Encyclopedia of Environmental Science and Engineering, Volume I and II

1268 WATER CHEMISTRY

K

P

 K

1

41 55

2

oH e^12 /{}{}^22 ; log.

c) g MnO 2 (s)  4H^ ^  2e

 Mn^2  2H 2 O(1);

KK





{}

{}{}

;log..

Mn

He

2

42 40 84

For the conditions stipulated the following p values are

obtained:

a) p

Fe

Fe

 



12 53  10 53

3

.log 2

{}

{}

.

b) poH
20 78 1 2. / log(P^122 /{ } )^2 13 11.

c) p

H

Mn

 



20 42 1 2  6 92

4

.log 2

{}

{}

.

Equilibrium Distribution in the Sulphur System

Figure 8 shows the p dependence of a 10^4 M SO 42 

HSsystem at pH  10 and 25C. The reaction is

SO^4 H e HS H O

2

(^98412)
  () (35)
and the redox equilibrium equation is
p
SO H
HS
ε 

(^1818) 
4
2
log log
[][]
[]
K (36)
where log K (for the reduction reaction) is 10^34. Hence,
ppHSOHS
  
425 1125 18  1
4

.. log[^2 ] log[ ]

or, for pH  10,

pSOHS
    

718  18

4

log[^2 ] log[ ].

HS^ ^ is the predominant S(II) species at pH  10. Figure 8

shows that the lines for [SO 42 ] and [HS^ ^ ] intersect at p  7.

The asymptotes for [SO 42 ] have slopes of 8 and 0, whereas

those for [HS^ ^ ] have slopes of 0 and 8.

Lines for the equilibrium partial pressure of O 2 and H 2 are

also given in the diagram. As the diagram shows, rather high

relative electron activities are necessary to reduce SO 42 . At

the pH value selected, the reduction takes place at p values

slightly less negative than for the reduction of water. Thus in

the presence of oxygen and at pH  10, only sulfate can exist;

its reduction is possible only at p values less than 6.

Equilibrium Constants for Redox Reactions

Equilibrium constants for some redox processes pertinent in

aquatic conditions are listed in Table 4. A quite comprehen-

sive reference source for such constants is Stability Constants

for Metal-Ion Complexes, L. G. Sillén and A. E. Martell, The

Chemical Society, London (1964) and its Supplement (1971).

Significantly, the first section of this reference deals with the

electron as a ligand, similarly to its treatment above. Another

compilation, somewhat outdated though still very useful, is

Oxidation Potentials, 2nd ed., W. M. Latimer, Prentice-Hall,

(1952). This treatise lists redox potentials rather than equi-

librium constants, but, as shown in the next section, the latter

are readily obtained from the former.

The Determination of p and Redox Potential

As with pH, p can be measured with a potentiometer using

an indicator electrode (e.g., a platinum or gold electrode)

and a reference electrode. The result is read as a potential

difference in volts. When a reversible hydrogen electrode, at

which the electrode reaction is H 2  2H^ ^  2e, is used as

the reference, the resulting potential difference is termed the

redox potential, E H , where the suffix H refers to the hydro-

gen electrode as the reference. Usually another reference

electrode is used, e.g., a calomel electrode, but the addition

of a constant factor (i.e., the potential difference between

the calomel electrode and the hydrogen electrode) to the

p

–12 –8 –4 0 4 812

–4HS–

–8

–12

–16 HS–

10 –92 10 –76 10 –60 10 –44 10 –28

10 –28

10 –12 10 +4 Po 2

PO 2

PH 2

pH 2

10 +4 10 –4^10 –12^10 –20^10 –36^10 –44

SO 4 –2

SO–2 4

pH=10

log CONC. (M or atm.)

FIGURE 8 Equilibrium distribution of sulfur compounds as a func-

tion of p at pH  10 and 25C. Total concentration is 10^4 M. The

dotted curve shows that solid sulphur cannot exist thermodynami-

cally at pH  10, since its activity never becomes unity. Figure from

Stumm, W. and J. Morgan, Aquatic Chemistry, Wiley-Interscience,

New York, 1970, p. 311.

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