P.A. Cox - Inorganic chemistry (793955), страница 63
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Figure 1 shows a summary of the main processes, with estimates of thereservoirs (square boxes) and annual fluxes (round-cornered boxes) in units of 1012 kg C.Fig. 1. The carbon cycle, showing reservoirs (square-cornered boxes) and annual fluxes (round-cornered boxes) in units of 1012 kg C.Atmospheric CO2 is cycled in about equal amounts by two different processes: (i) the conversion into solublebicarbonateand the subsequent regeneration of CO2 when water evaporates; (ii) the conversion into biologicalJ6—ENVIRONMENTAL CYCLING AND POLLUTION275carbon compounds by photosynthesis, and reoxidation to CO2 by respiration.
Anaerobic decay of vegetation, ruminantanimals such as cows, and other natural processes produce small amounts of CH4 and CO, which are oxidized in theatmosphere to CO2. Some parts of the cycle operate with much larger reservoirs of carbon, but also much more slowly:they include the mixing of surface bicarbonate with deep ocean waters, the production of sedimentary carbonate rocks(mostly from CaCO3 shells and skeletons of marine organisms) and the eventual decomposition of carbonates by heatingdeep in the crust to regenerate CO2.Different parts of the natural cycle must be very nearly in balance, although over a period of millions of years someorganic carbon has been buried before reoxidation, giving fossil fuels containing reduced carbon in the crust.Dioxygen from photosynthesis has passed into the atmosphere, but over geological time most of it has been used up inoxidizing surface rocks (principally FeII to FeIII compounds, and sulfides to sulfates), only a small fraction remaining asfree O2.
The burning of fossil fuels has reversed this natural trend and currently transfers around 5×1012 kg C per yearinto the atmosphere as CO2. Parts of the cycle outside human control may be responding to take up some of this extrainput, but the capacity of either surface ocean waters or life to accommodate it in the short term is very limited, and theburning of land vegetation contributes to the problem by reducing photosynthesis. Although excess CO2 mustultimately return to the crust as carbonate minerals, that can happen only over time scales measured in thousands oreven millions of years.Other nonmetallic elementsNitrogen and sulfurN and S have a diverse and important environmental chemistry, associated in both cases with the wide range ofoxidation states possible.
Biological nitrogen fixation converts atmospheric N2 into organic compounds needed bylife. A high proportion is recycled within the biosphere, but some microorganisms convert it into nitrate(nitrification) and others reduce nitrate to N2 (dinitrification), both processes being used to obtain metabolicenergy. Denitrification thus recycles N back into the atmosphere. The major human perturbations to the cycle comefrom the use of nitrate fertilizers (which can lead to undesirable concentrations ofin drinking water) and hightemperature burning of fossil fuels, which produce NO and NO2. These gases are air pollutants, locally because they aretoxic and take part in photochemical processes that generate other noxious compounds, and on a wider scale becausethey oxidize to nitric acid, which contributes to acid rain.The biological and atmospheric redox chemistry of sulfur is also complex.
The main natural inputs to the atmospherecome from biological decay (mostly H2S) and emissions of dimethyl sulfide (CH3)2S by marine organisms, together withvolcanic emissions (mostly SO2). These natural sources are now exceeded by the emission of SO2 from burning sulfurcontaining fossil fuels. Most atmospheric sulfur compounds oxidize rapidly to sulfuric acid, which is the majorcomponent of acid rain. The comparison of N and S is interesting, as the total atmospheric inputs of the two elementsare similar in magnitude (1–2×1011 kg per year).
The oxidation and removal of sulfur compounds is much more rapidthan for the very stable N2 molecule, and so the atmospheric concentrations are enormously different (about 1 p.p.b.for sulfur compounds, 78% for N2).Silicon and phosphorusSi and P occur naturally only in fully oxidized forms (SiO2 and silicates, phosphates), which are involatile and have lowsolubility in natural waters. Phosphorus is one of the most important elements of life (see Topic J3) and inaquatic environments the one that is often in shortest supply.
Pollution by soluble polyphosphates (e.g. from276SECTION J—ENVIRONMENTAL, BIOLOGICAL AND INDUSTRIAL ASPECTSdetergents; see Topic J4) can seriously upset the ecological balance of lakes, leading to uncontrolled growth of algae anddepletion in dissolved oxygen.HalogensThese elements occur naturally in halide minerals. CaF2 is very insoluble in water, but other halide ions are easilywashed out of rocks and are abundant in sea water. Volcanic emissions contain small amounts of HF and HCl but thesegases are very soluble and washed out of the atmosphere quickly.
Marine organisms produce small quantities of methylcompounds such as CH3Cl, which are oxidized and also washed out. Some synthetic organohalogen compounds poseenvironmental problems because natural chemical processes break them down very slowly. Organochlorine compoundsof concern include dioxins and persistent insecticides such as DDT, and volatile chlorofluorocarbons (CFCs) used asaerosol propellants and in refrigerators. CFCs resist photochemical breakdown in the lower atmosphere and can enterthe stratosphere where short-wavelength UV radiation splits them to produce Cl atoms, which then act as catalysts forthe decomposition of UV-absorbing ozone.Heavy metalsThe heavy post-transition metals such as Cd, Hg and Pb are toxic because of the very strong complexing ability of ‘soft’cations such as Hg2+ (see Topics G4, G5 and J3).
They have low concentrations in natural waters because they forminsoluble sulfides. Compounds that are either more soluble in water or volatile pose an environmental hazard. Of theseelements, lead has been the most widely used, in pipes for drinking water, in paints and (in the form of tetraethyl leadPb(C2H5)4 as a gasoline additive to improve combustion. As the toxic hazards have been more clearly recognized, theseuses have been phased out.Mercury also had many applications, including in hat-making (where the symptoms of mercury poisoning gave rise tothe saying ‘mad as a hatter’) but its industrial usage (e.g. for NaCl electrolysis; see Topic J4) has also declined.
Cases ofacute mercury poisoning have resulted from eating fish from water polluted by industrial Hg compounds. Someorganisms convert inorganic compounds into ones containing [CH3Hg]+, which are especially toxic as they pass more easilythrough the nonpolar constituents of cell membranes. It is likely that methylcobalamin (see Topic J3) is involved in thistransformation.FURTHER READINGText-books on inorganic chemistry differ greatly in their balance of conceptual and descriptive material. All the bookslisted under the General heading below (except that by Emsley, which is a useful compilation of data) include somediscussion of general concepts.GeneralCotton, F.A., Wilkinson, G.
and Gaus, P.L. (1995) Basic Inorganic Chemistry, 3rd edn., Wiley, New York, USA.Douglas, B., McDaniel, D.H. and Alexander, J.J. (1983) Concepts and Models of Inorganic Chemistry, 2nd edn., Wiley, New York,USA.Emsley, J. (1991) The Elements, 2nd edn., Clarendon Press, Oxford, UK.Huheey, J.E. (1993) Inorganic Chemistry: Principles of Structure and Reactivity, 4th edn., Harper Collins, New York, USA.Mackay, K.M. and Mackay, R.A. (1989) Introduction to Modern Inorganic Chemistry, 4th edn., Blackie, Glasgow, UK.Owen, S.M.
and Brooker, A.T. (1994) A Guide to Modern Inorganic Chemistry, Longman, Harlow, UK.Porterfield, W.W. (1984) Inorganic Chemistry: A Unified Approach, 2nd edn., Academic Press, San Diego, USA.Raynor-Canham, G. (1996) Descriptive Inorganic Chemistry, W.H.Freeman, New York, USA.Sharpe, A.G. (1992) Inorganic Chemistry, 3rd edn., Longman, Harlow, UK.Shriver, D.F.
and Atkins, P.W. (1999) Inorganic Chemistry, 3rd edn., Oxford University Press, Oxford, UK.Section AAtkins, P.W. (1998) Physical Chemistry, 5th edn. Oxford University Press, Oxford, (Ch. 11, 12).Cox, P.A. (1996) Introduction to Quantum Theory and Atomic Structure, Oxford University Press, Oxford.Cox, P.A. (1989) The Elements: Their Origin, Abundance and Distribution, Oxford University Press, Oxford. (Ch. 2).Whittaker, A.G., Mount, A.R. and Heal, M.R. (2000) Instant Notes in Physical Chemistry, BIOS Scientific Publishers, Oxford.Section BAlcock, N.W.
(1990) Bonding and Structure: Structural Principles in Inorganic and Organic Chemistry, Ellis Horwood, Chichester.Ebsworth, E.A.V., Rankin, D.W.H. and Cradock, S. (1991) Structural Methods in Inorganic Chemistry, 2nd edn. Blackwell ScientificPublications, Oxford.Kealey, D. and Haines, P.J. (2002) Instant Notes in Analytical Chemistry, BIOS Scientific Publishers, Oxford.Johnson, D.A. (1982) Some Thermodynamic Aspects of Inorganic Chemistry, 2nd edn. Cambridge University Press, Cambridge.Leigh, G.J. (1990) Nomenclature of Inorganic Chemistry: Recommendations 1990, Blackwell Scientific, Oxford.Mingos, D.M.P.
(1998) Essential Trends in Inorganic Chemistry, Oxford University Press, Oxford.Smith, D.W. (1990) Inorganic Substances: A Prelude to the Study of Descriptive Inorganic Chemistry, Cambridge University Press,Cambridge.278FURTHER READINGSection CDeKock, R.L. and Gray, H.B. (1980) Chemical Structure and Bonding, Benjamin-Cummings, Menlo Park, USA.Kettle, S.F.A. (1995) Symmetry and Structure: Readable Group Theory for Chemists, 2nd edn., Wiley, Chichester, UK.Murrel, J.N., Kettle, S.F.A. and Tedder, J.M. (1978) The Chemical Bond, Wiley, Chichester, UK.Section DCox, P.A.
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