In: Chemistry
Calculate the cell emf for the following reaction at 25°C: 2Ag+(0.010 M) + H2(1 atm) ? 2Ag(s) + 2H+(pH = 8.00)
First, some tabulated data:
Ag+ (aq) + e− ==> Ag(s), Eº = 0.80
2 H+ (aq) + 2 e− ==> H2(g), Eº = 0.00
Since your cell isn't in standard (all solutions 1 M, all gases 1
atm), you must use Nernst's equation to account for this
difference. It states that:
E = Eº - (RT/nF) ln([Red]/[Ox]),
where R and F are the ideal gas' and Faraday's, constants
respectively, n the ammount of electrons involved and [Red] and
[Ox] the concentrations of the oxidated and reduced compounds
involved, respectively. Since no temperature is stated, we'll
supose the standard 25ºC, and rewrite the equation as:
E = Eº - (0.0591/n) log([Red]/[Ox]).
Solids and the solvent have an activity of 1, and thus can be left
out of this equation. In the case of gases, their partial pressure
is used instead of their concentration.
So, for the silver's half-reaction:
Ag+ (aq) + e− ==> Ag(s), Eº = 0.80
And E = Eº - (0.0591/n) log([Red]/[Ox]) becomes:
E = 0.80 - (0.0591/1) log(1/[Ag+])
E = 0.80 + 0.0591 log([Ag+])
For [Ag+] = 0.010 M,
E = 0.80 + 0.0591 log(0.010)
E = 0.80 - 2*0.0591
E = 0.6818
And for hydrogen's:
2 H+ (aq) + 2 e− ==> H2(g), Eº = 0.00
And E = Eº - (0.0591/n) log([Red]/[Ox]) becomes:
E = - (0.0591/2) log((H2)/[H+]^2)
E = - 0.02955 log((H2)/[H+]^2)
E = - 0.02955 log((H2)) + 0.0591 log([H+])
E = - 0.02955 log((H2)) - 0.0591 pH
For (H2) = 1 atm and pH = 8.0,
E = - 0.02955 log(1) - 0.0591*8.0
E = - 0.4728
So, the cell's E is:
Ecell = Ered - Eox
Ecell = 0.6818 + 0.4728
Ecell = 1.1546 V