Physical Properties of Group 16 Elements
Overview of Group 16 (Chalcogens)
Group 16 elements: Oxygen (O), Sulfur (S), Selenium (Se), Tellurium (Te), Polonium (Po). All have outer electron configuration ns²np⁴ — 6 valence electrons, 2 unpaired p electrons available for bonding.
| Element | Z | Config | State | M.p. (°C) | Character | Common OS |
|---|---|---|---|---|---|---|
| Oxygen (O) | 8 | [He]2s²2p⁴ | Gas (O₂) | −218 | Non-metal | −2, −1, 0 |
| Sulfur (S) | 16 | [Ne]3s²3p⁴ | Solid (yellow) | 113 (rhombic) | Non-metal | −2, 0, +4, +6 |
| Selenium (Se) | 34 | [Ar]3d¹⁰4s²4p⁴ | Solid (grey/red) | 221 | Metalloid | −2, +4, +6 |
| Tellurium (Te) | 52 | [Kr]4d¹⁰5s²5p⁴ | Solid (silvery) | 450 | Metalloid | −2, +4, +6 |
| Polonium (Po) | 84 | [Xe]4f¹⁴5d¹⁰6s²6p⁴ | Solid | 254 | Metal (radioactive) | +2, +4 |
Key Trends Down Group 16
- Atomic radius — increases down group
- Ionisation energy — decreases
- Electronegativity — decreases: O (3.5) > S (2.5) > Se (2.4) > Te (2.1)
- Metallic character — increases (O = non-metal; Po = metal)
- Oxidising power — decreases: O₂ >> S > Se > Te
- Hydride boiling point — H₂O anomalously high (strong H-bonds); H₂S → H₂Te increases with van der Waals forces
- Oxide acidity — sulfur forms acidic oxides (SO₂, SO₃); tellurium oxides are amphoteric
Oxygen and Ozone
Dioxygen (O₂)
Oxygen occurs as O₂ — a colourless, odourless gas. It makes up 21% of the atmosphere. It is paramagnetic (two unpaired electrons in its molecular orbitals — best explained by MO theory). O₂ is the most abundant element on Earth's crust (by mass) when combined with other elements.
Industrial production: Fractional distillation of liquid air. Oxygen condenses at −183°C; nitrogen at −196°C — separated by fractional distillation. Oxygen is used in steelmaking (basic oxygen furnace), medical applications, and combustion.
Laboratory preparation:
Test for oxygen: A glowing splint rekindles (reignites) in the presence of oxygen — oxygen supports combustion.
Reactions of Oxygen
Ozone (O₃) — the Allotrope of Oxygen
Ozone is the second allotrope of oxygen. Structure: bent/angular molecule, bond angle = 117°, bond order = 1.5 (resonance between two structures). O₃ is formed from O₂ by UV radiation in the upper atmosphere:
Properties of O₃:
- Pale blue gas with distinctive sharp smell
- Much stronger oxidising agent than O₂ — used to sterilise water and bleach
- Decomposes to O₂: 2O₃ → 3O₂
- Reacts with alkenes via ozonolysis: C=C + O₃ → carbonyl compounds
- Test for ozone: turns moist starch–iodide paper blue (stronger oxidant than O₂)
- Ozone layer depletion: CFCs release Cl• radicals in the stratosphere → Cl• + O₃ → ClO• + O₂; ClO• + O• → Cl• + O₂ (catalytic cycle — each Cl• destroys thousands of O₃ molecules)
Sulfur — Allotropes and Properties
Allotropes of Sulfur
Sulfur exists in several allotropic forms. The two most important are rhombic (α) sulfur and monoclinic (β) sulfur, both consisting of S₈ rings (crown-shaped cyclic molecules) but with different crystal packing.
| Property | Rhombic Sulfur (α) | Monoclinic Sulfur (β) | Plastic Sulfur |
|---|---|---|---|
| Structure | S₈ rings, orthorhombic crystal | S₈ rings, monoclinic crystal | Amorphous chains (Sₙ) |
| Colour | Pale yellow | Amber/darker yellow | Brown, rubbery |
| Stable temp. | Below 96°C | 96–119°C (m.p.) | Above 160°C → reverses on cooling |
| M.p. | 113°C | 119°C | — |
| Solubility | Soluble in CS₂ | Soluble in CS₂ | Insoluble in CS₂ |
Transition temperature: 96°C — below this, rhombic is the stable form; above this (up to m.p. 119°C), monoclinic is stable. Above ~160°C, the S₈ rings break open → long chain polymers → dark viscous plastic sulfur. Rapid cooling of molten sulfur gives plastic sulfur (metastable).
Reactions of Sulfur
Sulfur is a good oxidising agent in reactions with metals and hydrogen, but can also act as a reducing agent in reactions with fluorine and oxygen.
Oxides of Sulfur: SO₂ and SO₃
Sulfur Dioxide (SO₂) — Sulfur in +4 Oxidation State
SO₂ is a bent molecule with bond angle ~119°. It has resonance structures (like O₃); both S–O bonds are equivalent with bond order ~1.5. SO₂ has a lone pair on S.
Production of SO₂:
Properties and reactions of SO₂:
Sulfur Trioxide (SO₃) — Sulfur in +6 Oxidation State
SO₃ is a planar trigonal molecule — sulfur uses sp² hybridisation with 3 S–O bonds (each with some double bond character via 3d–2p π bonding). Bond angle = 120°. It is the acid anhydride of sulfuric acid.
SO₃ reacts violently with water — produces an acid mist that is hard to handle. This is why SO₃ is first absorbed in concentrated H₂SO₄ to form oleum.
The Contact Process — Manufacture of H₂SO₄
The Four Stages of the Contact Process
The Contact Process manufactures sulfuric acid from sulfur (or pyrite). The name comes from the reactants "coming into contact" on the catalyst surface.
Stage 2 — Le Châtelier Analysis
| Factor | Effect on yield (Le Châtelier) | Effect on rate | Compromise chosen |
|---|---|---|---|
| Temperature | Lower T → better yield (exothermic → low T shifts right) | Lower T → slower rate | 450°C — fast enough rate, ~99.5% yield |
| Pressure | Higher P → better yield (2.5 mol gas → 1.5 mol gas, fewer moles on right — wait, let me recheck: SO₂+½O₂ → SO₃: 1.5 mol left, 1 mol right — high P favours right) | Higher P → faster rate | 1–2 atm — low P used; catalyst gives high yield anyway |
| V₂O₅ catalyst | No effect on equilibrium position | Greatly increases rate | Essential — allows 450°C to work effectively |
| Excess air (O₂) | Shifts equilibrium right (more O₂ → more SO₃) | Increases rate | Air used in excess |
Find the oxidation state of sulfur in: (a) H₂SO₄ (b) SO₂ (c) Na₂SO₃ (d) H₂S₂O₇ (e) S₂O₃²⁻
H₂SO₄: 2(+1) + S + 4(−2) = 0 → S = +8−2 = +6
SO₂: S + 2(−2) = 0 → S = +4
Na₂SO₃: 2(+1) + S + 3(−2) = 0 → S = 6−2 = +4
H₂S₂O₇: 2(+1) + 2S + 7(−2) = 0 → 2S = 14−2 = 12 → S = +6
S₂O₃²⁻ (thiosulfate): 2S + 3(−2) = −2 → 2S = −2+6 = 4 → S = +2 average (one S is 0, one is +4)
Sulfuric Acid (H₂SO₄)
Physical Properties
Concentrated sulfuric acid (98%) is a colourless, oily, very dense liquid (density 1.84 g cm⁻³). It is miscible with water in all proportions. Dilution is extremely exothermic — always add acid to water, never water to acid. Boiling point = 337°C (high — due to strong hydrogen bonding and H₂SO₄ → SO₃ + H₂O equilibrium).
Chemical Properties of H₂SO₄
1. As a Diprotic Strong Acid (dilute H₂SO₄)
2. As a Dehydrating Agent (concentrated H₂SO₄)
3. As an Oxidising Agent (hot concentrated H₂SO₄)
In these reactions, sulfur is reduced from +6 (in H₂SO₄) to +4 (in SO₂) — H₂SO₄ is the oxidising agent.
4. As a Non-volatile Acid
5. Important Uses of H₂SO₄
- Fertiliser manufacture (superphosphate, ammonium sulfate)
- Lead-acid car batteries (electrolyte)
- Detergent and dye manufacture
- Metal refining and pickling (removing oxide layers)
- Making other acids (HCl, HNO₃, HF)
- Oil refining and petrochemicals
Hydrides of Group 16: H₂O, H₂S, H₂Se, H₂Te
Comparison of Group 16 Hydrides
| Hydride | B.p. (°C) | Bond angle | Acid strength in water | Notes |
|---|---|---|---|---|
| H₂O | +100 | 104.5° | Very weak (Kw = 10⁻¹⁴) | Strong H-bonds; anomalously high b.p.; bent; amphiprotic |
| H₂S | −60 | 92° | Weak acid (Ka = 9×10⁻⁸) | Rotten egg smell; toxic; only van der Waals forces |
| H₂Se | −41 | 91° | Stronger than H₂S (Ka = 1×10⁻⁴) | Very toxic |
| H₂Te | −2 | 90° | Strongest (Ka = 2×10⁻³) | Extremely toxic; b.p. highest of H₂S–H₂Te (most e⁻, strongest vdW) |
Hydrogen Sulfide (H₂S) in Detail
H₂S is a colourless gas with a characteristic rotten egg smell, extremely toxic (as toxic as HCN at high concentrations). Bond angle = 92° (S uses near-pure p orbitals). It is a weak diprotic acid:
Test for H₂S / sulfide ions: Lead acetate paper turns black (Pb²⁺ + S²⁻ → PbS ↓ black). Also: H₂S turns aqueous silver nitrate solution black (Ag₂S).
Trends in Group 16 and Identification Tests
Trend in Oxide Acidity
| Oxide | OS of central atom | Type | Reaction with water |
|---|---|---|---|
| SO₂ | +4 | Acidic | SO₂ + H₂O ⇌ H₂SO₃ |
| SO₃ | +6 | Strongly acidic | SO₃ + H₂O → H₂SO₄ |
| SeO₂ | +4 | Acidic | SeO₂ + H₂O → H₂SeO₃ |
| TeO₂ | +4 | Amphoteric | Slight reaction — both acid and base |
| PoO₂ | +4 | Basic | Dissolves in acid; basic character |
As with all groups, oxide acidity decreases going down as metallic character increases.
Qualitative Tests for Sulfur-Containing Ions
| Ion | Test Reagent | Observation | Equation |
|---|---|---|---|
| SO₄²⁻ (sulfate) | Ba²⁺(aq) + dilute HCl (or HNO₃) | White precipitate of BaSO₄ — INSOLUBLE in dilute acids | Ba²⁺ + SO₄²⁻ → BaSO₄↓ (white) |
| SO₃²⁻ (sulfite) | Ba²⁺(aq) | White precipitate of BaSO₃ — dissolves in dilute HCl (unlike BaSO₄) | Ba²⁺ + SO₃²⁻ → BaSO₃↓; BaSO₃ + 2HCl → BaCl₂ + SO₂ + H₂O |
| SO₂ gas | Acidified K₂Cr₂O₇(aq) | Orange → green (Cr⁶⁺ → Cr³⁺); also decolourises KMnO₄ | SO₂ + K₂Cr₂O₇ + H₂SO₄ → Cr₂(SO₄)₃ + K₂SO₄ + H₂O |
| SO₂ gas | Moist filter paper + K₂Cr₂O₇ (orange) | Paper turns green | SO₂ reduces Cr(VI) to Cr(III) |
| H₂S / S²⁻ | Lead acetate paper OR Pb(NO₃)₂(aq) | Paper/solution turns black (PbS) | Pb²⁺ + S²⁻ → PbS↓ (black) |
| S²⁻ | AgNO₃(aq) | Black precipitate of Ag₂S | 2Ag⁺ + S²⁻ → Ag₂S↓ (black) |
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Exercises
- Explain why H₂O has a much higher boiling point (100°C) than H₂S (−60°C) despite having lower molar mass.
H₂O: oxygen is small (Period 2) and highly electronegative (3.5), making the O–H bonds very polar. H₂O molecules form strong O–H···O hydrogen bonds (requires ~20 kJ mol⁻¹ each to break). Many H-bonds must be broken to vaporise water → high b.p. (100°C). H₂S: sulfur is larger, less electronegative (2.5). S–H bonds are less polar — S–H···S H-bonds do not form (S not small/electronegative enough). Only weak London dispersion (van der Waals) forces act between H₂S molecules → much less energy needed to vaporise → low b.p. (−60°C).
- Write the four stages of the Contact Process with equations. Explain why SO₃ is absorbed in concentrated H₂SO₄ rather than water in Stage 3.
Stage 1: S + O₂ → SO₂ (or 4FeS₂ + 11O₂ → 2Fe₂O₃ + 8SO₂). Stage 2: 2SO₂ + O₂ ⇌ 2SO₃ (V₂O₅, 450°C, 1–2 atm). Stage 3: SO₃ + H₂SO₄ → H₂S₂O₇ (oleum). Stage 4: H₂S₂O₇ + H₂O → 2H₂SO₄. Why not water directly: SO₃ + H₂O is highly exothermic and produces an uncontrollable acid mist (fine aerosol of H₂SO₄ droplets) that cannot be collected efficiently and is hazardous. Absorbing in concentrated H₂SO₄ prevents mist formation and gives clean, concentrated product (oleum) that can then be safely diluted.
- How would you distinguish between SO₄²⁻ and SO₃²⁻ ions in solution? Give the reagents, observations, and equations.
Add excess Ba²⁺(aq) (barium chloride): both SO₄²⁻ and SO₃²⁻ give a white precipitate. Then add dilute HCl: BaSO₄ is insoluble in dilute acid — precipitate remains → confirms SO₄²⁻. BaSO₃ dissolves in dilute HCl → effervescence (SO₂ gas) and precipitate disappears → confirms SO₃²⁻. Equations: Ba²⁺ + SO₄²⁻ → BaSO₄↓; Ba²⁺ + SO₃²⁻ → BaSO₃↓; BaSO₃ + 2HCl → BaCl₂ + SO₂↑ + H₂O. (Note: acidify before testing SO₄²⁻ to remove interfering SO₃²⁻ first.)
- Explain the structure of SO₃ and predict its shape using VSEPR. State the hybridisation of S and the oxidation state of S in SO₃.
SO₃: S has 3 S–O bonds and no lone pairs (S uses 3d orbitals to expand its octet). VSEPR: 3 bond pairs → trigonal planar, bond angle = 120°. Each S–O bond has bond order ~1.33 (resonance between three structures, each with one S=O and two S–O). Hybridisation: S is sp² hybridised (three sp² orbitals for σ bonds; unhybridised p orbitals form π bonds with O using S 3d orbitals). Oxidation state: S + 3(−2) = 0 → S = +6.
- Give three different chemical reactions of concentrated H₂SO₄ that cannot be performed by dilute H₂SO₄. Write an equation for each.
(1) As oxidising agent (oxidises metals that do not react with dilute H₂SO₄): Cu + 2H₂SO₄(hot conc.) → CuSO₄ + SO₂↑ + 2H₂O. (2) As dehydrating agent (removes H₂O from compounds): C₁₂H₂₂O₁₁ + conc. H₂SO₄ → 12C + 11H₂O (sugar to carbon). (3) As non-volatile acid (displaces volatile acids from their salts): NaCl(s) + H₂SO₄(conc.) → NaHSO₄ + HCl↑ (releases HCl gas). Many other valid answers include: ethanol → ethene (dehydration); generating NO₂⁺ for nitration; HF production from CaF₂.
- Describe the environmental problem of acid rain caused by sulfur dioxide. Include the source of SO₂, the chemistry of acid rain formation, and two effects on the environment.
Source of SO₂: burning fossil fuels (coal, oil) containing sulfur impurities — S + O₂ → SO₂. Also: smelting metal ores (pyrite roasting). Chemistry: SO₂ + H₂O ⇌ H₂SO₃ (sulfurous acid, weak); 2SO₂ + O₂ → 2SO₃; SO₃ + H₂O → H₂SO₄ (sulfuric acid, strong) → rainfall with pH 4–5. Effects (any 2): (1) Corrodes limestone buildings and statues: CaCO₃ + H₂SO₄ → CaSO₄ + H₂O + CO₂; (2) Acidifies lakes and rivers — kills fish and aquatic invertebrates below pH ~5; (3) Damages forests by leaching mineral nutrients (Ca²⁺, Mg²⁺) from soil; (4) Corrodes metals (iron rusts faster). Solutions: desulfurisation of fuels before burning; flue gas desulfurisation (Ca(OH)₂ + SO₂ → CaSO₃); switching to low-sulfur fuels.
Multiple Choice Quiz — 25 Questions
Unit 11: Group 16 Elements and Compounds
25 Questions · Select one answer eachWhat is the outer electron configuration of all Group 16 elements?
The test for oxygen gas uses a:
Ozone (O₃) has a bond angle of ~117°. This is because:
The allotropy of sulfur that is stable at room temperature is:
Burning sulfur in air produces:
What is the correct order of the Contact Process stages?
The catalyst used in Stage 2 of the Contact Process is:
Why must acid always be added to water (not water to acid) when diluting H₂SO₄?
The oxidation state of sulfur in H₂SO₄ is:
Concentrated H₂SO₄ chars sugar (C₁₂H₂₂O₁₁ → 12C + 11H₂O). This demonstrates H₂SO₄ acting as:
What is observed when copper reacts with hot concentrated H₂SO₄?
The test for SO₄²⁻ ions in solution uses:
H₂S gas turns lead acetate paper:
In the Contact Process, why is a temperature of 450°C used (not 300°C or 600°C)?
What is the oxidation state of S in SO₂?
Acid rain forms when SO₂ from combustion dissolves in rainwater. The primary acid produced is:
Ozone (O₃) is a stronger oxidising agent than O₂ because:
The boiling point trend in Group 16 hydrides is: H₂O >> H₂Te > H₂Se > H₂S. The anomaly (H₂O being highest) is due to:
The Claus process converts H₂S to elemental sulfur. The reaction is:
When concentrated H₂SO₄ reacts with solid NaCl, the gas produced is:
The shape of SO₃ is:
Which of the following correctly summarises the trend in acid strength of Group 16 hydrides?
The laboratory preparation of H₂S gas from iron(II) sulfide uses:
CFCs cause ozone layer depletion by:
Which of these statements about SO₂ is correct?
Unit Test — 50 Marks
Section A — Short Answer
30 marksDescribe the Contact Process for manufacturing sulfuric acid. Write the equation for each stage, give the conditions for Stage 2, and explain why 450°C is used rather than a lower or higher temperature. [5]
Give three different chemical reactions of concentrated H₂SO₄ that demonstrate its three important chemical properties. For each reaction: state the property, write the equation, and name the other product(s). [5]
Explain why H₂O has a boiling point of 100°C while H₂S has a boiling point of −60°C. Explain the trend in boiling points H₂S < H₂Se < H₂Te. Why is the acid strength order the opposite: H₂O < H₂S < H₂Se < H₂Te? [5]
Describe the allotropes of sulfur — rhombic and monoclinic. State: the formula of the structural unit, the crystal system, the temperature of stability, and the melting point of each. Include the transition temperature and what happens above 160°C. [5]
Describe the environmental significance of: (a) the ozone layer and CFC depletion (b) acid rain from SO₂ emissions. For each, write the key equation(s) and state two consequences. [5]
Carry out the following calculations: (a) Find the oxidation state of S in: Na₂S₂O₃, H₂SO₃, SF₆, S₄O₆²⁻. [2] (b) Balance the following redox equation in acid solution: SO₂ + MnO₄⁻ → SO₄²⁻ + Mn²⁺. [3]
Section B — Extended Answer
20 marks(a) Describe the industrial importance of sulfuric acid (H₂SO₄). Give five specific uses with brief explanations for each. [4]
(b) Compare the reactions of SO₂ and SO₃ with water, and with sodium hydroxide. Write equations for all reactions. Classify the oxide type and explain any differences in reactivity. [6]
(a) Compare oxygen and sulfur: explain why O₂ is a gas at room temperature but S₈ is a solid, referring to structure and bonding. [3]
(b) Describe how you would test for the presence of: (i) SO₄²⁻ ions (ii) SO₃²⁻/SO₂ (iii) S²⁻/H₂S. State reagents, observations, equations, and possible interference. [4]
(c) Give three industrial or environmental uses of SO₂ (not its conversion to H₂SO₄). [3]