The Nernst Equation calculates cell potential under non-standard conditions
Concept #1: The Reaction Quotient is a ratio of product to reactant concentrations at a particular time.
Example #1: What is the reaction quotient for the following redox reaction with the given concentrations?
Pb2+ (aq) + 2 K (s) ⇌ Pb (s) + 2 K+(aq) [Pb2+] = 0.0880 M [K+] = 0.0015 M
Concept #2: The Nernst Equation is used to find the cell potential when ion concentration(s) do not equal 1 M.
Example #2: Calculate the cell potential for a reaction at 25.0ºC when given the following ionic concentrations and standard reduction potentials.
2 Co3+ (aq) + 3 Mg (s) ⇌ 2 Co (s) + 3 Mg2+(aq) [Co3+] = 1.0 M [Mg2+] = 0.0033 M
Standard Reduction Potentials
Co3+ (aq) + 3 e– →. Co (s) E°red = + 1.82 V
Mg2+ (aq) + 2 e– →. Mg (s) E°red = – 2.37 V
Practice: If [Br –] = 0.010 M and [Al3+] = 0.022 M, predict whether the following reaction would proceed spontaneously as written at 25ºC:
Al (s) + Br2 (l) ⇌. Al3+ (aq) + Br – (aq)
Standard Reduction Potentials
Al3+ (aq) + 3 e– → Al (s) E°red = – 1.66 V
Br2 (l) + 2 e – 2 → Br– (aq) E°red = + 1.09 V
Practice: Determine [Fe2+] for the following galvanic cell at 25ºC if given [Sn2+] = 0.072 M, [Fe3+] = 0.0219 M, and [Sn4+] = 0.00345 M.
Sn2+ (aq) + 2 Fe3+ (aq) ⇌. Sn4+ (aq) + 2 Fe2+ (aq) Ecell = + 0.68 V
Standard Reduction Potentials
Sn4+ (aq) + 2 e– →. Sn2+ (aq) E°red = + 0.151 V
Fe3+ (aq) + e– → Fe2+ (aq) E°red = + 0.771 V
