CHEM 001C Lecture Notes - Lecture 16: Nernst Equation, Faraday Constant, Electrochemistry
CHEM 001C Lecture 16: Electrochemistry
●Cell Potential and Free Energy
○Spontaneous cell = ( + ) voltage (discharge of cell or battery)
■Should be ( - ) ΔG
○ ΔG = -nFE or ΔG° = -nFE°
■F = Faraday constant: 96500 C/mol e-
■n = number of moles of electrons
■1 J = 1C * 1V
●Nernst Equation
○ ΔG = ΔG° + RTlnQ
■E = E° - lnQ
nF
RT
●E = E° - logQ (in volts, at 25℃)
n
0.0592 V
●Example
○Given the following information:
■Zn(s) + 2H+
(aq) → Zn2+
(aq) + H2(g)
●Zn(s) → Zn2+
(aq) + 2e- Ecell = +0.76 V
●2H+
(aq) + 2e- → H2(aq) Ecell = 0 V
○If the Zn-H2 cell has 0.10 atm H2, 1.0 M H+, 0.0010 M Zn2+, what is Ecell?
■Q = [1.0]2
[0.0010][0.10]
■Ecell = 0.76 V - [log ]
2 mol e−
0.0592 V
[1.0]2
[0.0010][0.10]
●Ecell = + 0.88 V
○More (+), more spontaneous, reaction shifts to products
●Nernst and pH
○H2(g) → 2H+
(aq) + 2e-Q = [H]
+ 2
[P H2] Assume PH2 = 1 atm
■EH2 = E° - (log )
2 mol e−
0.0592 V[H]
+ 2
[P H2]
■EH2 = - (log )
2 mol e−
0.0592 V[H]
+ 2
[P H2]
■EH2 = (log[H+]2)
2 mol e−
0.0592 V
