Precipitation

11.1 Principles of Precipitation

Precipitation in a strictly chemical sense is the transition of a substance from the dissolved state to the non-dissolved state upon the addition of other (dissolved) reagents that lead to the formation of precipitates.

However, if chemicals causing precipitation are added to water, other reactions may also take place such as for instance coagulation. Thus, in any practical application of the precipitation process it is often very difficult to distinguish between these reactions. Only on the basis of a detailed knowledge of the composition of the (waste-) water matrix is it possible to describe the direction into which the process advances, i.e., which reaction is favored or which reaction is suppressed.

Precipitation is accomplished by a reaction between a specific metal ion and an anion, for instance:

In surface water, and in the pore water, there is a predominance of the following anions: chloride, sulfate, carbonate, hydrogen carbonate, hydroxide, and under reducing conditions anionic species derived from hydrogen sulfide. The chlorides and sulfates of the common metals are readily soluble, whereas the carbonates, hydroxides and sulphides only dissolve with difficulty.

Hydroxides precipitate in several forms, which may behave quite differently with respect to the effects of co- precipitation or later redissolution. Precipitates may persist in metastable equilibrium with the solution and may slowly convert into the aged forms, thereby becoming more stable and Inactive.

The solubility is highly dependent on pH, as the concentration of the precipitating anions hydroxide, carbonate or sulfide, decreases with decreasing pH due to reaction with the hydrogen ion:

C032" + H+ <=> HC03-S2" + H+ <=> HS-

HC03- + H+ <=> H2C03 = H20 + C02 (g) HS" + H+ <=> H2S (g)

With increasing pH, first carbonates and then hydroxides become the stable phase for many metal Ions. For negative values of the reduction potential, the sulfide remains the stable phase over a wide pH range for many metal ions.

The various interacting processes, which determine the solubility at different pH values can conveniently be illustrated in a graphical double logarithmic representation; see below.

The concentrations of proteolytic species are characterized by the total alkalinity A, and pH. The total alkalinity is determined by adding an excess of a standard acid (e.g., 0.1 M), boiling off the carbon dioxide formed and back titrating to a pH of 6. During this process all the carbonate and hydrogen carbonate are converted to carbon dioxide which is expelled and all the borate is converted to boric acid. The amount of acid used (i.e., the acid added minus the base used for back titration) then corresponds to the alkalinity, At, and the following equation is valid;

where C = the concentration in moles per liter for the indicated species.

In other words the alkalinity is the concentration of hydrogen ions that can be taken up by proteolytic species present in the sample examined. Obviously, the higher the alkalinity, the better the solution is able to maintain a given pH value if acid is added. The buffering capacity and the alkalinity are proportional (see e.g. Stumm and Morgan, 1981).

Each of the proteolytic species in an aquatic system has an equilibrium constant. If we consider the acid HA and the dissociation process:

Al - CH2B03- + 2CC032- + CB03- + (C0H- - CH+) (11.7)

we have:

where Ka = the equilibrium constant.

It is possible, when the composition of the aquatic system is known, to calculate both the alkalinity and the buffering capacity, using the expression for the equilibrium constants. However, these expressions are more conveniently used in logarithmic form. If we consider the expression for Ka for a weak acid, the general expression (11.9), may be used in a logarithmic form:

pH = pKa + log_= pKa + log [A"] - log [HA] (11.10)

multiplying both sides of the equation by -1 and using the symbol p for -log and pH for -log H+.

It is often convenient to plot concentrations of HA and A" versus pH in a logarithmic diagram. If C denotes the total concentrations C=[HA]+[A"], we have at low pH:

This means that log[A~] increases linearly with increasing pH, the slope being +1. The line goes through (log C, pKa) as pH=pKa gives log[A"] = log C. Correspondingly, at high pH, [A"] = C and log [HA] = pKa - pH + logC (11.13)

which implies that log[HA] decreases with increasing pH, the slope being -1. This line also goes through (logC, pKa).

At pH = pKa, [A-] = [HA] = C/2 or log [A-] = log [HA] = log C - 0.3

Table 11.1 arid Fig. 11.1 show the result of these considerations for a single acid-base system.

Note that for H2A the slope will be -2 at pH>pK2, corresponding to the dissociation of 2H+:

H2A = 2H+ + A2" and for A2" the slope will be +2 at pH<pK. This is demonstrated in Fig. 11.2.

13, the buffer capacity, is defined as dC / dpH, where C is the added acid or base in moles of hydrogen or hydroxide ions respectively.

It can hence be shown that:

Figure 11.1. H30+, OH", HA and A" are plotted versus pH for a weak acid with pKa = 4.64 and C = 0.1 M.

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