete successfully with the hydroxyl ion at low pH values (see Section 3.3.1.2). For aqueous systems in contact with solid phases,

108

Chapter 3

the formation of such complexes increases mineral solubility by several orders of magnitude over that predicted by the solubility product, KSP. The presence of even trace concentrations of organic ligands can have an even greater eﬀect over a wide pH range on AlIII speciation (Example 3.9).12 More generally, complexation clearly has an important role to play, not only in reducing metal toxicity (see Section 3.3.1.2) but also in markedly increasing metal ion mobility.

Example 3.9: Determination of AlIII speciation at pH 2, 6, and 10 in an aqueous solution where {AlIIIT} ¼ 1 107 mol L1. Use log *b1 ¼ 4.95, log *b2 ¼ 10.55, log *b3 ¼ 17.25, log *b4 ¼ 22.85, log b1 ﬂuoride ¼ 7.00, log b2 ﬂuoride ¼ 12.70, log b1sulfate ¼ 3.50, log b2 sulfate ¼ 5.00, and log b1 org ¼ 5.9. Assume SF ¼ 5 108 mol L1, SSO42 ¼ 1 104 mol L1, SOrg ¼ 1 107 mol L1 and that no precipitation of Al(OH)3 occurs. Hydrolysis only: 2þ þ 3þ fAlIII T g ¼fAl g þ fAlðOHÞ g þ fAlðOHÞ2 g

þ fAlðOHÞ3 g þ fAlðOHÞ 4g ¼107 mol L1 ¼fAl3þ g 1 þ b1 =fHþ g þ b2 =fHþ g2 þ b3 =fHþ g3 þ b4 =fHþ g4 Main species: pH 2: Al31; pH 6: Al(OH)21, Al(OH)21; pH 10: Al(OH)4. Hydrolysis and one inorganic ligand, F: 2þ þ 3þ fAlIII T g ¼ fAl g þ fAlðOHÞ g þ fAlðOHÞ2 g þ fAlðOHÞ3 g 2þ þ 7 þ fAlðOHÞ mol L1 4 g þ fAlF g þ fAlF2 g ¼ 10

¼fAl3þ g ð1 þ b1 =fHþ g þ b2 =fHþ g2 þ b3 =fHþ g3 þ b4 =fHþ g4 þ b1 fluoride fF g þ b2 fluoride fF g2 Þ

Chemistry of Freshwaters

109

Main species: pH 2: Al31, AlF21; pH 6: Al(OH)21, Al(OH)21; pH 10: Al(OH)4. Hydrolysis and two inorganic ligands, F and SO2 4 : 2þ þ 3þ fAlIII T g ¼fAl g þ fAlðOHÞ g þ fAlðOHÞ2 g 2þ þ fAlðOHÞ3 g þ fAlðOHÞ 4 g þ fAlF g þ 7 mol L1 þ fAlFþ 2 g þ fAlSO4 g þ fAlðSO4 Þ2 g ¼ 10

¼fAl3þ g ð1 þ b1 =fHþ g þ b2 =fHþ g2 þ b3 =fHþ g3 þ b4 =fHþ g4 þ b1 fluoride fF g þ b2 fluoride fF g2 2 2 þ b1 sulfate fSO2 4 g þ b2 sulfate fSO4 g Þ

21 Main species: pH 2: Al31, AlF21, (AlSO1 4 ); pH 6: Al(OH) , 1 Al(OH)2 ; pH 10: Al(OH)4 .

Hydrolysis, two inorganic ligands, F and SO2 4 , and organic ligand, Org : 2þ þ 3þ fAlIII T g ¼fAl g þ fAlðOHÞ g þ fAlðOHÞ2 g 2þ þ fAlðOHÞ3 g þ fAlðOHÞ 4 g þ fAlF g þ 2þ þ fAlFþ 2 g þ fAlSO4 g þ fAlðSO4 Þ2 g þ fAlOrg g

¼ 107 mol L1 ¼fAl3þ g ð1 þ b1 =fHþ g þ b2 =fHþ g2 þ b3 =fHþ g3 þ b4 =fHþ g4 þ b1 fluoride fF g þ b2 fluoride fF g2 þ b1 sulfate fSO2 4 g 2 þ b2 sulfate fSO2 4 g þ b1 org fOrg gÞ

Main species: pH 2: Al31, AlF21, (AlSO41); pH 6: AlOrg21, Al(OH)21; pH 10: Al(OH)4.

110

Chapter 3

1.0

hydrolysis only

1.0

Al(OH) 4-

Al 3+

Al(OH) 4-

0.8

0.6

0.6 Al(OH) 2

0.4 0.2

Al(OH) 03

AlOH 2+ Al(OH) 2+

0.4

AlF 2+

0.2

AlSO4 +

0.0 0

Al 3+

AlOH 2+

0.0

hydrolysis + fluoride + sulfate

Al(OH) 30

AlF2+

1.0

hydrolysis + fluoride

1.0

Al(OH) 4-

0.4

Al(OH) 4-

0.8 Al 3+

AlOH 2+

0.6

Al(OH) 2

AlF 2+

0.4 Al(OH) 30

0.2 0.0

hydrolysis + fluoride + sulfate+ organic ligand

AlOrg

0.8 0.6

0.2

AlF2+

Al 3+ AlF 2+ +

AlOH 2+, Al (OH)2+, Al(O H)30

AlSO4 AlF2+

0.0

Inﬂuence of Ionic Strength. It should be noted that corrections to take account of ionic strength as discussed in Section 3.2.1.2 apply not only to the acid–base equilibrium constants but also to the stability constants for complex formation. 3.2.4.2 pe as a Master Variable. Chemical speciation is also inﬂuenced by the redox conditions prevailing in natural waters. Although redox reactions are often slow, and therefore species are present at activities far from equilibrium, they are commonly represented by thermodynamic equilibrium expressions, which can provide the boundary conditions towards which a system is proceeding. pe, a parameter describing redox intensity, gives the hypothetical electron activity at equilibrium. It measures the relative tendency of a solution to accept or donate electrons with a high pe being indicative of a tendency for oxidation, i.e. accepting electrons, while a low pe is indicative of a tendency for reduction, i.e. donating electrons. It is deﬁned as pe ¼ log{e}

(3.66)

The pe scale is thus analogous to the pH scale (pH ¼ log{H1}), since a low value for pe is obtained where the hypothetical {e} is large (pH is low where {H1} is large) and conversely a high value of pe is obtained

111

Chemistry of Freshwaters

where {e} is small (pH is high where {H1} is small). pe is related to the electrode potential, EH, by the expression pe ¼ 16.9 EH (V)

(3.67)

This relationship can be derived from the Nernst Equation EH ¼ EH0 þ (2.303 RT n/F) log(P (oxidized species)/ P (reduced species))

(3.68)

where R is the universal gas constant, T the absolute temperature, F the Faraday constant, and n the number of electrons. Starting with the equilibrium expression aA þ bB þ ne " cC þ dD

(3.69)

K ¼ {C}c{D}d/{A}a{B}b{e}n

(3.70)

we can write

Logarithmic expressions are then log K ¼ log ðPðreduced speciesÞ=Pðoxidized speciesÞÞ n logfe g pe ¼ ð1=nÞ log K þ ð1=nÞ log ðPðoxidized speciesÞ=Pðreduced speciesÞÞ Multiply by 2.303 RT/F to give ð2:303RT=FÞ pe ¼ ð2:303RT n=FÞ log K þ ð2:303RT n=FÞ log ðPðoxidized speciesÞ=Pðreduced speciesÞÞ Now use the standard Gibbs free energy, DG0 ¼ 2.303 RT logK ¼ –nFE H0 to give ð2:303RT=FÞ pe ¼ EH0 þ ð2:303RT n=FÞ logðPðoxidized speciesÞ=Pðreduced speciesÞÞ ð3:71Þ The right hand side of Equation (3.71) is the same as that of the Nernst Equation (3.68) and so the left hand side of Equation (3.71) must be equal to the left hand side of the Nernst Equation (3.68), thus leading to Equation (3.67). An additional deﬁnition is pe0 ¼ (1/n) log K

(3.72)

pe0 ¼ (2.303 RT/F ) EH0 ¼ 16.9 EH0

(3.73)

and so

112

Chapter 3

The usefulness of Equation (3.72) will be demonstrated further in Example 3.10. Although the electron activity is a hypothetical phenomenon, pe is a useful parameter to describe the redox intensity of natural systems and hence the species distribution under prevailing redox conditions (Example 3.10).

Example 3.10: Calculate the pe values for the following solutions (298 K, I ¼ 0) (i) a solution at pH 2 containing {Fe31} ¼ 104.5 mol L1 and {Fe21} ¼ 102.7 mol L1 where log K ¼ 13; (ii) a neutral solution containing {Mn21} ¼ 106 mol L1 in equilibrium with the solid phase, MnIVO2 and log K ¼ 40.84. (i) At pH 2, the hydrated metal ion has not undergone hydrolysis to any signiﬁcant extent: Fe3þ þ e Ð Fe2þ

K ¼ fFe2þ g=fFe3þ gfe g

pe ¼ log K þ logðfFe3þ g=fFe2þ gÞ ¼ 13 þ log ð104:5 =102:7 Þ ¼ 13 1:8 ¼ 11:2

(ii)

MnO2 þ 4H þ þ 2e Ð Mn2þ þ 2H2 O K ¼ fMn2þ g=fHþ g4 fe g2 pe ¼ 0:5 log K þ 0:5 log ðfHþ g4 =fMn2þ gÞ ¼ 20:42 þ 0:5 log ð1028 =106 Þ ¼ 20:42 11 ¼ 9:42

For a ﬁxed pH value (e.g. pH ¼ 2), the species distribution (e.g. where {FeT} ¼ 1 103 mol L1) at diﬀerent pe values can be illustrated by plotting log{ } against pe. This is analogous to the treatment of pH as a master variable. The pe range is split into two parts, peope0 and pe4pe0. Using pe0 ¼ (1/n) log K, this becomes {e}4K1 and {e}oK1 (because n ¼ 1 in this Example). Thereafter, construction of the graph is achieved by the same method used in Example 3.5.

113

Chemistry of Freshwaters 0

pε = 1/n log K {Fe2+}

{Fe3+}

-log { }

-5

pε

Redox intensity or electron activity in natural waters is usually determined by the balance between those processes which introduce oxygen (e.g. dissolution of atmospheric oxygen, photosynthesis) and those which remove oxygen (e.g. microbial decomposition of organic matter). Often these processes are controlled by the availability of inorganic nutrients such as phosphate and nitrate, e.g. as utilized in the formation of organic matter during photosynthesis (see Section 3.3.4). 2 1 106CO2 þ 16NO 3 þ HPO4 þ 122H2O þ 18H

" C106H263O110N16P1 þ 138O2

(3.74)

The decay of organic matter produced in this manner leads to the subsequent consumption of oxygen, e.g. by respiration. 2 C106H263O110N16P1 þ 138O2 " 106CO2 þ 16NO 3 þ HPO4 1 þ 122H2O þ 18H (3.75)

The decay of organic matter requires the presence of a terminal electron acceptor and in Equation (3.75) molecular oxygen is reduced to water. Other terminal electron acceptors present in natural waters include NO3, MnIV, FeIII, SO42, and CO2. Once all molecular oxygen has been consumed, organic matter is decomposed via reactions involving other terminal electron acceptors in a series determined by the pe intensity as shown in Table 3. The sequence of redox reactions involving organic matter, all of which are microbially mediated, can be thought of as progressing through

114

Chapter 3 Table 3

Redox processes – terminal electron acceptors

Terminal electron acceptor

O2 NO3 (reduction to N2) MnIV NO3 (reduction to NH41) FeIII Reducible organic matter SO42 CO2

13.75 12.65 8.9 6.15 0.8 3.01 3.75 4.13

Source: Adapted with permission from W. Stumm and J.J. Morgan, Aquatic Chemistry 3rd Edn,15 r Wiley, 1996.

decreasing levels of pe and so, for example, NO3 in denitriﬁcation reactions would be utilized as the electron acceptor before MnIV and SO42 would be utilized before CO2 (see Section 3.3.3.2). Redox processes are important for elements which can exist in more than one oxidation state in natural waters, e.g. FeII and FeIII, MnII, and MnIV. These are termed redox-sensitive elements. The redox conditions in natural waters often aﬀect the mobility of these elements since the inherent solubility of diﬀerent oxidation states of an element may vary considerably. For example, MnII is soluble whereas MnIV is highly insoluble. In oxic systems, MnIV is precipitated in the form of oxyhydroxides. In anoxic systems, MnII predominates and is able to diﬀuse along concentration gradients both upwards and downwards in a water column. This behaviour gives rise to the classic concentration proﬁles observed for Mn (and Fe) at oxic-anoxic interfaces as illustrated in Figure 2. Other examples of redox-sensitive elements include heavy elements such as uranium, plutonium, and neptunium, all of which can exist in multiple oxidation states in natural waters. Redox conditions in natural waters are also indirectly important for solute species associated with redox-sensitive elements. For example, dissolution of iron (hydr)oxides under reducing conditions may lead to the solubilization and hence mobilization of associated solid phase species, e.g. arsenate, phosphate (see Sections 3.3.2.1, 3.3.3.2, and 3.3.4.1). 3.2.4.3 pe – pH Relationships. Many redox equilibria also involve the transfer of protons and so chemical speciation depends on both pe and pH. Stability relationships involving redox processes can be investigated via pe–pH diagrams. The region of interest for natural waters displayed in these diagrams is deﬁned by the stability limits of water, i.e. where the total partial pressure of oxygen and hydrogen is no greater than 1 atm.

115

Chemistry of Freshwaters

Figure 2 Oxic-anoxic boundary in the water or sediment column15 (Reproduced with permission from ref 15, r John Wiley & Sons Inc. 1996)

Upper Limit O2(g) þ 4e þ 4H1 " 2H2O(1)

log K ¼ 83.1

(3.76)

Where pO2 ¼ 1, log pO2 ¼ 0 and log K ¼ 4 logfe g 4 logfHþ g pe ¼ 20:78 pH

ð3:77Þ

Lower Limit 2H1 þ 2e " H2(g)

log K ¼ 0

(3.78)

log K ¼ 28

(3.79)

or 2H2O þ 2e " H2(g) þ 2OH Where pH2 ¼ 1, log pH2 ¼ 0 and log K 2log Kw ¼ 2log{e} 2log {H1} or

log K ¼ 2log{e} 2log {H1}

In both cases: pe ¼ pH

(3.80)

For any system, stability relationships between solution phase species and solid phases can be used to construct pepH diagrams representing

116

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species distribution within the upper and lower stability limits of water. In general, the relationships between pe and pH for redox equilibria involving proton and electron transfer give lines, e.g. pe ¼ a pH þ b (where a and b are constants), with slope ¼ a. Redox equilibria involving only electron transfer give rise to horizontal lines, i.e. pe ¼ b, a ¼ 0. Vertical lines arise from equilibria involving only proton transfer and pH ¼ c (where c is a constant). The relationships between pe and pH for As–H2O and CrH2O systems at 298 K and 1 atm total are shown in Examples 3.11 and 3.12, respectively. The importance of both As and Cr speciation as represented in these pe–pH diagrams is discussed further in Sections 3.3.2.1 and 3.3.2.5, respectively. Example 3.11: Construction of a pe–pH graph for Cr-H2O. Assume that no solid phases are formed. Acid-base behaviour of CrVI: 2 1 HCrO K1 ¼ 106.45 4 " CrO4 þ H Hydrolysis behaviour of CrIII: Cr3þ þ H2 O Ð CrðOHÞ2þ þ Hþ þ CrðOHÞ2þ þ H2 O Ð CrðOHÞþ 2 þH þ 0 CrðOHÞ2 þ H2 O Ð CrðOHÞ3 þ Hþ þ CrðOHÞ03 þ H2 O Ð CrðOHÞ 4 þH

K2 K3 K4 K5

¼ 104:00 ¼ 105:62 ¼ 107:13 ¼ 1011:02

Reduction of CrVI to CrIII: 3þ þ HCrO þ 4H2 O 4 þ 7H þ 3e Ð Cr þ HCrO4 þ 6H þ 3e Ð CrðOHÞ2þ þ 3H2 O þ þ HCrO 4 þ 5H þ 3e Ð CrðOHÞ2 þ 2H2 O þ 2 þ CrO4 þ 6H þ 3e Ð CrðOHÞ2 þ 2H2 O 0 þ CrO2 4 þ 5H þ 3e Ð CrðOHÞ3 þ H2 O 2 þ CrO4 þ 4H þ 3e Ð CrðOHÞ 4

K6 ¼ 1070:55 K7 ¼ 1066:55 K8 ¼ 1060:93 K9 ¼ 1067:38 K10 ¼ 1060:25 K11 ¼ 1049:23

Construction of vertical lines: e.g. K1 ¼ {H1}{CrO2 4 }/{HCrO4 } On the line, the concentrations of the CrVI species are equal, so pH ¼ pK1 ¼ 6.45. Construction of sloped lines: 1 5 3 e.g. K8 ¼ {Cr(OH)þ 2 }/{HCrO4 }{H } {e } On the line, the concentrations of the CrIII and CrVI species are equal, so pe ¼ log K8/3 – 5/3 pH ¼ 20.31–1.6 pH. Crossing points: For the lines deﬁned by K1 and K8, substitute pH ¼ 6.45 in pe ¼ 20.31–1.6 pH, so pe ¼ 9.99.

117

Chemistry of Freshwaters 30

20 HCr O

Cr 3+

O2/H2O

Cr(OH) 2+

CrOH 2+

pε

CrO 42Cr(OH) 30 Cr(OH) 4-

H2O/H2

-10

Example 3.12: Construction of a pe–pH diagram for As–H2O. Assume that no solid phases are formed. Acid-base behaviour of AsV: þ H3 AsO4 Ð H2 AsO 4 þH 2 þ H2 AsO 4 Ð HAsO4 þ H 2 3 þ HAsO4 Ð AsO4 þ H

K1 ¼ 102:24 K2 ¼ 106:76 K3 ¼ 1011:60

Acid-base behaviour of AsIII: þ H3 AsO3 Ð H2 AsO 3 þH 2 H2 AsO3 Ð HAsO3 þ Hþ

K4 ¼ 109:23 K5 ¼ 1012:10

Reduction of AsV to AsIII: H3 AsO4 þ 2Hþ þ 2e Ð H3 AsO3 þ H2 O þ H2 AsO 4 þ 3H þ 2e Ð H3 AsO3 þ H2 O 2 þ HAsO4 þ 4H þ 2e Ð H3 AsO3 þ H2 O þ HAsO2 4 þ 3H þ 2e Ð H2 AsO3 þ H2 O þ AsO3 4 þ 4H þ 2e Ð H2 AsO3 þ H2 O 3 2 þ AsO4 þ 3H þ 2e Ð HAsO3 þ H2 O

K6 ¼ 1019:46 K7 ¼ 1021:70 K8 ¼ 1028:46 K9 ¼ 1019:23 K10 ¼ 1030:83 K11 ¼ 1018:73

Construction of vertical lines: e.g. K1 ¼ {H1}{H2AsO 4 }/{H3AsO4} On the line, the concentrations of the AsV species are equal, so pH ¼ pK1 ¼ 2.24. Construction of sloped lines: e.g. K6 ¼ {H3AsO3}/{H3AsO4}{H1}2 {e}2

118

Chapter 3

On the line, the concentrations of the AsIII and AsV species are equal. so pe ¼ log K/2 – pH ¼ 9.73–pH. Crossing points: For the lines deﬁned by K1 and K6, substitute pH ¼ 2.24 in pe ¼ 9.73– pH, so pe ¼ 7.49. 20 H3AsO4

O2/H2O H2AsO4

−

AsO43−

pε

HAsO42− H3AsO3

0 H2O/H2

H2AsO3−

−10

HAsO32−

In Examples 3.11 and 3.12 it has been assumed that no solid phases are formed. At higher solution concentrations, however, the formation of solid phases such as insoluble hydroxides, carbonates and sulﬁdes must be considered. The stability of these phases is often pe and/or pH dependent. Redox equilibria involving both aqueous and solid phase ‘S’ species are shown in Example 3.13.

Example 3.13: Construction of a pe–pH diagram for SO4–S(s)–H2S system. Use ST ¼ 102 mol L1. Acid-base behaviour of SVI: 2 1 HSO 4 " SO4 þ H

K1 ¼ 102.0

Acid-base behaviour of SII: H2S(aq) " HS þ H1

K2 ¼ 107.0

Reduction of SVI to S0: 1 HSO 4 þ 7H þ 6e " S(s) þ 4H2O

K3 ¼ 1034.2

119

Chemistry of Freshwaters 1 SO2 4 þ 8H þ 6e " S(s) þ 4H2O

K4 ¼ 1036.2

Reduction of SVI to SII: 1 SO2 4 þ 9H þ 8e " HS þ 4H2O

K5 ¼ 1034.0

Reduction of S0 to SII: S(s) þ 2H1 þ 2e " H2S(aq)

K6 ¼ 104.8

Construction of sloped lines involving a solid phase: pe ¼ log K6/2 – log{H2S(aq)}/2 – pH On the line, assume that ST ¼ log{H2S(aq)} ¼ 102 mol L1 so pe ¼ 3.4–pH. 20

HSO4-

SO42-

pε

S(s) 0

H2S HS-

-10

The information contained in this pe–pH diagram is particularly useful in the determination of stability zones for solid phase metal sulﬁdes (Example 3.14). For example, it is the reduction of sulfate to sulﬁde which determines the stability ﬁeld of pyrite (FeS2).

Example 3.14: Construction of a pe–pH diagram for the Fe–O2–‘S’– CO2–H2O open system where {FeT} is 1 106 mol L1, {S(aq) T} ¼ 1 102 mol L1 and pCO2 ¼ 103 atm. Reduction of FeIII to FeII: Fe31 þ e " Fe21 Fe(OH)21 þ e þ H1 " Fe21 þ H2O

K1 ¼ 1013.0 K2 ¼ 1015.2

120

Chapter 3

1 21 Fe(OH)þ þ 2H2O 2 þ e þ 2H " Fe

K3 ¼ 1018.7

Fe(OH)3(s) þ e þ 3H1 " Fe21 þ 3H2O

K4 ¼ 1017.9

Fe(OH)3(s) þ e þ H1 " Fe(OH)2(s) þ H2O Fe(OH) 4 þ e

þ 2H1 " Fe(OH)2(s) þ 2H2O

Reduction of FeII to Fe0: Fe21 þ 2e " Fe(s) Fe(OH)2(s) þ 2e þ 2H1 " Fe(s) þ 2H2O Hydrolysis of FeIII: Fe31 þ H2O " Fe(OH)21 þ H1 1 Fe(OH)21 þ H2O " Fe(OH)þ 2 þ H 1 Fe(OH)þ 2 þ H2O " Fe(OH)3(s) þ H

K5 ¼ 105.5 K6 ¼ 1022.2

K7 ¼ 1013.8 K8 ¼ 101.4

K9 ¼ 102.2 K10 ¼ 103.5 K11 ¼ 100.8

1 Fe(OH)3(s) þ H2O " Fe(OH) 4 þ H

K12 ¼ 1016.7

Hydrolysis of FeII: Fe21 þ 2H2O" Fe(OH)2(s) þ 2H1

K13 ¼ 1012.4

Carbonate complexation of FeII and additional reduction reactions: K14 ¼ 105.8 FeCO3(s) þ H1 " Fe21 þ HCO 3 1 12.1 Fe(OH)3(s) þ HCO 3 þ 2H þ e " FeCO3(s) þ 3H2O K15 ¼ 10 1 32.4 Fe(OH) 4 þ HCO3 þ 3H þ e " FeCO3(s) þ 4H2O K16 ¼ 10

In the stability ﬁeld of sulfate: 21 þ 16H1 þ 14e " FeS2(s) þ 8H2O K17 ¼ 1086.8 SO2 4 þ Fe 1 FeCO3(s) þ 2SO2 4 þ 14e þ 18H " FeS2(s) þ H2CO3 þ 8H2O K18 ¼ 1085.4

At lower pe values (in the H2S and HS stability ﬁelds): K19 ¼ 105.0 FeS2(s) þ 4H1 þ 2e " Fe21 þ 2H2S(aq) FeS2(s) þ H1 þ 2e " FeS(s) þ HS

K20 ¼ 1014.6

121

Chemistry of Freshwaters 3+

Fe Fe(OH)

Fe(OH)2

20 15

H2O/O2

pε

Fe2+ + H2S

Fe(OH)3 ferrihydrite

Fe2+

Fe(OH)4-

0 FeS2 pyrite

-5

FeCO

3 sideri

H2 /H2O

-10

FeS

-15

8 pH

It should be remembered that these diagrams are based on equilibrium conditions and that redox reactions in natural waters may be far from equilibrium. Additional assumptions are often made, e.g. the activities of the main components are ﬁxed and that those of all minor species are known, the activity of all solid phases is unity, and only a few reactive components need to be considered.34 As a consequence, straight line boundaries between species are generated. Importantly, adsorption, ion exchange, and metal-organic interactions are often omitted, limiting the usefulness of classical pe–pH diagrams (e.g. Examples 3.11–3.14).34 Generation of more environmentally representative diagrams requires the use of geochemical computer models (see Section 3.2.5). For example, Kinniburgh and Cooper34 demonstrate how PHREEQC35 and ORCHESTRA36 can be used to obtain predominance diagrams (including pe–pH diagrams) for a wide range of metals, metalloids, and anionic species. 3.2.5 Modelling Aquatic Systems This chapter has discussed the fundamental importance of the master variables pe and pH in determining species distribution under the conditions prevailing in natural systems (pe 10 to þ 17 and pH 4 to 10). The equilibria contained in Sections 3.2.2–3.2.4 also provide the basis for many of the currently available geochemical computer models (e.g. WATEQ4F,30 PHREEQC,35 and MINTEQA222), which are being

122

Chapter 3

used to predict the geochemical behaviour of metal ions, organic pollutants, etc. As discussed in Section 3.2.1.2, certain equations, e.g. Davies equation for activity coeﬃcient calculation, are contained in the source code of these models. In contrast, a recently developed model, ORCHESTRA (Objects Representing CHEmical Speciation and TRAnsport), enables the user to deﬁne equations in text format, which can then be read by the ORCHESTRA calculation kernel.36 The objectoriented structure of this model also allows the user to choose appropriate chemical models, e.g. for activity coeﬃcient calculation (see Section 3.2.1.2), adsorption at organic surfaces (MODEL V,37 NICA38), and adsorption at oxide surfaces (CD-MUSIC39). Overall, geochemical computer models can be extremely useful in the description of chemical equilibria occurring in the aquatic environment. In some cases, predictions about reaction kinetics and transport of species can also be made. The application of geochemical models is not limited to natural aquatic systems but has been usefully extended to predict the eﬀectiveness of certain remediation strategies in the treatment of waters emanating from contaminated sites.40 3.3 CASE STUDIES 3.3.1 Acidiﬁcation 3.3.1.1 Diatom Records. The onset and ﬂourishing of the Industrial Revolution, with its dependence for energy upon the combustion of fossils fuels, especially coal, released huge quantities of sulfur dioxide and nitrogen oxides to the atmosphere, which ultimately were returned to the land surface in the form of acid precipitation (H2SO4, HNO3). The eﬀect of this upon freshwater lakes, especially in poorly buﬀered catchment areas (e.g. low-carbonate soils, granitic bedrock) was a signiﬁcant decrease in pH and a concomitant decline in biological productivity, often leading to clear acid waters devoid of ﬁsh. That the acidiﬁcation of freshwater lakes is a comparatively recent (post1850) phenomenon linked to acid deposition has been demonstrated by the diatom record in the sediments at the bottom of acidiﬁed lakes.41 Diatoms are algae with siliciﬁed cell walls. Depending upon the acidity of the water, some species ﬂourish more than others. Upon deposition, these siliceous remains are preserved in the sediment column. The dramatic change in pH after 1850 inferred from the diatom record in dated sediments from the Round Loch of Glenhead in the Galloway region of south-west Scotland rules out long-term acidiﬁcation processes as the cause.42

Chemistry of Freshwaters

123

3.3.1.2 Aluminium. The eﬀect of such acidiﬁcation upon the mobilization, behaviour, and subsequent biological impact of aluminium in freshwaters has been spectacular. Constituting 8.13% of the Earth’s crust, aluminium is the third most abundant element and is largely associated with crystalline aluminosilicate minerals of low solubility. In the absence of strong acid inputs, processes of soil development normally lead to only a small fraction of aluminium becoming available to participate in biogeochemical reactions. Thus, in the Northern Temperate Region, mobilization of aluminium from upper to lower mineral horizons by organic acids (or H2CO3) leached from foliage or forest ﬂoors often leads to the formation of Al(OH)3 in a process known as podzolization. Under elevated partial pressures of CO2, however, dissociation of H2CO3 produces H1, which may solubilize aluminium, while HCO3 serves as a counterion for transport of cationic AlIII(aq) through soil. Upon discharge to a surface water, CO2 degasses, resulting in aluminium hydrolysis and precipitation as Al(OH)3. Thus, in most natural waters, concentrations of dissolved aluminium are generally low due to the relatively low solubility of natural aluminium minerals under circumneutral pH values. In the case of strong acid (H2SO4, HNO3) inputs, however, especially to sensitive regions with small pools of basic cations (Ca21, Mg21, Na1, K1) and an inability to retain inputs of strong acid anions (SO42, NO3, Cl), acidic cations (H1, AlIII) are transported along with acid anions from soil to surface water.43 The speciation of dissolved aluminium in surface waters is highly pH-dependent (see Section 3.2.3.2), with, in a simple system (considering only monomeric aluminium species) in equilibrium with solid phase Al(OH)3, the aquated ion Al31(aq) predominating at pHo3, progressive hydrolysis to Al(OH)21 and Al(OH)þ 2 in the pH range 4.5–6.5, followed by Al(OH)30 from pH 6.5 to 7, and then Al(OH)4 predominating as the pH increases beyond 7. From pH 4.5–7.5, the solubility of aluminium is low and in this range it is often precipitated as Al(OH)3. Below pH 4.5 and above pH 7.5 the concentration of aluminium in solution increases rapidly. In a part experimental fractionation, part speciation modelling study of aluminium in acidiﬁed Adirondack surface waters in the north-eastern USA, Driscoll and Schecher43 found that, in the absence of organic ligands, Al31(aq) was signiﬁcant at pHo5, and that the inorganic monomeric complexes of aluminium predominated at pH45. In particular, Al–F complexes were the dominant form at pH 5–6 whereas, at pH46, Al–OH species became the major form. In the presence of organic ligands, however, they found that, over a broad pH range of 4.3–7.0, alumino-organic complexes were a major component of the monomeric aluminium. It is the chemically labile inorganic

124

Chapter 3

monomeric aluminium that has been identiﬁed as the toxic agent for ﬁsh which, in acidiﬁed lakes and streams at aluminium concentrations of 100–200 mg L1, are unable to maintain their osmoregulatory balance and are susceptible to respiratory problems from coagulation of mucus and Al(OH)3 on their gills at their physiological pH of 7.2.44 In response to a decline in acid deposition as a result of improved atmospheric emission controls, there was a signiﬁcant decline from the early 1980s to the late 1990s in the concentrations of both inorganic and organic monomeric aluminium in soil solutions draining from mineral soils and in stream water in New Hampshire, USA (Figure 3).45 These decreases in aluminium concentration were accompanied by decreases in SO42 concentrations, but there was only a modest increase in stream water acid-neutralizing capacity and no change in pH. The hydroxylation of aluminium appears to have been critical in buﬀering stream waters against such increases. In a survey of Adirondack lakes in the north-eastern USA, however, Driscoll and co-workers measured a

Figure 3

Long-term declines in sulfate concentrations mitigate the mobilization of aluminium in soil solutions and stream water in the Hubbard Brook Experimental Forest, New Hampshire, USA.45 a–c, annual volume-weighted concentrations of a, inorganic monomeric aluminium (Ali); b, organic monomeric aluminium (Alo); and c, sulfate (SO42) in mineral soil solutions at 750 m (K) and 730 m (J), and in stream water (triangles) (Reprinted with permission from ref 45, r Macmillan Publishing, 2002)

Chemistry of Freshwaters

125

signiﬁcant increasing trend in acid-neutralizing capacity of 1.60 meq L1 per year from 1992 to 2000 for 60% of the monitored lakes.46 Furthermore, there was a growing concentration of dissolved organic carbon in 15% of the lakes, helping to shift the aluminium from its toxic inorganic form towards less toxic organic forms although likely to delay recovery of acid-neutralizing capacity. Overall, however, by 2000 both USA and European surface waters appeared to be on the way to long-term recovery from acidiﬁcation, with decreases in SO42 concentrations, increases in pH and acid-neutralizing capacity, and decreases in concentrations of labile toxic forms of aluminium.47–49 With respect to human exposure to aluminium in drinking water, it must be remembered that aluminium has often been deliberately added, in the form of Al2(SO4)3, to water supplies at treatment works to remove coloration by organic compounds in reservoirs in upland catchments. This it achieves by hydrolyzing to a gelatinous, high surface area, precipitate of Al(OH)3, which helps to remove the coloured organic colloids. It has been suggested, partially because of observed water aluminium-related dialysis dementia in some chronic renal patients on artiﬁcial kidney dialysis machines in the 1960s and 1970s, that exposure to low concentrations of aluminium in drinking water, for which there is a current EC Maximum Admissible Concentration of 200 mg L1, might be implicated in Alzheimer’s Disease.50 There is still no proof of this at present,51 although it is known that the incidence of senile dementia on the remote Paciﬁc island of Guam, where bauxite is mined and the local water is elevated in aluminium, is a hundred-fold greater than in the USA. Ten years after an accident at Camelford in south-west England in 1988, when 20 tonnes of Al2(SO4)3 were emptied directly into the main water supply, there did appear to be a lasting eﬀect on the cerebral function of those who had been exposed to greatly enhanced aluminium concentrations in the tap water.52 It has been mooted that, for humans, a high intake of silicic acid (H4SiO4) could reduce the bioavailability of aluminium by two orders of magnitude for, at slightly alkaline intestinal pH, hydroxyaluminosilicates are stable and unavailable for absorption.53 This would mimic the situation in nature, where, in the absence of a strong acid input, the bioavailability of aluminium is kept low by the formation of hydroxyaluminosilicate species. The latter has been described as a ‘geochemical brake on an aluminium juggernaut which would otherwise career into the biotic cycle’.54 More generally, in addition to aluminium, the role of other metals, including zinc, copper, iron, and manganese, in neurodegenerative disease is now under scrutiny, there being suggestions of adverse eﬀects upon protein folding and subsequent oxidative stress.55

126

Chapter 3

3.3.1.3 Acid Mine Drainage and Ochreous Deposits. One of the major pollution problems aﬀecting freshwaters is that of acid waters discharged from coal and metal mines.56 In an active underground coal mine, pumping lowers the water table. The deeper strata thus become exposed to air with the result that pyrite (FeS2) present is subject to oxidation, generating acid conditions. 2FeS2(s) þ 2H2O þ 7O2(g) " 2FeSO4(s) þ 2H2SO4(1)

(3.81)

When mining and pumping cease, the water table returns to its natural level. While the ﬂooding of the mine stops the direct oxidation of pyrite, it does bring the sulfuric acid and iron sulfates into solution. Some of the ferrous ion (Fe21) may be oxidized to the ferric ion (Fe31), a slow process at low pH, but one which can be catalysed by bacteria, and the ferric ion may react further with pyrite.57 4Fe21 þ O2(g) þ 4H1 " 4Fe31 þ 2H2O

(3.82)

1 FeS2(s) þ 14Fe31 þ 8H2O " 15Fe21 þ 2SO2 (3.83) 4 þ 16H

Underground in the Richmond Mine at Iron Mountain, California, extremely acidic mine waters have been encountered, with pH ranging from 1.5 to 3.6 and total dissolved metal (mainly iron) and sulfate concentrations as high as 200 g L1 and 760 g L1, respectively.58 These are the most acidic waters known. On reaching the surface, the acidic mine drainage water is mixed with air and oxygenated water, leading to rapid oxidation from the ferrous to the ferric form and the precipitation of the characteristic unsightly orange ochreous deposits of iron (hydr)oxides, e.g. ferrihydrite (Fe(OH)3), goethite (FeOOH), observed along many stream and river beds. 4Fe21 þ O2(g) þ 4H1 " 4Fe31 þ 2H2O

(3.84)

Fe3þ þ 3H2 O Ð FeðOHÞ3ðsÞ þ3Hþ

ð3:85Þ

ferrihydrite

Fe3þ þ 2H2 O Ð FeOOHðsÞ þ3Hþ

ð3:86Þ

goethite

The overall process may be represented as 2FeS2(s) þ 15/2O2(g) þ 7H2O " 2Fe(OH)3(s) þ 4H2SO4(1) (3.87) Microorganisms (e.g. Thiobacillus thiooxidans, Thiobacillus ferrooxidans, Metallogenium) are involved as catalysts in many of the oxidizing

Chemistry of Freshwaters

127

reactions, which would be extremely slow under the prevailing conditions of low pH (e.g. o4.5).59 The combination of acid waters and coatings can have devastating eﬀects upon aquatic biota, with depletion of free-swimming and bottom-dwelling organisms, the loss of spawning gravel for ﬁsh and direct ﬁsh mortalities. Typical approaches to treatment of acid mine drainage have included methods to remove the iron ﬂoc by oxidation and to adjust the pH through the use of limestone ﬁlter beds. 3.3.1.4 Acid Mine Drainage and Release of Heavy Metals. The production of sulfuric acid from sulﬁde oxidation in mines can also lead to the leaching of metals other than iron, with the result that the emerging acidic waters may be laden with heavy metals. Thus in the former metalliferous mining area of south-west England, where there are now many abandoned mine workings, the closure and ﬂooding of the famous Wheal Jane tin mine in 1992 led to a highly acidic co*cktail of dissolved metals (Cu, Pb, Cd, Sn, As) entering the Carnon and Fal Rivers at 7–15 million litres per day and spreading throughout the surrounding estuaries and coastal waters.60 In the USA, the EPA has identiﬁed over 31,000 hazardous waste sites, with the largest complex of ‘Superfund’ sites to be remediated in western Montana, in the Clark Fork River Basin where there have been more than 125 years of copper and silver mining and smelting activities. Moore and Luoma61 have characterized three types of contamination resulting from large-scale metal extraction: primary, consisting of wastes produced during mining, milling, and smelting and deposited near their source of origin; secondary, resulting from transport of contaminants away from these sites by rivers or through the atmosphere to soils, groundwaters, rivers, etc.; and tertiary, where contaminants may be remobilized many kilometres away from their point of origin. More generally, Furrer and co-workers,62 in an examination of ﬂocs from streams polluted from acid mine drainage in California and Germany, found that when aluminium-rich acid mine drainage mixes with nearneutral surface waters, ﬂuﬀy aluminium oxyhydroxide ﬂocs precipitate and move downstream as suspended solids, transporting adsorbed pollutants, e.g. heavy metal cations such as Pb21, Cu21, and Zn21. The ﬂocs appear to be formed from aggregation of the aqueous polyoxocation, AlO4Al12(OH)24(H2O)1271 [Al13], which is toxic and possibly responsible for the declining ﬁsh populations in rivers polluted by acid mine drainage. Some of the largest waste deposits occur in tailing ponds containing acid mine water. In the Clark Fork River Complex, it is estimated that

128

Chapter 3

ponds contain approximately 9000 t arsenic, 200 t cadmium, 90,000 t copper, 20,000 t lead, 200 t silver, and 50,000 t zinc. These metals enter streams and rivers as solutes and particulates and contaminate sediments in the river and reservoirs far downstream from the primary sources (Figure 4).61 Downstream concentrations follow an exponential decline when viewed over several hundred kilometres. The sediment of Milltown Reservoir, more than 200 km from the mines and smelters at Butte and Anaconda, is highly contaminated with various metals and arsenic. It is believed that oxidation-reduction processes release arsenic from the reservoir sediments and cause contamination of an aquifer from which water drawn through wells, now closed, contains arsenic higher than the EPA drinking water standards.61 Davis and Atkins,63 however, have pointed out that, in the bottom sediments of the Clark Fork River channels, the fraction of ﬁne-grained (o63 mm) clay/silt material, which contains high metal concentrations and to which benthic organisms are likely to be directly exposed, is small relative to sands and gravels. Furthermore, the metals occur predominantly in sulﬁdes frequently armoured with an oxide rim and other sparingly soluble phases.

Figure 4 Types of contamination resulting from large-scale metal extraction (Reprinted with permission from ref 61, r American Chemical Society, 1990)

Chemistry of Freshwaters

129

When open-pit mining ended in the Clark Fork River Complex in 1982, pumping was discontinued, with the result that water began ﬁlling underground shafts and tunnels and the 390 m deep Berkeley pit. Contaminated acid water containing individual metal and sulfate concentrations thousands of times those in uncontaminated water could ﬂow into an adjacent alluvial aquifer and eventually over the rim of the pit. Large-scale hard-rock mining in the western USA, especially Nevada, has greatly increased in recent years, with the result that deep ‘pit lakes’ are likely to form as open-pit metal mines intersecting groundwater are depleted and shut down.64 Pit lakes in high sulﬁde rock will tend to have poor quality acidic water, although oxidized rock that contains signiﬁcant carbonate will produce better quality, nearneutral pit lake water. Most of the larger existing pit lakes currently contain water that does not meet standards for drinking water, agricultural water quality, or aquatic life.64 Factors such as the oxygen status of the lake, pH, the hydrogeologic ﬂow system, composition of the wallrock, evapo-concentration, biological activity, and hydrothermal inputs are all important to the modelling of future water quality and impact. Remediation approaches in circ*mstances like those described here are based on a variety of physical, chemical, and biological systems. These include the construction of ponds where suﬃcient organic matter is available to establish anaerobic conditions and immobilize at least some of the metals (e.g. Cu, Pb, Cd, Zn) as sulﬁdes. Similar passive treatment of tailings-impacted groundwater has also employed precipitation of metal sulﬁdes as the key clean-up step.65 Active chemical treatment with lime, producing sludges, and biosorption by reeds, wetlands, etc., are other methods which have been tried. Subsequent to the ﬂooding of the Wheal Jane mine, an active treatment (chemical neutralization) plant and a passive treatment plant, the latter featuring three schemes diﬀering only in the pre-treatment method (lime dosing, anoxic limestone drain, lime-free) used to modify the pH of the inﬂuent mine water, were set up. All three systems at the Wheal Jane Passive Treatment Plant included (i) constructed aerobic reed beds designed to remove iron and arsenic, (ii) an anaerobic cell to encourage reduction of sulfate and facilitate removal of zinc, copper, cadmium, and the remaining iron as metal sulﬁdes, and (iii) aerobic rock ﬁlters designed to promote the growth of algae and facilitate the precipitation of manganese.66 The study showed considerable success in the removal of key toxic metals and clearly demonstrated the potential for natural attenuation of acid mine drainage, particularly iron oxidation, by microbial populations.

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Chapter 3

3.3.2 Metals and Metalloids in Water 3.3.2.1 Arsenic in Groundwater. In what has been described as the largest arsenic poisoning epidemic in the world, hundreds of thousands of people in West Bengal (India) have been seriously aﬀected by arsenic poisoning resulting from the consumption of water drawn from tube wells sunk some 20 to 150 m below the ground into aquifers.67 In West Bengal, Bangladesh, and other parts of the world (e.g. Chile, Mexico, Argentina, Ghana, Mongolia, and Taiwan) where inorganic arsenic concentrations in drinking water are elevated (mg L1 vs. WHO recommended limits of 10 mg L1), exposure has resulted in hyperpigmentation, hyperkeratosis, and cancer of the skin and various internal organs.68–72 Sulﬁde minerals are one of the most important natural sources of arsenic in groundwater. Oxidation of arsenopyrite (FeAsS), in analogous fashion to pyrite (FeS2), may release high concentrations of arsenic into solution.73,74 2 1 4FeAsS(s) þ 13O2(g) þ 6H2O " 4Fe21 þ 4AsO3 4 þ 4SO4 þ 12H (3.88)

In Bengal and Bangladesh, it has been suggested that large-scale withdrawal of groundwater for irrigation produces seasonal ﬂuctuation of the water table, which in turn results in intake of oxygen into the pore waters of sediments that are arsenic-rich in the form of arsenopyrite.73,74 The exact speciation of soluble inorganic arsenic, e.g. as the undissociated forms or diﬀerent oxyanions of the acids H3AsIIIO3 and H3AsVO4, will be dependent upon the prevailing pe and pH. Sorption of arsenic, especially pentavalent arsenate (e.g. AsO43, HAsO42, H2AsO4), onto ferric hydroxide (Fe(OH)3) produced under oxidizing conditions may, however, restrict its mobility and availability, although fertilizer phosphate present in groundwater could perhaps displace the arsenate. Arsenite, especially as H3AsIIIO3 (pK1 ¼ 9.2), the predominant form under reducing conditions at pHo9.2, is much less strongly sorbed. As an alternative to the theory of arsenopyrite oxidation, however, especially where there is an absence of elevated concentrations of sulfate, it has been proposed that elevated arsenic concentrations in Bengal and Bangladesh groundwater could result from the release of adsorbed arsenate during the dissolution of hydrous iron oxides (similar to ferric hydroxide) under reducing conditions (see Section 3.3.3.2).73–76 This phenomenon can be seen in those parts of Asia where sub-strata consist of recent alluvial sediments, rich in decaying organic matter, that were deposited and buried during the post-glacial period. Indeed, McArthur

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and co-workers77 believe that the proximity of soluble organic matter to drive FeOOH reduction is the key and that this can only occur at the margins of buried peat basins. There is an alternative theory, propounded by Harvey and co-workers,78 that the arsenic mobilization is associated with the inﬂow of younger carbon, probably untreated human waste in water that has been drawn into the aquifers of Bangladesh as a consequence of irrigation pumping. Although trivalent inorganic arsenic, with its propensity for binding to the SH group of enzymes, is acknowledged to be more toxic to humans than pentavalent inorganic arsenic, it must be recognized that AsV can be converted to AsIII in the human body as part of the reduction/biomethylation pathway of excretion AsV - AsIII - MMAA - DMAA

(3.89)

where MMAA (CH3AsO(OH)2, monomethylarsonic acid) and DMAA ((CH3)2AsO(OH), dimethylarsinic acid) are less toxic metabolites, the latter predominating as the major form of arsenic excreted in urine.68–72 There is evidence to suggest, however, that trivalent metabolic intermediates of MMAA and DMAA may be highly implicated in the initiation of cancer.70–72 Methylation could also occur naturally by microbial action in freshwaters, although reported occurrences suggest that the eﬀect is small.79 3.3.2.2 Lead in Drinking Water. The naturally soft, slightly acidic, plumbosolvent water of the Loch Katrine water supply for the Glasgow area was recognized many years ago to release lead from the lead pipes and tanks in the domestic plumbing of the Victorian and subsequent (even post-World War II) eras.80 O2(g) þ 4H1 þ 4e " 2H2O Pb21 þ 2e " Pb(s) O2(g) þ 4H1 þ 2Pb(s) " 2Pb21 þ 2H2O

E0H ¼ 1.229 V (3.90) E0H ¼ 0.126 V (3.91) E0H ¼ 1.355 V

(3.92)

Essentially, elemental lead becomes soluble in acidic conditions due to its oxidation by dioxygen. Furthermore, compounds such as carbonate and hydroxycarbonate compounds of lead, i.e. PbCO3 and Pb3 (CO3)2(OH)2, that may coat the pipes, will dissolve under acid conditions. In view of the concern over detrimental eﬀects of exposure to lead upon human health, in particular the possible impact upon intelligence

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and behaviour of young children, steps were taken in the mid-1970s to reduce the lead content of tap water in Glasgow and other susceptible areas, which often exceeded the WHO maximum guideline at the time of 100 mg L1. The method chosen was adjustment of pH to 8–9 by lime dosing. The eﬀects of liming the Glasgow water supply were quite dramatic. Whereas pre-1978, when the pH was 6.3, only 50% of random daytime samples were o100 mg L1 in lead, the ﬁgure increased to 80% during 1978–1980, when the adjusted pH was 7.8. After 1980, when the pH was increased further to 9.0, 95% of samples were o100 mg L1.80 It appears that carbonate and hydroxycarbonate lead compounds present in the coatings on the pipes were stabilized. Signiﬁcant reductions in blood lead levels of key exposed groups (e.g. pregnant women) were also observed. Since 1989, as regulatory upper limits for lead in drinking water have fallen, e.g. to 50 mg L1 (EC) and now to 10 mg L1 (WHO), orthophosphate has been added to the water supply in Glasgow to precipitate insoluble lead compounds such as Pb3(PO4)2 and Pb5(PO4)3OH. This has resulted in a fall in the proportion of households with water lead 410 mg L1 from 49% in 1981 to 17% in 1993.81 Despite this improvement, an estimated 13% of infants were still exposed via bottle feeds to tap water lead concentrations in excess of 10 mg L1 and it seems very unlikely that further treatment of the water supply will be able to guarantee water lead concentrations o10 mg L1. Occasionally, elevated concentrations of lead in tap water arise from the illegal use of lead solder in capillary joints to join copper pipes. Cu21 þ 2e " Cu(s)

E0H ¼ 0.337 V

(3.93)

Cu21 þ Pb(s) " Cu(s) þ Pb21

E0H ¼ 0.463 V

(3.94)

In Washington DC, USA, where the water disinfection programme was recently modiﬁed from chlorine to chloramines, an unexpected appearance of high levels of lead in the tap water has been partly attributed to similar galvanic corrosion when copper pipes are connected to leadcontaining brass ﬁttings in the presence of chloramines. With the shift to chloramines, the oxidizing potential of Washington DC’s water was lowered and not only could the brass become highly anodic and the copper cathodic, but also the previously deposited PbO2 scales on the inside of lead service pipes began to dissolve.82 3.3.2.3 Cadmium in Irrigation Water. In the 1950s in Japan, many people, especially menopausal women suﬀering from malnutrition, low

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vitamin D intake, and calcium deﬁciency, suﬀered a condition known as Itai-Itai (Ouch-Ouch) disease, with symptoms ranging from lumbagotype pains to multiple fractures of softened bones. The cause was irrigation water from a river (Jintsu) chronically contaminated with dissolved cadmium from a zinc mining and smelting operation. Contaminated rice from the irrigated paddy ﬁelds was eaten and the Ca21 in bones replaced by Cd21, an ion of the same charge and size. Although this is a particularly extreme example of acute eﬀects resulting from high exposure to this non-essential element, concern has been expressed about chronic eﬀects (e.g. kidney damage, hypertension) from possible enhanced exposure of humans through increased application of sewage sludge to agricultural land, in view of ECenforced cessation of dumping at sea by 1999. Compared with other heavy metals, cadmium exhibits an especially high mean sludge concentration (20 mg kg1) relative to mean soil concentration (0.4 mg kg1).83 In acidic soils the concentration of Cd21 available for uptake by plants can be substantial,84 as it adsorbs only weakly onto clays, whereas at pH 47 it readily precipitates, e.g. as CdCO3 for which the solubility product, KSP ¼ 1.8 1014, is indicative of the advantages of liming. Similarly, in drinking water, the presence of dissolved carbonate at a concentration of 5 104 mol L1 can reduce the solubility of cadmium from 637 to 0.11 mg L1, in line with evidence that hard water, with a high calcium content, can protect against cadmium.85 3.3.2.4 Selenium in Irrigation Water. In 1983, high rates of embryonic deformity and death, attributed to selenium toxicosis, were found in wild aquatic birds at the Kesterson National Wildlife Refuge in the San Joaquin Valley of California.86 Kesterson Reservoir was a regional evaporation pond facility where drainage waters, often containing high level of salts and trace elements (including selenium), from irrigated farmland had been collected since 1978. With evapotranspiration, greatly in excess of precipitation, bringing soluble salts to the surface of farmland in the arid climate of the westcentral San Joaquin Valley, and crop productivity, after irrigation, threatened by shallow saline groundwater near the root zone, grids of sub-surface tile drains were constructed to divert the saline waters to a collective drain (San Luis), which ﬂowed into Kesterson Reservoir. It was the geologic setting of the San Joaquin Valley as well as the climate, however, which led not only to the soil salinization but also to the presence of selenium (in the form of SeO42) in Kesterson inﬂow waters at concentrations well in excess of the USEPA designation of 1000 mg L1 for selenium as a toxic waste, never mind its 10 mg L1 limit for

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drinking water and the o2.3 mg L1 limit subsequently suggested for the protection of aquatic life. To the west of the San Joaquin River (Figure 5), selenium in the soils is believed to be of natural origin. During the Jurassic and Cretaceous periods, there was deposition of marine sediments comprising sandstones, shales, and conglomerates, including seleno-sulﬁdes of iron (FeS2 þ FeSe2). Subsequent uplifting of these sediments produced the Coast Ranges and

Figure 5

The Kesterson effect: biogeochemical cycling of selenium in the Coast Ranges, San Joaquin Valley, and Kesterson National Wildlife Refuge of California (Reproduced with permission from ref 86, r Springer Verlag, 1994)

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subjected the sediments to weathering under oxidizing conditions. With resultant acid neutralized by the carbonate component of the sediments, the predominant form of selenium produced from oxidation of selenide, SeII, under alkaline conditions, is selenate, SeVIO2 4 , in preference to selenite, IV 87 , or biselenite, HSe O . Thus the soils on the Coast Ranges, SeIVO2 3 3 which are alkaline, contain signiﬁcant amounts of soluble mineral salts (e.g. Na2SO4.10H2O, Na2Mg(SO4)2.4H2O), including selenates (e.g. Na2 SeO4.10H2O, Na2Mg(SeO4)2.4H2O). The chemistry and mobility of selenite and selenate diﬀer greatly in soils. Selenite is adsorbed by speciﬁc adsorption processes (e.g. on clays and hydrous metal oxides) and to a much greater extent than selenate, which is adsorbed, like sulfate, by comparatively weak non-speciﬁc processes. Thus, in aerated alkaline soils, such as those of semi-arid regions like the west-central San Joaquin Valley, mobile SeVIO2 4 will be the dominant form and, through runoﬀ, be capable of entering the groundwaters and sub-surface drainage waters, ultimately being removed to Kesterson Reservoir. There, it is taken up by the biota, with such devastating eﬀects, and also converted to a range of selenium0 containing species (e.g. SeIVO2 3 , Se , H2Se, Se-rich protein), with much ultimately being deposited in the sediments. There, over geological time, natural diagenetic processes would presumably lead again to the formation of reduced seleno-sulﬁde species. Considerable eﬀort has been devoted to the clean-up of Kesterson Reservoir, which was drained in 1987, one year after the input of drainage water was stopped. Natural processes, stimulated by the addition of nutrients, of summertime microbial conversion and volatilization to the atmosphere of methylated forms such as dimethylselenide, (CH3)2Se, and downwards leaching of SeVIO2 4 by percolating water in winter helped to dissipate 68–88% of total selenium from the topsoil (0–15 cm) over a period of eight years.88 Phytoremediation techniques were also tested. Similarly, pilot bioreactors were set up to convert drainage water SeO2 4 into commercially useful selenium-containing products. Although Kesterson Reservoir has been closed, selenium loading in the Central Valley of California, the San Joaquin River, the San Joaquin-Sacramento Delta, and San Francisco Bay still occurs and an emerging selenium contamination issue is developing in south-eastern Idaho, USA, in streams draining areas with phosphate mining activities.89 3.3.2.5 Chromium in Groundwater. Millions of tonnes of high-pH chromite ore processing residue (COPR) were deposited (e.g. as landﬁll material) in the past in urban areas such as Glasgow, Scotland,90 and Hudson County, New Jersey, USA.91 Even now the high-lime process

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Chapter 3

responsible for the generation of this waste continues to be used in countries such as China, Russia, Kazakhstan, India, and Pakistan.92 In Glasgow, the chromium content of this residue is high at 3–4% w/w but, more signiﬁcantly, elevated concentrations as high as B1% w/w of highly toxic and carcinogenic hexavalent chromium, which takes the form of chromate (CrVIO2 4 ) at the prevailing high pH of 9–12 characteristic of high-lime COPR, have been measured in the solid phase at the contaminated sites. Furthermore, on sites in the Glasgow area where COPR from a high-lime chemical works was last deposited in the 1960s, CrVI is still leaching out in such quantity that receiving groundwaters and streams in the area are seriously contaminated at concentrations of up to 100 mg L1, about 2000 times greater than Environmental Quality Standards (EQSs).93 There are many problems involved in the in situ application of remediation-by-reduction methods, e.g. employing Fe, FeII, or organic compounds, to reduce CrVI held in speciﬁc mineral phases within COPR to the much less harmful trivalent CrIII. For example, the use of ferrous iron, Fe21, from ferrous sulfate, rather than reduce the CrVIO2 4 held within the COPR via 21 2CrO2 þ 13H2O - Cr2O3(s) þ 6Fe(OH)3(s) þ 8H1 (3.95) 4 þ 6Fe

actually increases the release of CrO2 4 because of anion exchange of from the ferrous sulfate for CrO2 held in the layered double SO2 4 4 hydroxide mineral hydrocalumite within the COPR.94 In principle, many of the problems involved in the in situ application of methods designed to reduce CrVI held within COPR to CrIII can be avoided if the emphasis is switched to treatment of the CrVI released to groundwater and surface water from COPR. Where the objective is to treat groundwater either in a conventional pump and treat system or as it migrates through a permeable reactive barrier, there is much less impediment to chemical methods using FeII or Fe0 or S2:95–97 2CrO2 (3.96) 4 þ 3H2S(aq) þ 2H2O - 2Cr(OH)3(s) þ 3S(s) þ 4OH

More generally for CrVI-contaminated groundwater, which may also arise from other industrial activities such as tanning and electroplating, both in situ98 and pump and treat methods could employ biological as well as chemical methods for the reduction of CrVI to CrIII, for example in packed bed bioreactors.99 Another generic method for the reduction of CrVI to CrIII is the addition of organic matter, for example as horse dung or molasses. Such treatment, however, whether for direct chemical reduction (cf. Equation

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137

(3.111)) or, in conjunction with other nutrients, to stimulate bacterial reduction of CrVI, could provide a means of retention of CrIII in solution in preference to either precipitation of Cr(OH)3 or reoxidation to CrVI under the prevailing alkaline conditions of high-lime COPR sites. Farmer and co-workers93 found that o3% of dissolved chromium was present as CrVI in groundwater at one Glasgow borehole site with a high dissolved organic carbon content of 300 mg L1. Here the chromium occurred predominantly as CrIII in association with organic material of high molecular weight, with an implied reduction of CrVI by humic substances, perhaps via carboxylate groups, and the formation of CrIII-humic complexes. Although the drinking water supply of people in the Glasgow area is not threatened by the CrVI-contaminated groundwater, there are instances elsewhere of such problems. For example, a massive plume of CrVI-contaminated groundwater in the Mojave Desert, with concentrations as high as 12 mg L1, was reported to be within 40 m of the Colorado River, which supplies drinking water to California, USA.100 This contamination arises from when an electric utility used CrVI to control corrosion and mould in water-cooling towers and then dumped untreated wastewater into percolation beds in the 1950s and 1960s. Current short-term emergency pumping, transport, and oﬀsite treatment of the contaminated groundwater will have to be replaced by a longterm solution close to the river, possibly featuring in situ groundwater treatment or the installation of a permeable reactive barrier in which iron would reduce CrVI to CrIII. 3.3.2.6 Aquatic Contamination by Gold Ore Extractants. The separate use of mercury and cyanide has led to the contamination of freshwater systems. 3.3.2.6.1 Mercury in the Amazon Basin. With the price of gold soaring in the late 1970s and early 1980s, there developed a modernday gold rush in South America,101 which reached its peak during 1990– 1993.102 In the relatively inaccessible Amazon area in Brazil, there were up to a million people operating on an informal basis. The technique which is used to extract the placer gold which has accumulated in many stream and alluvial deposits is based on amalgamation with mercury. Typically, high-pressure water hoses are used to dislodge alluvium, which is then taken into a large sluice by a motor suction pump. The sluices are lined with sacking into which mercury is added to amalgamate the gold particles, which lodge in the sacking. In other versions, no mercury is used in the sluices. Instead, material is collected from the

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Chapter 3

sluice lining and crushed in an oil drum, where the mercury is then introduced, followed by manual panning. The eﬀects of all this activity can be seen in the increased turbidity, with clear rivers turning a muddy brown colour, and increased ﬁsh mortality. Furthermore, releases of mercury to the aquatic environment have resulted in elevated concentrations of mercury in sediments and ﬁsh. Commonly eaten ﬁsh species from the Madeira River were found to contain mean mercury levels of 0.7 mg kg1 and up to 3.8 mg kg1 in the Tapajos Valley.103 Given the bioaccumulation and biomagniﬁcation of mercury in many ﬁsh species and the possible microbial methylation of mercury deposited in sediments to the even more toxic lipid-soluble methyl mercury, CH3Hg1, there is a real prospect of enhanced dietary uptake of the most toxic form of mercury by the local populations. This would be in addition to exposure from inhalation of mercury vapour during the handling of mercury and from subsequent post-amalgamation treatment steps, which usually involve reﬁnement of the extracted gold by heating to drive the more volatile mercury oﬀ. It is thought that there is an annual discharge of about 100 tonnes of mercury into the Amazon ecosystem and perhaps the same amount released as mercury vapour. Thus there could be an increased risk of the neurological damage and foetal deformities commonly associated with mercury poisoning, as most notably exempliﬁed in the ﬁshing village population of Minimata, Japan, in the 1950s, when a chemical manufacturing plant discharged mercury salts and CH3Hg1 into the bay, to be taken up by ﬁsh and shellﬁsh. 3.3.2.6.2 Cyanide. Another method of extracting gold that can lead to contamination of freshwater is cyanide heap leaching, which is used to extract tiny amounts of gold from huge volumes of low-grade deposits of gold-bearing rock. In this case, ore is dug from open-pit mines, crushed, and spread on asphalt or plastic pads. The heaps are then sprayed with cyanide, which dissolves the gold. 4Au(s) þ 8NaCN þ 2H2O þ O2(g) - 4NaOH þ 4NaAu(CN)2 (3.97) The gold-bearing cyanide solutions run oﬀ on impermeable asphalt or plastic to collection vats and are then treated, for example with zinc, to extract gold. 2NaAu(CN)2 þ Zn(s) - Na2Zn(CN)4 þ 2Au(s)

(3.98)

One major environmental problem with this technique is that cyanide is highly toxic. Although attempts are made to loop the cyanide through a closed system to try to ensure that none is lost, cyanide does commonly

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139

escape to surface waters and groundwater. While cyanide quickly decomposes in oxygenated, acidic surface water, it can persist at toxic levels for much longer periods in groundwater and thus pose a longerterm threat. 3.3.3 Historical Pollution Records and Perturbatory Processes in Lakes 3.3.3.1 Records – Lead in Lake Sediments. Under ideal circ*mstances, freshwater lake sediments can preserve the record of temporal variations in anthropogenic input of contaminants via atmospheric deposition, catchment runoﬀ, eﬄuent inﬂow and dumping from industrial, transportation, mining, agricultural, and waste disposal sources.104 Prerequisites are transfer of contaminants associated with settling inorganic particulates and/or biotic detritus from the water column to the sediments, no disturbance of sediments by physical mixing, slumping or bioturbation after deposition, no post-depositional degradation or mobility of the contaminants, and the establishment of a reliable time axis (e.g. via the use of the naturally-occurring radionuclide 210Pb, half-life 22.35 years, and the nuclear fallout radionuclide 137Cs, half-life 30.2 years). The derived records of lead deposition to sediments in Loch Ness and the northern and southern basins of Loch Lomond, Scotland, over the past few hundred years since the onset of the Industrial Revolution are shown in Figure 6.105 The noticeable change in the 206Pb/207Pb ratio of the lead deposited from the late 1920s onwards is attributed to carexhaust emissions of lead of lower 206Pb/207Pb ratio as a result of the use of lead ores from Broken Hill, Australia, in the manufacture of alkyllead additives for the UK market. On the basis of the published literature, there appears to be reasonable conﬁdence that, in the absence of postdepositional mixing, sediments do, in general, preserve a record of the deposition of lead to the sediments.105–107 Many such records in Western Europe and North America now show concentrations of lead declining towards the sediment surface as a consequence of the withdrawal of leaded petrol and also a reduction in other emissions.105,108 Longer-term palaeolimnological records, stretching back over several thousand years, demonstrate the inﬂuence of lead mining and smelting in Greek–Roman times and the later impact of the Mediaeval mining industry.109 3.3.3.2 Perturbatory Processes in Lake Sediments. The extent to which conditions for preservation of records are met depends upon the characteristics of the speciﬁc individual systems and contaminants under study. In reviewing the perturbation of historical pollution records in aquatic sediments, Farmer110 has suggested that in the case

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Figure 6 Fluxes (J) and associated 206Pb/207Pb ratios (K) of anthropogenic lead deposited in the sediments of Loch Ness, northern Loch Lomond and southern Loch Lomond (Scotland), derived from analysis of sections of 210Pb-dated sediment cores, vs. calendar date since the 17th century. Dates prior to approximately 1860 should be considered as extrapolations (Reprinted with permission from ref 105, r Elsevier, 2002)

of heavy metal behaviour in freshwater lakes there should ideally be investigation of diﬀerent systems of varying status (e.g. oligotrophic, eutrophic, acid),

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seasonal inﬂuences, modes of element introduction and transport, element associations in the water column, depositional and redistributional processes and rates (e.g. via radionuclide and stable isotope studies) in the sediments, concentration proﬁles and speciation in the overlying water, solid sediment, and pore water, element/solid phase associations (e.g. via sequential chemical extraction procedures), and authigenic mineral identiﬁcation in the sediments. While opinions may diﬀer as to whether particular sediment columns have been disturbed by mixing or the extent to which certain elements are vertically mobile, it is accepted that post-depositional redox-controlled cycling can aﬀect the vertical proﬁles of manganese and iron.111,112 The driving force is the microbiological decomposition of organic matter (represented as CH2O) through bacterial utilization of O2 and inorganic oxidizing agents in thermodynamically favoured sequence.113 DG0 (kJ mol1 CH2O) CH2O þ O2(g) " CO2(g) þ H2O

475

(3.99)

5CH2O þ 4NO 3 " 2N2(g) þ 4HCO3 þ CO2(g) þ 3H2O

448

(3.100)

CH2O þ 3CO2(g) þ H2O þ 2MnO2(s) " 2Mn21 þ 4HCO 3

349

(3.101)

CH2O þ 7CO2(g) þ 4Fe(OH)3(s) " 4Fe21 þ 8HCO 3 þ 3H2O

114

(3.102)

2CH2O þ SO2 4 " H2S(g) þ 2HCO3

(3.103)

2CH2O " CH4(g) þ CO2(g)

(3.104)

Upon dissolution of oxides and (hydr)oxides of manganese and iron under reducing conditions at depth, divalent cations of these elements can diﬀuse upwards through the pore waters to be oxidized and precipitated in near-surface oxic layers, leading to the characteristic near-surface enrichment of manganese and iron in the sediments of many well oxygenated lakes (see Section 3.2.4.3). Phosphorus and

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arsenic are also recycled but in their case it is the negatively charged anions that are released into solution from an association with solid phase ferric (hydr)oxides, which dissolve under suﬃciently reducing conditions. Upward diﬀusion of reduced species leads ultimately to surface enrichment via co-precipitation or adsorption of phosphate and arsenate on iron (hydr)oxides in surface layers. As a result, arsenic concentrations as high as 675 mg kg1 can be found near the surface of Loch Lomond sediments, compared with background values of 15–50 mg kg1 at depth.114 Enrichments of phosphorus and arsenic have similarly been found in association with authigenic iron oxides in sedimentary Fe/Mn layers in Lake Baikal where, however, enrichments of molybdenum were found associated with manganese oxides, perhaps through mixed oxide formation.115 For lakes which have undergone signiﬁcant acidiﬁcation, it has been suggested that heavy metals could be released from surface sections by pH-dependent dissolution, resulting in sub-surface maxima in sedimentary heavy metal concentrations. In two Canadian acid lakes, however, Carignan and Tessier116 found that downward diﬀusive ﬂuxes of dissolved zinc from overlying waters into anoxic pore waters were responsible for the pronounced sub-surface sediment maxima in solid phase zinc, presumably as the insoluble sulﬁde. Traditionally, heavy metals such as lead, zinc, copper, and cadmium have been considered diagenetically (i.e. as a consequence of chemically and biologically induced changes in prevailing sedimentary conditions, e.g. pe, anionic composition, etc.) immobile and ﬁxed in the sediment after deposition, partly as a result of the formation of comparatively insoluble sulﬁdes under reducing conditions at depth. This may well be an oversimpliﬁed view, as interactions and formation of complexes between metals and dissolved organic matter (e.g. humic substances) may maintain dissolved metal concentrations at levels greater than those predicted by simple solubility product calculations (see Sections 3.2.3.5 and 3.2.4.1).117 A combination of geochemical modelling, incorporating interactions of metals with mineral phases and naturally occurring organic matter,34,118 and the development of increasingly sophisticated analytical technology, such as DET (diﬀusive equilibration in thin ﬁlms)119 and DGT (diﬀusive gradients in thin ﬁlms) for the high-resolution determination on a submillimetre scale of metals in sedimentary pore waters,120 oﬀers considerable promise for greater understanding of redox-driven cycling of trace elements in lakes. Not only have ﬁne structure and sharp-featured maxima been observed for manganese, iron, sulﬁde, and some trace metals in pore waters, but two-dimensional structure in concentration proﬁles has demonstrated the potential importance of sedimentary microniches for metal

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remobilization.119,121 The increasing sophistication of the DGT technique is now also permitting comparisons of measured metal speciation in freshwaters with model (e.g. WHAM, ECOSAT) predictions.122,123 Finally, it should be noted that sediment cores continue to be used to investigate historical perspectives of environmental contamination by anthropogenic organic compounds, such as organochlorine pesticides,124,125 polychlorinated biphenyls,124,125 polychlorinated naphthalenes,126 polycyclic aromatic hydrocarbons,124,127,128 aliphatic hydrocarbons,128,129 polychlorinated dibenzo-p-dioxins,130 polychlorinated dibenzofurans,130 and polybrominated diphenyl ethers125,131 in environments as diverse as the Great Lakes124,129–131 and Washington128 in North America, Esthwaite Water126 in England, remote mountain lakes127 in the Pyrenees and Alps in Europe and Greenland.125 3.3.3.3 Onondaga Lake. Onondaga Lake, near Syracuse, New York State, has been called ‘one of the world’s most polluted lakes’ and provides a textbook case of the impact of industrial processes on the environment.132 Rich natural resources, especially the local brine springs, made the Syracuse region an ideal location for chemical manufacturing. Unfortunately, the exploitation of these, despite the implementation of signiﬁcant technological advances and, for the time, environmentally responsible operation of factories had a devastating impact upon the chemistry of the lake. The Solvay Process, which was introduced for the manufacture of the industrially important chemical, ‘soda ash’, i.e. sodium carbonate (Na2CO3), along the west shore of Onondaga Lake in 1884, was cheaper and less polluting than the existing Leblanc Process. Essentially, it made use of two cheap and plentiful naturally occurring substances in the area – NaCl from the deep brine springs and CaCO3 from limestone outcroppings – in a simple reaction, which yielded two useful products. CaCO3 þ 2NaCl " Na2CO3 þ CaCl2

(3.105)

In practice, this overall reaction required a series of individual reactions and chemical intermediates, one of the most important involving ammonia in a step to separate NaHCO3 from NH4Cl by fractional crystallization at 273 K. NH3(g) þ CO2(g) þ NaCl(ag) þ H2O(1) " NaHCO3(s) þ NH4Cl(aq) (3.106) The NaHCO3 was subsequently heated to yield Na2CO3 while NH3, the most expensive compound used, was regenerated from the decomposition of NH4Cl. So the very eﬃcient Solvay Process, one of the ﬁrst industrial

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processes to regenerate intermediates, maximized proﬁts and, in theory, prevented environmental harm from unwelcome discharges. Unfortunately, however, sales of CaCl2 from the Onondaga Lake plant failed to match those of Na2CO3, with the result that excess CaCl2 was allowed to be released into a tributary of the lake. In addition, substantial quantities of unmarketable salts (mainly CaCl2) from the Solvay Process were dumped daily into the lake. Furthermore, waste slurries pumped into diked beds along the lake shoreline resulted in substantial leaching of Ca21, Na1, and Cl ions into the lake by rainwater runoﬀ. The Ca21 ions reacted with CO32 ions, resulting from 2 H2O(1) þ CO2(g) " H2CO3(aq) " 2Hþ (aq) þ CO3 (aq) (3.107)

to precipitate CaCO3 via 2 Ca2þ (aq) þ CO3(aq) " CaCO3(s)

(3.108)

The eﬀect of the deposition of CaCO3, at up to 1.7 104 t per year, was to increase the sedimentation rate of the lake several-fold, with a large CaCO3 delta where the major tributary, Ninemile Creek, ﬂows into the lake, and create a layer of CaCO3 about 1 m thick over the bottom sediments. As a result, Onondaga Lake became eﬀectively a saturated solution of CaCO3, with a pH of 7.6–8.2, and, most unusually for a lake in the north-eastern USA, immune to both acid precipitation and phosphate-induced eutrophication, the latter largely being avoided as a consequence of 2 3Ca2þ (aq) þ 2PO4 (aq) " Ca3(PO4)2(s)

(3.109)

Since the closure of the Solvay Plant in 1980, the water quality has improved as salt concentrations have decreased from 3500 to 1000 mg L1. The other major industrial process which has polluted Onondaga Lake has been the production of sodium hydroxide (NaOH) and chlorine, two of the world’s most important chemicals, from sodium chloride via the electrolysis of NaCl solution and subsequent amalgamation of the sodium metal produced with the mercury cathode [i.e. Na(Hg)], prior to spraying into water to yield NaOH. With a single electrolysis cell containing up to 4 t Hg and a single plant containing dozens of cells, the escape of mercury (e.g. through leakage) at a rate of 5–10 kg per day from 1946–1970 into the lake resulted in mercurycontaminated sediments (440 mg kg1) and ﬁsh, much of the mercury in the latter being in the form of the highly toxic CH3Hg1. Despite the closure of the chlor-alkali plant in the late 1980s, mercury is still ﬂowing into the lake from tributaries, although concentrations of soluble

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mercury compounds in the lake are declining.132 Kim and co-workers133 have simulated mercury transport and speciation (HgII, Hg0, CH3Hg1) in both the water column and the benthic sediment using a Water Quality Simulation Program. Model predictions generally agreed with measured values for the water column, in which advection, sorption, and settling were important mechanisms for mercury transport. Reduction, methylation, and demethylation aﬀected mercury speciation in the water column and the benthic sediment. Natural attenuation showed no positive impact for remediation when compared with dredging and capping. 3.3.4 Nutrients in Water and Sediments 3.3.4.1 Phosphorus and Eutrophication. Eutrophication can be considered as the excessive primary production of algae and higher plants through enrichment of waters by inorganic plant nutrients, usually nitrogen and phosphorus. The latter, in the form of phosphate, is normally the limiting nutrient because the amount of biologically available phosphorus is small in relation to the quantity required for algal growth.134 Sources of nutrients can be discrete (e.g. speciﬁc sewage outfall) or diﬀuse (e.g. farmland fertilizers). Eutrophic lakes, highly productive and often turbid owing to the presence of algae, can be contrasted with oligotrophic lakes, which exhibit low productivity and are clear in summer. There have been many examples of unsightly algal blooms aﬀecting freshwater bodies throughout the world, from Lake Erie in North America to the Norfolk Broads in East Anglia, England.59 Public concern has increased along with the reported incidences of toxicity of the bloom-forming organisms, in particular the cyanobacteria (blue-green algae), which have been implicated in ﬁsh fatalities, for example in Loch Leven, Scotland.135 The environmental damage costs of freshwater eutrophication in England and Wales have been estimated at d75–114.3 M per year.136 The chemical form of phosphorus in the water column available for uptake by biota is important. The biologically available phosphorus is usually taken to be ‘soluble reactive phosphorus (orthophosphate)’, i.e. which, upon acidiﬁcation of a water sample, reacts with added molybdate to yield molybdophosphoric acid, which is then reduced with SnCl2 to the intensely-coloured molybdenum blue complex and is determined spectrophotometrically (lmax ¼ 882 nm).137 Reduction in inputs of phosphate, for example from point sources or by creating water meadows and buﬀer strips to contain diﬀuse runoﬀ, has obviously been one of the major approaches to stemming eutrophication trends and

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encouraging the restoration of aﬀected lakes. That this has not always been successful, however, can be attributed in many cases to the release/ recycling of phosphorus previously deposited to and incorporated within the bottom sediments of the lake systems in question.138 The potential mobility and bioavailability of sedimentary phosphorus are to a large extent governed by the chemical associations and interactions of phosphorus with diﬀerent sedimentary components.139 Bearing in mind the removal or transport of phosphorus to the sediments, especially important phases are likely to be ‘organic phosphorus’ from deposited, dead, decaying biota, and ‘sorbed orthophosphate’ on inorganic particulates (e.g. iron (hydr)oxides). Many sequential extraction schemes have been developed to investigate phosphorus fractionation in lake sediments. Perhaps the most sophisticated is that of Psenner et al.140 who identify, operationally deﬁne, and separate the following fractions: ‘labile, loosely bound or adsorbed’ (NH4Cl-extractable); ‘reductant-soluble, mainly from iron hydroxide surfaces’ (buﬀered dithionite-extractable); ‘adsorbed to metal oxides (e.g. Al2O3)’ (NaOH-extractable), subsequently distinguishable from ‘organic’ (also NaOH-extractable); ‘apatite-bound’ (HCl-extractable); and ‘residual’ (persulfate-digestible). Pardo and co-workers141 have reviewed such approaches in the light of the development by the European Commission of a harmonized protocol for the fractionation of sedimentary phosphorus. Many such studies of sedimentary phosphorus proﬁles, also incorporating pore water measurement of soluble reactive phosphate, have demonstrated that redox-controlled dissolution of iron (hydr)oxides under reducing conditions at depth releases orthophosphate to solution. This then diﬀuses upwards (and downwards) from the pore water maximum to be re-adsorbed or co-precipitated with oxidized FeIII in near-surface oxic sections. The downwards decrease in solid phase ‘organic’ phosphorus indicates increasing release of phosphorus from deposited organic matter with depth, some of which will become associated with hydrous iron and other metal oxides, added to the pool of mobile phosphorus in pore water or contribute to ‘soluble unreactive phosphorus’. The characteristic reactions involving inorganic phosphorus in the sediments of Toolik Lake, Alaska, are shown in Figure 7.142 If, at depth, the concentrations of Fe21 and phosphate are high enough, authigenic vivianite (Fe3(PO4)2.8H2O) may precipitate out. With redox control largely responsible for phosphorus mobility in sediments, what might the consequences of oxygen depletion in the hypolimnion be? If conditions in the surface sediments are not suﬃciently oxidizing to precipitate iron (hydr)oxides and thereby adsorb the phosphate (i.e. the redox boundary for iron may be in the overlying

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Figure 7 Characteristic reactions involving phosphorus in the sediments of Toolik Lake, Alaska. The primary processes controlling porewater phosphorus concentrations are adsorption to and desorption from iron oxyhydroxides and the precipitation of authigenic vivianite (Reprinted with permission from ref 142, r E. Schweizerbart’sche Verlagsbuchhandlung, 1987)

water column), the phosphate from previously deposited sediments would stream oﬀ into the water column and promote eutrophication. This process is called internal loading of phosphorus. Redox control may not be the only process aﬀecting release of phosphorus from sediments. During the enhanced photosynthesis of algal blooms, the pH of lake water increases as CO2 is used up and HCO3 increases. Thus, in summer, both Lough Neagh143 in N. Ireland and Lake Glanningen144 in Sweden have shown an increase in water

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phosphate concentration. It seems likely that OH is exchanging with sorbed phosphate in alkaline lakes, thus releasing phosphorus from association with iron and aluminium (hydr)oxides (see Section 3.2.4.3). If the waters are calcium-rich, however, this could have the eﬀect of precipitating phosphate as hydroxyapatite (Ca10(PO4)6(OH)2) under the high-pH conditions prevailing. There can be other factors, such as temperature, which promote phosphorus release. An increase in temperature can lead to increased bacterial activity, which increases oxygen consumption and decreases the redox potential. Turbulence may also be a factor. For example, wind-induced bottom currents in shallow lakes could destroy any pH gradients across a buﬀered sediment-water interface and re-suspend P-rich sediment in water of high pH.145 There have been many proposals for restoration of eutrophic lakes.146 For example, 34 options have been put forward for Loch Leven,147 Scotland, with the aim of reducing algal biomass and the incidence of the bloom-forming cyanobacteria in particular. These strategies fall into two categories: (i) those aimed at stemming the production of algae in the ﬁrst place, including reduction in the supplies of light and nutrients and (ii) those aimed at reducing existing algal biomass, including physical methods such as increased ﬂushing of the loch and harvesting of blooms, chemical treatment with algicides, and biological methods involving viruses, parasitic fungi and grazing protozoans, rotifers, and micro-crustaceans. Thus far, progress has largely been restricted to reducing external inputs of phosphorus from point sources to the loch, an approach which could perhaps be supplemented by the creation of water meadows and/or buﬀer strips to reduce inputs from diﬀuse runoﬀ. Elsewhere, there was a surge in wetlands construction in the 1980s, using a mixture of plants to clean water contaminated with nitrates and phosphates.148 One of the largest wetlands restoration projects is planned for the Florida Everglades, to the south and east of the rich farmlands around Lake Okeechobee, where runoﬀ waters enriched in phosphorus from fertilizers have disrupted the ﬂora and fauna. As water ﬂows through cattails and sawgrass, phosphorus concentrations are expected to decline from 170 to 50 mg L1. An intriguing attempt to use ochre (cf. Section 3.3.1.3), discharged from abandoned mines, to strip phosphate from sewage and agricultural runoﬀ and then use the phosphate-saturated ochre as a slow-release phosphate fertilizer to fertilize a range of grasses and cereal crops has been reported.149 It should be noted that the addition of ferric sulfate or chloride to mop up phosphate in eutrophic lakes59 may not work in the long term as there may be subsequent release of deposited phosphate from sediments under reducing conditions.

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A novel use of phosphate to counter acidiﬁcation has emerged in recent years.150 Traditional remedial methods which have been adopted include the direct liming of lakes, e.g. in Sweden,151 or of catchments, e.g. around Loch Fleet in SW Scotland.152 Although helpful, such approaches based on neutralization are costly and usually need to be repeated at regular intervals. Furthermore, the resulting Ca-rich waters may turn out to support biota quite diﬀerent from those found in natural softwater lakes. An alternative approach has been tried by Davison and co-workers on Seathwaite Tarn, an upland reservoir in the English Lake District.131 Phosphate fertilizer was added to stimulate primary productivity and thereby increase the assimilation of nitrate. This generates base according to the equation 106CO2(g)þ138H2Oþ16NO 3 "(CH2O)106 (NH3)16þ16OH þ138O2(g)

(3.110) As concentrations of nitrate are increasingly high in acid waters, the addition of modest amounts of phosphate may generate suﬃcient base to combat acidity without inducing excessive productivity. An increase in pH of 0.5 and a marked increase in biological productivity at all levels were observed over the three-year period of the experiment. In the longer term, additional quantities of base should be generated by the anoxic decomposition of organic material accumulating on the lake-bed (through the dissimilative reduction of inorganic oxidants such as nitrate, sulfate, or iron hydr(oxides) present in the sediments – Equations (3.100) to (3.103)). If oxygen is the electron acceptor there is no net gain of base (Equation (3.99)), there is no advantage in adding nitrate because 1 mol of nitrate is required to generate 1 mol of base, and the generated base (contributing to alkalinity) should not be confused with the temporary rise in pH associated with CO2 consumption, which aﬀects neither alkalinity nor acidity (see Section 3.2.4.1). 3.3.4.2 Nitrate in Groundwater. Principal sources of nitrate in water are runoﬀ and drainage from land treated with agricultural fertilizers and also deposition from the atmosphere as a consequence of NOx released from fossil fuel combustion. The nitrogen present in soil organic matter may also be released as nitrate through microbial action. Nitrate’s role as a nutrient contributes signiﬁcantly to blooms of algae, which upon death are decomposed ﬁrst by aerobic bacteria, thereby depriving ﬁsh and other organisms of oxygen.59,147 There has long been concern expressed over the presence of nitrate in drinking water at concentrations exceeding the EC guideline of 50 mg L1 because of the risk of methaemoglobinaemia (blue baby syndrome). Here,

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NO3 is reduced in the baby’s stomach to NO2, which, on absorption, reacts with oxyhaemoglobin to form methaemoglobin. There have also been fears over the reaction of NO2 with secondary amines from the breakdown of meat or protein to produce carcinogenic N-nitroso compounds, but there is as yet no clear evidence of a link between stomach cancer and nitrate in water.153 Nevertheless, there is growing concern over the contamination of groundwater by nitrate, for example in regions such as the Sierra Pelona Basin, California.154 There the local groundwaters, the major source of drinking water from private water wells located near each private residence in the rural communities, often exceed the USEPA maximum contaminant level for drinking water of 10 mg L1 (NO3–N). Isotopic investigations, based upon 15N/14N, have conﬁrmed the predominance of anthropogenic, organic human, and/or animal waste and decay of irrigation-enhanced vegetation rather than natural nitrate sources. In south Dorset, UK, the cause of the increase in the nitrate concentrations of the Chalk groundwater from a baseline of 1.04 mg L1 (NO3-N) to 6.37 mg L1 has been attributed to increased fertilizer use and increased livestock numbers, mostly during the ﬁnal 25 years of the 20th century.155 It is feared that nitrate concentrations in both surface and groundwaters of the Chalk in southern England may approach or exceed the EC limit of 11.3 mg L1 (NO3–N) for potable supplies in future years. Nitrate in manure and fertilizer is also now under suspicion as an endocrine disruptor (see Section. 3.3.5.2.3).156 3.3.5 Organic Matter and Organic Chemicals in Water 3.3.5.1 BOD and COD. The solubility of oxygen in water in equilibrium with the atmosphere at 251C is 8.7 mg L1. Causes of oxygen depletion include decomposition of biomass (e.g. algal blooms) and the presence of oxidizable substances (e.g. sewage, agricultural runoﬀ, factory eﬄuents) in the water. The addition of oxidizable pollutants to streams produces a typical sag in the dissolved oxygen concentration. The degree of oxygen consumption by microbially mediated oxidation of organic matter in water is called the Biochemical (or Biological) Oxygen Demand (BOD) (cf. Equation (3.99)). Another index is the Chemical Oxygen Demand (COD), which is determined by using the powerful oxidizing agent, dichromate (Cr2O72), to oxidize organic matter 1 31 2Cr2O2 þ 3CO2(g) þ 11H2O (3.111) 7 þ 3CH2O þ 16H " 4Cr

followed by back-titration of excess added dichromate with Fe21 21 Cr2O2 þ 14H1 " 2Cr31 þ 6Fe31 þ 7H2O (3.112) 7 þ 6Fe

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As Cr2O72 oxidizes substances not oxidized by O2, the COD is usually greater than the BOD and to some extent overestimates the threat posed to oxygen content. 3.3.5.2 Synthetic Organic Chemicals. A large number of organic compounds are synthesized for agricultural use, mainly as pesticides, and for industrial use as solvents, cleaners, degreasers, petroleum products, plastics manufacture, etc.157 Many organic micropollutants percolate into the soil and accumulate in aquifers or surface waters. They can contaminate drinking water sources via agricultural runoﬀ to surface waters or percolation into groundwaters, industrial spillages to surface waters and groundwaters, runoﬀ from roads and paved areas, industrial waste water eﬄuents leaching from chemically treated surfaces, domestic sewage eﬄuents, atmospheric fallout, and as leachate from industrial and domestic landﬁll sites.157 In the last few years, emerging chemical contaminants such as pharmaceuticals and personal care products have also been detected in freshwaters.158 In a widely publicized study, Kolpin and co-workers159 investigated the occurrence of pharmaceuticals as well as other organic wastewater contaminants in a national reconnaissance study of 95 contaminants in 139 streams in 30 states of the USA in 1999–2000. As many as 38 of the 95 targeted compounds were found in a single water sample, although the average number of compounds in a given sample was seven and the concentrations of individual compounds were typically much less than 1 mg L1. Three classes of compounds (detergent metabolites, plasticizers, steroids) had the highest concentrations and the most frequently detected compounds were coprostanol (faecal steroid), cholesterol (plant and animal steroid), N,N-diethyltoluamide (insect repellent), tri(2-chloroethyl)phosphate (ﬁre retardant) and 4-nonylphenol (non-ionic detergent metabolite) (Figure 8).160 Although measured concentrations were generally low and rarely exceeded drinking water guidelines or aquaticlife criteria, little is known about the potential interactive eﬀects that may occur from complex mixtures of organic wastewater contaminants in the environment. 3.3.5.2.1 Pesticides. Many pesticides in aquifers have resisted degradation and are more likely to persist there because of reduced microbial activity, absence of light, and lower temperatures.161 Numerous aquifers, for example in eastern England, have been found to exceed the Maximum Admissible Concentrations (MAC) guidelines for total pesticides in drinking water (0.5 mg L1), due primarily to the presence of herbicides of the carboxy acid and basic triazine groups. Gray157 has listed the 12 pesticides most often found in UK drinking waters in two

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Frequency of Detection (%)

5 1

2 22

11 14

0 40

Total Concentration (%)

dr c ge t re ugs nt p m elie et ab nt D is olite in s fe c Pl tan a ts Fi stic re iz e re ta rs rd An ant tib s In ioti se c R ct s ep ic ro i de O duc s th PA er tive pr H s es hor m cr o ip n tio es n An dru tio gs xi da Fr ag nts ra nc e So s lv en t

tio n

es pr on N

St er

oi ds

Figure 8 Organic wastewater contaminants by general use category. Emerging organic contaminants in US streams, as reported by the US Geological Survey 159,160, can be broken down into 15 categories. A shows the frequency of detection, and B shows the fraction of the total measured concentration (%). The number of compounds in each category is shown above the bars in A (Reprinted with permission from ref 159, r American Chemical Society, 2002)

groups, frequently occurring (atrazine, chlortoluon, isoproturon, MCPA, mecoprop, simazine) and commonly occurring (2,4-D, dicamba, dichlorprop, dimethoate, linuron, 2,4,5-T). A major seasonal source of insecticide in freshwaters is sheep dipping. With organophosphorus compounds, which themselves replaced the more persistent organochlorine insecticides such as lindane, now falling out of favour because of health risks to users, the use of synthetic pyrethroids as alternatives

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has resulted in the death of aquatic organisms in UK rivers. Pyrethroids, which do not persist in the environment and are largely non-toxic to mammals, are toxic to invertebrates and are estimated to be at least 100 times more toxic in the aquatic environment than organophosphorus pesticides. It appears to be the pouring of waste dip into holes in the ground that has caused the problem.162 In order to protect the aquatic environment, the Environment Agency in England assesses water quality against EQSs. Deﬁned as the concentration of a substance which must not be exceeded within the aquatic environment in order to protect it for its recognized uses, EQSs are speciﬁc to individual substances, including pesticides. Atrazine, now banned in many European countries but which still accounts for 40% of all herbicide applied in the USA and is most often applied in the spring, is routinely present in streams, rivers, and reservoirs at levels of 1–10 mg L1, close to the USEPA standard of 3 mg L1.163,164 Runoﬀ peaks of 100–200 mg L1, however, have been recorded at the times when frogs are breeding and tadpoles are developing. Atrazine is therefore suspected of causing the gonadal abnormalities, such as retarded development and hermaphroditism, found in male wild leopard frogs in diﬀerent regions of the USA.163 (cf. Section 3.3.5.2.3) 3.3.5.2.2 Polychlorinated biphenyls (PCBs). The Great Lakes, the largest body of freshwater in the world, with long hydraulic residence times, long food chains, and multiple sources of PCBs, have been the focal point of PCB research in aquatic systems. The distribution of PCBs between the dissolved and particulate phases is dependent on the concentration of suspended particulate matter, dissolved and particulate organic carbon concentrations, and the extent to which the system is at equilibrium. In the 1990s, water concentrations of up to 0.6 ng L1 were observed in the most contaminated lakes (Michigan, Erie, and Ontario), which may be compared with the USEPA Great Lakes Water Quality Guidance criteria of 0.017 ng L1. Sediment concentrations peaked in the 1970s, the period of maximum PCB production in the USA. Following a ban upon their North American production during the 1970s, there was a signiﬁcant decline in the PCB concentrations of Great Lakes’ ﬁsh from the mid-1970s to the mid-1980s but the rate of decrease has since slowed or stopped for all lakes.165 In common with other organochlorine compounds, it is the stability, persistence, volatility, and lipophilicity of PCBs, which lead to considerable biomagniﬁcation along the food chain, often far from the place of release. To these traditional halogenated contaminants can now be added ﬁre-retardant polybrominated diphenyl

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ethers (PBDEs),166 and perﬂuorooctane sulfonates (PFOS),167 which have been found in Great Lakes’ water and biota. 3.3.5.2.3 Endocrine disruptors. Concern has recently been expressed over the possible role of various synthetic organic chemicals (e.g. organochlorine pesticides, PCBs, phthalates, alkylphenolethoxylates (APEs), alkylphenols, etc.) as disruptors of endocrine systems of wildlife and perhaps even of humans.168,169 Reproductive changes in male alligators from Lake Apopka, Florida, embryonic death, deformities, and abnormal nesting behaviour in ﬁsh-eating birds in the Great Lakes region and the occurrence of hermaphroditic ﬁsh near sewage outfalls on some British rivers have been attributed to postulated oestrogenic or anti-androgenic eﬀects of some of these chemicals.168,169 Most APEs, which are used as detergents, emulsiﬁers, wetting agents, and dispersing agents, enter the aquatic environment after disposal in wastewater. During biodegradation treatment of the latter, the APEs are transformed into more toxic, short-chain ethoxylates, alkylphenol carboxylic acids, and alkylphenols.170 The threshold concentration of nonylphenol in water for production of the female egg yolk protein, vitellogenin, in male rainbow trout is B10 mg L1.171 It also appears that nonylphenol concentrations in many European rivers are up to 10 times higher than those found in USA rivers, which are typically o1 mg L1,170,172 the EU limit. It has been suggested, however, that natural oestrogens (e.g. oestrone and 17b-oestradiol) may be responsible for the observed eﬀects, perhaps after conversion of inactive excreted metabolites to active forms by bacterial enzyme action during sewage treatment.173 The COMPREHEND (Community Programme of Research on Environmental Hormones and Disruptors) 1999–2001 survey found that in some domestic sewage eﬄuents there were oestrogenic steroids present, almost certainly from human excretion, at concentrations suﬃcient to account for most of the oestrogenicity, causing widespread sexual disruption, in ﬁsh in rivers throughout England and Wales.174 Maximum measured concentrations were 14 ng L1 for oestradiol, 51 ng L1 for its principal metabolite oestrone, 17 ng L1 for another metabolite oestriol, and 2 ng L1 for the synthetic steroid ethinylestradiol, the active ingredient in the birth control pill. As a result, the Environment Agency has set a threshold exposure limit for steroid estrogens of 1 ng L1 for oestradiol, 3 ng L1 for oestriol and 0.1 ng L1 for ethinylestradiol in sewage eﬄuent.175 In industrial eﬄuents, the oestrogenic steroids were not detected but nonylphenol (NP), NP mono- and diethoxylates, and bisphenol-A were found at concentrations up to 3, 7, and 1 mg L1, respectively.174 The relatively high hydrophobicity of alkylphenols has

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led to their accumulation in sludge at activated sludge treatment works.176 3.3.5.2.4 Pharmaceuticals and personal care products (PPCPs). In Europe, B3000 diﬀerent substances are used in medicines such as painkillers, antibiotics, contraceptives, beta-blockers, lipid regulators, tranquilizers, and impotence drugs, and there are also thousands of personal care products (e.g. sunscreens, cosmetics, shampoos, etc.).177 Municipal wastewater is one of the main entry routes of pharmaceuticals into the environment. Ternes158 began monitoring for pharmaceuticals in Germany in the mid-1990s. Since then, in Switzerland, the detection of macrolide antibiotics, among the most important anti-bacterial agents used, in wastewater treatment plants and in the Glatt Valley Watershed has demonstrated that they are not eliminated completely in the treatment plants and that residual amounts reach receiving waters.178 Also in Switzerland, chemicals from personal care products containing UV ﬁlters, such as benzophenone-3, 4-methylbenzylidene camphor (4-MBC), and ethyl-hexyl methoxy cinnamate (EHMC), have been found in untreated wastewater (maximum concentration 19 mg L1, EHMC), treated wastewater (max. 2.7 mg L1, 4-MBC) and, at much lower concentrations, lakes and rivers (o2–35 ng L1).179 In the UK, Ashton and co-workers180 investigated the environmental transport of 12 human pharmaceuticals to streams and found that while 10 were detected in sewage treatment works eﬄuent samples (e.g. ibuprofen in 84% of samples, median concentration 3086 ng L1), there were fewer and at lower concentrations in streams (e.g. ibuprofen in 69% of samples, median concentration 826 ng L1). In the USA, where similar studies have been carried out in treated waters and streams,159,181–183 Stackelberg and co-workers184 have demonstrated that many of the PPCPs can survive conventional drinking-water-treatment and occur in potable water supplies. Although concentrations in ﬁnished water were low (93%o0.5mg L1) and the maximum possible lifetime intake of, for example, carbamazepine would be only 13 mg, compared with a single therapeutic dose of B100 mg, little is known about potential health eﬀects associated with chronic ingestion via drinking water. It is not yet known whether the relatively low environmental concentrations found for PPCPs produce adverse eﬀects on aquatic and terrestrial biota or whether the toxicity of complex mixtures might be totally diﬀerent from that of individual compounds.177 Consideration should also be given to metabolites excreted by humans and to biological degradates at treatment plants and elsewhere in the environment.184,185 Many pharmaceuticals are conjugated with glucuronic acid or sulfate

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prior to excretion, but these may be cleaved in wastewater treatment plants, releasing active pharmaceuticals.177 Similarly, the disposal of contaminated sludge on land as a fertilizer should be viewed with concern.177 The use of veterinary medicines may also cause environmental risks.186 3.3.5.2.5 Industrial solvents. Six are widely used in the UK – dichloromethane (DCM), trichloromethane, 1,1,1-trichloroethane (TCA), tetrachloromethane trichloroethene (TCE), tetrachloroethene or perchloroethylene (PCE).157 Although there has been a steady decrease in use over the past 20 years, four (TCE, PCE, TCA, and DCM) are frequently found in drinking water. The major pollution threat is when solvents are discharged directly into or onto the ground, due to illegal disposal or accidental spillage. Since groundwater is not exposed directly to the atmosphere, the solvents are not able to escape by evaporation, so it is the aquifers that are most at risk from these chemicals. Pollution of groundwater by industrial solvents is a very widespread problem (e.g. Netherlands, Italy, USA, UK187), with aquifers underlying urbanized areas such as Milan, Birmingham, London, or New Jersey containing high concentrations of all solvents. The stability of TCE in particular and the inability of these solvents to readily evaporate mean that such contamination will last for many decades.157 A typical example of groundwater contaminated by synthetic organic chemicals from jet fuel and degreasing solvents (e.g. halogenated methanes, ethanes, and ethenes, including trichloroethylene, perchloroethylene, and vinyl chloride) can be found at Otis Air Force Base on Cape Cod, Massachusetts, USA, where the Cape’s sole-source aquifer has been aﬀected. The groundwater plumes, some in excess of 5 km in length, moving at B0.5 m per day, contaminate 30 million litres of the Cape’s drinking water every day. Reactive wall (Ni–Fe) technology is being employed to try to reduce concentrations of 5–150 mg L1 TCE and tetrachloroethylene below the local drinking water limits of 0.5 mg L1.188 The chlorinated ethenes are degraded in the presence of zerovalent iron by an abiotic reduction process. 3.3.5.2.6 MTBE. Methyl tert-butyl ether (MTBE), an oxygen-containing compound used as a fuel additive since the 1970s, moves quickly through soil, is highly soluble and does not biodegrade easily. It has been found in shallow groundwater in Denver, New England,189 and elsewhere in the USA, most notably in Santa Monica, California, where a 1996 spill of MTBE-containing gasoline resulted in the contamination of wells at concentrations as high as 610 mg L1, well above the state’s advisory limit of 35 mg L1.190 In a preliminary assessment of the

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occurrence and possible sources of MTBE in groundwater in the United States in 1993/1994, as part of the US Geological Survey’s National Water-Quality Assessment programme, Squillace and co-workers191 found that, out of 60 volatile organic chemicals determined, MTBE was the second most frequently detected chemical in samples of shallow ambient groundwater from urban areas. MTBE was detected above 0.2 mg L1 in 17% of shallow groundwater samples from urban areas, 1.3% of shallow groundwater from agricultural areas, and 1.0% of deeper groundwater samples from major aquifers. Only 3% of the shallow wells sampled in urban areas had MTBE concentrations 420 mg L1, the estimated lower limit of the USEPA draft drinking water health advisory. The 1999 Report of the Blue Ribbons Panel on oxygenates in gasoline found that 5–10% of community drinking water supplies in high-MTBE-use areas showed at least detectable concentrations of MTBE, with 1% 420 mg L1.192 As MTBE generally was not found, in shallow urban groundwater, along with benzene, toluene, ethylbenzene, or xylene, which are commonly associated with petrol spills, it was concluded that possible sources of MTBE in groundwater include leaking storage tanks and non-point sources such as recharge of precipitation and storm water runoﬀ. In the summer of 1999, 13 municipal wells on the south shore of Lake Tahoe, located close to commercial areas with gasoline stations, were shut down as a result of actual or threatened MTBE contamination.193 In Meyers, an unknown amount of gasoline leaked from a pipe into the soil for about two weeks before the release was repaired and ﬁve months later a groundwater sample collected at the gasoline station was 28,000 mg L1 in MTBE.193 As MTBE is thought to be potentially carcinogenic to humans, its use poses an interesting dilemma for regulators, given that it helps to reduce carbon monoxide emissions from cars. Regardless of what happens to MTBE use in the future, it is clear that signiﬁcant amounts of this compound are already present in the sub-surface.192 In Europe, however, the problem seems much less severe.194,195 QUESTIONS AND PROBLEMS (i) Calculate the ionic strength of a solution which is 0.01 mol L1 with respect to potassium sulfate and 0.002 mol L1 with respect to magnesium chloride. Use this value of ionic strength and the Gu¨ntelberg approximation to determine the mixed acidity constant, K 0 , for methanoic acid (K ¼ 103.75). (ii) Construct a plot to show how log{activity} varies with pH for all species in one litre of groundwater containing 3 104 moles

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of 4-chlorophenol (K 0 ¼ 109.18). Annotate the plot to show the equilibrium pH of the groundwater. (iii) Calculate the equilibrium pH of rainwater in equilibrium with CO2(g) at a partial pressure of 0.00035 atm (hint: the charge balance equation can be approximated by {H1} ¼ {HCO3}).

CO2ðgÞ þ H2 O Ð H2 CO3 H2 CO3 Ð Hþ þ HCO 3 2 þ HCO 3 Ð H þ CO3

log Kð288 KÞ 1:34 6:35 10:33

(iv) Construct a pe–pH diagram for selenium species present in natural waters. Assume that the total dissolved concentration of selenium is 106 mol L1. From the diagram, predict the major selenium species in (i) oxic and acidic waters (peB10, pH 4–7) (ii) oxic and alkaline waters (peB10, pH 47) (iii) anoxic groundwaters (peB3, pH 4–8). For each set of conditions, outline the implications for sorption onto solid phases.

þ SeO2 4 þ 4H þ 2e Ð H2 SeO3 þ H2 O 2 þ SeO4 þ 3H þ 2e Ð HSeO 3 þ H2 O 2 þ þ 2H þ 2e Ð SeO SeO2 4 3 þ H2 O þ þ H H2 SeO3 Ð HSeO 3 2 þ HSeO 3 Ð SeO3 þ H H2 SeO3 þ 4Hþ þ 4e Ð SeðsÞ þ 3H2 O þ HSeO 3 þ 5H þ 4e Ð SeðsÞ þ 3H2 O 2 þ SeO3 þ 6H þ 4e Ð SeðsÞ þ 3H2 O SeðsÞ þ 2Hþ þ 2e Ð H2 SeðaqÞ SeðsÞ þ Hþ þ 2e Ð HSe H2 Se Ð HSe þ Hþ

log K(298 K) 41.40 38.76 30.26 2.75 8.50 48.58 51.33 59.83 3.52 7.32 3.80

(v) Draw, describe, and compare the concentration proﬁles of the chemical species of manganese, iron, arsenic, and phosphorus likely to be found in the solid and solution phases of unmixed sediments at the bottom of a seasonally anoxic, eutrophic freshwater lake during (i) summer and (ii) winter.

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(vi) With respect to human health and ecological impact, discuss the short- and long-term consequences of the products of the synthetic organic chemical industry upon the quality of surface waters and groundwaters. (vii) Explain the chemistry underlying (i) the major detrimental eﬀects of metal mining upon rivers and groundwaters and (ii) associated preventive or remedial measures. (viii) Critically discuss the eﬀectiveness of measures to counter the eﬀects of (i) acidiﬁcation and (ii) eutrophication in freshwater lakes. FURTHER READING E.K. Berner and R.A. Berner, Global Environment: Water, Air and Geochemical Cycles, Prentice-Hall, Englewood Cliﬀs, NJ, 1996. P.J. Boon and D.L. Howell (eds), Freshwater Quality: Deﬁning the Indeﬁnable?, The Stationery Oﬃce, Edinburgh, 1997. R. Clarke and J. King, The Atlas of Water, Earthscan, London, 2004. R.A. Downing, Groundwater – Our Hidden Asset, British Geological Survey, NERC, Keyworth, 1998. J.I. Drever, The Geochemistry of Natural Waters, 3rd edn, Prentice-Hall, Englewood Cliﬀs, NJ, 1997. M.E. Essington, Soil and Water Chemistry: An Integrative Approach, CRC Press, Boca Raton, FL, 2004. A.J. Horne, Limnology, 2nd edn, McGraw-Hill, New York, 1994. D. Langmuir, Aqueous Environmental Geochemistry, Prentice-Hall, Englewood Cliﬀs, NJ, 1997. P.S. Maitland, P.J. Boon and D.S. McLusky (eds), The Fresh Waters of Scotland, Wiley, Chichester, 1994. C.F. Mason, Biology of Freshwater Pollution, 3rd edn, Longman, Harlow, 1996. J. Mather, D. Banks, S. Dumpleton and M. Fermor, Groundwater Contaminants and their Migration (eds), Geological Society, London, Special Publication No.128, 1998. F.M. Morel and J.G. Hering, Principles and Applications of Aquatic Chemistry, Wiley, New York, 1993. J. Pankow, Aquatic Chemistry Concepts, Lewis, Michigan, 1991. REFERENCES 1. J. Rees, Sci. Public Aﬀ., 1997, (Winter), 20. 2. R. Clarke and J. King, The Atlas of Water, Earthscan, London, 2004.

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167. B. Boulanger, J. Vargo, J.L. Schnoor and K.C. Hornbuckle, Environ. Sci. Technol., 2004, 38, 4064. 168. T. Colborn, J.P. Meyers and D. Dumanoski, Our Stolen Future, Little, Brown and Company, Boston, 1996. 169. R.E. Hester and R.M. Harrison (eds), Issues in Environmental Science and Technology, vol 12, Royal Society of Cambridge, Cambridge, 1999. 170. R. Renner, Environ. Sci. Technol., 1997, 31, 316A. 171. S. Jobling, D. Sheahan, J.A. Osborne, P. Matthiessen and J.P. Sumpter, Environ. Toxicol. Chem., 1996, 15, 194. 172. C.P. Rice, I. Schmitz-Afonso, J.E. Loyo-Rosales, E. Link, R. Thoma, L. Fay, D. Altfater and M.J. Camp, Environ. Sci. Technol., 2003, 37, 3747. 173. J. Kaiser, Science, 1996, 274, 1837. 174. A.D. Pickering and J.P. Sumpter, Environ. Sci. Technol., 2003, 37, 331A. 175. M. Burke, Environ. Sci. Technol., 2004, 38, 362A. 176. A.C. Johnston and J.P. Sumpter, Environ. Sci. Technol., 2001, 35, 4697. 177. T.A. Ternes, A. Joss and H. Siegrist, Environ. Sci. Technol., 2004, 38, 392A. 178. C.S. McArdell, E. Molnar, M.J.-F. Suter and W. Giger, Environ. Sci. Technol., 2003, 37, 5479. 179. M.E. Balmer, H.-R. Buser, M.D. Mu¨ller and T. Poiger, Environ. Sci. Technol., 2005, 39, 953. 180. D. Ashton, M. Hilton and K.V. Thomas, Sci. Total Environ., 2004, 333, 167. 181. D.W. Kolpin, M. Skopec, M.T. Meyer, E.T. Furlong and S.D. Zaugg, Sci. Total Environ., 2004, 328, 119. 182. G.R. Boyd, H. Reemtsma, D.A. Grimm and S. Mitra, Sci. Total Environ., 2003, 311, 135. 183. G.R. Boyd, J.M. Palmeri, S. Zhang and D.A. Grimm, Sci. Total Environ., 2004, 333, 137. 184. P.E. Stackelberg, E.T. Furlong, M.T. Meyer, S.D. Zaugg, A.K. Henderson and D.B. Reissman, Sci. Total Environ., 2004, 329, 99. 185. A.B.A. Boxall, C.J. Sinclair, K. Fenner, D. Kolpin and S.J. Maund, Environ. Sci. Technol., 2004, 38, 368A. 186. A.B.A. Boxall, D.W. Kolpin, B. Halling-Sørensen and J. Tolls, Environ. Sci. Technol., 2003, 37, 286A. 187. M.O. Rivett, D.N. Lerner and J.W. Lloyd, Water Environ. Management, 1990, 4, 242. 188. E.L. Appleton, Environ. Sci. Technol., 1996, 30, 536A.

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189. J.D. Ayotte, D.M. Argue and F.J. McGarry, Environ. Sci. Technol., 2005, 39, 9. 190. M. Cooney, Environ. Sci. Technol., 1997, 31, 269A. 191. P.J. Squillace, J.S. Zogorski, W.G. Wilber and C.V. Price, Environ. Sci. Technol., 1996, 30, 1721. 192. R. Johnson, J. Pankow, D. Bender, C. Price and J. Zogorski, Environ. Sci. Technol., 2000, 34, 210A. 193. L.S. Dernbach, Environ. Sci. Technol., 2000, 34, 516A. 194. C. Achten, A. Kolb, W. Pu¨ttmann, P. Seel and R. Gihr, Environ. Sci. Technol., 2002, 36, 3652. 195. C. Achten, A. Kolb and W. Pu¨ttmann, Environ. Sci. Technol., 2002, 36, 3662.

CHAPTER 4

Chemistry of the Oceans STEPHEN J. DE MORA International Atomic Energy Agency, Marine Environment Laboratory, 4 quai Antoine 1er, MC 98000, Principality of Monaco

4.1 INTRODUCTION The World Ocean is a complex mixture containing all the elements, albeit in dilute amounts in some cases. Seawater contains dissolved gases and, apart from some exceptional environments, is consequently both well oxygenated and buﬀered at a pH of about 8. There are electrolytic salts, the ionic strength of seawater being approximately 0.7, and many organic compounds in solution. At the same time, there is a wide range of inorganic and organic particles in suspension. However, these comfortable distinctions become quite confused in seawater because some molecules present in true solution are sufﬁciently large to be retained by a ﬁlter. Moreover, surface adsorption allows particles to scavenge-dissolved elements and accumulate coatings of organic material from solution. Some elements, particularly those with biochemical functions, may be rapidly removed from solution. Concurrently, reactions involving geological time scales are proceeding slowly. Yet despite this apparent complexity, many aspects of the composition of seawater and chemical oceanography can now be explained with recourse to the fundamental principles of chemistry. This chapter serves to bridge the gap between those with environmental expertise and those with a traditional chemical background. 4.1.1 The Ocean as a Biogeochemical Environment A traditional approach utilised in geochemistry, and now also in environmental chemistry, is to consider the system under investigation a reservoir. For a given component, the reservoir has sources (inputs) and sinks (outputs). The system is said to be at equilibrium, or operating 170

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under steady-state conditions, when a mass balance between inputs and outputs is achieved. An imbalance could signify that an important source or sink has been ignored. Alternatively, the system may be perturbed, possibly mediated through anthropogenic activities, and therefore be changing towards a new equilibrium state. Processes within the reservoir that affect the temporal and spatial distribution of a given component are transportation and transformations. Both physics and biology within the system play a role. Clearly, transport effects are dominated by the hydrodynamic regime. Although transformations involve both chemical (dissolution, redox reactions, speciation changes) and geological (sedimentation) processes, biological activity generally controls nutrient and trace metal distributions. Furthermore, the biota inﬂuences concentrations of O2 and CO2, which in turn determine the pH and pe (i.e., the redox potential), respectively. For these reasons, some fundamental aspects of descriptive physical and biological oceanography are included in this chapter. In terms of biogeochemical cycling, the ocean constitutes a large reservoir. The surface area is 361.11 106 km2, encompassing nearly 71% of the earth’s surface. The average depth is 3.7 km, but depths in the submarine trenches can exceed 10 km. The ocean contains about 97% of the water in the global hydrological cycle. A schematic representation of the oceanic reservoir is presented in Figure 1. The material within it can be operationally deﬁned, usually based on ﬁltration, as dissolved or particulate. For simplicity here, the ocean is divided into two layers, with distinct surface and deepwaters. The boundary regions are also distinguished, as the composition in these regions can be quite different to that of bulk seawater. Furthermore, interactions within these environments can alter the mass transfers across the boundary. The rationale for such features will be presented in subsequent sections. Material supplied to the ocean originates from the atmosphere, rivers, glaciers and hydrothermal waters. The relative importance of these pathways depends upon the component considered and geographic location. River runoﬀ commonly constitutes the most important source. Transported material may be either dissolved or particulate, but discharges are into surface waters and conﬁned to coastal regions. Hydrothermal waters are released from vents on the seaﬂoor. Such hydrothermal waters are formed when seawater circulates into the ﬁssured rock matrix, and under conditions of elevated temperature and pressure, compositional changes in the aqueous phase occur due to seawater – rock interactions. This is an important source of some elements, such as Li, Rb and Mn. The atmosphere supplies particulate material globally to the surface of the ocean. In recent years, this has been the most prominent pathway to the World

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Figure 1 A schematic representation of the ocean reservoir. The source and sink ﬂuxes are designated as g and n, referring to gross and net ﬂuxes, thereby indicating that interactions within the boundary regions can modify the mass transfer. Within seawater, the p " d term signiﬁes that substances can undergo particulate– dissolved interactions. However, it must be appreciated that several transportation and transformation processes might be operative (Adapted from Chester, 1990.1)

Ocean for Pb, identiﬁed by its isotopic signature as originating from petrol additives. Wind-borne transport is greatest in low latitudes and the Sahara Desert is known to act as an important source of dust. Also, the airborne ﬂux of nutrients, notably nitrogenous compounds, has become increasingly recognised as important both in some coastal waters, such as Chesapeake Bay, and large seas, including the Mediterranean Sea. In contrast, glacial activity makes little impact on the World Ocean. Glacier-derived material tends to be comprised of physically weathered rock residue, which is relatively insoluble. In addition, the input is largely conﬁned to Polar Regions, with Antarctica responsible for approximately 90% of the material. Although volatilisation and subsequent evasion to the atmosphere can be important for elements, such as Se and Hg that undergo biomethylation,

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sedimentation acts as the major removal mechanism. The mechanism is a geological/physical process, essentially the deposition of suspended particles in response to changes in the hydrodynamic regime. Such removal is especially important in coastal environments. Marine organisms play an important role in the open sea through the sinking of shells of microorganisms and faecal pellets. The chemical precipitation of salts from seawater occurs only under special conditions. CaCO3 can precipitate in warm tropical lagoons and Mn21 can be oxidised to the relatively insoluble MnO2, a process of importance near deep-sea hydrothermal vents. Some definitions help in the interpretation of chemical phenomena in the ocean. Conservative behaviour signiﬁes that the concentration of a constituent or absolute magnitude of a property varies only due to mixing processes. Components or parameters that behave in this manner can be used as conservative indices of mixing. Examples are salinity and potential temperature, the definitions for which are presented in subsequent sections. In contrast, non-conservative behaviour indicates that the concentration of a constituent may vary as a result of biological or chemical processes. Examples of parameters that behave non-conservatively are dissolved oxygen and pH. Residence time, t, is deﬁned as t¼

A ðdA=dT Þ

ð4:1Þ

where A is the total amount of constituent A in the reservoir and dA/dT can be either in rate of supply or the rate of removal of A. This represents the average lifetime of the component in the system and is, in effect, a reciprocal rate constant (his concept is discussed further in Chapter 7). The photic zone refers to the upper surface of the ocean in which photosynthesis can occur. This is typically taken to be the depth at which sunlight radiation has declined to 1% of the magnitude at the surface. This might typically be 4100 m for visible light or photosynthetically active radiation (PAR), but generally o20 m for UV wavelengths, of recent interest due to the enhanced input invoked by stratospheric ozone depletion. 4.1.2 Properties of Water and Seawater Water is a unique substance, with unusual attributes because of its structure. The molecule consists of a central oxygen atom with two attached hydrogen atoms forming a bond angle of about 1051. As oxygen is more

174

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electronegative than hydrogen, it attracts the shared electrons to a greater extent. In addition, the oxygen atom has a pair of lone orbitals. The overall effect produces a molecule with a very strong dipole moment, having distinct negative (O) and positive (H) ends. While there are several important consequences, two will be considered here. Firstly, the positive H atoms of one molecule are attracted towards the negative O atom in adjacent molecules giving rise to hydrogen bonding. This has important implications with respect to a number of physical properties, especially those relating to thermal characteristics. Secondly, the large dipole moment ensures that water is a very polar solvent. Considering ﬁrstly the physical properties, water has much higher freezing and melting points than would be expected for a molecule of molecular weight 18. Water has high latent heats of evaporation and fusion. Therefore, considerable energy is required to cause phase changes, the energy being utilised in hydrogen-bond rupturing. Moreover, it has a high speciﬁc heat and is a good conductor of heat. Consequently, heat transfer in water by advection and conduction gives rise to uniform temperatures. The density of pure water exhibits anomalous behaviour. In ice, O atoms have 4 H atoms orientated about them in a tetrahedral conﬁguration. These units are packed together with a hexagonal symmetry. At the freezing point, 0 1C, ice is less dense than water. Heating breaks some hydrogen bonds and the molecules can achieve slightly closer packing causing the density to increase. Thus, the maximum density occurs at 4 1C because thermal expansion at higher temperatures compensates for this compression effect. As will be discussed later, seawater differs in this respect. Thus, fresh ice ﬂoats on water, which in part explains how rivers and lakes can freeze over, while remaining liquid at depth. With respect to other properties, water has a high surface tension that is manifest in stable droplet formation and has a relatively low molecular viscosity and therefore is quite a mobile ﬂuid. Water is an excellent solvent. It is extremely polar and can dissolve a wider range of solutes and in greater amounts than any other substance. Water has a very high dielectric constant, a measure of the solvent’s ability to keep apart oppositely charged ions. The solvating characteristics of individual ions inﬂuence their behaviour in solution, i.e., in terms of hydration, hydrolysis and precipitation. Although water exhibits amphoteric behaviour, electrolytic dissociation is quite small. Furthermore, dissociation gives equal ion concentrations of both H3O1 and OH and so pure water is neutral. The amphoteric behaviour enhances dissolution of introduced particulate matter through surface hydrolysis reactions. While the concept will be considered in detail below, the term salinity (S%) is introduced here as a measure of the salt content of seawater, a

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typical value for oceanic waters being 35 g kg1. In an oceanographic context, the most important consequence of the addition of salt to water is the effect on density. However, many of the characteristics outlined above are also altered. The addition of electrolytes can cause a small increase in the surface tension. This effect is not commonly observed in seawater due to the presence of surfactants, which decrease the surface tension and so facilitate foam formation. As illustrated in Figure 2, the presence of salt does depress the temperature of maximum density and the freezing point of the solution relative to pure water. Thus, seawater with a typical salt content of 35 g kg1 freezes at approximately 1.9 1C and the resulting ice is denser than the solution. However, more often than the formation of sea ice itself, the freezing process tends to produce fresh ice overlying a more concentrated brine solution. Salts can be precipitated at much lower temperatures, i.e., mirabilite (Na2SO4.2H2O) at 8.2 1C and halite (NaCl) at 23 1C. Some brine inclusions and salt crystals can become incorporated into the ice. From an oceanographic perspective, the fundamental properties of seawater are temperature, salinity and pressure (i.e., depth dependent). Together, these parameters control the density of the water, which in turn determines the buoyancy of the water and pressure gradients. Small density differences integrated over oceanic scales cause considerable pressure gradients and result in currents.

Figure 2

The temperature of maximum density (—) and freezing point (- -) of seawater as a function of dissolved salt content (Adapted from Tchernia, 1980.2)

176

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Surface water temperatures are extremely variable, obviously inﬂuenced by location and season. The minimum temperature found in polar latitudes approaches the freezing point of nearly 2 1C. Equatorial oceanic waters can reach 30 1C. Temperature variations with depth are far from consistent. In a region where mixing prevails, as observed especially in the surface waters, a layer forms with a relatively uniform temperature. The zone immediately beneath normally exhibits a sharp change in temperature, known as the thermocline. The thermocline in the ocean extends down to about 1000 m within equatorial and temperate latitudes. It acts as an important boundary in the ocean, separating the surface and deep layers and limiting mixing between these two reservoirs. Below the thermocline, the temperature changes only little with depth. The temperature is a non-conservative property of seawater because adiabatic compression causes a slight increase in the in situ temperature measured at depth. For instance in the Mindanao Trench in the Paciﬁc Ocean, the temperature at 8500 and 10,000 m is 2.23 and 2.48 1C, respectively. The term potential temperature is deﬁned to be the temperature that the water parcel would have if raised adiabatically to the ocean surface. For the examples above, the potential temperatures are 1.22 and 1.16 1C, respectively. Potential temperature of seawater is a conservative index. Salinity in the surface waters in the open ocean range between 33 and 37 (Figure 3), the main control being the balance between evaporation and precipitation. The highest salinities occur in regional seas, where the evaporation rate is extremely high, namely the Mediterranean Sea (38– 39) and the Red Sea (40–41). Within the World Ocean, the salinity is greatest in latitudes of about 201 where the evaporation exceeds precipitation. Lower salinities occur poleward as evaporation diminishes and near the equator where precipitation is very high. Local effects can be important, as evident in the vicinity of large riverine discharges that dilute the salinity. Salinity variations with depth are related to the origin of the deepwaters and so will be considered in the section on oceanic circulation. A zone in which the salinity exhibits a marked gradient is known as a halocline. Whereas the density of pure water is 1.000 g ml1, the density of seawater (S% ¼ 35) is about 1.03 g ml1. The term ‘sigma-tee’, st, is used to denote the density (actually the speciﬁc gravity and hence it is a dimensionless number) of water at atmospheric pressure based on temperature and salinity in situ. Density increases, and so the buoyancy decreases, with an increase in st. It is deﬁned as st ¼ (speciﬁc gravityS%,T 1) 1000

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Figure 3 The distribution of mean annual salinity in the surface waters of the ocean (Adapted from The Open University, 1989.3)

In a plot of temperature against salinity (a T–S diagram), constant st appear as curved lines, which denote waters of constant pressure and are known as isopycnals. A zone in which the pressure changes greatly is known as a pycnocline. Within the water column, a pycnocline, therefore, separates water with distinctive temperature and salinity characteristics, usually indicative of different origins. A T–S diagram can also be used to estimate the properties resulting from the mixing of two water masses. As noted above, the temperature is not a conservative property, and therefore st is also non-conservative. To circumvent the associated difﬁculties of interpretation, an analogous term known as the potential density, sy, is deﬁned on the basis of potential temperature instead of in situ temperature. The sy is therefore a conservative index. 4.1.3 Salinity Concepts Salinity is a measure of the salt content of seawater. Developments in analytical chemistry have led to an historical evolution of the salinity concept. Intrinsically, it would seem to be a relatively straightforward task to measure. This is true for imprecise determinations that can be quickly performed using a hand-held refractometer. The salinity affects seawater density and thus, the impetus for high precision in salinity measurements came from physical oceanographers.

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The ﬁrst techniques utilised for the determination of salinity, involving the gravimetric analysis of salt left after evaporating seawater to dryness, were fraught with difﬁculties. Variable amounts of water of crystallisation might be retained. Some salts, such as MgCl2, can decompose leaving residues of uncertain composition. Other constituents, especially organic material, might be volatilised or oxidised. Overall, such methods led to considerable inconsistencies and inaccuracies. The second set of procedures for salinity measurement made use of the observation from the Challenger expedition of 1872–1876 that sea-salt composition was apparently invariant. Hence, the total salt content could be calculated from any individual constituent, such as Cl that could be readily determined by titration with Ag1. At the turn of the century, Knudsen deﬁned salinity to be the weight in grams of dissolved inorganic matter contained in 1 kg of seawater, after bromides and iodides were replaced by an equivalent amount of chloride and carbonate was converted to oxide. Clearly from the adopted definition, the analytical technique was not speciﬁc to Cl and so the term chlorinity was introduced. Chlorinity (Cl%) is the chloride concentration in seawater, expressed as g kg1, as measured by Ag1 titration (i.e., ignoring other halide contributions by assuming Cl to be the only reactant). The relationship of interest was that between S% and Cl%, given as S% ¼ 1.805Cl% þ 0.030 As a calibrant solution for the AgNO3 titrant, Standard Seawater was prepared that had certiﬁed values for both chlorinity and salinity. Unfortunately, the above salinity–chlorinity relationship was derived from only nine seawater samples that were somewhat atypical. It has since been redeﬁned using a much larger set of samples representative of oceanic waters to become S% ¼ 1.80655Cl% The third category of salinity methodologies was based on conductometry, as the conductivity of a solution is proportional to the total salt content. Standard Seawater, now also certiﬁed with respect to conductivity, provides the appropriate calibrant solution. The conductivity of a sample is measured relative to the standard and converted to salinity in practical salinity units (psu). Note that although psu has replaced the outmoded %, usually units are ignored altogether in modern usage. These techniques continue to be the most widely used methods because conductivity measurements can provide salinity values with a precision of 0.001 psu. Highly precise determinations require temperature control of samples and standards to within 0.001 1C. Application of a non-speciﬁc technique like conductometry relies upon the assumption that the sea-salt

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matrix is invariant, both spatially and temporally. Thus, the technique cannot be reliably employed in marine boundary environments where the seawater composition differs to the bulk characteristics. There are two types of conductometric procedures commonly used. Firstly, a Wheatstone Bridge circuit can be set up, whereby the ratio of the resistance of unknown seawater to standard seawater balances the ratio of a ﬁxed resistor to a variable resistor. The system uses alternating current to minimise electrode fouling. Alternatively, the conductivity can be measured by magnetic induction, in which case the sensor consists of a plastic tube containing sample seawater that links two transformers. An oscillator establishes a current in one transformer that induces current ﬂow within the tube, the magnitude of which depends upon the salinity of the sample. This in turn induces a current in the second transformer, which can then be measured. This design has been exploited for in situ conductivity measurements. 4.1.4 Oceanic Circulation The distribution of chemical components within the ocean is determined by both transportation and transformation processes. A brief outline of oceanic circulation is necessary to ascertain the relative inﬂuences. Two main ﬂow systems must be considered. Surface circulation is established by tides and the prevailing wind patterns and deep circulation is determined by gravitational forces. Both are modiﬁed by Coriolis force, the acceleration due to the earth’s rotation. It acts to deﬂect moving ﬂuids (i.e., both air and water) to the right in the northern hemisphere and to the left in the southern hemisphere. The magnitude of the effect is a function of latitude, being nil at the equator and increasing poleward. Surface oceanic circulation is depicted in Figure 4. For the most part, the circulation patterns describe gyres constrained by the continental boundaries. The prevailing winds acting under the inﬂuence of Coriolis force result in clockwise and counter-clockwise ﬂow in the northern and southern hemispheres, respectively. The ﬂow ﬁelds are non-uniform, exhibiting faster currents along the western margins. These are manifest, for example, as the Gulf Stream, Kuroshio Current and Brazil Current. Circulation within the Indian Ocean is exceptional in that there are distinct seasonal variations in accord with the monsoons. The absence of other continents within the immediate boundary region of Antarctica gives rise to a circumpolar current within the Southern Ocean. The surface circulation is restricted to the upper layer inﬂuenced by the wind, typically about 100 m. However, underlying water can be transported up into this zone when horizontal advection is insufﬁcient to

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Figure 4 The surface circulation in the ocean (Adapted from Stowe, 1979.4)

maintain the superimposed ﬂow ﬁelds. This process is called upwelling and is of considerable importance in that biochemical respiration of organic material at depth ensures that the ascending water is nutrientrich. Upwelling occurs notably in the eastern oceanic boundaries where longshore winds result in the oﬀshore transport of the surface water. Examples are found oﬀ the coasts of Peru and West Africa. Similar processes cause upwelling in the Arabian Sea, but this is seasonal due to the monsoon effect. A divergence is a zone in which the ﬂow ﬁelds separate. In such a case, upwelling may result as observed in the equatorial Paciﬁc. It should be noted that a region in which the streamlines come together is known as a convergence, and the water sinks in this zone. The density of the water controls the deepwater circulation. If the density of a water body increases, it has a tendency to sink. Subsequently, it will spread out over a horizon of uniform sy. As the density can be raised due to either an increase in the salinity or a decrease in the temperature, the deepwater circulatory system is also known as thermohaline circulation. As shown in Figure 5 of the ocean conveyor belt, the densest oceanic waters are formed in Polar Regions due to the relatively low temperatures and the salinity increase that results from ice formation. Antarctic Bottom Water (ABW) is generated in the Weddell Sea and ﬂows northward into the South Atlantic. North Atlantic Deep Water (NADW)

Chemistry of the Oceans

Figure 5

181

A schematic diagram of the thermohaline circulation of the world ocean, also know as the great ocean conveyor belt, highlighting polar regions of deepwater formation, deepwater circulation eastward from the poles and the returning westward surface water ﬂow (Adapted from IPCC, 2001.5)

is formed in the Norwegian Sea and oﬀ the southern coast of Greenland. The ﬂow of the NADW can be traced southwards through the Atlantic Ocean to Antarctica, where it is diverted eastward into the Southern Indian Ocean and South Paciﬁc. There it heads northwards and either enters the North Paciﬁc or becomes mixed upward into the surface layer in the equatorial region. The transit time is on the order of 1000 years. As noted previously, the thermocline acts as an effective barrier against mixing of dissolved components in the ocean. Consequently, this deepwater formation process in high latitudes is important because it facilitates the relatively rapid transport of material from the surface of the ocean down to great depths. The deep advection of atmospherically derived CO2 is a pertinent example. The formation of NADW could diminish with climate change because surface waters would be warmer and less dense due to enhanced melting of ice in the Arctic region. Diverse processes can form intermediate waters within the water column. In the southern South Atlantic, the NADW overrides the denser ABW. Antarctic Intermediate Water results from water sinking along the Antarctic Convergence (B501S). Relatively warm, saline

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water exits the Mediterranean Sea at depth and can be identiﬁed as a distinctive layer within the North and South Atlantic. 4.2 SEAWATER COMPOSITION AND CHEMISTRY 4.2.1 Major Constituents The major constituents in seawater are conventionally taken to be those elements present in typical oceanic water of salinity 35 that have a concentration greater than 1 mg kg1, excluding Si, which is an important nutrient in the marine environment. The concentrations and main species of these elements are presented in Table 1. One of the most significant observations from the Challenger expedition of 1872–1876 was that these major components existed in constant relative amounts. As already explained, this feature was exploited for salinity determinations. Inter-element ratios are generally constant, and often expressed as a ratio to Cl% as shown in Table 1. This implies conservative behaviour, with concentrations depending solely upon mixing processes, and indeed, salinity itself is a conservative index. Because of this behaviour, individual seawater constituents can be utilised for source apportionment studies in non-marine environments. For instance, an enrichment factor (EF) for a substance X is deﬁned as EFX ¼

ðX=Naþ Þsample

ð4:2Þ

ðX=Naþ Þseawater

An EF of 1 indicates that the substance exists in comparable relative amounts in the sample and in seawater, thereby giving a good indication Table 1

Chemical species and concentrations of the major elements in seawater

Element Na Mg Ca K Sr Cl S C (inorganic) Br B F

Chemical species Na1 Mg21 Ca21 K1 Sr21 Cl SO42, NaSO4 HCO3, CO32 Br B(OH)3, B(OH)4 F, MgF1

Concentration (mol L1) 4.79 5.44 1.05 1.05 9.51 5.59 2.89 2.35 8.62 4.21 7.51

Source: based on Dyrssen and Wedborg, 1974.6

101 102 102 102 105 101 102 103 104 104 105

For S ¼ 35 Ratio to chlorinity (g kg1) (Cl ¼ 19.374%) 10.77 1.29 0.4123 0.3991 0.00814 19.353 0.905 0.276 0.673 0.0445 0.00139

5.56 6.66 2.13 2.06 4.20 9.99 4.67 1.42 3.47 2.30 7.17

101 102 102 102 104 101 102 102 103 103 105

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of a marine origin. If EFX 4 1, then it is enriched with respect to seawater. Conversely, depletion is signiﬁed when values EFXo1. Another example of the application of inter-element ratios can be found in examining the geochemical cycle of sulfur. Concentrations of SO42 and Na1 in ice cores and marine aerosols exhibit a SO42/Na1 greater than that observed in seawater. This excess can be readily calculated and is known as non-sea salt sulfate (NSSS). Contributions to NSSS include SO2 derived from both volcanic and anthropogenic sources, together with dimethylsulﬁde (DMS) of marine biogenic origin. Not all the major constituents consistently exhibit conservative behaviour in the ocean. The most notable departures occur in deepwaters, where Ca21 and HCO3 exhibit anomalously high concentrations due to the dissolution of calcite. The concept of relative constant composition does not apply in a number of atypical environments associated with boundary regions. Inter-element ratios for major constituents can be quite different in estuaries and near hydrothermal vents. Obviously, these are not solutions of sea salt with the implication that accuracy of salinity measurements by chemical and conductometric means is limited. The residence times for some elements are presented in Table 2. The major constituents normally have long residence times. The residence time is a crude measure of a constituent’s reactivity in the reservoir. The aqueous behaviour and rank ordering can be appreciated simply in terms the ionic potential given by the ratio of electronic charge to ionic radius (Z/r). Elements with Z/ro3 are strongly cationic. The positive charge density is relatively diffuse, but sufﬁcient to attract an envelope of water molecules to form a hydrated cation. As the ionic potential increases, the force of attraction towards the water similarly rises to the extent that one oxygen – hydrogen bond in the molecule breaks. This causes the solution pH to fall and metal hydroxides to form. Neutral hydroxides tend to be relatively insoluble and so precipitate. However, in the more extreme case for which Z/r 4 12, the attraction towards the oxygen is so great that both bonds in the associated water molecules are broken. The reaction product is an oxyanion, usually quite soluble because of the associated anionic charge. Thus in seawater, those elements (Al, Fe) having a tendency to form insoluble hydroxides have short residence times. This is also true for elements that exist preferentially as neutral oxides (Mn, Ti). Hydrated cations (Na1, Ca21) and strongly anionic species (Cl, Br, UO2(CO3)24) have long residence times. This treatment is, of course, somewhat of an over-simpliﬁcation ignoring the rather significant role that biological organisms play in nutrient and trace element chemistry.

184 Table 2 Element Li B F Na Mg Al Si P Cl K Ca Sc Ti V Cr Mn Fe Co Ni Cu Zn Br Sr Ba La Hg Pb Th U

Chapter 4 The residence time and speciation of some elements in the ocean Principal species

Concentration (mol L1)

Residence time (year)

Li1 B(OH)3, B(OH)4 F, MgF1 Na1 Mg21 Al(OH)4, Al(OH)3 Si(OH)4 HPO42, PO43, MgHPO4 Cl K1 Ca21 Sc(OH)3 Ti(OH)4 H2VO4, HVO42, NaHVO2 CrO42, NaCrO4 Mn21, MnCl1 Fe(OH)3 Co21, CoCO3, CoCl1 Ni21, NiCO3, NiCl1 CuCO3, CuOH1, Cu21 ZnOH1, Zn21, ZnCO3 Br Sr21 Ba21 La31, LaCO31, LaCl21 HgCl42 PbCO3, Pb(CO3)22, PbCl1 Th(OH)4 UO2(CO3)24

2.6 105 4.1 104 6.8 105 4.68 101 5.32 102 7.4 108 7.1 105 2 106 5.46 101 1.02 102 1.02 102 1.3 1011 2 108 5 108 5.7 109 3.6 109 3.5 108 8 1010 2.8 108 8 109 7.6 108 8.4 104 9.1 105 1.5 107 2 1011 1.5 1010 2 1010 4 1011 1.4 108

2.3 106 1.3 107 5.2 105 6.8 107 1.2 107 1.0 102 1.8 104 1.8 105 1 108 7 106 1 106 4 104 1.3 104 8 104 6 103 1 104 2 102 3 104 9 104 2 104 2 104 1 108 4 106 4 104 6 102 8 104 4 102 2 102 3 106

Source: based on Brewer, 19757 and Bruland, 1983.8

4.2.2 Dissolved Gases 4.2.2.1 Gas Solubility and Air–Sea Exchange Processes. The ocean contains a vast array of dissolved gases. Some of the gases, such as Ar and chloroﬂuorocarbons (CFCs), behave conservatively and can be utilised as tracers for water mass movements and ventilation rates. Equilibrium processes at the air–sea interface generally lead to saturation, and the concentration then remains unchanged once the water sinks. Thus, the gas concentration is characteristic of the lost contact with the atmosphere. There are several important non-conservative gases, which exhibit wide variations in concentration due to biological activity. O2 determines the redox potential in seawater and CO2 buﬀers

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185

the ocean at pH 8. Ocean–atmosphere exchange processes for gases such as CO2 and DMS may play an important role in climate change. Both temperature and salinity affect the solubility of gases in water. Empirical relationships can be found elsewhere.9,10 The trends are such that gas solubility increases with a decrease in temperature or an increase in salinity. The changes in solubility are non-linear and differ dramatically for various gases. Whereas Figure 6 depicts the solubility of several gases as a function of temperature, Figure 7 shows the inﬂuence of both salinity and temperature on the solubility of molecular nitrogen. At the ocean–atmosphere interface, the exchange of gases occurs to achieve equilibrium between the two systems, and consequently gases become saturated in the water. However, supersaturation can be achieved by several mechanisms. Firstly, bubbles that form from white cap activity can be entrained and are dissolved at depth. The slight but significantly greater pressure at depth relative to the surface favours gas dissolution, and establishes a higher equilibrium concentration. Secondly as evident from Figure 6, if two water masses that have been

Figure 6 The solubility of various gases in seawater as a function of temperature (Adapted from Broecker and Peng, 1982.11)

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Chapter 4

Figure 7 The solubility of nitrogen in seawater as a function of temperature and salinity

equilibrated at different temperatures were mixed, then the resulting water body would be supersaturated. Thirdly, gases like oxygen that are produced in situ by biological activity may become supersaturated, particularly when evasion to the atmosphere is hindered. The gas solubility for a water body in equilibrium with the overlying air mass can be expressed in several ways. It is convenient to consider Henry’s Law that states H ¼ cacw1 where H is the Henry’s law constant and ca and cw refer to the concentration of a gas in air and water, respectively. As discussed by Liss (1983),12 air–sea exchange occurs when a concentration gradient exists (i.e., DC ¼ caH1cw) and the magnitude of the consequential ﬂux, F, is given as F ¼ K DC where the proportionality constant, K, has dimensions of velocity and so is frequently referred to as the transfer velocity. The concept is elaborated upon in Chapter 7. Air–sea exchange processes are consequently dependent upon the concentration gradient and the transfer velocity. The transfer velocity is not a constant, but rather depends upon several physical parameters, such as temperature, wind speed and wave state. The exchange can also be attenuated by the presence of a surface ﬁlm or slick. Alternatively, the exchange can be facilitated by bubble formation. The concentration

187

Chemistry of the Oceans Table 3 Gas CH4 CO2 N2O { CCl4 { CCl4 { CCl3F { CCl3F CH3I CO H2 Hg

The net global ﬂuxes of some trace gases across the air/sea interface Global air–sea directiona

Flux magnitude (g yr1)

þ þ ¼ ¼ þ þ þ þ

10121013 6 1015 6 1012 1010 B0 5 109 B0 3–13 1011 100 90 1012 42 1012 B2 109

þ sea to air; air to sea;¼no net ﬂux. Source: from Chester, 1990.1

gradient determines the direction of the ﬂux, into or out of the ocean. Net global ﬂuxes for some gases are presented in Table 3. The atmosphere serves as the source of material for conservative gases, especially those of anthropogenic origin, but several gases produced in situ by biological activity evade from the ocean. 4.2.2.2 Oxygen. Oxygen is a non-conservative gas and a typical oceanic proﬁle is shown in Figure 8. The concentration varies throughout the water column, its distribution being greatly inﬂuenced by biological activity. The generalised chemical equation for carbon ﬁxation is often given as nCO2 þ nH2O " (CH2O)n þ nO2 During photosynthesis this reaction proceeds to the right, thereby producing organic material, as designated by (CH2O)n, and O2. The surface waters become equilibrated with respect to atmospheric O2, but they can get supersaturated during periods of intense photosynthetic activity. Respiration occurs as the above reaction proceeds to the left and O2 is consumed. Photosynthesis is obviously restricted to the photic zone in the upper ocean and ordinarily exceeds respiration. However, the relative importance of the two processes changes with depth. The oxygen compensation depth is the horizon in the water column at which the rate of O2 production by photosynthesis equals the rate of respiratory O2 oxidation. Below the photic zone, O2 is utilised in chemical and biochemical oxidation reactions. As evident in Figure 8, the concentration diminishes

188

Chapter 4

Figure 8 A proﬁle of molecular oxygen in the North Paciﬁc Ocean (Data from Bruland, 1980.13)

with depth to develop an oxygen minimum zone. Thereafter, the O2 concentration in deeper waters begins to increase because these waters originated from Polar Regions. They were cold and in equilibrium with atmospheric gases at the time of sinking, but subsequently lost little of the dissolved O2 because the ﬂux of organic material to deepwaters is relatively small. The dissolved O2 content of seawater has a significant control on the redox potential, often designated in environmental chemistry by pe (see also Chapter 3). This is deﬁned with reference to electron activity in an analogous fashion to pH and thus pe ¼ logf e g The relationship between pe and the more familiar electrode potential E or EH is pe ¼

F E 2:303RT

ð4:3Þ

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Chemistry of the Oceans

and for the standard state

pe ¼

F E 2:303RT

ð4:4Þ

where F is Faraday’s constant, R the universal gas constant and T the absolute temperature in Kelvin. Whereas a high value of pe indicates oxidising conditions, a low value signiﬁes reducing conditions. Oxygen plays a role via the reaction O2 þ 4Hþ þ 4e Ð 2H2 O At 20 1C, K ¼ 1083.1 and so water of pH¼8.1 in equilibrium with atmospheric O2 (pO2 ¼ 0.21 atm) has pe ¼ 12.5. This conforms to surface conditions, but the pe decreases as the O2 content diminishes with depth. The oxygen minimum is particularly well developed beneath the highly productive surface waters of the eastern tropical Paciﬁc Ocean, where there is a large ﬂux of organic material to depth and subsequently considerable oxidation. The O2 becomes sufﬁciently depleted (i.e., hypoxia) that the resulting low redox conditions causes NO3 to be reduced to NO2. Aeolian transport of nitrate to Chesapeake Bay can lead to low O2 conditions. Similarly, intermittent hypoxia develops in parts of the Gulf of Mexico due to the riverine transport of nutrients derived from agricultural uses in the Mississippi catchment. When circulation is restricted vertically due to thermal or saline stratiﬁcation and horizontally by topographic boundaries, the water becomes stagnant and the oxygen may be completely utilised producing anoxic conditions. Such regions represent atypical marine environments where reducing conditions prevail. Well-known examples include the Black Sea, which is permanently anoxic below 200 m, and the Cariaco Trench, a depression in the Venezuelan continental shelf. Some fjords, such as Saanich Inlet in western Canada and Dramsfjord of Norway, may be intermittently anoxic. Periodic ﬂushing of these inlets by dense, oxygenated waters displaces deep anoxic water to the surface causing massive ﬁsh mortality. O2 can be used as a tracer to help identify the origin of water masses. The warm, saline intrusion into the Atlantic Ocean from the Mediterranean Sea is relatively O2 deﬁcient. Alternatively, the waters downwelling from Polar Regions have elevated O2 concentrations. 4.2.2.3 Carbon Dioxide and Alkalinity. Marine chemists sometimes adopt activity conventions quite different to those traditionally used in chemistry. It is useful to preface a discussion about the carbon dioxide– calcium carbonate system in the oceans with a brief outline of pH scales.

190

Chapter 4

Although originally introduced in terms of ion concentration, today the definition of pH is based on hydrogen ion activity and is pH ¼ logaH where aH refers to the relative hydrogen ion activity (i.e., dimensionless, as is pH). Deﬁned using concentration scales, the pH can be pH ¼ log cH gH =c or pH ¼ log mH gH m where cH and mH represent molar and molal concentrations, c1 and m1 are the respective standard state conditions (1 mol L1 and 1 mol kg1), and gH is the appropriate activity coefﬁcient. Obviously gH differs in these two expressions because c1 a m1. However, different activity scales may also be used. In the inﬁnite dilution activity scale, gH ) 1 as the concentration of hydrogen ions and all other ions approach 0. For analyses, pH meters are calibrated using dilute buﬀers prepared in pure water. Alternatively, in the constant ionic medium activity scale, gH ) 1 as the concentration of hydrogen ions approaches 0, while all other components are maintained at some constant level. Calibrant buﬀers are prepared in solutions of constant ionic composition, and in marine chemistry, this is often a solution of synthetic seawater. While these two methodologies are equally justiﬁable from a thermodynamic point of view, it is important to appreciate that pH scales so deﬁned are quite different. As a further consequence, the absolute values for dissociation constants also differ. The biogeochemical cycle of inorganic carbon in the ocean is extremely complicated. It involves the transfer of gaseous carbon dioxide from the atmosphere into solution. Not only is this a reactive gas that readily undergoes hydration in the ocean, but also it is ﬁxed as organic material by marine phytoplankton. Inorganic carbon can be regenerated either by photochemical oxidation in the photic zone or via respiratory oxidation of organic material at depth. Surface waters are supersaturated with respect to aragonite and calcite, different solid phases of CaCO3, but precipitation is limited to coastal lagoons such as found in the Bahamas. However, several marine organisms utilise calcium carbonate to form shells. Sinking shells can remove inorganic carbon from surface waters that is then regenerated following dissolution in the under-saturated waters found at depth. Nonetheless, calcitic oozes of biogenic origin constitute a major component in marine sediments.

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Chemistry of the Oceans

Finally, the inorganic carbon equilibrium is responsible for buﬀering seawater at a pH near 8 on timescales of centuries to millennia. There are several equilibria to be considered (see also Chapter 3). Firstly, CO2 is exchanged across the air–sea interface CO2g Ð CO2aq The equilibrium process obeys Henry’s Law, but the dissolved CO2 reacts rapidly with water to become hydrated as CO2aq þ H2 O Ð H2 CO3 Relative to the exchange process, the hydration reaction forming carbonic acid occurs quite quickly. This means that the concentration of dissolved CO2 is extremely low. The two processes can be considered together as CO2g þ H2 O Ð H2 CO3 The equilibrium constant is then fH2 CO3 g KCO2 ¼ p CO2 fH2 Og

ð4:5Þ

where pCO2 is the partial pressure of CO2 in the marine troposphere. Carbonic acid undergoes dissociation H2 CO3 þ H2 O Ð H3 Oþ þ HCO 3 þ 2 HCO 3 þ H2 O Ð H3 O þ HCO3

for which the ﬁrst and second dissociation constants (using {H1} rather than {H3O1}) are fH þ g HCO 3 ð4:6Þ K1 ¼ fH2 CO3 g fH þ g CO23 K2 ¼ HCO 3

ð4:7Þ

The hydrogen ion activity can be established with a pH meter. However, as discussed above, this measurement must be operationally deﬁned. On the other hand, the individual ion activities of bicarbonate and carbonate ions cannot be measured. Instead, ion concentrations are determined, as outlined below, by titration. Accordingly, the equilibrium constants are redeﬁned in terms of concentrations. These are then known as apparent, rather than true, equilibrium constants and

192

Chapter 4

distinguished using a prime notation. It must be appreciated that apparent equilibrium constants are not invariant, but rather are aﬀected by temperature, pressure, salinity, and, as outlined previously, the pH scale adopted. The apparent dissociation constants are fH þ g HCO 3 0 K1¼ ð4:8Þ ½H2 CO3 fH þ g CO23 K2¼ HCO 3 0

ð4:9Þ

It should be noted that whereas ion activities are denoted by curly brackets {}, concentrations are designated by square brackets []. Analogous to the pH, pK conventionally refers to logK. Numerical values for the constants pKCO2 , pK 0 1 and, pK 0 2 based on a constant ionic medium scale (i.e., seawater with chlorinity¼19%) are given in Table 4. This provides sufﬁcient information to calculate the speciation of carbonic acid in seawater at a given temperature as a function of pH. This is shown for carbonic acid at 20 1C in both pure water and seawater equilibrated with atmospheric carbon dioxide in Figure 9. While there are several confounding features, the pH of seawater can be considered to be buﬀered by the bicarbonate/carbonate pair. The pH is generally about 8, but is sensitive to the concentration ratio [HCO3]:[CO32] as evident from rearranging the expression for K 0 2 to become þ 0 HCO3 ð4:10Þ fH g ¼K2 2 CO3 To understand the response of the oceanic CO2 system to in situ biological activity or enhanced CO2 concentrations in the atmosphere, Table 4

Equilibrium constants for the carbonate system

T ( 1C)

pKCO2

pK 0 1

pK 0 2

0 5 10 15 20 25 30

1.19 1.27 1.34 1.41 1.47 1.53 1.58

6.15 6.11 6.08 6.05 6.02 6.00 5.98

9.40 9.34 9.28 9.23 9.17 9.10 9.02

Source: adapted from Stumm and Morgan, 1996.14

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Chemistry of the Oceans

Figure 9

The distribution of carbonic acid species at 20 1C as a function of pH in seawater of salinity 35(solid lines) and pure water (dotted lines)

it is necessary to consider in more detail the factors inﬂuencing the inorganic carbon cycle. Two useful parametersPcan be introduced. Firstly, the total concentration of inorganic carbon, CO2, in seawater is P CO2 ¼ [CO2] þ [H2CO3] þ [HCO3] þ [CO32] The ﬁrst term is negligible and as evident in Figure 9, the major species at pH 8 are HCO3 and CO32. Alkalinity is deﬁned as a measure of the proton deficit in solution and should not to be confused with basicity. Alkalinity is operationally deﬁned by titration with a strong acid to the carbonic acid end point. This is known as the titration alkalinity (TA). Seawater contains weak acids other than bicarbonate and carbonate that are titrated, and therefore TA is given as TA ¼ [HCO3] þ 2[CO32] þ [B(OH)4] þ [OH] [H1] This is a simpliﬁed version of the more general expression in Chapter 3. The inﬂuence of [OH] and [H1] on the TA is small and can often be ignored. The borate contributes about 3% of the TA, and if not determined independently, can be estimated from the apparent boric acid dissociation constants and the salinity, relying upon the relative constancy of composition of sea salt. This correction would give the carbonate alkalinity (CA) CA ¼ [HCO3] þ 2[CO32] Considering the dissociation constants above, this can be alternatively expressed as CA ¼

K 0 1 ½H2 CO3 2 K 0 1 K 0 2 ½H2 CO3 þ 2 fH þ g fH þ g

ð4:11Þ

194

Chapter 4

or CA ¼

K 0 1 KCO2 pCO2 2 K 0 1 K 0 2 KCO2 pCO2 þ 2 fH þ g fH þ g

ð4:12Þ

This equation can be rearranged to give the following quadratic expression that can be solved for the pH 2

CAfH þ g K 0 1 KCO2 pCO2 fHþ g 2 K 0 1 K 0 2 KCO2 pCO2 ¼ 0

Worked Example 1 For seawater (35 psu, 15 1C) with an alkalinity of 2.30 meq l1 and in equilibrium with atmospheric CO2 ¼ 3.65 104 atm, calculate (i) the pH and (ii) the speciation of carbonic acid. (i) pH calculation Data, including constants from Table 4 are KCO2 ¼ 101.41, K 0 1 ¼ 106.05, K 0 2 ¼ 109.23, CA¼2.30 103, pCO2¼3.65 104 The pH is obtained from the calculation of {H1} using the above quadratic equation. Thus pﬃﬃﬃﬃﬃﬃﬃﬃﬃﬃﬃﬃﬃﬃﬃﬃﬃﬃ b þ b2 4ac þ fH g ¼ 2a where a ¼CA ¼2:30 103 b ¼ K10 KCO2 pCO2 ¼ 106:5 101:41 103:46 ¼ 1:20 1011 c ¼ 2 K10 K20 KCO2 pCO2 ¼ 2106:05 109:23 101:41 103:46 ¼ 1:42 1020 Giving {H1}¼6.22 109 and pH¼8.21. (ii) Carbonic acid speciation calculations Knowing the pH, each of the three major species (H2CO3, HCO 3 , P CO32) can be calculated as a fraction (or percentage) of the CO2.

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Chemistry of the Oceans

Note that the negligible contribution due to dissolved CO2 is ignored. The necessary expressions are derived from the definitions of the K 0 1 and K 0 2 given previously. 2 CO3 ¼

HCO 3

K10 ½H2 CO3 fHþ g

ð4:13Þ

K 01 K 02 K20 HCO ½H2 CO3 3 ¼ þ 2 fH g fH þ g

Substituting into the expression for the summation of all carbonic speciesP gives CO2 ¼ [H2CO3] þ [HCO3] þ [CO32]¼ [H2CO3] þ K 0 1{H1}1[H2CO3]þK 0 1 K 0 2{H1}2[H2CO3]¼ [H2CO3](1þ K 0 1{H1}1þK 0 1 K 0 2{H1}2) Thereafter, the fractional contribution of each species to the total can be calculated using 1:

½H2 CO3 P ¼ ½CO2

ð1 þ K10 f Hþ g 1 þ K10 K20 f Hþ g 2 Þ1 ½HCO3 0 þ P ¼ K1 f H g ½CO2 1 þ2 1 1 þ K10 fHþ g þK10 K20 fHþ g

½CO3 P

2 ½CO2

¼ ðK10 K20 f Hþ g

ð1 þ K10 f Hþ g 0

þ K10 K20 f Hþ g

2 1

Substituting the values for K 1 and K 2 and using the previously calculated pH of 8.21, the fractional contribution of each species is 0.006, 0.907 and 0.087 for H2CO3, HCO3, and CO32, respectively. Consider now the effect of altering the pCO2 in the water. The alkalinity should not change in response to variations in CO2 alone because the hydration and dissociation reactions give rise to equivalent amounts of H1 and anions. CO2 can be lost by evasion to the atmosphere (a process usually conﬁned to equatorial regions) or by photoP synthesis. This causes the CO2 to diminish and the pH to rise, an effect that can be quite dramatic in tidal rock pools in which pH may then rise to 9. Conversely, an increase in pCO2, either by invasion from the atmosphere or release following respiration, prompts an increase in

196

Chapter 4

P CO2 and a fall P in pH. Thus, the depth proﬁles of pH would mimic that of O2 but the CO2 would exhibit a maximum at the oxygen minimum. There are further confounding inﬂuences, in particular concerning the solid phases of CaCO3(s). Many organisms make calcareous shells (testes) of CaCO3(s) in the form of either aragonite or calcite. The shells sink and dissolve when the organism dies. The solubility is governed by Ca2þ þ CO2 3 Ð CaCO3ðsÞ Surface waters are supersaturated with respect to CaCO3(s), but precipitation rarely occurs, possibly due to an inhibitory effect by Mg21aq forming ion pairs with CO32aq. The solubility of CaCO3(s) increases with depth, due to both a pressure effect and the decrease in pH following respiratory release of CO2, causing the shells to dissolve. This behaviour not only increases the alkalinity, but also accounts for the non-conservative nature of Ca21 and inorganic carbon in deepwaters. The depth at which appreciable dissolution begins is known as the lysocline. At a greater depth, designated as the carbonate compensation depth (CCD), no calcareous material is preserved in the sediments. The depths of the lysocline and CCD are inﬂuenced by the ﬂux of organic material and shells, and tend to be deeper under high productivity zones. The CO2 and CaCO3(s) systems are coupled in that the pH buﬀering in the ocean is due to the reaction CO2 þ H2 O þ CaCO3ðsÞ Ð Ca2þ þ 2HCO 3 In addition to the effects noted previously, an input of CO2 promotes the dissolution of CaCO3(s). The reaction does not proceed to the right without constraint, but rather meets a resistance given by the Revelle factor, R dpCO2 =pCO2 P R¼ P d CO2 = CO2

ð4:14Þ

This value is approximately 10, indicating that the ocean is relatively P well buﬀered against changes in CO2 in response to variations in atmospheric pCO2. Although the ocean does respond to an increase in the atmospheric burden of CO2, the timescales involved are quite considerable. The surface layer can become equilibrated on the order of decades, but as the thermocline inhibits exchange into deepwaters, the equilibration of the ocean as a whole with the atmosphere proceeds on the order of centuries. The ventilation of deepwater by downwelling

Chemistry of the Oceans

197

water masses in polar latitudes only partly accelerates the overall process.

4.2.2.4 Dimethylsulﬁde and Climatic Implications. The Gaia hypothesis of Lovelock15 states that the biosphere regulates the global environment for self-interest. This pre-supposes that controls, perhaps poorly understood or unknown, serve to maintain the present status quo. Charlson et al.16 have made use of this hypothesis to suggest that biogenic production of DMS and the consequent formation of atmospheric cloud condensation nuclei (CCN, i.e., small particles onto which water can condense) acts as a feedback mechanism to counteract the global warming resulting from elevated greenhouse gas concentrations in the atmosphere. The cycle is illustrated in Figure 10. Global warming, with concurrent warming of the ocean surface, leads to enhanced phytoplankton productivity. This promotes the production and evasion to the atmosphere of DMS. The DMS undergoes oxidation to form CCN that promote cloud formation and increase the planetary albedo (i.e., reﬂectivity with respect to sunlight) thereby causing a cooling effect. From a biogeochemical perspective, the two key features are the controls on the biogenic production of DMS and the formation of CCN following aerial oxidation of DMS. These will be considered below in more detail. With respect to the physics, the most important aspects of the proposed climate control mechanism are that the enhancement of the albedo is due to an increase in the number and type of CCN, and that this CCN production occurs in the marine boundary layer. The albedo of calm seawater is very low (B2%) in comparison to vegetated regions (10–25%), deserts (B35%), and snow-covered surfaces (B90%). That biological processes within the oceans act as a major source of reduced sulfur gases is well established,17 and of particular importance is the generation of DMS. Surface concentrations, approximately in the range of 0.7–17.8 nmol L1, exhibit large temporal and geographic variations. Oceanic distributions indicate that DMS is produced within the photic zone, that is consistent with a phytoplankton source, but DMS concentrations are poorly correlated with normal indicators of primary productivity. While Phaeocystis and Coccolithoporidae have been identiﬁed as important DMS producers, there is uncertainty as to the full potential for biological DMS formation. With respect to climate modiﬁcation, questions remain as to the biological response to global warming. For the model of Charlson et al.16 to hold, organisms might either increase DMS formation or biological succession could change in such a way as to favour DMS producers. Thus, marine biogenic source

198

Figure 10

Chapter 4

The possible climatic inﬂuence of DMS of marine biogenic origin (Adapted from Charlson et al., 1987.16)

strengths and the controlling factors remain important unresolved issues in sulfur biogeochemistry. DMS concentrations in the remote marine troposphere vary in the range of 0.03–32 nmol m3. Not surprisingly and as with the seawater concentrations, considerable temporal and geographic disparities occur. Furthermore, atmospheric DMS concentrations exhibit diurnal variations, with a night time maximum and an afternoon minimum consistent with a photochemical sink. Whereas oxidation involves HO free radicals

Chemistry of the Oceans

199

during the day, a reaction with NO3 may be important at night. Relatively low levels are associated with air masses derived from continental areas, owing to the enhanced concentrations of oxidants. While oceanic venting rates are dependent upon a number of meteorological and oceanographic conditions, there is no question that the marine photic zone acts as the major source of DMS to the overlying troposphere. The oxidation of DMS in the atmosphere (see also Chapter 2) could yield several products, namely dimethylsulfoxide (DMSO), methanesulfonate (MSA) or sulfate. Insofar as aerosol formation is concerned, the two key products are MSA and SO42. Atmospheric particles in the sub-micron size range exert a significant inﬂuence on the earth’s climate. The effect can be manifested via three mechanisms. Firstly, the particles themselves may enhance backscatter of solar radiation. Secondly, they act as CCN promoting cloud formation and so increasing the earth’s albedo. Thirdly, such clouds affect the hydrological cycle. The evidence for such a biofeedback mechanism limiting global warming remains circ*mstantial. 4.2.3 Nutrients Although several elements are necessary to sustain life, traditionally in oceanography the term ‘‘nutrients’’ has referred to nitrogen (notably, but not exclusively, nitrate), phosphorus (usually as phosphate), and silicon (as silicate). The rationale for this classiﬁcation was that analytical techniques had long been available that allowed the precise determination of these constituents despite their relatively low concentrations. They were observed to behave in a consistent and explicable manner, but quite differently to the major constituents in seawater. The distributions of these three nutrients are determined by biological activity. Nitrate and phosphate become incorporated into the soft parts of organisms. As evident in the modiﬁed carbon ﬁxation equation of Redﬁeld18 given below 3 106CO2 þ 16NO 3 þ HPO4 þ 122H2 O

þ 18Hþ Ð C106 H263 O110 N16 P þ 138O2 The uptake of these nutrients into tissues occurs in constant relative amounts. The ratio (i.e., Redﬁeld ratio) for C/N/P is 106:16:1. Silicate is utilised by some organisms, particularly diatoms (phytoplankton) and radiolaria (zooplankton), to form siliceous skeletons. Such skeletons

200

Chapter 4

consist of an amorphous, hydrated silicate, SiO2 nH2O, often called opaline silica. Photosynthetic carbon ﬁxation depends on the availability of the nutrients, and of course, it ceases when one of the nutrient supplies is exhausted. The limiting nutrient in the ocean is usually nitrogen, but there are conditions under which Fe may limit biological productivity. Biological productivity is greatest in regions with the greatest supply of nutrients, such as coastal regions and near zones of upwelling. Moreover, biological activity is seasonal at most latitudes, being inﬂuenced by PAR coupled with the depletion of nutrients from surface waters that become thermally stratiﬁed due to summer warming. The biogeochemical cycling of nitrogen in the marine environment is quite complex, largely owing to the diversity of forms in which it is present, coupled with the ease with which organisms can either assimilate or transform the various nitrogenous species. Nitrogen exhibits the range of oxidation states, from N(V) in nitrate to N(III) in NH3 and numerous organic compounds. Also, N occurs as dissolved gaseous molecules, including N2 and N2O; dissolved inorganic nitrogen (DIN) such as NO3 and NO2; dissolved organic nitrogen (DON) present in organic compounds like amino acids; and particulate organic nitrogen (PON) in organisms and detritus. Few marine organisms are capable of assimilating N2 directly, (known as nitrogen ﬁxation) despite its relative abundance in seawater. The nitrogenous species more commonly used are nitrate, nitrite, urea and ammonium. Bacteria play an important role in the redox chemistry of nitrogen species in seawater. Starting with PON, the ﬁrst step is remineralisation in which PON is converted to DON. The breakdown of some of the DON to DIN follows with the ﬁrst product being NH3; the process is relatively rapid and known as ammoniﬁcation. NH3 is protonated to a limited extent in seawater, giving rise to NH41 ions. Nitriﬁcation is the stepwise oxidation of NH41 to NO2 and eventually to NO3. Denitriﬁcation, the reduction of nitrogen species to N2, can occur under conditions of hypoxia or anoxia. In such cases, bacteria respire organic material using NO3 and NO2 as electron acceptors. A depth proﬁle of nitrate, phosphate and silicate in the North P Paciﬁc Ocean is presented in Figure 11. Nutrients behave much like CO2 and are removed in the surface layer, especially in the photic zone. Thus, concentrations can become quite low, and indeed sufﬁciently low to limit further photosynthetic carbon ﬁxation. The organisms sink following death. The highest concentrations occur where respiration and bacterial decomposition of the falling organic material are greatest, that is at the oxygen minimum. The nutrients, including silica, are consequently

Chemistry of the Oceans

201

Figure 11 The depth distribution of nitrate (K), phosphate (’), and silicate (m) in the North Paciﬁc Ocean (Data from Bruland, 1980.13)

regenerated and their concentrations in deepwaters are much greater than those observed in the surface waters, thereby accounting for the fertilising effect of upwelling. It should be noted that the siliceous remains behave differently than the calcareous shells discussed previously. The oceans everywhere are under-saturated with respect to silica. Its solubility exhibits no pronounced variation with depth and there is no horizon analogous to the CCD (see Section 4.2.2.3). Silica is preserved to any great extent only in deep-sea sediments associated with the highly productive upwelling zones in the ocean.

4.2.4 Trace Elements Trace elements in seawater are taken to be those that are present in quantities less than 1 mg L1, excluding the nutrient constituents. The distribution and behaviour of minor elements have been reviewed in the light of data that conform to an oceanographically consistent manner.1,8

202

Chapter 4

Analytical difﬁculties are readily comprehensible when it is appreciated that the concentration for some of these elements can be extremely low, i.e., a few pg L1 for platinum group metals.19 Some trace elements, such as Cs1, behave conservatively and therefore absolute concentrations depend upon salinity. More often, the elements are non-conservative and their distributions in both surface waters and the water column vary greatly, reﬂecting the differing source strengths and removal processes in operation. Generalisations regarding residence times cannot be made in many cases because biologically active elements are removed from seawater relatively rapidly. Nevertheless, conservative constituents and platinum group metals have rather long residence times on the order of 105 years. Considering ﬁrstly the distribution in surface waters, several elements exhibit high concentrations in coastal waters in comparison to levels in the centres of oceanic gyres. Typically, this distribution arises because the elements originate predominantly from riverine inputs or through diffusion from coastal sediments. However, as they are effectively removed from the surface waters in the coastal regions, little material is advected horizontally to the open sea. Examples of elements that behave in this way are Cd, Cu and Ni. In contrast, the concentration of Pb, including 210Pb, is greater in the gyres. This results from a strong widespread aeolian input coupled with less effective removal from surface waters in the gyres. Clearly, the removal mechanisms have an appreciable effect on dissolved elemental abundances. The two major processes in operation are uptake by biota and scavenging by suspended particulate material. In the ﬁrst instance, the constituent mimics the behaviour of nutrients. This is evident in the metal/nutrient correlation for Cd/P and Zn/Si (Figure 12). No consistent pattern for depth proﬁles of trace elements exists. Conservative elements trend with salinity variations provided they have no significant submarine sources. Non-conservative elements may exhibit peak concentrations at different depths in oxygenated waters as (i) (ii) (iii) (iv)

surface enrichment, maximum at the O2 minimum, mid-depth maximum not associated with the O2 minimum, and bottom enrichment.

The criteria for an element, such as Pb, to exhibit a maximum concentration in surface waters are that the only significant input must be at the surface (aeolian supply) and it must be effectively removed from the water column. Constituents such as As, Ba, Cd, Ni and Zn exhibit nutrient type behaviour. Those elements (Cd) associated with the soft parts of the organism are strongly correlated with phosphate and are

Chemistry of the Oceans

Figure 12 Metal/nutrient correlations in the North Paciﬁc for Cd/P and Zn/Si (Adapted from Bruland, 1980.13)

203

204

Chapter 4

regenerated at the O2 minimum. Elements (i.e., Zn) associated with the skeletal material may exhibit a smoothly increasing trend with depth. The third case pertains to elements, notably Mn, that have a substantial input from hydrothermal waters. These are released into oceanic waters from spreading ridges. Ocean topography is such once these waters are advected away from such regions towards the abyssal plains, they are then found at some intermediate depth. Bottom enrichment is observed for elements (Mn) that are remobilised from marine sediments. The behaviour of Al combine features outlined above, resulting in a middepth minimum concentration. Surface enrichment evident in mid (411 N) but not high (B601 N) latitudes in the North Atlantic results from the solubilisation of aeolian material. Removal occurs via scavenging and incorporation into siliceous skeletal material. Subsequent regeneration by shell dissolution increases deepwater Al levels. 4.2.5 Physico-Chemical Speciation Physico-chemical speciation refers to the various physical and chemical forms in which an element may exist in the system. In oceanic waters, it is difﬁcult to determine chemical species directly. Whereas some individual species can be analysed, others can only be inferred from thermodynamic equilibrium models as exempliﬁed by the speciation of carbonic acid in Figure 9. Often an element is fractionated into various forms that behave similarly under a given physical (e.g., ﬁltration) or chemical (e.g., ion exchange) operation. The resulting partition of the element is highly dependent upon the procedure utilised, and so known as operationally deﬁned. In the following discussion, speciation will be exempliﬁed with respect to size distribution, complexation characteristics, redox behaviour and methylation reactions. Physico-chemical speciation determines the environmental mobility of an element, especially with respect to partitioning between the water and sediment reservoirs. The inﬂuence can be manifested through various mechanisms as summarised in Figure 13. Settling velocities, and by implication the residence time, are controlled by the size of the particle. Thus, dissolved to particulate interactions involving adsorption, precipitation or biological uptake can readily remove a constituent from the water column. The redox state can have a comparable inﬂuence. Mn and Fe are reductively remobilised in sediments. Following their release in reduced forms from either hydrothermal sources or interstitial waters, they rapidly undergo oxidation to form colloids, which then are quickly removed from the water column. Speciation also determines the bioavailability of an element for marine organisms. It is generally accepted

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Figure 13 Speciation of metal ions in seawater and the main controlling mechanisms (Adapted from O¨hman and Sjo¨berg, 1988.20)

that the uptake of trace elements is limited to free ions and some types of lipid–soluble organic complexes. This is important in that some elements may be essential, but toxic at elevated concentrations. Via complexation reactions, organisms may be able to modify the free ion concentration. Thus, siderophores can sequester Fe, often in limited supply. Alternatively, metallothionein detoxify contaminants, notably Cd and Hg, through chelation. An element can exist in natural waters in a range of forms that exhibit a size distribution as indicated in Table 5. Entities in true solution include ions, ion pairs, complexes and a wide range of organic molecules that can span several size categories. At the smallest extreme are ions, which exist in solution with a co-ordinated sphere of water molecules as discussed previously. Na1, K1 and Cl exist predominantly as free hydrated ions. Collisions of oppositely charged ions occur due to electrostatic attraction and can produce an ion pair. An ion pair is the transient coupling of a cation and anion during which each retains its own co-ordinated water envelope. While impossible to measure directly, concentrations can be calculated with knowledge of the ion activities and the stability constant. For the formation of the ion pair NaSO4 via Naþ þ SO24 Ð NaSO 4

206 Table 5

Chapter 4 The size distribution of trace metal species in natural waters

Size range o 1 nm 1–10 nm

10–100 nm 100–1000 nm 4 1000 nm

Metal species

Examples

Free metal ions Inorganic ion pairs Inorganic complexes Low molecular mass organic complexes High molecular mass organic complexes Metal species adsorbed onto inorganic colloids Metals associated with detritus Metals absorbed into living cells Metals adsorbed onto or incorporated into mineral solids and precipitates

Phase state 21

Mc , Cd NiCl1 HgCl24 Zn-fulvates

Soluble Soluble

Pb-humates

Colloidal

Co–MnO2

Particulate

Pb–Fe(OH)3 Cu–clays

Particulate

PbCO3(s)

Source: adapted from de Mora and Harrison, 1984.21

the stability constant is deﬁned as K¼

fNaSO 4g þ fNa gfSO24 g

ð4:15Þ

Ion pair formation is important for Ca21, Mg21, SO42 and HCO3. If the attraction is sufﬁciently great, a dehydration reaction can occur leading to covalent bonding. A complex consists of a central metal ion sharing a pair of electrons donated by another constituent, termed a ligand, acting as a Lewis base. The metal ion and ligand share a single water envelope. Ligands can be neutral (e.g., H2O) or anionic (e.g., Cl, HCO3) species. A metal ion can co-ordinate with one or more ligands, which need not be the same chemical entity. Alternatively, the cation can share more than one electron pair with a given ligand thereby forming a ring structure. This type of complex, known as a chelate, exhibits enhanced stability largely due to the entropy effect of releasing large numbers of molecules from the co-ordinated water envelopes. Complex formation is an equilibrium process. Ignoring charges for the general case of a metal, M and ligand, L, complex formation occurs as M þ L " ML for which the formation constant, K1, is given by K1 ¼

fMLg fMgfLg

ð4:16Þ

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Chemistry of the Oceans

A second ligand may then be co-ordinated as M þ L " ML2 K2 ¼

fML2 g fMLgfLg

ð4:17Þ

The equilibrium constant for ML2 can be expressed solely in terms of the activities of the M and L b2 ¼

fML2 g fMgfL g2

ð4:18Þ

where ß2 is the product of K1K2 and is known as the stability constant. The case can be extended to include n ligands as bn ¼

fMLn g fMgfL gn

ð4:19Þ

Worked Example 2 Assuming all g 0 s ¼ 1, calculate the speciation of mercury in typical seawater (35 psu at 25 1C) given the following values for stepwise stability constants for successive chlorocomplexes (K1 ¼ 106.74, K2 ¼ 106.48, K3 ¼ 100.85, K4 ¼ 101.00). Note that [Cl] is 0.559 mmol L1 and that it is not necessary to know the mercury concentration in seawater. The total mercury concentration is given as the sum of all contributing species. Thus [Hg]T ¼ [Hg2 1] þ [HgCl1] þ [HgCl2]þ[HgCl3] þ [HgCl42] From the definition of the stability constants, we know that [HgCl1] ¼ K1[Hg21][Cl] and [HgCl2] ¼ b2[Hg21][Cl]2, etc. Thus, [Hg]T¼[Hg21](1þK1[Cl]þb2[Cl]2þb3[Cl]3þb4[Cl]4). Now let D ¼ (1 þ K1[Cl] þ b2[Cl]2 þ b3[Cl]3 þ b4[Cl]4)1, and note that this is a constant for a stipulated chloride concentration. Thus, at the seawater chloride concentration ([Cl] ¼ 100.25) this gives D ¼ (1þ106.74 100.25 þ 106.74106.48100.25 þ 106.7406.48100.85100.25 þ 106.74106.48100.85101.00 100.25)1 ¼ 7.27 1015 The fractional (or percentage) contribution of each species can be determined using

208

Fractional contribution ½Hg2þ : ½Hg2þ T ½HgCl þ : ½Hg2þ T ½HgCl2 : ½Hg2þ T ½HgCl3 : ½Hg2þ T ½HgCl24- : ½Hg2þ T

Chapter 4

Expression D D K1½Cl D b2 ½Cl 2 D b3 ½Cl 3 D b4 ½Cl 4

Value for ½Cl ¼ 0:559 mmol L-1 7:3 1015 2:2 108 3:7 102 ð4%Þ 1:5 101 ð15%Þ 8:2 101 ð82%Þ

Thus, at the high chloride concentration found in seawater, the mercury speciation is dominated by the tetra- and tri-chloro species. Note (1) If the total mercury content for a given seawater sample were known, the concentration of each species is readily calculated. That is [Hg21] ¼ fHG21 [Hg21]T; (2) This is a simpliﬁed example that has not considered other species (bromides, hydroxides, etc.) that might be important in seawater. The speciation of constituents in solution can be calculated if the individual ion activities and stability constants are known. This information is relatively well known with respect to the major constituents in seawater, but not for all trace elements. Some important confounding variables create considerable difﬁculties in speciation modelling. Firstly, it is assumed that equilibrium is achieved, meaning that neither biological interference nor kinetic hindrance prevents this state. Secondly, seawater contains appreciable amounts (at least in relation to the trace metals) of organic matter. However, the composition of the organic matrix, the number of available binding sites, and the appropriate stability constants are poorly known. Nevertheless, speciation models can include estimates of these parameters. Organic material can form chelates with relatively high-stability constants and dramatically decrease the free ion activity of both necessary and toxic trace elements. Organisms may make use of such chemistry, producing compounds either to sequester metals in limited supply (e.g., siderophores to complex Fe) or to detoxify contaminants (e.g., metallothionein to chelate Cd, Hg, etc.). Thirdly, surface adsorption of dissolved species onto colloids or suspended particles may remove them from solution. As with organic matter, an exact understanding of the complexation characteristics of the suspended particles is not available, but approximations can also be incorporated into speciation models. Elements may be present in a variety of phases other than in true solution. Colloidal formation is particularly important for elements such as Fe and Mn, which produce amorphous oxyhydroxides with very great

Chemistry of the Oceans

209

complexation characteristics. Adsorption processes cannot be ignored in biogeochemical cycling. Particles tend to have a much shorter residence time in the water column than do dissolved constituents. Scavenging of trace components by falling particles accelerates deposition to the sediment sink. Several elements in seawater may undergo alkylation via either chemical or biological mechanisms.22 Type I mechanisms involve methyl radical or carbonium ion transfer and no formal change in the oxidation state of the acceptor element. The incoming methyl group may be derived for example from methylcobalamin coenzyme, S-adenosylmethionine, betaine or iodomethane. Elements involved in Type I mechanisms include Pb, Tl, Se and Hg. Other reaction sequences involve the oxidation of the methylated element. The methyl source can be a carbanion from methylcobalamin coenzyme. Oxidative addition from iodomethane and enzymatic reactions has also been suggested. Some elements that can undergo such methylation processes are As, Sb, Ge, Sn and S. Methylation can enhance the toxicity of some elements, especially for Pb and Hg. The environmental mobility can also be affected. Methylation in the surface waters can enhance volatility and so favour evasion from the sea, as observed for S, Se and Hg. Methylation within the sediments may facilitate transfer back into overlying waters. Elements may exhibit multiple oxidation states in seawater. Redox processes can be modelled in an analogous manner to the ion pairing and complexation outlined previously. The information is often presented graphically in the form of a predominance area diagram, which is a plot of pe vs. pH showing the major species present for the designated conditions. Although a single oxidation state might be anticipated from equilibrium considerations, there are several ways in which multiple oxidation states might arise. Biological activity can produce non-equilibrium species, as evident in the alkylated metals discussed above. Whereas Mn(IV) and Cu(II) might be expected by thermodynamic reasoning, photochemical processes in the surface waters can lead to the formation of significant amounts of Mn(II) and Cu(I). Fe(III) is the favoured redox state of Fe in seawater, but it is relatively insoluble and exists predominantly in a colloidal phase. Photochemical reduction to Fe(II), which only slowly oxidises to Fe(III), might act as a very important mechanism rendering Fe bioavailable to marine organisms. Goldberg19 has presented impressive information (given that seawater concentrations are as low as 1.5 pg L1 for Ir and 2 pg L1 for Ru) on the speciation, including redox state, of platinum group metals as a means of interpreting distributions in seawater and marine sediments. Pt and Pd are stabilised in seawater as tetrachloro-divalent anions. Their

210

Chapter 4

relative abundance of 5 Pt:1 Pd agreeing with a factor of ﬁve difference in ß4. Pt is enriched in ferromanganese nodules following oxidation to a stable (IV) state, behaviour not observed for Pd. Rh exists predominantly in the heptavalent state, but accumulates in reducing sediments as lower valence sulfides. Au and Ag are present predominantly in solution as monovalent forms. Ag, but not Au, accumulates in anoxic coastal sediments. 4.3 SUSPENDED PARTICLES AND MARINE SEDIMENTS 4.3.1 Description of Sediments and Sedimentary Components The sediments represent the major sink for material in the oceans. The main pathway to the sediments is the deposition of suspended particles. Such particles may be only in transit through the ocean from a continental origin or be formed in situ by chemical and biological processes. Sinking particles can scavenge material from solution. Accordingly, this section introduces the components found in marine sediments, but emphasises processes that occur within the water column that lead to the formation and alteration of the deposited material. Marine sediments cover the ocean ﬂoor to a thickness averaging 500 m. The deposition rates vary with topography. The rate may be several millimetres per year in nearshore shelf regions, but is only from 0.2 to 7.5 mm per 1000 years on the abyssal plains. Oceanic crustal material is formed along spreading ridges and moves outwards eventually to be lost in subduction zones, the major trenches in the ocean. Because of this continual movement, the sediments on the seaﬂoor are no older than Jurassic in age, about 166 million years. The formation of marine sediments depends upon chemical, biological, geological and physical inﬂuences. There are four distinct processes that can be readily identiﬁed. Firstly, the source of the material obviously is important. This is usually the basis for classifying sediment components and will be considered below in more detail. Secondly, the material and its distribution on the ocean ﬂoor are inﬂuenced by its transportation history, both to and within the ocean. Thirdly, there is the deposition process that must include particle formation and alteration in the water column. Finally, the sediments may be altered after deposition, a process known as diagenesis. Of particular importance are reactions leading to changes in the redox state of the sediments. The components in marine sediments are classiﬁed according to origin. Examples are given in Table 6. Lithogenous (or terrigenous) material comes from the continents as a result of weathering processes.

211

Chemistry of the Oceans Table 6

The four categories of marine sedimentary components with examples of mineral phases

Classiﬁcation

Mineral example

Chemical formula

Lithogenous

Quartz Microcline Kaolinite Montmorillonite Illite Chlorite Fe–Mn mineral Carbonate ﬂuoroapatite Barite Pyrite Aragonite Dolomite Calcite Aragonite Opaline silica Apatite Barite Organic matter Cosmic spherules Meteoric dusts

SiO2 KA1Si3O8 Al4Si4O10(OH)8 Al4Si8O20(OH)4 nH2O K2Al4(Si, Al)8O20(OH)4 (Mg,Fe21)10Al2(Si,Al)8O20(OH,F)16 FeO(OH) – MnO2 Ca5(PO4)3x(CO3)xF11x BaSO4 FeS2 CaCO3 CaMg(CO3)2 CaCO3 CaCO3 SiO2 nH2O Ca5(F,Cl) (PO4)3 BaSO4

Hydrogenous

Biogenous

Cosmogenous

Source: adapted from Harrison and de Mora, 1996.23

The relative contribution of lithogenous material to the sediments will depend on the proximity to the continent and the source strength of material derived elsewhere. The most important components in the lithogenous fraction are quartz and the clay minerals (kaolinite, illite, montmorillonite and chlorite). The distribution of the clay minerals varies considerably. Illite and montmorillonite tend to be ubiquitous in terrestrial material, but the latter has a secondary origin associated with submarine volcanic activity. Kaolinite typiﬁes intense weathering observed in tropical and desert conditions. Therefore, it is relatively enriched in equatorial regions. On the other hand, chlorite is indicative of the high-latitude regimes where little chemical weathering occurs. The lithogenous components tend to be inert in the water column and represent detrital deposition. Nonetheless, the particle surfaces can act as important sites for adsorption of organic material and trace elements. Hydrogenous components, also known as chemogenous or halmeic material, are those produced abiotically within the water column. This may comprise primary material formed directly from seawater upon exceeding a given solubility product, termed authigenic precipitation. The best-known examples of authigenic material are the various types of ferromanganese nodules found throughout the oceans. Alternatively,

212

Chapter 4

secondary material may be formed as components of continental or volcanic origin become altered by low-temperature reactions in seawater, a mechanism known as halmyrolysis. Halmyrolysis reactions can occur in the estuarine environment, being essentially an extension of chemical weathering of lithogenous components. Such processes continue at the sediment–water interface. Accordingly, there are considerable overlaps between the terms weathering, halmyrolysis and diagenesis. Owing to the importance that surface chemistry has on the ﬁnal composition, authigenic precipitation and halmyrolysis are considered further in Section 4.3.2. Biogenous (or biotic) material is produced by the ﬁxation of mineral phases by marine organisms. The most important phases are calcite and opaline silica, although aragonite and magnesian calcite are also deposited. As indicated in Table 7, several plants and animals are involved, but the planktonic organisms are the most important with respect to the World Ocean. The source strength depends upon the biological species composition and productivity of the overlying oceanic waters. For instance, siliceous oozes are found in polar latitudes (diatoms) and along the equator (radiolaria). The relative contribution of biogenous material to the sediments depends upon its dilution by material from other sources and the extent to which the material can be dissolved in seawater. As noted previously, both calcareous and siliceous skeletons are subject to considerable dissolution in the water column and at the sediment–water interface. There are two sources that give rise to minor components in the marine sediments. Cosmogenous material is that derived from an extra-terrestrial source. Such material tends to comprise small (i.e., o0.5 mm) black Table 7

Quantitatively important plants and animals that secrete calcite, aragonite, Mg-calcite, and opaline silica

Mineral

Plants

Animals a

Calcite

Coccolithophorids

Aragonite

Green algae

Mg-calcite

Coralline (red) algae

Opaline silica

Diatomsa

Planktonic organisms Source: adapted from Berner and Berner, 1987.24

Foraminiferaa Molluscs Bryozoans Molluscs Corals Pteropodsa Bryozoans Benthic foraminifera Echinoderms Serpulids (tubes) Radiolariaa Sponges

Chemistry of the Oceans

213

micrometeorites or cosmic spherules. The composition is either magnetite or a silicate matrix including magnetite. They are ubiquitous but scarce, with relative contributions to the sediments decreasing with an increase in sedimentation rate. Finally, there are anthropogenic components, notably heavy metals and Sn, which can have a significant inﬂuence on sediments in coastal environments. As noted in Section 4.1.1, the principal modes of transport of particulate material to the ocean are by rivers or via the atmosphere. Within the oceans, distribution is further affected by ice rafting, turbidity currents, organisms and oceanic currents. Turbidity currents refer to the turbid and turbulent ﬂow of sediment-laden waters along the seaﬂoor caused by sediment slumping. They are especially important in submarine canyons and can transport copious amounts of material, including coarse-grained sediments, to the deep sea. Ice rafting can also transport substances to the deep sea. Although, ice rafting is presently conﬁned to the polar latitudes (401N and 551S), there have been considerable variations in ice limits within the geologic record. Organisms are notable not only for biogenous sedimentation, but also they can inﬂuence ﬁne-grained lithogenous material that becomes incorporated into faecal pellets consequently accelerating the settling rate. Ocean currents are important for the distribution of material with a long residence time. Major surface currents are zonal and tend to reinforce the pattern of aeolian supply. On the other hand, deepwater currents are of little consequence as velocities are slow relative to the settling rates. 4.3.2 Surface Chemistry of Particles 4.3.2.1 Surface Charge. Parcles in seawater tend to exhibit a negative surface charge. There are several mechanisms by which this might arise. Firstly, the negative charge can result from crystal defects (i.e., vacant cation positions) or cation substitution. Clay minerals are layered structures of octahedral AlO6 and tetrahedral SiO4. Either substitution of Mg(II) and Fe(II) for the Al(III) in octahedral sites or replacement of Si(IV) in tetrahedral location by Al(III) can cause a net negative charge. Secondly, a surface charge can result from the differential dissolution of an electrolytic salt such as barite (BaSO4). A charge will develop whenever the rate of dissolution of cations and anions differs. Thirdly, organic material can be negatively charged due to the dissociation of acidic functional groups. Adsorption processes can also lead to the development of a negatively charged particle surface. One example is the speciﬁc adsorption of anionic organic compounds onto the surfaces of particles. Another

214

Chapter 4

mechanism relates to the acid–base behaviour of oxides in suspension. Metal oxides (most commonly Fe, Mn) and clay minerals have frayed edges on account of broken metal–oxygen bonds. The surfaces can be hydrolysed and exhibit amphoteric behaviour þ X O ðsÞ þ Haq Ð X OHðsÞ þ þ X O ðsÞ þ Haq Ð X OH2ðsÞ

The hydroxide surface exhibits a different charge depending upon the pH. Cations other than H1 can act as the potential determining ion. The point of zero charge (PZC) is the negative log of the activity at which the surface exhibits no net surface charge. At the PZC þ ½X O ðsÞ ¼ ½X OH2ðsÞ

The PZC for some mineral solids found in natural waters are shown in Table 8. Clearly, the extent to which such surfaces can adsorb metal cations will be dependent upon the pH of the solution. At the pH typical of seawater, most of the surfaces indicated in Table 8 would be negatively charged and would readily adsorb metal cations. 4.3.2.2 Adsorption Processes. Physical or non-speciﬁc adsorption involves relatively weak attractive forces, such as electrostatic attraction and van der Waals forces. Adsorbed species retain their co-ordinated Table 8

The PZC for some mineral phases

Mineral a-Al2O3 a-Al2(OH)3 g-AlOOH CuO Fe3O4 a-FeOOH g-Fe2O3 ‘Fe(OH)3’ (amorphous) MgO d-MnO2 b-MnO2 SiO2 ZrSiO4 Feldspars Kaolinite Montmorillonite Albite Chrysotile

pHPZC 9.1 5.0 8.2 9.5 6.5 7.8 6.7 8.5 12.4 2.8 7.2 2.0 5 2–2.4 4.6 2.5 2.0 412

Source: adapted from Stumm and Morgan, 1996.14

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Chemistry of the Oceans

sphere of water and hence, cannot approach the surface closer than the radius of the hydrated ion. Adsorption is favoured by ions having a high charge density, i.e., trivalent ions in preference to univalent ones. Additionally, an entropy effect promotes the physical adsorption of polymeric species, such as Al and Fe oxides, because a large number of water molecules and monomeric species is displaced. Chemisorption or speciﬁc adsorption involves greater forces of attraction than physical adsorption. As hydrogen bonding or p–orbital interactions are utilised, the adsorbed species lose their hydrated spheres and can approach the surface as close as the ionic radius. Whereas multilayer adsorption is possible in physical adsorption, chemisorption is necessarily limited to monolayer coverage. As outlined previously, hydrated oxide surfaces have sites that are either negatively charged or readily deprotonated. The oxygen atoms tend to be available for bond formation, a favourable process for transition metals. Several mechanisms are possible. An incoming metal ion, Mz1, may eliminate an H1 ion as X O H þ Mzþ Ð X O Mðz-1Þþ þ Hþ Alternatively, two or more H1 ions may be displaced, thereby forming a chelate as —X – O – H

—X – O + Mz+

—X – O – H

M(z−2)+ + 2H+ —X – O

MLnz1,

A metal complex, may be co-ordinated instead of a free ion by displacement of one or more H1 ions in a manner analogous to the above reaction. In addition, the metal complex might eliminate a hydroxide group giving rise to a metal–metal bond as ðz-1Þþ þ OH X O H þ MLzþ n Ð X MLn

It should be noted that all of these reactions are equilibria for which an appropriate equilibrium constant can be deﬁned and measured. This data can then be incorporated into the speciation models discussed in Section 4.2.5. 4.3.2.3 Ion Exchange Reactions (see also Chapter 5). Both mineral particles and particulate organic material can take up cations from solution and release an equivalent amount of another cation into solution. This process is termed cation exchange and the cation

216

Chapter 4

exchange capacity (CEC) for a given phase is a measure of the number of exchange sites present per 100 g of material. This is operationally deﬁned by the uptake of ammonium ions from 1 mol L1 ammonium acetate at pH 7. The speciﬁc surface area and CEC are given in Table 9 for several sorption-active materials. There are several factors that inﬂuence the afﬁnity of cations towards a given surface. Firstly, the surface coverage will increase as a function of the cation concentration. Secondly, the afﬁnity for the exchange site is enhanced as the oxidation state increases. Finally, the higher the charge density of the hydrated cation, the greater will be its afﬁnity for the exchange site. In order of increasing charge density, the group I and II cations are BaoSroCaoMgoCsoRboKoNaoLi 4.3.2.4 Role of Surface Chemistry in Biogeochemical Cycling. Reactions at the aqueous–particle interface have several consequences for material in the marine environment, from estuaries to the deep-sea sediment–water boundary. Within estuarine waters, suspended particles experience a dramatic change in the composition and concentration of dissolved salts. A number of halmyrolysates can be formed. Clay minerals undergo cation exchange as Mg21 and Na1 replace Ca21 and K1. Alternatively, montmorillonite may take up K1 becoming transformed to illite. Hydrogenous components, in particular Mn and Fe oxides, may be precipitated onto the surfaces of suspended particles. The particulate material generally accumulates organic coatings within estuaries, which together with an increase in the ionic strength of the surrounding solution, leads to the formation of stable colloids. Both the oxide and organic coatings can subsequently scavenge other elements in the estuary. Table 9

The speciﬁc surface area and cation-exchange capacities of several sorption-active materials

Material

Speciﬁc surface area (mg2 g1)

Cation-exchange capacity (meq/100 g)

Calcite (o2 mm) Kaolinite Illite Chlorite Montmorillonite Freshly precipitated Fe(OH)3 Amorphous silicic acid Humic acids from soils

12.5 10–50 30–80 — 50–150 300 — 1900

— 3–15 10–40 20–50 80–120 10–25 11–34 170–590

Source: adapted from Fo¨rstner and Wittman, 1981.25

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Chemistry of the Oceans

Within the ocean, the exchange of material from the dissolved to the suspended particulate state inﬂuences the distribution of several elements. This scavenging process removes dissolved metals from solution and accelerates their deposition. The effectiveness of this process is obvious in the depth proﬁles of metals, especially those of the surface enrichment type. Furthermore, the removal can be expressed in terms of a deepwater scavenging residence time as indicated in Table 10. The scavenging mechanism can be particularly effective in the water– sediment boundary region. The re-suspension of ﬁne sediments generates a very large surface area for adsorption and ion exchange processes. Within the immediate vicinity of hydrothermal springs, reduced species of Mn and Fe are released and subsequently are oxidised to produce colloidal oxyhydroxides, which have a large surface area and very great sorptive characteristics. Finally, ferromanganese nodules form at the sediment–water interface and become considerably enriched in a number of trace metals via surface reactions. Ferromanganese nodules result from the authigenic precipitation of Fe and Mn oxides at the seaﬂoor. Two morphological types are recognised, depending upon the growth mechanism. Firstly, spherical encrustations produced atop oxic deep-sea sediments grow slowly, accumulating material from seawater. These seawater nodules exhibit a relatively low Mn:Fe ratio and are especially enriched with respect to Co, Fe and Pb. Secondly, discoid shaped nodules develop in nearshore environments deriving material via diffusion from the underlying anoxic sediments. Such diagenetic nodules grow faster than deep-sea varieties and metals tend to be in lower oxidation states. They have a high Mn:Fe ratio and enhanced content of Cu, Mn, Ni and Zn. Ferromanganese nodules have concentric light and dark bands in cross-section, related to

Table 10

The deepwater scavenging residence times of some trace elements in the oceans

Element

Scavenging residence time (year)

Element

Scavenging residence time (year)

Sn Th Fe Co Po Ce Pa Pb

10 22–33 40–77 40 27–40 50 31–67 47–54

Mn Al Sc Cu Be Ni Cd Particles

51–65 50–150 230 385–650 3700 15 850 117 800 0.365

Source: adapted from Chester, 1990.1

218

Chapter 4

Fe and Mn oxides, respectively. Patterns of trace element enrichment in the nodules are determined by mineralogy, the controlling mechanisms being related to cation substitution in the crystal structure. Mn phases preferentially accumulate Cu, Ni, Mo and Zn. Alternatively, Co, Pb, Sn, Ti and V are enriched in Fe phases. 4.3.3 DIAGENESIS Diagenesis refers to the collection of processes that alter the sediments following deposition. These mechanisms may be physical (compaction), chemical (cementation, mineral segregation, ion exchange reactions) or biological (respiration). The latter are of particular importance as the bacteria control pH and pe in the interstitial waters, which in turn affect a wide range of chemical equilibria. These master variables inﬂuence the composition of the interstitial water, and can exert a feedback effect on the overlying seawater. Also, they can ultimately control the mineralogical phases that are lost to the sediment sink. Organic material accumulates with other sedimentary components at the time of deposition. High biological productivity in surface waters and rapid sedimentation ensures that most nearshore and continental margin sediments contain significant amounts of organic matter. Biochemical oxidation of this material exhausts the available O2 creating anoxic conditions. The oxic/anoxic boundary occurs at the horizon where the respiratory consumption of O2 balances its downward diffusion. Upon depleting the O2, other constituents are used as oxidants leading to the stepwise depletion of NO3, NO2 and SO42. Thereafter, organic matter itself may be utilised with the concurrent production of CH4. This series of reactions causes progressively greater reducing conditions, with consequent inﬂuences on the chemistry of several elements. Metals are reduced and so are present in lower oxidation states. In particular, Mn undergoes reductive dissolution from MnO2s to Mnaq21. The divalent state being much more soluble, Mn is effectively remobilised under anoxic conditions and can be released back into overlying seawater. As seen in the previous section, this is one pathway to ferromanganese nodule formation. This can also be true for other elements that had been deposited following incorporation into the Fe and Mn oxide phases. In contrast, some elements can be preserved very effectively in anoxic sediments. Interstitial waters in marine sediments, in contrast to freshwater deposits, have high initial concentrations of SO42. Bacterial sulfate reduction proceeds via the reaction 2CH2 O þ SO2 4 Ð H2 S þ 2HCO3

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Thus, sulﬁde levels in interstitial waters increase. A number of elements form insoluble sulfides, which under these anoxic conditions are precipitated and retained within the sediments. A notable example is the accumulation of pyrite, FeS2, but also Ag, Cu, Pb and Zn are enriched in anoxic sediments in comparison with oxic ones. 4.4 PHYSICAL AND CHEMICAL PROCESSES IN ESTUARIES Rivers transport material in several phases: dissolved, suspended particulate and bed load. Physical and chemical processes within an estuary inﬂuence the transportation and transformation of this material, thereby affecting the net supply of material to the oceans. Several definitions and geomorphologic classiﬁcations of estuaries have been reviewed by Perillo (1995).26 From a chemical perspective, an estuary is most simply described as the mixing zone between river water and seawater characterised by sharp gradients in the ionic strength and chemical composition. Geographic distinctions can be made between drowned river valleys, fjords and bar-built estuaries. They can alternatively be classiﬁed in terms of the hydrodynamic regime as: (i) (ii) (iii) (iv)

salt wedge, highly stratiﬁed, partially stratiﬁed, and vertically well mixed.

The aqueous inputs into the system are the river ﬂow and the tidal prism. The series above is ranked according to the diminishing importance of the riverine ﬂow and the increasing marine contribution. Thus, a salt wedge estuary represents the extreme case dominated by river ﬂow, in which very little mixing occurs. A fresh, buoyant layer ﬂows outward over denser, saline waters. In contrast, the vertically well-mixed estuary is one dominated by the tidal prism. The inﬂowing river water mixes thoroughly with and dilutes the seawater, but the effective dilution diminishes with distance along the mixing zone. The position of the mixing zone in the estuary exhibits considerable temporal variations. There can be a strong seasonal effect, largely due to non-uniform river discharge. High winter rainfall leads to a wintertime discharge maximum. However, winter precipitation as snow creates a storage reservoir, such that the river ﬂow maximum occurs following snow melt in spring or even early summer if the catchment area is of high elevation as for the Fraser River in western Canada. On shorter timescales, the mixing zone is inﬂuenced by tidal cycles. Thus, the penetration of seawater into the estuary depends upon the spring–neap tidal

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cycle and the diurnal nature of the tides. Together these inﬂuences determine the geographic extent of sediments experiencing a variable salinity regime. Variations in the river discharge affect the mass loading of the discharge, both in terms of suspended sediment and bed load material. The hydrodynamic regime in the estuary inﬂuences the deposition of the riverine sedimentary material and the mixing of dissolved material. The estuary is a mixing zone for river water and seawater, the characteristics of which differ considerably. River water is slightly acidic and of low ionic strength with a salt matrix predominantly of Ca(HCO3)2 (B120 mg L1). In contrast, seawater has a higher pH (B8), higher ionic strength (B0.7) and consists primarily of NaCl (B35 g L1). As a consequence, the salt matrix within the estuary is dominated by the sea salt end member throughout the mixing zone except for a small proportion at the dilute extreme. Salinity can be used as a conservative index, although conductivity is better, not being subject to systematic conversion errors in the initial mixing region. In a plot of concentration vs. some conservative index (i.e., S%, Cl% or conductivity), the theoretical dilution curve would comprise a straight line between the river and seawater end members. A dissolved constituent that exhibits such a distribution is said to behave conservatively in the estuary. Whereas a negative slope shows that the riverine end member is progressively diluted during mixing with seawater, a positive slope indicates that the seawater end member has the greater concentration. Conservative behaviour is exhibited, for example, by Na1, K1 and SO42. Reactive silica may at times behave conservatively. Non-conservative behaviour can result from an additional supply of material (i.e., causing positive deviations from the theoretical dilution curve). Elements that may show a maximum concentration at some intermediate salinity are Mn and Ba. Alternatively, the removal of dissolved material during mixing (i.e., negative deviation from the theoretical dilution curve) can be caused by biological activity or by dissolved to particulate transformations. Biological activity can cause non-conservative behaviour for nutrient elements, including reactive silica. Dissolved constituents typically transformed to the particulate phase include Al, Mn and Fe (see Figure 14) in some estuaries. The pH distribution is usually characterised by a pH minimum in the initial mixing zone, resulting from the non-linear salinity dependence of the ﬁrst and second dissociation constants of H2CO3. Notwithstanding the obvious utility of component–conservative index plots, they can be applied and interpreted only with caution. Often it is difﬁcult to deﬁne the exact composition of the end members. Hence, a plot that apparently denotes non-conservative behaviour could arise if temporal

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ﬂuctuations in the concentration of the component of interest occur on the same relative time scales as estuarine ﬂushing. A component can undergo considerable physico-chemical speciation alterations in an estuary. With respect to dissolved constituents, the composition and concentration of available ligands changes. Depending upon the initial pH of the riverine water, OH may become markedly more important down the estuary. Similarly, chlorocomplexes for metals such as Cd, Hg and Zn become more prevalent as the salinity increases. Conversely, the competitive inﬂuence of seawater derived Ca and Mg for organic material decreases the relative importance of humic complexation for Mn and Zn. Estuaries are particularly well known for dissolved–particulate interactions. Such changes in phase come about via several mechanisms. Firstly, dissolution–precipitation processes may occur. This is especially important for the authigenic precipitation of Fe and Mn oxyhydroxides. Secondly, components may experience adsorption–desorption reactions. Desorption can occur in the initial mixing zone, partly in response to the pH minimum. Adsorption, particularly in association with the Fe and Mn phases, can accumulate material within the suspended sediments. Thirdly, ﬂocculation and aggregation processes can remove material from solution. This occurs as particulate material with negatively charged surfaces adsorb cations in the estuary. The surface charge diminishes and as the ionic strength increases, the particles experience less electrostatic repulsion. Eventually, the situation arises whereby particle collisions lead to aggregation due to weak bonding. This process can be facilitated if particle surfaces are coated with organic material (then known as ﬂocculation rather than aggregation). The three types of processes outlined above often happen simultaneously in the estuarine environment. Dissolved Fe, as shown in Figure 14, typiﬁes such non-conservative behaviour. The transformation of dissolved into particulate phases can then be followed by deposition to the estuarine sediments. Thus, the ﬂux of material to the ocean can be considerably modiﬁed, particularly as such sediments may be transported landward rather than seaward. Biological activity in the estuarine environment can also inﬂuence the speciation of constituents, notably dissolved–particulate partitioning. A complex regeneration cycle determines distributions and modiﬁes the riverine ﬂux. This is especially the case for nutrients, and estuaries are often termed a nutrient trap. Estuaries tend to be regions of high biological productivity as rivers have elevated nutrient concentrations. Moreover, several freshwater organisms die upon encountering brackish water with consequent cell rupture and the release of contents into solution. Regardless of the source, the nutrients stimulate phytoplankton productivity.

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6 Taieri estuary

Dissolved Fe (µmol/kg)

Salinity (‰)

Figure 14 The distribution of dissolved Fe vs. Salinity in the Taieri Estuary, New Zealand; 21 October 1980, 4 December 1980 (Adapted from Hunter, 1983.27)

Whereas some of the biological material (either transported via the freshwater or produced in situ within the estuarine zone) can be ﬂushed out to sea, the remainder settles out of the surface water to be deposited onto the ﬂoor of the estuary. Respiration of this debris by benthic organisms regenerates nutrients and any contaminants that have been accumulated, releasing them into the bottom water. In estuaries where a two-layer ﬂow is well deﬁned, these nutrients are transported upriver in the salt wedge and entrained into the exiting river water, thereby adding to the available nutrient pool. Polluted rivers having a high nutrient loading are subject to eutrophication due to over-stimulation of biological activity. Anoxic conditions within the bottom waters and/or underlying sediments can result, depending upon the organic loading. Aeolian transport of nitrate to Chesapeake Bay can lead to low O2 conditions. As mentioned previously, some fjords develop anoxic conditions when bottom waters stagnate due to limited mixing.

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4.5 MARINE CONTAMINATION AND POLLUTION Both contamination and pollution entail the perturbation of the natural state of the environment by anthropogenic activity. The two terms are distinguishable in terms of the severity of the effect, whereby pollution induces the loss of potential resources.28 Additionally, a clear cause– effect relationship must be established for a substance to be classiﬁed as a pollutant towards a particular organism. The human-induced disturbances take many forms, but the greatest effects tend to be in coastal environments due to the source strengths and pathways. Waters and sediments in coastal regions bear the brunt of industrial and sewage discharges, and are also subject to dredging and spoil dumping. Agricultural runoﬀ may contain pesticide residues and elevated nutrients, the latter of which may over-stimulate biological activity producing eutrophication and anoxic conditions. The deep sea has not escaped contamination. The most obvious manifestations comprise crude oil, petroleum products and plastic pollutants, but include the long-range transport of long-lived radionuclides from coastal sources. Additionally, the aeolian transport of heavy metals has enhanced natural ﬂuxes of some elements, particularly lead. Three case studies are introduced below to illustrate diverse aspects of marine contamination and pollution.

4.5.1 Oil Slicks Major releases of oil have been caused by the grounding of tankers (i.e., Torrey Canyon, Southwest England, 1967; Argo Merchant, Nantucket Shoals, USA, 1976; Amoco Cadiz, Northwest France, 1978; Exxon Valdiz, Alaska, 1990; Erika, France, 1999; Prestige, Spain, 2002) and by the accidental discharge from oﬀshore platforms (i.e., Chevron MP41C, Mississippi Delta, 1970; Ixtox I, Gulf of Mexico, 1979). Because oil spills receive considerable public attention and provoke substantial anxiety, oil pollution must be put into perspective. Crude oil has been habitually introduced into the marine environment from natural seeps at a rate of approximately 340 106 L y1. Anthropogenic activity has recently augmented this supply by an order of magnitude; however, most of this additional oil has originated from relatively diffuse sources relating to municipal runoﬀ and standard shipping operations. Exceptional episodes of pollution occurred in the Persian Gulf in 1991 (910 106 L y1) and due to the Ixtox I well in the Gulf of Mexico in 1979 (530 106 L y1). In contrast to such mishaps, the Amoco Cadiz discharged only 250 106 L y1 of oil in 1978 accounting for the largest spill from a tanker. The cumulative pollution from tanker accidents on an annual

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basis matches that emanating from natural seepage. Nevertheless, the impacts can be severe when the subsequent slick impinges on coastal ecosystems. Regardless of the source, the resultant oil slicks are essentially surface phenomena that are affected by several transportation and transformation processes.29 With respect to transportation, the principal agent for the movement of slicks is the wind, but length scales are important. Whereas small weather systems, such as thunderstorms, tend to disperse the slick, cyclonic systems can move the slick essentially intact. Waves and currents also affect the advection of an oil slick. To a more limited extent, diffusion can act to transport the oil. Transformation of the oil involves phase changes and degradation. Several physical processes can invoke phase changes. Evaporation of the more volatile components is a significant loss mechanism, especially for light crude oil. The oil slick spreads as a buoyant lens under the inﬂuence of gravitational forces, but generally separates into distinctive thick and thin regions. Such pancake formation is due to the fractionation of the components within the oil mixture. Sedimentation can play a role in coastal waters when rough seas bring dispersed oil droplets into contact with suspended particulate material and the density of the resulting aggregate exceeds the speciﬁc density of seawater. Colloidal suspensions can consist of either water-in-oil or oil-in-water emulsions, which behave distinctly differently. Water-in-oil emulsiﬁcation creates a thick, stable colloid that can persist at the surface for months. The volume of the slick increases and it aggregates into large lumps known as ‘‘mousse’’, thereby acting to retard weathering. Conversely, oil-in-water emulsions comprise small droplets of oil in seawater. This aids dispersion and increases the surface area of the slick, which can subsequently accelerate weathering processes. Chemical transformations of oil are evoked primarily through photochemical oxidation and microbial biodegradation. Not only is the latter more important in nature, but strategies can be adopted to stimulate biological degradation, consequently termed bioremediation. All marine environments contain microorganisms capable of degrading crude oil. Furthermore, most of the molecules in crude oils are susceptible to microbial consumption. Oil contains little nitrogen or phosphorus and therefore, microbial degradation of oil tends to be nutrient limited. Bioremediation often depends upon on the controlled and gradual delivery of these nutrients, while taking care to limit the concurrent stimulation of phytoplankton activity. Approaches that have been adopted are the utilisation of slow-release fertilisers, oleophilic nutrients and a urea-foam polymer fertiliser incorporating oil-degrading

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bacteria. Bioremediation techniques were successfully applied in the clean up of Prince William Sound and the Gulf of Alaska following the Exxon Valdez accident. Alternative bioremediation procedures relying on the addition of exogenous bacteria have still to be proved. Similarly, successful bioremediation of ﬂoating oil spills has yet to be demonstrated. Source apportionment of crude oil in seawater and monitoring the extent of weathering and biodegradation constitute important challenges in environmental analytical chemistry. As the concentration of individual compounds varies from one sample of crude oil to another, the relative amounts deﬁne a signature characteristic of the source. Compounds that degrade at the same rate stay at ﬁxed relative amounts throughout the lifetime of an oil slick. Hence, a ‘‘source ratio’’, which represents the concentration ratio for a pair of compounds exhibiting such behaviour remains constant. Conversely, a ‘‘weathering ratio’’ reﬂects the concentration ratio for two compounds that degrade at different rates and consequently continually changes. Oil spill monitoring programmes conventionally determine four fractions30: (i) (ii) (iii) (iv)

volatile hydrocarbons; alkanes; total petroleum hydrocarbons; and polycyclic aromatic hydrocarbons (PAHs).

The volatile hydrocarbons, albeit comparatively toxic to marine organisms, evaporate relatively quickly and hence serve little purpose as diagnostic aids. The alkanes and total petroleum hydrocarbons make up the bulk of the crude oil. They can be used to some extent for source identiﬁcation and monitoring weathering progress. The ﬁnal fraction, the PAHs, comprises only about 2% of the total content of crude oil but includes compounds that are toxic. Moreover, these components exhibit marked disparities in weathering behaviour due to differences in water solubility, volatility and susceptibility towards biodegradation. As demonstrated in Figure 15, both a source ratio (C3-dibenzothiophenes/ C3-phenanthrenes) and a weathering ratio (C3-dibenzothiophenes/ C3-chrysenes) have been deﬁned from among such compounds that enable the extent of crude oil degradation to be estimated in the marine environment, as well as for sub-tidal sediments and soils.30 The discussion above focuses on petroleum hydrocarbons from oil spills, representing just one source of contamination. Marine contaminant surveys are conducted not just in response to tanker accidents, but routinely as a means to assess the quality of the marine environment. Although analytical difﬁculties with environmental samples persist and

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Figure 15 Plot of weathering ratio (C3-dibenzothiophenes/C3-chrysenes) vs. source ratio (C3-dibenzothiophenes/C3-phenanthrenes) for fresh and degraded oil samples from three different crude oil spills (Adapted from Douglas et al., 1996.30)

much material can only be classiﬁed as an unresolved complex mixture (UCM), many individual aliphatic and aromatic compounds can be routinely determined. Total petroleum hydrocarbon distributions, as shown in Figure 16 for the Caspian Sea, provide evidence of pollution ‘‘hot spots’’.31 In this case, the highest concentrations were observed in Azerbaijan. Given that several individual PAHs can be ascribed to different origins, plots of the relative amounts allows an assessment to be made of the relative importance of petrogenic (fossil oil), pyrolytic (combustion products) and natural (in situ biological activity) sources. As shown in Figure 17, most of the PAHs in the Caspian Sea sediments were derived from fossil oil.31 This technique cannot distinguish between material derived from natural seeps or anthropogenic activity. Otherwise, the contribution of combustion products was evident in Russia and natural sources were only observed in the Iranian sector.

4.5.2 Plastic Debris The accumulation of litter and debris along shorelines epitomises a general deterioration of environmental quality on the high seas. The material originates not only from coastal sources, but also from the ancient custom of dumping garbage from ships. Drilling rigs and oﬀshore production platforms have similarly acted as sources of contamination. Some degree of protection in recent years has accrued from both the London Dumping Convention (LDC) and the International

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Figure 16

227

Distribution of total hydrocarbons (total aliphatic and aromatic) in sediments from the Caspian Sea (mg g1 dry wt) (From Tolosa et al., 2004.31)

Convention for the Prevention of Pollution from Ships (MARPOL) that outlaw such practices. However, the problem of seaborne litter remains global in extent and not even Antarctica has been left unaffected.32 The debris consists of many different materials, which tending to be non-degradable endures in the marine environment for many years. The most notorious are the plastics (e.g., bottles, sheets, ﬁshing gear, packaging materials and small pellets), but also includes glass bottles, tin cans and lumber. This litter constitutes an aesthetic eyesore on beaches, but more importantly can be potentially lethal to marine organisms. Deleterious impacts on marine birds and mammals result from entanglement and ingestion. Lost or discarded plastic ﬁshing nets remain functional and can continue ‘‘ghost ﬁshing’’ for several years. This is similarly true for traps and pots that go astray. Plastic debris settling on soft and hard bottoms can smother benthos and limit gas exchange with

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Figure 17 Ternary diagram showing the relative amounts of PAHs from different sources (fossil oil, combustion products and natural biological activity) and distinguished on a country basis (From Tolosa et al., 2004.31)

pore waters. Despite the negative effects of seaborne plastic debris, this material can have positive consequences serving as new habitats for opportunistic colonisers. 4.5.3 Tributyltin Tributyltin (TBT) provides an interesting case study of a pollutant in the marine environment.33 Because TBT compounds are extremely poisonous and exhibit broad-spectrum biocidal properties, they have been utilised as the active ingredient in marine anti-fouling paint formulations. Its potency and longevity ensures good fuel eﬃciencies for ship operations and guarantees a long lifetime between repainting. TBTbased paints have been used on boats of all sizes, from small yachts to supertankers, ensuring the global dispersion of TBT throughout the marine environment, from the coastal zone to the open ocean. Notwithstanding such beneﬁts, the extreme toxicity and environmental persistence has resulted in a wide range of deleterious biological effects on non-target organisms. TBT is lethal to some shellﬁsh at concentrations as low as 0.02 mg TBT-SnL1. Lower concentrations result in sub-lethal effects, such as poor growth rates and reduced recruitment leading to the decline of shellﬁsheries. The most obvious manifestations of TBT contamination have been shell deformation in Paciﬁc oysters (Crassostrea gigas), and the development of imposex (i.e., the imposition of male sex organs on females) in marine gastropods. The later effect has caused dramatic population decline of gastropods at locations throughout the

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world. TBT has been observed to accumulate in ﬁsh and various marine birds and mammals, with yet unknown consequences. Although it has not been shown to pose a public health risk, one recent study reported measurable butyltin concentrations in human liver.34 The economic consequences of the shellﬁsheries decline led to a rapid political response globally. The ﬁrst publication35 suggesting TBT to be the causative agent appeared in only 1982, and led to prompt legislative action. France was the ﬁrst country to control the use of TBT-based paints, in the ﬁrst instance on boats o25 m in length at some coastal locations. Oyster aquaculture in Arcachon Bay beneﬁted immediately, with a notably decline in shell deformations and TBT-body burdens, and the complete recovery of production within two years (Figure 18). Several other countries introduced various controls and bans on these marine anti-foulants. Comparable improvements in oyster conditions were subsequently observed in Great Britain and Australia. Similarly, there have been many reported instances of restoration of gastropod populations at previously impacted locations. Certainly the TBT ﬂux to the marine environment decreased, as manifested in sedimentary TBT proﬁles. However, most of the legislation ﬁrst introduced applied only to small vessels and large ships continued to act as a source of TBT to the marine environment. Imposexed gastropods were observed at sites (e.g., North Sea and Strait of Malacca) where the source of TBT could only

Figure 18 Annual oyster (Crassostrea gigas) production in Arcachon Bay, 1978–1985; Restrictions on TBT use were applied starting January 1982 (Data from Alzieu, 1991.36)

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have been attributed to shipping. Global concern was sufﬁciently great that in 2001 the International Maritime Organisation (IMO) successfully negotiated the International Convention on the Control of Harmful Anti-fouling Systems on Ships. The Convention aims to ban the application of organotin-based anti-fouling paints on ships from 2003 and their presence on ship hulls from 2008, but does require ratiﬁcation of 25 countries comprising 25% of the world’s merchant shipping tonnage before coming into effect. TBT exists in solution as a large univalent cation and forms a neutral complex with Cl or OH. It is extremely surface active and so is readily adsorbed onto suspended particulate material. Such adsorption and deposition to the sediments limits its lifetime in the water column. Degradation, via photochemical reactions or microbially mediated pathways, obeys ﬁrst-order kinetics. Several marine organisms, as diverse as phytoplankton to starﬁsh, debutylate TBT. Stepwise debutylation produces di- and mono-butyltin, which are much less toxic in the marine environment than is TBT. As degradation rates in the water column are on the order of days to weeks, they are slow relative to sedimentation. TBT accumulates in the sediments where degradation rates are much slower, with the half-life being on the order of years.37 Furthermore, concentrations are highest in those areas, such as marinas and harbours, which are most likely to undergo dredging. The intrinsic toxicity of TBT, its persistency in the sediments and its periodic remobilisation by anthropogenic activity are likely to retard the long-term recovery of the marine ecosystem. Although the unrestricted use of TBT should end, significant challenges remain in many parts of the world. For the most part, the coastal tropical ecosystems remain unprotected and the sensitivity of its indigenous organisms relatively poorly evaluated. TBT endures in sediments globally, with concentrations usually greatest in environments most likely to be perturbed, such as ports and marinas. The widespread introduction of TBT into seawater will continue from vessels not yet subject to legislation until at least 2008, depending on when the IMO Convention comes into force. However, the paramount lesson learned from TBT should be that potential replacement compounds must be properly investigated prior to their introduction in order to avoid another global pollution experiment. QUESTIONS (i) What are the main sources and sinks for dissolved and particulate metals entering the ocean?

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(ii) Deﬁne residence time and outline the factors that inﬂuence the residence time of an element in the ocean. Provide examples of oceanic residence times for elements that span the time scale. (iii) Deﬁne salinity. What are the notable variations in sea salt, in terms of concentration and composition, in the world ocean? (iv) Using data available in Table 4, calculate the concentration of H2CO3 in seawater at 51C in equilibrium with atmospheric CO2 at 370 ppm. (v) Calculate the speciation of H2S in seawater at pH 8.3 at 251C given that pK1 ¼ 7.1 and pK2 ¼ 17.0. (vi) Explain how DMS of marine origin might affect global climate. (vii) What are the major nutrients in the ocean? Describe their concentration proﬁles and account for any differences between them according to chemical behaviour. (viii) Describe the different types of depth proﬁles that various metals exhibit and explain how differences in their proﬁles originate. (ix) Assuming all g 0 s ¼ 1, calculate the speciation of lead in seawater at 251C given the following values for stability constants for the chlorocomplexes (K1 ¼ 107.82, b2¼1010.88, b3 ¼ 1013.94, b4 ¼ 1016.30). Note that [Cl] is 0.559 mmol L1. (x) With speciﬁc reference to Question 9, outline the potential limitations of using equilibrium models to explain chemical behaviour of trace metals in seawater. (xi) Describe the classiﬁcation of marine sediments and give examples of each sediment type. (xii) Why do suspended particles exhibit a surface charge and how could this characteristic moderate the composition of ocean’s waters? (xiii) With respect to estuarine chemistry, describe conservative and non-conservative behaviour and provide examples of cationic and anionic species for each category. (xiv) What natural processes are responsible for weathering an oil slick? (xv) Why is tributyltin considered such a potent pollutant in the marine environment?

REFERENCES 1. R. Chester, Marine Geochemistry, Unwin Hyman, London, 1990, 698. 2. P. Tchernia, Descriptive Regional Oceanography, Pergamon Press, Oxford, 1980, 253.

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3. The Open University, Seawater: Its Composition, Properties and Behaviour, Open University, Oxford, 1989, 165. 4. K.S. Stowe, Ocean Science, Wiley, New York, 1979. 5. International Panel on Climate Change (IPCC), Climate Change 2001: Synthesis Report, WMO–UNEP, Third Assessment Report, vol 4, Geneva, 2001, 397. 6. D. Dyrssen and M. Wedborg, in The Sea, E. Goldberg (ed), Wiley, New York, 1974, 181. 7. P. Brewer, in Chemical Oceanography, 2nd edn, vol. 1, J.P. Riley and G. Skirrow (eds), Academic Press, London, 1975, 415. 8. K.W. Bruland, in Chemical Oceanography, vol. 8, J.P. Riley and R. Chester (eds), Academic Press, London, 1983, 157. 9. R.F. Weiss, Deep-Sea Res., 1970, 17, 721. 10. D. Kester, Chemical Oceanography, 2nd edn, vol 1, J.P. Riley and G. Skirrow (eds), Academic Press, London, 1975, 497. 11. W.S. Broecker and T.H. Peng, Tracers in the Sea, Lamont-Doherty Geological Observatory, Palisades, 1982, 690. 12. P.S. Liss, in Air–Sea Exchange of Gases and Particles, P.S. Liss and W.G.N. Slinn (eds), Reidel, Dordrecht, 1983, 241. 13. K.W. Bruland, Earth Planet. Sci. Lett., 1980, 47, 176. 14. W. Stumm and J.J. Morgan, Aquatic Chemistry, 3rd edn, Wiley, New York, 1996, 1022. 15. J. Lovelock, Gaia. A New Look at Life on Earth, Oxford University Press, Oxford, 1979. 16. R.J. Charlson, J.E. Lovelock, M.O. Andreae and S.G. Warren, Nature, 1987, 326, 655. 17. M.O. Andreae, in The Role of Air–Sea Exchange in Geochemical Cycling, P. Buat-Me´nard (ed), Reidel, Dordrecht, 1986, 331. 18. A.C. Redﬁeld, Am. J. Sci., 1958, 46, 205. 19. E. Goldberg, M. Koide, J.S. Yang and K.K. Bertine, Metal Speciation: Theory, Analysis and Applications, J.R. Kramer and H.E. Allen (eds), Lewis Publishers, Chelsea, 1988. 20. L. O¨hman and S. Sjo¨berg, in Metal Speciation: Theory, Analysis and Applications, J.R. Kramer and H.E. Allen (eds), Lewis Publishers, Chelsea, 1988, 1. 21. S.J. de Mora and R.M. Harrison, in Hazard Assessment of Chemicals, Current Developments, vol 3, J. Saxena (ed), Academic Press, London, 1984, 1. 22. P. Craig, Organometallic Compounds in the Environment, P.S. Craig (ed), Longman, Harlow, 1986, 1.

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23. R.M. Harrison and S.J. de Mora, Introductory Chemistry for the Environmental Sciences, 2nd edn, Cambridge University Press, Cambridge, 1996, 373. 24. E.K. Berner and R.A. Berner, The Global Water Cycle, PrenticeHall, NJ, 1987, 397. 25. U. Fo¨rstner and G.T.W. Wittman, Metal Pollution in the Aquatic Environment, 2nd edn, Springer, Berlin, 1981, 486. 26. G.M.E. Perillo, in Geomorphology and Sedimentology of Estuaries, G.M.E. Perillo (ed), Elsevier Science, Amsterdam, 1995, 17. 27. K. Hunter, Geochim. Cosmochim. Acta, 1983, 47, 467. 28. E. Goldberg, Mar. Pollut. Bull., 1992, 25, 45. 29. S. Murray, in Pollutant Transfer and Transport in the Sea, vol. 2, G. Kullenberg (ed), CRC Press, Boca Raton, 1982, 169. 30. G.S. Douglas, A.E. Bence, R.C. Prince, S.J. McMillen and E.L. Butler, Environ. Sci. Technol., 1996, 30, 2332. 31. I. Tolosa, S.J. de Mora, M.R. Sheikholeslami, J.-P. Villeneuve, J. Bartocci and C. Cattini, Mar. Pollut. Bull., 2004, 48, 44–60. 32. M.R. Gregory and P.G. Ryan, in Marine Debris: Sources, Impacts and Solutions, J.M. Coe and D.B. Rogers (eds), Springer, New York, 1996, 49. 33. S.J. de Mora (ed.), Tributyltin: Case Study of an Environmental Contaminant, Cambridge University Press, Cambridge, 1996, 301. 34. K. Kannan and J. Falandysz, Mar. Pollut. Bull., 1997, 34, 203. 35. C. Alzieu, M. Heral, Y. Thibaud, M.J. Dardignac and M. Feuillet, Rev. Trav. Inst. Marit., 1982, 45, 100. 36. C. Alzieu, Mar. Environ. Res., 1991, 32, 7. 37. C. Stewart and S.J. de Mora, Environ. Technol., 1990, 11, 565.

CHAPTER 5

The Chemistry of the Solid Earth IAN D. PULFORD Chemistry Department, University of Glasgow, Joseph Black Building, Glasgow, G12 8QQ, UK

5.1 INTRODUCTION Soil is the product formed when the rocks of the earth’s crust are exposed at the surface and are subjected to various physical, chemical, and, eventually, biological weathering processes. The minerals in these rocks are predominantly silicates, which dominate the characteristics of most soils. Table 1 shows those elements that are found in the crust above an average concentration of 1% and their corresponding soil content. The importance of aluminosilicates in soil is clear from the enrichment factors of approximately 1 for O, Si, and Al. Some loss occurs of K, Fe, Ca, Na, and Mg as a result of soil processes. But two elements, C and N, show considerable enrichment in soil because of the crucial role played by organic matter. Table 1

Average elemental concentrations in the earth’s crust and in soil, and the enrichment factor in soil

Element Oxygen Silicon Aluminium Iron Calcium Magnesium Sodium Potassium Carbon Nitrogen

Average concentration in crust Average concentration in soil Soil:crust (%) (%) ratio 47 28 8.2 4.1 4.1 2.3 2.3 2.1 0.05 0.0025

49 31 7.2 2.6 2.4 1.2 1.2 1.5 2.5 0.20

Source: adapted from Sposito, The Chemistry of Soils, Oxford University Press, 1989.

234

1.0 1.1 0.88 0.63 0.59 0.52 0.52 0.71 50 80

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Soil is formed not just by the weathering of minerals derived from rocks, but also by the input of organic matter from the decomposition of plants and, to a lesser extent, the decomposition and waste products of animals and microorganisms. The relative amounts of the mineral and organic components can vary widely, from almost all mineral materials, such as in desert sand, to highly organic soil such as peat. It is the combination of the mineral and organic components that gives soil its characteristics and makes it a valuable resource, which can be used to grow crops for a number of uses, such as food, ﬁbre, and energy. These two aspects – soil development and soil use – are reﬂected in two deﬁnitions of soil given by the Soil Science Society of America (http://www.soils.org/sssagloss/): The unconsolidated mineral or organic material on the immediate surface of the earth that serves as a natural medium for the growth of land plants.

The unconsolidated mineral or organic material on the surface of the earth that has been subjected to and shows the eﬀects of genetic and environmental factors of: climate (including water and temperature eﬀects), and macro- and microorganisms, conditioned by relief, acting on parent material over a period of time. A product – soil diﬀers from the material from which it is derived in many physical, chemical, biological, and morphological properties and characteristics.

The ﬁve points underlined in the second deﬁnition are known as the soilforming factors, which interact to produce a soil of particular characteristics in any given place. The parent material is acted upon by a series of weathering processes, outlined in Figure 1. Initially these are mainly physical processes, which bring about a decrease in particle size, and hence an increase in surface area. Weathering is strongly inﬂuenced by climate, and, once living organisms colonize the developing soil, by biological processes. Climate and relief interact to inﬂuence the movement of water and weathering products through the soil. Most of these physical, chemical, and biological processes proceed slowly, and so soil is evolving with time. Soil is a combination of all of the major components of the surface environment: the atmosphere, hydrosphere, lithosphere, and biosphere (Figure 2). It is the mix and interaction of these four components that result in the particular properties of a speciﬁc soil. The solid components (mineral and organic matter) make up approximately 50% of soil by volume. Air and water occupy the pore space between the solid phase, and their relative amounts can vary considerably, resulting in important eﬀects on the chemical and biological processes in a soil. Various measurements can be made to express the air–water balance in a soil (Table 2).

236

Chapter 5 Exposed, unweathered parent material – consolidated bedrock or unconsolidated material (e.g. glacial drift; alluvium)

Physical weathering of parent material decrease of particle size, but little or no chemical alteration

Chemical / biological weathering of parent material – formation of secondary minerals

Soil formation – incorporation of organic matter into surface layers soil forming processes in subsoil –weathering, leaching, etc.

slowly weathering parent material

Figure 1 Weathering sequence in soil Atmosphere

Biosphere Biomass Humus

Soil air

Soil Rocks and minerals

Lithosphere

Figure 2 Components of soil

Soil water

Hydrosphere

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The Chemistry of the Solid Earth

Worked example 5.1 – soil water and air contents, and bulk density A soil core of 5 cm diameter and 8 cm height contains 257 g of fresh soil. After drying the soil at 1101C, the soil weighed 196 g. Calculate (i) the gravimetric soil water content, (ii) the volumetric soil water content, (iii) the soil bulk density, (iv) the % pore space, (v) the % water-ﬁlled pore space, and (vi) the % air-ﬁlled pore space. Density of water ¼ 0.997 g cm3 Volume of soil core ¼ p.r2.h ¼ 3.14 (2.5)2 8 ¼ 157 cm3 (i) Gravimetric soil water content ¼ [(257–196 g)/196 g] 100 ¼ 31% (ii) Volumetric soil water content ¼ [(257–196 g)/(0.997 g cm3 157 cm3)] 100 ¼ 39% (iii) Soil bulk density ¼ 196 g/157 cm3 ¼ 1.25 g cm3 Assuming a soil particle density of 2.6 g cm3, the volume of particles in the soil core ¼ 196 g/2.6 g cm3 ¼ 75.4 cm3 (iv) % pore space ¼ [(157 cm375.4 cm3)/157 cm3] 100 ¼ 52% (v) % water-ﬁlled pore space ¼ volumetric water content ¼ 39% (vi) % air-ﬁlled pore space ¼ 52%39% ¼ 13% What is the weight of this soil in 1 ha to a depth of 20 cm? 1 ha ¼ 10,000 m2, therefore volume of soil to 20 cm ¼ 10,000 m2 0.2 m ¼ 2000 m3 Bulk density of soil ¼ 1.25 g cm3R1.25 t m3 2000 m3 1.25 t m3 ¼ 2500 t of soil.

Table 2

Determination of soil water and air contents and bulk density

Gravimetric soil water content (%)

(Mass of water/mass of oven dry soil) 100

Volumetric soil water content (%)

(Volume of water/volume of soil core) 100

Soil bulk density (g cm3)

Mass of oven dry soil/volume of soil core

Soil particle-speciﬁc gravity (g cm3)

Mass of oven dry soil solids/volume of soil solids

Pore space (%)

((Soil core volume–volume of soil particles in core)/soil core volume) 100

Solid material (%)

(Volume of soil particles in core/total core volume) 100

Water-ﬁlled pore space (%)

Volumetric soil water content

Air-ﬁlled pore space (%)

% pore space–% water-ﬁlled pore space

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Chapter 5

5.2 MINERAL COMPONENTS OF SOIL 5.2.1 Inputs The nature of the parent material is the most important factor inﬂuencing the mineral components of a soil. In particular, the textural (particle size) properties and inherent fertility are directly aﬀected by the types of rocks and minerals found in the parent material. Parent material can be consolidated rock (igneous, sedimentary, or metamorphic) or unconsolidated superﬁcial deposits that have been transported by some agency, such as wind, ice, or water. Igneous rocks are formed from molten magma from the earth’s crust and are the ultimate source of all rocks. Extruded igneous rock is formed when molten magma appears at the surface and cools quickly, resulting in small grain size (e.g. basalt). Intruded igneous rock cools slowly on its way to the surface and results in large grain size (e.g. granite). Sedimentary rocks are formed from the weathering products of igneous, metamorphic, or older sedimentary rocks. Metamorphic rocks are formed by changes in igneous or sedimentary rocks owing to high temperature and/or high pressure, which usually has the eﬀect of making the rock more resistant to weathering. Unconsolidated materials are surface deposits of partially weathered rock that have been transported and deposited in various ways. The most important groups of such materials, from which much of the world’s major soils are formed, are: glacial deposits, resulting from the action of ice; alluvium, deposited from water; and aeolian or wind-blown deposits.w 5.2.2 Primary Minerals The primary soil minerals, sometimes called ‘inherited’ minerals, are those that are derived from the rocks of the parent material. They are distinguished from the secondary, or ‘pedogenic’ minerals, which are formed during the process of soil formation (see next section). The main primary minerals are silicates (Table 3), all based on the SiO44 tetrahedron, in which three oxygen atoms are in a triangular plane below the silicon atom and one oxygen is above the silicon (Figure 3a). There may be isomorphous substitution of silicon by aluminium. The charge is balanced by cations such as K1, Na1, Ca21, Mg21, and Fe21. The silica tetrahedra can be linked in a number of ways to form silicates of diﬀering structures, and hence susceptibility to weathering w

A description of the origins and morphology of speciﬁc soil types is outwith the scope of this chapter. An excellent description of the major soils of the world can be found at http:// soils.usda.gov/education/.

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Table 3

Structures of common primary soil minerals

Quartz

SiO2

Feldspars Microcline Albite Anorthite

KAlSi3O8 NaAlSi3O8 CaAlSi2O8

Micas Muscovite Biotite

KAl2(Si3Al)O10(OH)2 K(Mg, Fe)2(Si3Al)O10(OH)2

Amphiboles Hornblende

(Na, Ca)2(Mg, Fe, Al)5(Si, Al)8O22(OH)2

Pyroxenes Enstatite Augite

MgSiO3 Ca(Mg, Fe, Al)(Si, Al)2O6

Olivine

(Mg, Fe)2SiO4

(a)

(b)

Silicon

Aluminium

Oxygen or hydroxyl

Figure 3 Structure of (a) silica tetrahedron and (b) alumina octahedron

(Figure 4). Quartz and the feldspars are framework silicates, in which adjacent tetrahedra share all four oxygens. The micas form sheet silicates by the sharing of three oxygens by each silica tetrahedron. The amphiboles and pyroxenes are chain silicates: amphiboles are double chains, with the tetrahedra sharing two or three oxygens alternately; pyroxenes are single chains with sharing of two oxygens. Olivine is an isolated tetrahedon (or orthosilicate).

240

(a)

Chapter 5

(b)

(c)

(d)

Figure 4 Structures of (a) ortho (isolated); (b) single chain; (c) double chain; and (d) sheet silicates

In Table 3, susceptibility to weathering increases down the list as fewer silicon–oxygen bonds need to be broken to release silicate. Consequently, quartz and feldspars especially, but also mica in temperate soils, are common inherited minerals in the coarse particle size fractions of soil (the silt and sand fractions, 0.002–2 mm). The amphiboles, pyroxenes, and olivine are much more easily weathered. Thus, soils derived from parent material with rock containing a predominance of framework silicates (e.g. granite, sandstone) tend to be more sandy, while those derived from rocks containing the more easily weathered minerals tend to be more clayey. The chemical composition of the silicate minerals aﬀects the inherent nutrient properties of soil. Coming down the list of silicates in Table 3, there is a change from the dominance of Si, Al, and Na in quartz and the feldspars to greater contents of Ca, Mg, and Fe in the amphiboles, pyroxenes, and olivine. These latter make up the so-called ‘ferromagnesian’ group of minerals, and soils in which they are an important component of the parent material have better inherent nutrient properties. In particular, such soils are better buﬀered against the natural acidiﬁcation process that occurs due to loss of exchangeable cations, especially Ca21 and Mg21, by leaching (see Sections 5.4 and 5.5).

5.2.3 Secondary Minerals Secondary minerals (which dominate the clay-size fraction of soil, o0.002 mm) are formed in soil by the action of various weathering processes on

241

The Chemistry of the Solid Earth Table 4

Structures of common secondary soil minerals

Clay minerals (phyllosilicates) Kaolinite Illite Vermiculite Smectite Chlorite

Al4Si4O10(OH)8 K1.5(Al3.5Mg0.5)(Si7Al)O20(OH)4 M0.7Al2(Al0.7Si3.3)O10(OH)2 M¼Ca21, Mg21, K1, etc. M0.3Al2(Al0.3Si3.7)O10(OH)2 M¼Ca21, Mg21, K1, etc. (Mg, Fe)3(Si, Al)4O10(OH)2(Mg, Fe)3(OH)6

Amorphous aluminosilicates Allophane Imogolite

xSiO2 Al2O3 yH2O (x¼0.8–2, y Z 2.5) SiO2.Al2O3.2.5H2O

Hydrous oxides Goethite Haematite Ferrihydrite Gibbsite Birnessite

FeOOH Fe2O3 Fe5O7(OH).4H2O Al(OH)3 (Na, Ca)(Mn31, Mn41)7O14.2.8H2O

Carbonates Calcite Dolomite

CaCO3 CaMg(CO3)2

Others Gypsum

CaSO4.2H2O

the primary minerals derived from the parent material (see section 5.2.4). These may be formed by alteration of a primary mineral – e.g. the conversion of mica into a clay mineral – or by reactions of soluble products released into the soil environment – e.g. precipitation of iron oxide. The most important groups of secondary minerals are: (a) the aluminosilicate clay minerals; (b) short-range order (amorphous) aluminosilicates; and (c) the hydrous oxides of Al, Fe, and Mn (Table 4). In the study of soil science, most attention has historically been paid to the aluminosilicate clays, which dominate the properties of temperate soils, the ﬁrst to be scientiﬁcally studied. More recently, the importance of the amorphous aluminosilicates has been shown in young soils, in soils derived from volcanic ash and in leached, acidic soil (e.g. podzols or spodosols). The hydrous oxides are especially important components of old, highly weathered soils, such as those found in the tropics (e.g. oxisols). This is an important distinction as the charge on the aluminosilicate clays is predominantly a permanent negative charge, while the amorphous aluminosilicates and hydrous oxides have a variable,

242

Chapter 5

pH-dependent charge. This results in very diﬀerent properties of tropical soils compared to temperate soils. 5.2.3.1 Aluminosilicate Clay Minerals (Phyllosilicates). The aluminosilicate clay minerals are sheet silicates. They are sometimes referred to as phyllosilicates, but strictly this term should also include the micas, which are primary minerals. The basic building blocks of all the aluminosilicate clay minerals are the silica tetrahedron (Figure 3a) and alumina octahedron (Figure 3b). These build up into tetrahedral sheets by the sharing of oxygens between adjacent silicons, and into octahedral sheets by the sharing of oxygens or hydroxyls between adjacent aluminiums. The sharing of oxygens between tetrahedral and octahedral sheets leads to the formation of clay unit layers, of which there are two basic types (Table 5): 1:1, one tetrahedral Si sheet and one octahedral Al sheet, and 2:1, two tetrahedral Si sheets sandwiching one octahedral Al sheet. Unit layers are held together in various ways to produce a clay crystal. These are regular, rigid systems and the distance between equivalent points in adjacent unit layers, the c spacing or basal spacing, can be measured by X-ray diﬀraction and is used to identify the clay minerals (Figure 5). Within the sheets, Si or Al may be replaced by a diﬀerent element by the process of isomorphous substitution. Common replacements in the clay minerals are Al31 for Si41 in the tetrahedral sheet and Mg21 or Fe21 for Al31 in the octahedral sheet. Substitution of an element by one of a lower valency results in a deﬁcit of positive charge, and hence the development of a permanent negative charge on the clay mineral. In soils, this charge is neutralized by cations in solution being attracted to the negatively charged surface, and is called the cation-exchange capacity Table 5

Components of soil clays and how they combine to form a clay mineral

Basic building blocks

Si tetrahedra, Al octahedra

Sheets

Tetrahedral sheet, made up of Si tetrahedra sharing oxygen atoms; octahedral sheet, made up of Al octahedra sharing oxygens or hydroxyls

Unit layers

Sheets combined in the ratio of one tetrahedral:one octahedral (1:1), or two tetrahedral:one octahedral (2:1)

Clay crystal

Stacks of unit layers held together by various bonding mechanisms, depending on the type of clay

243

The Chemistry of the Solid Earth Si tet. sheet

basal spacing or c spacing

Al oct. sheet

2:1 unit layer

Si tet. sheet

interlayer space

clay edge

clay face

Figure 5 Structure of a 2:1 aluminosilicate clay mineral

(CEC). This permanent negative charge is the major characteristic of temperate soils, in which the aluminosilicate clays dominate the reactive fraction. The consequences of this are discussed in relation to ionexchange processes in Section 5.5. Kaolinite is the commonest 1:1 clay mineral. There is a small amount of isomorphous substitution of Al for Si in the tetrahedral sheet, resulting in a charge of r 0.005 mol negative charge per unit cell. The unit layers are bound by hydrogen bonds between the oxygen atoms on the tetrahedral face and the hydroxyls on the adjacent octahedral face, resulting in a low CEC and surface area (Table 6) and a ﬁxed c spacing of 0.7 nm. Unit layers stack up to form hexagonal crystals typically 0.05–2 mm thick. This means that the variable, pH-dependent charge on the clay edge is relatively more important in kaolinite than the other aluminosilicate clays. Illite (a hydrous mica) is a 2:1 clay mineral with isomorphous substitution occurring mainly in the tetrahedral sheet (Al31 for Si41), although a small amount also occurs in the octahedral sheet (Mg21 or Fe21 for Al31), resulting in a charge of 1.5 mol negative charge per unit cell. Most of this charge is neutralized by K1 ions in the interlayer space. The geometrical arrangement of oxygen atoms on the tetrahedral face allows the K1 ions to sit very close to the clay surface, forming a very strong bond between unit layers. This results in a low CEC and surface area (Table 6) and a ﬁxed c spacing of 1.0 nm. Vermiculite, a 2:1 clay mineral, is also a hydrous mica, with isomorphous substitution in the tetrahedral sheet, resulting in a charge of between 1.2 and 1.9 mol negative charge per unit cell. In this case the

244 Table 6

Chapter 5 Properties of the common aluminosilicate clay minerals

Clay mineral

CEC (cmolc Type kg1)

Surface area (m2 g1)

Kaolinite

1:1

3–20

5–100

Illite

2:1

10–40

100–200

Vermiculite 2:1

80–150

300–500

Smectite

2:1

80–120

700–800

Chlorite

2:1:1 10–40

300–500

Expanding/ c spacing Interlayer non-expanding nm binding Nonexpanding Nonexpanding Slight expansion Fully expanding Nonexpanding

0.7 1.0 1.4 1.4–1.9 1.4

Hydrogen bonds K1 ions K1, Ca21, Mg21 ions Interlayer cations Brucite sheet and H bonds

interlayer cations are mainly Mg21 with some Ca21, which are more strongly hydrated cations than K1 (see Section 5.5). The water of hydration associated with these cations widens the c spacing to 1.4 nm, and allows some degree of expansion. Smectites are a group of 2:1 clay minerals that have a low degree of isomorphous substitution and hence a low-layer charge. The commonest smectite found in soil is montmorillonite, in which substitution occurs in the octahedral layer giving 0.7 mol negative charge per unit cell. Because the charge originates further from the clay surface, interlayer bonding is weak, allowing expansion of the clay lattice and easy entry of cations and water molecules into the interlayer space. Therefore, the CEC and surface area are both high (Table 6). The c spacing is variable, up to 1.9 nm when Ca21 or Mg21 are the dominant interlayer cations. If sodium is the dominant cation, the smectites can be fully dispersed in suspension. Chlorite is a 2:1:1 clay mineral, with the usual 2:1 unit layer of two tetrahedral sheets and one octahedral held together by a sheet of brucite. Brucite is a magnesium hydroxide in which about two thirds of the Mg21 ions are substituted by Al31, resulting in a positive charge. This neutralizes much of the negative charge that arises due to isomorphous substitution in the tetrahedral sheets (2 mol negative charge per unit cell) and forms a strong electrostatic bond between unit layers. There is also hydrogen bonding between hydroxyls on the brucite sheet and oxygens on adjacent tetrahedral sheets. As a result of this strong interlayer bonding, the CEC and surface area are low, and the c spacing is ﬁxed at 1.4 nm. In reality, soil clays are often not so well deﬁned. Interstratiﬁcation is common, with, for example, intergrades of vermiculite and illite or smectite and illite being found. This means that some layers in the clay

245

The Chemistry of the Solid Earth

mineral are illitic in nature, while others more closely resemble vermiculite or smectite. Weathering at the edges of the clay mineral can also lead to changes. For example, at the edge of an illite crystal, there may be some loss of interlayer K1 ions and their replacement by other cations. This results in a widening of the interlayer space to form the socalled wedge site (Figure 6). 5.2.3.2 Short-Range Order (Amorphous) Aluminosilicates. This is a group of minerals that have small particle size and no regular crystal structure, and as a result produce poorly deﬁned X-ray diﬀraction patterns. They do, however, show some regular short-range structure. Thus, there has been a tendency over recent years to describe them as short-range order, rather than amorphous, aluminosilicates. They have a very high surface area, making them highly reactive. They have both a

K+

Kaolinite Illite

Mg2+ H2O

H2O Ca2+

Mg2+

H2O

Ca2+

H 2O

Mg2+

Vermiculite Smectite

Chlorite

Figure 6

Structures of common clay minerals octahedral sheet; brucite sheet.

Silica tetrahedral sheet;

alumina

246

Chapter 5

ﬁxed negative charge, arising from isomorphous substitution in Si tetrahedral sheets, and a variable, pH-dependent charge similar to that on clay edges and hydrous oxide surfaces. The two main minerals in this group, allophane and imogolite, are found in young soils and soils of volcanic origin. 5.2.3.3 Hydrous Oxides. This term is generally taken to include the oxides, hydroxides, and oxyhydroxides of aluminium, iron and manganese, which form in soil when these elements are released from primary minerals by weathering. They exist mainly as small particles in the claysized fraction of a soil (o2 mm), and also as coatings on other soil minerals or as components of larger aggregates. The most common aluminium oxide found in soil is gibbsite (gAl(OH)3), which may be the ﬁnal product resulting from the precipitation of amorphous aluminium hydroxide. A number of iron oxides are found in soil, the commonest being ferrihydrite, goethite, and haematite. Iron released from weathering of the ferromagnesian minerals precipitates out of solution as either ferrihydrite or goethite depending on soil conditions (low organic matter and a high rate of Fe release favour ferrihydrite formation). Ferrihydrite is a poorly ordered oxide of small particle size, whereas goethite is a well-structured mineral and is the commonest of the iron oxides. Haematite is derived by structural changes to ferrihydrite, and is commonly found in tropical soils. Manganese oxides occur in a large number of forms, often with variable valency, for example, birnessite, and in mixed oxides, especially with iron. They are the least well understood of this group of minerals. The hydrous oxides have a variable, pH-dependent charge and are important in soils as aggregating agents and as adsorptive surfaces (see Section 5.5) (Figure 7). 5.2.4 Weathering Processes (See Also Chapter 3) The processes that cause changes to the rocks and minerals of the soil parent material are collectively known as weathering. In the early stages, physical weathering dominates and results in fragmentation of rock, a decrease in particle size, and increase in surface area, but no chemical change. It is brought about by a number of actions: Abrasion of the rock caused by the action of ice, water, and wind leads to physical disintegration. Temperature change due to diurnal temperature variation and by freeze–thaw cycles can create signiﬁcant pressures that will cause rock to split along lines of weakness.

247

The Chemistry of the Solid Earth +

Fe - OH2 O

Fe - OH O

Fe - OH2

Iron oxide surface below pzc Overall positive charge

Fe - OH O

Fe - OH2

pzc = point of zero charge pH at which net surface charge is zero

Fe - OH

Iron oxide surface above pzc Overall negative charge

Figure 7 Variable, pH-dependent charge on an iron oxide surface

Roots tend to grow into existing ﬁssures and channels and can exert signiﬁcant pressure. Crystallization of some salts can also cause swelling, which can exert pressure. The key factor about physical weathering is the increase in surface area of rock that is then available for attack by various chemical-weathering processes. Chemical weathering is the reaction with natural waters, and dissolved components, resulting in partial or complete dissolution of the rock and formation of a new mineral phase. In most cases, it is the interaction of water with the minerals in the rock that is the crucial issue. 5.2.4.1 Dissolution. Some minerals, such as gypsum (CaSO4.2H2O), are readily soluble in water and so can persist only in soils of arid and semi-arid regions, where water is limiting. In temperate and tropical regions suﬃcient water will pass through the soil to cause dissolution and loss of gypsum: CaSO4.2H2O ! Ca21 þ SO2 4 þ 2H2O

(5.1)

Quartz is a much less soluble mineral than gypsum, but even this is slightly soluble: SiO2 þ 2H2O " H4SiO4

(5.2)

This reaction is driven by the loss of the soluble silicic acid from the soil. In temperate regions insuﬃcient water passes through the soil, and weathering has proceeded over a relatively short time since the last ice

248

Chapter 5

age, for there to have been signiﬁcant loss of silica, and so quartz is a dominant mineral in the coarse fraction of soils. In the humid tropics, however, the much larger volume of water passing through the soil, over a much longer timescale, has resulted in signiﬁcant, if not complete, loss of silica. 5.2.4.2 Acid Hydrolysis. The water that enters soil as rain or snow is in equilibrium with CO2 in the atmosphere, which dissolves to form carbonic acid. Unpolluted rainwater has a pH of approximately 5.7, whereas water in soil pores may be exposed to air containing a higher partial pressure of CO2 than the free atmosphere, and hence soil water may be more acidic (see Section 5.4). It is the attack on soil minerals by this weak carbonic acid that is the major chemical weathering process in most soils. For example, acid hydrolysis of calcium carbonate yields calcium and bicarbonate ions: CaCO3 þ H2CO3 " Ca21 þ 2HCO 3

(5.3)

Acid hydrolysis of the primary mineral microcline feldspar results in release of some soluble components (silicic acid, potassium ions, and bicarbonate ions) and alteration of the solid phase to kaolinite: 4KAlSi3O8 þ 22H2O þ 4CO2 " Al4Si4O10(OH)8 þ 8H4SiO4 þ 4K1 þ 4HCO 3

(5.4)

5.2.4.3 Oxidation. For those elements that can exist in more than one valence state oxidation, and indeed reduction, may be a major reaction in the chemical-weathering process. Iron and manganese are the most important elements that behave in this way. For example, iron in the ferromagnesian minerals is in the Fe(II) state, which is oxidized to Fe(III) when released from the mineral. This can cause changes in the charge balance, requiring other ions to be lost. Formation of iron oxide can cause physical disruption to the mineral. 5.2.4.4 Chelation. Once biological activity is established, an important reaction is the formation of a stable complex between a metal ion and an organic molecule due to release of organic compounds by colonizing organisms and plant roots. 5.3 ORGANIC COMPONENTS OF SOIL It cannot be emphasized enough that soil is formed by the interaction of the mineral weathering product of rocks and the organic material

249

The Chemistry of the Solid Earth

introduced as a result of biological activity. Much of this organic material is plant-derived litter and the amount added to soil per year depends very much on the dominant vegetation type (Table 7). These ﬁgures do not include inputs from roots, which may be signiﬁcant but diﬃcult to quantify, or more resistant inputs, such as wood in forests. In addition, animal and microbial excretion and decomposition must be taken into account. The values in Table 7 do, however, give a reasonable estimate of what may be described as degradable carbon entering soils per year. The organic materials added to soil have varying degrees of susceptibility to degradation. Simple compounds, such as sugars and amino acids, are rapidly decomposed by soil microorganisms, whereas complex material, such as lignin and hemicellulose, tends to persist in an undegraded or partially degraded form. About 70% of the carbon added to soil is lost as carbon dioxide, produced mainly by microbial respiration. The other 30% is incorporated into the microbial biomass and soil humus, a resistant material, which degrades only very slowly. Figure 8 shows typical values for the annual turnover of organic matter in soil. Humus formation proceeds broadly along two pathways (Figure 9). Readily decomposed material releases low molecular weight compounds, especially phenolic and amino compounds, which react and polymerize to form humic material. Alternatively, the more resistant organic material can undergo partial degradation, and this partially degraded material may also react with the low molecular weight compounds released into the soil, again leading to the formation of humus. It is likely that both of these processes operate in soils. The terms humus and humic substances are used interchangeably to refer to the resistant organic material that persists in soil over relatively long timescales (100s–1000s of years). Although the humic substances may be chemically diverse, being dependent on the dominant type of plant tissue input, they do have some overall similarities. They are amorphous, colloidal polymers built up mainly from aromatic units. There is a wide range of molecular weights, from 100s to 100,000s Da. Table 7

Inputs of carbon to soil by various types of land use

Dominant land use

Tonnes carbon added to soil ha1 year1

Alpine and arctic forest Arable agriculture Coniferous forest Deciduous forest Temperate grassland Tropical rain forest

0.1–0.4 1–2 1.5–3 1.5–4 2–4 4–10

Source: adapted from White, Principles and Practice of Soil Science, Blackwell Science, 1997.

250

Chapter 5 Input of plant litter 4.0 t C ha-1 (100%) Lost from soil as CO2 2.8 t C ha-1 (70%)

Incorporated into new humus 1.2 t C ha-1 (30%)

Reservoir of C in soil humus 72 t C ha-1

Loss of CO2 1.2 t C ha-1

Figure 8 Typical annual turnover of organic matter in a temperate grassland soil at steady state receiving 4 t C ha1

All of these compounds do, however, have similar overall properties due to the preponderance of carboxylic and phenolic functional groups on the humic polymer. Both of these groups can ionize by losing a hydrogen ion to give a negatively charged group. This is a variable charge, the magnitude of which varies with pH, but is always negative and forms part of the CEC of a soil (see Section 5.5). R-COOH " R-COO þ H1

(5.5)

251

The Chemistry of the Solid Earth Input of plant residues

Attack by microorganisms and soil fauna

Resistant materialpoorly degraded

Partially degraded material

Easily degraded material

Reaction of simple organics with partially degraded material

Loss of CO2 , simple inorganic ions (NO3etc.) and release of simple organic compounds, e.g. amino acids, sugars and polyphenols Microbial activity

Reaction between simple organic compounds to form more complex polymers

Soil humic material

Figure 9 Possible pathways for formation of humic material in soil from inputs of plant residues

O− + H +

ð5:6Þ

The pK of Equation (5.5) is in the range 3–5, and of Equation (5.6) in the range 7–8. Thus, carboxyl groups are important in acid soils, while phenolic groups become important above pH 7.

252

Chapter 5

Soil organic matter has been traditionally fractionated into a number of subgroups (Figure 10). Humin and humic acid are high molecular weight fractions (410,000 Da) and it is thought that humin is not extracted by alkali because it is too strongly bound to the soil mineral component. Fulvic acid is the low molecular weight fraction (o10,000 Da). There are some chemical diﬀerences between humic acid and fulvic acid, the latter being more acidic and having a higher oxygen content and lower carbon content than humic acid. This reﬂects the greater number of carboxylic and phenolic groups in the fulvic acid fraction. It is speculated that humic acid forms the less reactive backbone or core of the humic material, while the more reactive fulvic acid is present as side chains branching from the core. Although there is considerable variation between soils, the ratio of the major elements, C:N:P:S, in humic material is approximately 100:10:2:1, forming a major reservoir of N, P, and S that is made available to plants as soil organic matter, is broken down by soil micoorganisms.

Soil organic matter

Floatation

Sinks

Floats

Humic material

Non-humic material

Extraction with alkali Insoluble

Soluble

Humin

Humic colloids

Acidify to pH1

Insoluble Humic acid

Figure 10

Fractionation scheme for soil organic matter

Soluble Fulvic acid

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253

Identiﬁcation of N-, P-, and S-containing compounds is diﬃcult, and those that are identiﬁed tend to be metabolic products, such as nucleotides and vitamins, released into the soil following the death of cells. Broadly, nitrogen is an integral part of the humic molecule, and is released as NH41 ions when the humic material is degraded. Phosphorus and sulfur are more commonly found as P and S esters, which can be released as orthophosphate and sulfate ions by the action of phosphatase and sulfatase enzymes, respectively.

Worked example 5.2 – organic carbon content of soil Organic carbon in soil can be measured by oxidation with acid dichromate solution: 2K2Cr2O7 þ 3C þ 8H2SO4 " 2Cr2(SO4)3 þ 3CO2 þ 2K2SO4 þ 8H2O

(5.7)

The excess dichromate is measured by titration with ferrous sulfate 21 þ 14H1 " 2Cr31 þ 6Fe31 þ 7H2O Cr2O2 7 þ 6Fe

(5.8)

0.20 g soil was oxidized with 20 cm3 of M/6K2CrO7 in sulfuric acid. The excess dichromate was determined by titration with 0.5 M FeSO4 solution. A blank, no soil, control was carried out to assess loss of dichromate by thermal decomposition. The titration values obtained were 17.5 cm3 for the soil sample and 39.0 cm3 for the blank. What is the % organic carbon in the soil? Moles of Cr2O72 added initially ¼ 20 1/6 ¼ 3.334 mmol Moles of Cr2O72 remaining after reaction: 17.5 0.5 ¼ 8.75 mmol Fe21 From Equation (5.8), 1 mol of Cr2O72 reacts with 6 mol of Fe21 Therefore, 8.75/6 ¼ 1.458 mmol of Cr2O72 unused. Blank titration 39.0 0.5 ¼ 19.5 mmol Fe21 ¼ 3.25 mmol of Cr2O72 Therefore, 3.334–3.25 ¼ 0.084 mmol of Cr2O72 lost by thermal decomposition. Moles of Cr2O72 used in oxidation ¼ 3.334–1.458 ¼ 1.876 mmol Subtract Cr2O72 lost by thermal decomposition ¼ 1.876–0.084 ¼ 1.792 mmol of Cr2O72 From Equation (5.7) 2 mol of Cr2O72 react with 3 mol of C 1.792 3/2 ¼ 2.688 mmol C 2.688 12 ¼ 32.26 mg carbon (32.26 mg/200 mg) 100 ¼ 16.13% C in the soil.

254

Chapter 5

5.4 SOIL pH AND REDOX POTENTIAL Soil pH, which is a measure of the acidity or alkalinity, and redox potential, a measure of the aeration status, are the two crucial factors that control many of the chemical and biological processes in soil. 5.4.1 pH and Buﬀering pH is deﬁned as log[H1], i.e. the negative logarithm of the hydrogen ion concentration (or strictly speaking its activity) in solution, and is based on the concept of the partial ionization of water. In pure water, the H1 ion concentration is 107 mol L1 and so the pH is 7. If the concentration of H1 ions increases, the pH falls below 7 (acid conditions) and if the concentration decreases, the pH rises above 7 (alkaline conditions). In soil, however, factors other than the H1 ion concentration in solution aﬀect the pH. Measurement of pH in a soil is usually done using a glass pH electrode in a soil suspension. It is important to state the ratio of the soil weight to the solution volume used, because pH varies with dilution. Common ratios are 1:1 (10 g:10 cm3), 1:2.5 (10 g:25 cm3), and 1:5 (10 g:50 cm3). It is also necessary to state the nature of solution used, commonly deionized water or a dilute salt solution (e.g. 0.01 M CaCl2). What is actually being measured is the pH of a solution in equilibrium with negatively charged soil particles – thus, some H1 ions are in solution and some on soil surfaces (see ion exchange, Section 5.5). The pH measured depends on the ratio of these two pools of ions, and is aﬀected by dilution and the concentration of other ions in solution. If care is taken in measuring pH, it is a very useful parameter from which a lot can be deduced about a soil. However, it must be borne in mind that in the ﬁeld soil pH can vary signiﬁcantly over short distances and over relatively short timescales. H1 and Al31 ions in solution and on surface-exchange sites constitute the solution and exchangeable acidity in a soil, also called the active acidity. The non-exchangeable, or reserve, acidity is associated with carboyxl and phenolic groups on humic material, OH groups on hydrous oxide surfaces and clay edges, and polymeric aluminium hydroxides. The balance between solution and exchangeable acidity and non-exchangeable acidity can shift in response to a change in pH: Non-exchangeable acidity " Exchangeable and solution acidity (5.9) If soil pH is increased (H1 in solution decreased), the equilibrium shifts to the right – some non-exchangeable acidity released. If soil pH is

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255

decreased (H1 in solution increased), the equilibrium shifts to left – some acidity goes on to non-exchangeable acidity. This represents the buﬀer capacity of a soil, and acts to avoid large changes in pH. Acidity in soil is normally corrected by addition of a liming material such as CaCO3. The amount required can be calculated by measuring the buﬀering of a soil by titration with a dilute solution of an alkali (often Ca(OH)2 is used). The amount of alkali needed to reach the required pH can be found from the buﬀer curve.

Worked example 5.3 – soil buﬀering and liming To raise the pH of 5 g of the soil in Worked example 5.1 to 6.5, required addition of 4 cm3 of 0.005 mol L1 Ca(OH)2 solution. How much calcium carbonate would be required to be added to 1 ha? (Mol wt CaCO3 ¼ 100). 4 cm3 of 0:005 mol L1 ¼ 0:02 mmol CaðOHÞ2 solution for 5 g soil ¼ 0:004 mmol g1 ¼ 0:4 mg CaCO3 g1 ð1 mmol CaCO3 ¼ 100 mgÞ ¼ 0:4 kg CaCO3 t1 From Worked example 5.1, there are 2500 t ha1 in the top 20 cm, which approximates to rooting depth. 0.4 kg CaCO3 t1 2500 t ¼ 1000 kg CaCO3 required ha1.

5.4.2 Soil Acidity The atmosphere is a major source of soil acidity. Even in unpolluted environments rainwater is slightly acidic, having a pH of about 5.7 due to the dissolution of atmospheric CO2 to form the weak carbonic acid (see Worked example 5.4). The CO2 concentration in the partially enclosed soil pore system can be signiﬁcantly higher (typically up to about 10 times) than in the free atmosphere due to respiration of soil microorganisms and plant roots. This results in a lower pH. In areas aﬀected by industrial pollution, sulfur dioxide and nitrogen oxides dissolve in rainwater to produce sulfuric and nitric acids (acid rain), which are both strong acids and cause even more acidity.

256

Chapter 5

Worked example 5.4 – pH of soil solution Carbon dioxide dissolves in water to form carbonic acid, which is a weak acid. CO2 þ H2O " H2CO3

K ¼ 101.5

(5.10)

[H2CO3]/pCO2 ¼ 101.5 mol L1 atm1 Carbonic acid dissociates to produce bicarbonate and hydrogen ions H2CO3 " HCO3 þ H1 K ¼ 106.4

(5.11)

þ ½HCO 3 ½H ¼ 106:4 molL1 ½H2 CO3 þ ½HCO 3 ¼ ½H

½H þ 2 ¼ 106:4 ½H2 CO3 ¼ 106:4 101:5 pCO2 pCO2 in the atmosphere ¼ 103.5 atm ½Hþ 2 ¼106:4 mol L1 101:5 mol L1 atm1 103:5 atm ¼ 1011:4 ½Hþ ¼105:7 mol L1 pH ¼5:7 If the partial pressure of CO2 in the soil atmosphere is considered to be 10 times greater than the free atmosphere, ½Hþ 2 ¼106:4 mol L1 101:5 mol L1 atm1 102:5 atm ¼ 1010:4 ½Hþ ¼105:2 mol L1 pH ¼5:2 Sulfur dioxide dissolves to form the strong acid, sulfurous acid, so is strongly acidifying at very low concentrations. For example, for pSO2 ¼ 107 atm (a value much higher than concentrations now occurring owing to air pollution control measures in developed countries. Note, however, that low-pH rainfall still occurs as a result of incorporation of sulfuric and nitric acids). SO2 þ H2O " H2SO3

K¼1

[H2SO3]/pSO2 ¼ 1 mol L1 atm1

(5.12)

The Chemistry of the Solid Earth 1 H2SO3 " HSO K ¼ 1.7 102 3 þ H

257

(5.13)

þ ½HSO 3 ½H ¼ 1:7 102 mol L1 ½H2 SO3

½Hþ 2 ¼1:7 102 molL1 1molL1 atm1 107 atm ¼ 1:7 109 ½Hþ ¼4:12 105 log ½Hþ ¼ 4:4 pH ¼4:4 Various soil processes also contribute to soil acidity. Decaying organic matter releases a number of organic acids, as indeed do plant roots. Nitriﬁcation, the microbial oxidation of ammonium ions to nitrate by Nitrosomonas and Nitrobacter, occurs in slightly acidic to neutral soils and releases H1 ions: 1 2NHþ 4 þ 3O2 - 2NO2 þ 4H þ 2H2O

(5.14)

2NO 2 þ O2 - 2NO3

(5.15)

Sulﬁde oxidation, another microbially mediated process, also results in the production of acidity: 2FeS2 þ 2H2O þ 7O2 " 2FeSO4 þ 2H2SO4

(5.16)

4FeSO4 þ 10H2O þ O2 " 4Fe(OH)3 þ 4H2SO4

(5.17)

Note that in this case much of the acidity is due to the precipitation of iron(III) oxide. Leaching, the passage of water through the soil proﬁle, is a process that occurs in soils in areas of moderate to high rainfall, and which results in the loss of exchangeable cations (such as Ca21, Mg21, K1, and Na1 – see Section 5.5). Unless the rate of weathering of soil minerals is suﬃciently rapid to replace these losses, the exchange sites will become occupied by H1 and Al31, which are acidic cations. 5.4.3 Soil Alkalinity (see also Section 3.2.4) Alkaline soils (pH 4 7) contain solid phase carbonate, and bicarbonate is the dominant anion in solution. Calcium and magnesium carbonates

258

Chapter 5

produce a soil of pH between 7 and 8.5, depending upon the concentrations of CO2 and Ca21 or Mg21 ions. In arid regions, the lack of rainfall, and hence leaching, allows the more soluble Na2CO3 to accumulate and the pH rises to about 10.5. 5.4.4 Inﬂuence of pH on Soils A major inﬂuence of pH on soil is its eﬀect on biological activity. Many of the soil microorganisms function within a narrow optimal range of pH and their activity is inhibited in more acidic or alkaline conditions. For example, the nitrifying bacteria mentioned above, Nitrosomonas and Nitrobacter, have a pH optimum in the range 6–8, and are severely inhibited below pH 5.5. Conversely, the iron- and sulfur-oxidizing bacteria of the genus Thiobacillus are active only under acid conditions, and are inhibited at pH 4 5. In general, bacteria are less tolerant of acid conditions than fungi. Soil animals too are aﬀected by pH; earthworms, for example, cannot survive below a pH of about 4.5. Certain soil components have a variable, pH-dependent charge. The charge on the humiﬁed soil organic matter is negative overall due to dissociation of carboxyl and phenolic groups, but its magnitude varies, being greater at high pH (see Section 5.3). The hydrous oxides and edges of clay minerals are positively charged at low-pH values and negative at

Relative availability of nutrient

Ca, Mg

N, S

Mn, Fe, Cu, Zn 4.0

8.0

Figure 11 Relative availabilities of plant nutrients in soil as affected by pH (only relative amounts of individual nutrients should be compared)

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259

high pH (see Section 5.2). In general, therefore, low pH results in a lessening of the negative charge on the surfaces of soil particles. pH has a big inﬂuence on the solubility of many important elements in soil, and can have a big inﬂuence on the availability of plant nutrients (Figure 11). Aluminium, which is not required by plants, is relatively insoluble except under acid conditions when toxicity problems may occur. Using gibbsite, Al(OH)3, which is a common Al mineral in many soils, as an example, the eﬀect of pH on the concentration of aluminium in solution can be modelled.

Worked example 5.5 – eﬀect of pH on concentration of aluminium in soil solution What is the concentration of Al31 ions in soil solution at the pH values calculated in Worked example 5.4 for soil solution at (a) pCO2 ¼ 102.5 atm and (b) pSO2 ¼ 107 atm? Using the dissolution of gibbsite as a model, Al(OH)3 " Al31 þ 3OH Solubility product, KSP ¼ 5 1033 ½Al3 þ ½OH 3 ¼ 5 1033 ðaÞ At pCO2 ¼ 102:5 atm pH ¼ 5:2; pOH ¼ 8:8 ½OH ¼1:58 109 ½OH 3 ¼3:94 1027 ½Al3 þ ¼ð5 1033 Þ=ð3:94 1027 Þ ¼1:27 106 molL1 ðbÞ At pSO2 ¼ 107 atm pH ¼ 4:4; pOH ¼ 9:6 ½OH ¼2:51 1010 ½OH 3 ¼1:58 1029 ½Al3 þ ¼ð5 1033 Þ=ð1:58 1029 Þ ¼3:2 104 molL1 :

(5.18)

260

Chapter 5

5.4.5 Redox Potential Redox potential is the measure of the oxidation–reduction state of a soil, and is determined by redox reactions involving the transfer of electrons from one chemical species to another. A generalized redox reaction can be written as Ox þ mH1 þ ne " Red

(5.19)

where ‘Ox’ and ‘Red’ are the oxidized and reduced species, respectively. Note also that H1 ions are involved in the reaction, and so pH aﬀects the redox potential. The potential produced is a consequence of the ratio of oxidized to reduced species, and is expressed by the Nernst equation

E ¼E 0

RT ½Red ln nF ½Ox

ð5:20Þ

where E ¼ potential in volts E0 ¼ standard electrode potential R ¼ universal gas constant (8.314 J mol1 K1) T ¼ absolute temperature in K n ¼ number of electrons involved F ¼ Faraday constant 96,487 C mol1 (the charge when 1 mole of [Ox] is reduced). If a temperature of 251C (298 K) is assumed, the values of the constants R and F are used and natural logs converted to log10, the equation becomes E ¼E 0

0:0591 ½Red log10 n ½Ox

ð5:21Þ

Redox potential is measured using an inert platinum electrode, which acquires the electric potential (EH) of the soil when placed into the soil or a soil suspension. This potential is measured relative to a reference electrode with a known potential, such as the calomel electrode (Ecal), which has a potential of 0.248 V at 251C. This potential has to be added to the measured value EH ¼ Ecal þ 0.248 V

(5.22)

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The Chemistry of the Solid Earth

EH is measured at the pH of the sample, but is often expressed corrected to pH 7 by subtracting 0.0591 V per unit pH up to 7 for samples with a pH below 7 and adding 0.0591 V per unit pH down to 7 for samples above pH 7. EH can also be expressed as pe, which is the negative log of the electron activity (analogous with pH) (see also Chapter 3). This approach considers electrons as a reactant or product in the reaction. The two are related by EH (in volts) ¼ 0.0591 pe

(5.23)

Worked example 5.6 – redox limits in aqueous systems (see also Section 3.2.4.2) The limits of redox potential in aqueous systems are determined by the oxidation or reduction of water. Oxidation of water 2H2O " O2 þ 4H1 þ 4e

log K ¼ 83.1

(5.24)

log K ¼ log pO2 þ 4log H þ þ 4log e 83:1 ¼ log pO2 þ 4logHþ þ 4loge 83:1 ¼ log pO2 þ 4pH þ 4pe For pO2 ¼ 1atm; log pO2 ¼ 0 83:1 ¼ 4pH þ 4pe pe ¼ 20:78 pH Reduction of water 2H1 þ 2e " H2

log K ¼ 0

(5.25)

pK ¼ log pH2 2pH 2pe For pH2 ¼ 1atm; log pH2 ¼ 0 pe ¼ pH At pH 7, this gives theoretical limits for the soil environment at EH ¼ approximately þ0.8 V (oxidizing) and approximately 0.4 V (reducing). In practice soil atmospheres will not contain 1 atm of oxygen or hydrogen.

5.4.6 Reduction Processes in Soil Reduction is caused by a decrease in oxygen concentration in the soil atmosphere; for example, by waterlogging or compaction. Oxygen

262

Chapter 5

diﬀuses through water 104 times more slowly than through air, and so oxygen concentration falls. The microbial population changes in response to the decreased oxygen, from aerobic organisms, requiring oxygen for respiration, to facultative anaerobes, which can use sources of electrons other than oxygen, to obligate anaerobes, which cannot survive if oxygen is present. As the microbial population changes, their use of alternative electron sources to oxygen causes a series of important redox changes in soil (Table 8). As each of these reduction processes occurs, the EH is buﬀered, or poised, at a particular value. So, for example, while there is nitrate present, the system will be poised at around þ0.22 V; once the nitrate has been used up, reduction of MnIV will poise the system at around þ0.2 V, and so on. EH in oxidized systems is diﬃcult to measure because there is not a speciﬁc reaction controlling it. Generally, when oxygen is present, the EH is greater than þ0.3 V. The phase changes that occur are important: the gaseous products can be lost from the soil to the atmosphere; Mn and Fe become soluble and can be moved within or out of the system. Reduction of NO3, MnIV, and FeIII can occur under moderately reduced conditions, such as intermittently waterlogged soil or sediment. Reduction of nitrate causes the loss of a major nutrient. Often the reduction does not proceed through to the formation of N2 gas, and nitrous oxide (N2O) is formed, which is a potent greenhouse gas and can Redox reactions in soil

Table 8

Redox reaction þ

Range of soil EH values (V)

O2 þ 4H þ 4e Ð 2H2 O þ 2NO 3 þ 12H þ 10e Ð N2 þ 6H2 O solution gas

Disappearance of oxygen 0.6–0.4 Disappearance of nitrate 0.5–0.22

MnO2 þ 4Hþ þ 2e Ð Mn2þ þ 2H2 O solid solution

Appearance of MnII 0.4–0.2

FeðOHÞ3 þ 3Hþ þ e Ð Fe2þ þ 3H2 O solid solution

Appearance of FeII 0.3–0.1

þ SO2 4 þ 10H þ 8e Ð H2 S þ 4H2 O solution gas

CO2 þ 8Hþ þ 8e Ð CH4 þ 2H2 O gas gas

Disappearance of sulfate 0 to 0.15 Appearance of methane 0.15 to 0.22

Source: adapted from Bohn, McNeal and O’Connor, Soil Chemistry, Wiley Interscience, 1985.

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cause ozone depletion in the atmosphere. Redox changes of Mn and Fe are the main factor in the formation of gleys (waterlogged soils). Reoxidation to MnIV and FeIII results in precipitation of highly reactive oxides, which can sorb signiﬁcant amounts of anions such as phosphate and heavy metals. Sulﬁde and methane production occur only under highly reduced conditions, such as permanently waterlogged and peat soils. Sulﬁde may be lost as hydrogen sulﬁde gas, but often metals sulﬁdes (e.g. iron sulﬁde (FeS)) are precipitated. Methane is a potent greenhouse gas. 5.5 CHEMICAL REACTIONS IN SOIL Two of the important chemical functions of a soil are: (i) to supply nutrients to plants and (ii) to act as a ﬁlter or sink for pollutants. In both cases, the crucial chemistry governing the process occurs at the interface between soil solution and the solid phase. Various mechanisms control the equilibrium between chemical species in the two phases, with movement between them being controlled by changes in solution concentration. 5.5.1 Reactions in Soil Solution The soil solution contains dissolved gases (carbon dioxide being particularly important – see Section 5.4), soluble organic molecules from soil organic matter or exuded by plants and microorganisms, and inorganic ions released from minerals by weathering or, in agricultural soils, added as fertilizer. Typical ranges of ionic concentrations are given in Table 9. Inorganic ions can exist as hydrated ions in solution, with the form of the ion being determined by its charge (z) and size (r), expressed as the ionic potential z2/r. Ions with a small ionic potential (oapproximately six) are present in solution as the hydrated ion (e.g. in soil solution, K1 Table 9

Typical ionic concentrations in soil solution at pH 6–7

Ions , NO3 21 1

Ca Mg , K , Na1, Cl, HCO3 H2PO4 Trace elements (Zn, Cu, etc.)

Concentration range 102–103 mol L1 103–104 mol L1 105–106 mol L1 o106 mol L1

264

Chapter 5

and Na1 are weakly hydrated cations, while Ca21 and Mg21 are strongly hydrated). Ions with a medium ionic potential (10–20) tend to hydrolyze and precipitate as oxides and hydroxides (e.g. FeIII, AlIII, CrIII). Those with a high-ionic potential (420) form soluble oxyanions (e.g. phosphate, nitrate, sulfate). The chemical speciation of the ions is highly dependent on pH and redox potential. Phosphorus, for example, exists as orthophosphate anions: 1 H3PO4 " H2PO 4 þ H

pK 2.15

(5.26)

2 1 H2PO 4 " HPO4 þ H

pK 7.20

(5.27)

pK 12.35

(5.28)

3 1 HPO2 4 " PO4 þ H

Over the pH range of most soils, the dominant orthophosphate anion in soil solution will be H2PO4 (Figure 12). Only in alkaline soils will HPO42 be signiﬁcant, and PO43 ions are not important in soil solution. Interactions between ions in solution can be considered in two ways. Long-range interactions, those 40.5 nm, are accounted for by the

Figure 12

A distribution diagram showing the percentage composition of phosphoric acid speciﬁed as a function of pH (Reproduced from Harrison, R.M. et al., Introductory Chemistry for the Environmental Sciences, Cambridge University Press, Cambridge, 1991.)

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The Chemistry of the Solid Earth

concept of ion activity (see also Chapter 3) {i} ¼ gi[i]

(5.29)

where {i} ¼ activity of the ion [i] ¼ concentration of ion (mol L1) gi ¼ activity coeﬃcient. At inﬁnite dilution, {i} ¼ [i], i.e. gi ¼ 1. Ionic interaction increases with concentration and the square of the ionic charge, which is described by the ionic strength of the solution (I): I ¼

1X ½i z2i 2

ð5:30Þ

where

I ¼ ionic strength [i] ¼ molarity of each ion in soilution zi ¼ charge on each ion in solution. If all the ions in soil solution can be accounted for, I can be calculated, and then g. The relationship between the activity coeﬃcient of an ion and the ionic strength of the solution is given by the Debye–Hu¨ckel limiting law pﬃﬃﬃ ð5:31Þ Log10 gi ¼ 0:5 z2i I For soil solutions, the Davies equation, which is a modiﬁcation of the Debye–Hu¨ckel equation, is commonly used, and is applicable to solutions up to approximately I ¼ 0.7 mol L1 (see also Section 3.2.1): hpﬃﬃﬃ i pﬃﬃﬃ Log10 gi ¼ 0:5 z2i I= 1 þ I 0:3I ð5:32Þ

Worked example 5.7 – soil solution chemistry: ionic strength and activity coeﬃcients The following analytical data were obtained for an isolated soil solution: Ca21, 5.05 mmol L1; Mg21, 3.55 mmol L1; K1, 0.68 mmol L1; Na1, 1.80 mmol L1; NO3, 3.70 mmol L1; HCO3, 1.80 mmol L1; SO42, 6.25 mmol L1; and Cl, 1.40 mmol L1.

266

Chapter 5

How can we tell from these ﬁgures that the main ionic species have been accounted for? There are 19.68 mmol L1 of monovalent positive charge and 19.40 mmol L1 of monovalent negative charge. What is the ionic strength of this solution? I ¼1=2½ð5:05 22 Þ þ ð3:55 22 Þ þ ð0:68 12 Þ þ ð1:80 12 Þ þ ð3:70 12 Þ þ ð1:80 12 Þ þ ð6:25 22 Þ þ ð1:40 12 Þ ¼1=2 68:78 ¼34:39 in terms of concentrations in mmol L1 Ionic strength is expressed in terms of mol L1 I ¼ 0.0344. What is the activity coeﬃcient of each ion? By substituting I ¼ 0.0344 into the Davies equation, g monovalent ions ¼ 0.843, g divalent ions¼0.504.

Short-range (o0.5 nm) interactions between ions in solution result in the formation of ion pairs or complexes with other ions or with organic molecules. An ion pair is a transient entity formed by coulombic attraction between ions of opposite charge, with each ion retaining its own water of hydration. In a complex, the ions form covalent or coordinate bonds to form an entity with its own hydration shell. An example of an ion pair could be between calcium and sulfate ions in solution: 0 Ca21 þ SO2 4 " CaSO4

(5.33)

An example of a complex could be between copper ions and an organic molecule:

COOH

COO 2+

+ Cu COOH

Cu COO

(5.34)

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The Chemistry of the Solid Earth

Worked example 5.8 – activity of ions pairs in soil solution In a soil solution the activity of Ca21 ¼ 103 mol L1 and SO42 ions ¼ 5 104 mol L1, what is the activity of the ion pair CaSO40? CaSO04 " Ca21 þ SO2 4 ½Ca2 þ ½SO2 4 ½CaSO04

K ¼ 5.25 103

(5.35)

¼ 5:25 103

ð103 Þð5 104 Þ ¼ 9:5 105 molL1 5:25 103 Alternatively, this calculation can be done using logarithms.

½CaSO4 ¼

CaSO04 " Ca21 þ SO2 log K¼ 2.28 4

(5.36)

log½CaSO4 ¼log½Ca2 þ þ log½SO2 4 ð 2:28Þ ¼ð 3:00Þ þ ð 3:30Þ ð 2:28Þ ¼ 4:02 ½CaSO4 ¼9:5 105 molL1

5.5.2 Ion Exchange (Physisorption) Ion exchange is the process where ions held at a charged surface by coulombic bonding are exchangeable with other ions of the same charge in solution in contact with the surface. The reaction of major importance in soils is cation exchange – i.e. positively charged ions (cations) held at negatively charged surfaces. Anion exchange can occur, but is of minor importance. The solid phase surfaces on which cation exchange occurs are: Silicate clays

permanent negative charge on clay faces pH-dependent charge on clay edges

Humiﬁed organic matter

pH-dependent charge (always overall negative)

Hydrous oxides

pH-dependent charge (overall positive or negative)

268

Chapter 5

The strength of the attractive force between the cation in solution and the negatively charged surface of the soil component is governed by Coulomb’s law: q0 :q00 ðconstantÞ F ¼ ð5:37Þ D:r2 Where F is the attractive (or repulsive) force q 0 and q00 are the charges on the ion and the surface D the dielectric constant and r the distance of separation between the two charges. Thus, more highly charged ions are held at exchange surfaces in preference to lesser charged ions. Trivalent ions such as Al31 and Fe31 are most strongly held, then divalent ions such as Ca21 and Mg21, with monovalent ions being the least strongly held. For ions of the same charge, the strength of binding depends on the size of the hydrated ion, which is the species that exists in soil solution, and hence how close to the surface it can approach. So for monovalent ions, the relative strength of binding, or aﬃnity for the surface, is Cs1 4 Rb1 4 K1 4 Na1 4 Li1. Ionic concentration must also be taken into account, as these preferences can be overcome by high concentrations of an ion in soil solution. For example, sodium ions would dominate the exchange sites in soil ﬂooded by seawater. Ions are considered to be held at the charged surface in a diﬀuse double layer (the Guoy–Chapman model) in which the concentration of cations falls, and that of anions increases, with distance from the surface. In a reﬁnement of this, the Stern model introduces a layer of cations held directly on the surface of the soil component (Figure 13). Small cations, or dehydrated ions having lost their water of hydration, can sit at the surface and form a strong coulombic bond. This allows for the speciﬁcity described above. The exchange process can be written as a chemical reaction, so, for example, for K1 ions in soil solution exchanging with Ca21 on the surface (X) 1 ð5:38Þ CaX þ K þ Ð KX þ Ca2 þ 2 At equilibrium, the ratio of the ions on the surface is related to their ratio in solution: KX ðK þ Þ ¼ k: pﬃﬃﬃﬃﬃﬃ2þ CaX Ca

ð5:39Þ

269

The Chemistry of the Solid Earth Mn+

Mn+

Mn+ M

Mn+

An-

Mn+

A An-

Mn+

Mn+ AnMn+

M M

Mn+ Negative Surface

Stern layer

Mn+

Diffuse layer cations > anions

An-

Free solution cations = anions

Figure 13 The distribution of ions at a negatively charged surface of a soil component

The exchange is between equivalents of charge. Strictly, ionic activity, not concentration, should be considered, but as ionic concentration in soil solution is so low, it is commonly used. This is known as the Gapon equation. This approach stresses the importance of the ion activity ratio, which tends to remain constant. As a result, monovalent ions are lost from a soil by leaching in preference to divalent and trivalent ions. A consequence is the acidifcation of soil by leaching (see Section 5.4).

5.5.2.1 Cation-Exchange Capacity (See Also Chapter 4). The CEC of a soil is a measure of its ability to hold cations at negative sites by coulombic bonds. Because the exchange is by equivalents of charge, the units used for CEC are centimoles of monovalent charge per kilogram of soil (cmolc kg1). (Previously the units of milliequivalents per 100 g soil were used (meq 100 g1); the numerical value of CEC is the same whether expressed as cmolc kg1 or meq 100 g1.) Typical values of CEC for the aluminosilicate clays given in Section 5.2 are shown in Table 6. These values for clay minerals are accounted for mainly by the permanent negative charge formed owing to isomorphous substitution, with only a minor component coming from the variable, pHdependent charge on the clay edge. Humiﬁed organic matter has a CEC typically between 150 and 300 as cmolc kg1, all of which is pH dependent; the hydrous oxides also have a pH-dependent charge,

270

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but at most soil pHs the negative charge contributing to the CEC is small. In a soil of between pH 6 and 7, about 80% of the cation-exchange sites hold Ca21, about 15% Mg21, K1, and Na1 account for up to 5%, with small amounts of NH41 and H1. As the pH falls, Al31, Mn21, and H1 become relatively more important. The % base saturation is a measure of the amount of the CEC of a soil holding the major cations (also called exchangeable bases) Ca21, Mg21, K1, and Na1. P

% base saturation ¼

Ca2 þ ; Mg2 þ ; K þ ; Na þ 100 CEC

ð5:40Þ

where the units of exchangeable cations and CEC are cmolc kg1. % base saturation is high (approximately 80%) in fertile soils with basic minerals in the parent material and low (o30%) in highly leached acidic soils.

Worked example 5.9 – cation-exchange capacity and % base saturation One method to measure the CEC of a soil is to leach a column containing a known weight of soil with a solution of a salt so that the exchange sites are saturated with the cation (index ion). After washing out the excess salt solution, a diﬀerent salt solution is leached through the column and the exchangeable ions displaced, collected in a known volume and their concentration measured. 10 g soil was leached with 200 cm3 of 1 mol L1 CH3COOK solution in order to saturate the exchange sites with K1 ions. After washing out the excess CH3COOK solution, 200 cm3 of 1 mol L1 CH3COONH4 solution was passed through to displace the K1 ions by NH41. The solution that leached through the soil was made up to 250 cm3 in a volumetric ﬂask. Ten centimetre cube of this solution was diluted to 100 cm3 and the K1 concentration in the diluted solution measured by ﬂame photometry. If the K1 concentration in this solution is 41 mg cm3, what is the CEC of the soil in cmolc kg1? Concentration of Kþ in the diluted solution ¼ 41 mg cm3 ¼ 410 mg cm3 in the original 250 cm3 of solution:

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Weight of K in this solution ¼ 410 mg cm3 250 cm3 ¼ 102,500 mg or 102.5 mg Atomic weight of K ¼ 39:1 102:5=39:1 ¼ 2:62 mmol K in 10 g soil or 262 mmol kg1 ¼ 26:2 cmol kg1 As K1 is monovalent, this is 26.2 cmol of monovalent charge (cmolc) kg1 and equal to the CEC of the soil. Concentrations of exchangeable cations in the above soil were: 3480 mg Ca kg1; 474 mg Mg kg1; 430 mg K kg1; and 161 mg Na kg1. What is their concentration in cmolc kg1, and what is the % base saturation of the soil? Atomic weights: Ca, 40; Mg, 24.3; K, 39.1; and Na, 23. Ca :3480=40 ¼ 87 mmol Ca kg1 ¼ 8:7 cmol Ca kg1 ¼ 17:4 cmolc kg1 Mg :474=24:3 ¼ 19:5 mmol Mg kg1 ¼ 1:95 cmol Mg kg1 ¼ 3:9 cmolc kg1 K :430=39:1 ¼ 11 mmol K kg1 ¼ 1:1 cmol K kg1 ¼ 1:1 cmolc kg1 Na :161=23 ¼ 7 mmol Na kg1 ¼ 0:7 cmol Na kg1 ¼ 0:7 cmolc kg1 % base saturation ¼ ½ð17:4 þ 3:9 þ 1:1 þ 0:7Þ=26:2 100 ¼ 88%:

5.5.3 Ligand Exchange (Chemisorption) Ligand exchange is the process by which ions are held at a solid phase surface by covalent bonding. The term chemisorption is often used to describe this process in order to distinguish it from physisorption, where coulombic bonds are involved and the process is ion exchange (see previous section).

272

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Fe - O

Fe - OH2

Fe - O

Figure 14 Chemisorption of a phosphate anion on to an iron oxide surface by ligand exchange

The important ions held in this way are certain anions, especially phosphate, and some trace metals; the main surfaces involved are the iron and aluminium oxides. Unlike ion exchange, the surface charge plays no part in ligand exchange, which can take place onto positive, uncharged, or negative surfaces, as these surfaces have a variable charge, which depends on pH (Figure 7). The ion being chemisorbed replaces a water molecule or hydroxyl ion (OH) from the surface (Figure 14). Phosphate, silicate, borate, arsenate, selenite, chromate, and ﬂuoride are anions for which ligand exchange is important. Nitrate, chloride, bromide, and perchlorate are not held, while sulfate and selenate may be weakly held. As a consequence, leaching of nitrate and sulfate from soil in drainage water can be signiﬁcant, but very little phosphate is lost in solution. Of the trace metals, Co, Cu, Ni, and Pb are strongly held on oxide surfaces by chemisorption, but the process is much less important for Cd and Zn. Chemisorption is often described by use of the Langmuir equation: x/xm ¼ Kc/(1 þ Kc)

(5.41)

where c ¼ concentration of chemisorbed ion in solution at equilibrium x ¼ amount of the ion chemisorbed on to the soil xm ¼ maximum amount of ion chemisorbed onto the soil K ¼ a constant related to bonding. The equation is often used in the linearized form c/x ¼ c/xm þ 1/Kxm

(5.42)

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273

where c can be measured and, knowing the initial concentration and volume of the solution containing the chemisorbed ion, x can be calculated. By plotting c/x against c, a straight line will result (if the Langmuir equation is a suitable model for the process), which will have a slope of 1/xm and an intercept of 1/Kxm, allowing the theoretical maximum chemisorption of the ion to be calculated. In practice, application of the Langmuir equation is often not strictly valid, as a straight line relationship is not found. But it is commonly used in order to obtain a value for xm. 5.5.4 Complexation/Chelation Humic compounds can react with metal cations to form organometallic complexes: Mx1 þ OM-LHy " OM-LM(xy)1 þ yH1

(5.43)

These are particularly stable if more than one bond is formed (chelation) and are important for heavy metals such as Cu21 and Pb21, and also for Fe31 and Al31 for certain soil-forming processes. Complexation by low molecular weight organic compounds, such as those in root exudates, can act as a means of bringing metals into solution, making them more mobile and bioavailable, but reaction with high molecular weight humic material is an immobilizing process. The stability of the metal complex can be expressed as the equilibrium constant (see also Chapter 4). For example, if Equation (5.43) is simpliﬁed to M þ L " ML

(5.44)

(where M is the metal ion, L the organic ligand, and ML the metal complex) K ¼ [ML]/[M][L]

(5.45)

While experimentally [M] can be measured and [ML] calculated, it is diﬃcult to quantify [L], given the polymeric and chemically diverse nature of soil organic matter. 5.5.5 Precipitation/Dissolution The concentration of some ions in soil solution is controlled mainly by the presence of poorly soluble solid phase components. Good examples of this are iron and aluminium, both of which are highly insoluble at the pH of most soils, and which are controlled by the solubility of the

274

Chapter 5

hydrous oxides. Control of solution concentration by dissolution of an oxide does not apply to calcium, as the solid phases are too soluble. When calcium carbonate is present, which is insoluble, it can control the Ca21 concentration in soil solution. Worked example 5.10 – control of ion concentrations in soil solution by solubility of a solid phase (a) Amorphous iron oxide is commonly found in soil and its dissolution to release Fe31 ions can be represented by the equation Fe(OH)3(amorph) þ 3H1 " Fe31 þ 3H2O K ¼ 103.54 (5.46) ½Fe3 þ =½H þ 3 ¼ 103:54 log Fe3 þ ¼ 3:54 3pH At pH 6; log Fe3 þ ¼ 14:46; ½Fe3 þ ¼ 3:5 1015 mol L1 At pH 3; log Fe3 þ ¼ 5:46; ½Fe3 þ ¼ 3:5 106 mol L1 At pH 1; log Fe3þ ¼ 0:54; ½Fe3 þ ¼ 3:5 mol L1 agreeing well with the observed behaviour of Fe31, which is soluble only under very acid conditions. This can be conﬁrmed by using Ksp ¼ 1036 for Fe(OH)3 in Worked example 5.5. (b) Calcium carbonate dissolves to release Ca21 ions: CaCO3 þ 2H1 " Ca21 þ CO2 þ H2O K ¼ 109.74 (5.47) Log Ca21 þ 2pH ¼ 9.74 log CO2 For a soil atmosphere with pCO2¼102.5 (see Worked example 5.5) Log Ca2 þ ¼ 12:24 2pH At pH 7; log Ca2 þ ¼ 1:76; ½Ca2 þ ¼ 0:02 mol L1 At pH 5; log Ca2 þ ¼ 2:24; ½Ca2 þ ¼ 200 mol L1

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5.5.6 Soil Processes A soil receives inputs of ions from the weathering of parent material, mineralization of organic matter, pollution, and speciﬁc additions, such as fertilizers. The processes described above control the mobility, or bioavailability, of these ions (Figure 15). Certain ions, such as sodium and nitrate, are poorly held in soil, and so readily leached out. On the other hand, phosphate and copper are strongly held. From the point of view of soil as a source of plant nutrients, it is beneﬁcial for a soil to hold suﬃcient nutrients to meet plant requirements, and to hold the nutrients in a form that can be released to supply the plant over the growing season. Increasingly, we are using soils as a sink for pollutants, when the desirable result is that the pollutant is tightly held in order to prevent its dispersal into the wider environment. QUESTIONS (i) Explain the origin and characteristics of the charges on soil components that make up the CEC of a soil. (ii) Using the information in Figure 8, how much N, P, and S are stored per hectare in the soil humus? How much N, P, and S are released annually per hectare by mineralization of the humus? It is commonly found that about twice as much P and S is released from the soil than predicted, but that the amount of N agrees well with the prediction. What process could account for this? A soil solution contains 3 106 mol PO4-P L1, 5 104 mol SO4-S L1, and 2 103 mol NO3-N L1. If the annual drainage is 250 mm of water, how much P, S, and N are lost annually per hectare by leaching? (P ¼ 31; S ¼ 32; and N ¼ 14). (iii) A typical rate of addition of Pb in sewage sludge to agricultural soils in the UK is 1200 g Pb ha1 year1. If this amount is added to a soil of bulk density 1.2 g cm3 containing 5% humus with a binding capacity of 80 cmolc kg1, what percentage of the binding capacity to a depth of 20 cm is used annually to hold the lead? (Assume that all the added Pb is bound to the humus. Pb ¼ 207.) (iv) Using the values for solubility product below, and assuming that the solubility of the hydroxides controls the metal ion concentration in solution, calculate the theoretical concentrations of Fe21, Fe31, and Cr31 at pH 2, 5, and 7.

276

Chapter 5 Anthropogenic inputs:

Natural inputs: Atmospheric deposition Deposition from water Cycling through vegetation

Atmospheric contamination Water contamination Wastes added to soil

Losses: Volatilization to atmosphere Offtake in crops

Fertilizers

Soil Ion exchange physisorption

Ligand exchange chemisorption

Soil solution

Chelation

Precipitation

Losses: Leaching in drainage water Erosion of soil

Figure 15

An overview of soil processes

277

The Chemistry of the Solid Earth

Ksp Fe(OH)2 ¼ 1014; Fe(OH)3 ¼ 1036; Cr(OH)3 ¼ 1033 What does this tell you about the solubility of these ions? Does this agree with their observed behaviour in soil? (v) A soil with a bulk density of 1.2 g cm3 contains 1% FeS2. How much CaCO3 is required per hectare to neutralize the acid produced to a depth of 20 cm? (Assume that the FeS2 oxidizes completely. Fe ¼ 56; S ¼ 32; Ca ¼ 40; C ¼ 12; and O ¼ 16.). (Hint: look at Worked example 5.1 and combine Equations (5.16) and (5.17) to get an overall reaction for the oxidation of pyrite.) (vi) Iron oxides in soil undergo reduction to produce Fe21 ionsFe(OH)3 þ 3H1 þ e " Fe21 þ 3H2O E0 ¼ 1.057 V What is the concentration of Fe21 ions in solution in a soil of pH 7 at 100 and 0 mV? (vii) 1.00 g samples of a soil were shaken with 50 cm3 of diﬀerent solutions of known zinc concentration. After 1 h, the soil suspensions were ﬁltered and the concentration of zinc in the ﬁltrate measured, with the following results:

Initial concentration (mg Zn L1) 7.5 10 20 30 45 60

Final concentration (mg Zn L1) 1 2 5 10 21 34

In each case, calculate the weight of zinc held by 1 g of soil. Plot the graph of weight of Zn held per gram (y-axis) against the ﬁnal concentration of Zn. Fit the data to the Langmuir equation and plot the graph. Obtain values for xm and K. (viii) The equilibrium between Fe31 ions in soil solution and iron oxides in soil has been described by the equation Fe(OH)3

soil

þ 3H1 " Fe31 þ 3H2O log K ¼ 2.70

278

Chapter 5

(Note, this is a diﬀerent form of iron oxide to that used in Worked example 5.10, hence a diﬀerent value for K). Derive the relationship between the concentration of Fe31 ions in solution and pH. Use this equation to plot the concentration of Fe31 over the pH range 1–9. How well does this information agree with the Fe31 concentrations calculated in Question 4? (ix) What properties of a soil would you investigate in order to learn about its ability to hold and supply P and K to crops? (x) What properties of soil would you investigate in order to learn about the ability of a soil to act as a sink for pollutants?

REFERENCES 1. S.A. Barber, Soil Nutrient Availability, Wiley, New York, 1995. 2. J.B. Dixon and D.G. Schulze (eds), Soil Mineralogy with Environmental Applications, Soil Science Society of America, Madison, WI, 2002. 3. D.J. Greenland and M.H.B. Hayes (eds), The Chemistry of Soil Constituents, Wiley, Chichester, 1978. 4. D.J. Greenland and M.H.B. Hayes (eds), The Chemistry of Soil Processes, Wiley, Chichester, 1981. 5. W.L. Lindsay, Chemical Equilibria in Soils, Wiley, New York, 1979. 6. M.B. McBride, Environmental Chemistry of Soils, Oxford University Press, New York, 1994. 7. R.G. McLaren and K.C. Cameron, Soil Science, 2nd edn, Oxford University Press, Auckland, 1997. 8. J.J. Mortvedt, F.R. Cox, L.M. Shuman and R.M. Welch, Micronutrients in Agriculture, 2nd edn, Soil Science Society of America, Madison, WI, 1991. 9. D.L. Rowell, Soil Science: Methods and Applications, Longman Scientiﬁc & Technical, Harlow, 1994. 10. G. Sposito, The Chemistry of Soils, Oxford University Press, New York, 1989. 11. F.J. Stevenson, Humus Chemistry, Wiley, New York, 1982. 12. R.E. White, Principles and Practice of Soil Science, 3rd edn, Blackwell Science Ltd., Oxford, 1997. 13. A. Wild (ed), Russell’s Soil Conditions and Plant Growth, 11th edn, Longman Scientiﬁc & Technical, Harlow, 1988.

CHAPTER 6

Environmental Organic Chemistry CRISPIN J. HALSALL Environmental Science Department, Lancaster University, Lancaster LA1 4YQ, UK

6.1 INTRODUCTION Environmental Organic Chemistry (EOC) is an exciting branch of chemistry that has developed over the last few decades, and is probably best known for the study into the fate and behaviour of anthropogenic (manmade) pollutants. EOC, however, cannot be restricted to organic pollutants and their behaviour, but also encompasses the myriad of organic molecules produced by nature and their associated biogeochemistry. This area of chemistry is coupled to the recent advances in analytical techniques, particularly improvements in chromatography and the rise in ‘hyphenated’ techniques, such as GC-MS and LC-MS. The sensitivity of contemporary bench-top instruments allows detection of analytes in environmental samples, down to as low as a few picogrammes (pg) (1012 g) on the chromatography column. Recently, the improvements in LC-MS technology, has resulted in affordable, yet sensitive instrumentation that allows separation and detection of a mixture of polar, watersoluble analytes. The use of LC avoids derivatisation steps of highly polar analytes required as a pre-requisite to GC analysis. The different sources now available to LC-MS, including atmospheric pressure photoionisation and chemical ionisation, allows a high degree of selectivity for a wide number of analytes of varying molecular weights, polarity and aqueous solubility. This enables analysis of pollutant degradates and metabolites in aqueous systems and has led to the screening of pharmaceuticals, veterinary drugs and antibiotics in both surface and ground water.1–3 The traditional organic chemist often works in a highly controlled environment, where variables may be reduced to several reactants under carefully controlled temperatures and pressures; the goal to produce a 279

280

Chapter 6

product of the highest possible yield. The EOC, on the other hand, has to understand the chemistry and reactions of an organic molecule but under earth’s near-surface conditions, with varying temperature (50 1C to þ50 1C) and atmospheric pressures (950 to 41000 hectoPa (hPa)). In addition, the laboratory analyst in a chemical company, has the job of qualifying the purity of products to some standard speciﬁed by the customer. The environmental chemist however, has the difﬁcult task of extracting, purifying and quantifying the same chemicals, their unwanted by-products (and possibly their degradates) from environmental media (e.g. soil, water, biota, etc.) to a high degree of precision and accuracy. Measuring speciﬁc chemicals in different environmental media and understanding how the chemistry in these media ‘ﬁts together’ presents a considerable challenge, and often results in the environmental chemist returning to the laboratory to simulate a particular process. This chapter will explore aspects of EOC in relation to the environmental behaviour and fate of anthropogenic pollutants and will explore the physical and chemical processes that result in their environmental partitioning and degradation; specifically photodegradation for the latter. In each case practical examples and measurement techniques will be presented and illustrated. 6.2 THE DIVERSITY OF ORGANIC COMPOUNDS The wide array of organic compounds subject to research within EOC, encompass a broad range of physical–chemical properties, but can be categorised according to either molecular weight, volatility and/or reactivity. Figure 1 provides examples of a range of molecules that have vapour pressures spanning orders of magnitude, and are subject to research because of their interesting chemistry either in the atmosphere, aquatic systems and/or soil and sediments. Many of the more volatile compounds (C1C6) emitted into the atmosphere have an important impact on atmospheric photochemical processes. Oxidation of short-chain alkanes largely through reaction with the OH radical, form short-lived peroxy (HO2.) and alkoxy radicals (RO2.) that are important for converting NO to NO2 in the polluted atmosphere and hence allowing the build up of ground-level ozone (O3) during warm sunny weather (see Chapter 2). For higher molecular weight compounds ( Z C5) the isomerisation of alkoxy radical intermediates gives rise to hydroxyl carbonyl products containing OH, COOH and CQO groups. These products comprise the major fractions of organic aerosol, and a growing body of research literature is dedicated to understanding the oxidation products and reaction mechanisms of vapour hydrocarbons to less volatile,

281

Environmental Organic Chemistry P/Pa (at 298K) methane

CH4

ethene

107

propane

106

Cl 2,4- dichlorophenol

100

Cl H Cl H γ-HCH Cl Cl (organochlorine insecticide) H H Cl H 1,3- dimethylnaphthalene (alkylated-PAH)

10-1

isoprene dichloromethane

105

CH2Cl2

10-2

fluoranthene (PAH) benzene CH3

n-hexane

acetone

104

10-3

103

10-4

102

10-5 Cl

O toluene α-pinene (monoterpene) n-decane

pinonaldehyde (from α-pinene)

Figure 1

O O

n-heptadecane (n-C17)

trichlorinated naphthalene (PCN) pentachlorinated biphenyl Cl (PCB) benzo[a]pyrene (PAH)

101

-6

Approximate vapour pressures (P) (107106 Pa) for a diverse range of organic compounds, many of which are mentioned in this chapter. Note the use of common/trivial chemical names rather than IUPAC nomenclature

more polar condensates involved in aerosol formation and growth.4 Monoterpenes (C10) released from different types of vegetation, contain unsaturated double bonds and are therefore susceptible to reaction with O3. The reaction of a-pinene, for example, gives rise to a number of carbonyl and mono- and dicarboxylic acids, including pinonaldehyde and pinic acid which have been detected in forest aerosols.5 Semi-volatile organic compounds arise from a multitude of sources, both anthropogenic and natural, but include the persistent organic pollutants, many of which are halogenated aromatic compounds found at trace levels in the environment, but are dispersed widely across the globe. For aliphatic compounds such as petroleum hydrocarbons, the semivolatile constituents include longer chain alkanes (4C15), both normal (n-alkanes), branched (iso-alkanes) and cyclic alkanes (cyclo-alkanes or naphthenes) as well as a range of polycyclic compounds or terpanes including tricyclic carpanes, tetracyclic hopanes and pentacyclic steranes. Many of the hopanes and steranes are largely involatile, and in the atmosphere are predominantly associated with particles.6 They are important constituents of crude oils, arising principally from the decay of bacterial/biotic matter and serve as useful biomarkers for different fuel types, and are also widespread in sediments and soils from both biogenic

282

Chapter 6

and petrogenic sources. The semi-volatile polycyclic aromatic hydrocarbons (PAHs) and their alkylated derivatives are generally less abundant than the alkanes and terpanes of similar carbon number, but are present in oil and liquid fuels and are also produced during combustion processes (pyrolysis), and are present in soil and sediments from both petrogenic and pyrolytic sources. The ﬁve-ringed benzo[a]pyrene is a notable example, and is classed as a probable human carcinogen, present in urban environments chieﬂy from the combustion of fossil fuels. The involatile, complex polymeric organic matter present in soils and sediments is commonly referred to as ‘humic substances’, and probably forms through a series of oxidation and condensation reactions between polyphenols, polysaccharides and polyamino acids of plant and microbial origin.7 There is no commonly accepted structure of humic substances, with molecular weights and molecular formulae open to debate, and varying widely between different soil and sediment types. Extraction of soils with water and/or organic solvents generally yield only a small fraction of soilcarbon present as humic substances, but the accepted practice of extracting soils with aqueous alkali (0.5% NaOH), followed by centrifugation of the soil/sediment suspension with subsequent acidiﬁcation of the supernatant, yields the acid-insoluble humic fraction known as humic acid. Any organic material left in solution is referred to as fulvic acid and the remaining insoluble material is referred to as humin.8 These operationally deﬁned fractions prove very useful for conducting sorption and fate experiments for organic pollutants present in soil/sediment systems, where humic substances play an important role in the soil chemistry of organic compounds. Humic and fulvic acids give rise to dissolved organic material in water, and are complex organic macromolecules with a wide variety of functional groups, primarily carboxylic and phenolic groups, and to a lesser extent amino (e.g. -NH2) and thiol (-SH) groups. These macromolecules have a wide range of molecular weights (Eo500 to 410,000 g mol1) and various humic acid fractions have been separated by gel-permeation chromatography (or size-exclusion chromatography) with their structural and chemical characteristics investigated using 1H- and 13C-NMR, and pyrolysis-GC-MS.9 The large surface areas of these molecules provide ideal surfaces for sorption by other organic molecules, including semi-volatile organic contaminants such as PAHs and pesticides, providing active sites for surface-based, heterogeneous chemical reactions. 6.2.1 Identifying Sources of Hydrocarbons Distinguishing between different sources of hydrocarbons including alkanes and PAHs for example, is of fundamental importance for source

Environmental Organic Chemistry

283

apportionment of pollution, as well as in petroleum geochemistry to indicate the origin and thermodynamic maturity of organic matter in sediments. A commonly used index to broadly distinguish between anthropogenic and biogenic sources of hydrocarbons, present in any environmental compartment, is the Carbon Preference Index (CPI). CPI is usually applied to the hom*ologous series of semi-volatile n-alkanes with carbon chain lengths ranging from BC15 to C35, where the CPI is simply the concentration ratio of the summed ‘odd-numbered’ chain lengths over the ‘even-numbered’ chain lengths over a speciﬁed chain-length range. The premise here is that recent biogenically derived n-alkanes (e.g. plant waxes, biota exudates) tend to favour odd-carbon numbered chain lengths (typically within the range C25 to C35) whereas aged organic material, such as petrogenic hydrocarbons that have undergone ageing over geological timescales generally do not exhibit this odd-carbon number preference. Therefore, environmental samples with a CPI 41, suggest that the aliphatic hydrocarbons, have arisen through biogenic input, whereas a CPI of o1, suggests petrogenic (e.g. fossil fuel) sources. As an example, suspended particle matter (SPM) ﬁltered from seawater samples collected in the Black Sea, showed high CPI values of 3.81–13.1 in oﬀshore locations, where the SPM was considered to be of algal origin (where CPImarine ¼ SC1525/SC1424).10 Much lower CPI values were observed closer to the Danube estuary, and were indicative of petrogenic derived hydrocarbons, presumably due to river discharges of petroleum products. There are other useful aliphatic indices or markers that are used in conjunction with the CPI to provide evidence of oil ‘weathering’ or to apportion carbon sources as either petrogenic or of recent biogenic origin. The ubiquitous and rather persistent branched alkanes, pristane (C19) and phytane (C20), are often ratioed to n-heptadecane (C17) and n-octadecane (C18) respectively, as supporting evidence of either biogenically derived or petrogenic hydrocarbons to sediments and aqueous systems (both pristine and phytane closely elute with these two n-alkanes on a low-polarity GCcolumn). For instance, a high n-C17/pristane ratio (410), would indicate biogenically derived hydrocarbons, whereas a low ratio (o3) would indicate a carbon source depleted in n-C17, perhaps indicative of petrogenic carbon sources. PAHs and their alkylated derivatives are also used to distinguish between pyrolytic and petrogenic inputs to sediments and soils and are often used in conjunction with aliphatic indices to discriminate hydrocarbon sources. The lower molecular weight 2- to 4-ringed PAHs and their alkylated versions comprise a small, but significant, component of crude oils and liquid fuels. Ratios of parent PAHs (non-alkylated) as well as parent to alkylated versions are frequently used to distinguish

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petrogenic sources vs. pyrolytic sources (both natural and anthropogenic) for aromatic hydrocarbons in soil/sediment systems.11 6.3 THE FATE OF ORGANIC CONTAMINANTS There are thousands of organic chemicals entered into the international databases as part of legislative programmes aimed at assessing the risk they pose to human health and the environment. For example, in the European Union, the European Inventory for Existing Chemical Substances and the European List of Notiﬁed Chemical Substances include over 100,000 commercial substances.12 Similarly in the US, the Ofﬁce for Pollution Prevention and Toxics (OPPT) as part of the US-Environmental Protection Agency (EPA) maintains lists of high-production volume chemicals as well as persistent bioaccumlative toxins (PBTs). The increasing demand for risk assessment to examine substances with respect to their environmental persistence, ability to bioaccumulate and potential for toxicity, increases the demand for reliable physical–chemical property data, upon which to conduct environmental fate and degradation studies. In this respect, a risk assessment is conducted to ascertain the likelihood or probability of some harm arising to either humans or the environment due to exposure to that chemical substance. In order to assess exposure and understand chemical fate, then intimate knowledge of both the chemical use (quantity released) and behaviour (environmental transport, partitioning and transformation/degradation) are required. As Schwarzenbach et al.13 have so elegantly described in their comprehensive text on EOC; organic chemicals once present in the environment are subject to essentially two types of process. The ﬁrst involves their environmental transport within, and transfer between, environmental phases and the second involves their chemical alteration (transformation/degradation) driven by chemical and/or biological processes. In reality, within any given environmental system, these processes may occur simultaneously and different processes may strongly inﬂuence one another. 6.4 CHEMICAL PARTITIONING The transfer of chemicals between two or more environmental compartments or phases can be described by equilibrium partitioning, and knowledge of this partitioning is essential for understanding and describing chemical fate in the environment. Chemical partitioning takes place between adjacent phases such as between a solid and a liquid (dissolution), a liquid and a gas (volatilisation), a solution and a solid surface (adsorption) or a solution and an immiscible liquid (solvent

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partitioning). The net transport of an organic compound from one phase to another (e.g. air to water) is limited by equilibrium constraints and can be quantiﬁed according to a partitioning constant or coefﬁcient (K). This is simply deﬁned as the ratio of the concentration (C) of the chemical present in each phase (e.g. 1 and 2) at equilibrium; the point at which the chemical potential in each phase is equal and there is no longer net transfer in one direction. K can be represented as K1;2 ¼

C1 C2

ð6:1Þ

In strict terms, K is expressed as the ratio of the mole fractions (x) of a chemical within each phase (for example, this could be the mole fraction of a chemical present in a gaseous or liquid mixture). However, in environmental chemistry, the common way of expressing chemical concentration is not by mole fractions but as a molar concentration (mol L1 or molarity, M), although it is worth pointing out that the volume (L or dm3) may vary according to changes in temperature and pressure and should be normalised to standard temperature (298 K) and pressure (101,325 Pa (or 1 atm)). Furthermore, when dealing with one phase that is solid, it is often more appropriate to deal with concentrations as amount per mass to avoid having to estimate phase densities. A mole fraction can be converted to molar concentration (M) according to M¼

xðmol=molÞ VðL=molÞ

ð6:2Þ

where V is the molar volume of the mixture or solution. For aqueous solutions relevant for environmental scenarios, it is possible to ignore the contribution of the organic chemical to the molar volume of the mixture due to the typically low concentrations of a chemical pollutant present in a body of water. In this case, V is simply set as the molar volume of water (Vw ¼ 0.018 L mol1 at 25 oC).13 The term partitioning constant or coefﬁcient refers to one chemical species in each phase (for example, ionisable chemicals present in water may exist in both neutral and dissociated forms and therefore each would have a separate partitioning coefﬁcient). For well-deﬁned phases, such as pure water or the pure liquid state of the chemical, then partitioning with another equally well-deﬁned phase such as air results in the use of the term partitioning constant. The Henry’s Law constant describing chemical partitioning between pure air and water is one example. Distribution ratios on the other hand, such as the soil-water

286

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distribution ratio (Kd), describe partitioning between a heterogeneous solid phase (soil) and soil-water (containing a variety of dissolved solids and organic material). Furthermore, a distribution ratio accounts for all the speciated forms of a particular chemical, as well as the various different sorption mechanisms and correctly describes the distribution of a chemical rather than a particular partitioning process. An example of a useful partitioning coefﬁcient applied to soils and sediments is the organic-carbon to water partitioning coefﬁcient (Koc), which speciﬁcally describes chemical partitioning between the organic-carbon fraction of a soil or sediment and water and is frequently used to describe pesticide fate in agricultural soils. 6.4.1 Important Partitioning Coefﬁcients There are several important partitioning coefﬁcients, besides the fundamental properties of vapour pressure and aqueous solubility, which are essential for understanding chemical transfer in the environment and these are illustrated in Figure 2 and include the air/water (Kaw), n-octanol/ water (Kow) and n-octanol/air (Koa) partitioning coefﬁcients. Indeed, chemical vapour pressure (the ‘solubility’ of a chemical in air) and aqueous solubility explain the ‘partitioning’ of a chemical from the pure-liquid or solid-state into air and water, respectively. Transfer of either vapour or dissolved chemical to other environmental compartments, media or phases can be explained by speciﬁc partition coefﬁcients. Chemical P

AIR

WATER

Kaw

Koa

Kow

ORGANIC PHASE ‘Octanol like’ e.g. fats/lipids/waxes organic carbon in soil

Figure 2 Key physical–chemical properties that control the environmental partitioning of an organic chemical. The fundamental properties of vapour pressure (P) and aqueous solubility (S) determine the partitioning from the pure liquid or pure solid-state of the chemical into air and water, respectively

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6.4.1.1 Octanol-Water Partitioning Coefﬁcient (Kow). Octanol serves as a very useful surrogate of fats, lipids and organic phases present in the environment. This is largely attributable to octanol having a similar carbon to oxygen ratio (8:1) as fatty acids and lipids. As a result Kow is used to represent chemical partitioning between lipids and water (where lipids and fats can be viewed as ‘octanol-like solvents’) and is an important chemical property used in the ﬁelds of biology and pharmacy, where Kow is represented by the term ‘P’ for partitioning coefﬁcient. In essence, Kow is the ratio of the equilibrium concentration of a chemical between octanol (mol m3) and water (mol m3) and is a measure of a chemical’s ‘water hating’ tendency or hydrophobicity. As Mackay14 points out, most organic chemicals have a fairly constant solubility in octanol (200–2000 mol m3) but very different solubility in water and it may be misleading to describe Kow as a chemical’s ‘love of fats’ or lipophilicity. As the range of Kow for environmental contaminants is large (over orders of magnitude), log Kow values are often reported. Perhaps the most common application of Kow is for predicting the bioconcentration of low-polarity hydrophobic contaminants from water into aquatic biota. A bioconcentration factor (BCF) can be represented as BCF ¼

Corg Cw

ð6:3Þ

where Corg is the actual concentration determined in an organism (e.g. ﬁsh), or some ‘compartment’ of the organism (e.g. lipids), and Cw is the truly dissolved concentration of the chemical in water. From empirical studies that have examined a range of different organisms and chemicals, BCF may be predicted from simple linear relationships with Kow. For example, Mackay15 described the BCF in ﬁsh, by examining hundreds of organic chemicals, as BCF ¼ 0:048 Kow

ð6:4Þ

The constant 0.048 representing the E5% lipid content of ﬁsh, or as Mackay describes, about 5% octanol by volume. However, relationships between BCF and Kow appear to break down for chemicals with log Kow47, as these ‘superhydrophobic’ substances may have difﬁculty crossing cellular membranes (i.e. they have a low bioavailability) or are rapidly egested from the organism. There is, in fact, little consensus on the reasons for the breakdown in the linear relationship between BCF and chemicals with log Kow47, but both Borga˚ et al.16 and Boethling et al.17 point out that the experimental artefacts associated

288

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with measuring the freely dissolved chemical are likely to provide large uncertainties in the BCFs. Furthermore, bioconcentration is unlikely to be a significant exposure pathway for these ‘superhydrophobic’ chemicals relative to dietary intake. Kow is also used to describe partitioning to general organic phases within the environment and Koc (introduced above) can be directly estimated from Kow according to the empirically derived relationship derived initially by Karickhoﬀ18 and reﬁned by Seth et al.19 where Koc ¼ 0:35 Kow

ð6:5Þ

this relationship is useful for calculating the sorption of chemicals to soils and sediments from water; specifically to the organic carbon fraction. It is worth noting that Koc represents a solid/liquid partition coefﬁcient and as such has units of L kg1, whereby Koc can be represented as Koc ¼

Coc ðmoles=kgÞ Cw ðmoles=LÞ

ð6:6Þ

where Coc is the chemical concentration sorbed to the natural organic carbon fraction in soil (B0.02 or 2%) and Cw is the dissolved water concentration. As Kow is dimensionless, the constant in Equation 6.5 has units of L kg1. Experimentally, Koc is determined by introducing (or ‘spiking’) a chemical into a water/soil mixture (slurry) that is stirred at a constant temperature, followed by ﬁltration and analysis of the water at different time intervals. The resulting Kd value is then divided by the organic carbon fraction (foc) of the test soil, whereby Koc ¼

Kd foc

ð6:7Þ

Multiplying Kd (which also has units of L kg1) by soil density (B2.4 kg L1) allows the dimensionless form of Koc to be used. 6.4.1.2 Henry’s Law Constant and Air-Water Partitioning (Kaw). Airwater partitioning is described by the Henry’s Law constant (H), which is deﬁned as the ratio of the partial pressure (p) of a chemical in the air to its mole fraction (or molar concentration) dissolved in water (Cw) at equilibrium, according to H¼

pðPaÞ Cw ðmol=m3 Þ

ð6:8Þ

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289

H therefore has the units of Pa m3 mol1 and can be estimated by ratioing a chemical’s liquid-phase saturation vapour pressure (usually the sub-cooled liquid vapour pressure) over its aqueous solubility, at some reference temperature. Note that the vapour pressure (P) can be readily converted into ‘solubility’ with units of molar concentration by applying the ideal gas law PV ¼ nRT

ð6:9Þ

Rearranging this equation to n/V (mol m3) ¼ P/RT, where RT is the gas constant – temperature product (8.314 Pa m3 mol1 K K), allows the following expression P Ca ¼ ¼ H 0 or Kaw RT Cw Cw

ð6:10Þ

where P/RT is the air concentration (Ca) with the same units as Cw, and H 0 is the ‘dimensionless’ H or Kaw and can also be obtained by simply dividing H by RT (note that the units cancel), where T is the relevant environmental temperature in K. Like Kow, values of H vary tremendously for the range of organic compounds of interest in the environment. For example, short-chain alkanes with their high vapour pressures and low aqueous solubilities have values of H that are orders of magnitude higher than alcohols, which are highly soluble in water and have lower vapour pressures. For certain pesticides and other semi-volatile organic compounds, which may persist in the environment, their lower vapour pressures generally result in relatively low values of H, but often their very low solubility can result in higher than expected values of H. As a consequence, these chemicals (particularly sparingly soluble organochlorine pesticides) appreciably partition to air from bodies of water in which they may be residing (i.e. run-oﬀ from agricultural regions). In addition, they will not be efﬁciently removed from the atmosphere by rainfall and as a result may be transported long-distances by the prevailing wind. Table 1 provides examples of select chemicals and their calculated H values using both saturated vapour pressure and aqueous solubility data. Clearly, n-hexane has the highest value of H out of this group of chemicals, but interestingly H for benzene is 557, B200-fold greater than DDT and B4000-fold greater than chlorpyrifos, yet both pesticides are relatively involatile, with vapour pressures some 600 million and 85 million times less than benzene, respectively. The low solubility of these pesticides ensure that H is higher than may be expected and both pesticides partition appreciably out of water.

290 Table 1

Chapter 6 Select physical–chemical property data for four relatively volatile chemicals and two semi-volatile insecticides: p,p 0 -DDT and chlorpyrifos (a currently used pesticide). Note that the chemicals are arranged in order of volatility (Data sourced from Mackay14)

Chemical Chloroform n-hexane Benzene Phenol Chlorpyrifos p,p 0 -DDT

Molecular mass (g mol1) 119.4 86.2 78.1 94.1 350.6 354.5

Vapour pressure (P) (Pa)

Aq. solubility (S) (mol m3)

Henry’s Law constant (H) (Pa m3 mol1)

23,080 20,200 12,700 70.6 0.00015 0.00002

68.7 0.11 22.8 871 0.0011 0.0000087

336 18,364 557 0.081 0.14 2.3

Temperature is a key parameter for controlling H or Kaw, thereby strongly inﬂuencing air-water partitioning of chemicals in the environment. A good deal of effort is expended in deducing temperature-dependent H for contaminants of concern, which is not an easy task for low-polarity hydrophobic organic chemicals (HOCs) with poor solubility. Figure 3 shows a schematic diagram of the typical apparatus used to measure H for these types of chemical. In essence the major component is the reactor vessel, which contains a large volume of deionised water (B0.5–10 L) spiked with a known concentration of the HOC in question. The chemical is introduced by ﬁrst dissolving the chemical in an organic solvent which is miscible with water (e.g. acetone), and an aliquot added to the water in the reactor vessel. Care must be taken that the concentration does not exceed the aqueous solubility of the chemical and usually the concentration in the vessel is only a few percent of the aqueous solubility, resulting in a very low concentration, B1–2 mg L1. Clean, water-saturated air (or N2) is then bubbled or sparged through the reactor vessel and the exiting air passed through a vapour-trap to collect the chemical present in the air stream. The chemical quantity on the vapour trap can be turned into a time-averaged air concentration (Ca), by calculating the volume of air that has passed through the trap. H is calculated by ratioing Ca to the average water concentration in the reactor vessel (Cw) (see Equation 6.10) determined from analysis of water samples taken at the start and end of the experiment. Each gas-stripping run will be replicated to assess the precision, and new experiments conducted at different temperatures to provide the regression parameters to derive temperature-dependent values of H. Artefacts are always present, not least keeping the reactor vessel at a uniform temperature, but also preventing water from condensing in the vapour trap (particularly at

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Figure 3

Schematic diagram of the gas-stripping apparatus used to measure temperaturedependent Henry’s Law constants (H) for semi-volatile organic compounds (SVOCs) From Bamford et al.20 courtesy of SETAC press

higher temperatures), and accounting for the loss of water through evaporation. Furthermore, the change in water volume in the reactor vessel when extracting aliquots for chemical analysis should also be taken into account21 and efforts directed at shielding the apparatus from laboratory light when determining H for light-sensitive chemicals. 6.4.1.3 Octanol-Air Partitioning Coefﬁcient (Koa). Partitioning of a chemical between some organic phase and air is often described using Koa. Examples of terrestrial organic phases include organic carbon in soil, the waxy cuticle and lipid portion of vegetation and the organic ﬁlm that coats atmospheric particle matter and partitioning between these phases and air for a number of chemical groups have been successfully described using Koa.22–24 Laboratory-based measurements of temperature-dependent Koa values have been derived for a number of semi-volatile chemical groups, including PCBs25,26 PAHs,27 polychlorinated naphthalenes (PCNs)27 and organochlorine pesticides.28 Koa can be described as the ratio of the chemical (solute) concentration in octanol to the concentration in air at equilibrium, represented as Koa ¼

Co Ca

ð6:11Þ

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Chapter 6

where Co and Ca represent the equilibrium concentrations in octanol and air respectively (mol m3) and hence Koa is dimensionless. Koa may be estimated as the ratio of Kow over Kaw i.e. Koa¼Kow/Kaw. As Kaw¼H/ RT then Koa can be expressed as Koa ¼

Kow RT H

ð6:12Þ

As there are very few direct measurements of Koa, application of Equation 6.12 is useful, although this approach may have limited utility, as reliable values of H and Kow for a particular chemical may be scare. Furthermore, the validity of Equation 6.12 is brought into question when comparing calculated Koa values with actual measurements26 and extensive additional measurements of Koa are required for a wide number of chemicals that occur in the environment. As with Kow and Kaw, values of Koa may vary over orders of magnitude within any particular chemical class, and therefore Koa is often reported as a log value. 6.4.2 Temperature Dependence Before proceeding, it is worth turning brieﬂy to thermodynamics to provide some fundamental aspects to partitioning coefﬁcients and, importantly, their relationship to temperature. Wide variations in temperature experienced in the environment can greatly alter the partitioning behaviour of a chemical, and this will be demonstrated later when comparing chemical behaviour between different latitudes with notable temperature differences. As with any chemical reaction, movement of a chemical from one phase to another (i.e. liquid to gas) results in a change in the energy status of the system. A partitioning coefﬁcient is related to the (Gibbs) free energy of transfer of the chemical between the two phases (DG, J mol1). Where DG relates both the enthalpic and entropic effects that result from the changes occurring in either phase (i.e. intermolecular interactions), through both the removal and addition of molecules of the chemical. The relationship between DG and the change in enthalpy (or heat energy) of transfer (DH, J mol1) and entropy (DS, J mol1 K1) is given by the Gibbs–Helmholz equation: DG ¼ DH–TDS. Knowledge of the change in enthalpy for a certain partitioning process is particularly useful and is sought by environmental chemists to allow derivation of temperature-dependent partitioning. The relationship between K and DG is given by K ¼ constant:eDG=RT or ln K ¼

DG þ lnðconstantÞ RT

ð6:13Þ

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Environmental Organic Chemistry

K is therefore a function of temperature through the term RT, where R is the gas constant, 8.314 J mol1 K and T is the temperature expressed in Kelvin (K). The value of the constant given in Equation 6.13 depends on how the abundance of the solute or chemical is expressed i.e. partial pressure, mole fraction or molar concentration. For mole fraction the constant has a value of 1, whereby ln 1 ¼ 0 and can therefore be ignored. Rearranging Equation 6.13 to solve for DG, and then combining with the Gibbs–Helmholz equation results in RT ln K ¼ DH TDS Hence; ln K ¼

DH DS þ RT R

ð6:14Þ ð6:15Þ

It therefore follows that an experimentally determined plot of ln K against 1/T (1/K) will yield a slope of DH/R, and an intercept of DS/R (DH and DS are considered to be constant over a small temperature range). In laboratories measuring partition coefﬁcients for environmental fate, then 10 to 45 1C is a typical temperature range. Equation 6.15 is usually expressed in the form: ln K ¼

m þb T

ð6:16Þ

where m is the slope of the line and hence Rm ¼ DH. Once DH is established then K can be adjusted to any environmentally relevant temperature according to the integrated form of the van’t Hoﬀ equation DH 1 1 ln KðT2 Þ ¼ ln KðT1 Þ ð6:17Þ R T2 T1 where T1 is the reference temperature (298 K) at which K(T1) and DH have been measured and T2 is the temperature of interest. For many organic chemicals that persist in the environment, careful laboratory measurements of a particular partitioning coefﬁcient over a range of temperatures may not exist, therefore DH0(the enthalpy for the partitioning process at standard state, 101 kPa and 298 K) is often used in Equation 6.17.

Case study and worked example – Calculating water concentrations of Lindane To illustrate the effect that temperature has on a partition coefﬁcient and use some of the equations above, it is worth examining the

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air-water partitioning of a pesticide. Lindane is an organochlorine insecticide that comprises of almost 100% g-hexachlorocyclohexane (g-HCH). The chemical has been recently nominated as candidate POP chemical and it is widely distributed in the global environment. Indeed this chemical is routinely monitored in arctic air29,30 and is present in the polar environment largely through long-range atmospheric transport. Let’s assume that g-HCH has an air concentration of 100 pg m3 within an agricultural region during the summer, where both air and surface water (lakes, rivers coastal seas) have a temperature of 25 oC. Using Kaw we can estimate the water concentration from the air concentration assuming equilibrium has been achieved. In the Arctic, the air concentration of g-HCH is B10 pg m3 (10-fold lower), but the air and surface water temperatures in the high Arctic during the summer are B1 1C. Again, the water concentration of g-HCH can be estimated using Kaw, but this time adjusted to 1 1C. The surface water concentration in the temperate agricultural area is calculated from the air concentrations as follows: H for g-HCH has been measured as 0.24 Pa m3 mol1 at 298 K.21 Whereby H 0 or Kaw ¼ H/RT (T is 25 1C or B298 K) Therefore, Kaw ¼ 0.24/8.314 298 ¼ 9.69 105 or ln Kaw ¼ 9.24 As Kaw ¼ Ca/Cw, then Cw ¼ Ca/Kaw¼100/9.69 105 Cw in the temperate agricultural region is 1.03 106 pg m3 or B1000 ng m3 or B1 ng L1 The water concentration in the Arctic is calculated in the same way but using the Ca for the Arctic. First, Kaw has to be adjusted for 1 1C (B274 K) using Equation 6.17, where: DH 1 1 ln Kaw ð274 KÞ ¼ ln Kaw ð298 KÞ R 274 298 DH is 61400 J mol1 determined from the slope, m, in Equation 6.16 by Sahsuvar et al.21 Note that in this study, DH actually represented the enthalpy of water to air volatilisation, rather than the enthalpy of air to water dissolution, and we consider 61,400 as approximate only. Therefore: DH 1 1 ln Kaw ð274 KÞ ¼ ln Kaw ð298 KÞ R 274 298 The inverse of ln Kaw 11.4 ¼ 1.12 105. Therefore, the surface water concentration in the Arctic, attained through equilibrium partitioning with the air is Cw¼Ca/Kaw ¼ 10 pg m3/1.12 105 ¼ 892,857 pg m3 or B890 ng m3 or El ng L1

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The two water concentrations (temperate and Arctic) are similar, even though the air concentration in the Arctic is an order of magnitude less than temperate air. This temperature driven phenomenon that favours chemical partitioning to surfaces is commonly known as the ‘cold condensation effect’ and has resulted in the worrying rise in persistent organic pollutants in the polar environments.

6.4.3 Partition Maps The rise in legislation and the concern over the use and release of chemicals that possess physical–chemical properties that may be analogous to DDT and other persistent organic pollutants such as polychlorinated biphenyls (PCBs), has prompted demand for ways to rapidly screen or assess chemicals with respect to their environmental distribution and fate. The environmental transport, lifetime and accumulation of organic chemicals in different compartments, such as biota, can be quantiﬁed through the use of multi-compartment (media) chemical fate models. It is simply not feasible to measure all chemicals of concern in a large number of compartments over an extended period of time. Identifying chemicals of concern, and the designation of the term ‘POP’ for a chemical substance, occurs if a chemical meets the criteria speciﬁed in Table 2. Importantly, the designation of ‘persistence’ is based on the Table 2

Criteria used to designate a chemical substance as a persistent organic pollutant or ‘POP’. Information from the Stockholm Convention on POPs31

Persistence

1. 2. 3.

t1/2a in water 4 2 months t1/2 in soil/sediments 4 2 months Other evidence that the chemical is sufﬁciently persistent to justify consideration

Bioaccumulation

BCF or Bioaccumulation Factor (BAF) in aquatic species 4 5000 Log Kow 4 5 Evidence of higher bioaccumulation in other species

2. 3. Potential for long range transport

1. 2. 3.

t1/2¼half-life.

Measured levels of the chemical in locations distant from sources of release that are of potential concern Environmental fate properties and/or model results that demonstrate that the chemical has a potential for long-range environmental transport t1/2 in air 4 2 days

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Chapter 6

exceedence of speciﬁc environmental half-lives i.e. Z 6 months in soils/ sediments, Z 2 months in water, Z 20 days in air. The latter is particularly relevant for those chemicals, which may be released to the air and are resistant to atmospheric degradation and/or deposition, and hence able to undergo long-range atmospheric transport, with the potential to contaminate remote areas. This multiple half-life approach is one adopted by several countries including Canada,24 whereby a chemical that exceeds the criteria for any of these compartments may be declared persistent. However, if the chemical is unlikely to reside in that compartment then this may unfairly penalise a substance, because it fails the half-life criterion for a medium into which it does not appreciably partition. To examine this issue, Gouin et al.32 proposed a screening level method to allow the quick assessment of a large number of chemicals, based on their media-speciﬁc half-lives and their partitioning characteristics. An evaluative environmental fate model, the Equilibrium Criteria Model (EQM), was used for this study. EQM comprised of four hypothetical compartments air, water, soil and sediment, where octanol represented the organic matter found in soils and sediments. Each compartment was given a volume, in this case to represent the regional or country scale. For example, the air volume (Va) was 1014 m3, that is a land surface area of 100,000 m2, with an atmospheric height of 1000 m. 10% of this area was considered to be covered with water to a depth of 20 m, resulting in a water volume (Vw) of 2 1011 m3. Soil was considered to have a depth of 0.1 m with the soil volume (Vsoil) calculated as 9 109 m3, converted to an equivalent volume of octanol (Vo) according to Vo ¼ Vsoil F 0:35r

ð6:18Þ

where F is the fraction of organic carbon in the soil (0.02), r is the density of soil (2.4 kg L1 or 2400 kg m3), where 0.35 represents the constant relating the Koc to Kow (see Equation 6.5). Similarly for sediment, the equivalent Vo was 3.4 106 m3 calculated using Equation 6.18, but with Vsed as 108 m3, F as 0.04 and r the same as soil. For this scenario the ratios of the air/water/octanol volume was B650,000:1300:1. The EQM model is a steady-state, equilibrium model into which a chemical was constantly discharged and allowed to attain equilibrium between the various media with no advective losses. Steady state was reached when total reactive losses equalled the discharge rate and the model was run for 233 chemicals using respective physical–chemical property data including Kaw, Kow and Koa as well as reactivity data in the various media (i.e. water, air, soil/sediment). At equilibrium the total mass of a chemical (M, g) can be given by M ¼ Vw Cw þ Va Ca þ Vo Co ¼ Cw ðVw þ Kaw Va þ Kow Vo Þ

ð6:19Þ

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297

where Co is the chemical concentration in octanol (i.e. ‘soils/sediment’) (g m3). Note that the volume-concentration product (VxCx) gives the mass of a chemical for a particular environmental medium. The right-hand side of Equation 6.19 illustrates how using the partition coefﬁcients, the concentration of a chemical can be deduced for one compartment with knowledge of the chemical concentration in an adjacent compartment (e.g. KawCw ¼ Ca). The mass fractions of a chemical in each medium (F) are therefore Air ðaÞ : Fa ¼ Kaw Va =ðVw þ Kaw Va þ Kow Vo Þ

ð6:20Þ

Water ðwÞ : Fw ¼ Vw =ðVw þ Kaw Va þ Kow Vo Þ

ð6:21Þ

Octanol ðoÞ : Fo ¼ Vo Kow =ðVw þ Kaw Va þ Kow Vo Þ

ð6:22Þ

Figure 4 displays the plot of log Kaw vs. log Kow including points that represent all 233 chemicals used in the study. The chemicals were selected to represent environmental contaminants of concern and include a broad range of different chemical classes with varying properties. Figure 4 is effectively a partitioning map that illustrates the proportion of a chemical within each of the compartments. The 45o diagonals are lines of constant log Koa, as Koa ¼ Kow/Kaw (or log Koa ¼ log Kowlog Kaw). Lines of constant Fa, Fw and Fo were drawn using the volume ratios outlined above for this particular study. The 1% and 99% lines divide the Kaw/Kow space into regions into which chemical partitioning is predominantly into one medium. For example, the region to the upper left of the ﬁgure is where more than 99% of a chemical is in air, to the lower right (beyond the 1% air line) is where more than 99% of a chemical is in the octanol phase and to the lower left (beyond the 1% octanol line) is where more than 99% of a chemical is in water. For chemicals that fall into these ‘areas’ then their half-lives in the respective medium will likely control their overall persistence and the half-lives in the other media are largely irrelevant. For example, a chemical present in the lower right of the ﬁgure, within the octanol ‘area’, will be strongly sorbed and only degradation data for soils/sediments is likely to be needed. Importantly, half-lives in air are generally shorter than those in water or organic phases due to rapid reactivity with the hydroxyl radical (OH ), so that even 0.5% partitioning to air may result in appreciable degradation with respect to overall loss. It is therefore useful to include a line that represents the 0.1% air. The lines corresponding to one third (33%) in each compartment, converge where log Kow is 3.1 and log Kaw is 2.74; chemicals that fall within this area (‘A & W & O’) are truly multi-media and their half-lives

298

Chapter 6 -3

-2

-1

1% octanol

3 Air

7 Air & Octanol 33% air

1 0

8 9

1% air

10 0.1% air Air

-2

-3

11 12

A&W&O Water

-4

-5

Log K OA

-1

Log K AW

-6 Octanol -7

Water

-8

Water & Octanol 33% 33% octanol water

-9

17 1% water 18

-10 0

4 5 Log K OW

Figure 4 Kaw – Kow partitioning map with the distribution of 233 organic chemicals and lines of equal Fa, Fw and Fo; that is the chemical mass fractions in air (Fa), water (Fw) and octanol (Fo). The central region denoted ‘A&W&O’ is occupied by chemicals which have mass fractions (41%) in all compartments of Air, Water and Octanol and are therefore multi-media chemicals. Octanol represents organic-phases in the environment e.g. the organic carbon within soils and sediments From Gouin et al.32, courtesy of the American Chemical Society

for each compartment or medium would be required to determine their overall persistence. Obviously the convergence point would change if a different scenario was modelled, whereby the water compartment volume was increased for example. However, out of the 233 chemicals included in the ﬁgure, B34% had fractions 41% in each medium. If the chemicals

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between the 1% and 0.1% of air are included, then B40% of the chemicals become multi-media. This approach of mapping chemicals into groups according to their partitioning tendencies provides a rapid method of identifying relevant environmental compartments where a chemical may reside. Key degradation data relevant for that speciﬁc compartment can then be further examined without the need for irrelevant data for other compartments. 6.5 CHEMICAL TRANSFORMATION AND DEGRADATION Chemical degradation may proceed by a wide number of processes and the reactivity of a chemical is governed by its stereochemistry (3dimensional structure), bond strengths and the presence of functional groups. Both abiotic and biotic processes will transform a chemical pollutant and the multitude of abiotic processes can be broadly grouped under the following types of reaction: oxidation, reduction, hydrolysis and photolysis. Indeed these generic reactions are not mutually exclusive, as the generation of oxidants present in the atmosphere for example, requires the action of sunlight on precursor molecules (e.g. the formation of the OH radical from ozone). Several of the chlorinated aromatic chemical groups illustrated in Figure 1 (and mentioned in Section 6.4) are largely resistant to these transformation processes, whereby their rates of reaction are very slow. However, even these chemicals do degrade eventually, although their degradates may have a similar longevity in the environment! Perhaps the fastest route of degradation for these semi-volatile compounds is through oxidation by the OH radical in the atmosphere, with half-lives typically on the order of hours to days, compared to months or years in soils and sediments. However, at any one time, the atmosphere may contain only a very small fraction of the overall environmental ‘inventory’ so its relevance as an important ‘sink’ for these chemicals is greatly reduced compared to volatile organic compounds (VOCs). Nonetheless, oxidation of PAHs present in the atmosphere can give rise to toxic analogues. The reaction of higher volatility PAHs (generally the 2- to 4-ring PAHs present predominantly in the vapour phase) with OH radicals during daylight hours, or with nitrate (NO3) radicals during nighttime gives rise to more polar hydroxy- and nitro-PAHs. The atmospheric formation of nitro-PAHs may also occur via a daytime OH radical-initiated pathway in the presence of NOx, although the formation yields of nitro-PAHs (or nitroarenes) via this pathway are low compared to the yield of hydroxyPAHs.33 Figure 5 illustrates a simpliﬁed reaction pathway of ﬂuoranthene with the OH radical in the presence of NOx, resulting in the

300

Chapter 6 OH

H OH

fluoranthene

NO2

+ HONO

3-hydroxy-fluoranthene

NO2

H NO2 OH

NO2

H + H2O

2-nitrofluoranthene

Figure 5 The gas phase OH radical-initiated reaction of ﬂuoranthene (a 3-ring PAH) in the presence of NOx. The reaction results in the formation of both hydroxy- and nitro-ﬂuoranthene, although nitro-ﬂuoranthene can also form from the nighttime reaction with the NO3 radical Reaction scheme adapted from Arey et al.33

formation of both 3-hydroxyﬂuoranthene and 2-nitroﬂuoranthene (2NF). Indeed, some of the nitroarenes are strongly mutagenic and present a risk to human health and have been reported in several recent air sampling campaigns.34,35 Interestingly, 2-NF is one of the major particle-associated nitro-PAHs observed in a study conducted in southern California,35 presumably due to its lower volatility and higher polarity than the parent ﬂuoranthene, resulting in its condensation/sorption to atmospheric particles. An important transformation process for halogenated organic compounds is reductive dehalogenation, a process which largely occurs in anoxic (oxygen free) environments, such as sediments, sewage sludges, water-logged soils and ground-water and catalysed by the presence of metal containing minerals, where the metal is in a low state of oxidation e.g. zero-valent or divalent iron. The reductive potential of a halogenated molecule (i.e. ‘the willingness to accept electrons from the metal donor’) is based on many factors, not least the number and type of halogen atoms around the molecule. In one common type of reaction the breaking of the carbon-halogen bond (through the halogen atom accepting electrons) is followed by replacement with a H-atom (hydrogenolysis), resulting in the gradual ‘de-halogenation’ of the molecule. For chlorinated solvents and chlorinated aliphatics in general, then this

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process may be driven solely by abiotic processes,36 but for the more stable chlorinated aromatics, reductive dechlorination is likely to be assisted by microbial activity (or biotic processes) and is currently the focus of much research into the biologically driven remediation of contaminated soils and sediments. 6.6 CHEMICAL TRANSFORMATION THROUGH PHOTOCHEMISTRY Organic pollutants are susceptible to photochemical transformation and degradation in both the atmosphere and in surface waters. In fact, photochemical degradation can occur in any environmental medium that is subject to sunlight (or more appropriately solar actinic radiation), and includes surface reactions whereby chemicals sorbed to environmental surfaces such as plant material, airborne particles and even snow and ice are affected by sunlight. Both direct and indirect photochemistry (the latter brought about by chemical reaction with short-lived oxidants such as the OH radical), is of interest to the organic chemist, and can lead to the destruction of the pollutant and/or its phototransformation. In some cases this may lead to more persistent, or toxic, photoproducts such as the photochemical transformation of chlorobenzenes to polychlorinated biphenyls in ice37 or the soil-sorbed photooxidation of the organophosphorus pesticide parathion to the active toxin paraoxon.38 The assessment of photochemistry on the fate of organic chemicals is therefore a necessary step in risk assessment procedures (required for pesticide registration, for example), with protocols for conducting laboratory photochemical fate studies now provided by organisations like the US-EPA, the Organisation for Economic Cooperation and Development (OECD) and the European Chemical Bureau; the aim here is to provide standard operating procedures to enable reproducibility in test results and intercomparsion between different studies. These protocols concern direct and indirect photolysis in aqueous systems, as both fresh and marine waters are important sinks for organic pollutants. For example, agrochemicals residing on soil and plant surfaces often end up in surface waters, and the multitude of organic contaminants released from waste water treatment works (WWTW) (e.g. non-ionic surfactants, pharmaceuticals, personal care products) and their impact on receiving waters is of growing concern.39 However, atmospheric photochemistry relating to both vapour and particle-bound organic contaminants is perhaps less well studied, primarily due to the problem of maintaining a consistent, reproducible vapour in order to conduct gas-phase photochemical experiments. Nevertheless, there is a growing need to assess

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both direct photolysis and photooxidant chemistry to determine reaction rates to establish atmospheric half-lives and hence atmospheric transport potential – an important feature with respect to the contamination of remote environments by persistent chemicals (see Table 2). For the atmospheric chemist, much more emphasis is placed on photochemistry of VOCs, which play a major role in inﬂuencing air quality, by driving the formation of ground-level O340 as well as initiating particle formation and growth.41 In the following sections focus will be on semivolatile organic contaminants and their aqueous photochemistry. 6.6.1 Light Absorption and the Beer–Lambert Law For the environmental photodegradation of organic chemicals, direct photolysis is an important removal process for those pollutants that absorb light at wavelengths above 290 nm; the cut-oﬀ of solar irradiation at the earth’s surface. Visible light includes those wavelengths that cover the spectral range of 400–760 nm, whereas shorter wavelengths i.e. 290–400 nm are in the ultraviolet or UV region, and posses greater energy than light at longer wavelengths. Hence light in the UV-region is of most concern with respect to sunburn and skin cancers, and can be further divided into different wavelength regions, including UV-B (290–320 nm) and UV-A (320–400 nm). The degree of absorption of UV and visible wavelengths (UV/vis) by organic compounds is related to the molecular structure of the chemical, in particular, by the presence of chromophores; structural moieties that exhibit a characteristic UV/vis absorption spectrum. Aromatic rings and conjugated double bonds are good examples of chromophores, but others include double bonds containing, for example, nitrogen and oxygen and other heteroatoms present within the carbon skeleton of the molecule. Light absorption by organic compounds at the wavelengths of environmental relevance is usually associated with the delocalised p-electron system of double bonds (i.e. those electrons involved in the multiple bonding of CQC and CQO groups, for example), rather than electrons in sigma orbitals as part of single bonds, as the absorption maxima of these electrons generally lie at much shorter wavelengths than the solar wavelength minimum of 290 nm. Nearly all environmental photochemistry involves the excitation of p or non-bonding electrons (n) to higher-energy, antibonding p* orbitals, thereby destabilising the molecule. Examples of n electrons are the non-bonding electron pairs associated with oxygen atoms of carbonyl groups (CQO) or nitrogen atoms (CQN). Chromophores consisting of a series of conjugated carbon double bonds (CQC–CQC) are readily found in nature, and

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a good example are the porphyrins (of which chlorophyll is a member). These provide useful ‘light sensitive’ systems that can harvest light energy and in the case of chlorophyll perform the vital role of photosynthesis. Absorption of light at a given wavelength (Al) by a solution of a chromophoric molecule can be treated quantitatively according to the Beer–Lambert law. The ratio of the light intensity transmitted through the chemical solution (Il) to the incident light intensity (Io,l) can be related to the chemical concentration (C) according to, Al ¼ log

Il ¼eCl Io;l

ð6:23Þ

where C is expressed as molar concentration (M), e the decadic (log) molar absorption coefﬁcient (M1 cm1) and l the light path length in the chemical solution (cm). I is often expressed in units of Einstein cm2 s1, (where an Einstein is 6.02 1023 photons or a ‘mole’ of photons) with absolute light irradiance (expressed as Joules s1 m2 or Watts m2) measured using a chemical actinometer for a speciﬁed wavelength range. 6.6.2 Photolysis in Aqueous Systems Light absorbance studies can be conducted using a UV/vis spectrophotometer programmed to irradiate a solution at individual wavelengths over part, or all, of the solar spectrum. Quartz-glass cuvettes (l ¼ 1 cm) are usually employed which hold the chemical solution and a reference solution (i.e. typically the solvent but without the chemical to assess the absorbance artefacts of the holding solvent). Water is the solvent used for environmental fate studies, but apolar or monopolar organic compounds (including many pesticides) are only sparingly soluble in water so an organic solvent – water solution is often used, where the organic solvent is strongly polar and water-miscible such as methanol or acetonitrile. Figure 6 shows the spectral absorbance of chlorpyrifos (a currently used organophosphorus insecticide) in pure methanol between the wavelengths 200–400 nm. Chlorpyrifos displays several absorbance maxima, but within the UV-B/UV-A range of the solar spectrum the maximum absorbance occurs at 288 nm and e288 nm can be calculated by rearranging Equation 6.23 to: e288 nm ¼

A288 nm Cl

ð6:24Þ

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Chapter 6

Cl Cl

OCH2CH3 OCH2CH3

Cl Chlorpyrifos

λmax=288nm UV-region of natural sunlight

Wavelength(nm)

Figure 6 Spectral absorbance of the insecticide chlorpyrifos (in methanol) over UV wavelengths. For the environmentally relevant range of the UV spectrum (UVA/B), lmax occurs at 288 nm, with significant absorbance at higher wavelengths (to 310 nm), suggesting natural sunlight is capable of inducing phototransformation (thanks to J. Weber, Lancaster University, for generating the absorbance spectrum)

From Figure 6, A288 nm ¼ 0.49 and was generated with C as B 0.1 mM (or 1 104 M) and l ¼ 1 cm, the value of e288 nm is therefore: e288 nm ¼

0:49 ¼ 4900 M1 cm1 104 1

ð6:25Þ

Values of lmax and emax are useful to assess whether a compound might absorb ambient light and therefore be susceptible to environmental photodegradation. Values of e however need to be calculated for each wavelength if full quantiﬁcation of photochemical processes is required. Indeed for chlorpyrifos, significant absorbance occurs at wavelengths 4290 nm suggesting that photodegradation is worth considering with respect to its environmental fate. It is worth noting that the solvent in which the chemical resides may have a strong inﬂuence on the absorption spectrum, and is important when trying to replicate environmental conditions for a chemical with a low aqueous solubility. It is therefore relatively common practice to replicate photochemical experiments using a number of solventwater mixtures to assess the solvent effect on the spectral absorbance. Indeed, OECD guidelines with respect to aqueous photodegradation studies on compounds with low solubility, recommend that co-solvents

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should not comprise more than 1% by volume of the test solution. Furthermore, the co-solvent must not solvolyse the test chemical, must not absorb in the 290–800 nm region, and must not be a photosensitiser (that is promote chemical change through indirect photolysis due to exchange of molecular energy with the test chemical). To this end, acetonitrile and methanol (MeOH) are the co-solvents of choice (e.g. aqueous solutions comprising 99.5% H2O/0.5% CH3CN), but not acetone because of its strong photosensitizing effect. Natural waters contain dissolved organic carbon (DOC) and also vary in pH (typically between pH 5 and 8). Both DOC and pH can have a marked effect on the photochemical degradation of organic pollutants and for realistic photochemical studies to be undertaken both factors need to be taken into account, particularly for ionisable chemicals that may dissociate in water depending on their acid dissociation constant (pKa) and the pH of the water. DOC comprises the principal light absorbing component in surface waters (comprising humic and fulvic acids) that provide the natural colouration of waters arising in organic-rich environments. Direct photochemical reactions can be slowed by the shielding effect of DOC, but additionally sensitised reactions may occur by indirect light-induced reactions of the dissolved material. One example of this type of light-induced reaction with DOC, is the possible coupling reactions of triclosan with humic matter, to form insoluble particles.42 Triclosan is a relatively common anti-microbial agent added to hand-washes, toothpaste and other personal care products and is frequently detected in surface waters throughout Europe and North America. Direct phototransformation is responsible for the majority of triclosan loss in surface waters43 and photobyproducts include 2,4-dichlorophenol and 2,8dichlorodibenzo-p-dioxin.42 Only the anion of triclosan is susceptible to significant phototransformation and as its pKa is 7.9 (weak acid), then the anion will predominate in water with a pH of 48. The photochemical reaction pathway is illustrated in Figure 7 and provides a useful example of photochemistry that is pH dependent and where DOC appears to inﬂuence the reaction by-products. 6.6.3 Photochemistry of Brominated Flame Retardants (BFRs) Brominated ﬂame retardants (BFRs) are an interesting and important group of chemicals that are widely incorporated into plastics, foams, electrical items and furnishings to reduce the risk of ﬁre. Indeed, it is likely that the polyurethane foam used in an upholstered chair in which you maybe sitting, will have been treated with a BFR in order to conform to ﬁre-safety standards. However, their widespread use, release

306

Chapter 6 Triclosan Cl

OH O

Cl OH

pKa = 7.9 O-

2,4-DCP

O hv

Cl Triclosan-anion

hv DOC

2,8-DCDD

Coupling reaction with humic material

Figure 7

Direct photochemical reaction pathway of Triclosan (5-chloro-2-(2,4-dichlorophenoxy)phenol) in natural water (3–100 mM of Triclosan in pHbuﬀered solutions), irradiated with a Hg-vapour lamp. Only the triclosan phenolate anion (present in water at pH 48) is susceptible to significant photochemical transformation and will also react through indirect photolysis (not shown) with singlet oxygen (1O2) and OH radicals to form 2,4-dichlorophenol (2,4-DCP). The photoproduct 2,8-dichlorodibenzo-p-dioxin (2,8DCDD) belongs to a notoriously persistent group of chemicals commonly referred to as ‘dioxins’ From Latch et al.42 courtesy of SETAC press

into the environment and occurrence in wildlife and human blood has resulted in growing concern about these groups of chemical.44,45 BFRs function by decomposing under heat to give rise to Brd radicals that react with hydrocarbons (ﬂammable gases), given oﬀ by the burning substance, to form HBr molecules. These, in turn, ‘mop-up’ high energy OHdand Hd radicals formed as part of the combustion chain reaction, thereby effectively retarding or slowing the combustion process. The effectiveness of a BFR is therefore dependent on the quantity of Brd radicals that can be released and the control of their release. Ideally, for the latter, the BFR should decompose at a lower temperature than that of the polymer in which it is incorporated and this will determine, among other things, the type of BFR used in a particular application.45 Two important BFRs are the polybrominated diphenyl ethers (PBDEs) and tetrabromobisphenol A (TBBPA) and their respective structures are presented in Figure 8. PBDEs are utilised as additives for treating plastics and polymers used within electronic components and

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Environmental Organic Chemistry PBDE structure

TBBPA

xBr

yBr

CH3

OH Br

Figure 8 Generic structure of the polybrominated diphenyl ether (PBDE) and the structure of tetrabrominated bisphenol-A (TBBPA). Note the x and y on the PBDE structure represent the number of bromine atoms around each phenyl ring (where x ¼ 1–5 and y ¼ 1–5). The maximum number of 5 bromine atoms on each ring results in the fully brominated, decabrominated diphenyl ether

electrical equipment in general, and are also used extensively as additives in textiles and ﬂexible foams, whereby they are dissolved in the polymer before it is formed into shape or blown into its foam structure. As a result, PBDEs, will, over time, volatilise from the equipment or furnishing into which they have been added. TBBPA, on the other hand, is primarily used as a reactive ﬂame retardant i.e. it is covalently bound into mainly epoxy and polycarbonate resins used in the manufacture of printed circuit boards within electronic equipment. 6.6.3.1 Demand for BFRs. PBDEs are grouped into hom*ologues according to the number of bromine atoms around the phenyl rings and are available commercially as three different formulations; pentaBDE, octaBDE and decaBDE, whereby the penta- and octaBDE contain between 10 and 20 individual congeners and the decaBDE formulation largely comprising of the single congener, decabrominated diphenyl ether. Worldwide demand of PBDEs is high, and is estimated to be 7500 t for pentaBDE, 3760 t for octaPBDE and 56100 t for decaBDE for 2001,46 although quantities of the pentaBDE and octaPDE are likely to change following a ban by the European Union and a voluntary phase-out by one of the main industrial producers in the US. Concern arises due to PBDE occurrence in human blood and breastmilk, with evidence of increasing PBDE levels in wildlife.44 The physical–chemical properties of the PBDEs are analogous to ‘legacy’ persistent organic pollutants such as polychlorinated biphenyls (PCBs), previously used in the electrical industry. For example, PBDEs are semi-volatile (although the decabrominated congener is considered nonvolatile at ambient temperatures) with log Kow values of 5.9–6.2 for tetraBDE, 6.5–7.0 for pentaBDE, 8.4–8.9 for octaBDE and 10 for decaBDE;46 evidence of their propensity for movement in the environment and ability to strongly partition to fats/lipids. It is estimated that worldwide demand for TBBPA in 2001 was some 120,000 tonnes47 and concern over this chemical (and related phenols) has arisen due to its occurrence in sewage sludge, aquatic sediments and ﬁsh.

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6.6.3.2 PBDE and TBBPA Photochemistry. For aryl-halides, the bromine-carbon bond is not as strong as the chorine-carbon bond, with a typical bond-dissociation energy of the bromine-carbon bond of 301 kJ mol1 compared to 360 kJ mol1 for the chlorine-carbon bond.48 Therefore, unlike PCBs, many PBDEs are more susceptible to loss of the halogen atoms and photochemical decomposition has been recognised as an important route for both debromination and transformation for various PBDE congeners present in the wider environment. Several studies have investigated the photochemical degradation of PBDEs46,49 to determine rates of degradation and identify photodegradation byproducts in aqueous solutions under simulated light conditions. Figure 9 shows the absorption spectra (e) of seven PBDE congeners dissolved in tetrahydrofuran (several of the higher brominated congeners are only sparingly soluble in MeOH/H2O) and exposed to a UV-light source in the study of Eriksson et al.46 Clearly the higher brominated congeners absorb light at longer wavelengths, possessing higher lmax values than less brominated congeners. As such, the rate of degradation of PBDEs by UV light in the sunlight region was found to be dependent on the (×10-4) 2.5

3000

2500

2.0

1500

1.5 8

1.0

1000

0.5

500

0 290

Js-1cm-2

ε (cm-1 M -1)

2000

300

310

320 λ (nm)

330

340

0 350

Figure 9 Absorbance spectra (expressed as e, the molar absorbance coefﬁcient) of seven PBDE congeners. The line numbers ar: (1) BDE-209; (2) BDE-206; (3) BCDE-203; (4) BDE-183; (5) BDE-155; (6) BDE-85; (7) BDE-77; line 8 represents the UV light irradiance in units of J s1 cm2 (or W cm2) From Eriksson et al.46 courtesy of the American Chemical Society

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309

degree of bromination, whereby the lower brominated congeners generally degraded at a much slower rate than that of the higher brominated congeners. For example, the observed ﬁrst order rate constant (k) for the photodecomposition of BDE-47 (a tetrabrominated diphenyl ether) in MeOH/H2O solution was 0.07 105 s1 compared to 40 105 s1 for BDE-209 (the decabrominated diphenyl ether). The relatively slow rate of reaction of BDE-47 (approximately 600-fold less than BDE-209) can be largely attributed to its lower absorbance of light at longer wavelengths, which in turn may help to explain the prevalence of PBD-47 in environmental samples. Interestingly, Eriksson et al.46 observed that the decaBDE decomposed to lower brominated diphenyl ethers and also tentatively identiﬁed additional photodegradation by-products including methoxylated polybrominated dibenzofurans through interpretation of mass spectra following analysis of the photolysis solution by GC-MS. A similar aqueous photochemistry study on TBBPA47 has found much faster rates of transformation for this chemical relative to the PBDEs. However, the rate was dependent on the pH of the solution, with phototransformation favouring the anion of TBBPA, present at a significant fraction when the pH of the test solution was 48 (the pKa of TBBA is B7.5–8.5). At pH 10, lmax and emax were 310 nm and 9170 cm1 M1 respectively, compared to 290 nm and B2000 cm1 M1 at pH 5.5, illustrating that at the higher pH the TBBA-anion is absorbing strongly at a higher wavelength compared to TBBPA and therefore more likely to be photodegraded under natural sunlight. The value of k for the photodecomposition of TBBPA in an aqueous solution at pH 10 was measured as 0.7 103 s1, compared to 0.033 103 s1 at pH 5.5, and the photodegradates included a number of bromophenols and alkylated bromophenols.

6.7 CONCLUSIONS Knowledge of the partitioning properties of organic contaminants is fundamental to understanding their environmental distribution and hence fate. For many organic chemicals that pose a threat to the environment, there is a lack of robust temperature-dependent partition coefﬁcients, and in many instances these have to be estimated either from other partitioning descriptors or modelled from the molecular properties of the chemical. Partition ‘maps’ provide a useful way to visualise the environmental distribution of a wide range of chemical classes, and to identify those relevant media where a chemical may reside. Increasingly the results from multi-media chemical fate models

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are displayed in this way and partition maps will likely provide a useful tier for chemical risk assessment in the future. Unlike the older ‘legacy’ POPs, such as DDT and PCBs, many current chemicals used in agriculture, industry and for domestic purposes, are generally more susceptible to degradation processes once present in the environment (although, ironically, the persistent nature of DDT can be attributed to its biological degradate p,p 0 -DDE). In some instances these degradates may be more persistent and/or harmful to the environment than the parent chemicals. The photolytic degradation of Triclosan and certain BFRs are a case in point. Interestingly, a recent review on pesticide-degradates highlighted their higher frequency of detection over the parent compounds in surface water sampled across the US.50 These pesticide-degradates typically had longer half-lives and lower log Koc values than the parent pesticides, making them relatively long-lived and more mobile in the environment. Future research efforts will be required to understand the fate of these chemicals and the threat they pose to both humans and the environment.

6.8 QUESTIONS (i) A pesticide has an experimentally determined Kd value of 5.6, calculate its Koc value assuming the soil has a 2% organic carbon content. What type of polymeric material comprises a large proportion of the organic carbon fraction in soil? Describe how you could operationally separate its base-soluble fraction? (ii) The Henry’s Law constant for a pesticide at 25 oC has been experimentally determined as 0.008 Pa m3 mol1. Calculate Kaw (dimensionless Henry’s Law constant)? If this pesticide has a log Kow 2.9, which environmental compartment is the chemical likely to reside in (use the partition map of Figure 3 to help with your answer)? (iii) The pesticide outlined in Question (i) and (ii) does not absorb sunlight in aqueous solution, however upon dissociation its anion strongly absorbs wavelengths 4300 nm. If this pesticide has a pKa of 6 and is present in surface waters with a pH B8, describe the probable fate of this chemical.

REFERENCES 1. D. Ashton, M. Hilton and K.V. Thomas, Sci. Tot. Environ., 2004, 333, 167–184.

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2. A. Bruchet, C. Hochereau, C. Picard, V. Deottignies, J.M. Rodrigues and M.L. Janex-Habibi, Water Sci. Technol., 2005, 52, 53–61. 3. P.A. Blackwell, H.C. Lutzhoft, H.P. Ma, B. Halling-Sorensen, A.B.A. Boxall and P. Kay, J. Chromatog. A, 2004, 1045, 111–117. 4. J. Yu, R.C. Flagan and J.H. Seinfeld, Environ. Sci. Technol., 1998, 32, 2357–2370. 5. I.G. Kavouras, N. Mihalopoulos and E.G. Stephanou, Environ. Sci. Technol., 1999, 33, 1028–1037. 6. M.P. Fraser, G.R. Cass, B.R.T. Simoneit and R.A. Rasmussen, Environ. Sci. Technol., 1997, 31, 2356–2367. 7. R.A. Larson and E.J. Weber, Reaction Mechanisms in Environmental Organic Chemistry, Lewis Publishers, CRC Press LLC, Florida,1994. 8. C.T. Chiou, D.E. Kile, D.W. Rutherford, G. Sheng and S.A. Boyd, Environ. Sci. Technol., 2000, 34, 1254–1258. 9. L. Li, Z.Y. Zhao, W.L. Huang, P. Peng, G.Y. Sheng and J.M. Fu, Org. Geochem., 2004, 35, 1025–1037. 10. C. Maldonado, J.M. Bayona and L. Bodineau, Environ. Sci. Technol., 1999, 33, 2693–2702. 11. M.B. Yunker and R.W. Macdonald, Org. Geochem., 2003, 34, 1525–154. 12. European Chemicals Bureau (http://ecb.jrc.it/). 13. R.P. Schwarzenbach, P.M. Gschwend and D.M. Imboden, Environmental Organic Chemistry, 2nd edn, Wiley Interscience, Wiley, NJ, 2003. 14. D. Mackay, Multimedia Environmental Models – The Fugacity Approach, Lewis Publishers, Michigan, USA, 1991. 15. D. Mackay, Environ. Sci. Technol., 1982, 16, 274–278. 16. K. Borga˚, A.T. Fisk, P.F. Hoekstra and D.C.G. Muir, Environ. Toxicol. Chem., 2004, 23, 2367–2385. 17. R.S. Boethling, P.H. Howard and W.M. Meylan, Environ. Toxicol. Chem., 2004, 23, 2290–2308. 18. S.W. Karickhoﬀ, Chemosphere, 1981, 10, 833–849. 19. R. Seth, D. Mackay and J. Muncke, Environ. Sci. Technol., 1999, 33, 2390–2396. 20. H.A. Bamford, D.L. Poster and J.E. Baker, Environ. Toxicol. Chem., 1999, 18, 1905–1912. 21. L. Sahsuvar, P.A. Helm, L.M. Jantunen and T.F. Bidleman, Atmos. Environ., 2003, 37, 983–992. 22. J.L. Barber, G.O. Thomas, R. Bailey, G. Kerstiens and K.C. Jones, Environ. Sci. Technol., 2004, 38, 3892–3900. 23. S.N. Meijer, M. Shoeib, K.C. Jones and T. Harner, Environ. Sci. Technol., 2003, 37, 1300–1305.

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24. A. Finizio, D. Mackay, T. Bidleman and T. Harner, Atmos. Environ., 1997, 31, 2289–2296. 25. T. Harner and D. Mackay, Environ. Sci. Technol., 1995, 29, 1599– 1606. 26. P. Ko¨mp and M.S. McLachlan, Environ. Toxicol. Chem., 1997, 16, 2433–2437. 27. T. Harner and T.F. Bidleman, J. Chem. Eng. Data, 1998, 43, 40–46. 28. M. Shoeib and T. Harner, Environ. Toxicol. Chem., 2002, 21, 984–990. 29. H.H. Hung, C.J. Halsall, P. Blanchard, H.H. Li, P. Fellin, G. Stern and B. Rosenberg, Environ. Sci. Technol., 2002, 36, 862–868. 30. C.J. Halsall, Environ. Poll., 2004, 128, 163–175. 31. Stockholm Convention on Persistent Organic Pollutants (http:// www.pops.int/). 32. T. Gouin, D. Mackay, E. Webster and F. Wania, Environ. Sci. Technol., 2000, 34, 881–884. 33. J. Arey, B. Zielinska, R. Atkinson, A.M. Winer, T. Ramdahl and J.N. Pitts, Atmos. Env., 1986, 20, 2339–2345. 34. H. Soderstrom, J. Hajslova, V. Kocourek, B. Siegmund, A. Kocan, W. Obiedzinski, M. Tysklind and P.A. Bergqvist, Atmos. Env., 2005, 39, 1627–1640. 35. F. Reisen and J. Arey, Environ. Sci. Technol., 2005, 39, 64–73. 36. M.M. McGuire, D.L. Carlson, P.J. Vikesland, T. Kohn, A.C. Grenier, L.A. Langley, A.L. Roberts and D.H. Fairbrother, Anal. Chim. Acta, 2003, 496, 301–303. 37. P. Kla´n and I. Holoubek, Chemosphere, 2002, 46, 1201–1210. 38. W.F. Spencer, J.D. Adams, R.E. Hess, T.D. Shoup and R.C. Spear, Agric. Food Chem., 1980, 28, 366–371. 39. T.A. Ternes, A. Joss and H. Siegrist, Environ. Sci. Technol., 2004, 38, 393A–399A. 40. A. Hakami, M.S. Bergin and A.G. Russell, Environ. Sci. Technol., 2004, 38, 6748–6759. 41. J.H. Seinfeld and J.F. Pankow, Ann. Rev. Phys. Chem., 2003, 54, 121–140. 42. D.E. Latch, J.L. Packer, B.L. Stender, J. Vanoverbeke, W.A. Arnold and K. McNeill, Environ. Toxicol. Chem., 2005, 24, 517–525. 43. H. Singer, S. Muller, C. Tixier and L. Pillonel, Environ. Sci. Technol., 2002, 36, 4998–5004. 44. R.A. Hites, Environ. Sci. Technol., 2004, 38, 945–956. 45. F. Rahman, K.H. Langford, M.D. Scrimshaw and J.N. Lester, Sci. Tot. Environ., 2001, 275, 1–17.

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46. J. Eriksson, N. Green, G. Marsh and A. Bergman, Environ. Sci. Technol., 2004, 38, 3119–3125. 47. J. Eriksson, S. Rahm, N. Green, A. Bergman and E. Jakobsson, Chemosphere, 2004, 54, 117–1126. 48. R.T. Morrison and R.N. Boyd, Organic Chemsitry, 3rd edn, Allyn and Bacon Inc., USA, 1973. 49. I. Hua, N. Kang, C.T. Jafvert and J.R. Fabrega-Duque, Environ. Toxicol. Chem., 2003, 22, 798–804. 50. A.B.A. Boxall, C.J. Sinclair, K. Fenner, D. Kolpin and S.J. Maund, Environ. Sci. Technol., 2004, 38, 369A–375A.

CHAPTER 7

Biogeochemical Cycling of Chemicals ROY M. HARRISON Division of Environmental Health & Risk Management, School of Geography, Earth & Environmental Sciences, University of Birmingham, Edgbaston, Birmingham B15 2TT, UK

7.1 INTRODUCTION: BIOGEOCHEMICAL CYCLING The earlier chapters of this book have followed the traditional subdivision of the environment into compartments (e.g. atmosphere, oceans, etc.). While the sub-divisions accord with human perceptions and have certain scientiﬁc logic, they encourage the idea that each compartment is an entirely separate entity and that no exchanges occur between them. This, of course, is far from the truth. Important exchanges of mass and energy occur at the boundaries of the compartments and many processes of great scientiﬁc interest and environmental importance occur at these interfaces. A physical example is that of transfer of heat between the ocean surfaces and the atmosphere, which has a major impact upon climate and a great inﬂuence upon the general circulation of the atmosphere. A chemically based example is the oceanic release of dimethylsulﬁde to the atmosphere, which may, through its decomposition products, act as a climate regulator (see Chapters 2 and 4). Pollutants emitted into one environmental compartment will, unless carefully controlled, enter others. Figure 1 illustrates the processes affecting a pollutant discharged into the atmosphere.1 As mixing processes dilute it, it may undergo chemical and physical transformations before depositing in rain or snow (wet deposition) or as dry gas or particles (dry deposition). The deposition processes cause pollution of land, freshwater, or the seas, according to where they occur. Similarly, pollutants discharged into a river will, unless degraded, enter the seas. Solid wastes are often disposed into a landﬁll. Nowadays these are

314

315

Biogeochemical Cycling of Chemicals Dry transformations Air concentrations Initial mixing

Wet transformations

Transport and diffusion Scavenging

Total emissions Scavenging

Manmade

Natural

Dry deposition

Wet deposition

Figure 1 Schematic diagram of the atmospheric cycle of a pollutant1 (Reprinted from Environmental Science and Technology by permission of the American Chemical Society.)

carefully designed to avoid leaching by rain and dissemination of pollutants into groundwaters, which might subsequently be used for potable supply. In the past, however, instances have come to light where insufﬁcient attention was paid to the potential for groundwater contamination, and serious pollution has arisen as a result. Another important consideration regarding pollutant cycling is that of degradability, be it chemical or biological. Chemical elements (other than radioisotopic forms) are, of course, non-degradable and hence once dispersed in the environment will always be there, although they may move between compartments. Thus, lead, for example, after emission from industry or motor vehicles, has a rather short lifetime in the atmosphere, but upon deposition causes pollution of vegetation, soils, and waters.2 On a very long timescale, lead in these compartments will leach out from soils and transfer to oceans, where it will concentrate in bottom sediments. Some chemical elements undergo chemical changes during environmental cycling, which completely alter their properties. For example, nitrate added to soil as fertilizer can be converted to gaseous nitrous oxide by biological denitriﬁcation processes. Nitrous oxide is an unreactive gas with a long atmosphere lifetime, which is destroyed only by breakdown in the stratosphere. As will be seen later, nitrogen in the environment may be present in a wide range of valence states, each conferring different properties. Some chemical compounds are degradable in the environment. For example, methane (an important greenhouse gas) is oxidized via carbon

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monoxide to carbon dioxide and water. Thus, although the chemical elements are conserved, methane itself is destroyed and were it not continuously replenished would disappear from the environment. The breakdown of methane is an important source of water vapour in the stratosphere, illustrating another, perhaps less obvious, connection between the cycles of different compounds. Degradable chemicals that cease to be used will disappear from the environment. Polychlorinated biphenyls (PCBs) are no longer used industrially to any significant degree, having been replaced by more environmentally acceptable alternatives. Their concentrations in the environment are decreasing, although because of their slow degradability (i.e. persistence), it will take many years before their levels decrease below analytical detection limits. The transfer of an element between different environmental compartments, involving both chemical and biological processes, is termed biogeochemical cycling. The biogeochemical cycles of the elements lead and nitrogen will be discussed later in this chapter. 7.1.1 Environmental Reservoirs To understand pollutant behaviour and biogeochemical cycling on a global scale, it is important to appreciate the size and mixing times of the different reservoirs. These are given in Table 1. The mixing times are a very approximate indication of the timescale of vertical mixing of the reservoir.3 Global mixing can take very much longer as this involves some very slow processes. These mixing times should be treated with considerable caution as they oversimplify a complex system. Thus, for example, a pollutant gas emitted at ground level mixes in the boundary layer (approximately 0–1 km altitude) on a timescale typically of hours. Mixing into the free troposphere (1–10 km) takes days, while mixing into the stratosphere (10–50 km) is on the timescale of several years. Thus, no one timescale describes atmospheric vertical mixing, and the same applies to other reservoirs. Such concepts are useful, however, Table 1

Size and vertical mixing of various reservoirs (from Brimblecombe3) Mass (kg)

Biosphere Atmosphere Hydrosphere Crust Mantle Core

4.2 5.2 1.4 2.4 4.0 1.9

10 1018 1021 1022 1024 1024

Mixing time (years) 60 0.2 1600 43 107 4 108

317

Concentration of lead (pg/g)

Biogeochemical Cycling of Chemicals

LEAD 100

0 1800 1900 Date of deposition of snow or ice

2000

Figure 2 Changes in lead concentrations in snow/ice deposited in central Greenland from 1773 to 1992 (Adapted from Candelone et al.4)

when considering the behaviour of trace components. For example, a highly reactive hydrocarbon emitted at ground level will probably be decomposed in the boundary layer. Sulfur dioxide, with an atmospheric lifetime of days, may enter the free troposphere but it is unlikely to enter the stratosphere. Methane, with a lifetime of several years, extends through all of the three regions. It should be noted from Table 1 that the atmosphere is a much smaller reservoir in terms of mass than the others. The implication is that a given pollutant mass injected into the atmosphere will represent a much larger proportion of total mass than in other reservoirs. Because of this, and the rather rapid mixing of the atmosphere, global pollution problems have become serious in relation to the atmosphere before doing so in other environmental media. The converse also tends to be true, that once emissions into the atmosphere cease, or diminish, the beneﬁcial impact is seen on a relatively short timescale. This has been seen in relation to lead, for instance, where lead in Antarctic ice (derived from snow) has shown a major decrease resulting from diminishing emissions from industry and use of leaded petrol4 (Figure 2). Improved air quality in relation to CFCs will take longer to achieve because of the much longer atmospheric lifetimes (4100 years) of some species (see Chapter 1). 7.1.2 Lifetimes A very useful concept in the context of pollutant cycling is that of the lifetime or residence time of a substance in a given reservoir. We can think in terms of substances having sources, magnitude S, and sinks, magnitude R. At equilibrium R¼S

318

Chapter 7

An analogy is with a bath; the inﬂow from a tap (S) is equal to the outﬂow (R) when the bath is full. An increase in S is balanced by an increase in R. If the total amount of substance A in the reservoir (analogy ¼ mass of water in the bath) is A, then the lifetime, t is deﬁned by t¼

A ðkgÞ S ðkg s1 Þ

ð7:1Þ

In practical terms, the lifetime is equal to the time taken for the concentration to fall to 1/e (where e is the base of natural logarithms) of its initial concentration, if the source is turned oﬀ. If the removal mechanism is a chemical reaction, its rate may be described as follows R0 ¼

d½A ¼ k½A dt

ð7:2Þ

(In this case d[A]/dt describes the rate of loss of A if the source is switched oﬀ; obviously with the source on, at equilibrium d[A]/dt¼0.) The latter part of Equation (7.2) assumes ﬁrst order decay kinetics, i.e. the rate of decay is equal to the concentration of A, termed [A], multiplied by a rate constant, k. As discussed later this is often a reasonable approximation. Taking Equation (7.1) and dividing both numerator and denominator by the volume of the reservoir, allows it to be rewritten in terms of concentration. Thus t¼

½A ðkg m3 Þ S0 ðkg m3 s1 Þ

ð7:3Þ

½A ¼ k1 k½A

ð7:4Þ

since S 0 ¼ R 0 t¼

Thus the lifetime of a constituent with a ﬁrst order removal process is equal to the inverse of the ﬁrst order rate constant for its removal. Taking an example from atmospheric chemistry, the major removal mechanism for many trace gases is reaction with hydroxyl radical, OH. Considering two substances with very different rate constants5 for this reaction, methane and nitrogen dioxide CH4 þ OH - CH3 þ H2O d ½CH4 ¼ k2 ½CH4 ½OH dt

k2 ¼ 6:2 1015 cm3 molec s1

(7.5) ð7:6Þ

Biogeochemical Cycling of Chemicals

NO2 þ OH ! HNO3

k2 ¼ 1:4 1011 cm3 molec s1

319

ð7:7Þ

Making the crude assumption of a constant concentration of OH radical6 (more justiﬁable for the long-lived methane, for which ﬂuctuations in OH will average out, than for short-lived nitrogen dioxide) d ½CH4 ¼ k2 ½CH4 ½OH dt ¼k01 ½CH4 where k01 ¼ k2[OH]

Worked example What are the atmospheric lifetimes of CH4 and NO2 if the diurnally averaged concentration of OH radical is 1 106 molec cm3? k 01 ¼6:2 1015 1 106 ¼6:2 109 s1 Then from Equation (7.4) t ¼k1 ¼ð6:2 109 Þ1 s ¼5:1 years for CH4 By analogy, for nitrogen dioxide, the lifetime t ¼ 20 h This general approach to atmospheric chemical cycling has proved useful in many instances. For example, measurements of atmospheric concentration, [A], for a globally mixed component may be used to estimate source strength, since S 0 ¼ R0 ¼

d½A ¼ k2 ½A½OH dt

and S ¼ S0 V where V is the volume of atmosphere in which the component is mixed. Source strengths estimated in this way, for example, for the compound methyl chloroform, CH3CCl3, known to destroy stratospheric ozone, may be compared with known industrial emissions to deduce whether natural sources contribute to the atmospheric burden.

320

Chapter 7

7.1.2.1 Inﬂuence of Lifetime on Environmental Behaviour. Some knowledge of environmental lifetimes of chemicals is very valuable in predicting their environmental behaviour. In relation to the atmosphere, there is an interesting relationship between the spatial variability in the concentrations of an atmospheric trace species and its atmospheric lifetime.3 Compounds such as methane and carbon dioxide with a long lifetime with respect to removal from the atmosphere by chemical reactions or dry and wet deposition (see Section 7.2 of this chapter) show little spatial variability around the globe, as their atmospheric lifetime (several years) exceeds the timescale of mixing of the entire troposphere (of the order of a year). On the other hand, for a short-lived species such as nitrogen dioxide, removal by chemical means or dry or wet deposition occurs much more quickly than atmospheric mixing and hence there is very large spatial variability, with concentrations sometimes exceeding 100 ppb in urban areas, while remote atmosphere concentrations can be a the level of a few parts per trillion. By analogy, short-lived species also show a much greater hour-to-hour and day-today variation at a given measuring point than long-lived species for which local sources impact only to a modest degree on the existing background concentration. This illustration using the atmosphere can be taken somewhat further in relation to other environmental media. Lifetimes of highly soluble species such as sodium and chloride in the oceans are long compared to the mixing times and therefore variations in salinity across the world’s oceans are relatively small (see Chapter 4). In contrast, where soils are concerned, mixing times will generally far exceed lifetimes and extreme local hot spot concentrations can be found where soils have become polluted. Lifetime also inﬂuences the way in which we study the environmental cycles of pollutants. In the case of reactive atmospheric pollutants, it is the reaction rate, or rate of dry or wet deposition, which determines the lifetime. We are therefore concerned mainly with the rates of these processes in determining the atmospheric cycle. In the case of longerlived species, such as persistent organic compounds like PCBs and dioxins, chemical reaction rates are rather slow and these compounds can approach equilibrium between different environmental media such as the atmosphere and surface ocean or the atmosphere and surface soil, with evaporation exceeding deposition during warmer periods and wet and dry depositions replacing the contaminant in the soils or oceans in cooler weather conditions. Both the kinetic approach dealing with reaction rates and the thermodynamically based approach considering partition between environmental media will be introduced in this

321

Biogeochemical Cycling of Chemicals

chapter. In general the kinetic or reaction rate approach will be most appropriate to the study of short-lived reactive substances, while the equilibrium approach will be more applicable to long-lived substances. 7.2 RATES OF TRANSFER BETWEEN ENVIRONMENTAL COMPARTMENTS 7.2.1 Air–Land Exchange The land surface is an efﬁcient sink for many trace gases. These are absorbed or decomposed on contact with plants or soil surfaces. Plants can be particularly active because of their large surface area and ability to absorb water-soluble gases. The deposition process is crudely described by the deposition velocity, nd, n d ðcm s1 Þ ¼

Flux ðmg m2 s1 Þ Atmospheric concentration ðmg m3 Þ

The term ﬂux is analogous to a ﬂow of material, in this case expressed as micrograms of substance depositing per square metre of ground surface per unit time. In the case of rough surfaces the square metre of area refers to the area of hypothetical horizontal ﬂat surface beneath the true surface rather than the sum of the area of all the rough elements such as plant leaves, which make up the true surface. Since the deposition process itself causes a gradient in atmospheric concentration, nd is deﬁned in relation to a reference height, usually 1 m, at which the atmospheric concentration is measured. For reasons described later, nd is not constant for a given substance, but varies according to atmospheric and surface conditions. However, some typical values are given in Table 2, which exemplify the massive variability. Table 2

Some typical values of deposition velocity

Pollutant

Surfaces

Deposition velocity (cm s1)

SO2 SO2 SO2 SO2 O3 O3 O3 HNO3 CO Aerosol (42.5 mm)

Grass Ocean Soil Forest Dry grass Wet grass Snow Grass Soil Grass

1.0 0.5 0.7 2.0 0.5 0.1 0.1 2.0 0.05 0.1

322

Chapter 7

For some trace gases, for example, nitric acid vapour, dry deposition represents a major sink mechanism. In this case the process may have a major impact upon atmospheric lifetimes.

Worked example Dry deposition is frequently the main sink for ozone in the rural atmospheric boundary layer. What is the lifetime of ozone with respect to this process? Assuming a typical dry deposition velocity of 1 cm s1 and a boundary layer height of 1000 m, (H), d Fluxðmg m2 s1 Þ ½O3 ðmg m3 s1 Þ ¼ dt Mixing depth ðmÞ n d ½O3 ¼ H ¼k½O3 where k ¼ nd/H By analogy with Equation (7.4), H nd 1000 ¼ 0:01 s ¼28 h

t¼

Thus taking the boundary layer as a discrete compartment, the lifetime of ozone with respect to dry deposition is around 1 day. The lifetime in the free troposphere (the section of the atmosphere above the boundary layer) is longer, being controlled by transfer processes in and out, and chemical reactions. The stratosphere lifetime of ozone is controlled by photochemical and chemical reaction processes. Dry deposition processes are best understood by considering a resistance analogue. In direct analogy with electrical resistance theory, the major resistances to deposition are represented by three resistors in series. Considering the resistances in sequence, starting well above the ground, these are as follows: (i) ra, the aerodynamic resistance describes the resistance to transfer downwards towards the surface through normally turbulent air;

Biogeochemical Cycling of Chemicals

323

(ii) rb, the boundary layer resistance describes the transfer through a laminar boundary layer (approximately 1 mm thickness) at the surface; (iii) rs, the surface (or canopy) resistance is the resistance to uptake by the surface itself. This can vary enormously, from essentially zero for very sticky gases such as HNO3 vapour, which attach irreversibly to surfaces, to very high values for gases of low water solubility, which are not utilised by plants (e.g. CFCs). Since these resistances operate essentially in series, the total resistance, R, which is the inverse of the deposition velocity, is equal to the sum of the individual resistances. R¼

1 ¼ ra þ rb þ rs nd

ð7:8Þ

Some trace gases have a net source at the ground surface and diffuse upwards; an example is nitrous oxide. Whether the ﬂux is downward or upward, it is driven by a concentration gradient in the vertical, dc/dz. The relationship between ﬂux, F, and concentration gradient is dc dz where Kz is the eddy diffusivity in the vertical (a measure of the atmospheric conductance). Fluxes, and thus deposition velocities, can be estimated by measurement of a concentration gradient simultaneously with the eddy diffusivity.7 It is usually assumed that trace gases transfer in the same manner as sensible heat (i.e. convective heat transfer, not radiative or latent heat) or momentum. Thus, the eddy diffusivity for either of these parameters is measured usually from simple meteorological variables (gradients in temperature and wind speed). A few substances are capable of showing both upward and downward ﬂuxes. An example is ammonia. Ammonium in the soil, NH41, is in equilibrium with ammonia gas, NH3(g)

F ¼ Kz

NH41 þ H2O " NH3(g) þ H3O1

(7.9)

when atmospheric concentrations of ammonia exceed equilibrium concentrations at the soil surface (known as the compensation point), the net ﬂux of ammonia is downwards. When atmospheric concentrations are below the equilibrium value, ammonia is released into the air.8

324

Chapter 7

7.2.2 Air–Sea Exchange The oceans cover some two-thirds of the Earth’s surface and consequently provide a massive area for exchange of energy (climatologically important) and matter (an important component of geochemical cycles). The seas are a source of aerosol (i.e. small particles), which transfer to the atmosphere. These will subsequently deposit, possibly after chemical modiﬁcation, either back in the sea (the major part) or on land (the minor part). Marine aerosol comprises largely unfractionated seawater, but may also contain some abnormally enriched components. One example of abnormal enrichment occurs on the eastern coast of the Irish Sea. Liquid efﬂuents from the Sellaﬁeld nuclear fuel reprocessing plant in west Cumbria are discharged into the Irish Sea by pipeline. At one time, permitted discharges were appreciable and as a result radioisotopes such as 137Cs and several isotopes of plutonium have accumulated in the waters and sediments of the Irish Sea. A small fraction of these radioisotopes were carried back inland in marine aerosol and deposited predominantly in the coastal zone.9 While the abundance of 137 Cs in marine aerosol was reﬂective only of its abundance in seawater (an enrichment factor – see Chapter 4 – of close to unity), plutonium was abnormally enriched due to selective incorporation of small suspended sediment particles in the aerosol. This has manifested itself in enrichment of plutonium in soils on the west Cumbrian coast,10 shown as contours of 2391240Pu deposition (pCi cm2) to soil in Figure 3. The seas may also act as a receptor for depositing aerosol. Deposition velocities of particles to the sea are a function of particle size, density, and shape, as well as the state of the sea. Experimental determination of aerosol deposition velocities to the sea is almost impossible and has to rely upon data derived from wind tunnel studies and theoretical models. The results from two such models appear in Figure 4, in which particle size is expressed as aerodynamic diameter, or the diameter of an aerodynamically equivalent sphere of unit speciﬁc gravity.11,12 If the airborne concentration in size fraction of diameter di is ci, then Total Flux ¼

i X

n d ðdi Þci

where vd (di) is the mean value of deposition velocity appropriate to the size fraction di. Measurements show that while most of the lead, for example, is associated with small, sub-micrometre particles, the larger particles compose the major part of the ﬂux. Airborne concentrations of particulate pollutants are not uniform over the sea. The spatial distribution of zinc over the North Sea13

Biogeochemical Cycling of Chemicals

325

IRISH SEA

10 km

Figure 3

Concentrations of plutonium in soils of West Cumbria (239 1 240Pu to 15 cm depth; pCi cm2). The point marked S indicates the position of the Sellaﬁeld reprocessing works (From Cawse10)

averaged over a number of measurement cruises appears in Figure 5. Spatial patterns of other metals and many artiﬁcial pollutants are similar, reﬂecting the impact of land-based source regions, with concentrations falling towards the north and centre of the sea.

326

Chapter 7

Dry deposition velocity Vd, (cm s−1)

100

Mean windspeed

10 m s−1

Mean windspeed

5 m s−1

0.1

0.01

0.001 0.001

1 0.01 10 0.1 Diameter of unit density particle dp (µm)

100

Figure 4 Calculated values of deposition velocity to water surfaces as a function of particle size and wind speed

Because of its position and relatively high pollution loading, the North Sea is a focus of considerable interest. An inventory of inputs of trace metals (e.g. Pb, Cd, Zn, Cu, etc.) accords similar importance to riverine inputs and atmospheric deposition.14 Controls have been applied to many source categories and total inputs of the metals indicated have in general declined appreciably. One particular example is lead, for which most European countries introduced severe controls on use in gasoline (petrol) during the 1980s and a total ban in 2000; atmospheric concentrations have fallen accordingly. Although the data are less clear, it might be anticipated that concentrations in river water will also decline as a result of reduced inputs from direct atmospheric deposition and in runoﬀ waters from highways and land surfaces. As explained in Chapter 4, the sea may be both a source and a sink of trace gases. The direction of ﬂux is dependent upon the relative concentration in air and seawater.15 If the concentration in air is Ca, the equilibrium concentration in seawater, Cw (equ) is given by Cw(equ) ¼ CaH1

(7.10)

where H is the Henry’s Law Constant. The Henry’s Law Constant can be expressed as follows: H¼

ps Saq

327

Biogeochemical Cycling of Chemicals

Figure 5 Spatial distribution of zinc concentrations (in ng m3) in air over the North Sea during 1989 (From Ottley and Harrison.13)

where ps ¼ saturation vapour pressure and Saq ¼ equilibrium solubility in water.

Worked example For benzene ps ¼ 12.7 kPa at 251C 1:78 1000 ¼ mol m3 78 Calculate H for benzene at 251C. For benzene

Saq ¼ 1:78 g L1 ¼

H¼

ps Saq

12:7 103 ð1:78 103 Þ=78

¼556 Pa m3 mol1

328

Chapter 7

If Cw is the actual concentration of the dissolved gas in the surface seawater and Cw ¼ C(equ) the system is at equilibrium and no net transfer occurs. If, however, there is a concentration difference, DC, where DC ¼ CaH1 Cw there will be a net ﬂux. If CaH1 4 Cw the water is sub-saturated with regard to the trace gas and transfer occurs from air to water. Conversely, gas transfers from supersaturated water to the atmosphere if CaH1 o Cw The rate at which gas transfers occurs is expressed by F ¼ K(T)wDC where K(T)w is termed the total transfer velocity. This can be broken down into component parts as follows: 1 KðTÞw

1 1 þ ¼ rw þ ra akw Hka

where ka and kw are the individual transfer velocities for chemically unreactive gases in air and water phases, respectively and a (¼kreactive/ kinert) is a factor that quantiﬁes any enhancement of gas transfer in the water due to chemical reaction. The terms rw and ra are the resistances to transfer in the water and air phases, respectively, and are directly analogous to the resistance terms in Equation (7.8). For chemically reactive gases, usually racrw and atmospheric transfer limits the overall ﬂux. For less reactive gases the inverse is true and K(T)w R kw; the resistance in the water is the dominant term. Much research has gone into evaluating kw and K(T)w, both in theoretical models, and in wind tunnel and ﬁeld studies. The results are highly wind speed dependent due to the inﬂuence of wind upon the surface state of the sea. The results of some theoretical predictions and experimental studies16 for CO2 (a gas for which kw is dominant) are shown in Figure 6. In addition to dry deposition, trace gases and particles are also removed from the atmosphere by rainfall and other forms of precipitation (snow, hail, etc.), entering land and seas as a consequence. Wet deposition may be simply described in two ways. First,

329

Biogeochemical Cycling of Chemicals Beaufort 2

50 12 40

(10−5 m s−1)

Transfer velocity (cm h−1)

20 4 10

0 0

Wind speed (m s−1)

Figure 6 Air–sea transfer velocities for carbon dioxide at 201C as a function of wind speed at 10 m (m s1 or Beaufort Scale). The graph combines experimental data (points) and a theoretical line (From Watson et al.16) (Reprinted by permission from Nature (London), 349, 145; Copyrightr 1991 Macmillan Magazines Ltd.)

1 KðTÞw

1 1 þ ¼ rw þ ra akw Hka

Typical values of scavenging ratio17 lie within the range 300–2000. Scavenging ratios are rather variable, dependent upon the chemical nature of the trace substance (particle or gas, soluble or insoluble, etc.) and the type of atmospheric precipitation. Incorporation of gases and particles into rain can occur both by in-cloud scavenging (also termed rainout) and below-cloud scavenging (termed washout). Numerical modellers often ﬁnd it convenient to describe wet deposition by a scavenging coefﬁcient, actually a ﬁrst order rate constant for

330

Chapter 7

removal from the atmosphere. Thus, for trace substance A, d½A ¼ L½A dt where L is the washout coefﬁcient, with units of s1. A typical value of L for a soluble substance is 104 s1 although actual values are difﬁcult to measure and are highly dependent upon factors such as rainfall intensity. 7.3 TRANSFER IN AQUATIC SYSTEMS When rain falls over land some drain oﬀ the surface directly into surface water courses in surface runoﬀ. A further part of the incoming rainwater percolates into the soil and passes more slowly into either surface waters or underground reservoirs. Water held in rock below the surface is termed groundwater, and a rock formation that stores and transmits water in useful quantities is termed an aquifer. Water that passes through soil or rock on its way to a river is chemically modiﬁed during transit, generally by addition of soluble and colloidal substances washed out of the ground. Some substances are removed from the water; for example, river water often contains less lead than rainwater; one mechanism of removal is uptake by soil. River waters carry both dissolved and suspended substances to the sea. The concentrations and absolute ﬂuxes vary tremendously. The suspended solids load is largely a function of the ﬂow in the river, which inﬂuences the degree of turbulence and thus the extent to which solids are held in suspension and resuspended from the bed, once deposited. Table 3 shows a comparison of ‘average’ riverine suspended particulate Table 3

A comparison of the concentration of major elements in ‘average’ riverine particulate material and surﬁcial rocks Concentrations (g kg1)

Element

Riverine particulate material

Surface rocks

Al Ca Fe K Mg Mn Na P Si Ti

94.0 21.5 48.0 20.0 11.8 1.1 7.1 1.2 285.0 5.6

69.3 45.0 35.9 24.4 16.4 0.7 14.2 0.6 275.0 3.8

331

Biogeochemical Cycling of Chemicals

matter and surﬁcial rock composition18 for the major elements. Elements resistant to chemical weathering or biological activity (e.g. aluminium, titanium, iron, phosphorus) show some enrichment in the riverine solids, while more soluble elements are subject to weathering and are depleted in the solids, being transported largely in solution (sodium, calcium). Some pollutant elements such as the metals lead, cadmium, and zinc tend to be highly enriched in the solids relative to surﬁcial rocks or soils due to inputs from human activities. The dissolved components of river water typically exhibit significantly higher concentrations than in rainwater19 (Table 4), due to leaching from rocks and soils. Some insight into the processes governing river water composition may be gained from Figure 7. Starting from the point of lowest dissolved salts concentrations, the ratio of Na/(Na þ Ca) approaches one. This is similar to rainwater, and is termed the precipitation dominance regime. It is typiﬁed by rivers in humid tropical areas of the world with very high rainwater inputs and little evaporation. As the dissolved solids concentration increases the ratio Na/(Na þ Ca) declines, indicating an increasing importance for calcium in the rock dominance regime. Here, increased weathering of rock provides the major source of dissolved solids. As dissolved solids increase further, the abundance of calcium decreases relative to sodium as the water becomes saturated with respect to CaCO3, and this compound precipitates. Waters in the evaporation/precipitation regime are typiﬁed by rivers in very arid parts of the world (e.g. River Jordan) and the major seas and oceans of the world.20,21 In total, the rivers of the world carry around 4.2 1012 kg per year of dissolved solids to the oceans and 18.3 1012 kg per year of suspended solids. Table 4

Average concentrations of the major constituents dissolved in rain and river water Concentrations (mg L1)

Constituent 1

Na K1 Mg21 Ca21 Fe Al Cl SO42 HCO3 SiO2 pH

Rain water

River water

1.98 0.30 0.27 0.09

6.3 2.3 4.1 15 0.67 0.01 7.8 11.2 58.4 13.1

3.79 0.58 0.12 5.7

332

Chapter 7 Major oceans

(a)

Black

104

Jordan

Caspian Baltic

103 Colorado Volga Yukon

Mississippi

Columbia

Ganges

Total dissolved solids (mg L−1)

Congo

Orinoco

101

Negro Tefé

(b) Sea water

Evaporation crystallisation 3

102

Series

Rock dominance

Series

101

Precipitation dominance

0.2

0.4

0.6

0.8

1.0

Weight Na+ Weight (Na+ + Ca2+)

Figure 7 The chemistry of the Earth’s surface waters: (a) typical values of the ratio Na1/ (Na1 þ Ca21) as a function of dissolved solids concentrations from various major rivers and oceans and (b) the processes leading to the observed ratios (From Gibbs20) copyrightr 1970, American Association for the Advancement of Science

In slow-moving water bodies such as lakes and ocean basins, suspended solids falling to the bottom produce a well-stratiﬁed layer of bottom sediment. This is stratiﬁed in terms of age with the oldest sediment at the bottom (where when suitably pressurised in can form rock) and the

Biogeochemical Cycling of Chemicals

333

Figure 8 Lead proﬁle in a lake sediment in relation to depth and the year of incorporation (From Davies and Galloway.22)

newest at the top, in contact with the water. If burrowing organisms do not provide too much disturbance (termed bioturbation), the sediment can preserve a record of depositional inputs to the water body. An example is provided by Figure 8 in which lead is analysed in sediment core dated from its radioisotope content.22 The concentration rises from a background around the year 1800, corresponding to the onset of industrialisation. Considerably increased deposition is seen after 1930 due to the introduction of leaded petrol. While some of the lead input is via surface waters, the majority probably arises from atmospheric deposition. 7.4 BIOGEOCHEMICAL CYCLES A general model of a biogeochemical cycle appears in Figure 9. Although biota are not explicitly included, their role is a very important one in mediating transfers between the idealised compartments of the model. For example, the role of marine phytoplankton in transferring sulfur from the ocean to the atmosphere in the form of dimethyl sulﬁde has been highlighted in Chapter 4. Biota play a major role in determining atmospheric composition. Photosynthesis removes carbon dioxide from the atmosphere and replenishes oxygen. In a world without biota, lightning would progressively convert atmospheric oxygen into nitrogen oxides and thence to nitrate, which would reside in oceans. Biota also exert more subtle inﬂuences. In aquatic sediments, micro-organisms often deplete oxygen more quickly than it can be replenished from the overlying water, producing anoxic conditions. This leads to chemical

334

Chapter 7

Atmosphere

3,4

Land

3,4

Freshwater

Oceans

Marine sediment/ rock

Figure 9 Schematic diagram of the major ﬂuxes and compartments in a biogeochemical cycle: (1) runoﬀ; (2) streamﬂow; (3) degassing; (4) particle suspension; (5) wet and dry deposition; (6) sedimentation and (7) remobilisation

reduction of elements such as iron and manganese, which has implications for their mobility and bioavailability. Biological reduction processes in sediments may be viewed as the oxidation of carbohydrate (in its simplest form CH2O) with accompanying reduction of an oxygen carrier. In the ﬁrst instance, dissolved molecular oxygen is used. The reaction is thermodynamically favoured, as reﬂected by the strongly negative DG. CH2O þ O2 - CO2 þ H2O DG ¼ 125.5 kJ mol1 e When all of the dissolved oxygen is consumed, anaerobic organisms take over. Initially, nitrate-reducing bacteria are favoured 1 þ 2CH2O þ NO 3 þ 2H - 2CO2 þ H2O þ NH4 DG ¼ 82.2 kJ mol1 e

Once the nitrate is utilised, sulfate reduction takes over 1 SO2 4 þ H þ 2CH2O - HS þ 2H2O þ 2CO2 DG ¼ 25.6 kJ mol1 e

Biogeochemical Cycling of Chemicals

335

Finally, methane-producing organisms dominate in a sediment depleted in oxygen, nitrate, and sulfate. 2CH2O - CH4 þ CO2 DG ¼ 23.5 kJ mol1 e Thus highly anoxic waters are commonly sources of hydrogen sulﬁde, H2S, from sulfate reduction and of methane (marsh gas). The formation of sulﬁde in sediments has led to precipitation of metal sulfides over geological time, causing accumulations of sulﬁde minerals of many elements, e.g. PbS, ZnS, HgS, etc. 7.4.1 Case Study 1: The Biogeochemical Cycle Of Nitrogen Nitrogen has many valence states available and can exist in the environment in a number of forms, depending upon the oxidising ability of the environment. Figure 10 indicates the most important oxidation states and the relative stability (in terms of free energy formation).23 The oxides of nitrogen represent the most oxidised and least thermodynamically stable forms. These exist only in the atmosphere. Ammonia can exist in gaseous form in the atmosphere but rather rapidly returns to the soil and waters as ammonium, NH41. Fixation of atmosphere N2 by leguminous plants leads to ammonia, NH3. In aerobic soils and aquatic systems, NH3 and NH41 are progressively oxidised by micro-organisms via nitrite to nitrate. The latter is taken up by some biota and used as a nitrogen source in synthesising amino acids and proteins; the most thermodynamically stable form of nitrogen. After the death of the organism, microbiological processes will convert organic nitrogen to ammonium (ammoniﬁcation), which is then available for oxidation or use by plants. Conversion of ammonia to nitrate is termed nitriﬁcation, while denitriﬁcation involves conversion of nitrate to N2. Figure 11 shows an idealised nitrogen cycle. The numbers in boxes represent quantities of nitrogen in the various reservoirs, while the arrows show ﬂuxes.23 It is interesting to note that substances involving relatively small ﬂuxes and burdens can have a major impact upon people. Thus nitrogen oxides, NO, NO2, and N2O are very minor constituents relative to N2 but play major roles in photochemical air pollution (NO2), acid rain (HNO3 from NO2), and stratospheric ozone depletion (N2O). Nitrate from fertilisers represents a very small ﬂux but has major implications in terms of eutrophication of surface waters. 7.4.2 Case Study 2: Aspects of Biogeochemical Cycle of Lead Lead is a simpler case to study than nitrogen due to the small number of available valence states. The major use of lead until recently was a

336

Chapter 7 Oxidation Free energy state of formation ∆G Oxides of nitrogen N2O dinitrogen oxide NO nitric oxide nitrogen dioxide NO2

N2 dinitrogen

+1 +3 +4

less stable

positive

5 NH3 ammonia

−3

negative

NO−2 nitrite 3

+3 2

more stable NO−3 nitrate

+5 3

amino acids, proteins

−3

Reactions 1 fixation 2 nitrification 3 assimilation by piants

4 ammonification 5 denitrification 6 nitrate-containing precipitation often as nitric acid in acid rain

micro-organisms play a part in reactions 1, 2, 4 and 5

Figure 10

Chemical forms and cycle of nitrogen (From O’Neill.23)

tetraalkyl lead gasoline additives in which lead is present as PbIV. The predominant compounds used were tetramethyl lead, Pb(CH3)4, and tetraethyl lead, Pb(C2H5)4. These are lost to the atmosphere as vapour from fuel evaporation and exhaust emissions from cold vehicles, but comprised only 1–4% of lead in polluted air.2 Leaded gasoline also contains the scavengers 1,2-dibromoethane, CH2BrCH2Br, and 1,2-dichloroethane, CH2ClCH2Cl, which convert lead within the engine to lead halides, predominantly lead bromochloride, PbBrCl, in which lead is in the PbII valence state, its usual form in environmental media. About 75% of lead alkyl burned in the engine is emitted as ﬁne particles of inorganic lead halides. Atmospheric emissions of lead arise also from industry; both these and vehicle-emitted lead have declined massively. Figure 12 shows trends in United Kingdom emissions of lead to

337

Biogeochemical Cycling of Chemicals NH3 /NH4+

4.1 × 109

N2O / NO / NO2

1300 2 90

120

30 25

80 fertiliser

SEDIMENTS-

biomass

1.3 ×

104

Organic compounds

Inorganic compounds

biomass

5 × 102

dissolved N compounds

3 × 105

8 × 108 30

2 × 108

30 SEDIMENT

× 109 kg N

reservoirs

× 109 kg N a−1

fluxes

Figure 11 Schematic representation of the biogeochemical cycle of nitrogen, indicating the approximate magnitude of ﬂuxes and reservoirs (After O’Neill.23)

atmosphere from leaded petrol upto 1996, since when emissions from leaded petrol have declined almost to zero.24 Lead emitted to the atmosphere has a lifetime of around 7–30 days and hence may be subject to long-range transport. Concentrations of trace elements in polar ice provide a historical record of atmospheric deposition. Measurements (Figure 2) have shown a marked enhancement in lead accompanying the increase in leaded gasoline usage, and a major decline in recent years attributable to reduced emissions to atmosphere. Atmospheric lead is deposited in wet and dry deposition. Lead is relatively immobile in soil, and agricultural surface soils in the UK exhibit concentrations approximately double those of background soil, which contain ca. 15–20 mg kg1 derived from soil parent materials, other than in areas of lead mineralisation where far greater concentrations can be found. Local perturbations to the cycle of lead can be important. For instance, the lead content of garden soils correlates strongly with the age of the house. This is probably due to the deterioration of leaded paintwork on older houses and the former practices of disposing of household refuse and ﬁre ashes in the garden. Lead is also of low mobility in aquatic sediments and hence the sediment may provide a record of historical lead deposition (see Figure 8).

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Plants can take up lead from soil, thus providing a route of human exposure. Careful research in recent years has established transfer factors, terms the Concentration Factor (CF), where D Concentration of lead in plant ðmg kg1 dry wt:Þ D Concentration of lead in soil ðmg kg1 dry wt:Þ The value of CF for lead is lower than for most metals and is typically within the range 103 102. Much higher values had been estimated from earlier studies, which ignored the importance of direct atmospheric deposition as a pathway for contamination. The direct input from the air to leaves of plants is often as great, or greater than soil uptake.24,25 This pathway may be described by another transfer factor, termed the Air Accumulation Factor (AAF), where

CF ¼

D Concentration of lead in plant ðmg g1 dry wt:Þ D Concentration of lead in air ðmg m3 Þ Values of AAF are plant dependent, due to differences in surface characteristics, but values of 5–40 are typical.25,26 Thus, a plant grown on an agricultural soil with 50 mg kg1 lead will derive 0.25 mg kg1 dry weight lead from the soil (CF ¼ 5 103), while airborne lead of 0.1 mg m3 will contribute 2.0 mg g1 (Rmg kg1) of lead (AAF ¼ 20 m3 g1). Thus, in this instance airborne lead deposition is dominant. The air lead concentration of 0.1 mg m3 was typical of rural areas of the UK until 1985. Since that time, the drastic reduction of lead in gasoline has led to appreciably reduced lead-in-air concentrations in both urban and rural localities. Human exposure to lead arises from four main sources:2,27

AAFðm3 g1 Þ ¼

(i) Inhalation of airborne particles. The adult human respires approximately 20 m3 of air per day. Thus for an urban lead concentration of 0.1 mg m3, intake is 2 mg per day. This is rather efﬁciently absorbed (ca. 70%) and therefore uptake is around 1.4 mg per day in this instance. (ii) Indigestion of lead in foodstuﬀs. The concentrations of lead in food obviously vary between different foodstuﬀs and even between different batches of the same food. Typical fresh weight concentrations (much of the weight of some foods is water) are from 10 to 50 mg Pb kg1. Thus a food consumption of 1.5 kg per day represents an intake of around 50 mg per day and an uptake (10–15% efﬁcient) of around 6 mg per day. (iii) Drinking water and beverages. Concentrations of lead in drinking water vary greatly, related particularly to the presence or

339

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absence of lead in the household plumbing system. Most households in the UK conform to the EC standard of 50 mg L1 and a concentration of 4 mg L1 may be taken as representative. Gastrointestinal absorption of lead from water and other beverages is highly dependent upon food intake. After long fasting, absorptions of 60–70% have been recorded, 14–19% with a short period of fasting before and after the meal, and only 3–6% for drinks taken with a meal. If 15% is taken as typical, for a daily consumption of 1.5 L, intake is 7.5 mg and uptake 1.1 mg. (iv) Cigarette smoking exposes the individual to additional lead. While both individual exposure to lead and the uptake eﬃciencies of individuals are very variable, it is evident that exposure arises from a number of sources and control or human lead intake, if deemed to be desirable, requires attention to all of those sources. An additional pathway of exposure, not easily quantiﬁed, and not included above is ingestion of lead-rich surface dust by hand-to-mouth activity in young children. The above calculations estimate that for a typical adult in a developed country, daily uptake of lead from air, diet, and drinking water is, respectively 1.4 mg, 6 mg, and 1.1 mg. Exposure to lead from all of these sources has fallen rapidly over the past 20–30 years. Figure 12 contrasts the temporal trends in use of lead in petrol (gasoline) and blood leads in the general population of the UK over the period when much of this decline took place. It is interesting to note that from 1971 to 1985 use of lead in petrol was relatively steady, but blood leads declined by a factor of more than two over this period mainly as a response to reductions in dietary exposure, particularly associated with the cessation of use of

Blood Lead (µg / dL)

Blood Lead

Gasoline Lead

8 15 6 10

5 0 1970

1974

1978

1982 Year

1986

1990

1994

Gasoline Lead (Th. Tonnes)

Figure 12 Trends in lead use in petrol (gasoline) and of lead in the blood of the general population in the United Kingdom, 1970–1995

340

Chapter 7

leaded solder to seal food cans. A dramatic reduction in gasoline lead usage occurred at the end of 1985 when the maximum permissible lead content of petrol was reduced from 0.4 g L1 to 0.15 g L1, and there has been a steady reduction in lead use since, with the increases in market penetration of unleaded fuel, until it became mandatory in the year 2000. Despite the ability of a vehicle emitting lead to cause direct lead exposure through the atmosphere, as well as indirect exposure through contamination of food and water, the lack of any obvious step change in blood lead associated with the reduction of lead in petrol shows clearly that at that time leaded petrol was not a major source of direct exposure for the general population. 7.5 BEHAVIOUR OF LONG-LIVED ORGANIC CHEMICALS IN THE ENVIRONMENT A very large number of organic chemicals are manufactured and have a wide variety of uses. In most cases of manufacture, distribution or use, some of the chemical becomes released and dispersed in the environment. In some instances, the chemical is highly reactive and degrades very rapidly; often the main mechanism is breakdown by bacteria or other organisms and there are standard tests to evaluate the biodegradability of synthetic chemicals. On the other hand, many such chemicals, in order to be useful in products, are relatively long-lived and hence may be persistent within the environment. This raises the question of which environmental compartment (air, water, soil, etc.) the chemical will partition into predominantly. Such knowledge is critical to predicting possible environmental effects and directing monitoring programmes. Over the past few years Mackay and colleagues28 have developed methods for predicting the behaviour of chemicals released into the environment and at the simplest level, these are based solely upon an assumption of equilibrium partitioning between the various environmental compartments. The basic principles are described in Chapter 6, and this section is concerned with how these methods are applied in a practical way. The approach depends upon assuming a model environment consisting of speciﬁed volumes of air, water, soil, and sediment (see Table 5). Within those compartments, there are other components such as particles in water, ﬁsh, or soil solids, and some of these are assumed to contain an organic fraction of deﬁned abundance (see Table 5). The presence of this organic matter is important since most organic compounds will preferentially partition into an organic-rich medium. Having deﬁned a model environment, partitioning between the various compartments is predicted on the basis of equilibrium relationships,

341

Biogeochemical Cycling of Chemicals Table 5

Properties of a ‘‘model world’’ for modelling the behaviour of organic chemicals28 Area (m2)

Air Water Soil Sediment

10 1010 9 1010 1010

Depth (m) 1000 20 0.2 0.05

Volume (m3) 10 2 1011 1.8 1010 5 108

Volume fractions Aerosol Particles in water Fish Soil pore air Soil pore water Soil solids Sediment pore water Sediment solids

Advective residence Time (h)

2 1011 5 106 106 0.2 0.3 0.5 0.8 0.2

100 1000 50,000

Organic carbon or lipid mass fraction

Density, kg m3

0.2 0.5

2000 1500 1000

0.02

2400

0.04

2400

which relate to the fundamental physico–chemical properties of the organic chemical under consideration. Important in this context is the Henry’s Law constant, which describes the equilibrium partitioning between air and water that is deﬁned as H ¼ Kaw ¼

p ps ¼ CW Saq

where p is the vapour pressure in air CW, the concentration in water, Saq, the equilibrium aqueous solubility, and ps, the saturation vapour pressure for the pure substance. Partition between compartments within the aqueous environment is typically described through the octanol–water partition coefﬁcient, Kow, or the organic carbon–water partitioning coefﬁcient, Koc (see Chapter 5). This approach assumes that partitioning from water into ﬁsh or suspended solids within the water is determined by the availability of organic matter whose properties as a solvent can be described by a parallel with octanol for which partitioning data are widely available Kow ¼

Concentration in octanol Concentration in water at equilibrium

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Chapter 7

A further partition coefﬁcient is the octanol–air partition coefﬁcient, which can be used to describe the partitioning between the organic matter in soil, for which octanol is taken as a surrogate, and the atmosphere. Thus, concentration in octanol concentration in air at equilibrium Values of these partition coefﬁcients have been measured for a wide variety of substances, but for those which measurements have not been made, there are methods of calculation based upon fundamental properties of the molecule. Frequently, the organic matter comprises only a fraction, f, of the medium. Therefore for partition between air and soil containing 20% organic matter content

Koa ¼

concentration in organic matter concentration in air Com ¼ p Csoil ¼ fp ¼ Koa fp ¼ 0.2 Koa.p

Koa ¼

and Csoil

Worked example What concentration of naphthalene in a soil of 20% organic matter will be in equilibrium with a concentration of 10 ppb vapour in air (water solubility ¼ 31 g m3; saturation vapour pressure ¼ 10.4 Pa and log Kow ¼ 3.37). In a saturated system, Kow ¼ 2:35 103 ¼

Coctanol 31

and Coctanol ¼ 7.28 104 g m3 at equilibrium Then, 7:28 104 10:4 ¼7:00 103 g Pa1 m3

Koa ¼

Biogeochemical Cycling of Chemicals

343

A concentration of 10 ppb in air ¼ 1.013 105 108 Pa ¼ 1.013 103 Pa Coctanol ¼1:013 105 108 Pa ¼1:013 103 Pa ¼1:013 103 Koa ¼1:013 103 7:00 103 ¼7:09 g m3 and Csoil ¼0:2 Coctanol ¼0:2 7:09 ¼1:42 g m3

In order to calculate the equilibrium partitioning of a substance within the model world, a mass balance exercise is conducted. Thus, for a system comprising simply air, water, and soil, the mass balance equation would be M ¼V1 C1 þ V2 C2 þ V3 C3 ¼V1 C1 þ V2 C1 K21 þ V3 C1 K31 ¼C1 ðV1 þ V2 K21 þ V3 K31 Þ where M is the total mass of substance and Vi and Ci represent, respectively the volumes and concentrations of each medium, i. It may be seen that by rearranging this equation, the concentration C1 in one of the media can be estimated from knowledge of the total mass of compound in the system, the volume of each of the three media and the equilibrium partition coefﬁcients between them. It may also be seen that if the various partition coefﬁcients are known, then concentrations in all three compartments, C1, C2, and C3 can be calculated and therefore the overall distribution of the chemical within the environment inferred. For partitioning, for example, into soil, it is assumed that the substance partitions solely into the organic carbon fraction of the soil solids, and since this is a relatively small percentage of the total soil, the mass fraction of organic matter has to be included in the calculation.

344

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This describes one of the most basic kinds of calculations, which nonetheless can be quite powerful. More sophisticated methods, which take account of transport times within the environment and breakdown rates of chemicals are also available. Such models are widely used in predicting the environmental fate of organic chemicals. QUESTIONS (i) Discuss what is meant by a biogeochemical cycle, describing the major facets and illustrating the processes involved with examples. (ii) Discuss the temporal trends in lead emissions and concentrations in the environment and how environmental media can be used to elucidate historical trends in environmental lead. (iii) Explain what is meant by an environmental lifetime and derive an expression for environmental lifetime in terms of a chemical rate constant. Compare and contrast the typical atmospheric lifetimes of methane, nitrogen oxide, and the CFCs and explain how this relates to the atmospheric distribution and properties of these compounds. (iv) Explain the processes by which trace substances can exchange between the atmosphere and the oceans and show how rates of exchange can be calculated. Give examples of substances whose exchange between these media is important. (v) Explain why the waters in rivers in different parts of the world have differing composition and relate this to the climatology of the region. Explain carefully what is meant by dissolved and suspended solids and explain how both arise. (vi) Explain the environmental pathways followed by lead emissions from road trafﬁc after emission to the atmosphere and explain how this can lead to pollution of a range of environmental media. Indicate the quantitative ways in which such transfer can be expressed. (vii) Estimate atmospheric lifetimes for the following: (i) methane, if the globally and diurnally averaged concentration of hydroxyl radical is 5 105 cm3 (ii) nitrogen dioxide in the middle of a summer day when the concentration of hydroxyl radical is 8 106 cm3 (iii) nitrogen dioxide at nighttime if the sole mechanism of removal is dry deposition with a deposition velocity of 0.1 cm s1, and the mixing depth is 100 m.

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345

(viii) If the atmospheric concentration of sulfur dioxide is 10 ppb, calculate the following: (i) the atmospheric concentration expressed in mg m3 at one atmospheric pressure and 251C. (Use the fact that 1 mol of any gas (64 g in the case of sulfur dioxide) occupies 22.4 L at STP) (ii) the deposition ﬂux to the surface if the deposition velocity is 1.0 cm s1 (iii) the atmospheric lifetime with respect to dry deposition for a mixing depth of 800 m (iv) the atmospheric lifetime with respect to oxidation by hydroxyl radical if the diurnally averaged OH radical concentration is 8 105 cm3 and the rate constant for the SO2OH reaction is 9 1013 cm3 molec1 s1 (v) the lifetime with respect to wet deposition if the washout coefﬁcient is 104 s1. (ix) Explain the thermodynamic controls on biological reduction processes in aquatic sediments and explain how these inﬂuence the chemical forms of nitrogen in the environment. (x) If any organic chemical has a log Kow of 4.3 and a river water concentration of 10 mg L1, what concentration is predicted to occur in ﬁsh if assumed to have a 50% organic matter (lipid) content (assume that the ﬁsh and the water are of equal density). REFERENCES 1. W.H. Schroeder and D.A. Lane, Environ. Sci. Technol., 1988, 22, 240. 2. R.M. Harrison and D.P.H. Laxen, Lead Pollution: Causes and Control, Chapman & Hall, London, 1981. 3. P. Brimblecombe, Air Composition and Chemistry, 2nd edn, Cambridge University Press, Cambridge, 1996. 4. J.-P. Candelone, S. Hong, C. Pellone and C.F. Boutron, J. Geophys. Res., 1995, 100, 16605–16616. 5. B.J. Finlayson-Pitts and J.N. Pitts Jr., Atmospheric Chemistry, Wiley, Chichester, 1986. 6. C.N. Hewitt and R.M. Harrison, Atmos. Environ., 1985, 19, 545. 7. J.A. Garland, Proc. R. Soc. London, Ser. A, 1977, 354, 245. 8. S. Yamulki, R.M. Harrison and K.W.T. Goulding, Atmos. Environ., 1996, 30, 109–118. 9. R.S. Cambray and J.D. Eakins, Nature, 1982, 300, 46. 10. P.A. Cawse, UKAEA Report No. AERE-9851, 1980.

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11. S.A. Slinn and W.G.N. Slinn, Atmos. Environ., 1980, 14, 1013. 12. R.M. Williams, Atmos. Environ., 1982, 16, 1933. 13. C.R. Ottley and R.M. Harrison, Eurotrac ASE Annual Report, Garmisch-Partenkirchen, 1990. 14. R.F. Critchley, Proceedings of International Conference on Heavy Metals in the Environment, Heidelberg, Germany; CEP Consultants, Edinburgh, 1983, 1109. 15. P.S. Liss and L. Merlivat, The Role of Air-Sea Exchange in Geochemical Cycling, P. Buat-Menard (ed), Reidel, Dordrecht, 1986, 113. 16. A.J. Watson, R.C. Upstill-Goddard and P.S. Liss, Nature, 1991, 349, 145. 17. R.M. Harrison and A.G. Allen, Atmos. Environ., 1991, 25A, 1719. 18. J.M. Martin and M. Meybeck, Mar. Chem., 1979, 7, 177–206. 19. R.M. Garrels and F.T. MacKenzie, Evolution of Sedimentary Rocks, W.W. Norton (ed), New York, 1971. 20. R.J. Gibbs, Science, 1970, 170, 1088. 21. R.M. Harrison and S.J. de Mora, Introductory Chemistry for the Environmental Sciences, 2nd edn, Cambridge University Press, Cambridge, 1996. 22. A.O. Davies and J.N. Galloway, Atmospheric Pollutants in Natural Waters, S.J. Eisenreich (ed), Ann Arbor, MI, 1981, 401. 23. P. O’Neill, Environment Chemistry, George, Allen, and Unwin, London, 1985. 24. Department of Environment, Transport and the Regions, ‘Digest of Environmental Statistics’, No. 19, The Stationary Ofﬁce Ltd., Edinburgh, 1997. 25. R.M. Harrison and M.B. Chirgawi, Sci. Total Environ., 1989, 83, 13. 26. R.M. Harrison and M.B. Chirgawi, Sci. Total Environ., 1989, 83, 47. 27. Royal Commission on Environmental Pollution, ‘Ninth Report: Lead in the Environment’, HMSO, London, 1983. 28. D. Mackay, E. Webster and T. Gouin, Partitioning, persistence and long-range transport of chemicals in the environment, in: Chemicals in the Environment: Assessing and Managing Risk, R.E. Hester and R.M. Harrison (eds), Issues in Environmental Science and Technology, No. 22, Royal Society of Chemistry, Cambridge, pp 132– 153, 2006.

Glossary abiotic: physical or chemical processes in the environment as opposed to biologically driven processes (biotic). adiabatic compression: compression of a ﬂuid without extraction of heat resulting in increased temperature. advection: horizontal or vertical transfer of material, heat, etc., due to the mass movement of water in the ocean. aeolian transport: wind-borne transfer of material. aggregation: process of combining particulate material into larger particles. algae: simple unicellular or multi-cellular photosynthetic plants, which thrive in a wet environment, e.g., lakes. alkalinity: a measure of the proton deficit in solution, which is operationally deﬁned by titration with a strong acid to the carbonic acid end point. alluvium(al): unconsolidated material, consisting of weathered or eroded particles of minerals and rocks, i.e., transported by a ﬂowing river and deposited at points along the ﬂood plain of the river. ammoniﬁcation: the bacterial transformation of dissolved organic nitrogen (DON) into dissolved inorganic nitrogen (DIN) with the ﬁrst product being NH3. amphoteric: ability of a substance to act as an acid (proton donor) or base (proton acceptor). androgenic: general name for any substance with male sex hormone activity in vertebrates. anoxic: waters devoid of oxygen because circulation is restricted vertically due to thermal or saline stratiﬁcation and horizontally by topographic boundaries thereby producing stagnant conditions. Examples of this atypical marine environment are the Black Sea and Saanich Inlet (Canada).

347

348

Glossary

apolar compound: a molecule that does not possess functional groups and undergoes molecular interactions that are dominated by Van-derWaals interactions only (e.g., n-alkanes). aragonite: a solid phase of CaCO3. authigenic: formed, e.g., minerals, in place during or after deposition of sediment. authigenic precipitation: solid-phase production in seawater due to a substance exceeding its solubility product, such as the formation of ferromanganese nodules. bacteria: cellular microorganisms incapable of photosynthesis. benthic: living on the bottom of a water body, e.g., lake. bioaccumulation: the progressive increase in concentration of a substance in an organism because the rate of intake via the body surface (e.g., from water) or in food is greater than the organism’s ability to remove the substance from the body. bioavailable: the portion of a chemical substance (nutrient, metal, toxin) that can be absorbed, transported, and utilized physiologically. biogenic hydrocarbon: an organic compound originating from a recent biological source, e.g., biogenic volatile organic compounds (VOCs) emitted from vegetation. biogenous material: matter produced by the ﬁxation of mineral phases by marine organisms; notable examples being calcite and opaline silica in the marine environment. biomagniﬁcation: a food chain or food web phenomenon, whereby a substance or element increases in concentration at successive trophic levels. biomethylation: biologically induced reactions by which methyl groups are added to metals, metalloids, and organometallic complexes, with consequent and varied inﬂuence on the volatility, bioavailability, and toxicity of the element concerned. bioremediation: the use of living organisms, generally bacteria, to clean up oil spills or remove other pollutants from soil, water, and wastewater. calcite: a solid phase of CaCO3. carbonate compensation depth (CCD): the depth at the sediment–water interface at which no calcareous material is preserved in the sediments. chelate: a complex formed when a central metal atom shares more than one electron pair with a given ligand thereby forming a ring structure, which exhibits enhanced stability largely due to the entropy effect of releasing large numbers of molecules from the co-ordinated water envelopes. chelation: the formation of a chelate.

Glossary

349

chlorinity (Cl%): the chloride concentration in seawater, expressed in g kg 1, as measured by Ag1 titration (i.e., ignoring other halide contributions by assuming Cl to be the only reactant). chromophore: a chemical group or moiety within a molecule that absorbs light. cloud condensation nuclei (CCN): small particles in the atmosphere onto which water can condense. conservative behaviour: the concentration of a constituent or absolute magnitude of a property varies only due to mixing processes. conglomerates: a group of detrital sedimentary rocks consisting of rounded or sub-rounded fragments. convergence: a region in which the streamlines (i.e., currents) come together causing water to sink. Coriolis force: the acceleration due to the earth’s rotation deﬂecting moving ﬂuids (i.e., both air and water) to the right in the northern hemisphere and to the left in the southern hemisphere; the magnitude of the effect is a function of latitude, being nil at the equator and increasing towards the poles. cyanobacteria: commonly known as blue–green algae, which are capable of reproducing rapidly to form algal blooms in nutrient-enriched freshwater and, after death, of releasing toxins. denitriﬁcation: the reduction of nitrogen species to molecular nitrogen (N2), which can occur under conditions of hypoxia or anoxia. detritus: disintegrated material of both inorganic and biological origins. diagenesis: the collection of processes that alter the sediments following deposition; they may be physical (compaction), chemical (cementation, mineral segregation, ion exchange reactions), or biological (respiration). diatoms: a class of planktonic one-celled algae with skeletons of silica. dielectric constant: the ratio of electric ﬂux density to electric ﬁeld, which expresses the degree of non-conductivity of different substances. divergence: a zone in which the ﬂow ﬁelds (i.e., currents) separate. endocrine disruptor: an exogenous substance or mixture that alters function(s) of the endocrine system and subsequently causes adverse health effects in an intact organism, or its progeny, or (sub-) populations. enrichment factor: the ratio in a sample compared to a reference substance (i.e., a possible source) of concentration ratios of a given element, X, to a reference element; typically Al is used as a reference element to denote a terrestrial source and Na for a marine source. eutrophication: over-enrichment of a water body with nutrients, resulting in excessive growth of organisms (i.e., algal blooms) and depletion of

350

Glossary

oxygen concentrations (hypoxia or anoxia depending upon the organic loading) in sub-surface waters and/or sediments. ﬂocculation: process of combining particulate material into larger particles facilitated by particle surfaces being coated with organic material. halmyrolysis: low-temperature reactions in seawater reactions producing secondary material from components of continental or volcanic origin; essentially an extension of chemical weathering of lithogenous components. halocline: a region in the water column that exhibits a sharp change in salinity. hermaphroditic: having both male and female sexual organs. humic substances: large, refractory molecules formed as a result of decomposition of organic matter. hydrogenous material: matter produced abiotically within the water column. hydrophobic: the low afﬁnity for water or ‘water-hating tendency’ of a chemical as measured by its octanol/water partition coefﬁcient. hydrothermal waters: formed when seawater circulates into a ﬁssured rock matrix and, under conditions of elevated temperature and pressure, results in compositional changes in the aqueous phase occur due to seawater–rock interactions. hyperkeratosis: skin lesion condition, often in the form of nodules on the palms of the hands and soles of the feet, which can be brought about by exposure to inorganic arsenic in drinking water. hyperpigmentation: in the case of chronic arsenic poisoning, a ﬁnely freckled, ‘raindrop’ pattern of pigmentation that is particularly pronounced on the trunk and extremities. hypolimnion: the stratum of water below the thermocline. hypoxia: waters that are deﬁcient in oxygen. imposex: the imposition of male sex organs on female marine gastropods, notably resulting from exposure to tributyltin compounds. ion pair: the transient coupling of a cation and anion, formed by the collision of oppositely charged ions due to electrostatic attraction, during which each ion retains its own co-ordinated water envelope. isopycnal: line (imaginary or on a chart) or an imaginary surface connecting points, which have the same density. latent heat of evaporation: the heat required (or released) to convert a unit mass of a substance from a liquid to a gas state (or from gas to liquid) at the same temperature and pressure. latent heat of fusion: the heat required (or released) to convert a unit mass of a substance from a solid to the liquid state (or from liquid to solid) at the same temperature and pressure.

Glossary

351

ligand: a moiety acting as a Lewis base that shares a pair of electrons with a metal atom thereby forming a complex; ligands can be neutral (e.g., H2O) or anionic (e.g., Cl , HCO3 ) species. lithogenous material: substances coming from the continents as a result of weathering processes; the most important components being quartz and the clay minerals (kaolinite, illite, montmorillonite, and chlorite). lysocline: the depth at the sediment–water interface at which appreciable dissolution of CaCO3(s) starts to occur. marine boundary layer: lowest portion of the atmosphere (from 0.25 to 3.5 km) in contact with the ocean surface. metallothionein: A protective protein that binds heavy metals such as cadmium and lead thereby detoxifying such contaminants. molal concentration: the number of moles of solute per kg of solution. molar concentration: the number of moles of solute in a litre (dm3) of solution (i.e., molarity). molecular viscosity: small-scale internal ﬂuid friction that is due to the random motion of the molecules within a smooth-ﬂowing ﬂuid. mutagen: a chemical substance that induces a change in the of DNA of cells. nitriﬁcation: the stepwise oxidation of NH41 to NO2 and eventually to NO3 . nitrogen ﬁxation: the direct uptake and assimilation of molecular nitrogen (N2) by organisms. ocean conveyor belt: see thermohaline circulation. ochre: a natural hydrated form of ferric oxide, often precipitated from acid mine drainage. oestrogenic: related to, behaving like, or causing effects similar to the natural female sex hormone in vertebrates. oligotrophic: a term used of lakes that are poorly productive in terms of organic matter formed because of a low nutrient supply, i.e., the opposite of eutrophic. osmoregulatory: a term used of an organism that actively regulates the osmotic concentration of its internal ﬂuids. oxygen compensation depth: the horizon in the water column at which the rate of O2 production by photosynthesis equals the rate of respiratory O2 oxidation. permeability: the capacity, e.g., of a rock, to transmit water. petrogenic hydrocarbon: an organic compound originating from unburnt fossil fuels. photic zone: the upper surface of the ocean in which photosynthesis can occur, typically taken to be the depth at which sunlight radiation has declined to 1% of the magnitude at the surface.

352

Glossary

photosensitiser: a substance that absorbs light and subsequently initiates a photochemical reaction or photophysical alteration in the chemical of interest. photosynthetically active radiation (PAR): electromagnetic energy in the 400–700 nm wavelength range that can be absorbed by chlorophyll or other light-harvesting pigments in plants and is used for photosynthesis. physico-chemical speciation: the various physical and chemical forms in which an element may exist in a speciﬁed system. phytoplankton: microscopic plants living in the oceanic water column. phytoremediation: the use of living plants to reduce the risk posed by contaminated soil, sludges, sediments, and groundwater by removing, degrading, immobilizing, or containing contaminants in situ. planetary albedo: the reﬂectivity of the world with respect to incident sunlight: calm seawater (B2%), vegetated regions (10–25%), deserts (B35%), and snow covered surfaces (B90%). podsolization: the leaching of soluble complexes of aluminum and ferrous iron from the A horizon and the subsequent deposition of these metals, together with organic matter, in the B horizon. potential density, rh: the density deﬁned on the basis of potential temperature instead of in situ temperature. potential temperature: the temperature that a water parcel would have if raised adiabatically to the ocean surface. pycnocline: a region in the water column that exhibits a sharp change in density. pyrolytic hydrocarbon: an organic compound originating from combustion sources. quantum yield: number of molecules of reactant consumed per photon of light absorbed. radiolaria: single-celled planktonic organisms that build a skeleton of silica. Redﬁeld ratio: the constant relative amount of nutrients taken up by phytoplankton during photosynthesis, being 106C:16N:1P. refractometer: instrument for measuring the indices of refraction of various substances. remineralisation: bacterial degradation of particulate organic nitrogen (POC) into dissolved organic nitrogen (DON). residence time: the average lifetime of the component in a speciﬁed system, which is, in effect, a reciprocal rate constant. respiration: biochemical and cellular processes within an organism by which carbon is combined with absorbed oxygen to produce carbon dioxide. Revelle factor (R): a resistance to the dissolution of CaCO3(s) by carbonic acid indicating that the ocean is relatively well buﬀered against changes in SCO2 in response to variations in atmospheric pCO2.

Glossary

353

salinity (S%): a measure of the salt content of seawater; a typical value for oceanic waters being 35 g kg 1. scavenge: the process by which sinking particles remove dissolved substances (metals and organic material) by surface adsorption. siderophore: a low molecular weight substance synthesized by a variety of micro-organisms to sequester Fe that is often in limited supply in the environment. sigma-tee (rt): the density (actually the speciﬁc gravity and hence it is a dimensionless number) of water at atmospheric pressure based on temperature and salinity in situ. sorption: either or both adsorption or absorption of a chemical substance to a surface. speciﬁc heat: the amount of heat required to raise the temperature of 1 g of a substance by 1 1C. super-saturation: the unstable state of a solution or vapour that contains more solute or gas than its solubility or vapour pressure allows. surface tension: the work required to expand the surface of a liquid by unit area; the cohesive force exerted at the surface of a liquid that makes it tend to assume a spherical shape because molecules at the surface are not surrounded by molecules on all sides and accordingly they interact more strongly with those directly adjacent to them on and below the surface. surfactants: surface-active or wetting agents that decrease the surface tension of a liquid. thermocline: that zone between the warm surface water and colder deeper water, e.g., of a lake, in which the temperature changes most rapidly. thermohaline circulation: the global deepwater circulation driven by density differences, which in turn are caused by variations in salinity or temperature. tidal prism: the difference in the volume of water in a water body between low and high tides. troposphere: the lowest region of the atmosphere, extending to an altitude of 8–18 km. T–S diagram: a plot of temperature against salinity, with applications in deﬁning water masses of different origin. turbidity currents: the muddy and turbulent ﬂow of sediment-laden waters along the seaﬂoor caused by sediment slumping that transport copious amounts of material, including coarse-grained sediments, to the deep sea. upwelling: the transport of sub-surface water to the ocean surface. zooplankton: animal component of the planktonic community.

Subject Index A Glossary of terms can be found on pages 347–353.

acid dissociation constant (pKa) 305 acid hydrolysis 248 acid mine drainage 126–9 acid rain 255, 335 acidiﬁcation 122–9 acidity constants 90, 97 activity coeﬃcient 83, 265–6, 267 adiabatic compression 176 adsorption 214–15 advection 145, 179, 224 aeolian transport 189, 202 aerodynamic resistance 322–3 aerosol deposition 321 formation 199, 324 organic 280 aggregation 129, 221 air accumulation factor (AAF) 338 air-land exchange 321–3 air-sea exchange 172, 184–7, 191, 324–30 air-water partition coeﬃcient 286, 288–91, 298 airborne particles 324, 338 albedo 197, 199 algae 129, 145–50, 212, 283 blooms 145, 147–50 macroalga 53, 55 aliphatic hydrocarbons 143, 281 alkalinity 103–5, 193–5 alkanes 225, 280–1, 283 alkenes 46 alkoxy radicals 280 alkyllead additives 139, 336 alkylphenolethoxylates (APEs) 154 alkylphenols 154

alluvium 137, 236, 238 aluminium 108–10, 123–5 aluminium oxide 246 aluminosilicate clay minerals 123, 234, 241–6 ammonia 323, 335 ammoniﬁcation 335 amorphous aluminosilicates 245–6 amphiboles 239–40 amphoteric behaviour 174, 214 androgenic eﬀects 154 anion exchange 267 anoxic 115, 189, 218–19 Antarctic bromine explosion 52–4 currents 179–81 glacier-derived material 172 ice 317 ozone hole 63–71 seabourne litter 227 vortex 66, 68–71 APEs 154 apolar compound 303 aquatic systems, modelling 121–2 aquifer 128–31, 156 aragonite 190, 212 Arctic bromine explosion 52–4 ice 181, 317, 337 ozone hole 65, 71 water concentrations, lindane 293–5 arsenate 130, 272 arsenic 117–18, 130–1 arsenopyrite 130

354

355

Subject Index atmosphere chemistry 5 conductance 323 cycle 315, 320 oxidising capacity 20 trace gases 10–11, 12–14 vertical structure 8–10, 19, 316–17 atrazine 152, 153 authigenic precipitation 142–3, 146–7, 211–12 bacteria 257, 258 base saturation 270–1 bauxite 125 BCF 287–8 Beer-Lambert law 302–3 benthic 128, 145, 222 benzene 157, 327 benzophenone-3 155 bioaccumulation 138, 284, 295 bioavailability metal complexes 82, 125 phosphorus 146 biochemical (or biological) oxygen demand (BOD) 150–1 bioconcentration factor (BCF) 287–8 biogenic hydrocarbon sources 10, 35–6, 46, 61, 183, 197–8, 283 biogenous 212, 213 biogeochemical cycles 170–3, 197–8, 200, 209, 216, 314–16, 333–40 biological reduction processes, sediments 334 biomagniﬁcation 138, 153 biomethylation 131, 172 bioremediation 224–5 bisphenol-A 154 blood lead levels 132 PBDEs 307 BOD 150–1 borate 272 boundary layer resistance 323 breast milk, PDBE levels 307 brominated ﬂame retardants (BFRs) 305–9 bromine explosion 52–4 Brønsted–Guggenheim equation 87 buﬀer capacity, soil 255

buﬀer strips, water systems 145, 148 bulk density, soil 237 c spacing 243 cadmium 132–3, 202 caesium 202, 324 calcite 190, 212 calcium 143–4, 148, 206 calcium carbonate deposition 173, 190–1, 196–7 carbamazepine 155 carbon dioxide 100–3, 189–97 Carbon Preference Index (CPI) 283 carbonate alkalinity (CA) 104–5 carbonate compensation depth (CCD) 196, 201 carbonate equilibria 100–3, 192 care products 151, 155–6, 305 catalytic cycles 53, 56–63, 69 cation exchange capacity (CEC) 216, 242–3, 250, 269–71 CCD 196, 201 CCN 197, 199 CD-MUSIC 122 CFCs 184, 317 Chapman cycle 58–9 charge balance 98–100, 103–5, 248 chelate 206 chelation 205, 208, 248, 273 chemical oxygen demand (COD) 150–1 chemical processes 1–2, 3 chemical speciation, ocean 204–10 chemical weathering 91–3 chemisorption 215, 271–3 chlor-alkali plant 144 chloramines 132 chlorinated solvents 82, 300 chlorine 59–61, 64, 67–71 chlorinity (Cl%) 178 chlorite 211, 244 chlorpyrifos 303–4 chlortoluon 152 cholesterol 151 chromate 272 chromite ore processing residue (COPR) 135 chromium 106–7, 116–17, 135–7 chromophore 302 Cl% 178

356 Clark Fork river basin 127–8 clay 128, 133 minerals 238–45, 259, 269 climate 8, 72, 181, 185, 197, 199, 231, 235, 314 closed aqueous system 101–3 cloud condensation nuclei (CCN) 197, 199 COD 150–1 colloids 204, 208–9 complexation 273 complexes inorganic 96 organic 97 organometallic 273 outer and inner sphere 94–7 pH-dependent formation 105–10 soil solution 266 surface 96–7 concentration, units of 5–6 concentration factor (CF) 338 conglomerates 134 congruent dissolution 92 conservative behaviour 220 consolidated rock 238 constant ionic strength equilibrium constant (cK) 89 convergence 181 copper 127, 202 coprostanol 151 Coriolis force 179 cosmogenous 212–13 Coulomb’s law 268 Criegee intermediates 46, 49 cumulative stability constants (bn) 95 cyanide 138–9 cyanobacteria 145, 148 2,4-D 152 Davies equation 87–9, 265 DDT 289, 295 Debye-Hu¨ckel (limiting law) equation 85–7, 265 degradation 299–301, 315 denitriﬁcation 335 density 176–7 deposition process 34–5, 210, 314–15, 321–3 velocity 34, 321–2, 324 DET 142

Subject Index detergent metabolites 151 detritus 139, 200 DGT 142–3 diagenesis 210, 218–19 dialysis 125 diatoms 122, 199, 212 dibromoethane 336 dicamba 152 dichloroethane 336 dichloromethane (DCM) 156 dichlorprop 152 dichromate 150–1 dielectric constant 268 N,N-diethyltoluamide 151 diﬀuse double layer 268 dimethoate 152 dimethylarsinic acid 131 dimethylsulﬁde (DMS) 41–2, 44–5, 183 climatic implications 197–9 dinitrogen pentoxide 40 disinfection, water 132 dissolution 91–4, 196, 247–8 dissolved inorganic nitrogen (DIN) 200 organic carbon (DOC) 305 organic nitrogen (DON) 200 divergence 180 DMS 41–2, 44–5, 183, 197–9 drinking water 80–1, 125, 128, 131–2 lead 338–9 nitrate 149 pesticides 151–2 dry deposition 26, 34, 50, 314–15, 322, 326 earth’s crust 234–75 eddy diﬀusivity, Kz 323 EF 182–3, 234, 324 electrode potential (EH) 111, 188, 260–1 electron activity 302 electron activity, hypothetical 110, 112, 188, 261 emission inventories 13 emulsions 224 endocrine disruptor 154–5 enrichment factor (EF) seawater 182–3, 324 soil 234 environmental behaviour 320

357

Subject Index chemicals 143–5, 225, 279–310 compartments 314, 316–17, 321, 340 environmental organic chemistry (EOC) 5, 279–310 Environmental Quality Standards (EQSs) 136, 153 environmental reservoirs 170–3 equilibrium constant 89–91, 192 equilibrium criteria model (EQM) 296 estuaries 212, 216, 219–22 ethinylestradiol 154 ethylbenzene 157 ethylhexyl methoxy cinnamate 155 eutrophication 145–9 exchangeable acidity 254 exchangeable cations 240, 257, 270–1 extended DH equations 85–8 feldspars 91, 239–40 ferric sulfate 148 ferrihydrite 246 ‘ferromagnesian’ minerals 240 ferromanganese nodules 211, 217–18 ferrous sulfate 136, 253 ﬂocculation 127 ﬂuoride 272 ﬂux 321–4 foodstuﬀs, lead 338 formaldehyde 11, 24, 39 freshwater 3–4, 80–157 species distribution 97–121 frog development 153 fulvic acid 252, 282 Gapon equation 269 gastropods, marine 228–9 geochemical computer models 87, 121–2 geogenic sources 11 geothite 246 Gibbs free energy of transfer, DG 111, 292–3 Gibbs-Helmholz equation 292–3 gleys 263 global pollution problems 230, 317 gold 137–9 groundwater 80–2, 330 arsenic 130–1 chromium 135–7

nitrate 149–50 selenium 133–5 Gu¨ntelberg approximation 86–7 Guoy-Chapman model 268 haematite 246 halmyrolysis 212, 216 halocarbons 2, 18, 36 halocline 176 halogen chemistry 51–8 oxidation 52–3 heavy metals 127–9, 142 Henry’s Law 186, 341 constant (H) 186, 285, 288–91, 326–9, 341 hermaphroditic ﬁsh 154 hexavalent chromium 136 human lead exposure 338–9 humic acid 252, 282 humic substances 252–3, 267, 282 humin 252 humus 249 formation 249–51 fractionation 252 hydrocalumite 136 hydrocarbons 35–7 oxidation 35–7 sources 282–4 total petroleum 225 volatile (oil) 225 hydrogen bonding 174, 215, 244 hydrogenolysis 300 hydrogenous material 211–12, 216 hydrophobic 154 hydrophobic organic chemicals (HOCs) 290 hydrothermal vents 171, 204, 217 hydrous iron oxides 130 hydrous metal oxides 135 hydrous oxides 20–4, 246, 267 hydroxyapatite 148 hydroxyl radical 20–4, 297, 318–19 hyperkeratosis 130 hyperpigmentation 130 hypolimnion 146 hypoxia 189

358 ibuprofen 155 ice Antarctic 317 Arctic 181, 317, 337 rafting 213 illite 211, 243 imposex 228–9 incongruent dissolution 92 inherited minerals 238 inner sphere complexation 95–6 internal loading, phosphorus 147 iodine monoxide 54–5 ion(s) activity 83–8, 265–6 exchange 215–16, 267–71 ionic potential 183–4, 263 ionic strength 84–9, 110–21, 265–6 pair formation 266–7 Irish Sea 324, 325 iron 112–13, 119–21 biological productivity limiting 200 estuarine 221–2 iron oxides 115, 126–7, 272 hydrous oxides 130 irrigation 132–5, 150 isomorphous substitution 242 isopleth 32 isoproturon 152 isopycnals 177 Itai-Itai 133 kaolinite 211, 243 Kaw 286, 288–91, 298 Kd 286, 288 Kesterson Reservoir 135 Koa 286, 291–2, 298, 342–3 Kow 286, 287–8, 298, 342 lake 82, 123–5 sediment 139–49 Langmuir equation 272–3 latent heat 174, 323 leaching 257 lead 131–2, 139, 202, 317, 333 biogeochemical cycle 335–40 human exposure 338–9 lead-210 139 leaded petrol 317, 336, 339–40 lead halides 336

Subject Index Leblanc process 143 lifetime(s) 25, 317–21 behavioural inﬂuence 319–21 calculation 317–19 ligand 206–7 ligand exchange 271–3 liming 132, 149, 255 lindane 152, 293–5 linuron 152 lithogenous material 210–11 long range transport 295 lysocline 196 MacInnes convention 87 magnesium 206 manganese 112, 204 manganese oxides 246 marine boundary layer 21, 34, 171–2, 179, 197 marine sediment 210–13 MCPA 152 mecoprop 152 mercury 137–8, 144–5, 207–8 mesopause 9–10 mesosphere 9–10 metallothionein 205, 208 methaemoglobinaemia 149–50 methane 315–16, 317, 318–19 methyl chloroform 319 methyl mercury 138 methyl tert-butyl ether (MTBE) 156–7 methylation 209 4-methylbenzylidene camphor 155 micas 239–40 mineral components, soil 238–46 Minimata, Japan 138 MINTEQA2 87, 121 mixed equilibrium constant 89–91 mixing times 316 MODEL V 122 ‘model world’ environment 340–1 molal concentration 285 molecular viscosity 174 molybdenum 142, 145 monomethylarsonic acid 131 montmorillonite 211, 216 Montreal protocol 2, 63–4

Subject Index MTBE 156–7 mutagen 300 natural organic matter 94, 97 Nernst Equation 111 NICA 122 nickel 202 nitrate 40–3, 149–50, 199, 272 nitrate radical 40–5 nitric acid 27–30, 40–4, 67–9, 322 nitriﬁcation 257, 335 nitrite 150 nitroarenes 299–300 nitrogen biogeochemical cycling 200–1, 335, 337 eutrophication 81, 149 ﬁxation 200 oxidation, atmospheric 26–8 in seawater 185–6, 200 nitrogen dioxide 6, 23, 318–19, 320 nitrogen oxides 11, 15–16, 26–8, 28–33, 335, 336 N-nitroso compounds 150 nitrous oxide 315 non-conservative behaviour 220 non-exchangeable acidity 254 non-methane hydrocarbons (NMHC) 30–1, 35–9 non-sea salt sulfate (NSSS) 183 4-nonylphenol 151 nonylphenol (NP) 154 North Sea 324–6 NOx see nitrogen oxides NOy 28 O2 minimum 188–9, 196, 200 ocean(s) 170–230 circulation 179–82, 213 conveyor belt 179 see also seawater ochre 148 octanol-air partition coeﬃcient 286, 291–2, 298, 342–3 octanol-water partition coeﬃcient 286, 287–8, 298, 342 oestradiol 154 17b-oestradiol 154 oestrogenic steroids 154 oestrone 154

359 oil slicks 223–6 oligotrophic 140, 145 olivine 239–40 ORCHESTRA 122 organic chemicals 151, 280–4, 340–4 ‘model world’ environment 340–1 organic chemistry 5, 279–310 organic matter 113, 141, 150–7 natural 94, 97 soil 235, 248–53 turnover of 249–50 see also humus organochlorine pesticides 143, 152–3, 154 organophosphorus pesticides 153, 303–4 orthophosphate 146 osmoregulation 124 outer and inner sphere complexes 94–7 outer sphere complexation 94–5 oxic-anoxic boundary 115, 189, 218–19 oxidation atmospheric nitrogen 26–8 atmospheric sulfur 48–51 halogens 52–3 hydrocarbons 35–7 night-time 40–5 soil weathering 248 tropospheric 17, 20–5 oxidising capacity, atmosphere 20 oxygen 187–9 compensation depth 187 minimum 188–9, 196, 200 ozone alkene interactions 46, 47 halogen interactions 52–3, 56–8 layer 14, 17–18, 58–63 production and destruction 28–40 ozone hole Antarctic 63–71 Arctic 65, 71 PAHs 143, 225, 228, 282 PAN 28, 39–40 PAR 200 parent material, soil 235–41, 246, 270, 275, 337 particles airborne 324, 338 estuarine 220–2

360 marine 210–13 surface chemistry 213–18 partitioning 284–6, 340–3 coeﬃcients 285, 286–92, 293–5, 341 partition maps 295–9 temperature dependence 292–5 PCBs 143, 153–4, 295, 316, 320 PDBEs 143, 153–4 pe 110–21, 171, 188–9, 261 pe-pH diagrams 114–21, 209 perchloroethylene (PCE) 156 perﬂuorooctane sulfonates (PFOSs) 154 permanent negative charge, clay 242 permeability layers 82 permeable reactive barrier 136–8 peroxides 23, 34–5, 46 peroxy radicals 15–16, 19–23, 30, 280 peroxyacetlynitrate (PAN) 28, 39–40 persistence 295–6 persistent bioaccumlative toxins (PBTs) 284 persistent organic pollutant (POP) 295, 320 personal care products 151, 155–6, 305 pesticides 151–3 petrogenic hydrocarbon 283–4 PFOSs 154 pH 97–110, 171, 190, 194 minimum, estuary 220 pe-pH diagrams 114–21 scales 190 soil 254–9 pH-dependent charge 246, 247, 258, 267, 269 complex formation 105–10 dissolution 142 pharmaceuticals 151, 155–6 phosphate 199, 201, 203, 264, 272 phosphorus 145–9 internal loading 147 photic zone 173, 187 photochemical transformation 209, 299–301, 302–3, 335 photodissociation 11–17 photolysis 303–5 photosensitiser 305 photostationary state (PSS) 26–8 photosynthesis 113, 147, 173, 187, 195, 303, 333

Subject Index photosynthetically active radiation (PAR) 200 PHREEQC 121 PHREEQE 87 phthalates 154 physical weathering 91 physico-chemical speciation estuary 221 sea 204–10 physisorption 267–71 phytoplankton 221 phytoremediation 135 pit lakes 129 Pitzer model 88–9 planetary albedo 197, 199 plastic debris 226–8 plasticizers 151 plutonium 324–5 podsolization 123 point of zero charge (PZC) 214 pollution 222, 223–30 cycling 315 polybrominated diphenyl ethers (PDBEs) 143, 153–4, 306–9 polychlorinated biphenyls (PCBs) 143, 153–4, 295, 316, 320 polychlorinated dibenzo-p-dioxins 143 polychlorinated dibenzofurans 143 polychlorinated naphthalenes (PCNs) 143, 291 polycyclic aromatic hydrocarbons (PAHs) 143, 225, 228, 282 pore water 4, 130, 141–2, 146, 228, 341 potential density (sy) 177 potential temperature 176 PPCPs 155–6 precipitation/dissolution 91–4, 273–4 precipitation reactions 91–4 primary soil minerals 238–40 primary sources 10, 21, 61–2, 128 pycnocline 177 pyrethroids, synthetic 152–3 pyrite 126–7, 130, 219 pyrolytic hydrocarbon 283–4 pyroxenes 239–40 quantum yield 16 quartz 239–40

Subject Index radioisotopes 324–5, 333 radiolaria 199, 212 rain acidity 255, 335 dissolved constituents 331 Redﬁeld ratio 199 redox equilibria 114, 116, 118 redox intensity 110–14 redox potential 110–21, 171, 188–9, 260–1 redox reactions 110–14 seawater 204–5, 209 soil 260–1 redox-sensitive elements 114 refractometer 177 remineralisation 200 reservoirs, environmental 170–3, 316–21 residence time 173, 184, 202 resistance aerodynamic 322–3 boundary layer 323 surface 323 total 323 respiration 187 Revelle factor (R) 196 Richmond mine 126 riverine particulate material 283, 330–1 river(s) 332 dissolved constituents 331 water 219, 330–2 rocks 5 sy 177, 179 st 176–7 S% 174–9 salinity 174–9 salinization 81, 133 salt waters 4–5 saturation vapour pressure 289, 326–9 scavenging coeﬃcient 329–30 ratio 329 sea borne litter 227 sea(s) see ocean(s); seawater seawater 173–9 albedo 197 constituents 182–4 dissolved gases 184–99 dissolved solids 332 nutrients 199–201

361 physico-chemical speciation 204–10 Standard Seawater 178 suspended particle matter (SPM) 283 trace elements 201–4 secondary soil minerals 240–6 secondary sources 10, 21, 39 sediment estuarine 220–2 freshwater 4, 332–3 diatom records 122 metal and metalloid contamination 128, 130, 134–5, 138–45 nutrients 94, 145–9 oxic-anoxic boundary 115 perturbatory processes 139–45 pesticides 153 marine 210–13 diagenesis 218–19 silica 201 sources 171–3, 190 trace elements 201–4 water-sediment boundary 216–17 selenate 135, 272 selenide 135 selenite 135 selenium 133–5 Sellaﬁeld 324 Setschenow equation 88–9 settling 139, 145 sewage sludge 133, 300, 307 short-range order (amorphous) aluminosilicates 245–6 siderophores 205, 208 sigma-tee (st) 176–7 silicate(s) 199, 201, 203, 272 clays 267 primary soil minerals 238–40 silicic acid 93 silver 127 simazine 152 single-ion activity coeﬃcients 83–4, 86 sinks (biogeochemical outputs) 170–3, 317–18 SIT model 88 skin cancer 130 smectites 244 smog(s) 37, 39–40 sodium 144, 205–6 soil 234–46

362 acidity 255–7 alkalinity 257–8 forming factors 235 mineral components 238–46 organic carbon content 253 organic matter 235, 248–53, 282 parent material 235–41, 246, 270, 275, 337 pH 254–9 processes 5, 275, 276 redox potential 260–1 reduction processes 261–3 solution 256, 259, 263–7 water and air content 237 weathering sequence 236 see also humus solubility 93–4, 184–7, 274, 289 equilibrium solubility 327–8 solubility product 84, 93, 108, 133, 211, 259 Solvay Process 144 solvents, industrial 156 sorption 130, 145 sources (biogeochemical inputs) 170–3, 317–18 speciﬁc gravity 176–7 speciﬁc heat 174 Speciﬁc Ion Interaction Theory (SIT) 87–9 stability constant 95, 205–8 Stern model 268–9 steroids 151, 154 stratosphere 8–10, 58–71 source gases 62 sulfate 206 reduction 218–19 sulﬁde 126–7, 129, 130, 218–19 oxidation 257 sulfur chemistry 46–51, 118–19 geochemical cycle 183 oxidation, atmospheric 48–51 sulfur dioxide 10, 48–51, 317 sulfuric acid 48, 67, 126–7 ‘Superfund’ sites 127 supersaturation 185 surface charge 213–14 surface resistance 323 surface tension 174–5

Subject Index surfactants 175, 301 surﬁcial rocks 330–1 suspended particle matter (SPM) 283, 330–1 synthetic pyrethroids 152 2,4,5-T 152 T-S diagram 177 TBT 228–30 TDS 83 tetrabromobisphenol A (TBBPA) 306–9 tetrachloroethene 156 tetrachloroethylene 156 tetrachloromethane 156 thermocline 82, 176 thermohaline circulation 181 tidal prism 219 toluene 157 total dissolved solids (TDS) 83 transfer in aquatic systems 330–3 gas (water-air) 328–9 velocity 186, 328–9 transport, long range 295 tributyltin (TBT) 228–30 1,1,1-trichloroethane (TCA) 156 trichloroethene (TCE) 156 trichloroethylene 156 tri(2-chloroethyl)phosphate 151 trichloromethane 156 triclosan 305 troposphere 8–10, 51 chemistry 17–20, 21 halogens 56–8 marine 191, 198–9 oxidation chemistry 20–40 Truesdell-Jones equation 87 tube wells 130 turbidity currents 213 ultraviolet (UV) light 15, 20, 58–60, 302 unconsolidated superﬁcial deposits 238 unit layers, clay 242, 243 upwelling 180, 200–1 urban chemistry 37–40 van’t Hoﬀ equation 293 vermiculite 243–4 vertical mixing, atmosphere 316

363

Subject Index vinyl chloride 156 vitellogenin 154 vivianite 146 volatile organic compounds (VOC) 31–2, 37–40, 280, 302 volcanoes 211–12 volume-mixing ratio 6 vortex 66, 68–71 waste water 151–2, 155–6, 301 WATEQ4F 121–2 water inventory, Earth’s surface 81 properties 173–7 river 219, 330–2

water meadow 145, 148 water surfaces 326 weathering 210 physical 91 ratio, oil 225–6, 283 soil sequence 236, 246–8 wet deposition 314–15 wetlands 129–30 Wheal Jane tin mine 127, 129 xylene 157 zinc 204, 324–5, 327 zooplankton 199

SUBJECT INDEX A Glossary of terms can be found on pages 347–353.

Index Terms

Links

A acid dissociation constant (pKa)

305

acid hydrolysis

248

acid mine drainage

126

acid rain

255

acidification

122

335

acidity constants

activity coefficient

265

267

224

adiabatic compression

176

adsorption

214

advection

145

179

aeolian transport

189

202

aerodynamic resistance

322

aerosol deposition

321

formation

199

organic

280

aggregation

129

air accumulation factor (AAF)

338

air-land exchange

321

air-sea exchange

172

184

191

air-water partition coefficient

286

288

298

airborne particles

324

338

albedo

197

199

324

221

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324

Index Terms

Links

algae

129

145

147

aliphatic hydrocarbons

143

281

alkalinity

103

193

alkanes

225

280

alkenes

blooms macroalga

alkoxy radicals

280

alkyllead additives

139

alkylphenolethoxylates (APEs)

154

alkylphenols

154

alluvium

137

236

aluminium

108

123

aluminium oxide

246

aluminosilicate clay minerals

123

234

ammonia

323

335

ammonification

335

amorphous aluminosilicates

245

amphiboles

239

amphoteric behaviour

174

androgenic effects

154

anion exchange

267

anoxic

115

212

283

336

238

241

214

189

218

Antarctic bromine explosion

currents

179

glacier-derived material

172

ice

317

ozone hole

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283

Index Terms

Links

Antarctic (Cont.) seabourne litter vortex

227 66

APEs

154

apolar compound

303

aquatic systems, modelling

121

aquifer

128

156

aragonite

190

212

Arctic bromine explosion ice ozone hole water concentrations, lindane

52 181

317

337

293

arsenate

130

272

arsenic

117

130

arsenopyrite

130

atmosphere chemistry

conductance

323

cycle

315

320

oxidising capacity

trace gases

atrazine

152

153

authigenic precipitation

142

146

vertical structure

316

211

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Index Terms

Links

B bacteria

257

258

base saturation

270

bauxite

125

BCF

287

Beer-Lambert law

302

benthic

128

145

benzene

157

327

benzophenone-3

155

bioaccumulation

138

284

125

222

295

bioavailability metal complexes phosphorus

146

biochemical (or biological) oxygen demand (BOD)

150

bioconcentration factor (BCF)

287

biogenic hydrocarbon sources

183

197

283

biogenous

212

213

biogeochemical cycles

170

197

200

216

314

333

biological reduction processes, sediments

334

biomagnification

138

153

biomethylation

131

172

bioremediation

224

bisphenol-A

154

blood lead levels

132

PBDEs

307

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209

Index Terms

Links

BOD

150

borate

272

boundary layer resistance

323

breast milk, PDBE levels

307

brominated flame retardants (BFRs)

305

bromine explosion

Brønsted–Guggenheim equation

buffer capacity, soil

255

buffer strips, water systems

145

bulk density, soil

237

148

C c spacing

243

cadmium

132

202

caesium

202

324

calcite

190

212

calcium

143

148

206

calcium carbonate deposition

173

190

196

carbamazepine

155

carbon dioxide

100

Carbon Preference Index (CPI)

283

carbonate alkalinity (CA)

104

carbonate compensation depth (CCD)

196

201

carbonate equilibria

100

192

care products

151

155

305

cation exchange capacity (CEC)

216

242

250

CCD

196

201

CCN

197

199

CD-MUSIC

122

catalytic cycles

189

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269

Index Terms

Links

CFCs

184

Chapman cycle

charge balance

chelate

206

chelation

205

chemical oxygen demand (COD)

150

chemical processes chemical speciation, ocean chemical weathering

317

103

248

208

248

273

204 91

chemisorption

215

chlor-alkali plant

144

chloramines

132

271

chlorinated solvents

300

chlorine

chlorinity (Cl‰)

178

chlorite

211

chlorpyrifos

303

chlortoluon

152

cholesterol

151

chromate

272

chromite ore processing residue (COPR)

135

chromium

106

chromophore

302

Cl‰

178

Clark Fork river basin

127

clay

128

133

238

259

269

181

185

197

199

231

235

minerals climate

244

116

135

314 This page has been reformatted by Knovel to provide easier navigation.

Index Terms

Links

closed aqueous system

101

cloud condensation nuclei (CCN)

197

COD

150

colloids

204

complexation

273

199

208

complexes inorganic

organic

organometallic outer and inner sphere

273 94

pH-dependent formation

105

soil solution

266

surface concentration, units of

96 5

concentration factor (CF)

338

conglomerates

134

congruent dissolution

conservative behaviour

220

consolidated rock

238

constant ionic strength equilibrium constant (cK)

convergence

181

copper

127

coprostanol

151

Coriolis force

179

cosmogenous

212

Coulomb’s law

268

Criegee intermediates

cumulative stability constants (ȕn)

cyanide

202

138

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Index Terms

Links

cyanobacteria

145

148

D 2,4-D Davies equation

152 87

265

289

295

265

degradation

299

315

denitrification

335

density

176

DDT Debye-Hückel (limiting law) equation

deposition process

210

314

velocity

321

324

DET

142

detergent metabolites

151

detritus

139

DGT

142

diagenesis

210

dialysis

125

diatoms

122

dibromoethane

336

dicamba

152

dichloroethane

336

dichloromethane (DCM)

156

dichlorprop

152

dichromate

150

dielectric constant

268

N,N-diethyltoluamide

151

diffuse double layer

268

dimethoate

152

200

218

199

212

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321

Index Terms

Links

dimethylarsinic acid

131

dimethylsulfide (DMS) climatic implications dinitrogen pentoxide disinfection, water dissolution

183

196

247

197 40 132 91

dissolved inorganic nitrogen (DIN)

200

organic carbon (DOC)

305

organic nitrogen (DON)

200

divergence

180

DMS

183

197

drinking water

125

128

131

314

322

326

lead

338

nitrate

149

pesticides

151

dry deposition

E earth’s crust

234

eddy diffusivity, Kz

323

182

234

324

electrode potential (EH)

111

188

260

electron activity

302

electron activity, hypothetical

110

112

188

emission inventories

emulsions

224

endocrine disruptor

154

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261

Index Terms

Links

enrichment factor (EF) seawater

182

soil

234

324

environmental behaviour

320

chemicals

143

225

279

compartments

314

316

321

279

Environmental Quality Standards (EQSs)

136

153

environmental reservoirs

170

environmental organic chemistry (EOC)

equilibrium constant

equilibrium criteria model (EQM)

296

estuaries

212

ethinylestradiol

154

ethylbenzene

157

ethylhexyl methoxy cinnamate

155

eutrophication

145

exchangeable acidity

254

exchangeable cations

240

extended DH equations

192

216

219

257

270

F feldspars

239

ferric sulfate

148

ferrihydrite

246

‘ferromagnesian’ minerals

240

ferromanganese nodules

211

217

ferrous sulfate

136

253

flocculation

127

fluoride

272

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340

Index Terms

Links

flux

321

foodstuffs, lead

338

formaldehyde freshwater species distribution

frog development

153

fulvic acid

252

282

G Gapon equation

269

gastropods, marine

228

geochemical computer models

geogenic sources

geothite

246

Gibbs free energy of transfer, ǻG

111

Gibbs-Helmholz equation

292

gleys

263

global pollution problems

230

gold

137

groundwater

arsenic

130

chromium

135

nitrate

149

selenium

133

Güntelberg approximation Guoy-Chapman model

121

292

317

330

86 268

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Index Terms

Links

H haematite

246

halmyrolysis

212

216

halocarbons halocline

176

halogen chemistry

oxidation

heavy metals

127

142

Henry’s Law

186

341

186

285

constant (H)

288

341 hermaphroditic fish

154

hexavalent chromium

136

human lead exposure

338

humic acid

252

282

humic substances

252

267

282

humin

252

humus

249

215

244

formation

249

fractionation

252

hydrocalumite

136

hydrocarbons

oxidation

sources

282

total petroleum

225

volatile (oil)

225

hydrogen bonding

174

hydrogenolysis

300

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326

Index Terms

Links

hydrogenous material

211

hydrophobic

154

hydrophobic organic chemicals (HOCs)

290

hydrothermal vents

171

hydrous iron oxides

130

hydrous metal oxides

135

hydrous oxides

hydroxyapatite

148

hydroxyl radical

hyperkeratosis

130

hyperpigmentation

130

hypolimnion

146

hypoxia

189

216

204

217

246

267

297

318

317

337

I ibuprofen

155

ice Antarctic

317

Arctic

181

rafting

213

illite

211

imposex

228

incongruent dissolution inherited minerals inner sphere complexation internal loading, phosphorus iodine monoxide

243

92 238 95 147 54

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Index Terms

Links

ion(s) activity

265

exchange

215

267

ionic potential

183

263

ionic strength

110

pair formation

266

Irish Sea

324

325

iron

112

119

biological productivity limiting

200

estuarine

221

iron oxides hydrous oxides

115

272

130

irrigation

132

isomorphous substitution

242

isopleth

126

265

150

isoproturon

152

isopycnals

177

Itai-Itai

133

K kaolinite

211

243

Kaw

286

288

286

288

Kesterson Reservoir

135

Koa

286

291

298

342

Kow

286

287

298

342

298

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Index Terms

Links

L lake sediment

123

139

Langmuir equation

272

latent heat

174

leaching

257

lead

131

323

139

202

336

339

333 biogeochemical cycle

335

human exposure

338

lead-210

139

leaded petrol

317

lead halides

336

Leblanc process

143

lifetime(s)

behavioural influence

319

calculation

317

ligand

206

ligand exchange

271

liming

132

149

lindane

152

293

linuron

152

lithogenous material

210

long range transport

295

lysocline

196

255

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317

Index Terms

Links

M MacInnes convention

magnesium

206

manganese

112

manganese oxides

246

marine boundary layer

204

171

144

207

197 marine sediment

210

MCPA

152

mecoprop

152

mercury

137

mesopause

mesosphere

metallothionein

205

methaemoglobinaemia

149

methane

315

methyl chloroform

319

methyl mercury

138

methyl tert-butyl ether (MTBE)

156

methylation

209

4-methylbenzylidene camphor

155

micas

239

mineral components, soil

238

Minimata, Japan

138

MINTEQA2

mixed equilibrium constant

mixing times

316

MODEL V

122

‘model world’ environment

340

208

317

318

121

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179

Index Terms

Links

molal concentration

285

molecular viscosity

174

molybdenum

142

monomethylarsonic acid

131

montmorillonite

211

216

Montreal protocol MTBE

156

mutagen

300

145

N natural organic matter

Nernst Equation

111

NICA

122

nickel

202

nitrate

nitrate radical

nitric acid

149

199

272

322

nitrification

257

335

nitrite

150

nitroarenes

299

nitrogen biogeochemical cycling eutrophication fixation oxidation, atmospheric in seawater

200

335

149

337

200 26 185

200

nitrogen dioxide

318

320

nitrogen oxides

335

336

N-nitroso compounds

150

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Index Terms

Links

nitrous oxide

315

non-conservative behaviour

220

non-exchangeable acidity

254

non-methane hydrocarbons (NMHC)

non-sea salt sulfate (NSSS)

183

4-nonylphenol

151

nonylphenol (NP)

154

North Sea

324

NOx see nitrogen oxides NOy

O O2 minimum

188

ocean(s)

170

circulation

179

conveyor belt

179

196

200

213

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