Oxygen and sulfur (group VIA)

The Group VIA elements (chalcogens, "ore formers") are oxygen O, sulfur S, selenium Se, tellurium Te and polonium Po. Oxygen is the most abundant element in the Earth’s crust (almost half its mass) and the oxidising agent needed for breathing and burning. Sulfur gives the "bread" of the chemical industry — sulfuric acid — from which fertilisers, dyes, medicines, battery electrolyte and many other substances are made. Exam questions often cover the reactions of concentrated H₂SO₄ with metals and non-metals, the dual behaviour of SO₂ and the industrial production of sulfuric acid.

§1. The group, and oxygen

Down the group non-metallic character weakens: oxygen and sulfur are typical non-metals, while tellurium and polonium show metallic properties. Oxygen has no d-orbitals, so it is only divalent; sulfur in an excited state can have valency IV or VI.

Oxygen is a colourless, odourless gas, slightly soluble in water (enough for fish to breathe), making up 21 % of air by volume. It is released into the atmosphere by photosynthesis.
Preparation: in the laboratory 2KMnO₄ → K₂MnO₄ + MnO₂ + O₂; 2KClO₃ → 2KCl + 3O₂ (MnO₂ catalyst); 2H₂O₂ → 2H₂O + O₂ (MnO₂); industrially by fractional distillation of liquid air.
Properties: oxygen forms oxides with almost all elements (except the noble gases, gold and platinum). Fast oxidation is combustion (light and heat are released); slow oxidation includes respiration, decay and rusting. When compounds burn, the oxides of their elements form: 2H₂S + 3O₂ → 2SO₂ + 2H₂O.

Worked example

24.5 g of potassium chlorate KClO₃ decomposed completely. What volume of oxygen (STP) is released?

1. n(KClO₃) = 24.5/122.5 = 0.2 mol.
2. 2KClO₃ → 3O₂: n(O₂) = 0.2·3/2 = 0.3 mol.
3. V = 0.3·22.4 = 6.72 L.

§2. Ozone and hydrogen peroxide

Ozone O₃ is an allotrope of oxygen, a pale blue gas with a sharp smell. It forms in electric discharges (lightning): 3O₂ → 2O₃. It is a stronger oxidising agent than oxygen: O₃ + 2KI + H₂O → I₂ + 2KOH + O₂ — this reaction detects ozone (starch turns blue). The ozone layer in the stratosphere shields the Earth from the Sun’s harsh ultraviolet radiation. Ozone is used to disinfect water as an alternative to chlorine.

Hydrogen peroxide H₂O₂ contains oxygen in the −1 state, an intermediate one, so it can act as either an oxidising or a reducing agent; it slowly decomposes by itself. A 3 % solution is used in medicine to clean wounds.

§3. Sulfur and its compounds

Sulfur is a yellow crystalline solid that neither dissolves in nor is wetted by water (its powder floats). Allotropes: rhombic (stable), monoclinic and plastic. The crystalline forms consist of ring-shaped S₈ molecules.

• With metals and hydrogen sulfur is an oxidising agent: S + Fe → FeS, S + H₂ → H₂S; with oxygen it is a reducing agent: S + O₂ → SO₂. It reacts with mercury even in the cold, which is why spilt mercury is covered with sulfur powder.
• H₂S is a toxic gas smelling of rotten eggs; a strong reducing agent (S⁻²), and its aqueous solution is a weak acid. With Pb²⁺ and Cu²⁺ it gives black sulfide precipitates: Pb²⁺ + S²⁻ → PbS↓.
• SO₂ is a pungent gas and an acidic oxide (SO₂ + H₂O ⇌ H₂SO₃). Sulfur +4 is an intermediate state, so SO₂ is both a reducing agent (SO₂ + Br₂ + 2H₂O → H₂SO₄ + 2HBr — bromine water is decolourised) and an oxidising agent (SO₂ + 2H₂S → 3S + 2H₂O). It causes acid rain.
• SO₃ combines very violently with water, releasing much heat: SO₃ + H₂O → H₂SO₄.

§4. Sulfuric acid

Dilute H₂SO₄ is a strong diprotic acid whose oxidising species is H⁺: with metals before hydrogen it releases H₂ (Zn + H₂SO₄ → ZnSO₄ + H₂), and it reacts with basic oxides, bases and salts. It forms two series of salts: acid salts (NaHSO₄, hydrogensulfates) and normal salts (Na₂SO₄, sulfates).
Concentrated H₂SO₄ is a strong oxidising agent (the oxidant is S⁺⁶) and releases no hydrogen:

In the cold it passivates Fe, Al and Cr, so the concentrated acid is carried in steel tankers. It is a powerful dehydrating agent: it chars sugar and paper and is used to dry gases.

Safety: when diluting, pour the acid into water in a thin stream, never the reverse.

§5. Industrial production of sulfuric acid

The contact process has three stages:

SO₃ is not absorbed directly in water because an acid mist forms. The second stage is exothermic and reversible, so a compromise temperature and a catalyst are used.

Worked example

What mass of sulfuric acid can be obtained from 120 g of pyrite (FeS₂, M = 120) at 100 % yield?

1. n(FeS₂) = 120/120 = 1 mol.
2. Sulfur balance: 1 mol FeS₂ contains 2 mol S → n(H₂SO₄) = 2 mol.
3. m = 2·98 = 196 g.

§6. An official test item

What mass of 80 % sulfuric acid is needed to dissolve 54 g of a mixture of Fe₃O₄, C and Cu in the mole ratio 2 : 1 : 1 (the reduction product is SO₂)?

Worked example

1. 2x·232 + x·12 + x·64 = 540x = 54 → x = 0.1: Fe₃O₄ 0.2, C 0.1, Cu 0.1 mol.
2. 2Fe₃O₄ + 10H₂SO₄ → 3Fe₂(SO₄)₃ + SO₂ + 10H₂O → 0.2·5 = 1.0 mol of acid.
3. C + 2H₂SO₄ → 0.2 mol; Cu + 2H₂SO₄ → 0.2 mol.
4. Total 1.4 mol = 137.2 g → solution 137.2/0.8 = 171.5 g.

Check yourself

1. 6.4 g of copper dissolved in conc. H₂SO₄. What volume of SO₂ (STP) is released?

2. What properties does SO₂ show, and why?

3. What mass of NaOH neutralises 24.5 g of H₂SO₄ completely?

4. Which metal does not dissolve in cold conc. H₂SO₄: Cu, Fe, Zn?

5. Which salt forms when 0.1 mol SO₂ reacts with 0.1 mol NaOH?

6. In which substance is oxygen in the −1 state: H₂O, Na₂O₂, OF₂, O₃?

7. Which cation is used to detect the sulfate ion?

8. What volume of SO₂ (STP) forms when 9.6 g of sulfur burns?

Common mistakes

• Writing that conc. H₂SO₄ gives H₂ with a metal.
• Pouring water into the acid — dangerous!
• Thinking iron dissolves in cold conc. H₂SO₄: it is passivated.
• Treating the iron in Fe₃O₄ as +3 only and forgetting the SO₂: it also contains Fe²⁺.
• Writing that SO₃ is absorbed directly in water in industry.