Sulfur Dioxide Removal Revision Notes | A Level AQA Chemistry

1. Why Sulphur dioxide Is a Problem

Fossil fuels — coal, oil, and natural gas — are not pure hydrocarbons. They contain sulphur impurities, and when those fuels burn in power stations, the sulphur oxidises to produce sulphur dioxide ($SO_2$). That gas rises up the flue (chimney), and if nothing stops it, it enters the atmosphere.

Once in the atmosphere, $SO_2$ dissolves in water droplets to form sulphurous acid, which can then oxidise further to sulphuric acid. Additionally, $SO_2$ can be oxidised in the atmosphere to $SO_3$, which also reacts with water to form sulphuric acid directly. The result is acid rain — precipitation with a pH low enough to damage ecosystems, corrode buildings, and harm human respiratory systems. This is why removing $SO_2$ before it leaves the flue is a legal and environmental priority.

The combustion equation for a sulphur impurity is straightforward: $S(s) + O_2(g) \rightarrow SO_2(g)$. The sulphur burns completely in excess oxygen, releasing $SO_2$ directly into the flue gas stream alongside $CO_2$, water vapour, and nitrogen.

2. The Chemistry of Removal: Acid-Base Neutralisation

The key insight for understanding both removal methods is that $SO_2$ is an acidic oxide. It dissolves in water to form an acid, and it reacts directly with bases. Both calcium oxide and calcium carbonate are basic — they can neutralise the acidic $SO_2$. That single principle underpins everything that follows.

Think of it this way: you already know that acids react with bases to form a salt and water. Here, the 'acid' is the gaseous $SO_2$ and the 'base' is either $CaO$ or $CaCO_3$. The product in both cases is calcium sulfite ($CaSO_3$), a solid that can be collected and disposed of or used in industry.

2a. Removal Using Calcium Oxide (CaO)

Calcium oxide — also called quicklime — is a white solid produced by heating calcium carbonate strongly (thermal decomposition). When $CaO$ is mixed with the flue gases, it reacts directly with $SO_2$ in a straightforward acid-base reaction.

The equation is: $CaO(s) + SO_2(g) \rightarrow CaSO_3(s)$. Calcium oxide acts as the base, $SO_2$ acts as the acidic oxide, and the product is solid calcium sulfite. No water is produced here because this is a direct oxide-oxide reaction rather than a full neutralisation in aqueous solution.

In practice, powdered $CaO$ is injected into the flue gas stream or used in a scrubber unit. The solid $CaSO_3$ is then collected from the base of the scrubber. The reaction is efficient because $CaO$ is highly reactive — the O2- ion is a strong base with high charge density, making CaO highly reactive toward acidic oxides.

2b. Removal Using Calcium Carbonate ($CaCO_3$)

Calcium carbonate (limestone) is cheaper and more abundant than calcium oxide, making it the more widely used industrial option. The reaction with $SO_2$ produces calcium sulfite and carbon dioxide: $CaCO_3(s) + SO_2(g) \rightarrow CaSO_3(s) + CO_2(g)$.

Notice that this reaction releases $CO_2$ as a by-product. That is a trade-off worth acknowledging: you remove one pollutant ($SO_2$) but generate a greenhouse gas ($CO_2$) in the process. At A-Level, examiners may ask you to evaluate this, so keep it in mind.

In the wet scrubbing process used at many power stations, a slurry of calcium carbonate in water is sprayed into the flue gas. The $SO_2$ dissolves into the water and reacts with the $CaCO_3$. The $CaSO_3$ produced can be further oxidised to calcium sulphate ($CaSO_4$), which is gypsum — a commercially useful material used in plasterboard manufacture.

3. Why These Reactions Work: Deeper Reasoning

Both reactions work because $SO_2$ is an acidic oxide and both $CaO$ and $CaCO_3$ are basic. Calcium oxide is a metal oxide and therefore a base; it reacts with acidic oxides directly. Calcium carbonate contains the carbonate ion ($CO_3^{2-}$), which is a base — it accepts a proton from the acidic $SO_2$ (via the intermediate sulphurous acid), releasing $CO_2$ and forming $CaSO_3$.

At A-Level, you should be able to justify the choice of reagent in terms of both chemistry and economics. Calcium carbonate is preferred industrially because it is cheaper and does not require a prior energy-intensive step to produce. Calcium oxide requires thermal decomposition of limestone first ($CaCO_3 \rightarrow CaO + CO_2$), which itself releases $CO_2$ and consumes energy — adding cost and additional emissions before the scrubbing even begins.

4. Exam Technique and Mark-Winning Responses

AQA questions on this topic typically use the command words 'explain' or 'evaluate'. An 'explain' question wants you to give a reason — not just state what happens, but say why. For example, 'Explain why calcium oxide can be used to remove $SO_2$' requires you to identify $SO_2$ as an acidic oxide and $CaO$ as a basic oxide, then state that they react in an acid-base neutralisation.

An 'evaluate' question on the choice between $CaO$ and $CaCO_3$ demands a balanced argument. You must present evidence on both sides and reach a conclusion. A response that only lists advantages of one reagent will not access the top mark band. Structure your answer as: advantage of $CaCO_3$ (cost, availability) → disadvantage of $CaCO_3$ ($CO_2$ released) → advantage of $CaO$ (no $CO_2$ by-product) → disadvantage of $CaO$ (energy cost of production) → conclusion.

Worked example: Constructing an Explain Answer: $CaCO_3$ Removing $SO_2$

Explain why calcium carbonate can be used to remove sulphur dioxide from flue gases. Include a balanced equation. [3 marks]

1. $SO_2$ is an acidic oxide; $CaCO_3$ is a base (contains $CO_3^{2-}$, a basic ion). — [B1] Identifying $SO_2$ as acidic and $CaCO_3$ as basic is the core chemical reasoning the examiner wants.
2. The two species undergo an acid-base reaction (neutralisation of an acidic oxide by a carbonate). — [B1] Naming the reaction type demonstrates understanding of why the reaction occurs, not just that it does.
3. $CaCO_3(s) + SO_2(g) \rightarrow CaSO_3(s) + CO_2(g)$ — [B1] The equation must be balanced and include state symbols. Check: Ca: 1=1, C: 1=1, O: 5=5, S: 1=1. Balanced.

Answer: $SO_2$ is an acidic oxide and $CaCO_3$ is a base, so they react in an acid-base reaction. The equation is $CaCO_3(s) + SO_2(g) \rightarrow CaSO_3(s) + CO_2(g)$.

Both removal methods ultimately serve the same purpose: preventing acidic $SO_2$ from reaching the atmosphere and causing acid rain. The choice between $CaO$ and $CaCO_3$ is driven by economics and practicality, but the underlying chemistry — acid-base neutralisation of an acidic oxide — is identical in principle. Mastering that core idea means you can answer any question on this topic, regardless of how the examiner frames it.

Bridge: flue-gas desulphurisation is not just an equilibrium idea; it is an industrial acid-base process designed to remove acidic oxides before they enter the atmosphere.