5.4.2 Standard Electrode Potentials

Introduction to Electrode Potentials

In an electrochemical cell, each electrode's potential to either lose or gain electrons is termed its electrode potential. The tendency of a substance to undergo oxidation or reduction can be measured relative to other electrodes using a standard electrode potential ($E^\theta$). A highly reactive metal, which easily loses electrons, will have a negative electrode potential, while a less reactive metal or a reactive non-metal may have a positive potential.

Standard Conditions for Electrode Potentials

To ensure electrode potentials are measured consistently, standard conditions are applied:

• Temperature: 298 K (25°C)
• Pressure: 100 kPa
• Concentration: 1.00 mol dm$^{-3}$ for all ionic species Using standard conditions allows for reliable and comparable values across different cells, ensuring that each measured potential reflects the inherent properties of the electrode materials rather than external conditions like temperature or concentration changes.

Standard Hydrogen Electrode (SHE) as a Reference

Electrode potentials are measured by connecting the electrode to a standard hydrogen electrode (SHE), which has a defined standard potential of 0.00 V.

The SHE includes:

• Hydrogen gas at 100 kPa bubbling over a platinum electrode.
• 1.00 mol dm$^{-3}$ solution of $\text{H}^+$ ions. In measurements, the SHE serves as the anode (negative electrode), meaning oxidation occurs at the hydrogen electrode:

$$\text{H}_2(g) \rightarrow 2\text{H}^+(aq) + 2e^-$$

The electrode connected to the SHE will undergo reduction, and all standard electrode potentials $E^\theta$ are written as reduction reactions.

Electrochemical Series

The electrochemical series is a list of standard electrode potentials arranged from most negative (strongest reducing agents) to most positive (strongest oxidizing agents).

Key patterns in the series include:

• Metals: More reactive metals (those that more readily lose electrons) have more negative $E^\theta$ values.
• Non-metals: More reactive non-metals (those that more readily gain electrons) have more positive $E^\theta$ values. | Half-Cell Reaction | Standard Electrode Potential ($E^\theta$) ($V$) | |---|---| | $\text{F}_2 + 2e^• \rightarrow 2\text{F}^-$ | +2.87 | | $\text{Cl}_2 + 2e^• \rightarrow 2\text{Cl}^-$ | +1.36 | | $\text{Br}_2 + 2e^• \rightarrow 2\text{Br}^-$ | +1.07 | | $\text{Ag}^+ + e^• \rightarrow \text{Ag}$ | +0.80 | | $\text{Fe}^{3+} + e^• \rightarrow \text{Fe}^{2+}$ | +0.77 | | $\text{I}_2 + 2e^• \rightarrow 2\text{I}^-$ | +0.54 | | $\text{Cu}^{2+} + 2e^• \rightarrow \text{Cu}$ | +0.34 | | $\text{H}^+ + e^• \rightarrow \frac{1}{2} \text{H}_2$ | 0.00 | | $\text{Zn}^{2+} + 2e^• \rightarrow \text{Zn}$ | -0.76 | | $\text{Fe}^{2+} + 2e^• \rightarrow \text{Fe}$ | -0.44 | | $\text{Al}^{3+} + 3e^• \rightarrow \text{Al}$ | -1.66 | | $\text{Mg}^{2+} + 2e^• \rightarrow \text{Mg}$ | -2.38 | | $\text{Na}^+ + e^• \rightarrow \text{Na}$ | -2.71 | | $\text{K}^+ + e^• \rightarrow \text{K}$ | -2.93 | | $\text{Li}^+ + e^• \rightarrow \text{Li}$ | -3.04 |

Calculating Cell EMF (Electromotive Force)

The EMF (cell potential) of an electrochemical cell is determined by the difference in standard electrode potentials between the reduction and oxidation half-cells:

$$\text{EMF} = E^\theta_{\text{reduction}} - E^\theta_{\text{oxidation}} = E^\theta_{\text{RHS}} - E^\theta_{\text{LHS}}$$

The EMF is always positive, as the more negative potential is subtracted from the less negative potential.

Practical Use of Simple Cells to Measure Electrode Potentials

You can explore electrode potentials practically by constructing a simple electrochemical cell with two different metals as electrodes, connected by a salt bridge. This setup allows you to measure the voltage generated by the cell and calculate the unknown electrode potential of one metal by comparing it to a reference electrode.