We will now extend electrochemistry by determining the relationship between [latex]{E}_{\text{cell}}^{\circ }[/latex] and the thermodynamics quantities such as ΔG° (Gibbs free energy) and K (the equilibrium constant). In galvanic cells, chemical energy is converted into electrical energy, which can do work. The electrical work is the product of the charge transferred multiplied by the potential difference (voltage):

[latex]\text{electrical work}=\text{volts}\times \left(\text{charge in coulombs}\right)=\text{J}[/latex]

The charge on 1 mole of electrons is given by Faraday’s constant (F)

[latex]F=\frac{6.022\times{10}^{23}\text{e}^{-}}{\text{mol}}\times\frac{1.602\times{10}^{-19}\text{C}}{\text{e}^{-}}=9.648\times{10}^{4}\frac{\text{C}}{\text{mol}}=9.684\times {10}^{4}\frac{\text{J}}{\text{V}\cdot\text{mol}}[/latex]

[latex]\text{total charge}=\left(\text{number of moles of e}^{-}\right)\times F=nF[/latex]

In this equation, n is the number of moles of electrons for the balanced oxidation-reduction reaction. The measured cell potential is the maximum potential the cell can produce and is related to the electrical work (w_ele) by

[latex]{E}_{\text{cell}}=\frac{-{w}_{\text{ele}}}{nF}\,\,\,\,\,\,\,\,\text{or}\,\,\,\,\,\,\,\,{w}_{\text{ele}}=-nF{E}_{\text{cell}}[/latex]

The negative sign for the work indicates that the electrical work is done by the system (the galvanic cell) on the surroundings. In an earlier chapter, the free energy was defined as the energy that was available to do work. In particular, the change in free energy was defined in terms of the maximum work (w_max), which, for electrochemical systems, is w_ele.

[latex]\Delta G={w}_{\text{max}}={w}_{\text{ele}}[/latex]
[latex]\Delta G=-nF{E}_{\text{cell}}[/latex]

We can verify the signs are correct when we realize that n and F are positive constants and that galvanic cells, which have positive cell potentials, involve spontaneous reactions. Thus, spontaneous reactions, which have ΔG < 0, must have E_cell > 0. If all the reactants and products are in their standard states, this becomes

[latex]\Delta G^{\circ }=-nF{E}_{\text{cell}}^{\circ }[/latex]

This provides a way to relate standard cell potentials to equilibrium constants, since

[latex]\Delta G^{\circ }=-RT\ln{K}[/latex]
[latex]-nF{E}_{\text{cell}}^{\circ }=-RT\ln{K}\,\,\,\,\,\,\,\,\text{or}\,\,\,\,\,\,\,\,{E}_{\text{cell}}^{\circ }=\frac{RT}{nF}\ln{K}[/latex]

Most of the time, the electrochemical reactions are run at standard temperature (298.15 K). Collecting terms at this temperature yields

[latex]{E}_{\text{cell}}^{\circ }=\frac{RT}{nF}\ln{K}=\frac{\left(8.314\frac{\text{J}}{\text{K}\cdot \text{mol}}\right)\left(298.15K\right)}{n\times \text{96,485 C/V}\cdot \text{mol}}\ln{K}=\frac{\text{0.0257 V}}{n}\ln{K}[/latex]

where n is the number of moles of electrons. For historical reasons, the logarithm in equations involving cell potentials is often expressed using base 10 logarithms (log), which changes the constant by a factor of 2.303:

[latex]{E}_{\text{cell}}^{\circ }=\frac{0.0\text{592 V}}{n}\log{K}[/latex]

Thus, if ΔG°, K, or [latex]{E}_{\text{cell}}^{\circ }[/latex] is known or can be calculated, the other two quantities can be readily determined. The relationships are shown graphically in Figure 1.

Figure 1. The relationships between ΔG°, K, and Eº_cell. Given any one of the three quantities, the other two can be calculated, so any of the quantities could be used to determine whether a process was spontaneous.

Given any one of the quantities, the other two can be calculated.

Now that the connection has been made between the free energy and cell potentials, nonstandard concentrations follow. Recall that

[latex]\Delta G=\Delta G^{\circ }+RT\ln{Q}[/latex]

where Q is the reaction quotient (see the chapter on equilibrium fundamentals). Converting to cell potentials:

[latex]-nF{E}_{\text{cell}}=-nF{E}_{\text{cell}}^{\circ }+RT\ln{Q}\,\,\,\,\,\,\,\,\text{or}\,\,\,\,\,\,\,\,{E}_{\text{cell}}={E}_{\text{cell}}^{\circ }-\frac{RT}{nF}\ln{Q}[/latex]

This is the Nernst equation. At standard temperature (298.15 K), it is possible to write the above equations as

[latex]{E}_{\text{cell}}={E}_{\text{cell}}^{\circ }-\frac{0.0\text{257 V}}{n}\ln{Q}\,\,\,\,\,\,\,\,\text{or}\,\,\,\,\,\,\,\,{E}_{\text{cell}}={E}_{\text{cell}}^{\circ }-\frac{\text{0.0592 V}}{n}\log{Q}[/latex]

If the temperature is not 273.15 K, it is necessary to recalculate the value of the constant. With the Nernst equation, it is possible to calculate the cell potential at nonstandard conditions. This adjustment is necessary because potentials determined under different conditions will have different values.

Finally, we will take a brief look at a special type of cell called a concentration cell. In a concentration cell, the electrodes are the same material and the half-cells differ only in concentration. Since one or both compartments is not standard, the cell potentials will be unequal; therefore, there will be a potential difference, which can be determined with the aid of the Nernst equation.

Glossary

concentration cell: galvanic cell in which the two half-cells are the same except for the concentration of the solutes; spontaneous when the overall reaction is the dilution of the solute

electrical work (w_ele): negative of total charge times the cell potential; equal to w_max for the system, and so equals the free energy change (ΔG)

Faraday’s constant (F): charge on 1 mol of electrons; F = 96,485 C/mol e⁻

Nernst equation: equation that relates the logarithm of the reaction quotient (Q) to nonstandard cell potentials; can be used to relate equilibrium constants to standard cell potentials