Fuel Cells: Hydrogen-Oxygen and Societal Impact Revision Notes | A Level AQA Chemistry

1. What Makes a Fuel Cell Different?

Picture a conventional battery: it stores a fixed amount of chemical energy, and once that energy is used, you either recharge it or throw it away. A fuel cell works on a fundamentally different principle — it is an open system that converts chemical energy directly into electrical energy as long as fuel is continuously supplied.

This distinction is the first thing AQA wants you to nail. Fuel cells do not need to be electrically recharged because the reactants — hydrogen and oxygen — are fed in continuously from an external source. The cell itself is not depleted; it is the fuel supply that determines how long the cell operates.

The key contrast worth holding in your mind: a rechargeable battery stores reactants internally and must be recharged when they are consumed; a fuel cell is more like a combustion engine in that it needs a continuous fuel supply, but unlike an engine it produces electricity directly rather than via heat and mechanical work. That direct conversion is what makes fuel cells so efficient in principle.

2. Inside the Alkaline Hydrogen-Oxygen Fuel Cell

Now that you know what a fuel cell is, let's look at what is actually happening at each electrode. The AQA specification focuses on the alkaline version, where the electrolyte is concentrated aqueous potassium hydroxide, $\text{KOH}_{(aq)}$. The alkaline environment means hydroxide ions, $\text{OH}^-_{(aq)}$, are the charge carriers moving through the electrolyte.

Hydrogen gas is fed to the anode (the negative electrode), and oxygen gas is fed to the cathode (the positive electrode). Both electrodes are typically made from a porous material — often platinum or a platinum-group catalyst — that allows gas to contact the electrolyte at a three-phase boundary: gas, liquid electrolyte, and solid electrode.

The hydroxide ions regenerated at the cathode migrate through the electrolyte to the anode.

2a. The Electrode Half-Equations

These half-equations are the core of what AQA will ask you to write or complete. Learn them precisely — state symbols included.

To obtain the overall cell equation, double the anode equation (so that electrons cancel) and add it to the cathode equation. The $\text{OH}^-$ ions and $\text{H}_2\text{O}$ terms simplify, leaving the clean overall reaction.

Worked example: Deriving the Overall Cell Equation from Half-Equations

Show how the two electrode half-equations for an alkaline hydrogen-oxygen fuel cell combine to give the overall cell equation.

1. $$\text{H}_{2(g)} + 2\text{OH}^-_{(aq)} \rightarrow 2\text{H}_2\text{O}_{(l)} + 2e^-$$ — Hydrogen is oxidised at the anode. Two electrons are released per hydrogen molecule.
2. $$\text{O}_{2(g)} + 2\text{H}_2\text{O}_{(l)} + 4e^- \rightarrow 4\text{OH}^-_{(aq)}$$ — Oxygen is reduced at the cathode. Four electrons are consumed per oxygen molecule.
3. $$2\text{H}_{2(g)} + 4\text{OH}^-_{(aq)} \rightarrow 4\text{H}_2\text{O}_{(l)} + 4e^-$$ — Multiplying the anode equation by 2 gives 4 electrons on each side, so they cancel when the equations are added.
4. $$2\text{H}_{2(g)} + 4\text{OH}^-_{(aq)} + \text{O}_{2(g)} + 2\text{H}_2\text{O}_{(l)} + 4e^- \rightarrow 4\text{H}_2\text{O}_{(l)} + 4e^- + 4\text{OH}^-_{(aq)}$$ — Combine both sides. Electrons appear on both sides and will cancel. Hydroxide ions and water also appear on both sides.
5. Cancel $4e^-$, $4\text{OH}^-_{(aq)}$, and $2\text{H}_2\text{O}_{(l)}$ from each side. — After cancellation: $2\text{H}_{2(g)} + \text{O}_{2(g)} \rightarrow 2\text{H}_2\text{O}_{(l)}$. The electrolyte species are regenerated and do not appear in the overall equation.

Answer: $$2\text{H}_{2(g)} + \text{O}_{2(g)} \rightarrow 2\text{H}_2\text{O}_{(l)}$$

3. Oxidation States and Electrode Roles

It is worth pausing to confirm the oxidation and reduction assignments, because examiners frequently ask you to identify which electrode is which and why. At the anode, hydrogen is oxidised: the oxidation state of hydrogen in $\text{H}_2$ is 0, and in $\text{H}_2\text{O}$ it is +1. Oxidation state increases — that is oxidation. Electrons are released.

At the cathode, oxygen is reduced: the oxidation state of oxygen in $\text{O}_2$ is 0, and in $\text{OH}^-$ it is −2. Oxidation state decreases — that is reduction. Electrons are consumed. The mnemonic OIL RIG (Oxidation Is Loss, Reduction Is Gain of electrons) applies here exactly as it does in any redox context.

4. Benefits and Risks to Society

The third specification requirement asks you to evaluate the societal impact of hydrogen-oxygen fuel cells. This is a classic AQA evaluate question: you need to weigh genuine advantages against real limitations and reach a supported conclusion. Let's build the evidence base you need.

4a. Benefits

• Zero harmful emissions at point of use: The only product of the cell reaction is water, $\text{H}_2\text{O}_{(l)}$. No carbon dioxide, no nitrogen oxides, no particulates are emitted during operation. This is a significant advantage over internal combustion engines in urban air quality.
• High efficiency: Fuel cells convert chemical energy directly to electrical energy without an intermediate heat stage. Conventional combustion engines are limited by the Carnot efficiency ceiling; fuel cells are not subject to the same thermodynamic constraint and can achieve higher practical efficiencies.
• Continuous operation without recharging: Unlike a battery, a fuel cell operates as long as hydrogen and oxygen are supplied. This makes it attractive for applications requiring sustained power — vehicles, remote power stations, and spacecraft.
• Quiet operation: There are no moving parts in the electrochemical conversion process, reducing mechanical noise and vibration compared with combustion engines.
• Scalability: Fuel cell stacks can be scaled from small portable devices to large stationary power plants, offering flexibility across sectors.

4b. Risks and Limitations

• Hydrogen production is currently carbon-intensive: Most hydrogen today is produced by steam reforming of natural gas, which releases $\text{CO}_2$. Unless hydrogen is produced from renewable sources (electrolysis using renewable electricity — so-called 'green hydrogen'), the overall carbon footprint of a fuel cell vehicle may not be significantly lower than a petrol vehicle.
• Hydrogen storage and transport are hazardous: Hydrogen has a very low density and must be stored under high pressure or at cryogenic temperatures. It is highly flammable and has a wide explosive range in air (4–75% by volume). Infrastructure for safe storage and distribution is expensive and not yet widespread.
• High cost of catalysts: The electrodes typically require platinum-group metal catalysts, which are rare, expensive, and subject to supply chain risks. This raises the manufacturing cost of fuel cells significantly.
• Durability and membrane degradation: Components such as the polymer electrolyte membrane (in PEM cells) or the alkaline electrolyte can degrade over time, reducing cell lifetime and increasing maintenance costs.
• Infrastructure gap: The network of hydrogen refuelling stations is currently very limited compared with petrol stations or electric vehicle charging points, restricting practical adoption of fuel cell vehicles.

Notice that many of the risks are not intrinsic to the fuel cell itself but relate to the wider hydrogen economy — how hydrogen is produced, stored, and distributed. This distinction matters in an evaluate question: the cell reaction is clean, but the full lifecycle may not be, depending on the energy source used to generate the hydrogen.

5. Comparing Fuel Cells, Batteries, and Combustion Engines

AQA questions sometimes ask you to compare fuel cells with other energy conversion technologies. The table below consolidates the key distinctions you need for a compare or evaluate response.

One nuance worth emphasising: both fuel cells and rechargeable batteries convert chemical energy directly to electrical energy, which is why both are more efficient than combustion engines. The critical difference is that a battery is a closed system — it stores its reactants internally — whereas a fuel cell is an open system that requires a continuous external fuel supply.

6. Exam Technique and Common Mistakes

Fuel cell questions appear in the physical chemistry section of AQA Paper 2. They tend to combine half-equation writing with a short evaluate or discuss question on societal impact. Here is how to approach each type with precision.

When writing half-equations, the most common error is forgetting to include $\text{OH}^-$ ions in the anode equation or $\text{H}_2\text{O}$ in the cathode equation. These species are essential because the electrolyte is alkaline — they are not optional extras. Always check that your half-equations balance for both atoms and charge before moving on.

Fuel cells sit at the intersection of electrochemistry and contemporary energy policy, which is why AQA pairs the electrode chemistry with a societal evaluation. Master both halves — the precise half-equations and the balanced evaluate argument — and you have covered every mark available on this topic.