The biogeochemical cycling of sulfur involves substantial abiotic chemical as well as the biochemical conversions. There are five sulfur oxidation states of appreciable biochemical importance in the environment. With an atomic number of 16, sulfur's outer orbital shell has six electrons. By either depleting or filling this shell, therefore, sulfur may reach extreme oxidation states of -2 (sulfide, S2—) and +6 (sulfate, SO42-) (Figure 13.24), as well as residing at 0 (elemental sulfur, S0), +2 (thiosulfate, S2O32-), and +4 (sulfite, SO32-). (Note: In pyrite, FeS2, the sulfur has an average oxidation state of -1, although this may result from a combination of S2— and S0.)

As with reduced nitrogen in the form of ammonia, reduced sulfide can also be volatile (as hydrogen sulfide, H2S). However, these two gases have opposite responses to pH. Whereas ammonia gas is a weak base and tends to form at higher pH levels (pKa = 9.3), hydrogen sulfide is a weak acid (pKa = 7.2 for first H+, 11.9 for the second). Thus, it tends to be nonionized, and hence volatile, only at neutral or lower pH. Reduced sulfur in alkaline solutions tends to remain ionized, mostly as HS- (or S2- at very high pH). Sulfide also tends to react with metals to form insoluble precipitates. Ferrous sulfide (FeS), in particular, gives many anaerobic sediments and biofilms their black color.

The organic sulfur of the amino acids methionine and cysteine (Table 3.6) is also at an oxidation state of —2 and thus has two bonds with adjacent atoms (Section 3.2). One of these is to a carbon atom, but the other may be to a hydrogen (sulfhydryl bond) or to the sulfur (disulfide bond) of a second amino acid, an important factor in the three-dimensional configurations of proteins.

In comparison to some of the other elements, which are considerably more limiting, sulfur tends to be readily available in most environments. Fresh water usually contains at least 10 mg/L of sulfate, and sometimes much more (especially near the coast), and sewage usually has ^30 mg/L more sulfate than the drinking water from which it is derived. The concentration in seawater is ^2700 mg/L. The limit for sulfate in drinking water is 250 mg/L, but this is a secondary standard (protecting public welfare rather than health) because of taste and a laxative effect.

The reactions of the sulfur cycle are depicted schematically in Figure 13.25. As with the nitrogen cycle, some of the steps are carried out only by prokaryotes. Bacteria and

*- Sulfur Oxidation States

Organic Sulfur

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