Electrochemistry

Information should be enough.

The standard reduction potentials are measured under standard conditions: concentrations of all species are 1M, pressure of all gases are 1atm, with temperature at 298K.

For reaction mixtures with "non-standard" concentrations, one can use Nernst equation to calculate the reduction potential for a half-equation of reduction:

E = E0 - (RT/nF) ln ( [reduced form]^x / [oxidized form]^y )
that
E = new reduction potential
E0 = standard reduction potential
R = universal gas constant
T = temperature (in K)
n = number of electrons involved in that half equation
F = Faraday constant = 96485 Coulomb
= charge of 1 mole of electrons
= Avogadro's constant x elementary charge
ln(x) is the natural-log of x
[reduced form] is the concentration of the species on the reduced side of the half equation, and vice versa for [oxidized form]
(actually activity should be used, but let's assume activity = concentration for simplicity)
x, y = coefficients of the species in the half-equation

take the following equation as example
O2 + 4H(+) + 4e- ------> 2H2O
the new E = E0 - ((8.314 x 298) / (4 x 96485)) ln ( 1 / P(O2) [H(+)]^4 )

E0 = +1.23
number of electrons transferred in the half equation is 4, so n = 4
concentration of H2O is constant and can be ignored
pressure of O2 is kept at 1atm
coefficient for proton is 4, so the concentration term of proton has power of 4

pH = 3.60
[H(+)] = 10^(-3.60) = 2.512E-4 M

E = +1.23 - ((8.314 x 298) / (4 x 96485)) ln (1 / (2.512E-4)^4 )
= +1.23 - 0.213
= +1.02 V

for the overall cell reaction, E = (+1.02) - (+1.07)
= -0.05 V

With negative cell potential, probably the reaction is not spontaneous.

But well, while the reaction is energetically unfeasible, no information is given on its kinetic properties, so...

Can you also help with another question of mine? THX!

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