Cell Potential (Standard Conditions)

Core Concept

Cell potential (Ecell) is the voltage produced by a voltaic cell due to the spontaneous redox reaction occurring in the cell.

Unit: Volts (V).

Standard Conditions:

  • Temperature: 25∘C (298 K).

  • Pressure: 1 atm for gases.

  • Concentration: 1.0 M for solutions.

  • Ecell°​ determines whether a redox reaction is spontaneous.

  • Use the standard reduction potential table to calculate Ecell°​.

  • A positive Ecell°​ indicates a reaction that can generate electric current, while a negative value means it requires external energy.

Test Yourself

Assorted Multiple Choice
A constant current is passed through an electrolytic cell for 45.0 minutes, delivering a total charge of 8,100 Coulombs. How many moles of electrons were transferred during this process? (Faraday's constant = 96,485 C/mol e⁻)

Podcast Episode

Episode

Cell Potential (Standard Conditions)

Companion Guides

Coming soon

Student Guide Teacher Guide

The Standard Cell Potential ($E^\circ_{\text{cell}}$)

The standard cell potential represents the potential of a galvanic cell under standard conditions (1 M concentration, 1 atm pressure, and 25°C).

It is calculated using the reduction potentials of the half-reactions at the cathode and anode:

$$E^\circ_{\text{cell}} = E^\circ_{\text{cathode}} - E^\circ_{\text{anode}}$$

  • $E^\circ_{\text{cathode}}$: Standard reduction potential of the reduction half-reaction.

  • $E^\circ_{\text{anode}}$: Standard reduction potential of the oxidation half-reaction.

Interpreting Cell Potential

  • $E^\circ_{\text{cell}} > 0$: The reaction is spontaneous, and the cell can generate an electric current.

  • $E^\circ_{\text{cell}} = 0$: The cell is at equilibrium, and no net current flows.

  • $E^\circ_{\text{cell}} < 0$: The reaction is non-spontaneous, and an external voltage is required for it to occur (as in an electrolytic cell).

Significance of Standard Reduction Potentials

The standard reduction potential ($E^\circ$) measures the tendency of a species to gain electrons (be reduced).

  • Higher $E^\circ$: Stronger oxidizing agent (e.g., $\text{F}_2$, $+2.87$ V).

  • Lower $E^\circ$: Stronger reducing agent (e.g., $\text{Li (s)}$, $-3.04$ V).

Steps to Calculate $E^\circ_{\text{cell}}$

  1. Write the Half-Reactions: Identify the oxidation and reduction half-reactions.

  2. Find Standard Reduction Potentials ($E^\circ$): Locate the values in a standard reduction potential table.

  3. Assign Electrode Roles:

    • Cathode: The site of reduction ($E^\circ_{\text{cathode}}$).

    • Anode: The site of oxidation ($E^\circ_{\text{anode}}$).

  4. Apply the Formula:

    $$E^\circ_{\text{cell}} = E^\circ_{\text{cathode}} - E^\circ_{\text{anode}}$$

Example Calculation

Reaction:

$$\text{Zn (s)} + \text{Cu}^{2+} (aq) \rightarrow \text{Zn}^{2+} (aq) + \text{Cu (s)}$$

Steps:

  1. Write Half-Reactions:

    • Oxidation (Anode): $\text{Zn (s)} \rightarrow \text{Zn}^{2+} (aq) + 2e^-$, $E^\circ = -0.76$ V

    • Reduction (Cathode): $\text{Cu}^{2+} (aq) + 2e^- \rightarrow \text{Cu (s)}$, $E^\circ = +0.34$ V

  2. Calculate $E^\circ_{\text{cell}}$:

    $$E^\circ_{\text{cell}} = E^\circ_{\text{cathode}} - E^\circ_{\text{anode}}$$

    $$E^\circ_{\text{cell}} = +0.34 \text{ V} - (-0.76 \text{ V}) = +1.10 \text{ V}$$

Conclusion:

The positive $E^\circ_{\text{cell}}$ indicates a spontaneous reaction.

Video Resources