Page 124 - Soil and water contamination, 2nd edition

Major dissolved phase constituents 111

-
with cations . In natural waters, the most important sulphate complexes are NaSO and
4
0
-
CaSO (aq), and, under acid conditions (pH less than 4), HSO . The strongest complexes
4 4
are formed with divalent or trivalent cations. For aqueous calcium sulphate, the equilibrium
constant is:
[CaSO 4 0 ] 10 . 2 31 (5.16)
[Ca 2 ][SO 2 ]
4
-3
2-
-1
-2
-1
This implies that solutions containing 10 –10 mol l of SO (≈ 100–1000 mg l ) contain
4
considerable amounts of this complex, and so the solubility of CaSO is usually much greater
4
(up to more than three times) than can be expected from the solubility product of gypsum
+
2-
alone. In the absence of Na , the SO concentration in equilibrium with gypsum is about
4
+
-1
1480 mg l and increases with increasing Na concentrations (Hem, 1989).
Under reducing conditions, sulphate is reduced to sulphide (see Equation 4.7). The
reduction of sulphate to sulphide is often mediated by bacteria that use sulphate as an energy
source in anoxic environments. Hydrogen sulphide causes water to have a rotten egg odour;
this is noticeable in waters having only a few tenths of milligram H S per litre in solution.
2
Figure 5.8 shows the fields of dominance of sulphur species at equilibrium as function of
pH and Eh at 25 °C and 1 atmosphere pressure. Note that the redox reactions involving the
sulphur species are generally slow, unless governed by microorganisms, so the solution may
not necessarily be in equilibrium.
Figure 5.8 shows the sulphur species in absence of other constituents. As mentioned
above, metal sulphides are barely soluble, so they tend to precipitate rapidly if metal cations
are present in solution. Because iron is the most common metal in solution under reduced
conditions, the most abundant metal sulphide formed is iron(II) disulphide (pyrite ). In
sediments, pyrite is mostly formed in two steps: first metastable FeS is formed, which is
subsequently transformed to FeS by the overall reaction:
2
FeS + S 0 FeS (5.17)
2
The formation of metal sulphides also governs the solubility of many heavy metals under
reducing conditions (see Chapter 7). As the redox potential increases, the metal sulphides are
reoxidised and dissolved. The overall oxidation reaction of pyrite can be summarised by the
following equation:
FeS + 15 4 / O + 7 2 / H O Fe (OH ) + 4H + + 2SO 2 (5.18)
2 2 2 3 4
This equation illustrates the strong acidification that occurs as pyrite is oxidised. In deep
groundwater, the dissolved oxygen concentration is limited because oxygen cannot be
replenished by exchange with the free atmosphere. Since aerated groundwater in equilibrium
-1
-1
with the atmospheric partial O pressure contains about 0.33 mmol l (= 10.6 mg l )
2
O , the maximum increase in sulphate caused by the complete pyrite oxidation according
2
-1
2-
-1
to Equation (5.18) is 0.18 mmol l (= 17.3 mg l ) SO (Appelo and Postma, 1996).
4
Obviously, the oxygen is also consumed by the decomposition of sediment organic matter , so
these figures are maximum values. The pH usually remains unchanged because of sediment

buffering , i.e. cation exchange and dissolution of carbonates. Therefore, pyrite oxidation
2+
and the subsequent dissolution of calcite often cause an increase in both Ca and HCO 3 -
concentrations (see Equation 5.11). If the pyrite-containing sediments are drained and/
or exposed to air, the pyrite oxidation is not hampered by oxygen limitation. The slow
oxidation of mineral sulphides in these sediments is non-biological until the pH reaches a
value of about 4. Below this pH the Thiobacillus ferrooxidans bacteria are the most active

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