Sulfur and Compounds

Summary: Sulfur sources: natural (volcanoes) and fossil fuels. SO₂ produced by burning sulfur or roasting sulfide ores. Contact process: SO₂ + O₂ ⇌ SO₃ using V₂O₅ catalyst, then SO₃ + H₂SO₄ → oleum → H₂SO₄. Tags: igcse chemistry environmental-chem Created: 2026-07-14 Last Updated: 2026-07-16

Occurrence and Sources of Sulfur

Sulfur is a yellow, non-metallic solid at room temperature. It is found in the Earth’s crust and atmosphere in both elemental and combined forms.

Natural Sources

• Elemental sulfur deposits: Found in volcanic regions (e.g., Sicily, Indonesia, Japan). Sulfur is extracted from underground deposits using the Frasch process — superheated water is pumped underground to melt the sulfur, and compressed air forces the molten sulfur to the surface.
• Volcanic emissions: Volcanoes release sulfur dioxide (SO₂) and hydrogen sulfide (H₂S) gases directly into the atmosphere.
• Sulfide ores: Many metal ores exist as sulfides, e.g., zinc blende (ZnS), galena (PbS), iron pyrite (FeS₂, also known as ‘fool’s gold’), copper pyrite (CuFeS₂).

Industrial Sources

• Fossil fuels: Coal, crude oil, and natural gas all contain sulfur impurities (typically 0.5–5% by mass). When these fuels are burned without treatment, the sulfur is released as SO₂, contributing to acid rain.
• By-product of metal extraction: When sulfide ores are roasted to extract metals, SO₂ is produced as a by-product. Rather than releasing it, it is often used to manufacture sulfuric acid.

Production of Sulfur Dioxide (SO₂)

Burning Sulfur in Air

When sulfur is burned in air or oxygen, it produces sulfur dioxide gas:

S(s) + O₂(g) → SO₂(g)

Observations: Sulfur burns with a blue flame and produces a colourless, choking gas with a pungent, suffocating smell. Sulfur dioxide is denser than air.

Roasting Sulfide Ores

When sulfide ores are heated strongly in air (roasted), the sulfide is converted to the metal oxide and sulfur dioxide is released:

2ZnS(s) + 3O₂(g) → 2ZnO(s) + 2SO₂(g)

Example: Zinc blende (ZnS) is roasted to produce zinc oxide and sulfur dioxide. The SO₂ produced is collected and used in the Contact Process for manufacturing sulfuric acid, rather than being vented to the atmosphere.

2PbS(s) + 3O₂(g) → 2PbO(s) + 2SO₂(g)
4FeS₂(s) + 11O₂(g) → 2Fe₂O₃(s) + 8SO₂(g)

Properties of Sulfur Dioxide

SO₂ as a Reducing Agent

Sulfur dioxide is a reducing agent (it is itself oxidised). In aqueous solution, SO₂ can reduce:

• Acidified potassium manganate(VII): Purple MnO₄⁻ is reduced to colourless Mn²⁺. The purple colour is decolourised.
• Acidified potassium dichromate(VI): Orange Cr₂O₇²⁻ is reduced to green Cr³⁺.
• Iron(III) ions to iron(II) ions: 2Fe³⁺ + SO₂ + 2H₂O → 2Fe²⁺ + SO₄²⁻ + 4H⁺

SO₂ as a Bleach

Sulfur dioxide acts as a bleach by reduction (unlike chlorine, which bleaches by oxidation). It removes colour from substances such as wood pulp and is used to bleach paper. The bleaching action is often reversible — the colour may slowly return upon exposure to air.

The Contact Process: Manufacture of Sulfuric Acid

Sulfuric acid (H₂SO₄) is the most important industrial chemical and is manufactured by the Contact Process. This is the key exam topic for sulfur chemistry.

Overview

The Contact Process converts sulfur dioxide (SO₂) into sulfuric acid (H₂SO₄) in three main stages:

Stage 1: Production of Sulfur Trioxide (SO₃)

Sulfur dioxide is oxidised to sulfur trioxide in a reversible, exothermic reaction:

2SO₂(g) + O₂(g) ⇌ 2SO₃(g)     ΔH = −197 kJ mol⁻¹

Conditions used:

• Catalyst: Vanadium(V) oxide (V₂O₅) — provides an alternative pathway with lower activation energy
• Temperature: Approximately 450 °C
• Pressure: Approximately 2 atmospheres (just above atmospheric pressure)

Why these specific conditions?

Temperature (compromise condition):

• The forward reaction is exothermic, so a low temperature favours a high equilibrium yield of SO₃ (Le Chatelier’s principle)
• However, at low temperatures the rate of reaction is too slow to be economically viable
• 450 °C is a compromise: high enough for an acceptable rate, while still giving a reasonable equilibrium yield (about 97–98% conversion under these conditions)

Catalyst:

• Vanadium(V) oxide (V₂O₅) lowers the activation energy, increasing the rate of reaction
• The catalyst does not affect the position of equilibrium — it only allows equilibrium to be reached more quickly
• The catalyst is not consumed in the reaction

Pressure (just above atmospheric):

• The forward reaction: 3 moles of gas (2 SO₂ + 1 O₂) produces 2 moles of gas (2 SO₃)
• Fewer moles on the product side means higher pressure would favour the forward reaction (Le Chatelier’s principle)
• However, at 2 atm, the equilibrium yield is already high enough (~97–98%) that the extra cost of high-pressure equipment is not justified
• Operating at just above 1 atm keeps the gases flowing through the plant

Stage 2: Formation of Oleum

Sulfur trioxide is not added directly to water because the reaction is too violent and exothermic, producing a dangerous, corrosive mist of sulfuric acid that is difficult to condense. Instead, SO₃ is absorbed in concentrated sulfuric acid to form oleum (H₂S₂O₇):

SO₃(g) + H₂SO₄(l) → H₂S₂O₇(l)

Oleum is a fuming liquid and can be safely transported. It is safer to handle than SO₃.

Stage 3: Dilution of Oleum

Oleum is carefully diluted with water to produce concentrated sulfuric acid:

H₂S₂O₇(l) + H₂O(l) → 2H₂SO₄(aq)

The overall reaction of stages 2 and 3 is equivalent to: SO₃ + H₂O → H₂SO₄, but the two-stage process avoids the violent, dangerous direct reaction.

Properties of Sulfuric Acid (H₂SO₄)

Sulfuric acid is a strong dibasic acid — it donates two protons (H⁺) per molecule when it dissociates in water:

H₂SO₄(aq) → 2H⁺(aq) + SO₄²⁻(aq)

Acidic Properties

As a typical strong acid, dilute H₂SO₄:

• Reacts with bases to form sulfates: H₂SO₄ + CuO → CuSO₄ + H₂O
• Reacts with carbonates to form sulfates, CO₂, and water: H₂SO₄ + Na₂CO₃ → Na₂SO₄ + CO₂ + H₂O
• Reacts with reactive metals to form sulfates and hydrogen gas: H₂SO₄ + Zn → ZnSO₄ + H₂
• Turns blue litmus red

Dehydrating Agent

Concentrated sulfuric acid is a powerful dehydrating agent — it removes water (or the elements of water, H and O in a 2:1 ratio) from compounds.

Dehydration of carbohydrates:

• Concentrated H₂SO₄ removes water from sugar (C₁₂H₂₂O₁₁), leaving a black mass of carbon:

  C₁₂H₂₂O₁₁(s) → 12C(s) + 11H₂O(l)   [catalysed by H₂SO₄]
  

• This is an exothermic reaction — the mixture heats up and steam is produced

Dehydration of hydrated salts:

• Blue hydrated copper(II) sulfate crystals (CuSO₄·5H₂O) turn white (anhydrous CuSO₄) when concentrated H₂SO₄ removes the water of crystallisation

Oxidising Agent

Hot, concentrated sulfuric acid is an oxidising agent. It can oxidise:

• Bromide ions to bromine: 2HBr + H₂SO₄ → Br₂ + SO₂ + 2H₂O
• Iodide ions to iodine (and further to H₂S)
• Metals such as copper (which does not react with dilute H₂SO₄): Cu + 2H₂SO₄ → CuSO₄ + SO₂ + 2H₂O

Uses of Sulfur Dioxide

Sulfur dioxide has several important applications separate from its role in sulfuric acid production:

• Bleaching wood pulp for paper production — SO₂ reduces coloured compounds in the pulp
• Food preservative: SO₂ kills bacteria and fungi, and prevents oxidation that causes browning. Used in dried fruits (apricots, raisins), fruit juices, and wine. Labeled as E220 on food packaging
• Water treatment: SO₂ is used to remove excess chlorine from treated water
• Production of calcium hydrogen sulfite (Ca(HSO₃)₂) for paper making

Key Concepts from Past Papers

• add water to anhydrous cobalt sulfate / add water to CoSO4
• colour changes (from blue) to red / pink
• any 2 from:
• improve growth of plants
• increase protein (in plants)
• fertilisers add nitrogen / nitrates / phosphorous / phosphates / potassium
• to put back nitrogen / nitrates / phosphorous / phosphates / potassium into the soil
• sulfuric acid + magnesium carbonate → magnesium sulfate + carbon dioxide + water IF full credit is not awarded, for either magnesium sulfate OR carbon dioxide + water

Keywords from Past Papers

sulfate, dioxide, sulfur, ammonia, water, sulfuric, acid, ammonium, oxygen, add, magnesium, plants, nitrogen, nitrates, phosphorous

Related Notes

• Uses of Sulfur Compounds
• Acids and Bases
• Making Salts

Sources

• OpenStax Chemistry 2e, Chapter 18 (Representative Metals, Metalloids, and Nonmetals) — Section 18.10 on Occurrence, Preparation, and Properties of Sulfur — openstax.org/books/chemistry-2e/
• BBC Bitesize GCSE Chemistry, Sulfur and Sulfuric Acid — bbc.co.uk/bitesize/topics/
• Cambridge IGCSE Chemistry 0620 Syllabus, Topic 10: Chemistry of the Environment
• CK-12 Chemistry, Sulfur and the Contact Process — ck12.org/chemistry/

Past Paper Sources

• 0620/31 May/June 2015: Q66(b)(i) (0m), Q66(b)(i) (0m)
• 0620/32 Feb/March 2015: Q55(d)(ii) (2m)
• 0620/32 Feb/March 2018: Q66(b)(i) (1m), Q66(b)(iii) (1m), Q77(b)(iii) (1m)
• 0620/32 Feb/March 2019: Q33(c)(ii) (2m), Q33(c)(ii) (2m)
• 0620/32 Feb/March 2022: Q44(a)(iii) (2m), Q22(c)(ii) (1m)
• 0620/32 Feb/March 2023: Q22(b)(ii) (1m)
• 0620/32 May/June 2018: Q11(a)(ii) (1m)
• 0620/32 May/June 2020: Q22(a)(i) (1m), Q22(a)(ii) (1m), Q22(a)(iv) (0m)
• 0620/33 May/June 2016: Q77(b)(i) (1m)
• 0620/33 May/June 2022: Q22(a)(ii) (1m)
• 0620/33 May/June 2023: Q33(b)(ii) (1m)
• 0620/33 October/November 2015: Q55(b)(ii) (1m), Q55(b)(iv) (1m)

Common Misconceptions