An Introduction to Environmental Chemistry E-Book

Julian E. Andrews

• Herausgeber: John Wiley & Sons
• Kategorie: Wissenschaft und neue Technologien
• Sprache: Englisch

Seitenzahl: 505

Veröffentlichungsjahr: 2013

Beschreibung

This introductory text explains the fundamentals of the chemistry of the natural environment and the effects of mankind's activities on the earth's chemical systems.

• Retains an emphasis on describing how natural geochemical processes operate over a variety of scales in time and space, and how the effects of human perturbation can be measured.
• Topics range from familiar global issues such as atmospheric pollution and its effect on global warming and ozone destruction, to microbiological processes that cause pollution of drinking water deltas.
• Contains sections and information boxes that explain the basic chemistry underpinning the subject covered.
• Each chapter contains a list of further reading on the subject area. Updated case studies.
• No prior chemistry knowledge required.
• Suitable for introductory level courses.

Leseprobe

Contents

List of boxes

Preface to the second edition

Preface to the first edition

Acknowledgements

Symbols and abbreviations

1 Introduction

1.1 What is environmental chemistry?

1.2 In the beginning

1.3 Origin and evolution of the Earth

1.4 Human effects on biogeochemical cycles?

1.5 The structure of this book

1.6 Internet keywords

1.7 Further reading

1.8 Internet search keywords

2 Environmental chemist’s toolbox

2.1 About this chapter

2.2 Order in the elements?

2.3 Bonding

2.4 Using chemical equations

2.5 Describing amounts of substances: the mole

2.6 Concentration and activity

2.7 Organic molecules – structure and chemistry

2.8 Radioactivity of elements

2.9 Finding more chemical tools in this book

2.10 Further reading

2.11 Internet search keywords

3 The atmosphere

3.1 Introduction

3.2 Composition of the atmosphere

3.3 Steady state or equilibrium?

3.4 Natural sources

3.5 Reactivity of trace substances in the atmosphere

3.6 The urban atmosphere

3.7 Air pollution and health

3.8 Effects of air pollution

3.9 Removal processes

3.10 Chemistry of the stratosphere

3.11 Further reading

3.12 Internet search keywords

4 The chemistry of continental solids

4.1 The terrestrial environment, crust and material cycling

4.2 The structure of silicate minerals

4.3 Weathering processes

4.4 Mechanisms of chemical weathering

4.5 Clay minerals

4.6 Formation of soils

4.7 Wider controls on soil and clay mineral formation

4.8 Ion exchange and soil pH

4.9 Soil structure and classification

4.10 Contaminated land

4.11 Further reading

4.12 Internet search keywords

5 The chemistry of continental waters

5.1 Introduction

5.2 Element chemistry

5.3 Water chemistry and weathering regimes

5.4 Aluminium solubility and acidity

5.5 Biological processes

5.6 Heavy metal contamination

5.7 Contamination of groundwater

5.8 Further reading

5.9 Internet search keywords

6 The oceans

6.1 Introduction

6.2 Estuarine processes

6.3 Major ion chemistry of seawater

6.4 Chemical cycling of major ions

6.5 Minor chemical components in seawater

6.6 The role of iron as a nutrient in the oceans

6.7 Ocean circulation and its effects on trace element distribution

6.8 Anthropogenic effects on ocean chemistry

6.9 Further reading

6.10 Internet search keywords

7 Global change

7.1 Why study global-scale environmental chemistry?

7.2 The carbon cycle

7.3 The sulphur cycle

7.4 Persistent organic pollutants

7.5 Further reading

7.6 Internet search keywords

Supplemental images

Index

© 2004 by Blackwell Science Ltda Blackwell Publishing company

350 Main Street, Malden, MA 02148-5020, USA108 Cowley Road, Oxford OX4 1JF, UK550 Swanston Street, Carlton, Victoria 3053, Australia

The right of J.E. Andrews, P. Brimblecombe, T.D. Jickells, P.S. Liss and B. Reid to be identified as the Authors of this Work has been asserted in accordance with the UK Copyright, Designs, and Patents Act 1988.

All rights reserved. No part of this publication may be reproduced, stored in a retrieval system, or transmitted, in any form or by any means, electronic, mechanical, photocopying, recording or otherwise, except as permitted by the UK Copyright, Designs, and Patents Act 1988, without the prior permission of the publisher.

First published 1996 by Blackwell Science LtdSecond edition 2004

Library of Congress Cataloging-in-Publication Data

An introduction to environmental chemistry / J.E. Andrews … [et al.]. – 2nd ed.

p. cm.

Includes bibliographical references and index.

ISBN 0-632-05905-2 (pbk.: alk. paper)

1. Environmental geochemistry. I. Andrews, J.E. (Julian E.)

QE516.4.I57 2004

551.9 – dc21

2003002757

A catalogue record for this title is available from the British Library.

For further information onBlackwell Publishing, visit our website:http://www.blackwellpublishing.com

Boxes

1.1 Elements, atoms and isotopes

3.1 Partial pressure

3.2 Chemical equilibrium

3.3 Acids and bases

3.4 Gas solubility

3.5 The pH scale

3.6 Reactions in photochemical smog

3.7 Acidification of rain droplets

3.8 Removal of sulphur dioxide from an air parcel

4.1 Properties of water and hydrogen bonds

4.2 Electronegativity

4.3 Oxidation and reduction (redox)

4.4 Metastability, reaction kinetics, activation energy and catalysts

4.5 Dissociation

4.6 Isomorphous substitution

4.7 Van der Waals’ forces

4.8 Chemical energy

4.9 Mineral reaction kinetics and solution saturation

4.10 Biopolymers

4.11 Base cations

4.12 Solubility product, mineral solubility and saturation index

4.13 Radon gas: a natural environmental hazard

4.14 Physical and chemical properties that dictate the fate of organic contaminants

4.15 Use of clay catalysts in clean up of environmental contamination

4.16 Mechanisms of microbial degradation and transformation of organic contaminants

5.1 Ionic strength

5.2 Measuring alkalinity

5.3 Worked examples of pH buffering

5.4 Eh-pH diagrams

5.5 Essential and non-essential elements

6.1 Salinity

6.2 Salinity and major ion chemistry of seawater on geological timescales

6.3 Residence times of major ions in seawater

6.4 Ion interactions, ion pairing, ligands and chelation

6.5 Abiological precipitation of calcium carbonate

6.6 Oceanic primary productivity

7.1 Simple box model for ocean carbon dioxide uptake

7.2 The delta notation for expressing stable isotope ratio values

7.3 Chiral compounds

Preface to the Second Edition

In revision of this book we have tried to respond to constructive criticism from reviewers and students who have used the book and at the same time have pruned and grafted various sections where our own experience as teachers has prompted change. Not least, of course, science has moved on in the eight years since we prepared the first edition, so we have had to make some substantial changes to keep up with these developments, especially in the area of global change.

We have tried to retain the ethos of the first edition, using concise and clear examples of processes that emphasize the chemistry involved. We have also tried to highlight how the chemistry, processes or compounds interlink between the chapters and sections, so that no compartment of environmental science is viewed in isolation.

The substantial changes include more emphasis on organic chemistry, soils, contaminants in continental water and remediation of contaminated land. To do this effectively, the terrestrial environments chapter from the first edition has been split into two chapters dealing broadly with solids and water. We have reorganized the box structure of the book and have placed some of the original box material, augmented by new sections, to form a new chapter outlining some of the basic chemical principles that underpin most sections of the book.

Much of the new material has been prepared by Brian Reid, who, in 1999, joined us in the School of Environmental Sciences at the University of East Anglia. Brian has very much strengthened the organic chemistry dimension of the book and we are very pleased to welcome him to the team of authors.

Julian Andrews, Peter Brimblecombe, Tim Jickells, Peter Liss and Brian ReidUniversity of East Anglia, Norwich, UK

Preface to the First Edition

During the 1980s and 1990s environmental issues have attracted a great deal of scientific, political and media attention. Global and regional-scale issues have received much attention, for example, carbon dioxide (CO2) emissions linked with global warming, and the depletion of stratospheric ozone by chlorofluorocarbons (CFCs). Local issues, however, have been treated no less seriously, because their effects are more obvious and immediate. The contamination of water supplies by landfill leachate and the build up of radon gas in domestic dwellings are no longer the property of a few idiosyncratic specialists but the concern of a wide spectrum of the population. It is noteworthy that many of these issues involve understanding chemical reactions and this makes environmental chemistry a particularly important and topical discipline.

We decided the time was right for a new elementary text on environmental chemistry, mainly for students and other readers with little or no previous chemical background. Our aim has been to introduce some of the fundamental chemical principles which are used in studies of environmental chemistry and to illustrate how these apply in various cases, ranging from the global to the local scale. We see no clear boundary between the environmental chemistry of human issues (CO2 emissions, CFCs, etc.) and the environmental geochemistry of the Earth. A strong theme of this book is the importance of understanding how natural geochemical processes operate and have operated over a variety of timescales. Such an understanding provides baseline information against which the effects of human perturbations of chemical processes can be quantified. We have not attempted to be exhaustive in our coverage but have chosen themes which highlight underlying chemical principles.

We have some experience of teaching environmental chemistry to both chemists and non-chemists through our first-year course in Environmental Chemistry, part of our undergraduate degree in Environmental Sciences at the University of East Anglia. For 14 years we used the text by R.W. Raiswell, P. Brimblecombe, D.L. Dent and P.S. Liss, Environmental Chemistry, an earlier University of East Anglia collaborative effort published by Edward Arnold in 1980. The book has served well but is now dated, in part because of the many recent exciting discoveries in environmental chemistry and also partly because the emphasis of the subject has swung toward human concerns and timescales. We have, however, styled parts of the new book on its ‘older cousin’, particularly where the previous book worked well for our students.

In places the coverage of the present book goes beyond our first-year course and leads on towards honours-year courses. We hope that the material covered will be suitable for other introductory university and college courses in environmental science, earth sciences and geography. It may also be suitable for some courses in life and chemical sciences.

Julian Andrews, Peter Brimblecombe, Tim Jickells and Peter LissUniversity of East Anglia, Norwich, UK

Acknowledgements

We would like to thank the following friends and colleagues who have helped us with various aspects of the preparation of this book: Tim Atkinson, Rachel Cave, Tony Greenaway, Robin Haynes, Kevin Hiscock, Alan Kendall, Gill Malin, John McArthur, Rachel Mills, Willard Pinnock, Annika Swindell and Elvin Thurston. Special thanks are due to Nicola McArdle for permission to use some of her sulphur isotope data.

We have used data or modified tables and figures from various sources, which are quoted in the captions. We thank the various authors and publishers for permission to use this material, which has come from the following sources.

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Symbols and Abbreviations

Multiples and submultiples

Symbol

Name

Equivalent

T

tera

10

12

G

giga

10

9

M

mega

10

6

k

kilo

10

3

d

deci

10

−1

c

centi

10

−2

m

milli

10

−3

μ

micro

10

−6

n

nano

10

−9

p

pico

10

−12

Chemical symbols

Symbol

Description

Units

a

activity

mol l

−1

c

concentration

mol l

−1

eq

equivalents

eql

−1

I

ionic strength

mol l

−1

IAP

ion activity product

mol

n

l

−

n

K

equilibrium constant

mol

n

l

−

n

K

´

first dissociation constant

mol

n

l

−

n

K

a

equilibrium constant for acid

mol

n

l

−

n

K

b

equilibrium constant for base

mol

n

l

−

n

K

H

Henry’s law constant

mol l

−1

atm

K

sp

solubility product

mol

n

l

−

n

K

w

equilibrium constant for water

mol

2

l

−2

mol

mole (amount of substance – see Section 2.5)

p

partial pressure

atm

General symbols and abbreviations

Symbol

Description

A

total amount of gas in atmosphere

(aq)

aqueous species

atm

atmosphere (pressure)

ATP

adenosine triphosphate

B[

a

]P

benzo[

a

]pyrene

°C

degrees Celsius (temperature)

CCD

calcite compensation depth

CCN

cloud condensation nuclei

CDT

Canyon Diablo troilite

CEC

cation exchange capacity

CFC

chlorofluorocarbon

CIA

chemical index of alteration

D

deuterium

DDT

2,2-bis-(

p

-chlorophenyl)-1,1,1-trichloroethane

DIC

dissolved inorganic carbon

DIP

dissolved inorganic phosphorus

DMS

dimethyl sulphide

DMSP

beta-dimethylsulphoniopropionate

DNA

deoxyribonucleic acid

DSi

dissolved silicon

E

°

standard electrode potential (V)

e

−

electron

Eh

redox potential (V)

EPA

Environmental Protection Agency

F

flux

FACE

free-air CO

2

enrichment

FAO

Food and Agriculture Organization

Fs

furans

G

Gibbs free energy (kJ mol

−1

)

g

gram (weight)

(g)

gas

GEOSECS

US geochemical ocean sections programme

GtC

gigatonnes expressed as carbon

H

scale height

H

enthalpy (J mol

−1

)

HCFCs

hydrochlorofluorocarbons

HCH

hexachlorocyclohexane

hv

photon of light

IAP

ion activity product

IGBP

International Geosphere–Biosphere Programme

IPCC

Intergovernmental Panel on Climate Change

J

joule (energy, quantity of heat)

K

kelvin (temperature)

l

litre (volume)

(l)

liquid

ln

natural logarithm

log

10

base 10 logarithm

m

metre (length)

M

a third body

MSA

methanesulphonic acid

N

neutron number

n

an integer

NAPL

non-aqueous phase liquid

Pa

pascal (pressure)

PAH

polycyclic aromatic hydrocarbon

PAN

peroxyacetylnitrate

PCBs

polychlorinated biphenyls

PCDD

polychlorinated dibenzo-

p

-dioxin

PCDF

polychlorinated dibenzo-

p

-furan

PCP

pentachlorophenol

PM

particulate matter

POP

persistent organic pollutant

ppb

parts per 10

9

ppm

parts per million

r

ionic radius

S

entropy (J mol

−1

K

−1

)

s

second (time)

(s)

solid

SOM

soil organic matter

SRB

sulphate reducing bacteria

SVOC

semi-volatile organic compound

T

absolute temperature (kelvin)

TBT

tributyl tin

TCA

tricarboxylic acid

TDIC

total dissolved inorganic carbon

UNESCO

United Nations Educational, Scientific and Cultural Organisation

USDA

United States Department of Agriculture

UV

ultraviolet (radiation)

V

volt (electrical potential)

V

volume

W

watt (power – J s

−1

)

WHO

World Health Organization

wt%

weight per cent

Z

atomic number

z

charge

|

z

|

charge ignoring sign

Greek symbols

α

alpha particle (radiation)

γ

activity coefficient

γ

gamma particle (radiation)

δ

stable isotope notation (

Box 7.2

)

δ−

partial negative charge

δ+

partial positive charge

τ

change in sum of residence time

Ω

degree of saturation

Constants

F

Faraday constant (6.02 × 10

23

e

−

)

R

gas constant (8.314 J mol

−1

K

−1

)

1

Introduction

1.1 What is environmental chemistry?

It is probably true to say that the term environmental chemistry has no precise definition. It means different things to different people. We are not about to offer a new definition. It is clear that environmental chemists are playing their part in the big environmental issues—stratospheric ozone (O3) depletion, global warming and the like. Similarly, the role of environmental chemistry in regional-scale and local problems—for example, the effects of acid rain or contamination of water resources—is well established. This brief discussion illustrates the clear link in our minds between environmental chemistry and human beings. For many people, ‘environmental chemistry’ is implicitly linked to ‘pollution’. We hope this book demonstrates that such a view is limited and shows that ‘environmental chemistry’ has a much wider scope.

Terms like contamination and pollution have little meaning without a frame of reference for comparison. How can we hope to understand the behaviour and impacts of chemical contaminants without understanding how natural chemical systems work? For many years a relatively small group of scientists has been steadily unravelling how the chemical systems of the Earth work, both today and in the geological past. The discussions in this book draw on a small fraction of this material. Our aim is to demonstrate the various scales, rates and types of natural chemical processes that occur on Earth. We also attempt to show the actual or possible effects that humans may have on natural chemical systems. The importance of human influences is usually most clear when direct comparison with the unperturbed, natural systems is possible.

This book deals mainly with the Earth as it is today, or as it has been over the last few million years, with the chemistry of water on the planet’s surface a recurrent theme. This theme emphasizes the link between natural chemical systems and organisms, not least humans, since water is the key compound in sustaining life itself. We will start by explaining how the main components of the nearsurface Earth—the crust, oceans and atmosphere—originated and how their broad chemical composition evolved. Since all chemical compounds are built from atoms of individual elements (Box 1.1), we begin with the origin of these fundamental chemical components.

1.2 In the beginning

It is believed that the universe began at a single instant in an enormous explosion, often called the big bang. Astronomers still find evidence of this explosion in the movement of galaxies and the microwave background radiation once associated with the primeval fireball. In the first fractions of a second after the big bang, the amount of matter and radiation, at a ratio of about 1 in 108, was fixed. Minutes later the relative abundances of hydrogen (H), deuterium (D) and helium (He) were determined. Heavier elements had to await the formation and processing of these gases within stars. Elements as heavy as iron (Fe) can be made in the cores of stars, while stars which end their lives as explosive supernovae can produce much heavier elements.

Hydrogen and helium are the most abundant elements in the universe, relics of the earliest moments in element production. However, it is the stellar production process that led to the characteristic cosmic abundance of the elements (Fig. 1.1). Lithium (Li), beryllium (Be) and boron (B) are not very stable in stellar interiors, hence the low abundance of these light elements in the universe. Carbon (C), nitrogen (N) and oxygen (O) are formed in an efficient cyclic process in stars that leads to their relatively high abundance. Silicon (Si) is rather resistant to photodissociation (destruction by light) in stars, so it is also abundant and dominates the rocky world we see about us.

1.3 Origin and evolution of the Earth

The planets of our solar system probably formed from a disc-shaped cloud of hot gases, the remnants of a stellar supernova. Condensing vapours formed solids that coalesced into small bodies (planetesimals), and accretion of these built the dense inner planets (Mercury to Mars). The larger outer planets, being more distant from the sun, are composed of lower-density gases, which condensed at much cooler temperatures.

As the early Earth accreted to something like its present mass some 4.5 billion years ago, it heated up, mainly due to the radioactive decay of unstable isotopes (Box 1.1) and partly by trapping kinetic energy from planetesimal impacts. This heating melted iron and nickel (Ni) and their high densities allowed them to sink to the centre of the planet, forming the core. Subsequent cooling allowed solidification of the remaining material into the mantle of MgFeSiO3 composition (Fig. 1.2).

Fig. 1.1 The cosmic abundance of elements. The relative abundance of elements (vertical axis) is defined as the number of atoms of each element per 106 atoms of silicon and is plotted on a logarithmic scale.

1.3.1 Formation of the crust and atmosphere

The crust, hydrosphere and atmosphere formed mainly by release of materials from within the upper mantle of the early Earth. Today, ocean crust forms at midocean ridges, accompanied by the release of gases and small amounts of water. Similar processes probably accounted for crustal production on the early Earth, forming a shell of rock less than 0.0001% of the volume of the whole planet (Fig. 1.2). The composition of this shell, which makes up the continents and ocean crust, has evolved over time, essentially distilling elements from the mantle by partial melting at about 100 km depth. The average chemical composition of the present crust (Fig. 1.3) shows that oxygen is the most abundant element, combined in various ways with silicon, aluminium (Al) and other elements to form silicate minerals.

Fig. 1.2 Schematic cross-section of the Earth. Silica is concentrated in the crust relative to the mantle. After Raiswell et al. (1980).

Various lines of evidence suggest that volatile elements escaped (degassed) from the mantle by volcanic eruptions associated with crust building. Some of these gases were retained to form the atmosphere once surface temperatures were cool enough and gravitational attraction was strong enough. The primitive atmosphere was probably composed of carbon dioxide (CO2) and nitrogen gas (N2) with some hydrogen and water vapour. Evolution towards the modern oxidizing atmosphere did not occur until life began to develop.

1.3.2 The hydrosphere

Water, in its three phases, liquid water, ice and water vapour, is highly abundant at the Earth’s surface, having a volume of 1.4 billion km3. Nearly all of this water (>97%) is stored in the oceans, while most of the rest forms the polar ice-caps and glaciers (Table 1.1). Continental freshwaters represent less than 1% of the total volume, and most of this is groundwater. The atmosphere contains comparatively little water (as vapour) (Table 1.1). Collectively, these reservoirs of water are called the hydrosphere.

Fig. 1.3 Percentage of major elements in the Earth’s crust.

Table 1.1 Inventory of water at the Earth’s surface. After Speidel and Agnew (1982).

The source of water for the formation of the hydrosphere is problematical. Some meteorites contain up to 20% water in bonded hydroxyl (OH) groups, while bombardment of the proto-Earth by comets rich in water vapour is another possible source. Whatever the origin, once the Earth’s surface cooled to 100°C, water vapour, degassing from the mantle, was able to condense. Mineralogical evidence suggests water was present on the Earth’s surface by 4.4 billion years ago, soon after accretion, and we know from the existence of sedimentary rocks laid down in water that the oceans had formed by at least 3.8 billion years ago.

Fig. 1.4 Schematic diagram of the hydrological cycle. Numbers in parentheses are reservoir inventories (106 km3). Fluxes are in 106 km3 yr−1. After Speidel and Agnew (1982).

Very little water vapour escapes from the atmosphere to space because, at about 15 km height, the low temperature causes the vapour to condense and fall to lower levels. It is also thought that very little water degasses from the mantle today. These observations suggest that, after the main phase of degassing, the total volume of water at the Earth’s surface changed little over geological time.

Cycling between reservoirs in the hydrosphere is known as the hydrological cycle (shown schematically in Fig. 1.4). Although the volume of water vapour contained in the atmosphere is small, water is constantly moving through this reservoir. Water evaporates from the oceans and land surface and is transported within air masses. Despite a short residence time (see Section 3.3) in the atmosphere, typically 10 days, the average transport distance is about 1000 km. The water vapour is then returned to either the oceans or the continents as snow or rain. Most rain falling on the continents seeps into sediments and porous or fractured rock to form groundwater; the rest flows on the surface as rivers, or re-evaporates to the atmosphere. Since the total mass of water in the hydrosphere is relatively constant over time, evaporation and precipitation must balance for the Earth as a whole, despite locally large differences between wet and arid regions.

The rapid transport of water vapour in the atmosphere is driven by incoming solar radiation. Almost all the radiation that reaches the crust is used to evaporate liquid water to form atmospheric water vapour. The energy used in this transformation, which is then held in the vapour, is called latent heat. Most of the remaining radiation is absorbed into the crust with decreasing efficiency with increasing latitude, mainly because of the Earth’s spherical shape. Solar rays hit the Earth’s surface at 90 degrees at the equator, but at decreasing angles with increasing latitude, approaching 0 degrees at the poles. Thus, a similar amount of radiation is spread over a larger area at higher latitudes compared with the equator (Fig. 1.5). The variation of incoming radiation with latitude is not balanced by an opposite effect for radiation leaving the Earth, so the result is an overall radiation imbalance. The poles, however, do not get progressively colder and the equator warmer, because heat moves poleward in warm ocean currents and there is poleward movement of warm air and latent heat (water vapour).

Fig. 1.5 Variation in relative amounts of solar radiation (energy per unit area) with latitude. Equal amounts of energy A and B are spread over a larger area at higher latitude, resulting in reduced intensity of radiation.

1.3.3 The origin of life and evolution of the atmosphere

We do not know which chance events brought about the synthesis of organic molecules or the assembly of metabolizing, self-replicating structures we call organisms, but we can guess at some of the requirements and constraints. In the 1950s there was considerable optimism that the discovery of deoxyribonucleic acid (DNA) and the laboratory synthesis of likely primitive biomolecules from experimental atmospheres rich in methane (CH4) and ammonia (NH3) indicated a clear picture for the origin of life. However, it now seems more likely that the synthesis of biologically important molecules occurred in restricted, specialized environments, such as the surfaces of clay minerals, or in submarine volcanic vents.

Best guesses suggest that life began in the oceans some 4.2–3.8 billion years ago, but there is no fossil record. The oldest known fossils are bacteria, some 3.5 billion years old. In rocks of this age there is fossil evidence of quite advanced metabolisms which utilized solar energy to synthesize organic material. The very earliest of autotrophic (self-feeding) reactions were probably based on sulphur (S), supplied from volcanic vents.

eqn. 1.1

However, by 3.5 billion years ago photochemical splitting of water, or photosynthesis was happening.

eqn. 1.2

(If you are unfamiliar with chemical reactions and notation, see Chapter 2.)

The production of oxygen during photosynthesis had a profound effect. Initially, the oxygen gas (O2) was rapidly consumed, oxidizing reduced compounds and minerals. However, once the rate of supply exceeded consumption, O2 began to build up in the atmosphere. The primitive biosphere, mortally threatened by its own poisonous byproduct (O2), was forced to adapt to this change. It did so by evolving new biogeochemical metabolisms, those that today support the diversity of life on Earth. Gradually an atmosphere of modern composition evolved (see Table 3.1). In addition, oxygen in the stratosphere (see Chapter 3) underwent photochemical reactions, leading to the formation of ozone (O3), protecting the Earth from ultraviolet radiation. This shield allowed higher organisms to colonize the continental land surfaces.

In recent decades a few scientists have argued that the Earth acts like a single living entity rather than a randomly driven geochemical system. There has been much philosophical debate about this issue, often called the Gaia hypothesis, and more recently, Gaia theory. This view, suggested by James Lovelock, argues that biology controls the habitability of the planet, making the atmosphere, oceans and terrestrial environment comfortable to sustain and develop life. There is little consensus about these Gaian notions, but the ideas of Lovelock and others have stimulated active debate about the role of organisms in mediating geochemical cycles. Many scientists use the term ‘biogeochemical cycles’, which acknowledges the role of organisms in influencing geochemical systems.

1.4 Human effects on biogeochemical cycles?

In discussing the chemistry of near-surface environments on Earth it is important to distinguish between different types of alteration to Earth systems caused by humans. Two main categories can be distinguished:

1 Addition to the environment of exotic chemicals as a result of new substances synthesized and manufactured by industry.

2 Change to natural cycles by the addition or subtraction of existing chemicals by normal cyclical and/or human-induced effects.

The first category of chemical change is probably easiest to understand. Some examples of substances which are found in the environment only as a result of human activities are given in Table 1.2 and include pesticides, such as 2,2-bis(p-chlorophenyl)-1,1,1-trichloroethane (DDT), which is broken down by bacteria in the soil to produce a number of other exotic compounds; polychlorinated biphenyls (PCBs), which have many industrial uses and are slow to degrade in the environment; tributyl tin (TBT), which is used in marine paints to inhibit organisms from settling on the hulls of ships; many drugs; some radionuclides; and a range of chlorofluorocarbon compounds (CFCs), which were developed for use as aerosol propellants, as refrigerants and in the manufacture of solid foams.

The list in Table 1.2 is by no means complete. It has been calculated that the chemical industry has synthesized several million different chemicals (mainly organic) never previously seen on Earth. Although only a small fraction of these chemicals are manufactured in commercial quantities, it is estimated that approximately a third of the total production escapes to the environment.

The impact of these exotic substances on the environment is difficult to predict, since there are often no similar natural compounds whose behaviour is understood. A new substance may be benign, but our lack of knowledge can lead to unforeseen and sometimes harmful consequences. For example, because of the chemical inertness of the CFCs, when they were first introduced it was assumed that they would be completely harmless in the environment. This was true in all environmental reservoirs except the upper layers of the atmosphere (stratosphere), where they were broken down by solar radiation. The breakdown products of CFCs led to destruction of ozone (O3), which forms a natural barrier, protecting animal and plant life from harmful ultraviolet (UV) radiation coming from the sun (see Section 3.10).

Table 1.2 Examples of substances found in the environment only as a result of human activities.

The second category of chemical changes is concerned with natural or humaninduced alterations to existing cycles. These types of changes are illustrated in Chapter 7 with the elements carbon and sulphur. The cycling of these elements has occurred throughout the 4.5 billion years of Earth history. Furthermore, the appearance of life on the planet had a profound influence on both cycles. As well as being affected by biology, the cycles of carbon and sulphur are also influenced by alterations in physical properties, such as temperature, which have varied substantially during Earth history—for example, between glacial and interglacial periods. It is also clear that changes in the cycles of carbon and sulphur can influence climate, by affecting variables such as cloud cover and temperature. In the last few hundred years, the activities of humans have perturbed both these and other natural cycles. Such anthropogenic changes to natural cycles essentially mimic and in some cases enhance or speed up what nature does anyway.

In contrast to the situation for exotic chemicals described earlier, changes to natural cycles should be easier to predict, since the process is one of enhancement of what already occurs, rather than addition of something completely new. Thus, knowledge of how a natural system works now and has done in the past should be helpful in predicting the effects of human-induced changes. However, we are often less able at such predictions than we would like to be, because of our ignorance of the past and present mode of operation of natural chemical cycles.

1.5 The structure of this book

In the following chapters we describe how components of the Earth’s chemical systems operate. Chapter 2 is a ‘toolbox’ of fundamental concepts underpinning environmental chemistry. We do not expect all readers will need to pick up these ‘tools’, but they are available for those who need them. The emphasis in each of the following chapters is different, reflecting the wide range of chemical compositions and rates of reactions that occur in near-surface Earth environments. The modern atmosphere (see Chapter 3), where rates of reaction are rapid, is strongly influenced by human activities both at ground level, and way up in the stratosphere. In terrestrial environments (see Chapters 4 & 5), a huge range of solid and fluid processes interact. The emphasis here is on weathering processes and their influence on the chemical composition of sediments, soils and continental surface waters. Human influence in the contamination of soils and natural waters is also a strong theme. Terrestrial weathering links through to the oceans (see Chapter 6) as the major input of constituents to seawater. It soon becomes clear, however, that the chemical composition of this vast water reservoir is controlled by a host of other physical, biological and chemical processes. Chapter 7 examines environmental chemistry on a global scale, integrating information from earlier chapters and, in particular, focusing on the influence of humans on global chemical processes. The short-term carbon and sulphur cycles are examples of natural chemical cycles perturbed by human activities. Persistent organic pollutants (POPs) are used as examples of exotic chemicals that persist for years to decades in soils or sediments and for several days in the atmosphere. Their persistence has allowed them to be transported globally, often impacting environments remote from their place of manufacture and use. In all of these chapters we have chosen subjects and case studies that demonstrate the chemical principles involved. To help clarify our main themes we provide information boxes that describe, in simple terms, some of the laws, assumptions and techniques used by chemists.

1.6 Internet keywords

There is now a wealth of information available on the Internet (worldwide web, www). In an environmental chemistry context there are many thousands of sites that provide quality information. Information ranges from lecture notes and problems set by university and college staff, through society web pages, to pages managed by government institutions. These pages have the advantage of many excellent colour illustrations and photographs. The information can be used to consolidate on material covered in this book, or as way of starting to explore a subject in more depth. To help you find material on the Internet, at the end of each chapter we have included a list of keywords or phrases as input for search engines. We use keywords rather than specific site addresses as website addresses change rapidly and would soon become dated in a book. The keyword lists are not intended to be complete, but are based on the main themes discussed in each chapter. You will be able to adapt the keywords or think up your own. We have personally checked each of the keywords included in the lists and know they give sensible outcomes.

We do, however, ask you to take care in your Internet searches. Remember, unlike scientific books and papers, there has been no peer review of material. If you are unsure about the quality of information on a specific site do check with your course teachers. They will be able to advise you on the validity of information.

Finally, when using search engines we advise you to use a variety of search options. Advanced search options that search for exact word strings are better for finding specific factual sites, whereas wider, less-constrained searches, usually find more diverse sites. Be as specific as you can. For example, if you are interested in ion exchange in soils use the phrase ‘ion exchange soil’ rather than ‘ion exchange’. This will help you home in to the subject of interest much more efficiently.

1.7. Further reading

Allegre, C. (1992) From Stone to Star. Harvard University Press, Cambridge, Massachusetts.

Broecker, W.S. (1985) How to Build a Habitable Planet. Lamont-Doherty Geological Observatory, Columbia University, Palisades, New York.

Lovelock, J. (1982) Gaia: A New Look At Life on Earth. Oxford University Press, Oxford, 157pp.

Lovelock, J. (1988) The Ages of Gaia. Oxford University Press, Oxford, 252pp.

1.8 Internet search keywords

big bang

formation chemical elements

elements stars

elements isotopes

differentiation Earth

origin atmosphere

origin hydrosphere

hydrological cycle

origin life earth

photosynthesis

Gaia theory

2

Environmental Chemist’s Toolbox

2.1 About this chapter

Undergraduate students studying environmental science come from a wide variety of academic backgrounds. Some have quite advanced chemical knowledge, while others have almost none. Whatever your background, we want you to understand some of the chemical details encountered in environmental issues and problems. To do this you will need some basic understanding of fundamental chemistry. As a rule, we find most students like to learn a particular aspect of chemistry where they need it to understand a specific problem. Learning material for a specific application is much easier than wading through pages of what can seem rather dull or irrelevant facts. Consequently much of the basic chemistry is distributed throughout the book in boxes, sited where the concept is first needed to understand a term or process.

Some of the basic chemistry is, however, so fundamental—underpinning most sections of the book—that we describe it here in a dedicated chapter. We have laid out enough information for students with little or no chemistry background to get a foothold into the subject. You may only need to ‘dip’ into this material. We certainly don’t expect you to read this chapter from beginning to end. Imagine the contents here as tools in a toolbox. Take out the tool (= facts, laws, etc.) you need to get the job (= understanding an aspect of environmental chemistry) done. Some of you will not need to read this chapter at all, and can move on to the more exciting parts of the book!

2.2 Order in the elements?

Most of the chemistry in this book revolves around elements and isotopes (see Box 1.1). It is therefore helpful to understand how the atomic number (Z) of an element, and its electron energy levels allow an element to be classified. The electron is the component of the atom used in bonding (Section 2.3). During bonding, electrons are either donated from one atom to another, or shared; in either case the electron is prised away from the atom. One way of ordering the elements is therefore to determine how easy it is to remove an electron from its atom. Chemists call the energy input required to detach the loosest electron from atoms, the ionization energy. As explained in Box 1.1, the number of positively charged components (protons, Z) in an atom is balanced by the same number of negatively charged electrons that form a ‘cloud’ around the nucleus. Although electrons do not follow precise orbits around the nucleus, they do occupy specific spatial domains called orbitals. We need only think in terms of layers of these orbitals. Those electrons in orbitals nearest the nucleus are tightly held by electrostatic attraction-forming core electrons that never take part in chemical reactions. Those further away from the nucleus are less tightly held and may be used in ‘transactions’ with other atoms. These loosely held electrons are known as valence electrons. Electrons normally occupy spaces available in the lowest energy orbitals such that energy dictates the electron distribution around the nucleus. The valence electrons reside in the highest occupied energy levels and are thus the easiest to remove. For example, the element sodium (Na) has a Z number of 11. This means that sodium has 11 electrons, 10 of which are core electrons, and one valence electron. It is this single valence electron that dictates the way sodium behaves in chemical reactions.

If the ionization energy is corrected to account for nuclear charge (Fig. 2.1b)—because increasing nuclear charge makes electron removal more difficult—the energy pattern in each period becomes more like a ramp. Each ‘period’ begins with an element of conspicuously low ionization energy, the so-called alkali metals (Li, Na and K). Each of these elements readily lose their single valence electron to form singly charged or monovalent ions (Li+, Na+ and K+). The periods of elements are depicted as ‘rows’ in the Periodic Table (Fig. 2.2), and when these rows are stacked on top of one another a series of ‘columns’ result (Fig. 2.2). Column Ia depicts the alkali metals. Moving up the energy ramps in Fig. 2.1b, the alkali metals are followed by the elements beryllium (Be), magnesium (Mg) and calcium (Ca), each with two, relatively easily removed valence electrons. These elements form doubly charged or divalent ions (Be2+, Mg2+ and Ca2+) and are known as alkali earth metals (column IIa in the Periodic Table). Continued progression up each energy ramp in Fig. 2.1b results in predictable patterns. For example, Mg is followed by aluminium (Al) which has three valence electrons, and then silicon with four valence electrons. Progressively more energy is required to remove these electrons due to the increasing nuclear attraction. This means that aluminium will form trivalent cations whereas silicon typically will not: instead it shares its electrons in covalent bonds (Section 2.3), except in one special case (Section 2.3.2). At the top of each energy ramp are the elements He, Ne and Ar that cling tenaciously to all of their electrons. These elements have no valence electrons and therefore no significant chemical reactivity. These chemically inert elements are often called the inert or noble gases and form column O on the far right of the Periodic Table (Fig. 2.2).

Fig. 2.2 Periodic Table of the elements and their Z numbers. Note that the periodic pattern is complicated by the transition metals between columns II and III. *La and the lanthanides are known as the rare earth elements (REE). The table has been constructed using conventional terminology and further details can be found in basic chemistry textbooks. Gill (1996) gives an accessible summary with a strong applied earth science stance. Elements in bold are those most abundant in environmental materials (see Fig. 2.3). After Gill (1996), with kind permission of Kluwer Academic Publishers.

Although the periodic pattern becomes more complicated above Z values of 20, the overall ordering persists. Complications arise in the so-called transition elements that occupy a position between columns II and III of the Periodic Table (Fig. 2.2). These elements have between one and three valence electrons. Importantly, however, the electrons in the orbital below the valence electrons have almost the same energy as the valence electrons themselves. In some compounds, usually depending on oxidation state (see Box 4.3), these lower orbital electrons act as additional valence electrons. For example, the element iron (Fe) exists in compounds in a reduced (Fe2+ or ferrous iron) and oxidized (Fe3+ or ferric iron) state. In general, the transition metals are less regular in their atomic properties when compared to the main groups, which also makes their behaviour more complicated to predict in nature.

It is clear from the discussion above, and by looking at the Periodic Table (Fig. 2.2) that some elements are classed as metals, some as semi-metals and some as non-metals. In each row of the Periodic Table the degree of metallic character decreases progressively from left to right, i.e. up the energy ramps of Fig. 2.1b. In essence this is because those elements with low ionization energy hold electrons loosely. In an applied electrical voltage these excited electrons will flow, conducting the electricity, whereas in non-metals there is a gap in the electron configuration that will not allow passage of excited electrons. In the case of semimetals the gap in electron configuration is small enough that excited electrons can jump through, but only when activated by an external energy source. In effect the semi-metal flips between being an insulator (when not stimulated by external energy) and a conductor (when stimulated by external energy). Semi-metals such as silicon are also known as semi-conductors, and are used in various industrial applications to speed up electrical processes, most famously as the key component of the ‘silicon chip’ in computer microprocessors.

There have been many attempts to further classify the elements geologically and environmentally. In Fig. 2.3 we show the most abundant elements in four of the main environmental materials of the Earth. A glance at this figure shows that oxygen (O), and to a lesser extent hydrogen (H), are superabundant in most Earth surface materials such as air, water, organic matter and silicate minerals. In the lithosphere, silicon (Si) and aluminium (Al) are next most abundant forming the silicate minerals feldspar and quartz (see Chapter 4). In the hydrosphere it is the dissolved ions in seawater (see Chapter 6) that dominate the chemistry, particularly chloride (Cl−) and sodium (Na+), while the main atmospheric gases are nitrogen (N2), oxygen (O2), argon (Ar) and carbon dioxide (CO2), along with water vapour (see Chapter 3). The organic matter of the biosphere is made principally of carbon and hydrogen bonded in various combinations (Section 2.7), along with lesser amounts of oxygen and the nutrient elements nitrogen (N) and phosphorus (P). Based on the information in this diagram it might be tempting to conclude that we need only understand the behaviour of these elements in nature to understand environmental chemistry. In fact the reverse is true. Paradoxically, it is often the elements present in trace amounts in the solids and fluids of the environment that tell us most about chemical processes.

Fig. 2.3 Distribution of elements in the four main environmental materials, lithosphere, hydrosphere, atmosphere and biosphere. The elements are shown in their actual form as compounds, ions or molecules as appropriate. The main components of each material are shown in boxes, other major constituents are shown outside the boxes.

2.3 Bonding

Many elements do not normally exist as atoms, but are bonded together to form molecules. The major components of air, nitrogen and oxygen for example, are present in the lower atmosphere as the molecules N2 and O2. By contrast, argon is rather unusual because as an inert element (or noble gas—Section 2.2) it is found uncombined as single argon atoms. Inert elements are exceptions and most substances in the environment are in the form of molecules.

2.3.1 Covalent bonds

Molecular bonds are formed from the electrostatic interactions between electrons and the nuclei of atoms. There are many different electronic arrangements that lead to bond formation and the type of bond formed influences the properties of the compound that results. It is the outermost electrons of an atom that are involved in bond formation. The archetypical chemical bond is the covalent bond and we can probably best imagine this as formed from outer electrons shared between two atoms. Take the example of two fluorine atoms that form the fluorine molecule:

eqn. 2.1