Chapter 1 · Chemical Reactions and Equations

1.2 Types of Chemical Reactions

Combine, split, swap: sort reactions into four families — combination, decomposition, displacement and double displacement — and tell whether a reaction gives out or takes in energy: $\mathrm{CaO + H_{2}O \rightarrow Ca(OH)_{2}}$.

In the last section you learnt to write and balance a chemical equation. There are millions of different reactions, though, and nobody remembers them one by one. Chemists do what a librarian does with books: they sort reactions into a few families by asking a simple question — what happens to the partners? Do two substances join into one? Does one substance break into several? Does an element push another out of its compound? Or do two compounds swap partners?

Once you can recognise the family, you can often predict the products of a reaction you have never seen, and you know what to look for in the laboratory. In this section you will meet the four families — combination, decomposition, displacement and double displacement — and a second way of sorting reactions that matters just as much in everyday life: whether a reaction gives out energy (like the LPG flame and the respiration that keeps you warm) or takes in energy (like the lime kiln and the green leaf).

Combination Reactions

Put a few lumps of quicklime (calcium oxide) in a beaker and slowly add water. The lumps hiss, crumble and swell, and the beaker becomes too hot to hold. A white suspension of slaked lime (calcium hydroxide) is formed: $$\mathrm{CaO(s) + H_{2}O(l) \rightarrow Ca(OH)_{2}(aq) + \text{heat}}$$ Two reactants have joined to give a single product.

📘 Definition 1.5 — Combination reaction
A combination reaction is a reaction in which two or more reactants (elements or compounds) combine to form a single product. In symbols: $\mathrm{A + B \rightarrow AB}$.

The reactants may be two elements, an element and a compound, or two compounds:

Calcium hydroxide is the main ingredient of whitewash, the cheap white coating used on walls and tree trunks. After it is painted on, it slowly reacts with carbon dioxide in the air; over two or three days a thin, hard layer of calcium carbonate forms, which gives the wall a shiny finish: $$\mathrm{Ca(OH)_{2}(aq) + CO_{2}(g) \rightarrow CaCO_{3}(s) + H_{2}O(l)}$$ Curiously, calcium carbonate is the same substance as marble and limestone — the starting material from which quicklime is made, as you will see in the next subsection.

✏️ Example 1 — From quicklime to a whitewashed wall
A painter mixes quicklime with water in a bucket to prepare whitewash. (a) Name the substance formed and write the balanced equation with state symbols. (b) Why does the bucket become hot? (c) Write the equation for the reaction on the wall over the next few days. Is it also a combination reaction?
(a) Slaked lime, calcium hydroxide: $$\mathrm{CaO(s) + H_{2}O(l) \rightarrow Ca(OH)_{2}(aq)}$$ Check: Ca $1=1$; O $1+1=2$; H $2=2$. ✓
(b) The reaction gives out a large amount of heat to its surroundings (the water and the bucket). It is an exothermic reaction — a word defined properly in the third subsection.
(c) $$\mathrm{Ca(OH)_{2}(aq) + CO_{2}(g) \rightarrow CaCO_{3}(s) + H_{2}O(l)}$$
Type No. Two reactants give two products, so it is not a combination reaction. (It is the reaction of a base with a non-metal oxide, which you will meet again in Chapter 2 — it is also why lime water turns milky with carbon dioxide.) Only reactions with a single product are combinations.

Decomposition Reactions

Run a combination backwards and you get the opposite family. When limestone (calcium carbonate) is heated strongly in a lime kiln, it breaks down into quicklime and carbon dioxide — the first step in making cement and lime for whitewash: $$\mathrm{CaCO_{3}(s) \xrightarrow{\text{heat}} CaO(s) + CO_{2}(g)}$$

📘 Definition 1.6 — Decomposition reaction
A decomposition reaction is a reaction in which a single reactant breaks down to give two or more simpler products. In symbols: $\mathrm{AB \rightarrow A + B}$. Breaking the bonds of the reactant needs energy, which may be supplied as heat (thermal decomposition), electricity (electrolytic decomposition, or electrolysis) or light (photolytic decomposition).
🔥 Thermal decompositionenergy as heat$\mathrm{CaCO_{3}}$, $\mathrm{FeSO_{4}}$ and $\mathrm{Pb(NO_{3})_{2}}$ break down when heated strongly.
⚡ Electrolytic decompositionenergy as electricityWater, with a little acid added to make it conduct, splits into hydrogen and oxygen.
☀️ Photolytic decompositionenergy as lightSilver chloride and silver bromide turn grey in sunlight as silver metal forms.

Heating ferrous sulphate. Green crystals of ferrous sulphate (iron(II) sulphate) are heated in a dry boiling tube. First they lose their water of crystallisation and the green colour fades; on stronger heating a reddish-brown solid, ferric oxide (iron(III) oxide), is left, and the gases sulphur dioxide and sulphur trioxide are given off with the smell of burning sulphur. (Do this only in a fume cupboard or a well-ventilated room.) $$\mathrm{2FeSO_{4}(s) \xrightarrow{\text{heat}} Fe_{2}O_{3}(s) + SO_{2}(g) + SO_{3}(g)}$$

✏️ Example 2 — Balancing the ferrous sulphate equation
Balance the skeletal equation $\mathrm{FeSO_{4} \rightarrow Fe_{2}O_{3} + SO_{2} + SO_{3}}$ and state two observations that show a chemical change has taken place.
Atom table
ElementLeftRight
Fe12
S11 + 1 = 2
O43 + 2 + 3 = 8
Try 2 $\mathrm{Fe_{2}O_{3}}$ needs 2 Fe, so write $\mathrm{2FeSO_{4}}$. This also gives 2 S and 8 O on the left. $$\mathrm{2FeSO_{4} \rightarrow Fe_{2}O_{3} + SO_{2} + SO_{3}}$$
Check Fe $2=2$; S $2=1+1$; O $8=3+2+3$. ✓ One coefficient fixed everything.
Observations The green colour of the crystals changes (a reddish-brown solid remains), and a gas with a choking smell of burning sulphur is given off.

Heating lead nitrate. When white lead nitrate powder is heated in a boiling tube, brown fumes of nitrogen dioxide are seen, oxygen is released, and a yellowish residue of lead(II) oxide is left behind. (Nitrogen dioxide is poisonous; again, work in a fume cupboard.)

✏️ Example 3 — Lead nitrate: a decomposition with three products
Balance $\mathrm{Pb(NO_{3})_{2} \rightarrow PbO + NO_{2} + O_{2}}$ and add state symbols and the condition.
Nitrogen Here the nitrate group breaks apart, so it cannot be kept as a unit. N: 2 on the left, 1 on the right, so write $\mathrm{2NO_{2}}$: $\mathrm{Pb(NO_{3})_{2} \rightarrow PbO + 2NO_{2} + O_{2}}$.
Oxygen Left: 6 O. Right: $1+4+2=7$ O — one too many. $\mathrm{O_{2}}$ comes in pairs, so the quickest fix is to allow half a molecule: $\mathrm{Pb(NO_{3})_{2} \rightarrow PbO + 2NO_{2} + \tfrac{1}{2}O_{2}}$, which has $1+4+1=6$ O.
Clear the fraction Multiply every coefficient by 2: $$\mathrm{2Pb(NO_{3})_{2}(s) \xrightarrow{\text{heat}} 2PbO(s) + 4NO_{2}(g) + O_{2}(g)}$$
Check Pb $2=2$; N $4=4$; O $12=2+8+2$. ✓ The brown colour in the tube is the $\mathrm{NO_{2}}$.

Electrolysis of water. Pass an electric current through water containing a few drops of dilute sulphuric acid, using two carbon electrodes, and bubbles rise at both. The gas collected over the negative electrode is hydrogen (it burns with a pop when a lit splint is brought near), and the gas over the positive electrode is oxygen (it relights a glowing splint). The volume of hydrogen is twice the volume of oxygen, exactly as the coefficients predict: $$\mathrm{2H_{2}O(l) \xrightarrow{\text{electric current}} 2H_{2}(g) + O_{2}(g)}$$

Light and silver salts. White silver chloride spread on a watch glass and left in sunlight slowly turns grey, because silver metal is formed. Pale yellow silver bromide behaves in the same way: $$\mathrm{2AgCl(s) \xrightarrow{\text{sunlight}} 2Ag(s) + Cl_{2}(g)}$$ $$\mathrm{2AgBr(s) \xrightarrow{\text{sunlight}} 2Ag(s) + Br_{2}(g)}$$ This is the chemistry of black-and-white photography: wherever light falls on a film coated with silver bromide, tiny grains of dark silver form, building up the image. It is also why silver salts are stored in dark brown bottles.

⚠️ Count the reactants, not the atoms

The family name depends on the number of substances, not on how many molecules the equation shows. $\mathrm{2H_{2}O \rightarrow 2H_{2} + O_{2}}$ is a decomposition because only one substance, water, is on the left, even though its coefficient is 2. Likewise $\mathrm{2H_{2} + O_{2} \rightarrow 2H_{2}O}$ is a combination because there is only one product. And a decomposition does not need a burner: the energy may come from light or from electricity.

Energy Changes: Exothermic and Endothermic Reactions

The quicklime bucket became hot; the lime kiln has to be kept hot for the limestone to break down. Every reaction involves an energy change, because bonds are broken (which needs energy) and new bonds are formed (which releases energy). Which effect wins decides whether the surroundings warm up or cool down.

📘 Definition 1.7 — Exothermic and endothermic reactions
An exothermic reaction is one in which energy, usually as heat (sometimes also as light or sound), is given out to the surroundings; the surroundings get warmer. An endothermic reaction is one in which energy (as heat, light or electricity) is taken in from the surroundings; either the mixture gets colder, or energy must be supplied all the time for the reaction to continue.

Exothermic — energy out

  • Burning of natural gas: $$\mathrm{CH_{4}(g) + 2O_{2}(g) \rightarrow CO_{2}(g) + 2H_{2}O(g)}$$
  • Respiration: the glucose from digested food combines with oxygen in your cells and releases the energy that keeps you alive and warm: $$\mathrm{C_{6}H_{12}O_{6}(aq) + 6O_{2}(aq) \rightarrow 6CO_{2}(aq) + 6H_{2}O(l)}$$
  • The decomposition of vegetable matter into compost: a garden compost heap is warm inside.
  • Quicklime with water.

Endothermic — energy in

  • Most decomposition reactions: limestone in a kiln, ferrous sulphate and lead nitrate in a boiling tube (heat); water (electricity); silver chloride (light).
  • Photosynthesis: green leaves absorb the energy of sunlight to make glucose.
  • Mixing solid barium hydroxide with solid ammonium chloride in a test tube: the mixture turns slushy and the tube feels cold to the touch.

Notice that respiration is a kind of slow, controlled burning, which is why both give out carbon dioxide and water. Notice too how a combination and the reverse decomposition are linked: if forming calcium hydroxide releases energy, breaking it apart again would need energy. That is why decompositions are usually endothermic.

✏️ Example 4 — Exothermic or endothermic?
Classify each process as exothermic or endothermic, with a reason: (a) a matchstick burning; (b) dissolved glucose giving energy to a sprinter's muscles; (c) heating limestone to make lime; (d) silver bromide on a photographic film turning grey in light; (e) a heap of fallen leaves and kitchen waste turning into compost.
Test Is energy given out to the surroundings (exothermic), or must it be supplied (endothermic)?
(a), (b) Exothermic. Burning gives out heat and light. Respiration releases the energy the muscles use (and makes the sprinter hot).
(c), (d) Endothermic. Both are decompositions: the kiln must supply heat continuously, and silver bromide breaks down only while light falls on it.
(e) Exothermic. The decomposition of vegetable matter by microbes gives out heat; the inside of a compost heap is noticeably warm. (So not every decomposition is endothermic — this one is the decay of food material, which is chemically a slow oxidation, like respiration.)
Interactive 2 — Energy Out or Energy In?

Choose a reaction and press ▶ Play (or drag the slider). On the left, the energy-level diagram shows the energy stored in the reactants and in the products; the ball travels along the reaction path. On the right you see where the energy goes: red arrows carry energy out to the surroundings, blue arrows carry it in. The levels are drawn to show the idea, not to scale.

type—
energy—
form of energy—
what you notice—

Displacement and Double Displacement Reactions

Clean two iron nails with sandpaper. Dip one into a test tube of blue copper sulphate solution and keep the other aside for comparison. After about 20 minutes the dipped nail is coated with a reddish-brown layer, and the blue colour of the solution has faded towards pale green. Iron has pushed copper out of copper sulphate: $$\mathrm{Fe(s) + CuSO_{4}(aq) \rightarrow FeSO_{4}(aq) + Cu(s)}$$ The brown coat is copper metal; the pale green solution is iron sulphate.

📘 Definition 1.8 — Displacement reaction
A displacement reaction is a reaction in which a more reactive element takes the place of (displaces) a less reactive element from its compound, usually in solution. In symbols: $\mathrm{A + BC \rightarrow AC + B}$.

Zinc and lead also displace copper from its compounds: $$\mathrm{Zn(s) + CuSO_{4}(aq) \rightarrow ZnSO_{4}(aq) + Cu(s)}$$ $$\mathrm{Pb(s) + CuCl_{2}(aq) \rightarrow PbCl_{2}(aq) + Cu(s)}$$ But a copper wire left in iron sulphate solution stays clean for ever. Which way a displacement goes is decided by how reactive the two metals are.

📐 Principle 1.2 — Who displaces whom
A metal displaces another metal from a solution of its salt only if it is more reactive than that metal. The common metals, from most to least reactive, are $$\mathrm{K \gt Na \gt Ca \gt Mg \gt Al \gt Zn}$$ $$\mathrm{\gt Fe \gt Pb \gt Cu \gt Hg \gt Ag \gt Au}$$ (the reactivity series, studied in detail in Section 3.1). If the free metal is below the metal in the salt, there is no reaction. (Potassium, sodium and calcium are so reactive that they react with the water of the solution instead, so they are not used for these tests.)
✏️ Example 5 — Will it displace?
Predict whether a reaction occurs. If it does, write the balanced equation with state symbols and say what you would see. (a) A strip of zinc in copper sulphate solution. (b) A copper strip in zinc sulphate solution. (c) A copper wire in silver nitrate solution.
Rule Compare positions in the reactivity series: the free metal must be above the metal in the salt.
(a) Zn is above Cu — yes. $$\mathrm{Zn(s) + CuSO_{4}(aq) \rightarrow ZnSO_{4}(aq) + Cu(s)}$$ A reddish-brown coat of copper forms on the zinc and the blue solution fades to colourless (zinc sulphate solution is colourless).
(b) Cu is below Zn — no reaction. The strip and the solution are unchanged.
(c) Cu is above Ag — yes. Silver nitrate is $\mathrm{AgNO_{3}}$ and copper forms $\mathrm{Cu(NO_{3})_{2}}$, so two nitrate units are needed: $$\mathrm{Cu(s) + 2AgNO_{3}(aq) \rightarrow Cu(NO_{3})_{2}(aq) + 2Ag(s)}$$ Shiny grey crystals of silver grow on the wire and the colourless solution slowly turns blue. Check: Cu $1=1$; Ag $2=2$; N $2=2$; O $6=6$. ✓

In the last family, nobody is pushed out alone — both compounds swap partners. Mix colourless sodium sulphate solution with colourless barium chloride solution and a white precipitate of barium sulphate appears at once: $$\mathrm{Na_{2}SO_{4}(aq) + BaCl_{2}(aq) \rightarrow BaSO_{4}(s) + 2NaCl(aq)}$$ Barium has swapped its chloride partner for sulphate, and sodium has swapped sulphate for chloride. The yellow lead iodide reaction of Section 1.1 is another example.

📘 Definition 1.9 — Double displacement and precipitation reactions
A double displacement reaction is a reaction in which two compounds exchange their ions (their positive and negative parts) to form two new compounds. In symbols: $\mathrm{AB + CD \rightarrow AD + CB}$. A reaction that produces a precipitate is called a precipitation reaction; many double displacement reactions in solution are precipitation reactions.
CombinationA + B → ABone product
DecompositionAB → A + Bone reactant
DisplacementA + BC → AC + Ban element pushes another out
Double displacementAB + CD → AD + CBtwo compounds swap partners
✏️ Example 6 — Swapping partners
(a) Silver nitrate solution is added to sodium chloride solution and a white precipitate forms. Name it and write the equation. (b) Copper sulphate solution is mixed with sodium hydroxide solution and a pale blue precipitate of copper(II) hydroxide forms. Write the balanced equation. (c) Name the type of each reaction.
(a) Swap partners: silver goes with chloride, sodium with nitrate. The white precipitate is silver chloride: $$\mathrm{AgNO_{3}(aq) + NaCl(aq) \rightarrow AgCl(s) + NaNO_{3}(aq)}$$
(b) Swap Copper goes with hydroxide, sodium with sulphate: skeleton $\mathrm{CuSO_{4} + NaOH \rightarrow Cu(OH)_{2} + Na_{2}SO_{4}}$.
(b) Balance Na: 1 vs 2, so $\mathrm{2NaOH}$, which also gives the 2 OH in $\mathrm{Cu(OH)_{2}}$: $$\mathrm{CuSO_{4}(aq) + 2NaOH(aq) \rightarrow Cu(OH)_{2}(s) + Na_{2}SO_{4}(aq)}$$ Check: Cu 1, S 1, O $4+2=6$, Na 2, H 2 on each side. ✓
(c) Both are double displacement reactions; because each forms a precipitate, both are also precipitation reactions.
Interactive 1 — Reaction-Type Sorter: Join, Split or Swap?

Watch the atoms of the reactants change partners. Each ball is one atom; atoms that stay together as one unit (such as $\mathrm{SO_{4}}$ or $\mathrm{NO_{3}}$) move as a group. Decide which family the reaction belongs to and click your answer — the pattern and the explanation appear below.

reactant substances—
product substances—
pattern?
your score0 / 0
⚠️ Common mistakes in naming reaction types

1. Displacement or double displacement? Look at the reactants. If one of them is a free element (Fe, Zn, Cu…), it is a displacement. If both are compounds that exchange parts, it is a double displacement.

2. A precipitate is not a type of its own. “Precipitation reaction” describes what you see; the family is usually double displacement. Say both when asked.

3. Not every reaction fits exactly one box. Burning magnesium is a combination and exothermic; heating limestone is a decomposition and endothermic. The four families describe what happens to the partners; “exothermic/endothermic” describes the energy — they are two separate questions.

4. Writing products that do not exist. Before you write products for a displacement, check the reactivity series: $\mathrm{Cu + FeSO_{4}}$ gives no reaction, so there is no equation to balance.

Exercises

I
Combination and decomposition reactions.
  • 1 What is a combination reaction? Give an example, with a balanced equation, in which two compounds combine.
  • 2 Write balanced equations with state symbols for: (a) the burning of coal (carbon) in air; (b) sodium metal burning in chlorine gas to form solid sodium chloride; (c) aluminium powder heated in chlorine gas to form aluminium chloride, $\mathrm{AlCl_{3}}$.
  • 3 Why should sacks of quicklime be kept dry, away from rain?
  • 4 A wall is whitewashed with calcium hydroxide. A week later the coating is hard and shiny. Explain, with a balanced equation.
  • 5 Name the three kinds of decomposition reaction according to the form of energy supplied, and give one balanced equation for each.
  • 6 In the electrolysis of water, $24\ \mathrm{mL}$ of gas is collected over the negative electrode. Name this gas. Name the gas collected over the positive electrode and give its volume. Give a test for each gas.
  • 7 Why are silver chloride and silver bromide kept in dark-coloured bottles? Write the equation for what happens to silver chloride in sunlight, and the colour change you would see.
  • 8 Baking soda (sodium hydrogencarbonate) decomposes on heating, which is why it makes cakes rise. Balance $\mathrm{NaHCO_{3} \rightarrow Na_{2}CO_{3} + H_{2}O + CO_{2}}$.
  • 9 Balance $\mathrm{FeSO_{4} \rightarrow Fe_{2}O_{3} + SO_{2} + SO_{3}}$ and $\mathrm{Pb(NO_{3})_{2} \rightarrow PbO + NO_{2} + O_{2}}$ without looking back, then name the coloured gas given off in the second reaction.
  • 10 Applied. In a lime kiln, $50\ \mathrm{kg}$ of pure limestone gives $28\ \mathrm{kg}$ of quicklime. What mass of carbon dioxide goes up the chimney? At the same rate, how much carbon dioxide is released per tonne ($1000\ \mathrm{kg}$) of limestone?
Check: type your balanced equation and count the atoms on each side
Write formulas with ordinary digits, coefficients in front, and -> or = for the arrow, e.g. 2Pb(NO3)2 -> 2PbO + 4NO2 + O2.
1 Two or more reactants form a single product. Two compounds: $\mathrm{CaO(s) + H_{2}O(l) \rightarrow Ca(OH)_{2}(aq)}$, or $\mathrm{NH_{3}(g) + HCl(g) \rightarrow NH_{4}Cl(s)}$.
2 (a) $\mathrm{C(s) + O_{2}(g) \rightarrow CO_{2}(g)}$. (b) $\mathrm{2Na(s) + Cl_{2}(g) \rightarrow 2NaCl(s)}$. (c) Cl comes in 2s and 3s, LCM 6: $\mathrm{2Al(s) + 3Cl_{2}(g) \rightarrow 2AlCl_{3}(s)}$.
3 Quicklime reacts with water in a strongly exothermic combination reaction. Wet quicklime gets very hot (it can even set fire to the sacks) and turns into slaked lime, so it is no longer quicklime.
4 Calcium hydroxide reacts slowly with carbon dioxide in the air to form a thin hard layer of calcium carbonate, which is shiny: $$\mathrm{Ca(OH)_{2}(aq) + CO_{2}(g) \rightarrow CaCO_{3}(s) + H_{2}O(l)}$$
5 Thermal: $\mathrm{CaCO_{3}(s) \xrightarrow{\text{heat}} CaO(s) + CO_{2}(g)}$. Electrolytic: $\mathrm{2H_{2}O(l) \xrightarrow{\text{electric current}} 2H_{2}(g) + O_{2}(g)}$. Photolytic: $\mathrm{2AgCl(s) \xrightarrow{\text{sunlight}} 2Ag(s) + Cl_{2}(g)}$.
6 Hydrogen collects over the negative electrode; oxygen over the positive electrode, with half the volume, $12\ \mathrm{mL}$. Hydrogen burns with a pop when a burning splint is brought near; oxygen relights a glowing splint.
7 Light decomposes them into silver and the halogen, spoiling the chemical. $\mathrm{2AgCl(s) \xrightarrow{\text{sunlight}} 2Ag(s) + Cl_{2}(g)}$: the white solid turns grey.
8 Na: 1 vs 2, so $\mathrm{2NaHCO_{3}}$; then H $2=2$, C $2=1+1$, O $6=3+1+2$: $$\mathrm{2NaHCO_{3}(s) \xrightarrow{\text{heat}} Na_{2}CO_{3}(s) + H_{2}O(g) + CO_{2}(g)}$$ The carbon dioxide gas makes the cake rise.
9 $\mathrm{2FeSO_{4} \rightarrow Fe_{2}O_{3} + SO_{2} + SO_{3}}$ and $\mathrm{2Pb(NO_{3})_{2} \rightarrow 2PbO + 4NO_{2} + O_{2}}$. The brown gas is nitrogen dioxide, $\mathrm{NO_{2}}$.
10 By conservation of mass, carbon dioxide $=50-28=22\ \mathrm{kg}$. Per tonne: $\tfrac{22}{50}\times1000=440\ \mathrm{kg}$ of carbon dioxide.
II
Exothermic and endothermic reactions.
  • 1 Define an exothermic reaction and an endothermic reaction, with one example of each.
  • 2 Why is respiration called an exothermic reaction? Write its balanced equation.
  • 3 Classify as exothermic or endothermic: (a) burning LPG; (b) photosynthesis; (c) electrolysis of water; (d) quicklime with water; (e) heating lead nitrate.
  • 4 A student mixes solid barium hydroxide with solid ammonium chloride in a test tube and holds the bottom of the tube. What does she feel, and what type of reaction (in terms of energy) is this?
  • 5 Explain why most decomposition reactions are endothermic, using the decomposition of limestone as your example.
  • 6 Applied. A “self-heating” can of soup has a sealed inner compartment holding quicklime and a separate pouch of water. Pressing a button breaks the pouch. Explain how the soup is heated, name the reaction type, and suggest why quicklime is chosen.
1 Exothermic: energy (mainly heat) is given out, e.g. burning methane. Endothermic: energy is taken in, e.g. the decomposition of calcium carbonate on heating.
2 The oxidation of glucose in the cells releases energy: $$\mathrm{C_{6}H_{12}O_{6}(aq) + 6O_{2}(aq) \rightarrow 6CO_{2}(aq) + 6H_{2}O(l)} + \text{energy}$$
3 (a) exothermic; (b) endothermic (light absorbed); (c) endothermic (electrical energy needed); (d) exothermic; (e) endothermic (heat needed).
4 The tube feels cold: the reaction takes heat from its surroundings, including her hand. It is endothermic.
5 Breaking the reactant into simpler substances means breaking bonds, which needs energy. Limestone decomposes only while it is kept strongly heated in the kiln; when heating stops, the decomposition stops.
6 When the water reaches the quicklime they react, $\mathrm{CaO(s) + H_{2}O(l) \rightarrow Ca(OH)_{2}(s)}$, a combination reaction that is strongly exothermic; the heat passes through the inner wall into the soup. Quicklime is cheap, safe to store while dry, and gives out a lot of heat with a small amount of water.
III
Displacement and double displacement reactions.
  • 1 An iron nail is kept in copper sulphate solution for about 20 minutes. Describe two changes you would observe and write the balanced equation.
  • 2 Use the reactivity series to decide whether each reaction happens; write an equation where it does: (a) magnesium ribbon in copper sulphate solution; (b) iron nail in zinc sulphate solution; (c) zinc granules in iron(II) sulphate solution; (d) silver wire in copper sulphate solution.
  • 3 Applied. Why should copper sulphate solution (used as a fungicide spray in vineyards and orchards) never be stored in an iron bucket?
  • 4 What is a double displacement reaction? Illustrate with the reaction between sodium sulphate and barium chloride solutions, stating what you observe.
  • 5 Iron(III) chloride solution is added to sodium hydroxide solution and a reddish-brown precipitate of iron(III) hydroxide, $\mathrm{Fe(OH)_{3}}$, forms. Write the balanced equation with state symbols.
  • 6 Calcium chloride solution and sodium carbonate solution give a white precipitate of calcium carbonate. Balance $\mathrm{CaCl_{2} + Na_{2}CO_{3} \rightarrow CaCO_{3} + NaCl}$ and name the type of reaction in two ways.
  • 7 Sodium hydroxide reacts with hydrochloric acid to give sodium chloride and water. Explain why this can be called a double displacement reaction.
  • 8 A strip of a metal M is placed in blue copper sulphate solution. After an hour the solution is colourless and the strip is covered with a reddish-brown deposit. Is M zinc or silver? Explain, and write the equation.
1 A reddish-brown coat forms on the nail, and the blue colour of the solution fades (to pale green). $\mathrm{Fe(s) + CuSO_{4}(aq) \rightarrow FeSO_{4}(aq) + Cu(s)}$.
2 (a) Yes: $\mathrm{Mg(s) + CuSO_{4}(aq) \rightarrow MgSO_{4}(aq) + Cu(s)}$. (b) No — iron is below zinc. (c) Yes: $\mathrm{Zn(s) + FeSO_{4}(aq) \rightarrow ZnSO_{4}(aq) + Fe(s)}$. (d) No — silver is below copper.
3 Iron is more reactive than copper, so it would displace copper from the solution: the bucket would slowly dissolve (forming iron sulphate) and the spray would lose its copper sulphate.
4 Two compounds exchange ions to form two new compounds. A white precipitate of barium sulphate forms at once: $$\mathrm{Na_{2}SO_{4}(aq) + BaCl_{2}(aq) \rightarrow BaSO_{4}(s) + 2NaCl(aq)}$$
5 Swap partners, then balance OH and Cl in threes: $$\mathrm{FeCl_{3}(aq) + 3NaOH(aq) \rightarrow Fe(OH)_{3}(s) + 3NaCl(aq)}$$ Check: Fe 1, Cl 3, Na 3, O 3, H 3 on each side.
6 $$\mathrm{CaCl_{2}(aq) + Na_{2}CO_{3}(aq) \rightarrow CaCO_{3}(s) + 2NaCl(aq)}$$ This is a double displacement reaction and a precipitation reaction.
7 The two compounds exchange partners: the hydrogen of $\mathrm{HCl}$ joins the hydroxide of $\mathrm{NaOH}$ to form water, while sodium joins chloride. $\mathrm{NaOH(aq) + HCl(aq) \rightarrow NaCl(aq) + H_{2}O(l)}$. (This is the neutralisation reaction of Chapter 2.)
8 Zinc. Only a metal more reactive than copper can displace it; silver is less reactive than copper, zinc is more reactive. Zinc sulphate solution is colourless: $\mathrm{Zn(s) + CuSO_{4}(aq) \rightarrow ZnSO_{4}(aq) + Cu(s)}$.
IV
Name the type of each reaction (and say whether it is exothermic or endothermic where you can), then the mixed problems.
  • 1 $\mathrm{2Mg(s) + O_{2}(g) \rightarrow 2MgO(s)}$
  • 2 $$\mathrm{2Pb(NO_{3})_{2}(s) \xrightarrow{\text{heat}} 2PbO(s) + 4NO_{2}(g) + O_{2}(g)}$$
  • 3 $$\mathrm{Zn(s) + H_{2}SO_{4}(aq) \rightarrow ZnSO_{4}(aq) + H_{2}(g)}$$
  • 4 $$\mathrm{Pb(NO_{3})_{2}(aq) + 2KI(aq) \rightarrow PbI_{2}(s) + 2KNO_{3}(aq)}$$
  • 5 $\mathrm{2AgBr(s) \xrightarrow{\text{light}} 2Ag(s) + Br_{2}(g)}$
  • 6 $$\mathrm{Fe_{2}O_{3}(s) + 2Al(s) \rightarrow Al_{2}O_{3}(s) + 2Fe(l)}$$ (the thermite reaction, which gives so much heat that the iron formed is molten)
  • 7 Write one balanced equation for (a) a decomposition reaction that gives three products; (b) a combination reaction in which an element combines with a compound. (Hint: look back at Exercise II.4 of Section 1.1.)
  • 8 Challenging. A white solid A, heated strongly, gives a colourless gas B that turns lime water milky, and leaves a white solid C. C reacts vigorously with water, giving out heat, to form D, which is used for whitewashing. Identify A, B, C and D and write the three equations involved (including the lime-water test).
1 Combination (one product); exothermic (heat and dazzling light).
2 Thermal decomposition (one reactant, heat supplied); endothermic.
3 Displacement: zinc displaces hydrogen from the acid. (The flask warms slightly, so it is exothermic.)
4 Double displacement, and a precipitation reaction (yellow $\mathrm{PbI_{2}}$).
5 Photolytic decomposition; endothermic (light energy absorbed).
6 Displacement: aluminium, more reactive, displaces iron from its oxide; strongly exothermic.
7 (a) e.g. $\mathrm{2FeSO_{4}(s) \xrightarrow{\text{heat}} Fe_{2}O_{3}(s) + SO_{2}(g) + SO_{3}(g)}$, or the decomposition of baking soda. (b) e.g. $\mathrm{2SO_{2}(g) + O_{2}(g) \rightarrow 2SO_{3}(g)}$ (Section 1.1 Exercise II.4), or $\mathrm{2CO(g) + O_{2}(g) \rightarrow 2CO_{2}(g)}$.
8 A = calcium carbonate, B = carbon dioxide, C = calcium oxide (quicklime), D = calcium hydroxide (slaked lime). $$\mathrm{CaCO_{3}(s) \xrightarrow{\text{heat}} CaO(s) + CO_{2}(g)}$$ $$\mathrm{CaO(s) + H_{2}O(l) \rightarrow Ca(OH)_{2}(aq)}$$ $$\mathrm{Ca(OH)_{2}(aq) + CO_{2}(g) \rightarrow CaCO_{3}(s) + H_{2}O(l)}$$ The last equation is the lime-water test: insoluble calcium carbonate makes it milky.
🎯 Exam corner BoardBoard · A–RBoard · Case-basedJEE Main / NEET

E1. Board · 1 mark Which of the following is a photolytic decomposition reaction? (A) $\mathrm{CaCO_{3} \xrightarrow{\text{heat}} CaO + CO_{2}}$   (B) $\mathrm{2AgBr \xrightarrow{\text{light}} 2Ag + Br_{2}}$   (C) $\mathrm{2H_{2}O \xrightarrow{\text{current}} 2H_{2} + O_{2}}$   (D) $\mathrm{Fe + CuSO_{4} \rightarrow FeSO_{4} + Cu}$

Idea “Photolytic” means broken down by light. (A) is thermal and (C) electrolytic decomposition; (D) is a displacement.
Answer (B).

E2. Board · A–R Assertion (A): The blue colour of copper sulphate solution fades when an iron nail is kept in it. Reason (R): Iron is more reactive than copper, so it displaces copper from copper sulphate solution. Choose: (A) Both A and R are true and R is the correct explanation of A. (B) Both A and R are true but R is not the correct explanation of A. (C) A is true but R is false. (D) A is false but R is true.

Idea The blue colour is due to copper sulphate. As iron displaces copper, copper sulphate is used up and pale green iron sulphate forms — so R explains A.
Answer (A).

E3. Board · 3 marks A student heats a small quantity of white lead nitrate powder in a boiling tube held with a test-tube holder. (a) State two observations. (b) Write the balanced chemical equation with state symbols. (c) Name the type of reaction and say whether it is exothermic or endothermic.

(a) Brown fumes (nitrogen dioxide) are given off, and a yellowish solid (lead(II) oxide) is left in the tube.
(b) $$\mathrm{2Pb(NO_{3})_{2}(s) \xrightarrow{\text{heat}} 2PbO(s) + 4NO_{2}(g) + O_{2}(g)}$$
Answer (c) Thermal decomposition; endothermic, because heat must be supplied continuously.

E4. Board · Case-based Before Diwali, Meena's family whitewashes their house. Her uncle buys white lumps of a substance X from a hardware shop and drops them into a drum of water; the water hisses and becomes very hot, and a milky suspension Y is formed. After painting, the walls take two to three days to dry to a shiny white finish. (i) Identify X and Y and write the balanced equation for the reaction in the drum. (2 marks) (ii) What type of reaction is this, with respect to (a) the reactants and products, and (b) energy? (1 mark) (iii) Write the equation for the reaction that gives the shiny finish. (1 mark)

(i) X = calcium oxide (quicklime), $\mathrm{CaO}$; Y = calcium hydroxide (slaked lime), $\mathrm{Ca(OH)_{2}}$. $$\mathrm{CaO(s) + H_{2}O(l) \rightarrow Ca(OH)_{2}(aq)}$$
(ii) (a) Combination reaction (two reactants, one product); (b) exothermic (the water becomes very hot).
Answer (iii) $$\mathrm{Ca(OH)_{2}(aq) + CO_{2}(g) \rightarrow CaCO_{3}(s) + H_{2}O(l)}$$ The thin layer of calcium carbonate is shiny.

E5. JEE Main / NEET Foundation. Which pair will not react? (A) zinc and copper sulphate solution   (B) iron and copper sulphate solution   (C) copper and iron(II) sulphate solution   (D) copper and silver nitrate solution

Idea A free metal displaces a metal from its salt only if it is higher in the reactivity series: Zn and Fe are above Cu; Cu is above Ag; but Cu is below Fe.
Answer (C).

E6. Board · 2 marks In the electrolysis of water, why is the volume of gas collected at one electrode double that at the other? Name the two gases.

Idea The balanced equation $\mathrm{2H_{2}O(l) \rightarrow 2H_{2}(g) + O_{2}(g)}$ shows that two molecules of hydrogen are formed for every one molecule of oxygen. At the same temperature and pressure equal numbers of gas molecules occupy equal volumes.
Answer So hydrogen (at the negative electrode) has twice the volume of oxygen (at the positive electrode).
🔬 Research–Teaching NexusResearch-tutored · optional, about 30 minutes

Research question. Making cement starts with a decomposition reaction — heating limestone. How much carbon dioxide does that one reaction release, and why is it so hard to avoid?

Why it matters. Cement is the most-used manufactured material on Earth, and the cement industry produces roughly 7–8% of the world's carbon dioxide emissions. China and India are the two largest producers. More than half of a cement plant's carbon dioxide does not come from burning fuel at all: it comes from the decomposition $\mathrm{CaCO_{3} \rightarrow CaO + CO_{2}}$ itself. Engineers in Europe, India and elsewhere are testing carbon capture at cement kilns and “low-clinker” cements that need less lime.

What to do.

  1. Write the balanced equation for the thermal decomposition of calcium carbonate, with state symbols and the condition.
  2. Using atomic masses $\mathrm{Ca}=40\ \mathrm{u}$, $\mathrm{C}=12\ \mathrm{u}$, $\mathrm{O}=16\ \mathrm{u}$, find the formula masses of $\mathrm{CaCO_{3}}$, $\mathrm{CaO}$ and $\mathrm{CO_{2}}$, and check that mass is conserved.
  3. Find the mass of carbon dioxide released (a) per tonne of limestone heated, and (b) per tonne of quicklime produced.
  4. Cement “clinker” contains about 65% calcium oxide by mass, all of it made from limestone. Estimate the carbon dioxide released by this reaction alone for each tonne of clinker.
  5. Explain why switching the kiln to solar or electric heating would not remove this carbon dioxide.

Tools: Calculator; optionally a spreadsheet.

Hand in: A half-page note with your equation, the mass check, the three figures, and your answer to step 5.

Going further. Find out how a whitewashed wall or a lime-mortar building slowly takes carbon dioxide back from the air (the whitewashing reaction). Roughly what fraction of the kiln's carbon dioxide could be re-absorbed this way over many years?

Instructor note: $\mathrm{CaCO_{3}(s) \xrightarrow{\text{heat}} CaO(s) + CO_{2}(g)}$; $100\ \mathrm{u} = 56\ \mathrm{u} + 44\ \mathrm{u}$. Per tonne of limestone $0.44\ \mathrm{t}$ of carbon dioxide; per tonne of quicklime $44/56\approx0.79\ \mathrm{t}$. Per tonne of clinker $0.65\times0.79\approx0.51\ \mathrm{t}$ (published process-emission factors are about $0.5\ \mathrm{t}$). Step 5: the carbon dioxide comes from the carbonate itself, not from fuel, so it is released whatever the heat source; only capture, or using less lime, avoids it.

🧪 Self-check quiz

0/5 correct

Q1. The reaction $\mathrm{CaO(s) + H_{2}O(l) \rightarrow Ca(OH)_{2}(aq)}$ is

Q2. When ferrous sulphate crystals are heated strongly, the gases given off are

Q3. Which metal can displace copper from copper sulphate solution but cannot displace iron from iron(II) sulphate solution?

Q4. Sodium sulphate solution is mixed with barium chloride solution. You observe

Q5. Which of these is an endothermic process?

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