Chemical cells and fuel cells (chemistry only) revision notes

Chemical Cells and Fuel Cells

1. What is a chemical cell?

• A cell is a small electrochemical device that converts chemical energy into electrical energy.
• It contains two different metals (the electrodes) that are in contact with an electrolyte – a solution that can conduct ions.
• The reaction that occurs at the electrodes is an oxidation‑reduction (redox) reaction. Electrons are released at the anode (oxidation) and travel through the external circuit to the cathode (reduction), producing a flow of electric current.

2. Voltage in a simple cell

• The voltage (potential difference) produced by a cell depends on:

1. Electrode material – different metals have different tendencies to lose or gain electrons.
2. Electrolyte composition – the ions present and their concentration influence the cell’s internal resistance and the overall reaction.
3. Temperature – higher temperatures generally increase reaction rates but can also change the relative reactivity of the metals.

• The cell’s voltage is calculated from the difference in standard electrode potentials of the two electrodes. The more positive the cathode potential and the more negative the anode potential, the higher the cell voltage.

3. Why two different electrodes are required

• A single metal cannot produce a potential difference because the same reaction would occur at both ends, cancelling any net electron flow.
• Two different metals create a potential difference because each metal has a distinct standard reduction potential. This difference drives the electron flow.
• The electrolyte allows ions to move between the electrodes, completing the circuit internally.

4. Cells in series – building a battery

• A battery is simply two or more cells connected in series.
• In a series connection, the positive terminal of one cell is connected to the negative terminal of the next. The voltages of each cell add up, giving a higher overall voltage.
• The current (amperage) remains the same as that of a single cell because the same electrons flow through each cell.

5. Rechargeable vs. non‑rechargeable cells

• Non‑rechargeable cells (e.g., alkaline batteries) run until one reactant is exhausted. The reaction is irreversible under normal conditions.
• Rechargeable cells (e.g., lead‑acid, nickel‑metal hydride) can be re‑charged by applying an external electrical current that reverses the redox reaction, restoring the original reactants.
• The key difference is the ability to reverse the reaction, which is governed by the cell’s design and the stability of the electrode materials.

6. Fuel cells – a different approach

• A fuel cell is supplied with an external fuel (commonly hydrogen) and an oxidiser (usually oxygen from air). It does not store energy chemically; instead, it continuously converts the fuel into electricity as long as the fuel is supplied.
• In a hydrogen fuel cell:

1. Anode reaction (oxidation):

1. Cathode reaction (reduction):

1. Overall reaction:

• The product of the overall reaction is water, which is a clean by‑product.
• Because the fuel is continuously supplied, a fuel cell can operate for a long time without the need for a battery‑like storage system.

7. Comparing fuel cells with rechargeable batteries

| Feature | Hydrogen Fuel Cell | Rechargeable Battery | |---------|--------------------|----------------------| | Fuel supply | Continuous hydrogen gas | Stored chemical reactants | | Products | Water (clean) | Often metal oxides or salts | | Rechargeability | Not required – fuel is replaced | Requires external electricity | | Operating use | Vehicles, stationary power | Portable electronics, backup power | | Advantages | High energy density, zero emissions | Simple, low cost, widely available | | Disadvantages | Requires hydrogen infrastructure | Limited energy density, finite life |

8. Safety when working with cells

• Always wear protective gloves and goggles when handling electrolytes or discharging cells.
• Keep liquids away from direct contact with the electrodes to avoid accidental short circuits.
• Dispose of used cells according to local hazardous waste regulations.

9. Key terms

• Electrode – the metal surface where oxidation or reduction occurs.
• Electrolyte – a solution that conducts ions.
• Anode – the electrode where oxidation takes place.
• Cathode – the electrode where reduction takes place.
• Series connection – linking cells so that voltages add.
• Rechargeable – a cell whose reaction can be reversed.
• Fuel cell – an electrochemical device that continuously supplies fuel and oxidiser.
• Half‑equation – the reaction occurring at a single electrode.
• Overall reaction – the sum of the half‑equations.
• Potential difference – the voltage produced by a cell.

10. Exam tips

• Draw clear diagrams of cells, showing anode, cathode, electrolyte, and electron flow.
• Remember that voltage is the difference in electrode potentials, not the sum of individual potentials.
• Practice writing half‑equations for both anode and cathode reactions in fuel cells.
• Compare the advantages and disadvantages of fuel cells and batteries using a table.
• Check that you can explain why a single metal cannot produce a voltage.

11. Common mistakes

• Confusing the roles of anode and cathode.
• Assuming that a single metal can generate a voltage.
• Forgetting that series connections add voltages but not currents.
• Mixing up the products of a fuel cell (water) with those of a typical battery (metal oxides).
• Overlooking the need for an external source to reverse reactions in rechargeable cells.

12. Further reading

• Look at the standard electrode potential tables to predict which metal combinations will give the highest voltage.
• Explore the design of commercial hydrogen fuel cells to understand how they manage fuel and oxidiser supply.
• Review safety guidelines for handling electrolytes and hydrogen gas.

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> Tip: When you write the half‑equations for a hydrogen fuel cell, remember to balance both mass and charge. The overall reaction should be > > > > which produces water as the only product.

Energy changes focus for 720

This undefined is anchored to AQA GCSE Chemistry 8462 Unit 4.5. It separates exothermic reactions, endothermic reactions, reaction profiles, activation energy, bond-energy calculations, chemical cells and fuel cells so students do not collapse nearby ideas into one generic energy answer.

How to answer exam questions

Start by naming the energy-change idea being tested. For reaction profiles, label reactants, products, activation energy and overall energy change. For bond-energy calculations, add the energy needed to break bonds, subtract the energy released when bonds form, keep the sign, and state whether the result is exothermic or endothermic.

Common checks

Check whether the question asks for temperature change, energy transfer, a diagram label, a calculation, a cell voltage pattern, a fuel-cell comparison or an evaluation. Use the exact subtopic wording and avoid drifting into chemical changes, rates, equilibrium or electrolysis unless the question explicitly connects them.

Practice method

After reading the section, write a three-part response: define the key idea, apply it to the named reaction or device, then explain the evidence using the correct GCSE Chemistry term. For calculations, show formula, substitution, calculation, final answer, unit and conclusion.