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Elements of Groups 16, 17 and 18

Ch. 7Std 12

Easy Overview

Oxygen, sulfur, chlorine, neon — these elements share the same neighborhoods in the periodic table. Group 16 (oxygen family) includes oxygen and its buddies. Group 17 (halogens) are the reactive drama queens. Group 18 (noble gases) are the introverts who do not react with anyone. Together, they make up a big chunk of the p-block. Group 16 elements (O, S, Se, Te, Po) have 6 valence electrons (ns²np⁴). They need two more electrons to complete their octet. Going down the group: atomic radius increases, ionization enthalpy decreases, electronegativity decreases, and metallic character increases (oxygen is a gas, polonium is a metal). All form hydrides H₂E (H₂O, H₂S, H₂Se, H₂Te). Stability decreases down the group (H₂O is very stable, H₂Te decomposes easily). Acidity increases — H₂O is neutral, H₂S is a weak acid. Group 17 halogens (F, Cl, Br, I, At) have 7 valence electrons. They exist as diatomic molecules (X₂). Going down: atomic size increases, electronegativity decreases (F: 4.0 — highest of all elements), bond dissociation enthalpy decreases unusually (F-F bond is weak due to lone pair repulsion), reactivity decreases. Chlorine is prepared by electrolysis of brine and is used for water disinfection, bleaching, and PVC manufacture. Interhalogen compounds (like ICl, ClF₃) are formed between two different halogens. Group 18 noble gases (He, Ne, Ar, Kr, Xe, Rn) have complete octets and were once considered completely inert. In 1962, Neil Bartlett synthesized XePtF₆, shattering that belief. Xenon forms stable compounds with fluorine (XeF₂, XeF₄, XeF₆). Helium is used in balloons and cryogenics. The Montreal Protocol (1987) banned CFCs that were destroying the ozone layer — one of humanity's great environmental success stories.

Group 16 — electronic configuration and trends

Group 16 elements (chalcogens = ore-forming): O, S, Se, Te, Po. General electronic configuration: ns²np⁴. Oxygen is the most abundant element in Earth's crust (~46%). Sulfur occurs as native sulfur (volcanic regions) and as sulfides (FeS₂, ZnS) and sulfates (CaSO₄·2H₂O). Se and Te are less common. Po is radioactive and rare. Oxidation states: -2 (common), +4, +6. Oxygen shows -2 in most compounds but exceptions exist — peroxides (H₂O₂, -1 oxidation state), superoxides (KO₂, -½), and OF₂ (OF₂, +2). Going down the group: atomic radius increases, ionization enthalpy decreases, electronegativity decreases (O: 3.44, S: 2.58, Se: 2.55, Te: 2.10), metallic character increases. Melting and boiling points increase — O₂ is a gas, S₈ is a yellow solid at room temperature.

Oxygen and ozone

Oxygen (O₂): colorless, odorless, tasteless gas. Laboratory prep: 2KClO₃ → 2KCl + 3O₂ (MnO₂ catalyst), or thermal decomposition of KMnO₄, H₂O₂, nitrates. Industrial: fractional distillation of liquid air (O₂ boils at -183°C, N₂ at -196°C). Oxygen supports combustion and is essential for respiration. Uses: steelmaking, rocket oxidizer, medical oxygen, oxyacetylene welding. Ozone (O₃): allotrope of oxygen, pale blue gas with pungent smell. Prepared by passing silent electric discharge through O₂: 3O₂ → 2O₃. Ozone is thermodynamically unstable (decomposes back to O₂). Structure: bent (bond angle ~117°), O-O bond intermediate between single and double. The ozone layer (15-35 km altitude) absorbs UV radiation (200-315 nm). CFCs (CF₂Cl₂) release Cl atoms that catalytically destroy ozone: Cl + O₃ → ClO + O₂; ClO + O → Cl + O₂. Net: O₃ + O → 2O₂. Each Cl atom destroys ~100,000 ozone molecules. The Montreal Protocol (1987) banned CFCs, and the ozone layer is slowly recovering.

Sulfur and sulfuric acid

Sulfur exists mainly as S₈ rings. Rhombic sulfur (α-sulfur) is stable below 96°C; monoclinic (β-sulfur) stable above 96°C. Both are yellow solids. Sulfur dioxide (SO₂): colorless pungent gas, prepared by burning S or roasting sulfide ores. Used as bleaching agent and preservative. Sulfuric acid (H₂SO₄) — the 'king of chemicals' — is produced by the Contact Process: (1) S + O₂ → SO₂, (2) 2SO₂ + O₂ ⇌ 2SO₃ (V₂O₅ catalyst, 450°C, 1-2 atm), (3) SO₃ + H₂SO₄ → H₂S₂O₇ (oleum), (4) H₂S₂O₇ + H₂O → 2H₂SO₄. Properties of H₂SO₄: strong acid, powerful dehydrating agent (charcoal forms when sugar is added to conc. H₂SO₄ — removes H and O as water), and oxidizing agent (hot conc. H₂SO₄ oxidizes Cu to CuSO₄). Uses: fertilizers (superphosphate, ammonium sulfate), petroleum refining, car batteries, detergent manufacture.

Group 17 — halogens: general properties

F, Cl, Br, I, At. Electronic configuration: ns²np⁵. They exist as diatomic molecules (X₂). Physical states: F₂ (pale yellow gas), Cl₂ (greenish-yellow gas), Br₂ (reddish-brown liquid), I₂ (violet-black solid). Going down: atomic size increases, electronegativity decreases (F: 4.0 — highest), ionization enthalpy decreases, melting/boiling points increase, bond dissociation enthalpy decreases. The F-F bond is unusually weak (158 kJ/mol vs Cl-Cl at 243 kJ/mol) due to lone pair-lone pair repulsion between small F atoms. Reactivity decreases down the group — F₂ is the most reactive element known, reacting with almost everything. Oxidizing power decreases: F₂ > Cl₂ > Br₂ > I₂. Cl₂ can oxidize Br⁻ to Br₂ and I⁻ to I₂; Br₂ can oxidize I⁻ to I₂; I₂ is the weakest oxidizing agent.

Chlorine — preparation, properties, and uses

Preparation: lab — MnO₂ + 4HCl → MnCl₂ + Cl₂ + 2H₂O. Industrial — electrolysis of brine (chlor-alkali process): 2NaCl + 2H₂O → 2NaOH + Cl₂ + H₂. Properties: greenish-yellow gas, 2.5× heavier than air, choking odor. Chemical reactions: (1) With water: Cl₂ + H₂O ⇌ HCl + HOCl (hypochlorous acid, a disinfectant). (2) With cold dilute NaOH: 2NaOH + Cl₂ → NaCl + NaOCl + H₂O (bleaching solution). (3) With hot concentrated NaOH: 6NaOH + 3Cl₂ → 5NaCl + NaClO₃ + 3H₂O. (4) Oxidizing agent: Cl₂ + 2FeCl₂ → 2FeCl₃; Cl₂ + 2KI → 2KCl + I₂. (5) Reaction with excess NH₃: 3Cl₂ + 8NH₃ → N₂ + 6NH₄Cl (test for Cl₂ — dense white fumes of NH₄Cl). Uses: water disinfection, bleaching (paper and textiles), PVC manufacture, production of HCl, pesticides (DDT), and as a chemical warfare agent (WWI). Bleaching action of Cl₂ is due to nascent oxygen: HOCl → HCl + [O]; [O] oxidizes colored substances to colorless ones.

Hydrogen chloride and hydrochloric acid

HCl is a colorless pungent gas. Lab preparation: NaCl + H₂SO₄ (conc.) → NaHSO₄ + HCl (below 200°C); NaCl + NaHSO₄ → Na₂SO₄ + HCl (above 200°C). HCl gas is highly soluble in water (1 volume water dissolves ~450 volumes HCl). When HCl gas dissolves in water, it forms hydrochloric acid — it fizzes and gets hot. HCl gas forms dense white fumes with NH₃: HCl + NH₃ → NH₄Cl (this is a test for HCl). Hydrochloric acid is a strong acid, completely dissociated in water: HCl → H⁺ + Cl⁻. Uses: cleaning metals (pickling — removing oxide scale from steel), digestive aid in the stomach (gastric juice contains ~0.5% HCl), laboratory reagent, production of chlorides, and in the petroleum industry. The stomach lining is protected by mucus; when this protection fails, gastric ulcers can form.

Interhalogen compounds

Compounds formed between two different halogens: XYₙ where n = 1, 3, 5, or 7. The central halogen is the larger and less electronegative one. Types: (1) XY type: ClF, BrF, ICl, IBr, ICl. Iodine monochloride (ICl) is used as an iodinating agent. (2) XY₃ type: ClF₃, BrF₃, IF₃. ClF₃ is extremely reactive — it even reacts with glass and is used in the nuclear industry to separate uranium isotopes (converts U to UF₆). (3) XY₅ type: BrF₅, IF₅. (4) XY₇ type: IF₇. Interhalogen compounds are more reactive than the constituent halogens (except F₂), because the X-Y bond is weaker than X-X or Y-Y bonds. All are covalent, diamagnetic molecules. Hybridization: XY (sp³), XY₃ (sp³d), XY₅ (sp³d²), XY₇ (sp³d³). They undergo hydrolysis and act as strong fluorinating agents.

Oxoacids of halogens

Halogens form several oxoacids where the halogen is in positive oxidation state. Chlorine oxoacids are the most important: HOCl (hypochlorous, Cl⁺¹), HClO₂ (chlorous, Cl⁺³), HClO₃ (chloric, Cl⁺⁵), HClO₄ (perchloric, Cl⁺⁷). Acid strength INCREASES with oxidation number: HOCl < HClO₂ < HClO₃ < HClO₄ (HClO₄ is the strongest known acid). Oxidizing power DECREASES: HOCl > HClO₂ > HClO₃ > HClO₄ (HOCl is the strongest oxidizer). Hypochlorites (NaOCl, Ca(OCl)₂) are used as bleaching agents and disinfectants. Potassium chlorate (KClO₃) is used in fireworks, matches, and explosives. Perchloric acid (HClO₄) is used as a laboratory reagent. For bromine: HOBr, HBrO₃. For iodine: HOI, HIO₃, HIO₄, H₅IO₆ (periodic acid). The stability of oxoacids increases down the group.

Group 18 — noble gases

He, Ne, Ar, Kr, Xe, Rn. Electronic configuration: ns²np⁶ (He: 1s²). Discovered by Ramsay and Rayleigh (1890s). They are monatomic, colorless, odorless, tasteless gases. Going down: atomic size increases, ionization enthalpy decreases (but still very high), boiling point increases (He: 4.2 K, Rn: 211 K). Helium has the lowest boiling point of any known substance. Their complete octet (duplet for He) makes them extremely unreactive — they were called 'inert gases.' Uses: He — party balloons, airships (lighter than air, non-flammable), cryogenics (liquid He cools MRI superconducting magnets to 4 K), deep-sea diving (replaces N₂ to prevent the bends). Ne — neon signs (reddish-orange glow). Ar — shielding gas in arc welding, filling incandescent light bulbs, inert atmosphere. Kr — high-intensity lamps, flash for high-speed photography. Xe — flash bulbs, certain lasers, xenon arc lamps. Rn — radioactive, used in radiotherapy.

Noble gas compounds — breaking the inert myth

For decades, noble gases were considered completely incapable of forming compounds. In 1962, Neil Bartlett noticed that PtF₆ is a strong oxidizing agent and that the ionization energy of Xe is similar to that of O₂. He reacted Xe with PtF₆ and synthesized XePtF₆ — the first noble gas compound. This shattered the 'inert gas' myth. Xenon forms several stable compounds: XeF₂ (Xe + F₂ in 1:5 ratio at 400°C in a nickel vessel — linear, sp³d hybridization), XeF₄ (1:5 ratio at 600°C — square planar, sp³d²), XeF₆ (1:20 ratio at 300°C — distorted octahedral, sp³d³). XeF₆ hydrolyzes to give XeO₃ (a shock-sensitive explosive!). XeO₄ is also known. Xe compounds have practical uses: XeF₂ is used as a fluorinating agent. Krypton forms only KrF₂ (unstable, decomposes at room temperature). Radon forms RnF₂. Helium, neon, and argon do not form stable compounds under normal conditions (though some exotic compounds like HeH⁺ have been detected in mass spectrometers). The bonding in noble gas compounds is explained by the availability of d-orbitals in heavier noble gases.

Key Points

  • •Group 16: ns²np⁴, oxidation states -2, +4, +6
  • •Electronegativity and ionization enthalpy decrease down Group 16
  • •Oxygen: most abundant in Earth's crust; Ozone: UV shield, destroyed by CFCs
  • •Contact process: S → SO₂ → SO₃ → H₂SO₄ (king of chemicals)
  • •Group 17: ns²np⁵, diatomic, highly reactive; oxidizing power F₂ > Cl₂ > Br₂ > I₂
  • •F-F bond unusually weak due to lone pair repulsion
  • •Cl₂ + H₂O → HCl + HOCl; HOCl is the bleaching agent (nascent oxygen)
  • •Interhalogens: XY, XY₃, XY₅, XY₇; more reactive than parent halogens
  • •Oxoacids of Cl: acid strength HOCl < HClO₂ < HClO₃ < HClO₄ (strongest known acid)
  • •Oxidizing power of oxoacids: HOCl > HClO₂ > HClO₃ > HClO₄
  • •Group 18: ns²np⁶ (He: 1s²), monatomic, complete octet → unreactive
  • •Bartlett synthesized XePtF₆ in 1962 — first noble gas compound
  • •XeF₂, XeF₄, XeF₆ are stable; XeO₃ is explosive on hydrolysis of XeF₆
  • •He: balloons, cryogenics, deep-sea diving; Ar: welding, light bulbs
  • •Montreal Protocol (1987) banned CFCs to protect the ozone layer
  • •Boiling point of He (4.2 K) is the lowest of any substance

Practice Questions

  • Explain the anomalous behavior of oxygen compared to other Group 16 elements.
  • Write the preparation, properties, and uses of ozone. How do CFCs deplete the ozone layer?
  • Why are halogens strong oxidizing agents? Arrange F₂, Cl₂, Br₂, I₂ in decreasing order of oxidizing power.
  • Discuss reactions of chlorine with (a) water (b) cold dilute NaOH (c) hot concentrated NaOH.
  • Why were noble gases considered inert? How was this myth broken? Write the preparation and structure of XeF₄.
  • Describe the Contact Process for manufacturing sulfuric acid. Write all reactions.
  • What are interhalogen compounds? Classify them with examples.
  • Give two uses each of He, Ar, and Ne. Explain why noble gases have very low boiling points.