Trace Elements in Terrestrial Environments

Metals external to the organism are unlikely to cause any adverse effect in that organism but may do so once absorbed (or taken up) and assimilated. This implies that as a prelude step, metals have to come in contact with the organism to be of any benefit or consequence to that organism. In turn, metals have to be in a particular form to be able to enter an organism. In essence, for a contaminant to be assimilated, it will have to be mobile and transported and be bioavailable to the organism. Because bioavailability (also known as bioaccessibility) may have different meanings to different disciplines, it is prudent to define it according to the receptor organism.

Environmental contamination as a result of human activities is not a recent phenomenon. Cicero first related structural damage of buildings and statues in Rome to smoky rains of wood and charcoal burning about 2100 years BP (Eney and Petzold, 1987). Some of the detrimental effects by mining activities on human health had been recognized a long time ago. The Romans used slaves to extract cinnabar (a Hg-containing ore) at the Almaden mine in Spain. Due to acute Hg exposure, the miner’s life expectancy was only about three years (Wren et al., 1995). Peat cores from a Swiss bog indicate that As, Sb, and Pb fluxes due to anthropogenic activities have been exceeding natural fluxes for more than 2000 years (Shotyk et al., 1998). The present enrichment factors in this bog are in the order of 20 times for Sb, 70 for Sb, and 130 for Pb. Modifications of the natural cycles of metals have led to a situation in which the inputs of metals in soils generally exceed the removal due to harvest of agricultural crops and the losses by leaching, volatilization, etc. (Jones, 1991; Van Driel and Smilde, 1990). The Industrial Revolution started in the mid-1800s and the large use of coal to produce energy caused the release of considerable amounts of gas, e.g., CO₂, SO_x, NO_x, and fly ash into the atmosphere. Since then, the biogeochemical cycles of inorganic contaminants (e.g., metals) naturally present in the environment have been largely accelerated by human activities. The conversion of the world’s economy from coal to oil, initiated between the two world wars, enlarged the range of contaminants released in the environment to organic compounds, e.g., PAHs.

The dissemination of metals in the environment from industrial activities ensures that a significant portion of the population is exposed to these contaminants. Depending on the degree of exposure and environmental setting, a fraction of the population may be at risk. Another fraction of the population may be at risk due to occupational exposure. Toxicological effects of metals to humans, particularly those of Cd, As, Hg, and Pb, have been well documented (see also individual chapters on these elements later in this book). Their ecotoxicological effects on agricultural, forest, and aquatic ecosystems are also well documented.

How clean is clean? How clean is clean enough? These are indications of the changing cleanup paradigm being faced by the scientific community, policymakers, and industry. A decade or so ago policymakers believed that cleanup activities should use background (or baseline) values of contaminants in soils as an index of a successful cleanup (e.g., in the case of soil washing/flushing or excavation). Today, the cost and even the best available remedial technologies may not be able to achieve such rather wishful levels without destroying the integrity of the contaminated site. This has led to the current cleanup questions: How clean is clean enough? In other words, how much risk reduction is acceptable without reaching the background level? Or can standards be tailored to specific land use?

Arsenic (atom. no. 33) is a steel-gray, brittle, crystalline metalloid with three allotropic forms that are yellow, black, and gray. It tarnishes in air and when heated is rapidly oxidized to arsenous oxide (As₂O₃) with the odor of garlic. It belongs to Group V-A of the periodic table, has an atom. wt. of 74.92, and closely resembles phosphorus chemically. Gray As, the ordinary stable form, has a density of 5.73 g cm⁻³, a melting pt. of 817 ºC, and sublimes at 613 ºC. The more common oxidation states for As are —III, 0, III, and V. Arsines and methylarsines, which are —III species, are generally unstable in air. Elemental As (As⁰) is formed by the reduction of arsenic oxides. Arsenic trioxide (As³⁺) is a product of smelting operation and is the raw material used in forming most arsenicals. It is oxidized catalytically or by bacteria to arsenic pentoxide (As⁵⁺) or orthoarsenic acid (H₃AsO₄). Arsenic compounds compete with their phosphorus analogs for chemical binding sites. Arsenic covalently bonds with most nonmetals and metals, and forms stable organic compounds in both its trivalent and pentavalent states. The most important compounds are white As (As₂O₃), the sulfide, Paris green (copper acetoarsenite) [3Cu(AsO₂)₂·Cu(CH₃COO)₂], calcium arsenate, and lead arsenate, the last three being used as pesticides and poisons.

Cadmium is a soft, ductile, silver-white, lustrous, electropositive metal with an atom. wt. of 112.4, density of 8.64 g cm⁻³, and melting pt. of 321 ºC. It has eight stable isotopes with the following percentages of abundance: ¹⁰⁶Cd (1.22%), ¹⁰⁸Cd (0.88%), ¹¹⁰Cd (12.39%), ¹¹¹Cd (12.75%), ¹¹²Cd (24.07%), ¹¹³Cd (12.26%), ¹¹⁴Cd (28.86%), and ¹¹⁶Cd (7.58%). Like Zn and Hg, Cd is a transition metal in Group II-B of the periodic table. Cadmium and Zn however, differ from Hg in that the Hg forms particularly strong Hg—C bonds. Like Zn, Cd is almost always divalent in all stable compounds, and its ion is colorless. Its most common compound in nature is CdS. It forms hydroxides and complex ions with ammonia and cyanide, e.g., Cd(NH₃) ₆ ⁴⁻ and Cd(CN) ₄ ²⁻ .It also forms a variety of complex organic amines, sulfur complexes, chloro-complexes, and chelates. Cadmium ions form insoluble usually hydrated white compounds, with carbonates, arsenates, phosphates, oxalates, and ferrocyanides.

Chromium, a member of Group VI-B of the periodic table has atom. no. 24, atom. wt. of 52.0, specific gravity of 7.2 at 20 °C, melting pt. of 1857 °C, and four stable isotopes with the following percentages of abundance: ⁵⁰Cr (4.31%), ⁵²Cr (83.76%), ⁵³Cr (9.55%), and ⁵⁴Cr (2.38%). It has five radioactive isotopes but only ⁵¹Cr with a half-life of 27.8 days is commonly used for tracer studies. Chromium is a silvery, lustrous, malleable metal that takes a high polish. It dissolves readily in nonoxidizing mineral acids but not in cold aqua regia or HNO₃. Thus, it is resistant to attack by oxidizing acids and a range of other chemicals, hence its use in corrosion-resistant alloys.

Remember the consumers’ slogan “unsafe at any speed” regarding the fatalities associated with what many people considered to be an unsafe automobile design? With Pb, perhaps the public slogan should be “unsafe at any level in the blood” (i.e., above the background)—reminding us of the potential toxicity of this element to humans, especially the unborn.

Mercury, also called liquid silver, has atom. no. 80, atom. wt. of 200.6, melting pt. of −38.8 °C, specific gravity of 13.55, vapor pressure of 1.22 × 10⁻³ mm at 20 °C (2.8 × 10⁻³ mm at 30 °C), and solubility in water of 6 × 10⁻⁶ g/100 ml (25 °C). It is a heavy, glistening, silvery-white metal that is a liquid at room temperature, a rather poor conductor of heat but a fair conductor of electricity. It has seven stable isotopes with the following percentages of abundance: ¹⁹⁵Hg (0. 15%), ¹⁹⁸Hg (10.1%), ¹⁹⁹Hg (17.0%), ²⁰⁰Hg (23.3%), ²⁰¹Hg (13.2%), ²⁰²Hg (29.6%), and ²⁰⁴Hg (6.7%). There are many minerals of Hg; the most common are the sulfides cinnabar and metacinnabar. Mercury is recovered almost entirely from cinnabar (α-HgS, 86.2% Hg); less important sources are livingstonite (HgS · 2Sb₂S₃), metacinnabar (β-HgS), and about 25 other Hg-containing minerals. Its unusual high volatility, which increases with increasing temperature, accounts for its presence in the atmosphere in appreciable amounts.

Boron belongs to Group III-A of the periodic table and is the only nonmetal among plant micronutrients. It has an atom. wt. of 10.81, a melting pt. of 2080 °C, with specific gravity of 2.34 for crystals. Boron has two stable isotopes in nature, ¹⁰B (19.78%) and ¹¹B (80.22%). At room temperature, B is inert except to strong oxidizing agents such as HNO₃. When fused with oxidizing alkaline mixtures such as NaOH and NaNO₃, it forms borates. The only important oxide is boric oxide (B₂O₃), which is acidic, soluble in water, and forms boric acid B(OH)₃, a very weak acid. In nature, B is fairly rare and occurs primarily as the borates of Ca and Na. Borax (Na₂B₄O₇ · 10H₂O) is the most common compound along with boric or boracic acid. By far the most important source of B is the mineral kernite (Na₂B₄O₇ · 4H₂O) (also known as rasorite), an evaporite deposit found in the Mojave desert of California. Boron does not form B³⁺ cations because of its high ionization potential; thus, B complexes mainly involve covalent rather than ionic bonds. Boron chemistry more closely resembles Si than Al chemistry. Boron may be entrapped in the clay lattice by substituting for Al³⁺ and/or Si⁴⁺ ions. There are numerous borosilicate minerals in soils, but only tourmaline and axinite may be of significance. Boric acid has a trigonal structure while borate anion has a tetrahedral structure.

Copper (atom. no. 29), one of the most important metals to society, is reddish, takes on a bright metallic luster, is malleable, ductile, and a good conductor of heat and electricity (second only to silver in electrical conductivity). It belongs to Group I-B of the periodic table, has an atom. wt. of 63.55, a melting pt. of 1083 °C, and a specific gravity of 8.96. It consists of two natural isotopes, ⁶³Cu and ⁶⁵Cu, with relative abundances of 69.1 and 30.9%, respectively. The radioactive isotope ⁶⁴Cu with a half-life of 12.8 hr, is the most suitable for tracer work.

Manganese (atom. no. 25; atom. wt. 54.94; melting pt, 1244 ± 3 °C, and specific gravity, 7.2), is a member of Group VII-A of the periodic table. It is next to Fe in the atomic series, is similar to it in chemical behavior, and is often closely associated with it in its natural occurrence (i.e., Mn ores and Fe ores often coexist). Its compounds can exist in the oxidation states of I, II, III, IV, VI, and VII. Its most stable salts are those of oxidation states II, IV, VI, and VII. The lower oxides (MnO and Mn₂O₃) are basic; the higher oxides are acidic. Manganese is a whitish-gray metal, harder than Fe but quite brittle. Manganese metal oxidizes superficially in air and rusts in moist air.

Molybdenum (atom. no. 42, atom. wt. 95.94) is in the second row of the transition metal elements and occurs as five isotopes. It is in Group VI-B with Cr and W and shares some chemical properties with each of these elements. It is a silvery white metal that is very hard, although softer and more ductile than W. It has a density of 10.22 at 20 °C, and a melting pt. of 2617 °C. Molybdenum has five possible oxidation states (II, III, IV, V, and VI), but in nature the IV and VI oxidation states predominate, with the latter being the most stable. At high oxidation states Mo has an affinity for oxides, and for sulfur- and oxygen-containing groups. The affinity of Mo for oxygen-containing groups is the reason for its predominant presence as dissolved anionic species in aqueous systems. The most important compound is the trioxide (MoO₃) from which most of the known Mo compounds can be prepared. Molybdenum is resistant to HC1, H₂SO₄, H₃PO₄, and HF solutions under many conditions of concentration and temperature. However, the metal is attacked by oxidizing acids and fused alkalis. It is rapidly oxidized in air at >500 °C. At moderate to high concentrations in solution, molybdate readily polymerizes into polymolybdates with a wide variety of highly complex structures. However, in dilute solutions, such as those found in soils or in most natural waters, the predominant form of soluble Mo is the molybdate anion (MoO ₄ ²⁻ ). Only under unusual conditions of very high enrichment will Mo be found as soluble polymolybdate in waters. In nature, the only important ore is molybdenite (MoS₂); however, some powellite and deposits containing wulfenite also occur.

Zinc (atom. no. 30) is a bluish white, relatively soft metal with a density of 7.13 g cm⁻³. It belongs to Group II-B of the periodic table and has an atom. wt. of 65.38 and melting pt. of 420 °C. Zinc is divalent in all its compounds. It is a composite of five stable isotopes: ⁶⁴Zn, ⁶⁶Zn, ⁶⁷Zn, ⁶⁸Zn, and ⁷⁰Zn. Their relative abundances are: 48.89, 27.81, 4.11, 18.56, and 0.62%, respectively. Six radioactive isotopes have been identified: ⁶²Zn, ⁶³Zn, ⁶⁵Zn, ⁶⁹Zn, ⁷²Zn, and ⁷³Zn, with ⁶⁵Zn, (t_1/2 = 245 days) and ⁶⁹Zn (t_1/2 = 55 min) being the most commonly used.

Nickel (atom. no. 28) belongs in Group VIII of the periodic table, the so—called iron—cobalt group of metals, and has an atom. wt. of 58.71, specific gravity of 8.9, and melting pt. of 1453 °C. It is a silvery white, hard, malleable, ductile, ferromagnetic metal that maintains a high luster and is relatively resistant to corrosion. It is insoluble in water, but soluble in dilute HNO₃, slightly soluble in HC1 and H₂SO₄ and insoluble in NH₄OH.

Selenium (atom. no. 34) is a member of Group VI-A, also known as the S family in the periodic table. It has a specific gravity of 4.79 for the metallic (gray) form or 4.28 for the vitreous (black) form. It has an atom. wt. of 78.96 and melting pt. of 217 °C (gray form). Selenium is classified as a metalloid having properties of both a metal and a nonmetal. Because of its chemical similarity to S, it resembles S both in its form and compounds, and this accounts for their many interrelations in biology. It has six stable isotopes in nature with the following composition (in %): ⁷⁴Se, 0.87; ⁷⁶Se, 9.02; ⁷⁷Se, 7.58; ⁷⁸Se, 23.52; ⁸⁰Se, 49.82; and ⁸²Se, 9.19. The most important oxidation states of Se are —II, 0, IV, and VI. Elemental Se (Se⁰) is often associated with S in compounds such as selenium sulfide (Se₂S₂) and polysulfides. Selenite (SeO ₃ ²⁻ ) and selenate (SeO ₄ ²⁻ )are common ions in soils and natural waters and are highly water soluble; Se⁰ is much less soluble in water. Selenium can be easily oxidized from elemental Se [Se(0)], to SeO ₃ ²⁻ [Se(IV)], and to Se O ₄ ²⁻ [Se(VI)]. Some of the better-known commercial Se compounds include H₂Se, metallic selenides [Se(—II)], Se0₂, H₂SeO₃, SeF₄, Se₂Cl₂, and H₂SeO₄ (selenic acid). Selenium also forms a large number of organic compounds that are similar to those of S.

The bluish white, lustrous, very brittle metal (atom. no. 51, atom. wt. 121.75, specific gravity 6.69, melting pt. 631 °C) is found in nature in more than 100 minerals. However, only about one dozen ores are commercially important, such as Sb oxides and sulfides, and complex Cu—, Pb—, and Hg—Sb sulfides, the most important of which is stibnite (Sb₂S₃). It also occurs with sphalerite, pyrite, and galena and is found in minor amounts in Hg deposits. In nature, it occurs primarily in the III oxidation state.