Electrode potentials and the standard hydrogen electrode, electrochemical cells and cell EMF, using standard electrode potentials to predict feasibility, and commercial cells and fuel cells.

What this dot point is asking

AQA wants you to explain electrode potentials and the standard hydrogen electrode, describe electrochemical cells and calculate EMF, use standard electrode potentials to predict reaction feasibility, and describe commercial cells and fuel cells.

Electrode potentials and the standard hydrogen electrode

A more positive $E^\ominus$ means the half-cell is more readily reduced (a better oxidising agent). A more negative $E^\ominus$ means it is more readily oxidised (a better reducing agent).

Electrochemical cells and EMF

Two half-cells joined by a wire (for electrons) and a salt bridge (for ions) form an electrochemical cell.

Predicting feasibility

A reaction is feasible if it gives a positive cell EMF. Combine the two half-equations so that the species with the more positive $E^\ominus$ is reduced and the other is oxidised; a positive overall $E_{cell}$ means the reaction can occur. Remember feasibility says nothing about rate; a slow reaction may still have a positive EMF.

Commercial cells and fuel cells

Commercial cells store chemical energy and release it as electricity. Non-rechargeable cells (such as the alkaline cell) run down when a reactant is used up and cannot be restored. Rechargeable cells (such as the lithium-ion cell in phones and electric cars) reverse the cell reaction when an external voltage is applied during charging. A fuel cell uses a continuous external supply of fuel (commonly hydrogen) and oxygen, converting chemical energy directly into electrical energy with water as the only product, so it does not run down while fuel is supplied.

In a hydrogen-oxygen fuel cell the half-reactions (in acidic conditions) are oxidation of hydrogen at the negative electrode, $\text{H}_2 \rightarrow 2\text{H}^+ + 2\text{e}^-$, and reduction of oxygen at the positive electrode, $\text{O}_2 + 4\text{H}^+ + 4\text{e}^- \rightarrow 2\text{H}_2\text{O}$, giving an overall reaction of $2\text{H}_2 + \text{O}_2 \rightarrow 2\text{H}_2\text{O}$. Fuel cells are more efficient than burning the fuel in an engine and produce only water at the point of use, but AQA expects you to note the drawbacks too: hydrogen is hard to store and transport safely, and the energy used to manufacture the hydrogen (often from hydrocarbons or electrolysis) may itself release carbon dioxide, so the cell is only as clean as its fuel source.

Try this

Q1. State the defined electrode potential of the standard hydrogen electrode. [1 mark]

• Cue. $0.00 \text{ V}$.

Q2. State the condition on cell EMF for a reaction to be feasible. [1 mark]

• Cue. A positive cell EMF.