Chemical Equilibria

Redox Equilibria

There are two sources of logK° values:

logK° values from thermodynamic data Gf° 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)
 =  Gf°(H2O) Gf°(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.