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How is sulfuric acid manufactured, and how do sulfur oxides affect the environment?

Describe the formation of sulfur dioxide and its role in acid rain, outline the Contact process for manufacturing sulfuric acid, and explain the use of sulfur dioxide as a preservative and the methods used to control sulfur emissions

A focused answer to the H2 Chemistry learning outcome on sulfur. The formation of sulfur dioxide and its role in acid rain, the Contact process for making sulfuric acid with its equilibrium reasoning, the use of sulfur dioxide as a preservative, and the control of sulfur emissions by flue-gas desulfurisation.

Generated by Claude Opus 4.89 min answerUpdated 2026-06-06

Reviewed by: AI editorial process; not yet individually human-reviewed

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What this dot point is asking

SEAB wants you to describe how sulfur dioxide forms and its role in acid rain, outline the Contact process for manufacturing sulfuric acid (with the equilibrium reasoning behind the conditions), and explain the use of sulfur dioxide as a preservative and the methods used to control sulfur emissions. The Contact process and the acid-rain and emission-control questions are dependable Paper 2 and Paper 3 content.

The answer

Formation of sulfur dioxide

Fossil fuels contain sulfur impurities. When they burn, the sulfur is oxidised:

\text{S} + \text{O}_2 \rightarrow \text{SO}_2

Sulfur dioxide is also produced when sulfide ores are roasted in metal extraction. It is a colourless, choking, acidic gas.

Sulfur dioxide and acid rain

In the atmosphere, sulfur dioxide is oxidised (a process catalysed by $\text{NO}_2$ and particulates) and dissolves to form acids:

\text{SO}_2 + \text{H}_2\text{O} \rightarrow \text{H}_2\text{SO}_3 \quad\text{(sulfurous acid)}

2\text{SO}_2 + \text{O}_2 \rightarrow 2\text{SO}_3, \qquad \text{SO}_3 + \text{H}_2\text{O} \rightarrow \text{H}_2\text{SO}_4 \quad\text{(sulfuric acid)}

The resulting acid rain lowers the pH of rainwater, damaging limestone buildings, acidifying lakes and soils, and harming plant and aquatic life.

The Contact process

Sulfuric acid is made in three stages:

Make $\text{SO}_2$ : burn sulfur in air, $\text{S} + \text{O}_2 \rightarrow \text{SO}_2$ .
Oxidise to $\text{SO}_3$ (the key equilibrium): $2\text{SO}_2 + \text{O}_2 \rightleftharpoons 2\text{SO}_3$ , exothermic, over a vanadium(V) oxide catalyst.
Absorb $\text{SO}_3$ : dissolve in concentrated sulfuric acid to form oleum, then add water: $\text{SO}_3 + \text{H}_2\text{SO}_4 \rightarrow \text{H}_2\text{S}_2\text{O}_7$ , then $\text{H}_2\text{S}_2\text{O}_7 + \text{H}_2\text{O} \rightarrow 2\text{H}_2\text{SO}_4$ . (Direct addition of $\text{SO}_3$ to water makes an uncontrollable acid mist.)

Conditions for step 2: about $450$ degrees Celsius (a compromise between yield, favoured by low temperature because the reaction is exothermic, and rate, favoured by high temperature), about $2$ atm (the yield is already high so high pressure is not worth the cost), and a $\text{V}_2\text{O}_5$ catalyst.

Sulfur dioxide as a preservative

Sulfur dioxide and sulfites are used to preserve foods and wine because they inhibit the growth of bacteria and moulds and act as antioxidants, preventing browning. Their use is regulated because some people are sensitive to sulfites.

Controlling sulfur emissions

Flue-gas desulfurisation removes $\text{SO}_2$ from power-station emissions by passing the flue gases through a spray or slurry of a base, usually calcium oxide or calcium carbonate:

\text{CaO} + \text{SO}_2 \rightarrow \text{CaSO}_3

The calcium sulfite can be oxidised to calcium sulfate (gypsum) and sold, turning a pollutant into a useful product. Using low-sulfur fuels also reduces emissions at source.

Worked example

Explain why, in the Contact process, sulfur trioxide is absorbed into concentrated sulfuric acid rather than directly into water, and write the equations.

Adding $\text{SO}_3$ directly to water is highly exothermic and produces a dense, uncontrollable mist of sulfuric acid droplets that is hard to condense and wasteful.

Instead, $\text{SO}_3$ is absorbed into concentrated sulfuric acid to form oleum (disulfuric acid):

\text{SO}_3 + \text{H}_2\text{SO}_4 \rightarrow \text{H}_2\text{S}_2\text{O}_7

The oleum is then diluted carefully with water to give sulfuric acid:

\text{H}_2\text{S}_2\text{O}_7 + \text{H}_2\text{O} \rightarrow 2\text{H}_2\text{SO}_4

This two-step absorption controls the reaction and avoids the acid mist.

Try it: Contact process explorer - vary temperature and pressure to see the predicted effect on the $\text{SO}_3$ yield.

Common traps

Adding $\text{SO}_3$ to water in the Contact process: It is absorbed in concentrated sulfuric acid first to avoid an acid mist, then diluted.
Forgetting the compromise reasoning: State that the moderate temperature balances yield (low $T$ ) against rate (high $T$ ), not simply "high temperature".
Confusing sulfur dioxide and nitrogen dioxide in acid rain: $\text{SO}_2$ gives sulfurous and sulfuric acids; $\text{NO}_2$ gives nitric acid (though $\text{NO}_2$ helps catalyse $\text{SO}_2$ oxidation).
Saying high pressure is used in the Contact process: Only a slightly raised pressure (about $2$ atm) is used because the yield is already high; high pressure is not cost-effective.
Naming the wrong catalyst: The Contact process uses vanadium(V) oxide, $\text{V}_2\text{O}_5$ , not iron.

Examples in context

Example 1. Why sulfuric acid is the most-produced industrial chemical. Sulfuric acid is central to fertiliser manufacture, metal processing and many industrial syntheses, which is why its annual production is sometimes used as an economic indicator. SEAB frames the Contact process not just as equations but as an optimisation problem where conditions are chosen for economic as well as chemical reasons.

Example 2. Turning a pollutant into gypsum. Modern coal-fired power stations capture $\text{SO}_2$ with limestone and oxidise the product to calcium sulfate, which is sold as gypsum for plasterboard. This is a textbook example of green chemistry that SEAB uses to test both the desulfurisation reaction and the idea of by-product valorisation.

Try this

Q1. Write the equation for the key equilibrium step of the Contact process and name the catalyst. [2 marks]

Cue. $2\text{SO}_2 + \text{O}_2 \rightleftharpoons 2\text{SO}_3$ ; catalyst vanadium(V) oxide, $\text{V}_2\text{O}_5$ .

Q2. Explain why sulfur dioxide is added to some foods and drinks. [2 marks]

Cue. It acts as a preservative (inhibits bacteria and moulds) and an antioxidant (prevents browning).

Q3. Write the equation for the removal of sulfur dioxide by calcium oxide in flue-gas desulfurisation. [1 mark]

Cue. $\text{CaO} + \text{SO}_2 \rightarrow \text{CaSO}_3$ .

Exam-style practice questions

Practice questions written in the style of SEAB exam questions on this dot point, with worked answer explainers. The year tag is the paper they imitate, not the source.

Specimen (9729)4 marksOutline the three main reactions of the Contact process for the manufacture of sulfuric acid, and state the conditions used for the key equilibrium step and why they are chosen.

Show worked answer →

Step 1: burn sulfur (or roast sulfide ore) in air to form sulfur dioxide.
$\text{S} + \text{O}_2 \rightarrow \text{SO}_2$ .

Step 2 (key equilibrium): oxidise sulfur dioxide to sulfur trioxide over a vanadium(V) oxide catalyst.
$2\text{SO}_2 + \text{O}_2 \rightleftharpoons 2\text{SO}_3$ , exothermic.

Conditions: about 2 atm (slightly above atmospheric), about 450 degrees Celsius, V2O5 catalyst. A modest temperature is a compromise between yield (favoured by low T, exothermic) and rate (favoured by high T). Pressure is only slightly raised because the yield is already high at 2 atm, so high pressure is not worth the cost.

Step 3: absorb SO3 in concentrated sulfuric acid to form oleum, then add water to give sulfuric acid.
$\text{SO}_3 + \text{H}_2\text{SO}_4 \rightarrow \text{H}_2\text{S}_2\text{O}_7$ ; then $\text{H}_2\text{S}_2\text{O}_7 + \text{H}_2\text{O} \rightarrow 2\text{H}_2\text{SO}_4$ .

Markers reward the three steps with equations, the catalyst and temperature, and the compromise reasoning for the conditions.

2022 (style)3 marksExplain how sulfur dioxide is formed when fossil fuels are burned and how it contributes to acid rain. State one method used to remove sulfur dioxide from power-station emissions.

Show worked answer →

Fossil fuels contain sulfur impurities. When they burn, the sulfur is oxidised: S + O2 -> SO2.

Sulfur dioxide reacts with water and oxygen in the atmosphere to form sulfurous and then sulfuric acid, lowering the pH of rain. This acid rain damages buildings (especially limestone), acidifies lakes and soils, and harms plant and aquatic life.

Removal method (flue-gas desulfurisation): the SO2 is passed through a slurry or spray of calcium oxide or calcium carbonate, which reacts with the acidic SO2 to form calcium sulfite (or sulfate), removing it from the flue gases.

Markers reward the oxidation of sulfur impurities, the formation of acid in rain with effects, and a valid desulfurisation method.