Ways Of Shifting Chemical Equilibria

Environmental engineers deal routinely with materials that are in either homogeneous or heterogeneous equilibrium. They, like analytical chemists, must be able to apply stresses to their systems in accordance with Le Chatelier's principle to bring about desired changes. Many of the stresses that they apply are exactly the same in character as those used by chemists. Therefore, it is important to consider the ways in which equilibria can be shifted to bring about essentially complete reactions. Five methods are commonly employed.

Formation of Insoluble Substances

All precipitation reactions are examples of this method of equilibrium shift. In this case a knowledge of the solubility-product principle, solubility-product constants, and the common ion effect is brought into service. The removal of metal ions from industrial wastes, such as copper and brass wastes, by precipitation with calcium hydroxide, and the softening of hard waters by lime-soda ash treatment are excellent examples of how this method of shifting a chemical equilibrium is applied to gain an objective. Equations associated with these precipitation reactions are listed in Table 2.5.

Formation of a Weakly Ionized Compound

Certain systems that are in equilibrium can be destroyed by adding a reagent that will combine with one of the ions to form a poorly ionized compound. The neutralization of acid and of caustic wastes is based upon such information, since the chapter 2 Basic Concepts from General Chemistry

reaction involved is between hydrogen ions and hydroxyl ions to form poorly ionized water:

Similar reactions are frequently used to bring the pH of industrial wastes into a favorable range for subsequent biological treatment

This method of equilibrium shifting routinely is used to dissolve precipitates of the metallic hydroxides, such as ferric and aluminum hydroxide:

Formation of Complex Ion

Complex-ion formation can be used to dissolve insoluble salts and hydroxides. Silver chloride, for example, dissolves readily in ammonium hydroxide solution. This occurs because silver ion combines with molecular NH3 contained in the ammonium hydroxide to form [Ag(NH3)2]+; as a result the solution becomes unsaturated with respect to silver and chloride ions and solid silver chloride passes into solution in an attempt to form a saturated solution. The net effect is as follows:

If enough ammonium hydroxide is present, all silver chloride passes into solution, The equilibrium relationship for Eq. (2.62) may be written as


A value for the equilibrium constant K may be determined from solubility-product data for silver chloride and from the equilibrium constant for the ammonia-silver complex. Illustrations of such computations are given in Section 4.9.

Zinc and copper hydroxide dissolve in ammonium hydroxide for the same reason given in the preceding paragraph. The complex ions formed are [Zn(NH3)4]2+ and [Cu(NH3)4]2*. These reactions illustrate why ammonium hydroxide would not be a good reagent for precipitating copper and zinc ions from a brass-mill waste.

Industrial wastes containing sodium cyanide are particularly toxic to fish, even though the concentration of cyanide ion may be reduced to very low levels by dilution with river water. The destruction or inactivation of cyanide ion may be accomplished by complex-ion formation. Formerly, before development of more efficient methods, cyanide ion was reduced to low levels by treatment with ferrous sulfate. In the reaction, Fe2+ combined with CN~ to give the complex ion Fe(CN)$~, which could be precipitated as Prussian blue by subsequent oxidation of excess Fe2+ to Fe3*, in the presence of potassium ions,

Formation of a Gaseous Product

In reactions involving the formation of a gaseous product, the reactions go to practical completion because the gas escapes from the sphere of the reaction. This method of forcing a reaction to completion is used when dissolving metallic sulfides, such as ferrous sulfide, in hydrochloric acid:9

This does not necessarily mean that all metallic sulfides are soluble in hydrochloric acid, for the sulfides of copper and mercury are not. This can be explained by considering that copper and mercury sulfides are so insoluble that the sulfide ion liberated by them at equilibrium is of such small concentration that not enough un-ionized hydrogen sulfide (H2S) is formed, in the presence of concentrated hydrochloric acid, to be released as a gas. Consequently, the sulfides do not dissolve.

In industrial waste treatment, cyanides were formerly removed from aqueous solution by treatment with sulfuric acid:

The hydrogen cyanide released as a gas was diluted with large volumes of air and forced up through tall stacks to get dispersion and hopefully to avoid serious atmospheric-pollution problems. More modern methods accomplish destruction of cyanide ion by oxidation or other techniques.

Oxidation and Reduction

A very sure method of sending reactions to completion is oxidation and reduction. In this way one or more of the ions involved in the equilibrium reaction can be destroyed, and the reaction will proceed to completion. A classic example is the destruction of cyanide ion by chlorination according to the following equation:

2CN~ + 5C12 + 80H" 10CF + 2C02 + N2 f + 4H20 (2.67)

From the discussion already presented, this reaction would go to completion for two reasons.

1. Cyanide ion is oxidized to carbon dioxide and nitrogen.

2. The nitrogen escapes as a gas.

Since formation of nitrogen is dependent upon oxidation of cyanide ion, the reaction is considered complete whether the nitrogen escapes or not.

*The arrow represents production and release of gas from solution.

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