Redox Equilibria
There are two sources of logK° values:
- Thermodynamic data, that is, Gibbs free energies of formation
(ΔGf°)
- Standard redox potentials
logK° values from thermodynamic data
(ΔGf° values):
See the topic “sources of thermodynamic data”
for a general description.
From the definition of the standard hydrogen electrode:
½ H2(g)
H+ + e−
logK° = 0 and
E° = 0.
it follows that ΔGf°(e−) = 0, because
ΔGf°(H+) = 0 and
ΔGf°(H2(g)) = 0
Example: SO42−
+ 9 H+
+ 8 e−
HS−
+ 4 H2O
| |
ΔGr° |
= |
−RT lnK° |
| |
= |
∑ νi
ΔGf°(i) |
| |
= |
4ΔGf°(H2O)
+ΔGf°(HS−)
−9ΔGf°(H+)
−8ΔGf°(e−)
−ΔGf°(SO42−)
|
Remember that ΔGf°(H+)
= zero and ΔGf°(e−)
= zero at all temperatures. This gives:
| |
ΔGr° |
= |
4(−237.14) +12.24 −(−744.00) |
| |
= |
−192.32 kJ/mol |
| logK° |
= |
ΔGr°/(−RT ln(10)) |
| |
= |
33.7 |
logK° and Standard Redox Potentials
(E°):
The following equations are used:
ΔGr°
= − ne F E°
= −RT lnK°
logK° = E° ne F
/ (RT ln(10))
where F and R
are the Faraday and gas constants; T is the temperature in Kelvin;
ne is the number of electrons in the reaction;
and E° is the potential of a cell reaction involving the
standard hydrogen electrode as reference. This means that that
E° values correspond to reduction reactions.
At 25°C one has (2.303RT / F) = 0.05916 V.
For the example given above (sulfate reduction to sulfide) E°
= +0.25 V at 25°C, which corresponds to
logK° = 34.
