In this chapter, you will learn
- —Identify signs and evidence of a chemical reaction (colour change, gas evolution, precipitate, temperature change)
- —Write and balance chemical equations using the hit-and-trial method with state symbols
- —Explain combination, decomposition, displacement, and double displacement reactions with examples
- —Distinguish between oxidation and reduction, and identify redox reactions
- —Differentiate between exothermic and endothermic reactions with examples
- —Understand corrosion and rancidity as effects of oxidation in daily life
- —Write balanced chemical equations for all NCERT activities in this chapter
Chemical Reactions — Signs and Evidence
A chemical reaction is a process in which one or more substances (reactants) are transformed into new substances (products) with different properties.
How do we know a chemical reaction has taken place? We look for observable changes:
| Sign / Evidence | Example |
|---|---|
| Change in colour | Copper sulphate (blue) + iron → iron sulphate (green) + copper (brown) |
| Evolution of gas | Zinc + dilute H₂SO₄ → zinc sulphate + hydrogen gas (bubbles) |
| Formation of precipitate | Na₂SO₄ + BaCl₂ → BaSO₄ (white precipitate) + NaCl |
| Change in temperature | Quicklime + water → slaked lime (heat released, beaker becomes hot) |
| Change in state | Wax (solid) burns to form CO₂ (gas) and H₂O (vapour) |
Key Points:
- Reactants: Substances that undergo change in a chemical reaction (written on the left side)
- Products: New substances formed after the reaction (written on the right side)
- A chemical reaction involves breaking of bonds in reactants and making of new bonds in products
- Not all changes are chemical — melting ice (physical change) does not form a new substance
Exam Tip
Board exams often ask: 'List any two observations that indicate a chemical reaction has taken place.' Give two distinct signs with examples from NCERT activities.
Common Mistake
Students confuse physical changes (like melting, dissolving) with chemical changes. A chemical reaction always produces a NEW substance with different chemical properties.
Chemical Equations — Writing and Balancing
A chemical equation is a shorthand representation of a chemical reaction using symbols and formulae of the substances involved.
Steps to write a chemical equation:
- Word equation: Magnesium + Oxygen → Magnesium oxide
- Skeletal equation: Mg + O₂ → MgO (unbalanced)
- Balanced equation: 2Mg + O₂ → 2MgO
Law of Conservation of Mass:
The total mass of reactants = total mass of products.
Therefore, atoms of each element must be equal on both sides.
Hit-and-Trial Method of Balancing:
- Write the skeletal equation with correct formulae
- Count atoms of each element on both sides
- Start balancing with the element that appears in the fewest formulae, or the element with the highest number of atoms
- Use coefficients (numbers before formulae) — never change the subscripts in a formula
- Check that all elements are balanced
Example — Balancing Fe + H₂O → Fe₃O₄ + H₂:
- Fe: 1 on left, 3 on right → put 3 before Fe: 3Fe + H₂O → Fe₃O₄ + H₂
- O: 1 on left, 4 on right → put 4 before H₂O: 3Fe + 4H₂O → Fe₃O₄ + H₂
- H: 8 on left, 2 on right → put 4 before H₂: 3Fe + 4H₂O → Fe₃O₄ + 4H₂
- Check: Fe(3=3), H(8=8), O(4=4) ✓ Balanced!
State Symbols:
- (s) — solid
- (l) — liquid
- (g) — gas
- (aq) — aqueous (dissolved in water)
Complete balanced equation with state symbols:
3Fe(s) + 4H₂O(g) → Fe₃O₄(s) + 4H₂(g)
Exam Tip
Balancing equations is tested every year! Always use the hit-and-trial method. Never change subscripts — only add coefficients. Include state symbols (s), (l), (g), (aq) for full marks.
Common Mistake
Students change the formula of a compound to balance the equation (e.g., writing H₃ instead of adding a coefficient before H₂). You must NEVER change subscripts — only adjust the coefficients.
Combination Reactions
A combination reaction is a reaction in which two or more substances combine to form a single product.
General Form: A + B → AB
Important Examples:
1. Burning of magnesium ribbon:
2Mg(s) + O₂(g) → 2MgO(s)
Magnesium burns with a dazzling white flame to form white ash (MgO)
2. Quicklime reacting with water (slaking of lime):
CaO(s) + H₂O(l) → Ca(OH)₂(aq)
Large amount of heat released — exothermic combination reaction
3. Burning of coal:
C(s) + O₂(g) → CO₂(g)
4. Formation of water:
2H₂(g) + O₂(g) → 2H₂O(l)
Key Points:
- Two or more reactants combine to form a single product
- Most combination reactions are exothermic (release heat)
- The slaking of lime (CaO + H₂O) is a classic NCERT example — the beaker becomes very hot
- Combination reactions are the opposite of decomposition reactions
Exam Tip
CaO + H₂O → Ca(OH)₂ is the most frequently asked combination reaction. Remember: it is exothermic (releases heat). Magnesium burning is another favourite.
Common Mistake
Students write Ca(OH) instead of Ca(OH)₂. Calcium has a valency of 2, and hydroxide (OH) has a valency of 1, so calcium hydroxide is Ca(OH)₂.
Decomposition Reactions
A decomposition reaction is one in which a single compound breaks down into two or more simpler substances. Energy (heat, light, or electricity) is required to break the bonds.
General Form: AB → A + B
Types of Decomposition Reactions:
1. Thermal Decomposition (by heat):
a) Ferrous sulphate:
2FeSO₄(s) →Heat Fe₂O₃(s) + SO₂(g) + SO₃(g)
Green crystals → reddish-brown residue + gases with smell of burning sulphur
b) Calcium carbonate (limestone):
CaCO₃(s) →Heat CaO(s) + CO₂(g)
This is how quicklime (CaO) is manufactured
c) Lead nitrate:
2Pb(NO₃)₂(s) →Heat 2PbO(s) + 4NO₂(g) + O₂(g)
Emits brown fumes of nitrogen dioxide (NO₂)
2. Electrolytic Decomposition (by electricity):
Electrolysis of water:
2H₂O(l) →Electricity 2H₂(g) + O₂(g)
Hydrogen collects at cathode, oxygen at anode. Volume ratio H₂:O₂ = 2:1
3. Photolytic Decomposition (by light):
a) Silver chloride:
2AgCl(s) →Sunlight 2Ag(s) + Cl₂(g)
White silver chloride turns grey in sunlight
b) Silver bromide:
2AgBr(s) →Sunlight 2Ag(s) + Br₂(g)
Used in black and white photography
Key Points:
- Decomposition is the opposite of combination
- Most decomposition reactions are endothermic (absorb energy)
- Decomposition of AgCl and AgBr is used in black and white photography
- Electrolysis of water proves that water is made of hydrogen and oxygen in the ratio 2:1
Exam Tip
The three types of decomposition (thermal, electrolytic, photolytic) are asked very frequently. Memorize FeSO₄ (green→brown), Pb(NO₃)₂ (brown fumes), AgCl (grey in sunlight), and electrolysis of water (2:1 ratio).
Common Mistake
Students forget to specify the type of decomposition. If heat is used → thermal, electricity → electrolytic, sunlight → photolytic. Always mention the energy source.
Displacement Reactions
A displacement reaction is one in which a more reactive element displaces a less reactive element from its compound (salt solution).
General Form: A + BC → AC + B
(A is more reactive than B)
Important Examples:
1. Iron nail in copper sulphate solution:
Fe(s) + CuSO₄(aq) → FeSO₄(aq) + Cu(s)
Blue colour of CuSO₄ fades → green FeSO₄ forms. Brown copper deposits on iron nail.
2. Zinc in copper sulphate solution:
Zn(s) + CuSO₄(aq) → ZnSO₄(aq) + Cu(s)
Blue colour disappears, copper deposits on zinc.
3. Lead in copper chloride solution:
Pb(s) + CuCl₂(aq) → PbCl₂(aq) + Cu(s)
Reactivity Series (Metals):
A more reactive metal can displace a less reactive metal from its salt solution:
K > Na > Ca > Mg > Al > Zn > Fe > Ni > Sn > Pb > H > Cu > Hg > Ag > Au
(Most reactive → Least reactive)
Key Points:
- A more reactive element displaces a less reactive element
- Iron displaces copper because iron is more reactive than copper in the reactivity series
- Copper cannot displace iron from FeSO₄ because copper is less reactive
- The iron nail in CuSO₄ activity (NCERT Activity 1.10) is a very important exam question
Exam Tip
The iron + CuSO₄ reaction is asked almost every year. Remember the observations: blue colour fades, green colour appears, brown deposit on iron nail. Know why — iron is MORE reactive than copper.
Common Mistake
Students try to show copper displacing iron (Cu + FeSO₄). This is WRONG because copper is less reactive than iron. Only a more reactive metal can displace a less reactive one.
Double Displacement Reactions
A double displacement reaction is one in which two compounds in solution exchange their ions to form two new compounds. One of the products is usually an insoluble precipitate, water, or a gas.
General Form: AB + CD → AD + CB
Important Examples:
1. Precipitation reaction:
Na₂SO₄(aq) + BaCl₂(aq) → BaSO₄(s)↓ + 2NaCl(aq)
White precipitate of BaSO₄ (barium sulphate) is formed
2. Another precipitation reaction:
NaOH(aq) + HCl(aq) → NaCl(aq) + H₂O(l)
This is also a neutralization reaction (acid + base → salt + water)
Types of Double Displacement:
- Precipitation reaction: When an insoluble compound (precipitate) is formed. Example: Na₂SO₄ + BaCl₂ → BaSO₄↓ + 2NaCl
- Neutralization reaction: When an acid reacts with a base to form salt and water. Example: NaOH + HCl → NaCl + H₂O
Key Points:
- Both reactants are usually in aqueous solution
- Ions are exchanged between the two compounds
- The ↓ symbol indicates a precipitate (insoluble product that settles down)
- Precipitation reactions are used as a test to detect certain ions in solution
Exam Tip
Na₂SO₄ + BaCl₂ is the standard NCERT example of a double displacement (precipitation) reaction. Remember: both ions are exchanged. The insoluble product (BaSO₄) forms as a white precipitate.
Common Mistake
Students confuse displacement with double displacement. In displacement, ONE element replaces another (A + BC → AC + B). In double displacement, TWO compounds exchange ions (AB + CD → AD + CB).
Oxidation and Reduction — Redox Reactions
Oxidation and reduction always occur simultaneously. A reaction in which both oxidation and reduction take place is called a redox reaction.
Oxidation: Gain of oxygen OR loss of hydrogen by a substance
Reduction: Loss of oxygen OR gain of hydrogen by a substance
Example — Reduction of copper oxide:
CuO(s) + H₂(g) → Cu(s) + H₂O(g)
- CuO is reduced — it loses oxygen (CuO → Cu)
- H₂ is oxidised — it gains oxygen (H₂ → H₂O)
- CuO is the oxidising agent (provides oxygen, gets reduced itself)
- H₂ is the reducing agent (removes oxygen, gets oxidised itself)
Another Example — Reduction of zinc oxide:
ZnO(s) + C(s) → Zn(s) + CO(g)
- ZnO is reduced (loses oxygen) — oxidising agent
- C is oxidised (gains oxygen) — reducing agent
Example involving hydrogen:
MnO₂ + 4HCl → MnCl₂ + 2H₂O + Cl₂
- HCl is oxidised (loses hydrogen)
- MnO₂ is reduced (gains hydrogen to form H₂O)
Key Definitions:
| Term | Definition |
|---|---|
| Oxidation | Gain of oxygen or loss of hydrogen |
| Reduction | Loss of oxygen or gain of hydrogen |
| Oxidising agent | Substance that gives oxygen / removes hydrogen (gets reduced itself) |
| Reducing agent | Substance that gives hydrogen / removes oxygen (gets oxidised itself) |
Exam Tip
CuO + H₂ → Cu + H₂O is the most important redox reaction. Identify which substance is oxidised, reduced, and name the oxidising and reducing agents — all four are commonly asked.
Common Mistake
Students confuse oxidising and reducing agents. The oxidising agent itself gets REDUCED (it gives oxygen to the other substance). The reducing agent itself gets OXIDISED.
Exothermic and Endothermic Reactions
Chemical reactions can be classified based on whether they release or absorb energy (heat).
Exothermic Reactions — reactions that release heat (products have less energy):
Examples:
1. Burning of natural gas: CH₄(g) + 2O₂(g) → CO₂(g) + 2H₂O(g) + heat
2. Respiration: C₆H₁₂O₆ + 6O₂ → 6CO₂ + 6H₂O + energy
3. Slaking of lime: CaO + H₂O → Ca(OH)₂ + heat
4. Burning of magnesium: 2Mg + O₂ → 2MgO + heat and light
5. Dissolving acid in water releases heat
Endothermic Reactions — reactions that absorb heat (products have more energy):
Examples:
1. Decomposition of CaCO₃: CaCO₃ → CaO + CO₂ (needs continuous heating)
2. Decomposition of FeSO₄: 2FeSO₄ → Fe₂O₃ + SO₂ + SO₃ (needs heat)
3. Electrolysis of water: 2H₂O → 2H₂ + O₂ (needs electrical energy)
4. Photosynthesis: 6CO₂ + 6H₂O → C₆H₁₂O₆ + 6O₂ (needs sunlight)
5. Dissolving NH₄Cl in water makes the beaker cold
Key Points:
- Exothermic: Temperature of surroundings increases; beaker feels hot
- Endothermic: Temperature of surroundings decreases; beaker feels cold
- Respiration is an exothermic process (releases energy from food)
- Photosynthesis is an endothermic process (absorbs light energy)
- Most combination reactions are exothermic
- Most decomposition reactions are endothermic
Exam Tip
Remember: Combination = exothermic (usually), Decomposition = endothermic (usually). Respiration = exothermic. Photosynthesis = endothermic. These are frequently asked in 1-mark questions.
Common Mistake
Students sometimes think all reactions that need heat to start are endothermic. Burning magnesium needs a flame to start, but it IS exothermic because it releases much more heat than it absorbs.
Effects of Oxidation — Corrosion and Rancidity
Oxidation reactions affect our daily life in two important ways: corrosion of metals and rancidity of food.
Corrosion:
When a metal is attacked by substances like moisture, acids, or gases in its environment, it is said to corrode. The process is called corrosion.
| Metal | Corrosion Product | Appearance |
|---|---|---|
| Iron | Rust (Fe₂O₃·xH₂O) | Reddish-brown |
| Silver | Silver sulphide (Ag₂S) | Black tarnish |
| Copper | Copper carbonate (CuCO₃) | Green coating (patina) |
Conditions for rusting of iron:
- Presence of moisture (water)
- Presence of oxygen (air)
- Both are required — iron does not rust in dry air or in boiled (oxygen-free) water
Prevention of Corrosion:
- Painting / oiling / greasing — prevents contact with air and moisture
- Galvanisation — coating iron with a layer of zinc
- Electroplating — coating with chromium, nickel, or tin
- Alloying — making stainless steel (iron + chromium + nickel)
Rancidity:
When fats and oils in food are oxidised, they become rancid — the food develops a bad taste and smell. This is called rancidity.
Prevention of Rancidity:
- Adding antioxidants (like BHA, BHT, vitamin E)
- Storing food in airtight containers — reduces contact with oxygen
- Flushing with nitrogen gas — chip packets are filled with nitrogen to prevent oxidation
- Refrigeration — low temperature slows down oxidation
- Storing in dark — light accelerates oxidation
Key Points:
- Both corrosion and rancidity are caused by oxidation
- Galvanisation (zinc coating) is the most common method to prevent rusting of iron
- Chip packets are filled with nitrogen gas, not air — to prevent rancidity
Exam Tip
Iron rusting needs BOTH moisture AND oxygen. Prevention methods (galvanisation, painting, oiling) are frequently asked. Rancidity prevention — remember nitrogen flushing in chip packets!
Common Mistake
Students write that rusting needs only water OR only air. Both moisture AND oxygen are required for rusting. Iron does not rust in dry air or in oxygen-free water.