Sulfate in Water
Sulfur is not a major constituent of the Earth’s outer crust, but is widely distributed in reduced form, both in igneous and sedimentary rocks, as metallic sulfides.
In weathering or contact with aerated water, the sulfides are oxidized to yield sulfate ions which are carried off in the water. Sulfate is carried to bodies of water by rain and through solution of sulfate compounds in sedimentary geologic formations within the drainage basin.
The most frequently encountered forms of sulfur in fresh waters are as the anion sulfate (SO4-2) in combination with the common cations and as hydrogen sulfide (H2S).
The thresholds of taste and smell were reported to be 0.2mg/L of sulfides in pulp mill wastes.
Sulfate is not highly soluble. Surface waters are generally low in sulfate, except in regions locally rich in iron or in closed lake basins where evaporation raises the concentration.
Pyrite crystals often occur in sedimentary rock and constitute a source of both ferrous iron and sulfate in ground water. Pyrite is commonly associated with biogenic coal deposits that were deposited under strongly reducing conditions. When these deposits are exposed to oxygen, hydrogen ions are produced in considerable quantity in the oxidation of sulfide to sulfate.
Sulfate occurs in certain igneous rock minerals of the fieldspathoid group, but the most extensive and important occurrences are in the evaporite sediments. Calcium sulfate as gypsum (CASO4.2H20) or as anhydrite, which contains no water or crystallization, make up a considerable part of many evaporite rock sequences. Barium and strontium sulfates are less soluble than calcium sulfate, but are relatively rare. Sodium sulfate occurs in some closed basins.
In regions where the rock was initially well supplied with sulfides, as most shales and fine grained sediments are when freshly raised above sea level. The natural processes of weathering bring about oxidation of the surface down to or below the water table and the sulfate produced is available for transport away from the area.
In humid regions, the upper layers of soil and rock are kept thoroughly leached and as fast as they are formed. The soluble products of sulfur are removed from the area in solution because the amount of water available is large in proportion to the supply of sulfate ions.
In semi-arid and arid regions on the other hand, the soils are usually not fully leached and surplus sulfates may accumulate near the surface.
The amount of drainage water that leaves such an area is a small fraction of the total received in precipitation. Because of these factors, the supply of solutes is relatively large in proportion to the water volume in which it can be carried away. As a result, the surface and the underground waters in semi-arid regions tend to be comparatively high in dissolved sulfate. Sulfate is the predominant anion in many of these areas.
One of the most interesting aspects of sulfate circulation is the occurrence of sulfate ion in rainfall and other forms of precipitation. Concentrations frequently exceed 1 mg/L and generally are considerably greater than chloride concentration except in rain falling on or near the ocean. The sulfate in rainfall has been attributed to a number of factors.
It reaches the atmosphere through emission of H2S from shallow ocean water near the continental margins. The H2S that reaches the atmosphere is ultimately oxidized to sulfur dioxide and then to sulfate.
The effect on air pollution, especially the contribution from combustion of fossil fuels, is noticeable in many places. Up to about 30% of the sulfate in rainfall is attributed to this source.
Among the land based factors which may be significant, are the sulfur emitted by volcanoes, fumaroles and springs, and the solution of dust particles.
Sulfate is ecologically important in natural waters in several ways.
It is necessary for plant growth. Short supplies of the material can inhibit the development of phytoplankton populations and therefore reduce primary production. Sulfur is important in protein metabolism and is supplied to the organisms originally as sulfate.
Under anaerobic conditions, sulfate is utilized in chemosynthetic processes of sulfur bacteria. The sulfate is reduced to hydrogen sulfide. This reduction may result in the liberation of other nutrients held in ferric complexes such as phosphates.
Seasonal pulses in sulfate concentration have been demonstrated in certain lakes. In some, sulfate reaches a high in the spring, followed by decreasing concentrations to a low in autumn.
Bicarbonate variation in the same lakes behaves in a reverse fashion. The sulfate decrease is due to reduction to sulfide that is taken up in the bottom lake muds. The bicarbonate fluctuation apparently results partly from carbon dioxide removal by photosynthesis and partly from sulfuric acid activity in the winter oxidation of sulfide.
When an area of low rainfall and accumulated solutes is reclaimed by irrigation, the increased water supply tends to leach away the solutes and they appear in drainage water or return flow. The process is an acceleration of natural leaching and will increase dissolved solids concentrations.
Even where the soil is relatively free from soluble salts, the sulfate concentrations and chloride concentrations in the residual water draining from an irrigated area are generally much higher than they were in the original water supply.
Owing to the unpleasant taste and odor which result when sulfides occur in water, it is unlikely that any person or animal will consume a harmful dose.
The toxicity of solutions of sulfides towards fish increases as the pH value is lowered. The H2S or HS- (anion) rather than the sulfide ion appears to be the toxicity principle. Inorganic sulfides have proved fatal to sensitive fishes such as trout at concentrations between .5 and 1 mg/L as sulfide, even in neutral and somewhat alkaline solutions.
In general, aquatic biota seems to thrive best in water containing less than 50 mg/L of sulfate ion. The toxicity of sulfuric acid towards aquatic life is a function of the resulting pH. In other words, a dose that would be lethal in distilled water or soft water might be harmless in highly buffered water. Strong mineral acids such as sulfuric acid can be directly lethal to fully developed fish in most natural waters only when the pH is reduced to 5.0 or lower.
For example, sulfuric acid must be diluted so as to give a pH of 4.5 or higher in order for fish to survive, and a pH of 5.5 to make it possible-for other aquatic organisms to thrive and provide food for fish. When sulfuric acid in streams depresses the pH below 5.0, specialized flora and fauna begin to develop.
In waters that support good game fish, 5% of the waters contain less than 11 mg/L of sulfates, 50% less than 32 mg/L of sulfate and 95% less than 90 mg/L of sulfate. Experiments indicate that water containing less than .5 mg/L of sulfate will not support the growth of algae.