Chemical cells convert chemical energy into electrical energy. Understanding how voltaic cells generate a voltage, how hydrogen fuel cells work, and the differences between rechargeable and non-rechargeable cells is essential for this topic.
A simple chemical cell is made by connecting two different metals in a solution of their own ions (an electrolyte). The difference in reactivity between the two metals determines the voltage produced by the cell.
Essential components of a cell:
Electrodes: Zinc and Copper
Electrolytes: Zinc sulfate solution and Copper sulfate solution (connected by a salt bridge)
Zinc is more reactive than copper, so zinc is the negative electrode and copper is the positive electrode.
At the zinc electrode (negative):
Zn(s) → Zn²⁺(aq) + 2e⁻ (oxidation)
Zinc atoms lose electrons and go into solution as zinc ions. The electrons flow through the external circuit to the copper electrode.
At the copper electrode (positive):
Cu²⁺(aq) + 2e⁻ → Cu(s) (reduction)
Copper ions from the solution gain electrons and deposit as copper metal on the electrode.
The zinc electrode decreases in mass (it dissolves) and the copper electrode increases in mass (copper deposits). The cell produces a voltage of about 1.1 V.
| Cell | Negative Electrode | Positive Electrode | Relative Voltage |
|---|---|---|---|
| Magnesium-Copper | Mg (very reactive) | Cu (very unreactive) | High (large reactivity difference) |
| Zinc-Copper | Zn (fairly reactive) | Cu (very unreactive) | Medium |
| Iron-Copper | Fe (moderately reactive) | Cu (very unreactive) | Lower than Zn-Cu |
| Iron-Zinc | Fe (less reactive) | Zn (more reactive) | Low (small reactivity difference) |
Other factors that affect cell voltage:
Non-rechargeable cells contain chemicals that react irreversibly. When the chemical reactions have gone to completion, no more electrical energy can be produced and the cell must be replaced.
Common types of non-rechargeable cells:
Rechargeable cells can be used hundreds of times because the chemical reactions can be reversed by passing an electric current through the cell in the opposite direction during recharging.
Common types of rechargeable cells:
| Feature | Non-Rechargeable Cells | Rechargeable Cells |
|---|---|---|
| Chemical reactions | Irreversible | Reversible |
| Lifespan | Used once then discarded | Can be recharged hundreds of times |
| Initial cost | Low | Higher |
| Long-term cost | Higher (need constant replacement) | Lower (reuse many times) |
| Environmental impact | More waste, toxic chemicals in landfill | Less waste, but eventual disposal still needed |
| Convenience | Ready to use, no recharging needed | Need access to electricity for recharging |
| Examples | Alkaline, zinc-carbon | Lithium-ion, lead-acid, NiMH |
A hydrogen fuel cell works like a battery, but it does not run down or need recharging as long as hydrogen and oxygen are continuously supplied. The fuel cell converts the chemical energy of hydrogen directly into electrical energy without combustion.
Structure of a hydrogen fuel cell:
At the anode (negative electrode):
H₂(g) → 2H⁺(aq) + 2e⁻
Hydrogen gas is oxidised. The hydrogen molecules split into hydrogen ions (protons) and electrons. The protons pass through the electrolyte membrane to the cathode. The electrons flow through the external circuit, providing the electric current.
At the cathode (positive electrode):
½O₂(g) + 2H⁺(aq) + 2e⁻ → H₂O(l)
Oxygen gas is reduced. The oxygen combines with the hydrogen ions and electrons to form water.
Overall reaction:
H₂(g) + ½O₂(g) → H₂O(l)
The only product is water, making this a very clean process.
There are two common versions of the half equations depending on whether the product is liquid water or water vapour:
Anode: H₂(g) → 2H⁺(aq) + 2e⁻
Cathode: ½O₂(g) + 2H⁺(aq) + 2e⁻ → H₂O(l)
Overall: H₂(g) + ½O₂(g) → H₂O(l)
Anode: H₂(g) → 2H⁺(aq) + 2e⁻
Cathode: ½O₂(g) + 2H⁺(aq) + 2e⁻ → H₂O(g)
Overall: H₂(g) + ½O₂(g) → H₂O(g)
| Advantages | Disadvantages |
|---|---|
| Only product is water — no pollution or greenhouse gases at point of use | Hydrogen is a gas and is difficult to store and transport safely |
| Higher efficiency than combustion engines (convert more chemical energy to useful energy) | Hydrogen is often produced from fossil fuels (e.g. by steam reforming of methane), which releases CO₂ |
| No moving parts — quiet operation and low maintenance | Producing hydrogen by electrolysis of water requires large amounts of electricity |
| Can be used in vehicles, buildings, and portable devices | Fuel cells and hydrogen storage systems are expensive |
| Continuous energy supply as long as fuel is provided (unlike batteries that run down) | Limited refuelling infrastructure for hydrogen vehicles |
| Lighter than batteries for the same energy output (important for vehicles) | Hydrogen is highly flammable and must be stored under high pressure or at very low temperatures |
Hydrogen combustion: 2H₂(g) + O₂(g) → 2H₂O(g)
This is a highly exothermic reaction and the only product is water vapour. There are no carbon dioxide emissions, no carbon monoxide, and no particulates.
| Feature | Hydrogen Fuel Cell | Hydrogen Combustion Engine |
|---|---|---|
| How it works | Electrochemical reaction converts chemical energy directly to electrical energy | Hydrogen is burned in air to release heat energy |
| Efficiency | Higher (up to 60-70%) | Lower (about 25-30%, similar to petrol engines) |
| Products | Water | Water (and some nitrogen oxides from air at high temperatures) |
| Noise | Very quiet | Noisy (like a conventional engine) |
| Moving parts | Few or none | Many moving parts |
Method 1: Electrolysis of water
2H₂O(l) → 2H₂(g) + O₂(g)
An electric current is passed through water, decomposing it into hydrogen and oxygen gases. If the electricity comes from renewable sources (wind, solar, hydroelectric), the hydrogen produced is truly clean with zero carbon emissions. However, electrolysis is expensive and requires a lot of electricity.
Method 2: Steam reforming of methane (natural gas)
CH₄(g) + H₂O(g) → CO(g) + 3H₂(g)
This is the most common industrial method for producing hydrogen. Methane reacts with steam at high temperature (about 700°C) with a nickel catalyst. The carbon monoxide produced can then react with more steam:
CO(g) + H₂O(g) → CO₂(g) + H₂(g)
The overall result is that large amounts of CO₂ are produced, meaning the hydrogen is not truly carbon-free unless carbon capture and storage is used.
| Feature | Simple Voltaic Cell | Commercial Battery | Fuel Cell |
|---|---|---|---|
| Energy source | Chemical reactions between metals and electrolyte | Chemical reactions between stored reactants | Chemical reactions between externally supplied fuel and oxygen |
| Rechargeable? | Depends on the chemistry | Some are, some are not | Not needed — runs continuously while fuel is supplied |
| Runs out? | Yes — reactants are consumed | Yes — stored chemicals are used up | Only if fuel supply stops |
| Voltage | Depends on metals used (typically 0.5-2 V) | Depends on chemistry (typically 1.2-3.7 V per cell) | About 0.7 V per cell (multiple cells stacked for more voltage) |
| Environmental impact | Moderate — contains metal salts | Varies — some contain toxic metals | Low if hydrogen is produced cleanly |
Key limitations in detail:
1. Explain why a zinc-copper cell produces a higher voltage than an iron-copper cell.
The voltage of a cell depends on the difference in reactivity between the two metals used as electrodes. Zinc is more reactive than iron, so there is a greater reactivity difference between zinc and copper than between iron and copper. A greater reactivity difference means more energy is available to push electrons through the circuit, resulting in a higher voltage.
2. Write the half equations for a hydrogen fuel cell and explain why the only product is water.
Anode: H₂(g) → 2H⁺(aq) + 2e⁻
Cathode: ½O₂(g) + 2H⁺(aq) + 2e⁻ → H₂O(l)
Overall: H₂(g) + ½O₂(g) → H₂O(l)
The only product is water because hydrogen is oxidised to H⁺ ions, which then combine with oxygen and electrons at the cathode to form water. No carbon-containing compounds are involved, so no CO₂ or other pollutants are produced.
3. Explain the difference between a rechargeable and a non-rechargeable cell.
In a non-rechargeable cell, the chemical reactions are irreversible. Once the reactants have been converted to products, the cell is dead and cannot be used again. In a rechargeable cell, the chemical reactions are reversible. When the cell discharges, the forward reactions produce electricity. When an external current is applied (recharging), the reverse reactions occur, restoring the original reactants so the cell can be used again.
4. Describe two advantages and two disadvantages of using hydrogen fuel cells in vehicles compared to petrol engines.
Advantages: (1) The only product is water, so there are no harmful emissions or greenhouse gases at the point of use. (2) Fuel cells are more efficient than petrol engines, converting more of the chemical energy into useful energy.
Disadvantages: (1) Hydrogen is difficult to store — it must be kept under high pressure or at very low temperatures, requiring heavy and expensive storage tanks. (2) There is limited refuelling infrastructure for hydrogen vehicles, with very few hydrogen stations available.
5. A student sets up a cell with a magnesium electrode and a copper electrode. Explain which metal forms the negative electrode and write a half equation for the reaction at this electrode.
Magnesium is more reactive than copper, so magnesium forms the negative electrode. At the negative electrode, magnesium is oxidised (loses electrons): Mg(s) → Mg²⁺(aq) + 2e⁻. The electrons flow through the external circuit to the copper electrode, creating an electric current.
6. Explain why hydrogen produced by electrolysis of water using renewable electricity is considered more environmentally friendly than hydrogen produced by steam reforming of methane.
When hydrogen is produced by electrolysis using renewable electricity (from wind, solar, or hydroelectric power), no fossil fuels are burned and no CO₂ is released at any stage of the process. The overall process is carbon-neutral. In contrast, steam reforming of methane (CH₄ + H₂O → CO + 3H₂) produces carbon monoxide, which is then converted to CO₂. This releases significant amounts of CO₂, a greenhouse gas that contributes to climate change. Therefore, hydrogen from steam reforming is not truly clean even though the fuel cell itself only produces water.
When comparing voltage data from different cells, identify patterns. A greater difference in reactivity between the two metals produces a higher voltage. Use data tables to calculate the reactivity difference and correlate it with the measured voltage.
Efficiency = (useful energy output / total energy input) × 100. Hydrogen fuel cells have an efficiency of about 60–70%, compared to petrol engines at about 25–30%. This means fuel cells convert a greater proportion of the chemical energy into useful work, wasting less as heat.
The half equation at the anode shows oxidation (electrons on the product side), e.g. Zn → Zn2+ + 2e−. At the cathode, reduction (electrons on the reactant side), e.g. Cu2+ + 2e− → Cu. The overall equation is obtained by combining half equations so that electrons cancel.
Wrong: Fuel cells burn hydrogen Correct: Fuel cells do not burn hydrogen. They convert chemical energy directly to electrical energy through electrochemical reactions. Hydrogen is oxidised at the anode and oxygen is reduced at the cathode — no combustion takes place
Wrong: Rechargeable batteries last forever Correct: Rechargeable batteries have a limited number of charge-discharge cycles (typically 500-1000 for lithium-ion). Over time, the chemical reactions become less reversible and the capacity decreases until the battery must be replaced
Evaluate the advantages and disadvantages of hydrogen fuel cells compared to petrol engines.
Advantages of hydrogen fuel cells: the only product is water, so there are no greenhouse gases or pollutants at the point of use (no CO2, NOx, or particulates). Fuel cells are more efficient than petrol engines (about 60–70% vs 25–30%), converting more chemical energy into useful work. They have no moving parts, so they are quiet and require less maintenance. [2 marks]
Disadvantages of hydrogen fuel cells: hydrogen is difficult to store and transport — it must be stored under high pressure (up to 700 atm) or at very low temperatures (−253°C), requiring heavy and expensive tanks. The refuelling infrastructure is very limited compared to petrol stations. Most hydrogen is currently produced from fossil fuels (steam reforming of methane), which releases CO2, reducing the environmental benefit. Fuel cells use expensive platinum catalysts. [2 marks]
Overall, hydrogen fuel cells offer significant environmental advantages over petrol engines at the point of use, but the full life cycle must be considered. If hydrogen is produced by electrolysis using renewable electricity, fuel cells are truly clean. If hydrogen comes from fossil fuels, the overall carbon footprint may not be much better than petrol. Petrol engines benefit from established infrastructure and lower upfront costs, but they produce CO2 and other pollutants every time they are used. [2 marks]
A city is choosing between three options to power its new fleet of 50 buses: (A) diesel engines, (B) hydrogen fuel cells, or (C) lithium-ion rechargeable batteries. The city has 3 hydrogen refuelling stations and 200 electric charging points. Average daily route is 180 km. Diesel costs £1.50/L, hydrogen costs £8/kg, electricity costs £0.15/kWh. Evaluate which option is most suitable, considering environmental impact, practicality, and cost.
Option A (diesel): Most practical — existing refuelling infrastructure, lowest upfront cost, proven technology. However, diesel buses produce CO2, NOx and particulates, contributing to air pollution and climate change. Running costs are moderate.
Option B (hydrogen fuel cells): Best environmental performance at point of use — only water emitted. Fast refuelling (5–10 minutes) suits buses with long routes (180 km). However, only 3 hydrogen stations is very limited for 50 buses. Hydrogen is expensive (£8/kg), making running costs high. If hydrogen is from fossil fuels, the environmental benefit is reduced.
Option C (lithium-ion batteries): Zero emissions at point of use. 200 charging points provides adequate infrastructure. Electricity is cheap (£0.15/kWh). However, 180 km daily routes may stretch battery range, especially in cold weather. Charging takes hours, requiring overnight charging or battery swaps. Batteries degrade over time (5–8 years), adding replacement costs.
Recommendation: Option C is most suitable for this city. The 200 charging points exceed the limited 3 hydrogen stations, electricity is cheapest, and the environmental impact is lowest (assuming renewable electricity). Battery range may need careful scheduling but is achievable for 180 km with modern buses.
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