Types Of Chemical Reactions

Watch a wall being whitewashed before a festival: the coating goes on dull, yet two or three days later the wall gleams with a shiny finish. Drop a little water on a lump of quicklime and it hisses, crackles and turns dangerously hot. Dip an iron nail into blue copper sulphate solution and the nail slowly turns brown while the blue itself fades away. Each of these is a chemical reaction — and every reaction you will meet in Class 10 fits into a small set of patterns that this lesson trains you to recognise on sight.

Definition: A combination reaction is a reaction in which two or more reactants combine to form a single product (pattern: A + B → AB). A decomposition reaction is its exact reverse — a single compound breaks down into two or more simpler products (pattern: AB → A + B) using heat, light or electricity.

Definition: In a displacement reaction, a more reactive element displaces (pushes out) a less reactive element from its compound (pattern: A + BC → AC + B). In a double displacement reaction, two compounds exchange their ions (pattern: AB + CD → AD + CB), often producing an insoluble solid called a precipitate.

The four patterns at a glance — count the reactants and products first, then look for free elements.

Combination reactions — many reactants, one product

The give-away of a combination reaction is the product count: however many substances go in, only one product comes out.

The NCERT flagship example is quicklime reacting with water:

CaO(s) + H2O(l) → Ca(OH)2(aq) + heat

Calcium oxide (quicklime) reacts vigorously with water to form calcium hydroxide (slaked lime), releasing such a large amount of heat that the container becomes hot to touch. Two reactants, one product — a combination reaction that is also exothermic.

This reaction has a beautiful follow-up that boards love. A suspension of slaked lime is used for whitewashing walls. The calcium hydroxide then reacts slowly with the carbon dioxide present in air to form a thin layer of calcium carbonate on the walls:

Ca(OH)2(aq) + CO2(g) → CaCO3(s) + H2O(l)

Calcium carbonate is formed two to three days after whitewashing and gives the walls their shiny finish. Interestingly, marble has the same chemical formula, CaCO3.

Other combination reactions to remember:

• Burning of coal: C(s) + O2(g) → CO2(g)
• Formation of water: 2H2(g) + O2(g) → 2H2O(l)
• Burning of magnesium ribbon: 2Mg(s) + O2(g) → 2MgO(s) — it burns with a dazzling white flame, leaving white magnesium oxide powder

A combination can be element + element, element + compound, or compound + compound. The only rule that never bends is the single product.

Decomposition reactions — one compound splits apart

A decomposition reaction is the mirror image of combination: one reactant, two or more products. Breaking chemical bonds needs an energy supply, so NCERT names each kind of decomposition after the energy used — heat, electricity or light.

1. Thermal decomposition (energy from heat). Heat light green ferrous sulphate crystals in a dry boiling tube:

2FeSO4(s) --heat--> Fe2O3(s) + SO2(g) + SO3(g)

The crystals first lose their water of crystallisation and their green colour changes; on stronger heating they decompose to a brown residue of ferric oxide (Fe2O3), while sulphur dioxide and sulphur trioxide escape with the characteristic smell of burning sulphur.

Heating limestone is a thermal decomposition used on an industrial scale (for example in cement manufacture):

CaCO3(s) --heat--> CaO(s) + CO2(g)

Heating lead nitrate gives a dramatic observation:

2Pb(NO3)2(s) --heat--> 2PbO(s) + 4NO2(g) + O2(g)

The brown fumes you see are of nitrogen dioxide (NO2) — a favourite one-mark observation question.

2. Electrolytic decomposition (energy from electricity). Pass an electric current through water (with a few drops of dilute sulphuric acid added to make the water a good conductor):

2H2O(l) --electric current--> 2H2(g) + O2(g)

Hydrogen gas collects at the cathode and oxygen at the anode, and the volume of hydrogen is double the volume of oxygen — exactly the 2 : 1 ratio the balanced equation predicts.

3. Photolytic decomposition (energy from sunlight). Keep white silver chloride in sunlight and it slowly turns grey as silver metal forms:

2AgCl(s) --sunlight--> 2Ag(s) + Cl2(g)

Light yellow silver bromide behaves the same way: 2AgBr(s) --sunlight--> 2Ag(s) + Br2(g). These light-sensitive reactions are used in black and white photography.

Because decomposition reactions must absorb energy to happen, they are typically endothermic.

Displacement reactions — the more reactive element wins

In a displacement reaction a free element attacks a compound and kicks out a less reactive element, taking its place: A + BC → AC + B.

The classic NCERT activity: dip an iron nail in copper sulphate solution and leave it for about 20 minutes.

Fe(s) + CuSO4(aq) → FeSO4(aq) + Cu(s)

Two observations fetch the marks: (i) the nail acquires a brownish coating of copper, and (ii) the blue colour of the solution fades towards light green, because blue copper sulphate is being replaced by light green ferrous sulphate. Iron has displaced copper from its salt, which proves iron is more reactive than copper.

Other standard examples:

• Zn(s) + CuSO4(aq) → ZnSO4(aq) + Cu(s)
• Pb(s) + CuCl2(aq) → PbCl2(aq) + Cu(s)
• Zinc with dilute acid: Zn(s) + 2HCl(aq) → ZnCl2(aq) + H2(g) — here zinc displaces hydrogen, and brisk bubbles of hydrogen gas appear

Zinc and lead are more reactive elements than copper, so both displace copper from its compounds. The reverse never happens: a copper strip left in zinc sulphate or ferrous sulphate solution shows no reaction, because a less reactive element cannot displace a more reactive one. (You will meet the full reactivity series in the chapter on Metals and Non-metals.)

Double displacement reactions — an exchange of ions

Here two compounds react by exchanging their ions: AB + CD → AD + CB. Mix solutions of sodium sulphate and barium chloride:

Na2SO4(aq) + BaCl2(aq) → BaSO4(s) + 2NaCl(aq)

A white precipitate of barium sulphate appears instantly, formed when the barium ions meet the sulphate ions; sodium chloride stays dissolved in the solution. Any reaction in which an insoluble solid (precipitate) forms is called a precipitation reaction, and such precipitation reactions are double displacement reactions.

Common mistake alert: students regularly confuse displacement with double displacement. In displacement, a free element replaces another element in a single compound. In double displacement, two compounds swap their ions and no free element needs to appear at all. Scan the equation: a lone element on the reactant side signals displacement; two compounds facing two compounds signals double displacement.

The energy view — exothermic and endothermic reactions

Cutting across all four patterns is a second, independent classification: does the reaction give out heat or take in energy?

Exothermic reactions release heat along with the products. NCERT's own examples:

• Burning of natural gas: CH4(g) + 2O2(g) → CO2(g) + 2H2O(g) + energy
• Respiration: the glucose from digested food combines with oxygen in the cells of our body and provides energy — C6H12O6(aq) + 6O2(aq) → 6CO2(aq) + 6H2O(l) + energy
• The decomposition of vegetable matter into compost
• Quicklime reacting with water (seen above)

Endothermic reactions absorb energy. Most decomposition reactions are endothermic — that is exactly why they must be fed heat, light or electricity to keep going. Photosynthesis, which absorbs sunlight, is nature's great endothermic reaction.

On an energy diagram, an exothermic reaction ends with products at lower energy than the reactants — the difference escapes as heat. An endothermic reaction ends with products at higher energy — the difference was absorbed from the surroundings.

Exothermic: products settle at lower energy and the difference is released as heat. Endothermic: products sit at higher energy because energy was absorbed.

Exam tip: "exothermic" is a property, not a fifth type. Every reaction is classified twice — once by pattern (combination, decomposition, displacement, double displacement) and once by energy (exothermic or endothermic). Burning coal, for instance, is both a combination reaction and exothermic.

Work through these examples in order. The early ones train the basic "count and classify" habit; the Tricky ones rehearse the observation-plus-reason questions that board papers love.

Example 1 — Quicklime meets water

Q: What is observed when water is added to a small amount of quicklime? Write the chemical equation and name the type of reaction.
Given: Reactants — calcium oxide (quicklime) and water.
Formula: Combination pattern A + B → AB.
Solve: CaO(s) + H2O(l) → Ca(OH)2(aq) + heat. The mixture hisses and the vessel becomes very hot; a single product, calcium hydroxide (slaked lime), is formed from two reactants.
Answer: It is a combination reaction, and since a large amount of heat is released it is also exothermic. Check: exactly one product on the right — combination confirmed.

Example 2 — Limestone in the kiln

Q: Limestone is heated strongly. Classify the reaction, write the balanced equation, and give one industrial use of the solid product.
Given: Reactant — calcium carbonate (limestone); energy supplied — heat.
Formula: Decomposition pattern AB → A + B.
Solve: CaCO3(s) --heat--> CaO(s) + CO2(g). One compound has split into two simpler products with the help of heat.
Answer: Thermal decomposition (an endothermic change). The product calcium oxide (quicklime) is used in the manufacture of cement. Check: one reactant, two products — decomposition confirmed.

Example 3 — The iron nail in blue solution

Q: An iron nail is kept immersed in copper sulphate solution for about 20 minutes. State two observations, write the equation, and state which of the two metals is more reactive.
Given: Iron nail; blue copper sulphate solution.
Solve: Fe(s) + CuSO4(aq) → FeSO4(aq) + Cu(s). Observation 1: the nail acquires a brownish coating of copper. Observation 2: the blue colour of the solution fades towards light green (ferrous sulphate forms).
Answer: A displacement reaction; iron is more reactive than copper, because it displaces copper from its salt solution. Check: a free element (Fe) stands on the reactant side — displacement confirmed.

Example 4 — Spot the precipitate

Q: Solutions of sodium sulphate and barium chloride are mixed. What is observed? Write the balanced equation and classify the reaction.
Given: Na2SO4 solution and BaCl2 solution.
Formula: Double displacement pattern AB + CD → AD + CB.
Solve: Na2SO4(aq) + BaCl2(aq) → BaSO4(s) + 2NaCl(aq). The barium ions and sulphate ions meet to form insoluble barium sulphate, while sodium chloride remains dissolved.
Answer: A white precipitate of BaSO4 appears — a double displacement reaction, specifically a precipitation reaction. Check: two compounds exchanged partners and an insoluble solid formed.

Example 5 — Tricky: classify all four in one go

Q: Classify each reaction, giving a reason:
(A) 2H2O(l) --electricity--> 2H2(g) + O2(g)
(B) Zn(s) + 2HCl(aq) → ZnCl2(aq) + H2(g)
(C) 2Mg(s) + O2(g) → 2MgO(s)
(D) AgNO3(aq) + NaCl(aq) → AgCl(s) + NaNO3(aq)
Solve: (A) One compound splits into two products using electricity — electrolytic decomposition. (B) The free element zinc replaces hydrogen from the acid — displacement. (C) Two reactants give a single product — combination (also exothermic). (D) Two compounds exchange ions and insoluble white silver chloride precipitates — double displacement (precipitation).

Reaction	Type	Key clue
A	Decomposition	one reactant → many products
B	Displacement	free element replaces an element
C	Combination	many reactants → one product
D	Double displacement	ions exchanged, precipitate forms

Answer: A — decomposition; B — displacement; C — combination; D — double displacement. Tip: count reactants and products first, then look for free elements.

Example 6 — Tricky: the green crystals puzzle

Q: Light green crystals are heated in a dry boiling tube. The colour changes, a brown solid is left behind, and a gas with the smell of burning sulphur is released. Identify the crystals, write the balanced equation, and classify the reaction.
Given: Colour clues — green reactant, brown residue; smell of burning sulphur.
Solve: The clues fit ferrous sulphate crystals (FeSO4·7H2O). On heating they first lose their water of crystallisation; on stronger heating: 2FeSO4(s) --heat--> Fe2O3(s) + SO2(g) + SO3(g). The brown residue is ferric oxide, and the smell comes from the sulphur oxides.
Answer: Thermal decomposition of ferrous sulphate. Watch the trap: two different gases (SO2 and SO3) are released, not one — many students lose a mark by writing only SO2.

Example 7 — Tricky: gas volumes in electrolysis

Q: During the electrolysis of water, 40 mL of gas collects over one electrode and 20 mL over the other. Identify each gas with its electrode, explain the 2 : 1 ratio, and state why a few drops of dilute sulphuric acid are added to the water.
Given: Gas volumes — 40 mL and 20 mL.
Formula: 2H2O(l) --electric current--> 2H2(g) + O2(g).
Solve: The balanced equation gives 2 volumes of hydrogen for every 1 volume of oxygen. So the 40 mL gas is hydrogen, collected at the cathode, and the 20 mL gas is oxygen, collected at the anode. Dilute sulphuric acid is added to make water a good conductor of electricity.
Answer: 40 mL of H2 at the cathode and 20 mL of O2 at the anode; the ratio simply follows the 2 : 1 mole ratio in the equation. Check: 40 = 2 × 20 — consistent.

Example 8 — Tricky: exothermic is not a type of its own

Q: A student writes: "CH4 + 2O2 → CO2 + 2H2O releases a lot of heat, so it is a combination reaction." Point out and correct the error. Also state, with reason, whether respiration is exothermic or endothermic.
Solve: Releasing heat only makes a reaction exothermic; the pattern decides its type. Here the products are TWO substances (CO2 and H2O), so it cannot be a combination reaction — it is the burning (combustion) of natural gas, an exothermic oxidation. For the second part: during respiration, glucose combines with oxygen in our cells — C6H12O6 + 6O2 → 6CO2 + 6H2O + energy — and energy is released.
Answer: The reaction is exothermic but not a combination reaction, because more than one product forms. Respiration is exothermic, since the breakdown of glucose releases energy. Check: combination demands a single product — this has two.

The five reaction types you must classify and give examples for in the exam.

• Combination: two or more reactants form a single product, e.g. CaO + H2O → Ca(OH)2 (releases heat — exothermic).
• Decomposition: a single reactant breaks into two or more products — thermal (CaCO3 → CaO + CO2), electrolytic (2H2O → 2H2 + O2) or photolytic (2AgCl → 2Ag + Cl2).
• Displacement: a more reactive element displaces a less reactive one, e.g. Fe + CuSO4 → FeSO4 + Cu.
• Double displacement: two compounds exchange ions, e.g. Na2SO4 + BaCl2 → BaSO4↓ + 2NaCl (a precipitation reaction).
• Precipitation: a double-displacement giving an insoluble salt (precipitate).
• Exothermic reactions release heat (respiration, combustion); endothermic reactions absorb heat (decomposition of CaCO3).

Cover the answer, recall, then check. 8 cards on classifying reactions with examples.

Q1. Define a combination reaction with an example.
A1. Two or more substances combine to form a single product, e.g. CaO + H2O → Ca(OH)2.

Q2. What are the three types of decomposition reaction?
A2. Thermal (by heat), electrolytic (by electricity) and photolytic (by light).

Q3. Give the reaction when quicklime is added to water.
A3. CaO + H2O → Ca(OH)2 + heat — a combination reaction that is strongly exothermic.

Q4. Why is respiration called an exothermic reaction?
A4. Glucose is oxidised to CO2 and water, releasing energy to keep the body warm.

Q5. Write a displacement reaction of iron with copper sulphate.
A5. Fe + CuSO4 → FeSO4 + Cu; the blue solution turns pale green and copper deposits.

Q6. What is a double displacement reaction?
A6. A reaction in which two compounds exchange their ions to form two new compounds, e.g. Na2SO4 + BaCl2 → BaSO4↓ + 2NaCl.

Q7. What is a precipitation reaction?
A7. A reaction that produces an insoluble solid (precipitate), such as white BaSO4.

Q8. On heating, ferrous sulphate crystals change colour. Why?
A8. 2FeSO4 → Fe2O3 + SO2 + SO3; the green crystals decompose to reddish-brown ferric oxide.