Chemical Cells and Fuel CellsAQA GCSE Chemistry: Revision notes
Section 1
How do simple cells convert chemical energy into electrical energy?
A simple cell is a device that converts chemical energy directly into electrical energy through a spontaneous redox reaction. The cell consists of two different metals (or electrodes) connected by an external circuit and immersed in an electrolyte solution.
The process works as follows:
- The more reactive metal acts as the negative electrode (anode) and is oxidised, losing electrons
- The less reactive metal acts as the positive electrode (cathode) where reduction occurs
- Electrons flow through the external circuit from the negative to positive electrode, creating an electric current
- The electrolyte allows ions to flow internally, completing the circuit
This spontaneous electron transfer is what generates electrical energy from the chemical reaction occurring inside the cell.
Think of a simple cell like a water slide: electrons flow downhill from high energy (negative electrode) to low energy (positive electrode), releasing energy as they go.
Section 2
How does metal reactivity affect the voltage produced by a simple cell?
The voltage (potential difference) produced by a simple cell depends directly on the difference in reactivity between the two metals used.
Key principle: The greater the reactivity difference between the metals, the greater the voltage produced.
This is because:
- More reactive metals are more easily oxidised and lose electrons more readily
- The larger the reactivity gap, the greater the driving force for the redox reaction
- Electrons flow more vigorously from a highly reactive metal to a less reactive one
Examples of reactivity differences:
| Metal pair | Reactivity difference | Voltage produced |
|---|---|---|
| Zinc and copper | Large | Relatively high |
| Magnesium and copper | Very large | Higher |
| Zinc and lead | Small | Relatively low |
To increase voltage in a simple cell, either use metals further apart in the reactivity series, or connect multiple cells in series (positive terminal of one to negative terminal of the next).
Examiners expect you to explain voltage in terms of reactivity difference, not just state that high reactivity = high voltage. Always reference how easily the metals are oxidised.
A zinc-copper cell produces a higher voltage than a zinc-lead cell because zinc and copper are further apart in the reactivity series, creating a greater driving force for electron transfer.
Section 3
What is the difference between rechargeable and non-rechargeable batteries?
Batteries are made up of multiple cells connected in series. They are classified based on whether their chemical reactions are reversible.
Non-rechargeable (primary) batteries:
- The chemical reactions inside are irreversible
- Once the reactants are used up, the battery cannot be restored
- Current can only flow in one direction
- Examples: alkaline cells, zinc-carbon batteries
- Must be disposed of when spent
Rechargeable (secondary) batteries:
- The chemical reactions inside are reversible
- When connected to an external power source, the direction of the reaction can be reversed
- Electrons are forced to flow backwards through the cell, restoring the original reactants
- Can be recharged many times
- Examples: lithium-ion batteries, lead-acid batteries, nickel-cadmium batteries
- More economical and environmentally friendly over time
How recharging works: When a reversible battery is recharged, an external voltage is applied that is greater than the cell's own voltage. This forces electrons to flow backwards, pushing the reaction in reverse and regenerating the original reactants. Once the external power is removed, the cell can produce electricity again.
Students often confuse 'reversible' with 'reversing polarity'. Reversible means the chemical reaction itself can go backwards, not just the direction of current flow.
When explaining the difference, always reference reversibility of the reaction itself. Mark schemes require explicit mention that recharging reverses the chemical reaction, not just that current flows backwards.
Section 4
How do hydrogen fuel cells generate electricity?
A hydrogen fuel cell is an electrochemical cell that generates electricity through the controlled reaction of hydrogen gas with oxygen. Unlike batteries, fuel cells continuously produce electricity as long as fuel is supplied, rather than storing a fixed amount of chemical energy.
Overall equation for a hydrogen fuel cell:
2H₂ + O₂ → 2H₂O
How a hydrogen fuel cell works:
- Hydrogen gas enters the negative electrode (anode)
- Oxygen gas enters the positive electrode (cathode)
- Inside the cell, hydrogen is oxidised and oxygen is reduced
- Electrons flow through an external circuit, generating an electric current
- The only product is water (and heat)
Key advantages:
- Clean: produces only water as a by-product, no harmful emissions
- Efficient: converts chemical energy to electrical energy very efficiently
- Renewable: if hydrogen is produced from renewable sources, the fuel is renewable
- Quiet operation: no combustion, so very quiet
Key disadvantages:
- Infrastructure: requires a hydrogen supply chain and fuelling stations, which don't yet exist widely
- Storage: hydrogen is difficult and dangerous to store
- Production: most hydrogen is currently produced from fossil fuels
- Cost: fuel cell technology is currently expensive
- Durability: fuel cells have limited lifespan
In a hydrogen fuel cell, H₂ is oxidised at the anode (loses electrons) and O₂ is reduced at the cathode (gains electrons). The electrons flow through the external circuit, powering devices, before combining with oxygen and hydrogen ions to form water.
Section 5
What are the half-equations for a hydrogen fuel cell? (Higher Tier)
Hydrogen fuel cells involve redox reactions that can be represented as separate half-equations for oxidation and reduction.
At the negative electrode (anode) – oxidation:
2H₂ - 4e⁻ → 4H⁺
or in acidic solution:
2H₂ → 4H⁺ + 4e⁻
At the positive electrode (cathode) – reduction:
O₂ + 4H⁺ + 4e⁻ → 2H₂O
or in alkaline solution:
O₂ + 2H₂O + 4e⁻ → 4OH⁻
How to combine half-equations:
- Ensure the electrons lost in oxidation equal the electrons gained in reduction
- Add the two half-equations together
- Cancel any species that appear on both sides
Verification: When the anode and cathode equations above are combined, they give: 2H₂ + O₂ → 2H₂O ✓
Note: The exact form of the half-equations depends on whether the fuel cell operates in acidic or alkaline conditions, but the overall equation remains the same.
When writing half-equations for fuel cells, check that electrons are balanced. The number of electrons lost at the anode must equal the number gained at the cathode before you add them together.
Section 6
How do hydrogen fuel cells compare to rechargeable batteries?
Advantages of hydrogen fuel cells over rechargeable batteries:
| Feature | Hydrogen Fuel Cell | Rechargeable Battery |
|---|---|---|
| By-products | Only water (clean) | Depends on chemistry; often involves heavy metals |
| Energy density | Very high – can store more energy per kg | Lower – heavier for equivalent power |
| Refuelling time | Fast – minutes | Hours (charging time) |
| Operating time | Continuous as long as fuel supplied | Limited to stored charge |
| Emissions | Zero harmful emissions | Zero direct emissions but depends on electricity source |
| Scalability | Suitable for large vehicles and power plants | Better for portable devices |
Disadvantages of hydrogen fuel cells compared to rechargeable batteries:
| Feature | Hydrogen Fuel Cell | Rechargeable Battery |
|---|---|---|
| Infrastructure | Requires hydrogen production, storage, and distribution networks – not yet widespread | Established electricity grid and charging infrastructure |
| Cost | Expensive technology; high manufacturing cost | More affordable and mass-produced |
| Durability | Limited lifespan; electrodes degrade | Can last many charge cycles |
| Hydrogen source | Most hydrogen from fossil fuels currently | Can use renewable electricity |
| Storage | Hydrogen difficult and dangerous to store | Batteries safer and easier to store |
| Efficiency | Good but losses in hydrogen production | Direct conversion of stored energy |
Current applications:
- Fuel cells: Research vehicles, some buses, future heavy transport
- Batteries: Electric cars, phones, laptops, portable electronics
Examiners want balanced evaluation: mention both advantages AND disadvantages. For fuel cells, always emphasise the 'clean water' advantage but also the 'lack of hydrogen infrastructure' disadvantage.
Students sometimes claim fuel cells produce 'zero emissions globally'. This is misleading – most current hydrogen production uses fossil fuels, so emissions are shifted upstream. Only mention zero direct/local emissions.
Must Know
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Simple cells convert chemical energy to electrical energy using two different metals in an electrolyte; the greater the reactivity difference between metals, the greater the voltage produced
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Non-rechargeable batteries have irreversible reactions and cannot be reused; rechargeable batteries have reversible reactions and can be restored by applying external voltage to reverse the chemical reaction
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Hydrogen fuel cells are electrochemical cells that react hydrogen and oxygen to produce electricity and water: 2H₂ + O₂ → 2H₂O
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Hydrogen fuel cells produce only water (clean advantage) but require hydrogen infrastructure development (disadvantage vs. existing electricity grid)
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Half-equations for hydrogen fuel cells: anode: 2H₂ → 4H⁺ + 4e⁻ and cathode: O₂ + 4H⁺ + 4e⁻ → 2H₂O (electrons must be balanced when combining)
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Fuel cells offer high energy density and fast refuelling but are more expensive than batteries; batteries have established infrastructure but fuel cells produce zero harmful emissions locally
That's the notes covered.
Carry on to the next subtopic.