Sulfur: the element
Sulfur is a non-metal. Its symbol is S and its atomic number is 16, so its electrons are arranged 2, 8, 6. It sits in group 16 (the oxygen family, also called the chalcogens: O, S, Se, Te, Po) and period 3.
- It is a yellow, brittle solid that does not dissolve in water but dissolves in carbon disulfide.
- It is found free near volcanoes and in compounds such as metal sulfides and gypsum (CaSO₄·2H₂O). Most sulfur today is recovered from crude oil and natural gas.
- Common oxidation states: −2 (H₂S), +4 (SO₂), +6 (SO₃, H₂SO₄).
Allotropes
Allotropes are different forms of the same element. Solid sulfur is made of S₈ rings (crown shape). Rhombic sulfur is stable at room temperature. Above about 96 °C the rings pack differently to give needle-like monoclinic sulfur. Pouring boiling sulfur into cold water gives rubbery plastic sulfur made of long chains.
Sulfur dioxide and hydrogen sulfide
Sulfur dioxide (SO₂) forms when sulfur or sulfide ores burn: S + O₂ → SO₂. It is a colourless gas with a choking smell and a bent shape.
- Acidic oxide: SO₂ + H₂O ⇌ H₂SO₃ (sulfurous acid). It turns blue litmus red.
- Bleaching: moist SO₂ removes colour by taking oxygen away (reduction). The colour can slowly come back in air.
- Reducing agent: it turns purple acidified potassium permanganate colourless. This is a lab test for SO₂.
- Uses: making sulfuric acid, preserving dried fruit and juices, bleaching wool and paper.
- Harm: SO₂ from burning coal and oil makes acid rain, which damages lungs, plants, lakes and limestone or marble.
Hydrogen sulfide (H₂S) is a poisonous gas that smells of rotten eggs. It is made when sulfides react with acid: FeS + 2HCl → FeCl₂ + H₂S. It turns lead acetate paper black (PbS) and acts as a reducing agent.
Sulfuric acid and the contact process
Sulfuric acid is made in the contact process:
- Burn sulfur: S + O₂ → SO₂.
- Oxidise: 2SO₂ + O₂ ⇌ 2SO₃. This needs a V₂O₅ catalyst, about 450 °C and 1–2 atm. It is reversible, so the conditions are a balance between fast rate and good yield.
- Absorb SO₃ in concentrated acid to form oleum: SO₃ + H₂SO₄ → H₂S₂O₇. (SO₃ is not added straight to water because it makes a mist of acid that is hard to collect.)
- Dilute: H₂S₂O₇ + H₂O → 2H₂SO₄.
Properties of sulfuric acid
- Strong acid when dilute: reacts with metals (Zn + H₂SO₄ → ZnSO₄ + H₂), bases and carbonates.
- Dehydrating agent when concentrated: it pulls water out of sugar, leaving black carbon (C₁₂H₂₂O₁₁ → 12C + 11H₂O). It also dries gases.
- Oxidising agent when hot and concentrated: it reacts with copper to give SO₂.
- Safety: mixing with water gives out a lot of heat. Always add acid to water, slowly, never water to acid.
Uses: fertilisers, detergents, dyes, car batteries, cleaning metals.
Testing for sulfate ions
To check a solution for sulfate (SO₄²⁻):
- Add dilute hydrochloric acid (this removes carbonate and sulfite ions, which could also give a white solid).
- Add barium chloride (or barium nitrate) solution.
- A white precipitate of barium sulfate shows sulfate: Ba²⁺ + SO₄²⁻ → BaSO₄↓.
BaSO₄ does not dissolve in acid. Doctors even give it as a 'barium meal' for X-rays because it is so insoluble that it is not absorbed.
How sulfur compounds change into each other
Follow the oxidation number of sulfur:
H₂S (−2) → S (0) → SO₂ (+4) → SO₃ (+6) → H₂SO₄ (+6) → sulfate salts (+6)
- Moving right means oxidation (adding oxygen or losing hydrogen).
- 2H₂S + SO₂ → 3S + 2H₂O: here H₂S is oxidised and SO₂ is reduced, both to sulfur.
- Metal sulfides roasted in air give SO₂: 2ZnS + 3O₂ → 2ZnO + 2SO₂.
Try it: Sprinkle a spoon of sugar in a glass dish. Do not add acid at home! Instead, watch the 3D (step 4) and the table here, then predict: what is left when all the water is taken out of sugar? (Answer: black carbon.)
Key formulas and definitions
- S + O₂ → SO₂
- 2SO₂ + O₂ ⇌ 2SO₃ (V₂O₅, 450 °C)
- SO₃ + H₂SO₄ → H₂S₂O₇; H₂S₂O₇ + H₂O → 2H₂SO₄
- SO₂ + H₂O ⇌ H₂SO₃
- BaCl₂ + Na₂SO₄ → BaSO₄↓ (white) + 2NaCl
- FeS + 2HCl → FeCl₂ + H₂S
- C₁₂H₂₂O₁₁ → 12C + 11H₂O (conc. H₂SO₄)
Worked examples
1. Write the electron arrangement of sulfur and say why it is in group 16.
Z = 16, so 2, 8, 6. It has 6 electrons in its outer shell, the same as oxygen (2, 6), so both are in group 16.
2. What is the oxidation number of S in H₂SO₄?
H is +1, O is −2. 2(+1) + S + 4(−2) = 0 → S − 6 = 0 → S = +6.
3. How many moles of SO₂ form when 64 g of sulfur burns completely? (S = 32 g/mol)
Moles of S = 64 ÷ 32 = 2 mol. S + O₂ → SO₂ is 1 : 1, so 2 mol SO₂ (= 2 × 64 = 128 g).
4. A white solid forms when barium chloride is added to a solution. Is it surely sulfate?
Not yet. Carbonate also gives a white solid (BaCO₃). Add dilute HCl: BaCO₃ dissolves with fizzing, BaSO₄ stays. If it stays, the ion is sulfate.
5. Why is SO₃ dissolved in sulfuric acid, not in water, in the contact process?
SO₃ reacts with water very violently and forms a fine acid mist that escapes. Dissolving it in H₂SO₄ to form oleum is smooth; the oleum is then diluted safely.
6. How much sulfuric acid can be made from 1.6 tonnes of sulfur? (S = 32, H₂SO₄ = 98)
Each S atom ends up in one H₂SO₄. 32 t S → 98 t acid, so 1.6 t → 1.6 × 98 ÷ 32 = 4.9 tonnes of H₂SO₄.
Common mistakes
- Adding water to concentrated acid. It boils and spits. Always add acid to water.
- Thinking the bleaching by SO₂ is the same as by chlorine. SO₂ bleaches by reduction and the colour can return; chlorine bleaches by oxidation and is permanent.
- Forgetting to add HCl before BaCl₂ in the sulfate test, so carbonate is mistaken for sulfate.
- Writing S for the solid but forgetting it is really S₈ rings in allotropes.