## Carbonic Acid

The most important environmental acid is carbonic acid, which is produced in natural waters by the dissolution of CO2 gas. The dissolution of gaseous CO2 is comparatively slow; the average lifetime of a CO2 molecule in the atmosphere is about seven years. The air-solution transfer process may be catalyzed by the enzyme carbonic anhydrase in waters containing biological materials. When CO2 enters solution, two steps are actually involved:

The species CO2(aq) (molecular CO2 associated with water of hydration) and H2CO3(aq) are not normally distinguished, so that an overall reaction can be written:

where

3 i CO2

Here, the Henry's law constant is in units of moles per liter per atmosphere (mol liter— 1atm—x). At a partial pressure of CO2 in the atmosphere of 3 x 10—4 atm, this gives an equilibrium concentration of H2CO3 of 1.03 x 10—5mol/ liter.

The ionization equilibria, which involve fast reactions, are

[the ionization constant for H2CO3(aq) alone is 3.16 x 10—4; the value in equation (9-23) is smaller because only a part of the dissolved CO2 is in the form of H2CO3] and

We can obtain the following expressions from the equilibrium constants and

We can find the concentration of any species as a function of pH at any value of Ct, or find slopes of the log concentration plots in various regions as illustrated in the general case just presented. Figure 9-9 is a plot of the fraction of CT that exists in the form of each of the species as a function of pH. Note that at the pH values of most natural water systems, almost all the carbonate is in solution as the bicarbonate ion.

The various equilibrium constants on which Figure 9-9 is based are for pure water containing only CO2. In seawater, the much higher ionic strength affects activities, and if concentration expressions are used, the corresponding

FIGURE 9-9 Plot of the carbonate species present in solution in pure water as a function of pH; a is the fraction of the total for each of the species.

stoichiometric equilibrium constants are different (see the list of values in Table 9-2). The net effect is that in seawater all the curves in Figure 9-9 would be shifted by 0.5-1 pH unit to the left, for the same CT. The concentration of species in a C02 solution as a function of pH is illustrated in Figure 9-9 for a fixed total carbonate concentration CT (CT = [H2C03]*+ [HCO" ] + [CO2"]) in pure water.

The pH of a solution containing only C02 as the carbonate source would be given from [H30+] = [HCO"] + 2[C02"] + [OH"] ~ [HCO"] (the concentration of all of the cations, only H3O+ in this case, must be electrically balanced by the anions). This gives a pH of about 5.6 for water in equilibrium with the atmosphere. Natural waters typically are more basic; for example, seawater has a value near 8.1. This is the result of reactions of C02 with

TABLE 9-2

Equilibrium Constants of Some Carbonate Reactions

"Pure" water Seawater

TABLE 9-2

Equilibrium Constants of Some Carbonate Reactions

"Pure" water Seawater

 Reaction 0°C 25°C 0°C 25°C co2 + h2o ^ h2co3 8 x 10"2 2.8 10 2 6.5 x 10"2 3 x 10" 2 h2co; + h2o ^ hco" + h3o 2.6 x 10-7 4.5 > 10 7 7.2 x 10"7 9 x 10" 7 HCO" + H20 ^ C02" + H20 3 x 10"11 4.8 > 10 -11 4 x 10"10 8 x 10" -10 CaC03(calcite) ^ Ca2+ + CO2 4.2 10 -9 6.4 x 10" FIGURE 9-9 Plot of the carbonate species present in solution in pure water as a function of pH; a is the fraction of the total for each of the species. bases from rocks. Indeed, a pH of 8 is remarkably near the equivalence point for HC0— given in Figure 9-10. It is evident from the figure that an isolated sample of seawater is not buffered, since it is near an equivalence point and holds its pH of about 8.1 rather precariously. The buffer capacity is very small. Natural water not only can react with sediments, but it is in equilibrium with atmospheric C02 that will serve to keep the concentration of H2C03 fixed according to the solubility relationships [equations (9-1) and (9-3)]. If PC02 (g) is constant, then [H2C03] and not CT is constant within the time scale permitted by the rate of the gas-solution exchange reaction and mixing. Figure 9-11 gives the log concentration versus pH curves for this case, with H2C03 fixed by a partial pressure of C02 of 3 x 10—4 atm. The pH of a solution of only C02 in water is not much different from the situation of the total concentration CT being constant, but the titration curve is very different, as shown in Figure 9-12. This system at pH 8 is strongly buffered toward added base, much less so toward acid. At this pH, most of the dissolved C02 is present as HC0—, and the amount of this in the oceans exceeds the amount of gaseous C02 in the atmosphere. As noted, the pH of the ocean and most other natural waters is far from that predicted by C02 alone, and additional base must be present. This may be provided by soluble carbonate salts. In addition, equilibria with insoluble carbonates also contribute to establishment of the pH. The most obvious of these carbonates is CaC03, since calcium is a very common element in aqueous systems, although others such as MgC03 are important as well. If CaC03 precipitation controls the carbonate ion concentration through the Ca2+ concentration of the water, it in fact controls the pH and acts as a buffer system for changes in carbon dioxide pressure. Increased [H2C03] Equivalents of base FIGURE 9-10 The titration curve for carbon dioxide. Equivalents of base FIGURE 9-10 The titration curve for carbon dioxide. FIGURE 9-11 Log concentration versus pH for carbonate species in equilibrium with a fixed atmospheric pressure of C02: 3 x 10~4 atm. FIGURE 9-11 Log concentration versus pH for carbonate species in equilibrium with a fixed atmospheric pressure of C02: 3 x 10~4 atm. 1 2 Equivalent fraction of base added FIGURE 9-12 Titration curve for a solution of CO2 in equilibrium with a fixed atmospheric pressure of C02 in freshwater. 1 2 Equivalent fraction of base added FIGURE 9-12 Titration curve for a solution of CO2 in equilibrium with a fixed atmospheric pressure of C02 in freshwater. arising from a higher C02 pressure would first cause the pH to drop as HC03 and H30+ form. This lower pH would also lower [C03— ], since the reaction C02— + H+ ^ HC0— is favored at lower pH values. The carbonate ion in turn is replaced in the solution as CaCO3 dissolves; that is, the extra H2CO3 is titrated by the basic carbonate. The sum of these reaction is The relation of these reactions to the carbon cycle and the increase in atmospheric carbon dioxide is discussed further in Section 10.2. The pH and calcium ion values in many freshwater systems are consistent with CaCO3 acting as a CO2-buffering system. However, there are a number of complications in the quantitative treatment of calcium carbonate equilibria, such as the existence of different crystalline forms (calcite and aragonite) with slightly different solubility constants, and the problem of supersaturation; that is, the solubility product concentrations may be exceeded without precipitation taking place. Surface ocean waters are often supersaturated with calcium carbonate, while deep waters often are undersaturated. There is evidence that supersaturation may be encouraged by organic materials, which hinder the nucleation process that is essential for initiating formation of the solid phase. Supersaturation is a metastable condition. That is, it is not one of thermo-dynamic equilibrium, but the rate at which the system moves toward equilibrium is slow. Variation of carbonate solubility with water C02 content is responsible for various deposits encountered in freshwater systems. Many soil waters have higher than normal concentrations of dissolved CO2, because they absorb this gas as it is formed from reactions of microorganisms in soils at partial pressures exceeding the normal atmospheric value and ground-water often is under pressure. CaC03 is then more soluble (it is also more soluble under pressure, see Section 9.3). When the water comes to the surface, however, the pressure is released, excess CO2 leaves the solution, and CaCO3 can deposit. This result is illustrated by stalactite and stalagmite formation in caves in limestone regions. Carbonic acid in rainwater also contributes to the weathering of buildings and monuments made of carbonate rock; this is discussed further in Section 12.4. Waters containing Ca2+ (or Mg2+) and HCO— are called temporarily hard waters, since heating will drive CO2 from solution and cause the soluble calcium bicarbonate to convert to the insoluble carbonate in the reverse of reaction (9-29). Permanent hardness, caused by other Ca2+ or Mg2+ salts such as sulfate or chloride, is not eliminated by heating. An important property often measured for waters to be used industrially or domestically is alkalinity, which is a measure of water quality. Titration of a water sample with acid to a pH of 8 (the phenolphthalein end point) converts all carbonate to bicarbonate, and the amount of acid required for this is a measure of the carbonate or hydroxide content (carbonate and caustic alkalinity, respectively). Large amounts of carbonate alkalinity, and any caustic alkalinity, are undesirable because such waters are unpalatable and may be corrosive in some uses. Titration to pH 5 (methyl orange end point) converts bicarbonate to C02, and is taken as a measure of bicarbonate content. Caustic alkalinity is present whenever the amount of acid required to reach pH 8 (the phenolphthalein end point) exceeds that needed to go from pH 8 to pH 5 (the methyl orange end point). Waters with excess acidity, rather than alkalinity, are encountered occasionally, usually in polluted systems such as those associated with acid mine drainage (Section 10.6.3). From the preceding discussion of the solubility of CaC03 in the presence of acid, it is clear that an increase in acidity will lead to an increase in solubility of a sparingly soluble salt when the anion is the conjugate base of a weak acid. Heavy metal carbonates (e.g., PbC03) are examples of compounds for which the solubility increase with decreasing pH can be of environmental concern. Hydroxides and oxides are other common materials that will be affected similarly. It does not follow that increasing pH will not be of concern. Ampho-teric substances may exhibit increased solubility in both acidic or basic conditions, and in many cases there is a particular pH that gives a minimum solubility. This is illustrated by alumina as discussed in Section 12.3.1. Figures 9-13 and 9-14 illustrate the solubility behavior of calcium carbonate
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