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Chapter 1. Chemical Reactions

Grade 10 Science  ·  CBSE / NCERT  ·  Chapter 1

Chemical Reactions
and Equations

Discover how matter transforms — from burning magnesium to rusting iron — and learn the symbolic language chemists use to describe these changes.

⚗ Chemical Equations ⚡ Types of Reactions 🔄 Oxidation & Reduction 🔩 Corrosion & Rancidity

  📋 Chapter at a Glance
Section Topics Covered Key Terms
1.1 Chemical Equations Word equations, skeletal equations, balancing equations, physical state symbols, reaction conditions Reactant, Product, Balanced Equation, Coefficient, LHS, RHS
1.2 Types of Reactions Combination, Decomposition (thermal/electrolytic/photolytic), Displacement, Double Displacement, Oxidation-Reduction Exothermic, Endothermic, Precipitation, Redox, Electrolysis
1.3 Oxidation in Daily Life Corrosion of metals, rusting of iron, rancidity of food, prevention methods Corrosion, Rancidity, Antioxidant, Oxidation, Tarnishing

  🔑 Key Definitions
Chemical Reaction: A process in which one or more substances (reactants) are converted into one or more new substances (products) with different properties. Indicators include change in state, colour, temperature or evolution of a gas.
Chemical Equation: A symbolic representation of a chemical reaction using chemical formulae and symbols, showing reactants on the left-hand side (LHS) and products on the right-hand side (RHS) separated by an arrow (→).
Balanced Chemical Equation: An equation in which the number of atoms of each element is equal on both the reactant and product sides, satisfying the Law of Conservation of Mass (mass can neither be created nor destroyed in a chemical reaction).
Combination Reaction: A reaction in which two or more substances (elements or compounds) combine together to form a single new product. General form: A + B → AB.
Decomposition Reaction: A reaction in which a single compound breaks down into two or more simpler products. It is the opposite of a combination reaction. General form: AB → A + B.
Displacement Reaction: A reaction in which a more reactive element displaces a less reactive element from its compound in a solution. General form: A + BC → AC + B (where A is more reactive than B).
Double Displacement Reaction: A reaction in which two compounds react by exchanging their ions to form two new compounds. One product is often an insoluble precipitate. General form: AB + CD → AD + CB.
Exothermic Reaction: A chemical reaction in which heat energy is released to the surroundings along with the formation of products, making the reaction mixture warm.
Endothermic Reaction: A chemical reaction in which energy is absorbed from the surroundings. Decomposition reactions typically require energy input in the form of heat, light, or electricity.
Oxidation: The gain of oxygen or loss of hydrogen by a substance during a chemical reaction. Example: 2Cu + O2 → 2CuO (copper gains oxygen, so it is oxidised).
Reduction: The loss of oxygen or gain of hydrogen by a substance during a chemical reaction. In a redox reaction, one substance is always oxidised while another is simultaneously reduced.
Corrosion: The slow and gradual deterioration of a metal due to attack by moisture, oxygen, acids or other substances in its environment. Rusting of iron (forming Fe2O3·xH2O) is the most common example.

  1.1 — Chemical Equations

A chemical reaction takes place around us every day — milk souring, iron rusting, food cooking, and even the process of breathing. Whenever a chemical change occurs, the nature and identity of the initial substances change. To represent these changes concisely, chemists use chemical equations.

A chemical reaction can first be written as a word equation, then converted to a more concise symbolic equation using chemical formulae.

Word Equation Example:
Magnesium + Oxygen  →  Magnesium Oxide
(Reactants)                 (Product)
Symbolic Equation:   Mg + O2 → MgO    (skeletal — needs balancing)

  1.1.1 — Writing a Chemical Equation

Chemical equations use chemical formulae instead of names. The reactants are written on the left of the arrow (→) and the products on the right. A plus sign (+) separates multiple reactants or products. The arrowhead points towards the products.

Symbol Meaning / Usage
→ Arrow showing direction of reaction (reactants → products)
+ Separates multiple reactants (LHS) or multiple products (RHS)
(s) Solid state
(l) Liquid state
(g) Gaseous state
(aq) Aqueous — dissolved in water

  📐 Diagram 1: Burning of Magnesium Ribbon (Activity 1.1)
TONG ↓ Magnesium Ribbon
✦ DAZZLING WHITE FLAME
↓ ash collected
Watch-Glass
White powder = MgO
🔥 BURNER

Magnesium ribbon burns in air. The white ash collected in the watch-glass is magnesium oxide (MgO). Reaction: 2Mg(s) + O2(g) → 2MgO(s)

  1.1.2 — Balanced Chemical Equations

The Law of Conservation of Mass states that mass can neither be created nor destroyed in a chemical reaction. Therefore, the total number of atoms of each element must be the same on both sides of a chemical equation. A skeletal (unbalanced) equation must be balanced by placing appropriate coefficients in front of the formulae — we never change the subscripts inside formulae.

⚠ Rule: Only change COEFFICIENTS (numbers in front of formulae). NEVER alter subscripts inside the formula itself.
Balancing Steps — Example: Fe + H2O → Fe3O4 + H2 Fe atoms H atoms O atoms
Step 1 — Skeletal (unbalanced): Fe + H2O → Fe3O4 + H2 1 vs 3 ✗ 2 vs 2 ✓ 1 vs 4 ✗
Step 2 — Balance O: Fe + 4H2O → Fe3O4 + H2 1 vs 3 ✗ 8 vs 2 ✗ 4 vs 4 ✓
Step 3 — Balance H: Fe + 4H2O → Fe3O4 + 4H2 1 vs 3 ✗ 8 vs 8 ✓ 4 vs 4 ✓
Step 4 — Balance Fe: 3Fe + 4H2O → Fe3O4 + 4H2 ✅ BALANCED 3 = 3 ✓ 8 = 8 ✓ 4 = 4 ✓

  📐 Diagram 2: Zinc + Dilute Sulphuric Acid (Activity 1.3)
Cork
Glass tube
H2 gas ↑

Dilute H2SO4
•••
Zn granules
Observation: Bubbles of H2 gas form around Zn granules.
The flask warms up — exothermic.

Equation:
Zn(s) + H2SO4(aq) → ZnSO4(aq) + H2(g) ✓ Balanced

  1.2 — Types of Chemical Reactions

Chemical reactions involve the breaking and making of bonds between atoms to produce new substances. There are five main types of chemical reactions you need to know for Class 10: combination, decomposition, displacement, double displacement, and oxidation-reduction (redox).

1. Combination

A + B → AB
Two or more substances combine into one single product.

  
2. Decomposition

AB → A + B
A single substance breaks down into simpler products.

  
3. Displacement

A + BC → AC + B
A more reactive element displaces a less reactive one.

  
4. Double Displacement

AB + CD → AD + CB
Ions are exchanged; often forms a precipitate.

  
5. Redox

Oxidation + Reduction happen simultaneously. One substance gains O2, another loses O2.

  1.2.1 — Combination Reaction

In a combination reaction, two or more substances combine to form a single product. These reactions are often exothermic (release heat). A classic example is calcium oxide (quick lime) reacting vigorously with water to produce slaked lime, releasing a large amount of heat.

  📐 Diagram 3: Slaked Lime Formation (Activity 1.4)
Beaker
💧 Water being added slowly
〇〇〇
CaO (Quick Lime)
🌡️ Beaker becomes HOT!		 Reaction:
CaO(s) + H2O(l) → Ca(OH)2(aq) + Heat
(Quick lime)      (Slaked lime)

This is a COMBINATION + EXOTHERMIC reaction.
💡 Do You Know? Slaked lime [Ca(OH)2] is used for whitewashing walls. It reacts slowly with CO2 in air to form calcium carbonate (CaCO3), giving walls a shiny finish after 2–3 days. Interestingly, marble also has the formula CaCO3!
Reaction: Ca(OH)2(aq) + CO2(g) → CaCO3(s) + H2O(l)

  1.2.2 — Decomposition Reaction

Decomposition is the reverse of combination. A single compound breaks down into two or more simpler substances. Energy must be supplied — in the form of heat (thermal decomposition), light (photolytic decomposition), or electricity (electrolytic decomposition). These are all endothermic reactions.

Type Energy Source Example
Thermal Decomposition Heat CaCO3(s) → CaO(s) + CO2(g)  |  2FeSO4(s) → Fe2O3(s) + SO2(g) + SO3(g)
Electrolytic Decomposition Electricity 2H2O(l) → 2H2(g) + O2(g) [Electrolysis of water]
Photolytic Decomposition Sunlight 2AgCl(s) → 2Ag(s) + Cl2(g)  |  2AgBr(s) → 2Ag(s) + Br2(g)

  📐 Diagram 4: Heating Ferrous Sulphate (Activity 1.5)
FeSO4·7H2O
GREEN crystals
HEAT ↑
Fe2O3
RED-BROWN powder
Colour change: Green → Reddish-brown
Pungent smell of burning sulphur

2FeSO4(s) →Heat Fe2O3(s) + SO2(g) + SO3(g)
Thermal decomposition; endothermic.

  📐 Diagram 5: Electrolysis of Water (Activity 1.7)
O2 gas
(½ volume)		   H2 gas
(double volume)
Anode (+)
Graphite rod		 Water +
dil. H2SO4		 Cathode (-)
Graphite rod
⚡ 6V Battery
Reaction: 2H2O(l) → 2H2(g) + O2(g)  |  H2 volume is double that of O2 (2:1 ratio).
H2 burns with a pop sound; O2 relights a glowing splint.

  1.2.3 — Displacement Reaction

In a displacement reaction, a more reactive element replaces a less reactive element from its compound in solution. The more reactive element “pushes out” the less reactive one. Reactivity order plays a key role: Zn > Fe > Cu.

  📐 Diagram 6: Iron Nails in Copper Sulphate Solution (Activity 1.9)
Test Tube A (No Fe nail)
CuSO4 solution
BLUE (intense)
→
20 min
Test Tube B (With Fe nail)
FeSO4 solution
PALE BLUE/GREEN
Fe nail turns BROWN (Cu deposited)
Fe(s) + CuSO4(aq) → FeSO4(aq) + Cu(s)
Iron displaces copper (Fe is more reactive than Cu). Copper is deposited on the iron nail.

  1.2.4 — Double Displacement Reaction

In a double displacement reaction, the positive and negative ions of two compounds exchange places to form two new compounds. These reactions often produce an insoluble precipitate (a solid formed in solution) — such reactions are also called precipitation reactions.

Na2SO4(aq)
Sodium sulphate
(colourless)		 + BaCl2(aq)
Barium chloride
(colourless)		 → BaSO4(s)
WHITE PRECIPITATE		 + 2NaCl(aq)
Sodium chloride
(stays in solution)

Na2SO4(aq) + BaCl2(aq) → BaSO4(s)↓ + 2NaCl(aq) — SO42− and Ba2+ ions combine to form the insoluble BaSO4 precipitate.

  1.2.5 — Oxidation and Reduction (Redox)

Oxidation and reduction always occur together in a redox reaction — one substance is oxidised while another is simultaneously reduced.

Process In terms of Oxygen In terms of Hydrogen Example
Oxidation GAINS oxygen LOSES hydrogen Cu + O2 → CuO (Cu is oxidised)
Reduction LOSES oxygen GAINS hydrogen CuO + H2 → Cu + H2O (CuO is reduced)

  📐 Diagram 7: Redox Reaction — Copper and Copper Oxide (Activity 1.11)
Cu (shiny brown) +  O2  →Heat CuO (black) ← OXIDATION of Cu
↕
CuO (black) +  H2  →Heat Cu (brown) + H2O ← REDUCTION of CuO

In CuO + H2 → Cu + H2O: CuO loses oxygen (reduced); H2 gains oxygen (oxidised). This is a redox reaction.

  1.3 — Effects of Oxidation Reactions in Everyday Life
🔩 1.3.1 Corrosion

When a metal is attacked by substances around it (moisture, oxygen, acids), it corrodes. Rusting of iron is the most familiar example — iron articles that are shiny when new develop a reddish-brown coating (hydrated iron(III) oxide, Fe2O3·xH2O) when left in humid air.

Other examples: black coating on silver (Ag2S), green coating on copper (basic copper carbonate).

Prevention: Painting, galvanising, alloying, applying oil/grease, electroplating.

  
🍟 1.3.2 Rancidity

When fats and oils in food are oxidised, the food becomes rancid — its taste and smell change unpleasantly. This is why old cooking oil smells bad.

Substances that prevent oxidation are called antioxidants (e.g., Vitamin E, BHA, BHT) and are added to packaged foods.

Prevention: Store in airtight containers; flush food packets with nitrogen gas (as done for chips) to replace oxygen and prevent oxidation.

  🔬 Activities (1.1 to 1.11) — Lab Experiments
▶ Activity 1.1 — Burning of Magnesium Ribbon
Materials: Magnesium ribbon (3–4 cm), sandpaper, tongs, spirit lamp, watch-glass.
Procedure: Clean the Mg ribbon with sandpaper. Hold with tongs. Burn over a flame (away from eyes!). Collect ash in watch-glass.
Observation: Mg burns with a dazzling white flame. White powder (MgO) forms in the watch-glass.
Conclusion: Chemical reaction has taken place — change in colour and state observed.
Equation: 2Mg(s) + O2(g) → 2MgO(s) [Combination + Exothermic + Oxidation of Mg]
▶ Activity 1.2 — Lead Nitrate + Potassium Iodide
Materials: Lead nitrate solution, potassium iodide solution, test tube.
Procedure: Add potassium iodide solution to lead nitrate solution in a test tube.
Observation: A bright yellow precipitate of lead iodide (PbI2) forms immediately.
Conclusion: Chemical reaction indicated by formation of a precipitate (change in state).
Equation: Pb(NO3)2(aq) + 2KI(aq) → PbI2(s)↓ + 2KNO3(aq) [Double Displacement / Precipitation reaction]
▶ Activity 1.3 — Zinc + Dilute Sulphuric Acid
Materials: Zinc granules, dilute H2SO4 (or HCl), conical flask. [CAUTION: Handle acid with care.]
Procedure: Add dilute acid to zinc granules in a conical flask.
Observation: Bubbles of gas form around zinc granules. The flask warms up.
Conclusion: Gas evolution and temperature change indicate a chemical reaction. H2 gas is evolved; exothermic reaction.
Equation: Zn(s) + H2SO4(aq) → ZnSO4(aq) + H2(g) [Displacement reaction]
▶ Activity 1.4 — Calcium Oxide + Water (Slaked Lime)
Materials: Calcium oxide (quick lime), water, beaker.
Procedure: Take CaO in a beaker. Slowly add water to it. Touch the beaker.
Observation: The reaction is vigorous. The beaker becomes very hot. Slaked lime (milky white solution) forms.
Conclusion: Combination reaction and exothermic reaction. Large amount of heat released.
Equation: CaO(s) + H2O(l) → Ca(OH)2(aq) + Heat
▶ Activity 1.5 — Heating Ferrous Sulphate Crystals
Materials: FeSO4·7H2O crystals (~2g), dry boiling tube, burner.
Procedure: Note colour of crystals. Heat the boiling tube over a burner. Waft gas gently towards nose.
Observation: Green crystals turn reddish-brown. Pungent smell of burning sulphur (SO2/SO3).
Conclusion: Thermal decomposition (endothermic).
Equation: 2FeSO4(s) →Heat Fe2O3(s) + SO2(g) + SO3(g)
▶ Activity 1.6 — Heating Lead Nitrate
Materials: Lead nitrate powder (~2g), boiling tube, tongs, burner.
Procedure: Hold boiling tube with tongs. Heat lead nitrate powder over a flame.
Observation: Brown fumes of nitrogen dioxide (NO2) are emitted. Yellow lead oxide (PbO) solid remains.
Conclusion: Thermal decomposition reaction.
Equation: 2Pb(NO3)2(s) →Heat 2PbO(s) + 4NO2(g) + O2(g)
▶ Activity 1.7 — Electrolysis of Water
Materials: Plastic mug with carbon electrodes, 6V battery, dilute H2SO4, two inverted test tubes filled with water. [CAUTION: Teacher supervision needed for flame test.]
Procedure: Connect electrodes to 6V battery. Pass electricity through water + acid solution. Collect gas in inverted test tubes.
Observation: Bubbles form at both electrodes. Gas at cathode (H2) is double the volume of gas at anode (O2). H2 burns with a pop; O2 relights a glowing splint.
Equation: 2H2O(l) →Electricity 2H2(g) + O2(g) [Electrolytic decomposition — endothermic]
▶ Activity 1.8 — Silver Chloride in Sunlight
Materials: Silver chloride (~2g), china dish, sunlight.
Procedure: Place white AgCl in a china dish. Leave in sunlight for some time.
Observation: White silver chloride slowly turns grey in sunlight.
Conclusion: Photolytic decomposition — light energy causes decomposition. Silver metal (grey) is formed along with chlorine gas.
Equation: 2AgCl(s) →Sunlight 2Ag(s) + Cl2(g) [Used in black and white photography!]
▶ Activity 1.9 — Iron Nails in Copper Sulphate Solution
Materials: 3 iron nails, copper sulphate solution (10 mL), 2 test tubes (A and B).
Procedure: Take CuSO4 solution in both test tubes. Dip 2 iron nails in test tube B for 20 minutes. Keep one nail aside for comparison.
Observation: In test tube B, the blue colour of CuSO4 fades. The iron nails turn brownish (copper deposits on them).
Conclusion: Displacement reaction — Fe displaces Cu because Fe is more reactive than Cu.
Equation: Fe(s) + CuSO4(aq) → FeSO4(aq) + Cu(s)
▶ Activity 1.10 — Sodium Sulphate + Barium Chloride
Materials: Sodium sulphate solution (3 mL), barium chloride solution (3 mL), test tubes.
Procedure: Take Na2SO4 solution in a test tube. Add BaCl2 solution to it.
Observation: A white, chalky insoluble precipitate immediately forms.
Conclusion: Double displacement and precipitation reaction. BaSO4 is insoluble.
Equation: Na2SO4(aq) + BaCl2(aq) → BaSO4(s)↓ + 2NaCl(aq)
▶ Activity 1.11 — Heating Copper Powder (Oxidation)
Materials: Copper powder (~1g), china dish, wire gauze, tripod stand, burner.
Procedure: Heat copper powder in a china dish.
Observation: Shiny brown copper powder turns black (copper oxide).
Conclusion: Oxidation of copper — copper gains oxygen to form black CuO. If H2 gas is now passed over the heated CuO, the black coating turns brown again (reduction).
Equations: 2Cu(s) + O2(g) →Heat 2CuO(s) [Oxidation]  |  CuO(s) + H2(g) →Heat Cu(s) + H2O(l) [Reduction]

  ✏ Worked Examples
EXAMPLE 1

Balance the equation: Mg + O2 → MgO

Show Solution ▶
Skeletal: Mg + O2 → MgO — O has 2 on LHS, 1 on RHS. Put coefficient 2 for MgO: Mg + O2 → 2MgO — now Mg has 1 on LHS but 2 on RHS. Put coefficient 2 for Mg.
Balanced: 2Mg(s) + O2(g) → 2MgO(s)
Check: Mg = 2=2 ✓, O = 2=2 ✓
EXAMPLE 2

Balance: H2 + Cl2 → HCl

Show Solution ▶
H has 2 on LHS, 1 on RHS. Cl has 2 on LHS, 1 on RHS. Put coefficient 2 for HCl.
Balanced: H2(g) + Cl2(g) → 2HCl(g)
Check: H = 2=2 ✓, Cl = 2=2 ✓ — Combination reaction.
EXAMPLE 3

Balance: Na + H2O → NaOH + H2

Show Solution ▶
H: LHS = 2, RHS = 3. O: LHS = 1, RHS = 1. Na: LHS = 1, RHS = 1. Try 2Na: 2Na + 2H2O → 2NaOH + H2. Check: Na=2=2✓, H=4=4✓, O=2=2✓
Balanced: 2Na(s) + 2H2O(l) → 2NaOH(aq) + H2(g)
EXAMPLE 4

Balance: BaCl2 + Al2(SO4)3 → BaSO4 + AlCl3

Show Solution ▶
Al: 2 on LHS, 1 on RHS → put 2AlCl3. SO4: 3 on LHS, 1 on RHS → put 3BaSO4. Ba needs 3BaCl2. Cl: 6 on LHS, 6 on RHS ✓
Balanced: 3BaCl2(aq) + Al2(SO4)3(aq) → 3BaSO4(s) + 2AlCl3(aq)
Ba=3=3✓, Cl=6=6✓, Al=2=2✓, SO4=3=3✓
EXAMPLE 5

Identify the type: CaO(s) + H2O(l) → Ca(OH)2(aq) + Heat

Show Solution ▶
Two reactants (CaO and H2O) combine to form a single product (Ca(OH)2) → Combination Reaction. Heat is released → also Exothermic.
EXAMPLE 6

Identify: 2FeSO4(s) →Heat Fe2O3(s) + SO2(g) + SO3(g)

Show Solution ▶
One reactant (FeSO4) breaks into multiple products → Decomposition Reaction. Energy (heat) is needed → Endothermic. Specifically, thermal decomposition.
EXAMPLE 7

Identify and explain: Fe(s) + CuSO4(aq) → FeSO4(aq) + Cu(s)

Show Solution ▶
Iron (Fe) displaces copper (Cu) from copper sulphate solution because Fe is more reactive than Cu → Displacement Reaction. The blue colour of CuSO4 fades and the iron nail becomes brownish (copper deposited on it).
EXAMPLE 8

Identify and name: Na2SO4(aq) + BaCl2(aq) → BaSO4(s)↓ + 2NaCl(aq)

Show Solution ▶
Ions exchange: SO42− and Ba2+ combine; Na+ and Cl− combine → Double Displacement Reaction. An insoluble white solid (BaSO4) forms → also a Precipitation Reaction.
EXAMPLE 9

In CuO + H2 → Cu + H2O, identify what is oxidised and what is reduced.

Show Solution ▶
CuO → loses oxygen → CuO is REDUCED (becomes Cu).
H2 → gains oxygen → H2 is OXIDISED (becomes H2O).
This is a Redox Reaction — both oxidation and reduction occur simultaneously.
EXAMPLE 10

Write balanced equations with state symbols: (a) Calcium hydroxide + CO2 → Calcium carbonate + Water; (b) Zinc + Silver nitrate → Zinc nitrate + Silver

Show Solution ▶
(a) Ca(OH)2(aq) + CO2(g) → CaCO3(s) + H2O(l) — Ca=1=1✓, O=3=3✓, H=2=2✓, C=1=1✓
(b) Zn(s) + 2AgNO3(aq) → Zn(NO3)2(aq) + 2Ag(s) — Zn=1=1✓, Ag=2=2✓, N=2=2✓, O=6=6✓
Both are displacement reactions (Zn is more reactive than Ag).
EXAMPLE 11

Why does magnesium ribbon need to be cleaned with sandpaper before burning?

Show Solution ▶
Magnesium ribbon is coated with a thin white layer of magnesium oxide (MgO) on its surface due to contact with air (oxidation in open atmosphere). This layer prevents the Mg from burning. Cleaning with sandpaper removes the MgO layer, exposing clean magnesium metal, which then burns brightly in air.
EXAMPLE 12

Why is the amount of H2 gas collected in Activity 1.7 double that of O2?

Show Solution ▶
From the balanced equation: 2H2O(l) → 2H2(g) + O2(g). For every 2 moles of H2 produced, only 1 mole of O2 is produced. The mole ratio is 2:1. By Avogadro’s law, equal moles of gases at same temp and pressure occupy equal volumes. Therefore, H2 gas produced is exactly double the volume of O2.
EXAMPLE 13

Why is respiration considered an exothermic reaction?

Show Solution ▶
In respiration, glucose (from food) reacts with oxygen to produce carbon dioxide, water and energy:
C6H12O6(aq) + 6O2(aq) → 6CO2(aq) + 6H2O(l) + energy
Since energy is released (as heat and ATP) during this reaction, respiration is an exothermic reaction. This energy keeps our body warm and powers all life processes.

  📝 Practice Set A — Multiple Choice Questions (8 Questions)
Q1. Which of the following is NOT an indicator of a chemical reaction?
(a) Change in state   (b) Change in colour   (c) Change in shape   (d) Evolution of a gas

Show Answer ▶
(c) Change in shape — A change in shape (like cutting, bending) is a physical change, not a chemical reaction indicator.
Q2. The balanced equation for burning of magnesium in air is:
(a) Mg + O2 → MgO   (b) 2Mg + O2 → 2MgO   (c) Mg + 2O → MgO2   (d) 4Mg + 2O2 → 4MgO

Show Answer ▶
(b) 2Mg(s) + O2(g) → 2MgO(s) — Mg=2=2✓, O=2=2✓. Option (d) is not wrong mathematically but uses unnecessarily large coefficients; the simplest ratio is always preferred.
Q3. In the reaction Fe2O3 + 2Al → Al2O3 + 2Fe, what type of reaction is this?
(a) Combination   (b) Double displacement   (c) Decomposition   (d) Displacement

Show Answer ▶
(d) Displacement — Aluminium (more reactive) displaces iron from iron oxide. It is also a redox reaction (Al is oxidised; Fe2O3 is reduced). This is the thermite reaction.
Q4. Which type of energy causes the decomposition of silver chloride into silver and chlorine?
(a) Heat   (b) Electricity   (c) Light   (d) Sound

Show Answer ▶
(c) Light (sunlight) — 2AgCl(s) →Sunlight 2Ag(s) + Cl2(g). This is photolytic decomposition. It is used in black and white photography.
Q5. In the reaction 2PbO(s) + C(s) → 2Pb(s) + CO2(g), which of the following is correct?
(a) Carbon is reduced   (b) Lead oxide is oxidised   (c) Carbon is oxidised   (d) Lead is oxidised

Show Answer ▶
(c) Carbon is oxidised — C gains oxygen to become CO2 (oxidation = gain of oxygen). PbO is reduced (Pb2+ loses oxygen, becomes Pb metal).
Q6. What type of reaction is respiration?
(a) Combination   (b) Endothermic   (c) Exothermic   (d) Photolytic decomposition

Show Answer ▶
(c) Exothermic — Respiration releases energy (heat + ATP): C6H12O6 + 6O2 → 6CO2 + 6H2O + energy. It is also a combination-type reaction (glucose combines with oxygen).
Q7. Which of the following prevents rancidity of food?
(a) Exposure to air   (b) Storing in airtight containers / using antioxidants   (c) Heating regularly   (d) Adding water

Show Answer ▶
(b) Storing in airtight containers and using antioxidants — Rancidity is caused by oxidation of fats/oils. Removing oxygen (airtight) or blocking oxidation (antioxidants like Vit E) prevents it. Chips bags are flushed with nitrogen to remove oxygen.
Q8. When iron nails are dipped in copper sulphate solution, which observation is correct?
(a) No change occurs   (b) The solution turns more blue   (c) Blue colour fades; iron nail turns brownish   (d) Iron nail dissolves completely

Show Answer ▶
(c) Blue colour fades; iron nail turns brownish — Fe displaces Cu from CuSO4. CuSO4 (blue) becomes FeSO4 (pale green). Cu deposits on the iron nail, making it brownish.

  📝 Practice Set B — Short Answer Questions (6 Questions)
Q1. What is a balanced chemical equation? Why should chemical equations be balanced?

Show Answer ▶
A balanced chemical equation has equal numbers of atoms of each element on both the reactant side (LHS) and the product side (RHS). Equations must be balanced to comply with the Law of Conservation of Mass — mass cannot be created or destroyed in a chemical reaction, so the total mass of reactants must equal the total mass of products.
Q2. Explain the difference between exothermic and endothermic reactions. Give one example of each.

Show Answer ▶
Exothermic reactions release heat energy to the surroundings — the reaction mixture becomes warm. Example: CaO + H2O → Ca(OH)2 + Heat (burning, respiration).
Endothermic reactions absorb heat energy from the surroundings — the reaction mixture cools down. Example: CaCO3 →Heat CaO + CO2 (thermal decomposition reactions).
Q3. Why are decomposition reactions called the opposite of combination reactions? Write equations for both.

Show Answer ▶
In a combination reaction, two or more substances combine to give a single product: CaO(s) + H2O(l) → Ca(OH)2(aq). In a decomposition reaction, a single substance breaks down into two or more products: Ca(OH)2(aq) →Heat CaO(s) + H2O(g). They are exact opposites: in one, reactants unite; in the other, a compound splits apart.
Q4. What is corrosion? How does it differ from rancidity? How can both be prevented?

Show Answer ▶
Corrosion is the slow deterioration of metals due to chemical attack by moisture, acids, or oxygen (e.g., rusting of iron, tarnishing of silver). Rancidity is the oxidation of fats/oils in food, causing unpleasant smell and taste. Prevention of corrosion: painting, galvanising, electroplating. Prevention of rancidity: storing in airtight containers, using antioxidants, flushing with nitrogen gas.
Q5. What is a precipitation reaction? Give an example with its equation.

Show Answer ▶
A precipitation reaction is a double displacement reaction in which one of the products formed is an insoluble solid called a precipitate. Example: When sodium sulphate solution is mixed with barium chloride solution, a white precipitate of barium sulphate forms instantly:
Na2SO4(aq) + BaCl2(aq) → BaSO4(s)↓ + 2NaCl(aq)
Q6. Identify the substance oxidised and the substance reduced in: 4Na(s) + O2(g) → 2Na2O(s)

Show Answer ▶
Sodium (Na) is OXIDISED — Na gains oxygen to become Na2O (oxidation = gain of oxygen).
Oxygen (O2) is REDUCED — O2 loses its elemental state by combining with sodium, and in one sense “accepts” electrons/gains a negative character in the compound (in terms of oxygen, it is incorporated into the product). Note: In the simpler Class 10 definition, Na is oxidised (gains O), and O2 acts as the oxidising agent.

  📝 Practice Set C — Long Answer Questions (4 Questions)
Q1. Describe in detail the different types of decomposition reactions with one example, one balanced equation, and one real-world application for each type.

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1. Thermal Decomposition — Energy source: heat. Example: Limestone heated in a kiln: CaCO3(s) →Heat CaO(s) + CO2(g). Application: CaO (quick lime) is used in cement manufacturing and as a building material.

2. Electrolytic Decomposition — Energy source: electricity. Example: Electrolysis of water: 2H2O(l) →Electric current 2H2(g) + O2(g). Application: H2 produced is used as a clean fuel; O2 is used in hospitals and welding.

3. Photolytic (Photochemical) Decomposition — Energy source: sunlight/light. Example: Silver chloride decomposed by light: 2AgCl(s) →Sunlight 2Ag(s) + Cl2(g). Application: This principle is used in black-and-white photography — AgBr on film decomposes where light hits it, recording an image.

Q2. What is oxidation? What is reduction? Explain with two examples each in terms of both oxygen and hydrogen. What are redox reactions?

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Oxidation: A substance is oxidised if it (a) gains oxygen — e.g., 2Mg + O2 → 2MgO (Mg is oxidised); or (b) loses hydrogen — e.g., H2S + Cl2 → S + 2HCl (H2S loses H, so S is oxidised).

Reduction: A substance is reduced if it (a) loses oxygen — e.g., CuO + H2 → Cu + H2O (CuO loses O, so Cu is reduced); or (b) gains hydrogen — e.g., S + H2 → H2S (S gains H, so S is reduced).

Redox Reactions: In a redox reaction, oxidation and reduction always take place simultaneously. One substance is oxidised (acts as reducing agent) while another is reduced (acts as oxidising agent). Example: CuO + H2 → Cu + H2O — CuO is reduced, H2 is oxidised. This is a redox reaction.

Q3. Explain the steps to balance a chemical equation using the hit-and-trial method, with a complete worked example of: Al + Cl2 → AlCl3

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Steps of Hit-and-Trial Method:
Step 1: Write the skeletal (unbalanced) equation: Al + Cl2 → AlCl3.
Step 2: Count atoms — Al: 1 LHS, 1 RHS ✓ | Cl: 2 LHS, 3 RHS ✗
Step 3: Balance chlorine — LCM of 2 and 3 is 6. Put 3Cl2 on LHS and 2AlCl3 on RHS: Al + 3Cl2 → 2AlCl3. Now Cl: 6=6 ✓
Step 4: Balance Al — 2AlCl3 needs 2Al. Put 2Al on LHS: 2Al + 3Cl2 → 2AlCl3. Al: 2=2 ✓
Step 5: Check: Al=2=2 ✓, Cl=6=6 ✓
Balanced: 2Al(s) + 3Cl2(g) → 2AlCl3(s) — This is a combination reaction.
Q4. What are displacement and double displacement reactions? Explain with two examples each and write balanced equations.

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Displacement Reactions: A more reactive element displaces a less reactive element from its compound.
Example 1: Fe(s) + CuSO4(aq) → FeSO4(aq) + Cu(s) — Fe is more reactive than Cu; blue colour fades, brownish Cu deposits on Fe nail.
Example 2: Zn(s) + CuSO4(aq) → ZnSO4(aq) + Cu(s) — Zn displaces Cu; zinc is more reactive.

Double Displacement Reactions: Two compounds react by exchanging ions to form two new compounds. One product is usually a precipitate, gas or water.
Example 1: Na2SO4(aq) + BaCl2(aq) → BaSO4(s)↓ + 2NaCl(aq) — White BaSO4 precipitate forms.
Example 2: Pb(NO3)2(aq) + 2KI(aq) → PbI2(s)↓ + 2KNO3(aq) — Yellow PbI2 precipitate forms.

  📝 Practice Set D — Numerical / Balancing Problems (4 Questions)
Q1. Balance the following equations and identify the type of reaction:
(a) HNO3 + Ca(OH)2 → Ca(NO3)2 + H2O
(b) NaOH + H2SO4 → Na2SO4 + H2O

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(a) 2HNO3(aq) + Ca(OH)2(aq) → Ca(NO3)2(aq) + 2H2O(l)
Check: H=2+2=4=4✓, N=2=2✓, O=6+2=8=8✓, Ca=1=1✓ — Double displacement (neutralisation) reaction.

(b) 2NaOH(aq) + H2SO4(aq) → Na2SO4(aq) + 2H2O(l)
Check: Na=2=2✓, O=2+4=6=6+2=6✓, H=2+2=4=4✓, S=1=1✓ — Double displacement (acid-base neutralisation) reaction.

Q2. Write balanced equations for the following and identify the type:
(a) Zinc carbonate → Zinc oxide + Carbon dioxide
(b) Potassium bromide(aq) + Barium iodide(aq) → Potassium iodide(aq) + Barium bromide(s)

Show Answer ▶
(a) ZnCO3(s) →Heat ZnO(s) + CO2(g) — Already balanced. Zn=1✓, C=1✓, O=3✓. Thermal decomposition (endothermic).

(b) 2KBr(aq) + BaI2(aq) → 2KI(aq) + BaBr2(s)
Check: K=2=2✓, Br=2=2✓, Ba=1=1✓, I=2=2✓ — Double displacement / precipitation reaction. BaBr2 precipitates.

Q3. In the refining of silver, copper is used to displace silver from silver nitrate solution. Write the balanced reaction. How many grams of copper are consumed to produce 108g of silver? (Atomic mass: Cu = 63.5, Ag = 108)

Show Answer ▶
Reaction: Cu(s) + 2AgNO3(aq) → Cu(NO3)2(aq) + 2Ag(s)
Cu=1=1✓, Ag=2=2✓, N=2=2✓, O=6=6✓ — Displacement reaction (Cu is more reactive than Ag).

Calculation:
From the equation: 1 mole Cu → 2 moles Ag
63.5g of Cu → 2 × 108 = 216g of Ag
For 108g of Ag: Cu needed = (63.5 × 108) / 216 = 6858/216 = 31.75g of copper

Q4. Balance: MnO2 + HCl → MnCl2 + H2O + Cl2. Identify what is oxidised and reduced. (This is from the NCERT text.)

Show Answer ▶
Balanced: MnO2(s) + 4HCl(aq) → MnCl2(aq) + 2H2O(l) + Cl2(g)
Check: Mn=1=1✓, O=2=2✓, H=4=4✓, Cl=4=4✓

Redox analysis:
HCl → Cl2: HCl loses H and Cl changes to Cl2 (oxidation state increases) → HCl is OXIDISED
MnO2 → MnCl2: MnO2 loses oxygen → MnO2 is REDUCED

  📚 Chapter Summary

Chapter 1 — Chemical Reactions and Equations: Key Points

Indicators of a Chemical Reaction
Change in state • Change in colour • Evolution of gas • Change in temperature		 Combination Reaction
A + B → AB • Two or more → single product • Often exothermic • Example: CaO + H2O → Ca(OH)2
Decomposition Reaction
AB → A + B • Three types: thermal, electrolytic, photolytic • Endothermic • Example: CaCO3 → CaO + CO2		 Displacement Reaction
A + BC → AC + B • More reactive displaces less reactive • Example: Fe + CuSO4 → FeSO4 + Cu
Double Displacement / Precipitation
AB + CD → AD + CB • Ions exchange • Forms precipitate • Example: Na2SO4 + BaCl2 → BaSO4↓ + 2NaCl		 Oxidation and Reduction (Redox)
Oxidation = gain O / lose H • Reduction = lose O / gain H • Always simultaneous in redox
Important Formulae / Equations Notes
2Mg + O2 → 2MgO Combination + Exothermic
3Fe + 4H2O → Fe3O4 + 4H2 Combination of iron with steam
2H2O → 2H2 + O2 Electrolytic decomposition
2AgCl →Sunlight 2Ag + Cl2 Photolytic decomposition
C6H12O6 + 6O2 → 6CO2 + 6H2O + energy Respiration — exothermic
CuO + H2 → Cu + H2O Redox: CuO reduced, H2 oxidised

  🏆 8-Point Exam Quick-Check
✅ Must-Know Facts

1. Four indicators of a chemical reaction: change in state, colour, temperature, gas evolution.

2. Law of Conservation of Mass: mass cannot be created or destroyed — equations must always be balanced.

3. In electrolysis of water: H2 volume is always double O2 volume (2:1 ratio).

4. AgCl and AgBr both decompose by light — used in black & white photography.

5. Respiration, burning, and slaked lime formation are all exothermic reactions.
⚠ Common Exam Traps

Trap 1: Never change subscripts to balance equations — only change coefficients!

Trap 2: Fe2O3 + 2Al → Al2O3 + 2Fe is a displacement reaction, NOT combination.

Trap 3: CaO + H2O → Ca(OH)2 is BOTH combination AND exothermic — give both types.

Trap 4: In redox: oxidation = gain O (not lose O). Students often confuse these.

Trap 5: Rusting forms Fe2O3·xH2O (hydrated), not just Fe2O3.

  💡 Important Facts to Remember
⚡ Nitrogen Gas in Chips Packets: Chips manufacturers flush bags with nitrogen (N2) to prevent oxidation of the fat/oil in chips. Nitrogen is inert and does not react with food. This prevents rancidity and extends shelf life.
⚡ Thermite Reaction: Fe2O3 + 2Al → Al2O3 + 2Fe is an extremely exothermic displacement reaction (thermite reaction). Temperatures reach 2000°C+. Used in welding railway tracks.
⚡ Whitewashing Walls: Ca(OH)2 + CO2 → CaCO3 + H2O. The shiny finish on walls appears after 2–3 days because CaCO3 (which is also the formula for marble!) forms a thin hard layer.
⚡ Why is endothermic not same as cold? Endothermic reactions absorb heat from the surroundings, so the surroundings (the beaker, the water around it) actually get colder — but the reaction itself consumes energy, it doesn’t produce cold. Example: Dissolving ammonium nitrate in water absorbs heat (feels cold to touch).

This comprehensive study page for Class 10 Science Chapter 1 — Chemical Reactions and Equations covers all CBSE and NCERT topics including word equations, balanced chemical equations, types of chemical reactions (combination, decomposition, displacement, double displacement, and redox reactions), oxidation and reduction, corrosion, and rancidity. Includes 13 fully solved worked examples, 4 practice sets with answers, interactive activity accordions, and HTML table diagrams of all key experiments. Perfect for CBSE Class 10 board exam preparation. Visit School Revise for more Grade 10 NCERT Science study materials.

© 2025 School Revise — All Rights Reserved — CBSE / NCERT Science Class 10 — Chapter 1: Chemical Reactions and Equations. Curriculum reference: NCERT Science Textbook, Class 10, Chapter 1. All content original and independently written by School Revise.

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