14.15-14.17 - Storage And Fuel Cells | Edexcel A-Level Chemistry

Explain how rechargeable cells store energy and fuel cells convert a continuing fuel supply into electrical work. Balance acidic and alkaline hydrogen–oxygen half-equations and compare devices for a stated use.

Separated Redox Makes A Voltage

A flame releases energy because a fuel and oxygen react directly. An electrochemical cell is more controlled: oxidation and reduction are separated so electrons are forced through an external circuit before the redox reaction is complete.

In a cell that is delivering electrical energy:

Process	Where it happens	Electron movement
Oxidation	negative electrode, the anode	electrons are produced
Reduction	positive electrode, the cathode	electrons are accepted

The electron flow through the external circuit is therefore:

negative electrode → external circuit → positive electrode

That direction is not a naming convention to memorise. It follows from the half-equations: the electrode that has electrons on the right-hand side of its half-equation is releasing electrons, so electrons leave that electrode.

Electrode potentials let chemists predict which reduction is more likely under standard conditions. A more positive reduction potential is more likely to run as a reduction; a less positive reduction potential is more likely to be reversed as an oxidation. For a discharging cell:

Cell potential of a discharging cell

E_{\mathrm{cell}} = E_{\text{positive electrode}} - E_{\text{negative electrode}}

Both values are reduction potentials in V; the positive electrode is where reduction occurs.

A positive $E_{\mathrm{cell}}$ means the overall redox reaction is thermodynamically feasible under the conditions being considered. The actual voltage of a working device can be lower because concentration, temperature, internal resistance and rate limitations are not always the same as standard electrode-potential conditions.

Worked example: choosing the direction from electrode potentials

A rechargeable cell contains two reversible half-cell systems:

$\ce{X+(aq) + e- <=>X(s)}$ — $E^{\circ} = +0.80\,\mathrm{V}$

$\ce{Y2+(aq) + 2e- <=> Y(s)}$ — $E^{\circ} = -0.25\,\mathrm{V}$

During discharge, the more positive reduction potential is used as the reduction:

positive electrode: $\ce{X+(aq) + e- -> X(s)}$

The less positive system is reversed to supply electrons:

negative electrode: $\ce{Y(s) -> Y2+(aq) + 2e-}$

To combine the equations, multiply the $\ce{X+}$ half-equation by 2 so the electrons cancel:

$\ce{2X+(aq) + 2e- -> 2X(s)}$

$\ce{Y(s) -> Y2+(aq) + 2e-}$

overall: $\ce{2X+(aq) + Y(s) -> 2X(s) + Y2+(aq)}$

The standard cell potential is:

E^{\circ}_{\mathrm{cell}} = +0.80 - (-0.25) = +1.05\,\mathrm{V}

Chemical interpretation: the cell voltage comes from a difference in tendency to accept electrons. Electrons do not travel through the electrolyte; ions move in the electrolyte to maintain charge balance while electrons travel through the external circuit.

The sign trap

A positive $E_{\mathrm{cell}}$ does not mean "the positive electrode loses electrons". The positive electrode accepts electrons: reduction happens there during discharge.

Storage Cells Use Reversible Redox

A rechargeable cell retains discharge products that can be converted back into reactants.

Storage cell

A rechargeable cell: the reacting chemicals are stored inside, and the discharge reaction can be reversed by applying an external voltage.

The chemical story has two directions:

Direction	Energy change	Redox direction
Discharge	chemical energy is transferred electrically to the circuit	spontaneous redox direction, positive $E_{\mathrm{cell}}$
Charge	electrical energy is supplied to the cell	reverse redox direction is forced

Electrode potentials are useful because they explain the voltage available on discharge and the direction that must be driven during charging. A storage cell is not making energy from nothing; charging stores energy by rebuilding higher-energy reactants from the discharge products.

Lead-acid storage cell as an application

A lead-acid cell is a useful A level example because the half-equations show why both electrodes change during discharge. In a simplified discharge model, the lead electrode is oxidised:

\ce{Pb(s) + SO4^{2-}(aq) -> PbSO4(s) + 2e-}

The lead(IV) oxide electrode is reduced:

\ce{PbO2(s) + 4H+(aq) + SO4^{2-}(aq) + 2e- -> PbSO4(s) + 2H2O(l)}

The electrons produced at the lead electrode travel through the external circuit to the lead(IV) oxide electrode. The sulfuric acid concentration decreases during discharge because $\ce{H+(aq)}$ and $\ce{SO4^{2-}(aq)}$ are used to form water and solid lead(II) sulfate.

When the cell is recharged, an external power supply forces these changes in the reverse direction. Solid $\ce{PbSO4}$ is converted back into $\ce{Pb}$ at one electrode and $\ce{PbO2}$ at the other, and the acid concentration rises again.

The charging supply must oppose the discharge direction and exceed the cell’s opposing emf, with additional voltage needed in practice because of resistance and electrode kinetics. During charging, oxidation and reduction swap electrodes: anode always means oxidation and cathode always means reduction, whereas their positive/negative signs depend on whether the cell is discharging or being driven electrolytically.

Worked example: deducing what happens on discharge

A storage cell has this oxidation half-equation during discharge:

$\ce{M(s) + A^{2-}(aq) -> MA(s) + 2e-}$

State the sign of this electrode and explain the electron flow.

Answer: this is the negative electrode during discharge because oxidation produces electrons. Electrons leave this electrode, travel through the external circuit and enter the positive electrode, where reduction occurs.

Common-error contrast

Tempting statement	Why it fails	Better statement
A rechargeable cell contains electricity	Electrical work during charging changes the chemical substances	A charged cell stores chemical energy
Both electrodes do the same thing in a storage cell	One electrode is oxidised while the other is reduced during discharge	Use the half-equations to identify each electrode
Charging is just a faster discharge	Charging drives the reverse redox reaction	The external supply forces electrons in the opposite direction
Standard $E$ values give the exact car-battery voltage at all times	Real cells are not always at standard conditions and have internal resistance	Electrode potentials predict direction and an ideal voltage tendency

Fuel Cells Use A Continuous Fuel Reaction

A fuel cell receives reactants from an external supply while it delivers electrical energy.

Fuel cell

A cell in which a fuel and oxygen, supplied continuously from outside, react by separated oxidation and reduction to generate a voltage.

A fuel cell is not normally a storage cell. The fuel and oxygen are supplied continuously, and products are removed, so the cell can keep producing a voltage as long as the reactants and operating conditions are maintained.

The essential model is:

fuel is oxidised at the negative electrode
oxygen is reduced at the positive electrode
electrons flow through the external circuit
ions move through the electrolyte to complete the circuit

For a hydrogen-oxygen fuel cell, the overall reaction is:

$\ce{2H2(g) + O2(g) -> 2H2O(l)}$

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Foundation Chemistry (Topics 1-10, shared with AS)

Advanced Chemistry (Topics 11-19)

14.15-14.17 - Storage And Fuel Cells

Separated Redox Makes A Voltage

Storage Cells Use Reversible Redox

Fuel Cells Use A Continuous Fuel Reaction

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