Chemistry

Chemistry

Problem

A galvanic cell operates via the half-reaction Cu2++2 e−→Cu(s)(E∘=0.34 V).\text{Cu}^{2+}+2\,e^-\rightarrow\text{Cu}(s) \quad (E^{\circ}=0.34\,\text{V}). At 37 °C the ion concentrations give a reaction quotient Q=0.050Q = 0.050. What is the cell potential at this temperature? (Use R=8.314 J mol−1K−1R = 8.314\,\text{J mol}^{-1}\text{K}^{-1} and F=9.65×104 C mol−1F = 9.65\times10^{4}\,\text{C mol}^{-1}.) a) 0.34 V0.34\,\text{V} b) 0.42 V0.42\,\text{V} c) 0.38 V0.38\,\text{V} d) 0.30 V0.30\,\text{V}
Intuitive explanation

The Nernst equation connects the standard cell potential to real conditions by accounting for how concentration gradients drive the reaction. You'll use the relationship E=E°−RTnFln⁡QE = E° - \frac{RT}{nF}\ln Q to see how the actual cell potential decreases when the product concentration is lower than equilibrium (Q is small), since the reaction still has a strong driving force. Plug in the given values carefully, converting temperature to Kelvin and keeping track of the number of electrons transferred.

Intuitive explanation

The Nernst equation connects the standard cell potential to real conditions by accounting for how concentration gradients drive the reaction. You'll use the relationship E=E°−RTnFln⁡QE = E° - \frac{RT}{nF}\ln Q to see how the actual cell potential decreases when the product concentration is lower than equilibrium (Q is small), since the reaction still has a strong driving force. Plug in the given values carefully, converting temperature to Kelvin and keeping track of the number of electrons transferred.

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