Standard Reduction Potentials Reference Cheat Sheet

A printable reference covering standard reduction potentials, cell voltage, spontaneity, Gibbs free energy, and the Nernst equation for grades 11-12.

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Standard reduction potentials help students predict how electrons move in electrochemical cells. This reference explains how to read a reduction potential table, identify the cathode and anode, and calculate standard cell voltage. It is useful for solving galvanic cell problems, ranking oxidizing agents, and connecting voltage to spontaneity. The core idea is that each half-reaction is written as a reduction and assigned an $E^\circ$ value under standard conditions. The more positive reduction potential occurs at the cathode, while the half-reaction with the lower reduction potential is reversed for oxidation at the anode. Important formulas include $E^\circ_{\text{cell}} = E^\circ_{\text{cathode}} - E^\circ_{\text{anode}}$ , $\Delta G^\circ = -nFE^\circ_{\text{cell}}$ , and the Nernst equation for nonstandard conditions.

Key Facts

Standard reduction potentials, $E^\circ$ , are measured in volts under standard conditions of $1\,\text{M}$ solutions, $1\,\text{atm}$ gases, and $25^\circ\text{C}$ .
The standard hydrogen electrode is the reference half-cell and is assigned $E^\circ = 0.00\,\text{V}$ for $2\text{H}^+ + 2e^- \rightarrow \text{H}_2$ .
For a galvanic cell, the half-reaction with the more positive $E^\circ$ is reduced at the cathode.
When using a standard reduction potential table, calculate cell voltage with $E^\circ_{\text{cell}} = E^\circ_{\text{cathode}} - E^\circ_{\text{anode}}$ .
If a reduction half-reaction is reversed to show oxidation, the sign of its $E^\circ$ value changes.
Do not multiply $E^\circ$ values by coefficients because voltage is an intensive property.
A reaction is spontaneous under standard conditions when $E^\circ_{\text{cell}} > 0$ and $\Delta G^\circ < 0$ .
Free energy and voltage are related by $\Delta G^\circ = -nFE^\circ_{\text{cell}}$ , where $F = 96485\,\text{C}\,\text{mol}^{-1}$ .

Vocabulary

Standard reduction potential: The voltage of a reduction half-reaction measured against the standard hydrogen electrode under standard conditions.
Cathode: The electrode where reduction occurs and electrons are gained by a chemical species.
Anode: The electrode where oxidation occurs and electrons are lost by a chemical species.
Oxidizing agent: A substance that causes another substance to be oxidized by accepting electrons and being reduced.
Reducing agent: A substance that causes another substance to be reduced by donating electrons and being oxidized.
Nernst equation: An equation that adjusts cell voltage for nonstandard concentrations using $E = E^\circ - \frac{0.0592}{n}\log Q$ at $25^\circ\text{C}$ .

Common Mistakes to Avoid

Adding reduction potentials instead of using $E^\circ_{\text{cell}} = E^\circ_{\text{cathode}} - E^\circ_{\text{anode}}$ is wrong because the table lists both half-reactions as reductions.
Multiplying $E^\circ$ by reaction coefficients is wrong because voltage does not depend on the amount of substance reacting.
Forgetting to reverse the sign when a half-reaction is written as oxidation is wrong because the listed value applies only to the reduction direction.
Choosing the anode as the half-reaction with the larger $E^\circ$ is wrong in a galvanic cell because the larger reduction potential is reduced at the cathode.
Using the Nernst equation without balancing electrons is wrong because $n$ must equal the number of electrons transferred in the balanced redox reaction.

Practice Questions

1 Given $E^\circ(\text{Cu}^{2+} + 2e^- \rightarrow \text{Cu}) = +0.34\,\text{V}$ and $E^\circ(\text{Zn}^{2+} + 2e^- \rightarrow \text{Zn}) = -0.76\,\text{V}$ , calculate $E^\circ_{\text{cell}}$ for a zinc-copper galvanic cell.
2 Given $E^\circ(\text{Ag}^+ + e^- \rightarrow \text{Ag}) = +0.80\,\text{V}$ and $E^\circ(\text{Fe}^{2+} + 2e^- \rightarrow \text{Fe}) = -0.44\,\text{V}$ , identify the cathode and calculate $E^\circ_{\text{cell}}$ .
3 For a cell with $E^\circ_{\text{cell}} = +1.10\,\text{V}$ and $n = 2$ , calculate $\Delta G^\circ$ using $\Delta G^\circ = -nFE^\circ_{\text{cell}}$ and $F = 96485\,\text{C}\,\text{mol}^{-1}$ .
4 Explain why $\text{Ag}^+$ is a stronger oxidizing agent than $\text{Zn}^{2+}$ when $E^\circ(\text{Ag}^+ / \text{Ag}) = +0.80\,\text{V}$ and $E^\circ(\text{Zn}^{2+} / \text{Zn}) = -0.76\,\text{V}$ .

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