Chapter OneNon-Metals and Their Compounds

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Section 1Key terms and concepts

Learn these definitions word-for-word. NECTA regularly awards one full mark for a correct definition, and examiners look for the key phrase in each one.

Section 2Meaning and general properties of non-metals

Non-metals are elements which form negative ions by gaining electrons. The number of electrons a non-metal gains is called its valency.

Chlorine gains only one electron, so it is said to be monovalent. Because all non-metals form negative ions, they are described as electronegative elements.

1. Oxidising properties

All non-metals are oxidising agents because they accept electrons. The strength of the oxidising power differs from one non-metal to another and decreases in the order:

Fluorine is therefore the strongest oxidising agent and the most electronegative element known.

2. Displacement of one non-metal by another

Depending on their oxidising power, non-metals displace each other from solutions of their compounds. Chlorine displaces both bromine and iodine, while bromine displaces only iodine. Iodine displaces neither.

Observation: when chlorine is bubbled through colourless potassium iodide solution, the solution turns dark brown and a dark grey-black solid (iodine) is deposited.

Section 3Chlorine

Chlorine is a member of Group VII of the periodic table — the halogens. It is so reactive that it never occurs free in nature; it is always found combined, mainly as sodium chloride (rock salt) in the sea and in underground deposits.

Laboratory preparation of chlorine gas

Chlorine is prepared by the action of concentrated hydrochloric acid on an oxidising agent such as potassium permanganate (potassium manganate(VII)) or manganese(IV) oxide. The oxidising agent supplies the oxygen atoms needed to oxidise the acid.

How the apparatus works

• Potassium permanganate is placed in the flask and concentrated hydrochloric acid is added slowly from the dropping funnel.
• Effervescence occurs and a greenish-yellow gas is given off.
• The gas is passed through water, which dissolves the hydrogen chloride fumes carried over with the chlorine.
• The gas then passes through concentrated sulphuric acid, which dries it.
• It is collected by downward delivery (upward displacement of air) because chlorine is denser than air. It cannot be collected over water because it dissolves in water.
• The whole preparation is carried out in a fume cupboard because chlorine is poisonous.

Using manganese(IV) oxide instead, the mixture must be heated:

Physical properties of chlorine

• Greenish-yellow gas with a sharp, choking (pungent) smell.
• Denser than air (relative molecular mass 71 against air's 29).
• Fairly soluble in water, forming chlorine water.
• Poisonous; easily liquefied under pressure.

Chemical properties of chlorine

1. Action on metals

Chlorine reacts with nearly all metals to form metallic chlorides.

2. Action on non-metals

3. Bleaching action

Chlorine bleaches damp litmus paper and wet coloured flowers — but it has no effect on dry materials. The gas first reacts with water to form a mixture of hydrochloric acid and chloric(I) acid (hypochlorous acid).

Chloric(I) acid is unstable and decomposes, releasing an atom of nascent oxygen:

Chlorine therefore bleaches by oxidation, and the bleaching is permanent.

4. Chlorine as an oxidising agent

Chlorine reacts by accepting electrons:

(a) With hydrogen sulphide. Chlorine oxidises hydrogen sulphide to sulphur at room temperature, and is itself reduced to chloride ions. A pale yellow deposit of sulphur is seen.

(b) With sulphur(IV) oxide. Chlorine oxidises sulphurous acid to sulphuric(VI) acid.

(c) With iron(II) chloride. Pale green iron(II) chloride is oxidised to yellow-brown iron(III) chloride.

5. Reaction with sodium hydroxide

The products depend on the temperature and concentration of the alkali.

Sodium chlorate(V) is used as a weed killer and in the manufacture of matches and fireworks.

Uses of chlorine

• As a bleaching agent in the paper and textile industries.
• As a germicide and weed killer.
• In the manufacture of hydrogen chloride gas.
• In the manufacture of plastics such as PVC.
• In the treatment of drinking water and sewage.
• As a disinfectant in swimming pools.

Test for chlorine gas

Chlorine is a greenish-yellow gas which rapidly bleaches damp litmus paper (blue litmus turns red first, then white). It also turns damp starch-iodide paper blue-black.

Section 4Hydrogen chloride gas and hydrochloric acid

Laboratory preparation of hydrogen chloride

Hydrogen chloride gas is prepared by reacting concentrated sulphuric acid with sodium chloride (rock salt) in a flask. The gas is dried by passing it through concentrated sulphuric acid.

Physical properties of hydrogen chloride

• Colourless gas with a sharp, irritating smell.
• Denser than air — collected by downward delivery.
• Extremely soluble in water — cannot be collected over water.
• Forms white misty fumes in moist air.

Dissolving hydrogen chloride in water

Because the gas is so soluble, water would be sucked back into the hot flask and crack it. An inverted filter funnel is therefore used, with its wide rim only just touching the water surface. If suck-back begins, the water level falls below the rim and air enters, breaking the vacuum.

Chemical properties of hydrogen chloride gas

1. It is an acidic gas: it turns damp blue litmus paper red (dry litmus is unaffected).
2. It does not burn and does not support combustion.
3. It forms white fumes when exposed to moist air.
4. It reacts with ammonia gas to give dense white fumes of ammonium chloride — this is the test for the gas.
   HCl_(g) + NH_3(g) → NH₄Cl_(s)dense white fumes
5. Its solution reacts with metals above hydrogen in the reactivity series, liberating hydrogen gas.
   Zn_(s) + 2HCl_(aq) → ZnCl_2(aq) + H_2(g)
6. Its solution reacts with carbonates, giving carbon dioxide.
   CaCO_3(s) + 2HCl_(aq) → CaCl_2(aq) + H₂O_(l) + CO_2(g)

Test for chloride ions

Hydrochloric acid

When hydrogen chloride gas dissolves in water, hydrochloric acid is formed. Industrially the acid is made by burning hydrogen in chlorine. Hydrogen comes from the cracking of alkanes and chlorine from the electrolysis of brine in the Downs cell. To avoid explosion, a jet of hydrogen is burned in an atmosphere of chlorine.

The product is dissolved in water to give hydrochloric acid of about 35 % concentration.

Uses of hydrochloric acid

• Removing rust from iron before galvanising (pickling).
• Chlorination and treatment of water and sewage.
• Manufacture of silver chloride used in photographic films.
• Manufacture of chlorides, dyes and in the food industry.

Section 5Chlorine — questions and answers

Attempt each question in your exercise book first. Only then tap Show suggested answer to compare your work.

Section 6Nitrogen and its compounds

Nitrogen is a non-metal in Group V of the periodic table. It is a gas at room temperature and is remarkably unreactive — the two atoms in N≡N are held by a very strong triple covalent bond that takes a great deal of energy to break.

Occurrence

• Free in the atmosphere, where it makes up about 78 % by volume.
• Combined as Chile saltpetre (sodium nitrate, NaNO₃), a mineral deposit found in Chile.
• In small quantities in compounds such as ammonium sulphate and potassium nitrate.
• As an essential constituent of proteins in all living matter.

Obtaining nitrogen from the atmosphere

Nitrogen occurs free in air and can be separated by removing all the other constituents one at a time.

1. Water in the aspirator bottle drives the air steadily through the apparatus.
2. The air bubbles through concentrated sodium hydroxide, which absorbs carbon dioxide.
   2NaOH_(aq) + CO_2(g) → Na₂CO_3(aq) + H₂O_(l)
3. The gas then passes over heated copper, which removes the oxygen. The pink-brown copper turns black.
   2Cu_(s) + O_2(g) → 2CuO_(s)
4. The remaining gas is nitrogen, collected over water.

The nitrogen obtained this way is not pure because it still contains traces of the noble (rare) gases — argon, neon, helium — which are unreactive and are not absorbed. Their presence makes atmospheric nitrogen slightly denser than pure nitrogen. It was exactly this small density difference that led Lord Rayleigh to discover argon in 1894.

Laboratory preparation of nitrogen

Pure nitrogen is prepared by gently heating a solution of ammonium nitrite, which is made in the tube by mixing ammonium chloride and sodium nitrite.

Properties of nitrogen

• Colourless, odourless and tasteless gas; only slightly soluble in water.
• Slightly less dense than air.
• It is neutral to litmus and gives a negative result to all the common gas tests — it neither burns nor supports combustion, and a burning splint is simply extinguished.
• It is very unreactive at room temperature because of the strong triple covalent bond (N≡N) between the two atoms.

Reaction with burning magnesium

Burning magnesium continues to burn in nitrogen, forming a pale yellow solid, magnesium nitride.

When magnesium nitride is warmed with water, ammonia gas is produced — a useful confirmation that nitrogen was present.

Reaction with oxygen

At the very high temperature of a lightning flash or an electric arc, nitrogen combines with oxygen to form nitrogen(II) oxide, which then turns brown in air.

Uses of nitrogen

• Manufacture of ammonia by the Haber process.
• Dilutes atmospheric oxygen so that burning is controlled rather than violent.
• A source of fertiliser once fixed by nitrogen-fixing bacteria in root nodules.
• Liquid nitrogen is used as a refrigerant, for example in preserving semen and other biological samples.
• Filling electric light bulbs to prevent oxidation of the filament.
• Filling food packets and vehicle tyres, where an inert atmosphere is needed.

Section 7Ammonia

Ammonia (NH₃) is a compound of nitrogen and hydrogen. The nitrogen atom carries a lone pair of electrons, which lets ammonia form hydrogen bonds with water — this is why it is the most soluble gas known. It is a colourless gas with a sharp, pungent smell.

Laboratory preparation of ammonia

Ammonia is prepared by heating an ammonium salt with an alkali. The usual pair is ammonium chloride and calcium hydroxide (slaked lime).

Using sodium hydroxide instead:

The ionic equation for any ammonium salt with any alkali is:

Why calcium oxide is used as the drying agent

Ammonia is an alkaline gas, so it reacts with the usual drying agents and cannot be dried by them:

Only calcium oxide (quicklime), which is itself basic, dries ammonia without reacting with it. The gas is collected by upward delivery in an inverted jar because it is less dense than air, and it cannot be collected over water because it is far too soluble.

Physical properties of ammonia

• Colourless gas with a choking, pungent smell.
• Less dense than air.
• Extremely soluble in water — the most soluble of all common gases.
• Easily liquefied; the liquid absorbs a lot of heat as it evaporates, which is why it is used as a refrigerant.

The fountain experiment

This experiment demonstrates how soluble ammonia is. A dry round-bottomed flask is filled with ammonia and fitted with two tubes carrying clips. It is clamped above a trough of water to which a little red litmus solution has been added.

What happens and why

1. One clip is opened and a few drops of water are allowed into the flask.
2. That small amount of water dissolves a very large volume of ammonia, so a partial vacuum is created in the flask.
3. When the second clip is opened, atmospheric pressure forces water up the jet tube in a fountain until the flask is full.
4. The water turns blue, proving that ammonia solution is alkaline.

Chemical properties of ammonia

1. Alkaline nature

Ammonia is the only common alkaline gas. It turns damp red litmus paper blue.

2. Reaction with hydrogen chloride — the test for ammonia

3. Reaction with acids (neutralisation)

Ammonia neutralises acids to form ammonium salts, many of which are fertilisers.

4. Combustion

Ammonia does not burn in air, but it burns in pure oxygen with a greenish-yellow flame:

In the presence of a hot platinum catalyst, the products are different — nitrogen(II) oxide is formed instead of nitrogen. This is catalytic oxidation, the first step in making nitric acid.

The nitrogen(II) oxide is then oxidised in air to brown nitrogen(IV) oxide, and finally absorbed in water to give nitric acid (the Ostwald process).

5. Reaction with heated copper(II) oxide

Ammonia acts as a reducing agent: black copper(II) oxide is reduced to brown copper metal.

Observations: the black solid turns brown, colourless droplets of water condense at the cool end of the tube, and a colourless gas that does not support combustion is collected.

6. Reaction with metal-ion solutions

Aqueous ammonia supplies OH⁻ ions and so precipitates metal hydroxides.

Section 8The Haber process

This is the industrial manufacture of ammonia by the direct combination of nitrogen and hydrogen. Nitrogen is obtained by the fractional distillation of liquid air; hydrogen comes from the cracking of alkanes or from natural gas and steam.

Conditions and the reasons behind them

The process step by step

1. The mixture of nitrogen and hydrogen is passed through a purifier to remove impurities that would poison the catalyst.
2. It is compressed to about 200–500 atmospheres.
3. It passes into the catalytic chamber over heated iron at about 450 °C.
4. The hot mixture passes through a heat exchanger and cooler, where the ammonia condenses to a liquid and is run off.
5. Unreacted nitrogen and hydrogen are recycled, so there are no dangerous by-products and little waste.

Uses of ammonia

• Manufacture of nitrogenous fertilisers — ammonium sulphate, ammonium nitrate, urea.
• Manufacture of nitric acid by the Ostwald process.
• In laundry work, where it removes grease and softens water.
• As a refrigerant in large cold stores.
• Manufacture of hydrazine, used as a rocket fuel.
• Manufacture of nylon, dyes and explosives.

Oxides of nitrogen and nitrates — a quick reference

Action of heat on nitrates depends on the reactivity of the metal:

All nitrates are soluble in water, and nitrate ions are identified by the brown ring test.

Section 9Nitrogen and ammonia — questions and answers

Section 10Sulphur and its allotropes

Occurrence

Sulphur occurs both free and combined. It is found free in large underground deposits in Louisiana and Texas (USA), Sicily and Poland. Its important mineral compounds are:

• Iron pyrites — FeS₂ ("fool's gold")
• Gypsum — CaSO₄·2H₂O
• Mirabilite (Glauber's salt) — Na₂SO₄·10H₂O
• Galena — PbS, and zinc blende — ZnS
• As hydrogen sulphide in natural gas and volcanic springs

Sulphur is also a constituent of plant and animal protein — it is the sulphur in hair and eggs that gives burning hair and rotten eggs their characteristic smell. Elemental sulphur is not toxic to humans, but it is a good fungicide. Pure sulphur is a brittle, crystalline, yellow solid whose molecules are puckered rings of eight atoms, S₈.

Extraction of sulphur: the Frasch process

The deposits lie about 200 metres underground beneath layers of quicksand, so ordinary mining is impossible. Three concentric pipes are sunk down to the sulphur bed.

1. Superheated water at 170 °C is forced down the outermost pipe. It melts the sulphur (melting point 119 °C).
2. Hot compressed air at about 16 atmospheres is blown down the innermost pipe.
3. The air whips the molten sulphur into a light froth and the pressure forces it up the middle pipe into a reservoir, where it solidifies.

Sulphur obtained by this process is about 99.5 % pure and needs no further refining.

Allotropes of sulphur

Sulphur has two main crystalline allotropes — rhombic and monoclinic — and two non-crystalline forms, plastic and amorphous sulphur.

Plastic sulphur

Formed when boiling sulphur is poured in a thin stream into a beaker of cold water. It is a brown, rubbery, elastic solid. It is unstable and slowly changes back into rhombic sulphur at room temperature.

The temperature of 96 °C, above which monoclinic sulphur is the stable form and below which rhombic sulphur is stable, is called the transition temperature.

Effect of heat on sulphur

• At 119 °C sulphur melts to a mobile, pale yellow liquid of S₈ rings.
• Above about 160 °C the rings open up and join into very long chains. The liquid darkens and becomes so viscous that the test tube can be turned upside down without it pouring out — the maximum is near 187 °C.
• Above 200 °C the long chains are broken by the heat, so the liquid becomes mobile again and turns dark reddish-brown.
• At 444 °C sulphur boils, giving a reddish-brown vapour.

Chemical properties of sulphur

1. Combustion

Sulphur burns in air with a blue flame to form sulphur(IV) oxide. In pure oxygen the flame is brighter and a little sulphur(VI) oxide is also formed.

2. Reaction with metals

Sulphur combines directly with heated metals to form sulphides. Both products below are black.

3. Reaction with non-metals

4. Reaction with acids

With hot concentrated nitric acid, sulphur is oxidised to sulphuric acid and the acid is reduced to brown nitrogen(IV) oxide:

Hot concentrated sulphuric acid oxidises sulphur to sulphur(IV) oxide:

Hydrochloric acid has no effect on sulphur, because it is not an oxidising agent.

Uses of sulphur

• Manufacture of sulphuric acid by the contact process — by far the largest use.
• As a fungicide, especially for dusting vines and controlling powdery mildew.
• Vulcanisation (hardening) of rubber for tyres.
• Making bleaching agents for the paper industry.
• Manufacture of fireworks, matches, gunpowder, dyes and sulphur ointments for skin disease.

Section 11The contact process and sulphuric acid

The contact process is the industrial manufacture of sulphuric(VI) acid. It has four stages.

Stage 1: production of sulphur(IV) oxide

Stage 2: purification

The mixture of sulphur(IV) oxide and excess air is washed and passed through electrostatic precipitators to remove dust and arsenic compounds, which would otherwise poison the catalyst.

Stage 3: conversion to sulphur(VI) oxide

Stage 4: conversion to sulphuric acid

Sulphur(VI) oxide is passed up an absorption tower against a shower of concentrated sulphuric acid, forming oleum:

The oleum is then carefully diluted with the correct amount of water in the dilution chamber:

Properties of concentrated sulphuric acid

1. Dehydrating property

This is the ability of the acid to remove the elements of water — hydrogen and oxygen in the ratio 2:1 — from a compound that contains no water at all.

Added to sugar in a beaker, the acid removes hydrogen and oxygen, leaving a black, spongy column of carbon that rises out of the beaker with a hiss of steam.

Blue crystals of hydrated copper(II) sulphate turn white as the water of crystallisation is removed:

2. Drying agent (hygroscopic property)

The acid absorbs free moisture from the atmosphere. If it is left in an open beaker for two days, its volume is found to have increased and it becomes more dilute. This is why it is used to dry gases that do not react with it.

3. Oxidising property

Hot concentrated sulphuric acid oxidises metals and non-metals, and is itself reduced to sulphur(IV) oxide.

Properties of dilute sulphuric acid

Dilute sulphuric acid shows the ordinary reactions of an acid.

The reaction of dilute sulphuric acid with calcium and lead stops after a short time, because the insoluble sulphates CaSO₄ and PbSO₄ coat the metal and prevent further attack.

Uses of sulphuric acid

• Manufacture of fertilisers — ammonium sulphate and superphosphate (the biggest single use).
• Making synthetic fibres, detergents, soaps, paints, pigments and dyestuffs.
• Refining petroleum and making petrochemicals.
• Extraction of metals (metallurgy) and pickling — cleaning metal surfaces before plating.
• Electrolyte in lead-acid car batteries and accumulators.
• As a drying agent for gases in the laboratory.

Section 12Sulphur(IV) oxide

Laboratory preparation

Sulphur(IV) oxide is usually prepared by adding dilute hydrochloric acid to sodium sulphite. It can also be made by warming copper turnings with concentrated sulphuric acid.

The gas is dried over concentrated sulphuric acid and collected by downward delivery because it is denser than air.

Physical properties

• Colourless gas with a strong, irritating, choking smell (like a struck match).
• Denser than air.
• Readily soluble in water, forming sulphurous acid, H₂SO₃.
• It turns damp blue litmus paper red and then slowly bleaches it.

Chemical properties

1. Acidic nature

2. Reducing property — and the test for the gas

Sulphur(IV) oxide reduces orange acidified potassium dichromate(VI) to green chromium(III) ions. This is the test for the gas.

It also decolourises purple acidified potassium manganate(VII):

3. Bleaching property

Sulphurous acid bleaches by reduction: it takes oxygen away from the coloured dye.

Because oxygen from the air slowly restores the dye, this bleaching is temporary. Old newspapers and straw hats bleached with sulphur dioxide turn yellow-brown again after some time in the sun.

4. Oxidising property

With stronger reducing agents, sulphur(IV) oxide itself acts as an oxidising agent.

Distinguishing sulphates from sulphites

Uses of sulphur(IV) oxide

• Intermediate in the manufacture of sulphuric acid by the contact process.
• Fumigation, to kill insect pests, because the gas is poisonous.
• Bleaching straw, sponges, wool, silk and paper.
• Preservation of fruit juices, dried fruit and wine (it kills bacteria and prevents browning).

Sulphur(IV) oxide as a pollutant

• It dissolves in rain water to form acid rain, which damages crops, kills fish in lakes and corrodes iron roofs and limestone buildings.
• Acid rain lowers the pH of the soil, so plants cannot absorb nutrients properly and growth is affected.
• Inhaled, it irritates the eyes and throat and causes bronchitis and other respiratory infections.

Section 13Sulphur — questions and answers

Section 14Carbon and its allotropes

Carbon occurs free as graphite and diamond, and in impure forms such as wood charcoal, animal charcoal, coke and lampblack. Combined, it is present in carbon dioxide, carbonates, shells and every organic compound.

Allotropy is the existence of an element in more than one form in the same physical state. Carbon has three allotropes: diamond and graphite, which are crystalline, and amorphous carbon, which is non-crystalline.

Diamond

Diamond is the hardest natural substance known. Each carbon atom is joined by strong covalent bonds to four other carbon atoms, forming a rigid tetrahedral giant structure. The bonding forces are equally strong in every direction, and the atoms are closely packed, so diamond is denser than graphite. There are no free electrons, so diamond does not conduct electricity.

Physical properties of diamond

• Very high melting and boiling point, because the many strong covalent bonds need a great deal of energy to break.
• Bad conductor of electricity — no free electrons. (It is, however, an excellent conductor of heat.)
• Colourless, transparent and sparkling, with a very high refractive index.
• Extremely hard and dense.

Uses of diamond

• As a gemstone in jewellery, because of its high refractive index and sparkle.
• For cutting glass, drilling rocks and in grinding and abrasive tools, because of its hardness.
• In heat sinks for electronic devices, because it conducts heat so well.

Graphite

In graphite, each carbon atom uses only three of its four outer electrons in bonding. The fourth is delocalised and free to move, so graphite conducts electricity. The atoms form flat hexagonal rings arranged in layers, and the layers are held together only by weak van der Waals forces, so they slide over each other easily.

Physical properties of graphite

• Good conductor of electricity, because of its delocalised electrons.
• Soft, slippery and greasy, because the layers slide over each other.
• High melting and boiling point, because the covalent bonds within each layer are very strong.
• Dark grey, opaque, and less dense than diamond.

Uses of graphite

• Mixed with clay to make the "lead" in pencils — the layers rub off onto the paper.
• As electrodes in electrolytic cells and dry cells, because it conducts and is chemically inert.
• As a lubricant for machinery working at high temperatures, where oil would burn away.
• As a moderator in nuclear reactors.

Amorphous carbon

In amorphous carbon the atoms are arranged disorderly, with no regular crystal shape. It does not occur naturally but is made in several forms.

1. Charcoal

• Animal charcoal — made by heating bones in a limited supply of air. Used to decolourise crude sugar by adsorbing the brown impurities.
• Sugar charcoal — made by the destructive distillation of sugar cane, or when concentrated sulphuric acid dehydrates sugar. Used to adsorb gases.
• Wood charcoal — made by the destructive distillation of wood. Used in gas masks and water filters because it adsorbs poisonous gases and impurities, and as a domestic fuel.

2. Coke

Formed when coal is heated in the absence of air. It is used as a fuel in furnaces and ovens, as a reducing agent in the extraction of metals from their ores, and as a source of gaseous fuel.

3. Lampblack (soot)

Formed when petroleum, kerosene, turpentine, natural gas and other hydrocarbons burn in a limited supply of air. It is used in making Indian ink, black shoe polish, carbon paper and printer's ink, and is added to rubber to strengthen vehicle tyres.

Reducing properties of carbon

Carbon is a good reducing agent.

Section 15Carbon(IV) oxide

Carbon dioxide makes up about 0.03 % of the atmosphere. It is also produced by respiration, combustion, fermentation and the decay of organic matter.

Laboratory preparation

Carbon dioxide is prepared by the action of dilute hydrochloric acid on marble chips (calcium carbonate).

Physical properties

• Colourless, odourless gas with a faintly acidic taste.
• Denser than air — collected by downward delivery.
• Fairly soluble in water, forming weak carbonic acid.
• Easily solidified under pressure to "dry ice", which sublimes directly to a gas.

Chemical properties

1. It neither burns nor supports combustion — it puts out a burning splint.
2. It dissolves in water to form weak carbonic acid, turning damp blue litmus faintly red.
   CO_2(g) + H₂O_(l) ⇌ H₂CO_3(aq)
3. As an acidic oxide it reacts with alkalis to form carbonates.
   2NaOH_(aq) + CO_2(g) → Na₂CO_3(aq) + H₂O_(l)
4. Burning magnesium continues to burn in carbon dioxide. The heat of the burning magnesium decomposes the gas; magnesium is oxidised to white magnesium oxide, and black specks of carbon are deposited on the sides of the jar.
   2Mg_(s) + CO_2(g) → 2MgO_(s) + C_(s)

Test for carbon dioxide

Bubble the gas through lime water (calcium hydroxide solution). A white precipitate of calcium carbonate forms, so the lime water turns milky.

If excess carbon dioxide is passed through, the milkiness disappears because soluble calcium hydrogen carbonate is formed.

Uses of carbon(IV) oxide

• In fire extinguishers, because it is denser than air and does not support combustion, so it smothers the flame.
• Making fizzy drinks and soda water.
• Solid carbon dioxide (dry ice) as a refrigerant for transporting food, ice cream and vaccines.
• In baking bread and cakes — yeast or baking powder releases the gas, which makes the dough rise.
• In fruit preservation, by creating an atmosphere low in oxygen and high in carbon dioxide.

A warning about carbon(II) oxide

When charcoal burns in a limited supply of air, the poisonous gas carbon monoxide is formed instead of carbon dioxide.

Carbon monoxide is colourless and odourless, and it combines with haemoglobin in red blood cells about 200 times more readily than oxygen does, forming carboxyhaemoglobin. The blood can then no longer carry oxygen, and the person suffocates without ever feeling short of breath. This is why a charcoal stove must never be left burning in a closed, unventilated room.

Section 16Chapter summary

Section 17End of topic questions

Work through all four parts under exam conditions — about 1 hour 30 minutes. The suggested answers follow in Section 18, so resist the temptation to scroll ahead.

Section 18Suggested answers to the end of topic questions