Chemistry of Non-metals

This is the longest unit in the course and the one that rewards organised revision most. For eight different non-metals it asks the same four things: how the element or its compound is prepared, what its properties are, how it reacts, and what it is used for. Once you see that pattern, twenty-two hours of syllabus collapses into a set of tables you can rebuild from memory.

9.1 Hydrogen

Hydrogen has the simplest possible atom — one proton, one electron, 1s¹. It sits awkwardly in the periodic table because it resembles both group 1 (one valence electron, forms H⁺) and group 17 (one short of a duplet, forms H⁻ in NaH).

Atomic and nascent hydrogen

Ordinary hydrogen is molecular, H₂, and is unreactive at room temperature because the H–H bond enthalpy is a large 436 kJ mol⁻¹.

• Atomic hydrogen is free H atoms, made by passing H₂ through an electric arc between tungsten electrodes at about 4000 °C. It lasts a fraction of a second and recombines, releasing enough heat for the atomic hydrogen torch.
• Nascent hydrogen is hydrogen released inside a reacting mixture, for example by adding zinc to dilute sulphuric acid in the presence of the substance to be reduced. It turns acidified KMnO₄ colourless and reduces FeCl₃ to FeCl₂, whereas bubbling H₂ through the same solutions does nothing.

The difference is energy: nascent hydrogen carries the energy released by the generating reaction and never has to pay the 436 kJ mol⁻¹ to split H₂.

Isotopes of hydrogen

Isotope	Symbol	Protons	Neutrons	Mass number	Abundance	Note
Protium	¹₁H	1	0	1	99.985 %	ordinary hydrogen
Deuterium	²₁H or D	1	1	2	0.015 %	stable, heavy water
Tritium	³₁H or T	1	2	3	trace	radioactive, β-emitter, half-life 12.3 yr

Uses. Deuterium is a tracer for reaction mechanisms and, as D₂O, a reactor moderator; tritium is a radioactive tracer in hydrology and is used in luminous signs.

Preparation of hydrogen

Method	Equation
Active metal + dilute acid (the lab method, Kipp's apparatus)	Zn + H₂SO₄(dil) → ZnSO₄ + H₂↑
Very active metal + cold water (violent)	2Na + 2H₂O → 2NaOH + H₂↑
Amphoteric metal + alkali	Zn + 2NaOH → Na₂ZnO₂ + H₂↑
Steam over red-hot coke (water gas)	C + H₂O --1273 K--> CO + H₂
Steam reforming of methane (main industrial route)	CH₄ + H₂O --Ni, 1100 K--> CO + 3H₂
Electrolysis of acidified water (purest H₂)	2H₂O --electricity--> 2H₂↑ + O₂↑

Properties of hydrogen

Physical	Chemical
Colourless, odourless, tasteless, neutral to litmus	Burns in air: 2H₂ + O₂ → 2H₂O, ΔH = −286 kJ per mol H₂O
Lightest known gas, 0.09 g L⁻¹ at STP	With halogens: H₂ + Cl₂ --sunlight--> 2HCl (explosive)
Very slightly soluble; b.p. −252.8 °C	With N₂: N₂ + 3H₂ ⇌ 2NH₃ (Haber process); with Na: 2Na + H₂ → 2NaH
Diffuses faster than any other gas	Reduces hot metal oxides: CuO + H₂ → Cu + H₂O

Hydrogen as a fuel

Hydrogen has the highest energy content per gram of any chemical fuel — about 143 kJ g⁻¹, against 48 kJ g⁻¹ for petrol — and its only combustion product is water. A hydrogen fuel cell runs the same reaction electrochemically and delivers electricity far more efficiently than a petrol engine.

Advantages	Problems
Highest calorific value per unit mass; product is only water — no CO₂, no SO₂, no soot	Very low density, so storage needs compression, liquefaction at −253 °C or metal hydrides
Renewable if electrolysed using hydropower; fuel cells are silent and efficient	Wide explosive range in air (4–75 %); most hydrogen today still comes from fossil methane

For Nepal the attraction is clear: surplus monsoon hydro-electricity could electrolyse water and be stored as hydrogen instead of being spilled.

Heavy water, D₂O

Heavy water is deuterium oxide, D₂O, present to about 1 part in 6000 in ordinary water. It is obtained by prolonged electrolysis of alkaline water: H₂O is electrolysed faster than D₂O, so D₂O concentrates in the residue — about 30 litres of water leaves roughly 1 mL of nearly pure D₂O.

Property	H₂O	D₂O
Molar mass (g mol⁻¹)	18	20
Melting point	0 °C	3.8 °C
Boiling point	100 °C	101.4 °C
Density at 25 °C	0.997 g cm⁻³	1.104 g cm⁻³

Uses. As a moderator in nuclear reactors — it slows fast neutrons without absorbing them; as a source of deuterium compounds (Ca₃N₂ + 6D₂O → 3Ca(OD)₂ + 2ND₃); and as a tracer. It is not hard water or water containing heavy metals — only the hydrogen is the heavier isotope.

9.2 Allotropes of Oxygen

Allotropy

Allotropy is the existence of an element in two or more forms in the same physical state, differing in physical properties but giving the same chemical products. Oxygen has two allotropes: colourless, odourless, stable dioxygen O₂ (b.p. −183 °C) and pale blue, pungent, unstable ozone O₃ (b.p. −112 °C), which is a far stronger oxidising agent. They differ in the number of atoms per molecule, so this is molecular allotropy; diamond and graphite show crystalline allotropy.

Types of oxides

An oxide is a binary compound of oxygen with another element. Oxides are classified by how they behave towards acids and alkalis.

Type	Behaviour	Examples
Acidic	dissolve in water to give acids; react with alkalis	CO₂, SO₂, SO₃, N₂O₅, P₄O₁₀; SO₂ + 2NaOH → Na₂SO₃ + H₂O
Basic	metal oxides; react with acids to give salt and water	Na₂O, CaO, MgO, FeO; CaO + 2HCl → CaCl₂ + H₂O
Amphoteric	react with both acids and alkalis	ZnO, Al₂O₃, PbO, SnO; ZnO + 2HCl → ZnCl₂ + H₂O and ZnO + 2NaOH → Na₂ZnO₂ + H₂O
Neutral	neither acidic nor basic; no salt with acid or base	CO, NO, N₂O, H₂O
Peroxide	contain the O–O linkage (O₂²⁻); give H₂O₂ with dilute acid	Na₂O₂, BaO₂, H₂O₂; BaO₂ + H₂SO₄ → BaSO₄↓ + H₂O₂
Mixed	behave as a mixture of two simpler oxides	Fe₃O₄ (FeO·Fe₂O₃), Pb₃O₄ (2PbO·PbO₂), Mn₃O₄

Applications of hydrogen peroxide

Hydrogen peroxide, H₂O₂, is a pale blue syrupy liquid sold as a 3 % or 6 % solution or labelled by volume strength. Its only reduction product is water, which makes it an environmentally clean oxidant.

Application	Basis
Antiseptic for wounds (3 %)	liberates nascent oxygen, kills anaerobic bacteria
Bleaching silk, wool, hair, ivory	mild oxidant, does not damage delicate fibre
Restoring old oil paintings	black PbS + 4H₂O₂ → white PbSO₄ + 4H₂O
Rocket propellant; effluent treatment; lab oxidant	decomposes to hot O₂; oxidises cyanides; 2KI + H₂O₂ → 2KOH + I₂

Medical and industrial applications of oxygen

Medical	Industrial
Oxygen therapy for pneumonia, asthma, COVID-19 patients	Oxy-acetylene flame for welding and cutting (≈3200 °C)
Oxygen cylinders for high-altitude climbing — above 8000 m on Everest, supplementary oxygen is standard	Oxy-hydrogen flame (≈2800 °C) for melting quartz and platinum
Artificial respiration, anaesthesia, hyperbaric chambers for CO poisoning	Steel making — oxygen blown through molten pig iron burns off carbon; LOX as rocket oxidiser

9.3 Ozone

Occurrence

Ozone occurs only in traces near the ground but forms a layer in the stratosphere between about 15 km and 35 km, densest near 25 km, where ultraviolet light splits O₂ into atoms that attach to further O₂ molecules. Brought to sea-level pressure all of it would form a layer only 3 mm thick — and that 3 mm absorbs the ultraviolet that would otherwise sterilise the land.

Preparation from oxygen

In the laboratory ozone is made by passing cold, dry oxygen through a silent electric discharge in a Siemens' ozoniser. A spark is deliberately avoided: the heat of a spark would decompose the ozone as fast as it formed.

3O₂ ⇌ 2O₃, ΔH = +284 kJ

Three features of the equation control the yield: the reaction is endothermic, so a low temperature favours ozone; the number of molecules falls (3 → 2), so higher pressure favours ozone; and the discharge must be silent so the product is not destroyed as it forms. Even so the product is only about 10 % ozone by volume — ozonised oxygen — from which pure ozone is separated by liquefaction (b.p. −112 °C against −183 °C for O₂).

Structure of ozone

The molecule is angular with a bond angle of 116.8°; the central oxygen is sp² hybridised and carries a positive formal charge. Experimentally both O–O bonds are identical at 128 pm, between a single bond (148 pm) and a double bond (121 pm). No single Lewis structure explains this, so ozone is a resonance hybrid of the two structures shown above.

Test for ozone

Test	Observation	Equation
Moist starch–iodide paper	turns blue-black	2KI + H₂O + O₃ → 2KOH + I₂ + O₂
Shaken with mercury	mercury loses its meniscus and sticks to glass ("tailing")	a surface film of Hg₂O forms

Ozone layer depletion

Chlorofluorocarbons (CFCs, e.g. CF₂Cl₂ used as refrigerants and aerosol propellants) are inert in the troposphere and drift into the stratosphere. There ultraviolet light breaks a C–Cl bond and releases chlorine free radicals:

CF₂Cl₂ --UV--> •CF₂Cl + •Cl

The chlorine radical then destroys ozone in a chain reaction in which it is regenerated, so one Cl radical can destroy of the order of 10⁵ ozone molecules:

•Cl + O₃ → •ClO + O₂ •ClO + O → •Cl + O₂

Nitric oxide from high-flying aircraft acts in the same way. The thinning is worst over Antarctica each spring — the "ozone hole" — and lets through UV-B that causes skin cancer, cataracts, weakened immunity and reduced crop yields. The Montreal Protocol (1987), which Nepal has ratified, phased out CFCs, and the layer is slowly recovering.

Uses of ozone

Use	Why ozone is chosen
Sterilising drinking water	powerful germicide; unlike chlorine it leaves no taste or smell
Bleaching flour, ivory, wax, delicate fabrics	strong but leaves only O₂ behind
Purifying air in cinemas, tunnels and cold stores	oxidises odours and kills microbes
Ozonolysis; manufacture of KMnO₄, artificial silk, camphor	locates C=C double bonds; general oxidising agent

9.4 Nitrogen

Inertness and active nitrogen

Nitrogen is 78 % of the air by volume, yet almost nothing reacts with it. The reason is the triple bond N≡N, dissociation enthalpy 941 kJ mol⁻¹ — the strongest bond between two like atoms. Nitrogen therefore reacts only under forcing conditions: with Mg at red heat, with H₂ at 450 °C over a catalyst, with O₂ in a lightning flash.

Active nitrogen is made by passing nitrogen at about 1 mm Hg through a high-voltage discharge; a golden-yellow afterglow persists for minutes after the discharge stops. It contains free N atoms and combines directly with S, P, Hg, Na and I₂ at room temperature to give nitrides such as Hg₃N₂ and Na₃N, which ordinary nitrogen never forms.

Ammonia

Laboratory preparation. Ammonia is prepared by heating any ammonium salt with a strong, non-volatile base — usually slaked lime:

2NH₄Cl + Ca(OH)₂ --Δ--> CaCl₂ + 2NH₃↑ + 2H₂O

The tube slopes down at the mouth so that condensed water cannot run back and crack the hot glass. Ammonia is dried over quicklime, CaO and collected by downward displacement of air, being lighter than air and very soluble.

Manufacture: the Haber process. N₂ + 3H₂ ⇌ 2NH₃, ΔH = −92 kJ, run at about 200 atm and 450–500 °C over finely divided iron promoted with molybdenum or K₂O/Al₂O₃.

Chemical properties of ammonia.

Reaction	Equation	Observation / use
Basic character	NH₃ + H₂O ⇌ NH₄⁺ + OH⁻	turns red litmus blue; weak base, Kb = 1.8 × 10⁻⁵
With hydrogen chloride	NH₃ + HCl → NH₄Cl	dense white fumes — the standard test
Combustion in oxygen	4NH₃ + 3O₂ → 2N₂ + 6H₂O	burns with a pale yellow flame; will not burn in air
Catalytic oxidation (Ostwald)	4NH₃ + 5O₂ --Pt, 1073 K--> 4NO + 6H₂O	first step in nitric acid manufacture
With copper(II) sulphate solution	CuSO₄ + 4NH₃ → [Cu(NH₃)₄]SO₄	deep blue solution
With Nessler's reagent	2K₂[HgI₄] + NH₃ + 3KOH → HgO·Hg(NH₂)I↓ + 7KI + 2H₂O	brown precipitate — test for NH₄⁺
Reducing action	3CuO + 2NH₃ → 3Cu + N₂ + 3H₂O	black CuO turns reddish-brown

Applications. Nitrogenous fertilisers — urea, ammonium sulphate and DAP, used across the Terai; nitric acid by the Ostwald process; refrigerant in ice plants (high latent heat of vaporisation); the Solvay process; cleaning agents; explosives and rayon.

Harmful effects. The gas is intensely pungent and corrosive, inflaming eyes, nose and lungs, and is fatal above about 2500 ppm; liquid-ammonia leaks in cold stores are a recognised hazard. Excess ammonium fertiliser washed into rivers causes eutrophication — algal blooms, oxygen depletion and fish kills — and ammonia is directly toxic to fish below 1 mg L⁻¹.

Oxy-acids of nitrogen

Name	Formula	Oxidation state of N	Anhydride	Salt
Hyponitrous acid	H₂N₂O₂	+1	N₂O	Na₂N₂O₂
Nitrous acid	HNO₂	+3	N₂O₃	NaNO₂
Nitric acid	HNO₃	+5	N₂O₅	NaNO₃

Only nitric acid can be isolated pure; nitrous acid exists only in cold dilute solution and disproportionates on warming, 3HNO₂ → HNO₃ + 2NO + H₂O.

Chemical properties of nitric acid

Nitric acid is a colourless fuming liquid (b.p. 83 °C) that turns yellow on standing through dissolved NO₂ from its own decomposition, 4HNO₃ → 4NO₂ + O₂ + 2H₂O.

Behaviour	Equation and note
Strong monobasic acid	NaOH + HNO₃ → NaNO₃ + H₂O
With copper, dilute	3Cu + 8HNO₃ → 3Cu(NO₃)₂ + 2NO↑ + 4H₂O (colourless NO, browns in air)
With copper, concentrated	Cu + 4HNO₃ → Cu(NO₃)₂ + 2NO₂↑ + 2H₂O (brown NO₂)
With magnesium, very dilute (1 %)	Mg + 2HNO₃ → Mg(NO₃)₂ + H₂↑ — the only case giving hydrogen
Passivity	Fe, Al and Cr are coated with an impervious oxide film by conc. HNO₃
With non-metals	C + 4HNO₃ → CO₂↑ + 4NO₂↑ + 2H₂O; S + 6HNO₃ → H₂SO₄ + 6NO₂↑ + 2H₂O; P₄ + 20HNO₃ → 4H₃PO₄ + 20NO₂↑ + 4H₂O
Aqua regia (3 conc. HCl : 1 conc. HNO₃)	HNO₃ + 3HCl → NOCl + 2H₂O + 2[Cl]; the nascent chlorine dissolves Au and Pt

Ring test for the nitrate ion

Add freshly prepared iron(II) sulphate solution to the unknown, then pour concentrated sulphuric acid slowly down the side of the inclined tube so that it forms a separate lower layer. A dark brown ring at the junction of the two liquids confirms a nitrate.

6FeSO₄ + 3H₂SO₄ + 2HNO₃ → 3Fe₂(SO₄)₃ + 4H₂O + 2NO FeSO₄ + NO → [Fe(NO)]SO₄ (brown ring, a nitrosyl complex)

The ring disappears on shaking or warming, as the loosely bound NO escapes.

9.5 Halogens

General characteristics

The halogens are group 17: fluorine, chlorine, bromine, iodine and astatine. Each has the valence configuration ns²np⁵, one electron short of the noble-gas octet, so all are strong oxidising agents and all are diatomic.

Property	F₂	Cl₂	Br₂	I₂
State at 25 °C	pale yellow gas	greenish-yellow gas	red-brown liquid	violet-black solid
m.p. / b.p. (°C)	−219.6 / −188.1	−101.5 / −34.0	−7.2 / 58.8	113.7 / 184.3
Atomic radius (pm)	64	99	114	133
Electronegativity	4.0	3.0	2.8	2.5
Bond enthalpy X–X (kJ mol⁻¹)	158	242	193	151
Oxidising power	strongest			weakest

Melting and boiling points increase down the group as the molecules become larger and more polarisable, strengthening the van der Waals forces. Electronegativity and oxidising power decrease down the group as the atom grows and the incoming electron sits further from the nucleus.

Comparative preparation of Cl₂, Br₂ and I₂

All three are made by oxidising the halide ion, and the oxidising agent needed gets milder down the group as the halide becomes easier to oxidise.

	Chlorine	Bromine	Iodine
Laboratory method	MnO₂ + 4HCl(conc) --Δ--> MnCl₂ + Cl₂↑ + 2H₂O	MnO₂ + 2NaBr + 2H₂SO₄ → MnSO₄ + Na₂SO₄ + Br₂ + 2H₂O	MnO₂ + 2KI + 2H₂SO₄ → MnSO₄ + K₂SO₄ + I₂ + 2H₂O
Milder alternative	2KMnO₄ + 16HCl → 2KCl + 2MnCl₂ + 5Cl₂↑ + 8H₂O (no heat)	KBr + conc. H₂SO₄ (H₂SO₄ itself oxidises Br⁻)	KI + conc. H₂SO₄ (even H₂SO₄ suffices)
Industrial source	electrolysis of brine; Deacon's process 4HCl + O₂ --CuCl₂, 723 K--> 2Cl₂ + 2H₂O	sea bittern: 2NaBr + Cl₂ → 2NaCl + Br₂	Chile saltpetre: 2NaIO₃ + 5NaHSO₃ → 2Na₂SO₄ + 3NaHSO₄ + H₂O + I₂
Collection	downward displacement of air, dried over conc. H₂SO₄	condensed as a liquid, distilled	sublimed and condensed as violet crystals

Comparative chemical properties

Reaction	Chlorine	Bromine	Iodine
With hydrogen	explosive in sunlight: H₂ + Cl₂ → 2HCl	needs 200 °C and a catalyst	reversible and incomplete: H₂ + I₂ ⇌ 2HI
With water	Cl₂ + H₂O → HCl + HOCl (bleaching)	slow, similar	almost insoluble; no reaction
With cold dilute NaOH	Cl₂ + 2NaOH → NaCl + NaOCl + H₂O	Br₂ + 2NaOH → NaBr + NaOBr + H₂O	I₂ + 2NaOH → NaI + NaOI + H₂O
With hot concentrated NaOH	3Cl₂ + 6NaOH → 5NaCl + NaClO₃ + 3H₂O	3Br₂ + 6NaOH → 5NaBr + NaBrO₃ + 3H₂O	3I₂ + 6NaOH → 5NaI + NaIO₃ + 3H₂O
Displacement	displaces Br₂ and I₂ (and gives bleaching powder with slaked lime, 2Cl₂ + 2Ca(OH)₂ → Ca(OCl)₂ + CaCl₂ + 2H₂O)	displaces I₂ only	displaces none
Bleaching action	strong, permanent (oxidation by HOCl)	weaker	negligible

All three reactions with alkali are disproportionations: the halogen is simultaneously reduced (to −1) and oxidised (to +1 or +5).

Tests and uses

Substance	Test	Observation
Cl₂ gas	moist blue litmus	turns red, then bleached white
Br₂ vapour	colour and smell	reddish-brown, choking
I₂	starch solution	deep blue-black
Cl⁻ / Br⁻ / I⁻	AgNO₃ solution, then NH₄OH	curdy white AgCl, soluble in NH₄OH; pale yellow AgBr, sparingly soluble; deep yellow AgI, insoluble

Uses. Chlorine: disinfecting drinking water, bleaching paper and textiles, manufacture of HCl, PVC and bleaching powder. Bromine: AgBr for photographic film, dyes, 1,2-dibromoethane. Iodine: tincture of iodine as an antiseptic, iodised salt to prevent goitre — the programme that has almost eliminated iodine deficiency in the Nepali hills — AgI emulsions and iodoform.

Haloacids: HCl, HBr, HI

Property	HCl	HBr	HI
Bond enthalpy (kJ mol⁻¹)	431	366	299
Boiling point (°C)	−85	−67	−35
Thermal stability	high	moderate	low, decomposes to I₂
Acid strength / reducing power	least	intermediate	greatest

Down the group the H–X bond lengthens and weakens, so the acid ionises more readily, is more easily oxidised and is thermally less stable. HCl is made by warming a chloride with concentrated sulphuric acid: NaCl + H₂SO₄ --Δ--> NaHSO₄ + HCl↑.

9.6 Carbon

Allotropes of carbon

Carbon has more allotropes than any other element, because a carbon atom can hybridise as sp³, sp² or sp and so build completely different frameworks.

	Diamond	Graphite	Fullerene C₆₀
Hybridisation	sp³	sp²	sp² (slightly pyramidal)
Structure	3-D tetrahedral network	flat hexagonal layers 335 pm apart	closed cage, 20 hexagons + 12 pentagons
C–C distance	154 pm	142 pm in the layer	140 and 145 pm
Hardness / density	hardest known; 3.51 g cm⁻³	soft, slippery; 2.25 g cm⁻³	soft molecular solid; 1.65 g cm⁻³
Conduction	insulator	good conductor along the layers	semiconductor; K₃C₆₀ superconducts
Uses	cutting tools, abrasives, gem stones	pencil lead, electrodes, lubricant, reactor moderator	drug-delivery cages, lubricants, nanotechnology

Graphite layers are held together only by weak van der Waals forces, so they slide over one another — hence the lubricating action and the mark a pencil leaves. Within a layer each carbon uses three of its four valence electrons for σ bonds; the fourth is delocalised and carries current. Fullerenes were discovered in 1985; C₆₀ is football-shaped and is the only allotrope of carbon that dissolves in ordinary solvents such as benzene, giving a magenta solution. Coal, coke, charcoal and lamp black are called amorphous carbon, but X-ray study shows them to be microcrystalline graphite.

Carbon monoxide

	Detail
Preparation (lab)	HCOOH --conc. H₂SO₄, 373 K--> CO↑ + H₂O (dehydration of formic acid)
Preparation (industrial)	C + H₂O --1273 K--> CO + H₂ (water gas); 2C + O₂ + 4N₂ → 2CO + 4N₂ (producer gas)
Physical properties	colourless, odourless, almost insoluble, slightly lighter than air, b.p. −192 °C
Neutral oxide	no reaction with acids or alkalis; does not turn litmus
Combustion	2CO + O₂ → 2CO₂, blue flame, ΔH = −566 kJ — a useful fuel
Reducing agent	Fe₂O₃ + 3CO → 2Fe + 3CO₂ (blast furnace); CuO + CO → Cu + CO₂
Carbonyl formation	Ni + 4CO --330 K--> Ni(CO)₄; heating to 450 K reverses it — the Mond process for pure nickel
Toxicity	binds haemoglobin about 300 times more strongly than O₂ as carboxyhaemoglobin, so the blood can no longer carry oxygen; with Cl₂ it gives phosgene, COCl₂
Uses	fuel gas; reducing agent in metallurgy; Mond process; manufacture of methanol, CO + 2H₂ --ZnO/Cr₂O₃--> CH₃OH

9.7 Phosphorus

Allotropes of phosphorus

Allotrope	Appearance	Note
White (yellow) phosphorus	waxy white solid, P₄ tetrahedra	most reactive, glows in the dark, stored under water
Red phosphorus	dull red powder, polymeric chains	stable, used on the striking surface of matchboxes
Scarlet and violet phosphorus	scarlet powder; violet crystals	from white P in PBr₃; by heating red P with sodium
Black phosphorus (α and β)	grey-black flakes, layered like graphite	the thermodynamically most stable form; conducts

(The syllabus asks for the names only.)

Phosphine, PH₃

Preparation. White phosphorus is boiled with concentrated sodium hydroxide in an atmosphere of coal gas or CO₂, to keep out air:

P₄ + 3NaOH + 3H₂O --boil--> PH₃↑ + 3NaH₂PO₂

Pure phosphine is obtained from a metal phosphide: Ca₃P₂ + 6H₂O → 3Ca(OH)₂ + 2PH₃↑

Property	Detail
Physical	colourless gas, smell of rotten fish, very poisonous, sparingly soluble
Basic character	PH₃ + HI → PH₄I; weaker base than NH₃, because the P–H bond is longer and the lone pair is in a larger, more diffuse orbital
Inflammability	pure PH₃ is not spontaneously inflammable, but traces of P₂H₄ make it catch fire in air, giving vortex smoke rings of P₄O₁₀
Reactions	3CuSO₄ + 2PH₃ → Cu₃P₂↓ + 3H₂SO₄ (black ppt — the test); PH₃ + 4Cl₂ → PCl₅ + 3HCl
Uses	Holme's signals — calcium phosphide and carbide thrown into the sea release PH₃, which ignites and marks the position; smoke screens; grain fumigant

9.8 Sulphur

Allotropes of sulphur and uses

Allotrope	Crystal form	Stable range	Density (g cm⁻³)	Made by
Rhombic (α)	octahedral, yellow	below 95.6 °C	2.06	slow evaporation of a solution of S in CS₂
Monoclinic (β)	long needles, pale yellow	95.6 °C to 119 °C	1.98	slow cooling of molten sulphur, pouring off the liquid
Plastic (γ)	amorphous, rubbery	unstable	1.95	pouring boiling sulphur into cold water

95.6 °C is the transition temperature: below it monoclinic sulphur changes to rhombic, above it the change reverses. Both crystalline forms are built from the same puckered S₈ crown ring and differ only in packing, which is why their chemistry is identical.

Uses. Manufacture of sulphuric acid (by far the largest use); vulcanisation of rubber; gunpowder and matches; fungicides and insecticides; sulphur ointments for skin disease; bleaching.

Hydrogen sulphide and Kipp's apparatus

H₂S is prepared by the action of dilute sulphuric acid on iron(II) sulphide:

FeS + H₂SO₄(dil) → FeSO₄ + H₂S↑

Because qualitative analysis needs a steady, controllable stream, the gas is generated in Kipp's apparatus, a three-bulb vessel in which the acid is pushed away from the solid the moment the stopcock is closed.

With the stopcock open, acid from the reservoir rises into the middle bulb and attacks the FeS. With it closed, the gas pressure forces the acid back down into the bottom bulb and up the stem, contact with the solid is broken and the reaction stops by itself, so no reagent is wasted.

Property	Detail
Physical	colourless gas, rotten-egg smell, highly poisonous, heavier than air, fairly soluble
Acidic	weak dibasic acid, H₂S ⇌ H⁺ + HS⁻ ⇌ 2H⁺ + S²⁻; turns blue litmus faintly red
Combustion	2H₂S + 3O₂ → 2SO₂ + 2H₂O (excess air); 2H₂S + O₂ → 2S + 2H₂O (limited air)
Reducing agent	Cl₂ + H₂S → 2HCl + S↓; 2FeCl₃ + H₂S → 2FeCl₂ + 2HCl + S↓; decolourises acidified KMnO₄
Group reagent (its main use)	precipitates coloured sulphides — CuS and PbS black, CdS yellow, ZnS white, MnS buff
Test	blackens lead-acetate paper: Pb(CH₃COO)₂ + H₂S → PbS↓ + 2CH₃COOH

Sulphur dioxide

	Detail
Preparation (lab)	Na₂SO₃ + H₂SO₄(dil) → Na₂SO₄ + SO₂↑ + H₂O
	Cu + 2H₂SO₄(conc) --Δ--> CuSO₄ + SO₂↑ + 2H₂O
Preparation (industrial)	S + O₂ → SO₂; 4FeS₂ + 11O₂ → 2Fe₂O₃ + 8SO₂ (roasting of pyrites)
Physical	colourless, suffocating smell of a burning matchstick, 2.2 times heavier than air, very soluble, easily liquefied (b.p. −10 °C)
Acidic oxide	SO₂ + H₂O → H₂SO₃; SO₂ + 2NaOH → Na₂SO₃ + H₂O
Reducing agent	decolourises acidified KMnO₄; turns acidified K₂Cr₂O₇ paper green (the standard test); 2FeCl₃ + SO₂ + 2H₂O → 2FeCl₂ + H₂SO₄ + 2HCl
Oxidising agent / bleach	SO₂ + 2H₂S → 3S↓ + 2H₂O; bleaches by reduction, so the colour slowly returns in air — a temporary bleach, unlike chlorine
Uses	manufacture of H₂SO₄; bleaching wool, silk and straw; preservative; disinfectant; refrigerant

Sulphuric acid — the Contact process

1. SO₂ is produced by burning sulphur or roasting iron pyrites.
2. Purification in a dust chamber, washing tower, drying tower and arsenic purifier — arsenic compounds would poison the catalyst.
3. Catalytic oxidation: 2SO₂ + O₂ ⇌ 2SO₃, ΔH = −196 kJ, over V₂O₅ at 450 °C and about 2 atm. The reaction is exothermic and 3 molecules become 2, so a low temperature and a raised pressure both favour SO₃; 450 °C is the compromise with rate.
4. Absorption in 98 % H₂SO₄ to give oleum, SO₃ + H₂SO₄ → H₂S₂O₇; direct absorption in water gives an unmanageable acid mist.
5. Dilution: H₂S₂O₇ + H₂O → 2H₂SO₄.

Property of H₂SO₄	Illustration
Strong dibasic acid	H₂SO₄ + 2NaOH → Na₂SO₄ + 2H₂O
High boiling point (338 °C), non-volatile	displaces volatile acids: NaCl + H₂SO₄ → NaHSO₄ + HCl↑
Powerful dehydrating agent	C₁₂H₂₂O₁₁ --H₂SO₄--> 12C + 11H₂O (sugar chars to a black column); CuSO₄·5H₂O → CuSO₄
Oxidising agent when hot and concentrated	Cu + 2H₂SO₄ → CuSO₄ + SO₂↑ + 2H₂O; C + 2H₂SO₄ → CO₂ + 2SO₂ + 2H₂O
Uses	fertilisers (superphosphate, ammonium sulphate), lead–acid batteries, petroleum refining, dyes, detergents, pickling of steel — its consumption is taken as a measure of a country's industrial development, hence "king of chemicals"

Sodium thiosulphate, Na₂S₂O₃·5H₂O

Commonly called hypo, it is made by boiling sodium sulphite solution with flowers of sulphur: Na₂SO₃ + S --boil--> Na₂S₂O₃

Reaction	Equation	Use
With dilute acid	Na₂S₂O₃ + 2HCl → 2NaCl + S↓ + SO₂↑ + H₂O	milky turbidity; a kinetics experiment
With iodine	2Na₂S₂O₃ + I₂ → 2NaI + Na₂S₄O₆	iodometric titration; the blue starch colour vanishes at the end point
With silver bromide	AgBr + 2Na₂S₂O₃ → Na₃[Ag(S₂O₃)₂] + NaBr	fixing a photographic film — unexposed AgBr is dissolved away
With chlorine	Na₂S₂O₃ + 4Cl₂ + 5H₂O → 2NaHSO₄ + 8HCl	antichlor — removes excess chlorine after bleaching

It is also a cyanide antidote in medicine and is used to extract silver from its ores.

Chapter summary

• Hydrogen: nascent hydrogen is more reactive than H₂ because the 436 kJ mol⁻¹ H–H bond need not be broken. Isotopes protium, deuterium (D₂O, reactor moderator) and tritium. As a fuel: 143 kJ g⁻¹, only water, but hard to store.
• Oxides: acidic (CO₂, SO₂), basic (Na₂O, CaO), amphoteric (ZnO, Al₂O₃), neutral (CO, NO), peroxide (Na₂O₂, BaO₂), mixed (Fe₃O₄, Pb₃O₄). H₂O₂ strength in g L⁻¹ = volume strength × 3.04.
• Ozone: 3O₂ ⇌ 2O₃, ΔH = +284 kJ, made by silent electric discharge; angular, 116.8°, both O–O bonds 128 pm, a resonance hybrid; turns starch–iodide paper blue; destroyed catalytically by Cl radicals from CFCs.
• Nitrogen is inert because of the 941 kJ mol⁻¹ N≡N bond; active nitrogen is atomic and very reactive. Ammonia: 2NH₄Cl + Ca(OH)₂ → CaCl₂ + 2NH₃↑ + 2H₂O, dried over quicklime only, tested with HCl (white fumes) and Nessler's reagent (brown ppt).
• Nitric acid gives NO₂ (conc.) or NO (dil.) with metals, never H₂. Nitrate is confirmed by the brown ring [Fe(NO)]SO₄. Oxy-acids: H₂N₂O₂, HNO₂, HNO₃.
• Halogens: oxidising power and electronegativity fall down the group; the F–F bond is anomalously weak (158 kJ mol⁻¹). MnO₂ + 4HCl → MnCl₂ + Cl₂ + 2H₂O. Cold dilute alkali gives X⁻ + XO⁻, hot concentrated alkali X⁻ + XO₃⁻. Acid strength and reducing power HCl < HBr < HI, so conc. H₂SO₄ cannot make HBr/HI.
• Carbon: diamond sp³ network, graphite sp² layers 335 pm apart, C₆₀ cage of 20 hexagons and 12 pentagons. CO is neutral, a strong reducing agent (Fe₂O₃ + 3CO → 2Fe + 3CO₂) and a poison through carboxyhaemoglobin.
• Phosphorus: white, red, scarlet, violet and black allotropes. PH₃ from P₄ + 3NaOH + 3H₂O → PH₃ + 3NaH₂PO₂; weaker base than NH₃; Holme's signals.
• Sulphur: rhombic ⇌ monoclinic at 95.6 °C, both S₈ rings. H₂S from Kipp's apparatus blackens lead-acetate paper; SO₂ is a temporary bleach and turns acidified K₂Cr₂O₇ green; Contact process 2SO₂ + O₂ ⇌ 2SO₃ over V₂O₅ at 450 °C, then oleum; hypo is the photographic fixer and the iodometric titrant.

Practice questions

Group A — Multiple choice (1 mark each)

1. Which of the following is used to dry ammonia gas? [1]aconc. H₂SO₄banhydrous CaCl₂cquicklimedP₄O₁₀
2. The bond angle in the ozone molecule is [1]a104.5°b109.5°c116.8°d120°
3. The bond dissociation enthalpy of F₂ is less than that of Cl₂ because [1]afluorine is more electronegativeblone-pair repulsion is large in the small F₂ moleculecfluorine has a higher electron affinitydF₂ is a gas
4. The transition temperature of rhombic and monoclinic sulphur is [1]a95.6 °Cb113.7 °Cc119 °Cd444.6 °C
5. Sulphur dioxide is passed over a paper moistened with acidified potassium dichromate. The paper turns [1]abluebblackcgreendpink

Answers to Group A

1. (c) — quicklime is basic, so it does not react with the basic gas; the other three form ammonium salts or CaCl₂·8NH₃.

2. (c) — ozone is angular with a bond angle of 116.8°.

3. (b) — the F atom is very small, so the three lone pairs on each atom repel across the short bond and weaken it.

4. (a) — below 95.6 °C rhombic is stable, above it monoclinic.

5. (c) — SO₂ reduces orange Cr(VI) to green Cr(III).

Group B — Short answer (5 marks each)

1. Distinguish between atomic hydrogen and nascent hydrogen. Why is nascent hydrogen a better reducing agent than ordinary hydrogen gas? [5]
2. How is ozone prepared from oxygen in the laboratory? Draw the essential features of the apparatus and explain why a silent electric discharge and a low temperature are used. [5]
3. Give reasons: (i) ammonia cannot be dried over concentrated sulphuric acid; (ii) concentrated sulphuric acid cannot be used to prepare HI; (iii) nitric acid does not liberate hydrogen with most metals. [5]
4. Classify the following oxides and justify each answer with one equation: CO, ZnO, N₂O₅, Na₂O₂, Pb₃O₄. [5]
5. Describe the working of Kipp's apparatus for the preparation of hydrogen sulphide, with the equation, and give the test for the gas. [5]

Answers to Group B

1. Atomic hydrogen is free H atoms from passing H₂ through an electric arc at about 4000 °C; it lasts a fraction of a second and recombines with great heat (atomic hydrogen torch). Nascent hydrogen is hydrogen at the moment of generation inside a reacting mixture, e.g. Zn + dil. H₂SO₄. It reduces acidified KMnO₄ and FeCl₃, which H₂ gas will not, because it need not first supply the 436 kJ mol⁻¹ H–H bond enthalpy and retains the energy of the generating reaction.

2. Cold dry oxygen is passed through a Siemens' ozoniser — a silent electric discharge between tin-foil electrodes on concentric glass tubes: 3O₂ ⇌ 2O₃, ΔH = +284 kJ. A spark would heat the gas and destroy the ozone as fast as it formed, so the discharge must be silent; and since the reaction is endothermic, a low temperature shifts the equilibrium right (Le Chatelier). The product is about 10 % ozonised oxygen.

3. (i) Conc. H₂SO₄ is acidic and ammonia basic: 2NH₃ + H₂SO₄ → (NH₄)₂SO₄; quicklime is used instead. (ii) Conc. H₂SO₄ oxidises the HI as it forms: 8HI + H₂SO₄ → 4I₂ + H₂S + 4H₂O; non-oxidising phosphoric acid is used instead. (iii) NO₃⁻ is a better oxidising agent than H⁺, so hydrogen is oxidised to water while NO₃⁻ is reduced to NO₂ (conc.) or NO (dil.).

4. CO — neutral. ZnO — amphoteric: ZnO + 2HCl → ZnCl₂ + H₂O and ZnO + 2NaOH → Na₂ZnO₂ + H₂O. N₂O₅ — acidic: N₂O₅ + H₂O → 2HNO₃. Na₂O₂ — peroxide: Na₂O₂ + H₂SO₄ → Na₂SO₄ + H₂O₂. Pb₃O₄ — mixed, 2PbO·PbO₂.

5. Lumps of FeS rest on the constriction in the middle bulb; dilute H₂SO₄ poured into the top funnel collects in the bottom bulb. Opening the stopcock lets acid rise into the middle bulb: FeS + H₂SO₄ → FeSO₄ + H₂S↑. Closing it lets the gas pressure force the acid back down out of contact with the solid, so the reaction stops automatically. Test: the gas blackens lead-acetate paper, Pb(CH₃COO)₂ + H₂S → PbS↓ + 2CH₃COOH.

Group C — Long answer (8 marks each)

1. (a) Describe with equations the laboratory preparation of chlorine, bromine and iodine from a suitable halide, and explain why a weaker oxidising agent suffices as the group is descended.(b) Write the equations for the action of chlorine on cold dilute and on hot concentrated sodium hydroxide, and state the type of reaction involved.(c) How would you distinguish between solutions of NaCl, NaBr and NaI in the laboratory? [8]
2. (a) Describe the manufacture of sulphuric acid by the Contact process, giving the conditions used in the catalytic converter and explaining the choice of each condition.(b) Why is sulphur trioxide absorbed in concentrated sulphuric acid rather than in water?(c) Calculate the mass of sulphuric acid that can be obtained from 1.00 tonne of sulphur if the overall yield is 90 %. (S = 32, H = 1, O = 16) [8]
3. (a) Give the names and formulae of three oxy-acids of nitrogen with the oxidation state of nitrogen in each.(b) Write balanced equations for the action of nitric acid on copper (dilute and concentrated), on carbon and on sulphur.(c) Describe the ring test for the nitrate ion, with equations. [8]

Answers to Group C

1. (a) MnO₂ + 4HCl(conc) --Δ--> MnCl₂ + Cl₂↑ + 2H₂O; MnO₂ + 2NaBr + 2H₂SO₄ → MnSO₄ + Na₂SO₄ + Br₂ + 2H₂O; MnO₂ + 2KI + 2H₂SO₄ → MnSO₄ + K₂SO₄ + I₂ + 2H₂O (conc. H₂SO₄ alone oxidises I⁻). Down the group the halide ion is larger and holds its outer electron less firmly, so it is oxidised more easily and a milder oxidant suffices. (b) Cold dilute: Cl₂ + 2NaOH → NaCl + NaOCl + H₂O. Hot concentrated: 3Cl₂ + 6NaOH → 5NaCl + NaClO₃ + 3H₂O. Both are disproportionations — chlorine goes from 0 to −1 and simultaneously to +1 (or +5). (c) Add AgNO₃ solution: NaCl gives curdy white AgCl, readily soluble in NH₄OH; NaBr pale yellow AgBr, sparingly soluble; NaI deep yellow AgI, insoluble.

2. (a) Burn sulphur or roast pyrites, 4FeS₂ + 11O₂ → 2Fe₂O₃ + 8SO₂; purify the gas of dust, moisture and arsenic (a catalyst poison); then in the contact tower 2SO₂ + O₂ ⇌ 2SO₃, ΔH = −196 kJ, over V₂O₅ at 450 °C and about 2 atm. Exothermic, so a low temperature favours SO₃ but the rate would be too slow — 450 °C is the compromise; 3 molecules become 2, so pressure helps, but only a modest pressure is needed since the yield is already about 97 %. Finally SO₃ + H₂SO₄ → H₂S₂O₇ and H₂S₂O₇ + H₂O → 2H₂SO₄. (b) Water absorption is strongly exothermic and gives a dense acid mist that is neither condensed nor absorbed, so most of the product would be lost. (c) 1 mol S gives 1 mol H₂SO₄, i.e. 32 g S gives 98 g acid. Theoretical mass = 1.00 × 10⁶ × 98/32 = 3.0625 × 10⁶ g = 3062.5 kg; actual = 0.90 × 3062.5 = 2756 kg ≈ 2.76 tonne.

3. (a) Hyponitrous acid H₂N₂O₂ (N = +1); nitrous acid HNO₂ (+3); nitric acid HNO₃ (+5). (b) 3Cu + 8HNO₃(dil) → 3Cu(NO₃)₂ + 2NO↑ + 4H₂O; Cu + 4HNO₃(conc) → Cu(NO₃)₂ + 2NO₂↑ + 2H₂O; C + 4HNO₃ → CO₂↑ + 4NO₂↑ + 2H₂O; S + 6HNO₃ → H₂SO₄ + 6NO₂↑ + 2H₂O. (c) Add fresh FeSO₄ solution, then pour conc. H₂SO₄ gently down the side of the inclined tube so that it forms a separate lower layer; a dark brown ring appears at the junction: 6FeSO₄ + 3H₂SO₄ + 2HNO₃ → 3Fe₂(SO₄)₃ + 4H₂O + 2NO, then FeSO₄ + NO → [Fe(NO)]SO₄.