5.2.2a Fuel cells
A fuel cell makes electricity from a fuel supplied from outside
Definition
Potential difference
The difference in electrical energy transferred per unit charge between two points in a circuit.
- A fuel cell changes chemical energy into electrical energy while fuel and oxygen are fed in from outside the cell.
- The oxygen can be supplied as pure oxygen or as part of the air.
- The fuel is oxidised electrochemically inside the cell.
- This controlled reaction sets up a potential difference across the terminals of the cell.
- An electric current flows once the terminals are joined in a complete circuit.
- The cell keeps working for as long as it is supplied with fuel and oxygen.
Key Idea
- A fuel cell never needs recharging because its reactants are fed in continuously from outside.
- It stops producing a potential difference if either the fuel or the oxygen supply stops.
In a hydrogen fuel cell, hydrogen is oxidised to form water
- A hydrogen fuel cell is fed with hydrogen as the fuel and oxygen, usually from the air.
- Hydrogen is oxidised when it reacts with the oxygen.
- The overall reaction is 2H2(g)+O2(g)→2H2O(l)2\text{H}_2\text{(g)} + \text{O}_2\text{(g)} \rightarrow 2\text{H}_2\text{O(l)}2H2(g)+O2(g)→2H2O(l).
- Water is the only chemical product of the overall reaction.
- Hydrogen contains no carbon, so the cell produces no carbon dioxide while it is running.
Common Mistake
- Do not claim that hydrogen fuel cells cause no environmental damage at all.
- The cell makes only water in use, but making and transporting the hydrogen can use energy and release pollutants.
A fuel cell is refuelled, while a rechargeable battery is recharged
Definition
Rechargeable cell
A cell that can be restored by passing an electric current through it so that it can be used again.
- Keeping a device running: a fuel cell runs while fuel and oxygen are supplied, whereas a rechargeable battery holds a fixed amount of chemicals and eventually goes flat.
- Returning it to use: a fuel cell is refilled with hydrogen, while a rechargeable battery must be connected to an electricity supply to charge.
- Time: refilling a hydrogen tank can be quicker than charging a large battery, though charging time depends on the battery and charger.
- Equipment: a fuel cell needs a tank, safe pipes and a hydrogen supply, while a battery needs a charger and access to mains electricity.
Example
- A hydrogen bus can be refuelled quickly and sent straight back into service.
- A battery bus may take longer to charge, but it does not carry a tank of flammable hydrogen.
Neither option is impact-free once you include making and storing the fuel
Definition
Battery
Two or more cells connected together to provide electrical energy.
- Emissions in use: a hydrogen fuel cell gives out only water, and a rechargeable battery gives out no exhaust gases either.
- Making the hydrogen or electricity: producing hydrogen can release carbon dioxide if fossil fuels are used, and charging a battery is only as clean as the electricity used.
- Storage and safety: hydrogen is highly flammable and is often stored under high pressure, needing strong tanks, while batteries avoid this but can be heavy.
- Materials and waste: batteries need extracted raw materials and lose capacity over time, and fuel cell systems also need materials, manufacture and replacement.
- Cost and infrastructure: hydrogen production and refuelling stations can be expensive, whereas batteries can use the existing electricity grid.
Note
- Compare both options against the same criteria: operating time, refuelling or charging time, emissions, safety, cost and waste.
- A claim that fuel cells are cleaner needs evidence about how the hydrogen was produced.
The better power source depends on the job and how the energy is supplied
- An evaluation weighs up the benefits and drawbacks before reaching a conclusion backed by evidence.
- Hydrogen fuel cells can suit vehicles that need long running times and quick refuelling, if hydrogen is available and stored safely.
- Rechargeable batteries can suit devices that are easy to plug into the grid and leave to charge.
- A hydrogen fuel cell has a stronger environmental case when the hydrogen is made without burning fossil fuels.
- Your conclusion should say which option is better for the stated use, not that one option is always better.
Exam technique
- To say why a fuel cell runs longer than a battery, state that fuel and oxygen are supplied continuously from outside the cell.
- In a comparison, judge fuel cells and rechargeable batteries on the same points and reach a decision for the named use.
- Say water is the only product in use, but avoid claiming the fuel cell has zero overall environmental impact.
Self review
- What two reactants must be supplied to a hydrogen fuel cell?
- What happens to the hydrogen inside a hydrogen fuel cell?
- What is the only product of the overall reaction in a hydrogen fuel cell?
- Why can a fuel cell run longer than a rechargeable cell without stopping?
- Why is a hydrogen fuel cell not automatically better for the environment?
5.2.2b Half equations for hydrogen fuel cells
Hydrogen is oxidised at one electrode and oxygen is reduced at the other
Definition
Half equation
An equation that shows the reaction at one electrode, including the electrons that are transferred.
Definition
Electrode
A conducting material through which electrical current enters or leaves a cell, with a simple cell using two different metal electrodes.
- A half equation shows the reaction at one electrode and includes the electrons transferred.
- At the hydrogen electrode, hydrogen atoms lose electrons, so hydrogen is oxidised.
- At the oxygen electrode, oxygen gains electrons, so oxygen is reduced.
- Oxidation is the loss of electrons, while reduction is the gain of electrons.
Key Idea
- Electrons appear on the products side of an oxidation half equation.
- Electrons appear on the reactants side of a reduction half equation.
Balancing the hydrogen half equation puts the electrons on the right
- Start with one hydrogen molecule, H2(g)\text{H}_2\text{(g)}H2(g), which has two hydrogen atoms.
- Write two hydrogen ions on the right to balance the atoms: H2(g)→2H+(aq)\text{H}_2\text{(g)} \rightarrow 2\text{H}^+\text{(aq)}H2(g)→2H+(aq).
- The two hydrogen ions carry a total charge of +2+2+2, while the molecule has no charge.
- Add two electrons to the right so that each side has a total charge of zero.
- The balanced half equation is H2(g)→2H+(aq)+2e−\text{H}_2\text{(g)} \rightarrow 2\text{H}^+\text{(aq)} + 2e^-H2(g)→2H+(aq)+2e−.
Example
- The left side has two hydrogen atoms and a total charge of zero.
- The right side has two hydrogen atoms and a total charge of 2(+1)+2(−1)=02(+1) + 2(-1) = 02(+1)+2(−1)=0.
Balancing the oxygen half equation puts the electrons on the left
- Start with one oxygen molecule and form two water molecules to balance the oxygen: O2(g)→2H2O(l)\text{O}_2\text{(g)} \rightarrow 2\text{H}_2\text{O(l)}O2(g)→2H2O(l).
- Add four hydrogen ions on the left to supply the four hydrogen atoms in the water.
- The four hydrogen ions carry a total charge of +4+4+4, so add four electrons on the left to make the charge zero.
- The balanced half equation is O2(g)+4H+(aq)+4e−→2H2O(l)\text{O}_2\text{(g)} + 4\text{H}^+\text{(aq)} + 4e^- \rightarrow 2\text{H}_2\text{O(l)}O2(g)+4H+(aq)+4e−→2H2O(l).
Common Mistake
- Do not put the electrons on the right of the oxygen half equation; oxygen is reduced, so it gains electrons.
- Check both the number of each atom and the total charge when you balance a half equation.
Matching the electrons lets the half equations combine into the overall reaction
- The hydrogen half equation releases two electrons, while the oxygen half equation uses four.
- Multiply the whole hydrogen half equation by two: 2H2(g)→4H+(aq)+4e−2\text{H}_2\text{(g)} \rightarrow 4\text{H}^+\text{(aq)} + 4e^-2H2(g)→4H+(aq)+4e−.
- Add the two half equations and cancel the four electrons and four hydrogen ions that appear on both sides.
- This gives the overall reaction 2H2(g)+O2(g)→2H2O(l)2\text{H}_2\text{(g)} + \text{O}_2\text{(g)} \rightarrow 2\text{H}_2\text{O(l)}2H2(g)+O2(g)→2H2O(l).
- Electrons must not appear in the overall equation because they are transferred between the electrodes, not created or destroyed.
Exam technique
- Balance the atoms first, then add electrons until the total charge is the same on both sides.
- Put electrons on the products side for the oxidation (hydrogen) and on the reactants side for the reduction (oxygen).
- To combine the half equations, scale them so the electrons match, add them, then cancel the electrons and any H+\text{H}^+H+ on both sides.
Self review
- What does a half equation show?
- On which side do the electrons appear in the hydrogen half equation?
- Write the balanced half equation for the reaction at the hydrogen electrode.
- Write the balanced half equation for the reaction at the oxygen electrode.
- What two things must you check when balancing a half equation?
