An Introduction to Environmental Chemistry

Julian E. Andrews, Peter Brimblecombe, Tim D. Jickells, Peter S. Liss, Brian Reid

• Publisher: Wiley-Blackwell
• Year: 2013
• ISBN:
  9781118685471
• Edition: 2
• Topic: Physical Sciences
• Subtopic: Organic Chemistry

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.

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).

Box 1.1 Elements, atoms and isotopes

Elements are made from atoms—the smallest particle of an element that can take part in chemical reactions. Atoms have three main components: protons, neutrons and electrons. Protons are positively charged, with a mass similar to that of the hydrogen atom.

Neutrons are uncharged and of equal mass to protons. Electrons are about 1/1836 the mass of protons, with a negative charge of equal value to the (positive) charge of protons.

Atoms are electrically neutral because they have an equal number (Z) of protons and electrons. Z is known as the atomic number and it characterizes the chemical properties of the element.

The atomic weight of an atom is defined by its mass number and most of the mass is present in the nucleus.

eqn. 1

Equation 1 shows that the mass of an element can be changed by altering the number of neutrons. This does not affect the chemical properties of the element (which are determined by Z). Atoms of an element which differ in mass (i.e. N) are called isotopes. For example, all carbon atoms have a Z number of 6, but mass numbers of 12, 13 and 14, written:

In general, when the number of protons and neutrons in the nucleus are almost the same (i.e. differ by one or two), the isotopes are stable. As Z and N numbers become more dissimilar, isotopes tend to be unstable and break down by radioactive decay (usually liberating heat) to a more stable isotope. Unstable isotopes are called radioactive isotopes (see Section 2.8).

Fig. 1 Representation of the hydrogen atom. The dots represent the position of the electron with respect to the nucleus. The electron moves in a wave motion. It has no fixed position relative to the nucleus, but the probability of finding the electron at a given radius (the Bohr radius, r) can be calculated; r = 5.3 × 10⁻⁵μm for hydrogen.

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 ...

Table of contents

• Cover
• Contents
• Title Page
• Copyright
• List of Boxes
• Preface to the Second Edition
• Preface to the First Edition
• Acknowledgements
• Symbols and Abbreviations
• 1 Introduction
• 2 Environmental Chemist’s Toolbox
• 3 The Atmosphere
• 4 The Chemistry of Continental Solids
• 5 The Chemistry of Continental Waters
• 6 The Oceans
• 7 Global Change
• Supplemental images
• Index