b Mass median diameters determined by centrifugal sedimentation.
b Mass median diameters determined by centrifugal sedimentation.
elements, metal ions are almost always assigned a positive oxidation number and are somewhat electrophilic. Because of this they are stabilized by association with electron-rich atoms. In particular, atoms that have a free electron pair can "donate" some of their electron density to the metal to form a bond. Such atoms are Lewis bases. The most common and environmentally important donor atoms are oxygen, nitrogen, and sulfur. The bonds they form with metal ions range in strength from relatively weak associations such as those between a dissolved metal ion and water to very strong covalent bonds. These types of bonds are significant in both aqueous phase reactions and in the formation of insoluble compounds.
The characteristic affinity of metals for electron-rich donor atoms leads to an important distinction between the geochemical behavior of metal ions and that of the elements discussed earlier in this text. Specifically, metal ions in a single oxidation state can bind to donor atoms that are part of a variety of cationic, anionic, or neutral molecules. Thus, the metals can be found in and can move through the environment as parts of molecules spanning the complete range of charge, molecular weight, bioavailability, and other chemical characteristics. In addition, since metals constitute an extremely small fraction of biological organisms, the connection between metal transport through the environment and biological activity is generally less direct and of less quantitative importance than for, say, oxygen or phosphorus. (However, recent evidence is pointing to a more significant role for biota in controlling metal concentrations and transport in the ocean.) In a relative sense, then, purely chemical reactions are of more importance for metals than for the elements discussed earlier. The strength of the chemical bonding between the donor atoms O, N, and S and metals is the overriding factor controlling the geochemical cycling of metals. Once these reactions are understood, the behavior of metals, with respect to both transport and biological impacts, can be placed in a more logical framework.
Metals exist in nature primarily in positive oxidation states, and many form stable compounds in more than one oxidation state. The formal oxidation number of the most common form can range from +1 to +6. The stable form in a given environment depends on the oxidation potential and chemical composition of that environment. Often the stable form at the Earth's surface in the presence of molecular oxygen is different from that which is stable in anoxic sediments or waters.
Thermodynamically, virtually all metals in the elemental form are unstable with respect to redox reactions in environments where they are exposed to air and water, i.e., virtually all environments where they are used. Those metals least likely to oxidize (corrode) were long ago given the distinguished title "noble metals." Efforts to prevent metals from corroding, and the cost of repairing and replacing metal structures that have done so, runs into the billions of dollars annually. Thus, one characteristic feature of the society's use of metals is that the metals are continuously, albeit slowly, "degrading" to a less useful form from the moment they are put into use.
The different oxidation states of a metal can have dramatically different chemical properties, which in turn affect their biogeochemical forms and significance. For example, almost 4 g/L ferrous iron, Fe(II), can dissolve in distilled water maintained at pH 7.0. However, if the water is exposed to air and the iron is oxidized to Fe(III) essentially all the iron will precipitate, reducing the soluble Fe concentration by more than eight orders of magnitude. Oxidation state can also affect a metal ion's toxicity. For instance, the toxicity of As(III) results from its ability to inactivate enzymes, while As(V) interferes with ATP synthesis. The former is considerably more toxic to both aquatic organisms and humans.
The thermodynamically stable oxidation state of a metal in a given environment is a function of the prevailing oxidation potential. The value of the potential is given by the Nernst equation, which is described in Chapter 5 for the generic reduction half-cell:
Reactants + ne~ -> Products
The electrochemical potential that would cause the reactants and products to be equilibrated with each other is
If the equilibrium half-cell potentials for two redox reactions are different, electrons will be transferred from the reduced species in the reaction with the less positive potential to the oxidized form with the more positive potential. The process is repeated until all exchangeable electrons have the same equilibrium potential. Water chemistry texts describe rapid graphical or computerized approaches to solve for the concentrations of all species once this equilibrium condition has been attained.
While these calculations provide information about the ultimate equilibrium conditions, redox reactions are often slow on human time scales, and sometimes even on geological time scales. Furthermore, the reactions in natural systems are complex and may be catalyzed or inhibited by the solids or trace constituents present. There is a dearth of information on the kinetics of redox reactions in such systems, but it is clear that many chemical species commonly found in environmental samples would not be present if equilibrium were attained. Furthermore, the conditions at equilibrium depend on the concentration of other species in the system, many of which are difficult or impossible to determine analytically. Morgan and Stone (1985) reviewed the kinetics of many environmentally important reactions and pointed out that determination of whether an equilibrium model is appropriate in a given situation depends on the relative time constants of the chemical reactions of interest and the physical processes governing the movement of material through the system. This point is discussed in some detail in Section 15.3.8. In the absence of detailed information with which to evaluate these time constants, chemical analysis for metals in each of their oxidation states, rather than equilibrium calculations, must be conducted to evaluate the current state of a system and the biological or geochemical importance of the metals it contains.
To summarize, an evaluation of the oxidation state of metals in an environment is central to determining their probable fate and biological significance. Redox reactions can lead to orders of magnitude changes in the concentration of metals in various phases, and hence in their mode and rate of transport. While equilibrium calculations are a valuable tool for understanding the direction in which changes are likely to occur, field measurements of the concentrations of metals in their various oxidation states are always needed to evaluate metal speciation, since equilibrium has often not yet been attained.
The extent to which any chemical species is volatilized is governed by its vapor pressure, which is sensitive to temperature. Most metals and their compounds have very low vapor pressures at normal temperatures, low enough that their tendency to vaporize can be ignored. The major exceptions are metallic Hg and organometallic compounds. Nevertheless, in some environments significant quantities of metals can be volatilized either as elements or inorganic compounds such as oxides or carbonates. The most obvious such environments are high-temperature furnaces such as in smelters or fossil-fuel-burning power plants and in regions of geothermal activity or vulcanism. While the oxides, sulfates, carbonates, and sulfides of a metal all have somewhat different volatilities, the most volatile metals, regardless of the anion with which they are associated, are Hg, As, Cd, Pb, and Zn. Metallic Hg and organometallic compounds may be transported significant distances as gases and eventually be removed from the atmosphere by dissolution in rain droplets. However, most volatilized metals condense rapidly as they cool and fall to the surface associated with particulate matter, either as dry deposition or scavenged by precipitation. Some of the condensed particles are light enough to carry long distances, and those that fall to Earth may be resuspended by wind action or washed into a water body by surface runoff. Particles produced by high-temperature combustion processes are mostly in the < 2 //m size range and typically have atmospheric residence times of 7 to 14 days, while those generated by soil erosion are larger (> 5 /im) (Hardy et ai, 1985). Anthropogenic particles are typically enriched in trace metals (normalized to the concentration of Al) by a factor of 100 to 10 000 compared with atmospheric particulates gener ated by natural erosion and wind action. As an indication of the importance of volatilization of metals to their overall biogeochemical budgets, Galloway (1979) estimated that volatilization of As, Hg, and Se overwhelms total dust and volcanic emanation rates by factors of 7.5, 625, and 7.3, respectively.
The equilibrium volatility of a species dissolved in water is characterized by its equilibrium constant for the reaction X(g) X(aq), i.e., its Henry's Law constant. These equilibrium constants are related, but are not directly comparable to vapor pressures of the corresponding dry solids because of the effect of the solvent, water. (The most appropriate direct comparison involves the calculation of fugacities. The fugacity, or escaping tendency of a chemical species from the environment in which it exists, depends upon both the concentration of the chemical species of interest and the strength of its interactions with the surrounding (solvent) molecules. Details of the calculation are provided in chemical thermodynamics textbooks and in a number of articles by Mackay and coworkers (e.g. Mackay, 1979).) Suffice it to say that when volatile metal species, such as methyl-mercury, are present in water, there virtually always exists a driving force to strip them out of the aqueous phase, since their partial pressures in air are essentially zero. Thus, equilibrium between gaseous and aqueous phases is rarely a limiting factor. Rather, the factor limiting volatilization of these species is usually the rate at which they move through the water column and across the water-air interface.
The ratio of anthropogenic emissions to total natural emissions is highest for the atmophilic elements Sn, Cu, Cd, Zn, As, Se, Mo, Hg, and Pb (Lantzy and Mackenzie, 1979). In the case of lead, atmospheric concentrations are primarily the consequence of the combustion of leaded gasoline. For many years, lead was used as a gasoline additive, in the form of an organometal compound, tetraethyl lead. When the fuel was burned, most of the lead was converted to inorganic forms and released in the exhaust. The widespread use of this compound as an anti-knock agent in gasoline engines led to its dispersion everywhere automobiles travel, and from there to very remote locations. As with metals near smelters, lead concentrations in the soil near major roadways decrease rapidly with distance from the source, but significant amounts of the metal are transported by wind, either directly after being emitted or by resuspension after a period of deposition.
Atmospheric fluxes of lead in the United States rose steadily from the first decades of this century, reaching a maximum in the early 1970s (see Eisenrich et al., 1986 and references therein). Passage of the Clean Air Act of 1972 and its subsequent amendments resulted in dramatic reductions in atmospheric lead concentrations, although lead fluxes worldwide still remain 10-1000 times background levels (Settle et al., 1982; Settle and Patterson, 1982).
Volatilization is also a dominant transport mode for mercury, which is the most volatile metal in its elemental state. As with lead, a key reaction that can increase the volatility of mercury is formation of an organometallic compound. In this case, the reactions take place in water and are primarily biological, being mediated by bacteria commonly found in the upper levels of sediments. These reactions and their importance in the global mercury cycle are discussed in some detail later in the chapter.
To summarize, metals can be transferred into the gas phase in high-temperature processes either in their elemental form or as inorganic compounds, and these compounds can then be transported long distances as gases or in other physical-chemical forms. Many organometallic compounds are also volatile. In a few cases, natural organometallic compounds may be formed that are volatile. These compounds are formed by microorganisms in mildly reducing aquatic environments and are then transported to the surface and across the air-water interface to enter the gas phase. Methylmercury compounds are probably the most important and certainly have been the most widely studied of these because of their central role in the bioaccumulation of mercury; others, such as methylar-
sine and organotin compounds, are environmentally important as well. Anthropogenic release to the atmosphere overwhelms natural sources for Hg, Cd, Cu, Ag, Zn, Pb, As, and Se, and possibly other metals.
The stability of liquid water is due in large part to the ability of water molecules to form hydrogen bonds with one another. Such bonds tend to stabilize the molecules in a pattern where the hydrogens of one water molecule are adjacent to oxygens of other water molecules. When chemical species dissolve, they must insert themselves into this matrix, and in the process break some of the bonds that exist between the water molecules. If a substance can form strong bonds with water, its dissolution will be thermodynamically favored, i.e., it will be highly soluble. Similarly, dissolution of a molecule that breaks water-to-water bonds and replaces these with weaker water-to-solute bonds will be energetically unfavorable, i.e., it will be relatively insoluble. These principles are presented schematically in Fig. 15-1.
Cationic metal ions form strong bonds with water molecules. By orienting themselves in such a way that the metal "faces" an oxygen atom, the negative charge on the oxygen is partially distributed onto the metal, forming the analog of a strong hydrogen bond. Similarly, some of the charge on the metal ion is neutralized. (Recall the electrophilic nature of metal ions.) These bonds are strong enough that, in most aqueous environments, most metal ions are surrounded by an "inner hydration sphere" of four to eight strongly bound water molecules, as well as a loosely attached "outer hydration sphere" of variable size. Although chemical convention represents dissolved metal ions as Me"+, a more accurate designation is Me(H20)"+. The strength of the metal-to-water bond increases with decreasing size and increasing charge of the metal ion, and also depends on the distribution of electrons around it.
The water molecules in the inner hydration sphere can undergo dissociation reactions just as water molecules far from a dissolved metal ion
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