Chemical Cells, Batteries and Fuel Cells

(Separate Chemistry only) Chemical cells and fuel cells (4.5.2) are assessed only in the separate AQA GCSE Chemistry course, not in Combined Science: Trilogy.

A cell contains chemicals that react to produce electricity. The reaction transfers energy in the form of a voltage (a potential difference) rather than as heat, which is what makes a cell useful for powering devices.

A simple cell is made from two different metals dipped into an electrolyte (a solution that conducts electricity because it contains ions). The two metals act as electrodes, and connecting them with a wire lets a current flow.

The voltage produced depends on several factors:

• the type of the two electrodes (which metals are used), and
• the type of electrolyte.

The bigger the difference in reactivity between the two metals, the greater the voltage the cell produces. Two metals with a large reactivity gap, such as magnesium and copper, give a higher voltage than two close together.

You may be asked to interpret data on how metal combinations affect the voltage. The pattern to look for is that voltage rises as the reactivity difference between the electrodes increases.

A single cell produces a limited voltage. To get a larger voltage, cells are combined.

A battery is two or more cells connected in series. Connecting cells in series adds their voltages together, so a battery produces a greater voltage than one of its cells alone.

Worked example. Three identical 1.5 V cells are connected in series to make a battery. The total voltage is:

$1.5+1.5+1.5=4.5\text{V}$

This is why a torch that needs 4.5 V uses three 1.5 V cells stacked end to end. In everyday language people call a single unit a "battery", but in chemistry a battery is strictly two or more cells joined in series.

Term	Meaning
Cell	One unit of two electrodes in an electrolyte
Battery	Two or more cells connected in series

Connecting cells in series is what increases the voltage. A battery of several cells gives more voltage than a single cell of the same type.

Cells and batteries fall into two groups depending on whether their reactions can be reversed.

Non-rechargeable cells and batteries produce a voltage until one of the reactants is used up, and then they stop. The reaction cannot be reversed, so the cell must be thrown away. Alkaline batteries are non-rechargeable.

Rechargeable cells and batteries can be recharged because the chemical reactions are reversed when an external electric current is passed through them. This rebuilds the original reactants so the cell can be used again.

Feature	Non-rechargeable	Rechargeable
What happens at the end	Reactant used up, reaction stops	Reaction reversed by external current
Example	Alkaline battery	Lithium-ion, car battery
Reused?	No, discarded	Yes, recharged and reused

A cell stops working when a reactant is used up. A rechargeable cell is different only in that an external current can reverse the reaction and restore the reactants.

You may be asked to evaluate the use of cells for a given situation, weighing cost, lifetime and whether recharging is practical.

A fuel cell is different from an ordinary cell: it is supplied with an external source of fuel (such as hydrogen) and oxygen or air, so it does not run down as long as fuel keeps flowing. The fuel is oxidised electrochemically inside the cell, producing a potential difference.

In a hydrogen fuel cell, the overall reaction is the oxidation of hydrogen to produce water:

$2H_2+O_2\to 2H_{2}O$

Energy is released as a voltage that can drive an electric motor, and the only product is water. Fuel cells offer a potential alternative to rechargeable cells and batteries, for example in vehicles.

Points to use when evaluating hydrogen fuel cells against rechargeable batteries:

• For fuel cells: they run continuously while fuel is supplied, are not thrown away, and their only product is water (no carbon dioxide at the point of use).
• Against fuel cells: hydrogen is a gas that is difficult and expensive to store safely, and producing hydrogen often uses energy from fossil fuels. Batteries can simply be recharged from mains electricity.

(Higher Tier only) Writing the electrode half equations for the hydrogen fuel cell is assessed only at Higher Tier.

The overall reaction is made of two half reactions, one at each electrode. At the negative electrode, hydrogen is oxidised (it loses electrons); at the positive electrode, oxygen is reduced (it gains electrons).

At the negative electrode (oxidation of hydrogen):

$2H_2\to 4H^{+}+4e^{-}$

At the positive electrode (reduction of oxygen):

$O_2+4H^{+}+4e^{-}\to 2H_{2}O$

Checking they combine. Add the two half equations. The 4H⁺ and 4e⁻ appear on both sides and cancel:

$2H_2+O_2\to 2H_{2}O$

This is the overall reaction from the previous slide, which confirms the half equations are balanced for both atoms and charge. The electrons lost by hydrogen are exactly the electrons gained by oxygen, which is why the same number, four, appears in each half equation.

1. Confusing a cell and a battery

A cell is a single unit; a battery is two or more cells in series. When a question asks why a battery gives more voltage than a cell, the answer is that the cell voltages add together in series.

2. Saying voltage depends only on the metals

Voltage depends on the electrodes and the electrolyte. When comparing cells, the largest voltage comes from the pair of metals with the greatest difference in reactivity, but the electrolyte matters too.

3. Getting the fuel cell product wrong

The only product of a hydrogen fuel cell is water. It does not produce carbon dioxide, so do not write CO₂ among the products of a hydrogen fuel cell.

4. Unbalanced fuel cell half equations

Charge as well as atoms must balance. At the negative electrode, $2H_2\to 4H^{+}+4e^{-}$ has zero charge on the left and $4(+)+4(-)=0$ on the right. Check that the electrons in the two half equations match before combining them.

5. Claiming rechargeable batteries never wear out

Rechargeable cells reverse their reactions with an external current, but they do degrade over many cycles. State that recharging reverses the reaction, without claiming an unlimited lifetime.