The value of electrode potential depends on the condition and so will cell potential, because the cell potential is the difference between the electrode potential of the two half-cells.
Electrode potential and cell potential in non-standard condition:
- The electrode potential value depends on the position of the redox equilibrium.
- Factors that affect the position of a redox equilibrium reaction:
  - Concentration
  - Temperature
  - Pressure
- The value of electrode potential varies with the concentration of the aqueous solution, both qualitatively and quantitatively.

Effect of the concentration of an aqueous on electrode potential value qualitatively:
- A redox equilibrium and electrode potential value involving zinc and zinc ions:
- If the concentration of zinc ion is decreased, the electrode potential will be more negative than the E^⦵ value.
  - Le-Chatelier principle: The position of the equilibrium shifts to the direction that resists the changes that disturb the equilibrium of a reaction.
  - When [Zn²⁺] < 1.00 mol dm^-3,
    - More electrons are lost,which means the tendency of zinc metal to undergo oxidation also increases.
    - The electrode potential becomes more negative with a decrease in the concentration of the ion in its aqueous solution.
    - Hence, [Zn²⁺] < 1.00 mol dm^-3 , E < E^⦵
  - When [Zn²⁺] > 1.00 mol dm^-3,
    - This leads to the formation of more zinc as a result of more electrons being gained.
    - The tendency of zinc ions to undergo reduction also increases.
    - The electrode potential becomes more positive with the increase in the concentration of the ions in its aqueous solution.
    - Hence, [Zn²⁺] > 1.00 mol dm^-3 , E > E^⦵
- The electrode potential is directly proportional to the concentration of the oxidised species and inversely proportional to that of the reduced species.
  - For example: The reduction of iron(III) ions to iron(II) ions:, two ionic species are involved.
- The concentration of ions in an aqueous solution also affects the oxidising power of the oxidising agents and the reducing power of the reducing agents.
  - Dichromate six ions are good oxidising agents, and their oxidising power depends on ion concentration in its solution.
  - If, [Cr₂O₇^–] > 1.0 mol dm^-3 or [H⁺] > 1.0 mol/dm³, E > +1.33V
    The equilibrium shifts towards the right.
  - If, [Cr₂O₇^–] < 1.0 mol dm^-3 and [H⁺] < 1.0 mol dm^-3, E < +1.33V
    The equilibrium shifts towards the left.
  - If, [Cr³⁺]>1.0 mol/dm³ E < +1.33V
    The equilibrium shifts towards the left.
  - If, [Cr³⁺]<1.0 mol/dm³ E > +1.33V
    The equilibrium shifts towards the right.
  - Effect of concentration of ions on the oxidising power:
    - Increasing [Cr₂O₇^–] and [H⁺] increases the oxidising power.
    - Decreasing [Cr₂O₇^–] and [H⁺] decreases the oxidising power.
    - Increasing [Cr³⁺] decreases the oxidising power.
    - Decreasing [Cr³⁺] increases the oxidising power.
  - The concentration of hydrogen ions has a direct effect on the electrode potential values so, the more acidic the solution, the stronger the oxidising agent.
  - The pH of the solution has an inverse effect on the electrode potential values and hence, the oxidising power only when hydrogen ions are involved in the reaction and appear on the left side of the equilibrium.
    - If the hydrogen ion appears in the right side of the equilibrium, the pH will have a direct effect.
Effect of the concentration of an aqueous on electrode potential value quantitatively:
- Nernst equation:
- Under standard temperature and pressure:
  - The greater the concentration of the oxidised species, the more positive the electrode potential value
  - The greater the concentration of the reduced species, the less positive the electrode potential value.
- Electrode potential under non-standard concentration:
- Sample problem:
  Calculate the value of electrode potential at 298 K of a Cu(s)/Cu²⁺(aq) electrode that has a concentration of Cu²⁺(aq) ions of 0.01 mol dm^-3. E^⦵= +0.34V.
  The half-cell reaction:
  Cu²⁺(aq) + 2e^–⇌ Cu(s)
  
  The concentration of solid metals does not change. So,
  
  E is less positive than E^⦵ as the [Cu²⁺] lower than 1.00 mol dm^-3.
- When [oxidised species] = [reduced species], E = E^⦵ value
  Because,

Gibbs free energy change:
- It is the energy that is available to do the work.
- The change in Gibbs free energy is used to determine the spontaneity of a reaction.
- Formula:
- In the case of a redox reaction, the change in free energy in standard condition:
- For non-standard condition:
- The Gibbs free energy change depends on the cell potential.
  - If E_cell = +ve, ΔG = -ve, the reaction will be spontaneous.
    - Only spontaneous redox reactions can produce voltages in the electrochemical cells.
  - If E_cell = -ve, ΔG = +ve, the reaction will be non-spontaneous.
    - Non-spontaneous redox reactions cannot produce voltages in the electrochemical cells.
- For example:
  - Since, the E^⦵ value is negative, the ΔG^⦵ value will be positive, and the redox reaction is non-spontaneous.
  - If reversed,
  - Since, the E^⦵ value is positive, the ΔG^⦵ value will be negative, and the redox reaction is spontaneous.

24.05 - Nernst Equation

Learn Step-By-Step Video Explanation For This Topic

Explore our Revision Resources

Mock Exams

Tutor Support

Create an account