Multiple Choice
Calculate the standard cell potential for the following reaction: Mg (s) + Cu²⁺ (aq) → Cu (s) + Mg²⁺ (aq). Given the standard reduction potentials: Cu²⁺ (aq) + 2 e⁻ → Cu (s) is 0.34 V and Mg²⁺ (aq) + 2 e⁻ → Mg (s) is -2.37 V.
20. Electrochemistry
Cell Potential: Standard
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Calculate the equilibrium constant (K) for an oxidation–reduction reaction (n = 2) with a standard cell potential (E°) of -0.51 V.
A
K = 1.2 x 10^-17
B
K = 4.8 x 10^-10
C
K = 2.1 x 10^3
D
K = 3.5 x 10^5
Verified step by step guidance1
Identify the relationship between the standard cell potential (E°) and the equilibrium constant (K) using the Nernst equation at standard conditions: E° = (RT/nF) * ln(K), where R is the gas constant, T is the temperature in Kelvin, n is the number of moles of electrons transferred, and F is Faraday's constant.
Since the problem involves standard conditions, we can simplify the Nernst equation to: E° = (0.0592/n) * log(K) at 25°C (298 K).
Substitute the given values into the simplified equation: E° = -0.51 V and n = 2.
Rearrange the equation to solve for log(K): log(K) = (n * E°) / 0.0592.
Calculate the value of log(K) and then find K by taking the antilogarithm (10^log(K)) to determine the equilibrium constant.
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