Commercial applications of electrochemical cells (A-level only) — AQA A-Level Chemistry
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Commercial applications of electrochemical cells (A-level only) explained
Any pair of half-cells with different electrode potentials can deliver electrical energy, and that is all a commercial cell is.
Read the full explanation
The electrode with the more negative potential is the negative terminal, where oxidation releases electrons into the external circuit; the more positive electrode is where reduction takes them in. The EMF is fixed by the difference between the two electrode potentials, so making a cell physically larger does not raise its voltage. It raises how long the cell lasts, because there is more reactant. A non-rechargeable cell is finished when a reactant is used up. A rechargeable cell, such as a lithium-ion cell, has both electrode reactions driven backwards when an external voltage larger than its EMF is applied. Inside the cell, ions move through the electrolyte or porous separator; electrons never do.
Fuel cells are used to generate an electric current and do not need to be electrically recharged.
A fuel cell is not a store of energy but a converter. Fuel and oxygen are fed in continuously from outside, so the electrode reactions never run down and the cell needs no electrical recharging; it needs refuelling. A hydrogen and oxygen fuel cell converts the chemical energy of the reaction directly into electrical energy, skipping the heat and mechanical stages of an engine, so it reaches a much higher efficiency than burning the same hydrogen. The only product at the point of use is water, so no carbon dioxide, nitrogen oxides or particulates are released in the street. The difficulties are upstream and practical: hydrogen is usually manufactured from methane, so carbon dioxide is released at the plant; it is hard to store and transport safely; and the electrodes need expensive catalysts.
The electrode reactions in an alkaline hydrogen–oxygen fuel cell.
In the alkaline cell the electrolyte supplies hydroxide ions, so both half-equations must contain OH⁻. At the negative electrode hydrogen is oxidised: H₂ + 2OH⁻ → 2H₂O + 2e⁻, for which E⦵ = −0.83 V. At the positive electrode oxygen is reduced: ½O₂ + H₂O + 2e⁻ → 2OH⁻, E⦵ = +0.40 V. Adding them cancels the electrons, one water and the hydroxide to give the overall reaction H₂ + ½O₂ → H₂O, and the EMF is +0.40 − (−0.83) = +1.23 V. Electrons leave the negative electrode and travel round the external circuit; the hydroxide ions travel the other way through the electrolyte, from the oxygen electrode to the hydrogen electrode, so charge does not build up at either end.
The benefits and risks to society associated with using these cells.
Judging a cell means looking beyond the device itself. The benefits are portable, reliable power; fuel cells release only water at the point of use, so no carbon dioxide or particulates enter a city street; rechargeable cells reduce the number of cells thrown away. The risks are real too. Cells contain toxic or corrosive materials, including cadmium, lithium compounds and concentrated alkali, which can leach from landfill into groundwater; lithium cells can catch fire if damaged or overcharged; hydrogen is extremely flammable and is stored at very high pressure. There is also the wider accounting: most hydrogen is made from methane, so the carbon dioxide appears at the plant rather than the exhaust, and mining lithium and cobalt carries its own environmental and human cost.
Students should be able to: use given electrode data to deduce the reactions occurring in non-rechargeable and rechargeable cells deduce the EMF of a cell explain how the electrode reactions can be used to generate an electric current.
Given two half-equations with their E⦵ values, the more negative one is reversed and becomes the oxidation at the negative electrode; the more positive one runs as written and is the reduction at the positive electrode. Scale them so the electrons cancel, add them, and you have the discharge reaction. The EMF is the difference between the two potentials. Electrons are pushed out of the negative electrode, through the external circuit and into the positive electrode, and it is that flow that does the useful work. Inside the cell the circuit is completed by ions moving through the electrolyte or porous separator, never by electrons. Recharging applies a larger external voltage, which drives both electrode reactions backwards and regenerates the original reactants.
Your focus
- Deduce the EMF and the polarity of a commercial cell from two given electrode potentials.
- Explain why a lithium-ion cell can be recharged whereas an alkaline zinc cell cannot.
- Describe the role of the electrolyte in allowing a cell to deliver current, without referring to electron flow through it.
Show all 15 objectives
- Explain why a hydrogen and oxygen fuel cell never needs recharging, in terms of the supply of reactants.
- Compare a hydrogen fuel cell with a petrol engine for efficiency and for what is released at the point of use.
- Suggest two practical difficulties in storing hydrogen for use in a vehicle.
- Write the half-equation occurring at the negative electrode of an alkaline hydrogen and oxygen fuel cell and check that charge balances.
- Combine the two electrode half-equations to show that water is the only product.
- Calculate the EMF of the alkaline fuel cell from E⦵ values of −0.83 V and +0.40 V.
- Evaluate the use of hydrogen fuel cells in buses, naming one benefit and one risk.
- Explain why the carbon footprint of a fuel-cell vehicle depends on how its hydrogen was manufactured.
- Justify why used lithium cells are collected separately rather than sent to landfill.
- Deduce from two given E⦵ values which electrode of a cell is negative and write the reaction occurring there.
- Write the charging equations for a rechargeable cell as the reverse of its discharge equations.
- Explain, in terms of electrons and ions, how the electrode reactions produce a current in an external circuit.
Commercial applications of electrochemical cells (A-level only) exam tips
Marking Points
- one mark for stating that the EMF equals the difference between the two electrode potentials
- one mark for identifying the negative electrode as the one where oxidation occurs
- one mark for stating that the electrolyte or porous separator allows ions to move so the circuit is complete
- one mark for distinguishing a rechargeable cell, whose reactions are reversed by an applied voltage, from a non-rechargeable one
- one mark for stating that the fuel and the oxygen are supplied continuously from outside the cell
- one mark for stating that the cell does not need to be electrically recharged because the reactants are supplied continuously
- one mark for identifying water as the only product at the point of use
- one mark for a named limitation, such as storing or transporting hydrogen, or making hydrogen from a fossil fuel
- one mark for H₂ + 2OH⁻ → 2H₂O + 2e⁻ as the oxidation at the negative electrode
- one mark for ½O₂ + H₂O + 2e⁻ → 2OH⁻ as the reduction at the positive electrode
- one mark for the overall equation H₂ + ½O₂ → H₂O with the electrons cancelling
- one mark for an EMF of +1.23 V from the E⦵ values −0.83 V and +0.40 V
- one mark for a named benefit, such as only water being released at the point of use
- one mark for a named risk, such as the flammability of hydrogen or the toxicity of a metal in the cell
- one mark for referring to how the hydrogen or the electricity was produced, rather than only the cell itself
- one mark for a conclusion that weighs the stated benefit against the stated risk
- one mark for identifying the half-equation with the more negative E⦵ as the one that is reversed
- one mark for stating that electrons flow through the external circuit from the negative to the positive electrode
- one mark for the EMF as the difference between the two electrode potentials
- one mark for stating that the electrolyte or porous separator allows ions to move and complete the circuit
- one mark for writing the recharging reactions as the reverse of the discharge reactions
Examiner Tips
- 💡Identify the negative electrode first by finding the more negative E⦵; everything else follows from that.
- 💡Keep electron flow in the external wire and ion flow inside the cell strictly separate in your answer.
- 💡'Does not need recharging' earns the mark only when paired with the reason: the reactants are supplied continuously.
- 💡For any comparison with a petrol engine, name the efficiency advantage and the emission at the point of use separately.
- 💡Read the electrolyte before writing anything; alkaline and acidic fuel cells have completely different half-equations.
- 💡Check each half-equation for charge as well as atoms; in the alkaline cell it is the OH⁻ that makes the charges work.
- 💡Name a specific substance in every point; 'chemicals leak out' scores nothing, 'cadmium leaches into groundwater' scores.
- 💡An evaluation needs a stated conclusion at the end, not just a list of two columns.
- 💡Underline the more negative E⦵ on the question paper; that electrode is negative and its equation is the one you flip.
- 💡For a 'how does this generate a current' mark, you must mention electrons in the external circuit and ions inside the cell.
Common Mistakes
- writing that the separator or electrolyte allows electrons to flow, which is not credited
- calling the electrode at which oxidation happens the positive electrode
- assuming a larger cell has a higher EMF rather than a greater capacity
- saying a flat cell has 'run out of voltage' rather than run out of a reactant
- describing a fuel cell as storing energy in the same way as a battery
- claiming a hydrogen fuel cell is carbon neutral without considering how the hydrogen was manufactured
- stating that the fuel cell is 100% efficient
- saying the cell must be recharged when the electrolyte is used up, when the electrolyte is not consumed
- using the acidic half-equations H₂ → 2H⁺ + 2e⁻ and O₂ + 4H⁺ + 4e⁻ → 2H₂O when the electrolyte is stated to be alkaline
- leaving OH⁻ out of the half-equations, so the charges do not balance
- writing hydrogen as reduced at the negative electrode instead of oxidised
- adding the two E⦵ values to give −0.43 V instead of subtracting to give +1.23 V
- writing 'better for the environment' with no named substance or effect
- ignoring that hydrogen is usually manufactured from methane, so carbon dioxide is still released elsewhere
- treating the disposal of lithium or cadmium cells as harmless
- listing only advantages when the command word is 'evaluate' or 'discuss'
- saying the electrolyte or separator allows electrons to move, which examiners refuse
- reversing the wrong half-equation, so the EMF comes out negative
- routing electrons through the electrolyte instead of the external wire
- reversing only one electrode reaction when describing recharging