Iron in Water

Iron

The element iron is abundant and widespread constituent of rocks and soils. Concentrations of only a few tenths of a mg/L of iron in water can make it unsuitable for some uses. For this reason, a determination of iron is frequently included in water analysis. Information on iron contents of water is voluminous even though the amounts present in most waters are small.

In igneous rocks, the principle minerals containing iron as an essential component include the pyroxenes, amphiboles, biotite, magnetite and olivine. Most commonly the iron in igneous rocks is in the ferrous form, but may be mixed with ferric iron as in magnetite.
In sediments, iron occurs in the ferrous form in such species as the polysulfides pyrite, or marcasite, the carbonate siderite, and in the mixed oxide magnetite which also is an igneous or metamorphic mineral.

Ferric iron is commonly mixed with ferrous in glauconite and ferrous and ferric iron are widespread minor components of most sediments. The ferric oxides and hydroxides are very important iron bearing minerals. The hydrous ferric oxide content of rocks and soils is commonly responsible for their red or yellow color.

The solution of iron from silicate minerals is a slow process normally, but surface weathering of iron bearing silicates may produce an accumulation of ferric oxide or hydroxide. The oxide and sulfide species of iron minerals are usually the principle sources from which the dissolved iron of ground water is derived.

Iron is an essential element in both plant and animal metabolism. Iron therefore is to be expected in organic wastes and in plant debris in soils. The activities in the biosphere may have a strong influence on the occurrence of iron in water.

Iron is of particular interest because of its importance as a vital element in many chemical reactions in water. Some of the various aspects of iron in relation to oxidation-reduction systems and its reactions with carbonate and phosphate compounds have been discussed. Whether or not these reactions enter into prominence in the particular bodies of water depends upon the presence of compounds other than iron and the form in which the iron occurs.

Iron is found widely in nature, usually as either bivalent (Fe+2) or trivalent (Fe+3). The bivalent ferrous state is soluble, but only under anaerobic conditions. In the presence of oxygen, the trivalent ferric form is present as a colloidal complex in combination with other inorganic ions and simple decomposition products. With oxygen depletion, as for example, in the hypolimneon of a lake in the summer, the ferric form is reduced to ferrous, the latter going into solution. As a result of the breakdown of the ferric complex, the concentration of silicate, phosphate, bicarbonate or iron is often increased depending, of course, on the original chemical nature of the water.

The mere absence of oxygen will not bring about the transformation from ferric to ferrous iron. The depletion of oxygen is a result of organic decomposition that also forms organic compounds that reduce the ferric iron. In some instances, that of ferrous bicarbonate for example, high carbon dioxide content in near neutral or acid conditions is necessary in addition to absence of oxygen and presence of reducing substances of organic origin to reduce iron.

The vertical distribution of iron in lakes is primarily a composite picture of several forms of the element influenced by the solubility factors just considered. In view of the chemical conditions prevailing in the epilimnetic regions of most lakes, the iron content is quite low, usually less than 0.2 mg/L. In the deeper zones, the iron concentration is a function of oxygen content and redox potential. Where the hypolimneon is enriched with oxygen, as in oligotrophic lakes, the iron is normally contained as the nonsoluble ferric complex.

If oxygen becomes deficient at the mud-water interface, ferrous iron may go into solution. In eutrophic lakes, the hypolimneon typically contains iron in solution during the later stages of summer stagnation. Even during these stages the content is normally low, generally only a few mg/L.

One of the forms of iron which appears to be most readily available to phytoplankton is ferric hydroxide Fe(0H)3. This form is the usual product of oxidation of ferrous iron in waters containing dissolved ferrous salts. It normally occurs as a mass of a ferruginous organic complex. In the utilization of the organic component of the complex by bacteria, ferric hydroxide is precipitated as a byproduct.

Iron occurs naturally in streams although in relatively small proportions. In unpolluted flowing waters, this ion usually takes the ferric form because of continual aeration and the presence of oxygen.

Where organic decomposition is great, oxygen depletion may result in the transformation to the ferrous state and the precipitation of iron hydroxide. Similarly, in stagnant stream pools, particularly following subsidence of flood conditions, iron bacteria may grow rapidly and form masses of ferruginous substances in the pools.

Human physiological considerations for iron in trace amounts are essential for nutrition. The daily nutritional requirement is 1 to 2 mg and most diets contain 7 to 35 mg per day with an average of 16 mg/day.

Consequently, drinking water containing iron in unpalatable, unaesthetic concentrations, say of 1 mg/L, would have little effect on the daily total iron intake. Instead of physiological reasons, therefore, the limit for iron concentration is based on aesthetic and taste considerations.

The taste threshold of iron in water has been given as 0.1 and 0.2 mg/L of iron from ferrous sulfate and ferrous chloride, respectively.

In relation to fish and other aquatic life, as was already stated, when iron is added to water in the form of chloride, sulfate or nitrate, the salt disassociates, but the resulting ferrous or ferric irons combine with hydroxyl ions to form precipitates. Hence very little of the iron remains in solution, but if the dosage is sufficient and the water is not strongly buffered, the addition of a soluble iron salt may lower the pH of the water to a toxic level.

Furthermore, the deposition of iron hydroxides on the gills of fish may cause an irritation and a blocking of the respiratory channels. Finally, heavy precipitates of ferric hydroxide may smother fish eggs.

The toxicity of iron in iron salts depends on whether the iron is present in the ferrous or ferric state and whether it is in solution or suspension. The following limiting concentrations have been noted.
Waters that support good fish fauna in U.S. have the following concentrations of iron.

Percent of Waters Having This
Concentration of Iron in Mg/L Concentration or Less
0.0 5%
0.3 50%
0.7 95%

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