CHEM 001C Lecture Notes - Lecture 17: Gibbs Free Energy, Nernst Equation, Redox
CHEM 001C Lecture 17: Electrochemistry
●Nernst Equation
○Derived from Gibbs free energy under standard conditions
■E°cell = E°red - E°ox
■ΔG° = -nFE°
●n = moles of e-
●F = Faraday’s constant (96500 C/mol*e-)
■ΔG = ΔG° + RTlnQ
●-nFEcell = -nFE°cell + RTlnQ
○At 298 K: E = E° - logQ
n
0.0592 V
■E = E° - lnQ
nF
RT
●Example
○The E°cell for the Zn-Cu redox reaction = +1.10 V
■Zn(s) + Cu2+
(aq) → Zn2+
(aq) + Cu(s)
○What is the equilibrium constant under standard conditions?
■E= E° - logK
n
0.0592 V
■0 = 1.10 - logK
2
0.0592 V
●logK = 0.0592 V
2(1.10 V)
●Keq = 1.58 x 1037
●Example
○What is the standard Gibbs free energy change and the equilibrium constant for
the following reaction under standard conditions?
■Sn(s) + 2Cu2+
(aq) → Sn2+
(aq) + 2Cu+
(aq)
●Sn2+
(aq) + 2e- → Sn(s) Ecell = -0.14 V
●Cu2+
(aq) + e- → Cu+
(aq) Ecell = +0.15 V
■ΔG° = -2(96500)(0.29 V)
●ΔG° = -55970 J
■0 = 0.29 V - logK
2
0.0592 V
