7.6 Fuel Cells

A free GCSE Chemistry lesson for AQA (8462) from OBA School: hydrogen fuel cells, the reactions at the electrodes, and comparing fuel cells with rechargeable batteries. Watch the video, read the explanation and worked examples, then try the revision questions. You can work through the full lesson, with every video, practice quizzes and printable worksheets with mark schemes, in our free online course.

What you will learn

• Describe what a fuel cell is and how a hydrogen fuel cell produces electricity
• Write the overall equation for the reaction in a hydrogen fuel cell
• (Higher tier) write half equations for the reactions at the electrodes
• Evaluate the use of hydrogen fuel cells compared with rechargeable cells and batteries

Key words

• Fuel cell: A cell supplied by an external source of fuel (e.g. hydrogen) and oxygen or air. The fuel is oxidised electrochemically to produce a potential difference.
• Electro- chemical oxidation: Oxidation that happens at an electrode, so the electrons lost travel round an external circuit instead of being released as heat.
• Proton exchange membrane: A type of electrolyte in many hydrogen fuel cells that lets H⁺ ions pass through.
• Oxidation / reduction: Loss / gain of electrons. Hydrogen is oxidised; oxygen is reduced.

What is a fuel cell?

A fuel cell is supplied with a fuel and oxygen (or air) from outside. The fuel is oxidised electrochemically inside the cell to produce a potential difference. Unlike a battery, a fuel cell does not run flat: it keeps working as long as fuel and oxygen are supplied. The most important type uses hydrogen as the fuel.

Overall reaction: 2H₂(g) + O₂(g) → 2H₂O(l). This is the same reaction as burning hydrogen, but the energy is transferred as electricity instead of mainly as heat. The only product is water.

• Energy transferred as; Burning hydrogen in an engine mostly heat, then some is turned into movement; Hydrogen fuel cell mostly electricity, directly
• Efficiency; Burning hydrogen in an engine lower – much energy is wasted as heat; Hydrogen fuel cell higher – more of the energy does useful work
• Products; Burning hydrogen in an engine water (plus some nitrogen oxides from the hot air); Hydrogen fuel cell water only

Reactions at the electrodes (Higher tier)

• Anode (negative); What happens hydrogen molecules lose electrons and form hydrogen ions (oxidation); the electrons flow through the external circuit; Half equation 2H₂ → 4H⁺ + 4e−
• Cathode (positive); What happens oxygen gains electrons and reacts with hydrogen ions that have moved through the electrolyte, forming water (reduction); Half equation O₂ + 4H⁺ + 4e− → 2H₂O

Worked example: adding the half equations

2H₂ → 4H⁺ + 4e− and O₂ + 4H⁺ + 4e− → 2H₂O
Add them and cancel 4H⁺ and 4e− from both sides:
2H₂ + O₂ → 2H₂O

Answer: the overall equation, with no ions or electrons left over

• Negative electrode: hydrogen is oxidised; Half equation 2H₂(g) + 4OH−(aq) → 4H₂O(l) + 4e−
• Positive electrode: oxygen is reduced; Half equation O₂(g) + 2H₂O(l) + 4e− → 4OH−(aq)

Some fuel cells use an alkaline electrolyte instead of an acid. The half equations look different, but adding them and cancelling the electrons, OH− ions and water that appear on both sides gives the same overall equation, 2H₂ + O₂ → 2H₂O. Either version is accepted if it is balanced for atoms and charge.

Exam tip: In a fuel cell the anode is negative (it releases electrons into the circuit). In electrolysis the anode is positive. In both, oxidation happens at the anode.

Fuel cells or rechargeable batteries?

• For; Hydrogen fuel cells only product is water, so no pollution where used; keep working as long as fuel is supplied; refuel in minutes; fewer parts to wear out; more energy per kg, so longer range in vehicles; Rechargeable batteries store electricity for cars and phones; no explosive gas to store; charging network already exists; can be recharged many times
• Against; Hydrogen fuel cells hydrogen is a flammable gas, hard to store (high pressure or very cold); most hydrogen is made from fossil fuels or by electrolysis using electricity; few refuelling stations; expensive catalysts (e.g. platinum); Rechargeable batteries take hours to recharge; limited number of charge cycles; contain toxic metals that are hard to dispose of; energy to charge may come from fossil fuels

Worked example: mass of water produced

2H₂ + O₂ → 2H₂O. Mr: H₂ = 2, H₂O = 18.
4 g of hydrogen (2 moles) makes 36 g of water (2 moles), so mass of water = 9 × mass of hydrogen.
A fuel-cell car uses 1.2 kg of hydrogen: water produced = 1.2 × 9 = 10.8 kg.

Answer: 10.8 kg of water

Common mistakes

• Saying a fuel cell “burns” hydrogen. The hydrogen is oxidised electrochemically; there is no flame.
• Saying fuel cells cause no pollution at all. Making and storing the hydrogen can cause pollution.
• Writing the anode half equation with electrons on the left. Hydrogen loses electrons: they go on the right.
• Thinking a fuel cell needs recharging. It needs refuelling with hydrogen.

Revision questions

Try these, then check your answers below each one.

Question 1. What is supplied to a hydrogen fuel cell? (Choose from: Hydrogen and oxygen (or air), Hydrogen and water, Oxygen and water, Hydrogen and carbon dioxide.)

Answer: Hydrogen and oxygen (or air)

Working: Hydrogen (the fuel) and oxygen from the air are supplied.

Question 2. Name the only chemical product of a hydrogen fuel cell.

Answer: water

Working: 2H₂ + O₂ → 2H₂O. The only product is water.

Question 3. (Higher tier) Which half equation shows the reaction at the anode of a hydrogen fuel cell with an acidic electrolyte? (Choose from: 2H₂ → 4H⁺ + 4e−, O₂ + 4H⁺ + 4e− → 2H₂O, 4H⁺ + 4e− → 2H₂, 2H₂O → O₂ + 4H⁺ + 4e−.)

Answer: 2H₂ → 4H⁺ + 4e−

Working: At the anode hydrogen loses electrons: 2H₂ → 4H⁺ + 4e−. The oxygen equation happens at the cathode.

Related lessons

• Previous lesson: 7.5 Chemical Cells and Batteries
• Next lesson: 7.7 Required Practicals: Energy Changes
• In this chapter: 7.1 Exothermic and Endothermic Reactions
• In this chapter: 7.2 Using Energy Transfers from Reactions
• In this chapter: 7.3 Reaction Profiles

Continue in the full course

This lesson is section 7.6 (AGCCH0706) of our free GCSE Chemistry course for AQA. In the full course you get every video for this section, a self-marking practice quiz, a printable worksheet with a mark scheme, chapter tests, and your progress saved. Go to the full GCSE Chemistry course.