Info

~i—I—r—|—I I I I I I—I I I I—I I I I—1—|—i—I—r

330 ppm CO

500

450

400

350

2

300

i

250

CM

O

200

150

100

50

0

FIGURE 7.4 Effective Henry's law constant for C02 as a function of the solution pH. Also shown is the corresponding equilibrium total dissolved C02 concentration [COj] for a C02 mixing ratio of 330 ppm.

which can be used to calculate [H + ], Using (7.12), (7.21), and (7.22) for [OH"], [HCOj], and [CO2 ], respectively, we can place (7.27) in the form of an equation with a single unknown, the hydrogen ion concentration

Hco2KclPco2_ ^ 2 Hco2KclKc2Pc0l

or after rearranging:

Even if some C02 dissolves in the water, this amount as we saw will be small and will not change the gas-phase concentration of C02. Given the temperature, which determines the values of Kw, HCq2, Kc\, and Kc2, [H+] can be computed from (7.29), from which all other ion concentrations can be obtained. At an ambient C02 mixing ratio of 350 ppm at 298 K the solution pH is 5.6. This value is often cited as the pH of "pure" rainwater.

7.3.3 Sulfur Dioxide-Water Equilibrium

The behavior of S02 in aqueous solution is qualitatively similar to that of carbon dioxide. Absorption of S02 in water results in

with

"so2 -

(7.33)

Pso2

[h+][hso3-] sl [so2 • h2o]

52 ~ [hso3-]

(7.35)

where Ksl = 1.3 x 10"2M and Ks2 = 6.6 x 10"8M at 298 K. The concentrations of the dissolved species are given by

Ks1[S02-H20} _Hs02KslPs02

Hso2Ks\Ks2pSOl

The corresponding ionic concentrations in equilibrium with 1 ppb of S02 are shown as functions of the pH in Figure 7.5.

Hso3 Temp Rature

FIGURE 7.5 Concentrations of S(IV) species and total S(IV) as a function of solution pH for an S02 mixing ratio of 1 ppb at 298 K.

FIGURE 7.5 Concentrations of S(IV) species and total S(IV) as a function of solution pH for an S02 mixing ratio of 1 ppb at 298 K.

FIGURE 7.6 Effective Henry's law constant for S02 as a function of solution pH at 298 K.

FIGURE 7.6 Effective Henry's law constant for S02 as a function of solution pH at 298 K.

The total dissolved sulfur in solution in oxidation state 4, referred to as S(IV) (see Section 2.2), is equal to

The total dissolved sulfur, S(IV), can be related to the partial pressure of S02 over the solution by or, if we define the effective Henry's law coefficient for S02, //^,rVj, as

the total dissolved sulfur dioxide is given by

The effective Henry's law constant for S02 increases by almost seven orders of magnitude as the pH increases from 1 to 8 (Figure 7.6). The effect of the acid-base equilibria is to "pull" more material into solution than predicted on the basis of Henry's law alone. The Henry's law coefficient for S02 alone, //so,, is 1.23 M atm 1 at 298 K, while for the same

FIGURE 7.7 Equilibrium dissolved S(IV) as a function of pH, gas-phase mixing ratio of S02, and temperature, pH

FIGURE 7.7 Equilibrium dissolved S(IV) as a function of pH, gas-phase mixing ratio of S02, and temperature, temperature, the effective Henry's law coefficient for S{1V), is 16.4 Matm-1 for pH = 3,152Matm 1 for pH = 4, and 1524Matm^1 for pH = 5. Equilibrium S(IV) concentrations for S02 gas-phase mixing ratios of 0.2 to 200 ppb, and over a pH range of 1-6, vary approximately from 0.001 to 1000 pm (Figure 7.7).

S(1V) Concentrations in Different Environments Calculate [S{IV)j as a function of pH for S02 mixing ratios of 0.2 and 200 ppb over the range from pH = 0-6 and for temperatures of 0°C and 25°C.

The concentration of S(IV) is given by (7.40). As we have seen, at 298 K, H&(h = 1.23Matm"1, Ks] = 1.3 x 10"2 M, and ff^ = 6.6 x 10"* M. We can calculate the values of these constants for T = 273 K using (7.5). The required parameters are given in Table 7. A. 1 in Appendix 7 in the end of ihis chapter. We find that at 273 K, tfSO;=3.2Matm Ks] — 2.55 x !0~2 M, and Ka = I0-7 M. Figure 7.7 shows [S(IV)] as a function of pH for these two S02 concentrations at 0°C and 25CC. The concentration of [S(IV)J increases dramatically as pH increases. The concentration of S02 H20 does not depend on the pH, and the abovementioned increase (for constant 7 and £,SQj ) is due exclusively to the increased concentrations of HSOj and S02~. Temperature has a significant effect on the jS(IV)] concentration. For example for pH = 5 and 0.2 ppb S02, the equilibrium concentration increases from 0.16 to 0.82 pM, that is by a factor of ^5 as the temperature decreases from 298 to 273 K.

S(IV) Composition and pH Let us compute the mole fractions of the three S(IV) species as a function of solution pH. The moie fractions are found by combining (7.36H7.38) with (7.40):'

Figure 7.8 shows these three mole fractions as a function of pH at 298 K. At pH values lower than 2, S(IV) is mainly in the form of S02 • H20. At higher pH values the HSO^ fraction increases, and in the pH range from 3 to 6 practically all S(IV) occurs as HSO,. At pH values higher than 7, SO2- dominates. Since these different S(1V) species can be expected to have different chemical reactivities, if a chemical reaction occurs in solution involving either HSOj or SO2-, we expect that the rate of the reaction will depend on pH since the concentration of these species depends on pH.

Figure 7.8 shows these three mole fractions as a function of pH at 298 K. At pH values lower than 2, S(IV) is mainly in the form of S02 • H20. At higher pH values the HSO^ fraction increases, and in the pH range from 3 to 6 practically all S(IV) occurs as HSO,. At pH values higher than 7, SO2- dominates. Since these different S(1V) species can be expected to have different chemical reactivities, if a chemical reaction occurs in solution involving either HSOj or SO2-, we expect that the rate of the reaction will depend on pH since the concentration of these species depends on pH.

1 In aquatic chemistry a common notation for these mole fractions is a«, a,. and a;, rcspcctivcly (Stumm and Morgan J 996).

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