CBSE Class 10 · Science · board exam 2027
Chapter 1: Chemical Reactions and Equations — notes
The chapter’s 11 parts in the book’s order — the first idea of each and the points Medha Ma’am writes on the board — and the summary at the end. The full lesson and a check question for each part are in the app.
1. How do we know a chemical reaction has taken place?
First, think of a few everyday things.
- • Milk turning sour, rusting of iron, fermentation of grapes, digestion of food
- • In all of these, the nature and identity of the substance changes
- • A chemical change means a chemical reaction has taken place
- • How do we know? Three activities
- • Activity 1.1: burning a magnesium ribbon
- • Materials: Mg ribbon, sandpaper, tongs, burner, watch-glass
- • Method: clean the ribbon → hold with tongs and burn → collect the ash
- • Observation: dazzling white flame; a white powder (magnesium oxide) forms
- • Conclusion: a new substance formed, so a reaction took place
- • Precaution: clean the ribbon with sandpaper (to remove the layer of MgO)
- • Precaution: keep the burning ribbon away from the eyes
- • Activity 1.2: lead nitrate + potassium iodide
- • Observation: a yellow precipitate (lead iodide) forms
- • Activity 1.3: zinc + dilute sulphuric acid
- • Observation: bubbles of hydrogen gas; the flask becomes warm
- • Burning splinter near the gas: a pop sound, so the gas is hydrogen
- • Signs of a chemical reaction
- • ① change in state ② change in colour
- • ③ evolution of a gas ④ change in temperature
2. Writing a chemical equation
Writing every reaction as a full sentence becomes very long.
- • Word-equation
- • Magnesium + Oxygen → Magnesium oxide
- • Reactants: take part in the reaction (left side)
- • Products: the new substances formed (right side)
- • Arrow: the direction of the reaction, from reactants to products
- • Chemical equation: a reaction written with symbols and formulae
- • Mg + O₂ → MgO (unbalanced)
- • | Element | LHS | RHS |
- • | Mg | 1 | 1 ✓ |
- • | O | 2 | 1 ✗ |
- • Atoms not equal: a skeletal (unbalanced) equation
3. Balancing a chemical equation
Recall a law from Class IX, the law of conservation of mass.
- • Law of conservation of mass: mass can neither be created nor destroyed in a chemical reaction
- • So: the atoms of each element must be equal on both sides
- • Balanced equation: the atoms of each element are equal on both sides
- • Zn + H₂SO₄ → ZnSO₄ + H₂
- • Zn: 1 = 1, H: 2 = 2, S: 1 = 1, O: 4 = 4 ✓
- • Example: Fe + H₂O → Fe₃O₄ + H₂ (unbalanced)
- • | Element | LHS | RHS |
- • | Fe | 1 | 3 |
- • | H | 2 | 2 |
- • | O | 1 | 4 |
- • Step 2: oxygen — put 4 in front of H₂O
- • Fe + 4H₂O → Fe₃O₄ + H₂ (unbalanced)
- • Rule: never change a formula, only put coefficients (4H₂O, not H₂O₄)
- • Step 3: hydrogen — LHS 4 × 2 = 8, so put 4 in front of H₂
- • Fe + 4H₂O → Fe₃O₄ + 4H₂ (unbalanced)
- • Step 4: iron — put 3 in front of Fe
- • Check: Fe: 3 = 3, H: 8 = 8, O: 4 = 4 ✓
- • 3Fe + 4H₂O → Fe₃O₄ + 4H₂
- • This is the hit-and-trial method
- • Mg + O₂ → MgO (unbalanced)
- • Oxygen: 2 in front of MgO ⇒ magnesium: 2 in front of Mg
- • 2Mg + O₂ → 2MgO (Mg: 2 = 2, O: 2 = 2 ✓)
4. Physical states and the conditions of a reaction
To make a balanced equation tell us even more, we write the physical state of each substance with it.
- • Symbols of physical states
- • (s) solid (l) liquid (g) gas (aq) aqueous solution
- • 3Fe(s) + 4H₂O(g) → Fe₃O₄(s) + 4H₂(g)
- • H₂O(g): water used as steam
- • Over the arrow: the condition of the reaction (temperature, pressure, light, catalyst)
- • CO(g) + 2H₂(g) →[340 atm] CH₃OH(l)
- • 6CO₂(aq) + 12H₂O(l) →[sunlight, chlorophyll] C₆H₁₂O₆(aq) + 6O₂(aq) + 6H₂O(l)
5. Combination reactions and exothermic reactions
The types of reactions.
- • Activity 1.4: quick lime + water
- • CaO(s) + H₂O(l) → Ca(OH)₂(aq) + heat
- • quick lime slaked lime
- • Combination reaction: two or more reactants → a single product
- • Ca: 1 = 1, O: 2 = 2, H: 2 = 2 ✓
- • Whitewash: Ca(OH)₂(aq) + CO₂(g) → CaCO₃(s) + H₂O(l)
- • The layer of CaCO₃ makes the walls shine (marble is also CaCO₃)
- • C(s) + O₂(g) → CO₂(g)
- • H₂ + O₂ → H₂O (unbalanced)
- • 2H₂(g) + O₂(g) → 2H₂O(l) (H: 4 = 4, O: 2 = 2 ✓)
- • Exothermic reaction: heat is released along with the products
- • CH₄(g) + 2O₂(g) → CO₂(g) + 2H₂O(g) + energy
- • Respiration: C₆H₁₂O₆(aq) + 6O₂(aq) → 6CO₂(aq) + 6H₂O(l) + energy
- • Vegetable matter into compost: also exothermic
6. Decomposition reactions: by heat
The exact opposite.
- • Activity 1.5: heating ferrous sulphate crystals
- • Observation: the green colour changes; smell of burning sulphur
- • FeSO₄ → Fe₂O₃ + SO₂ + SO₃ (unbalanced)
- • Fe: 2 on the RHS, so put 2 in front of FeSO₄
- • O: LHS 2 × 4 = 8, RHS 3 + 2 + 3 = 8 ✓
- • 2FeSO₄(s) →[heat] Fe₂O₃(s) + SO₂(g) + SO₃(g)
- • Decomposition reaction: a single reactant → two or more simpler products
- • Broken by heat: thermal decomposition
- • CaCO₃(s) →[heat] CaO(s) + CO₂(g)
- • limestone quick lime (used to make cement)
- • Activity 1.6: heating lead nitrate
- • Observation: brown fumes (NO₂) are given off
- • Pb(NO₃)₂ → PbO + NO₂ + O₂ (unbalanced)
- • LHS 2Pb(NO₃)₂: Pb 2, N 2 × 2 = 4, O 2 × 6 = 12
- • RHS 2PbO + 4NO₂ + O₂: O 2 + 8 + 2 = 12 ✓
- • 2Pb(NO₃)₂(s) →[heat] 2PbO(s) + 4NO₂(g) + O₂(g)
7. Decomposition by electricity and light; endothermic reactions
Decomposition does not always need heat.
- • Activity 1.7: electrolysis of water
- • A few drops of dilute H₂SO₄: to make water conduct electricity
- • Cathode (−): hydrogen Anode (+): oxygen
- • 2H₂O(l) →[electric current] 2H₂(g) + O₂(g)
- • Volume: hydrogen : oxygen = 2 : 1
- • Activity 1.8: silver chloride in sunlight
- • 2AgCl(s) →[sunlight] 2Ag(s) + Cl₂(g)
- • 2AgBr(s) →[sunlight] 2Ag(s) + Br₂(g)
- • Use: black and white photography
- • Endothermic reaction: energy is absorbed
- • Energy for decomposition: heat, light or electricity
- • Most decompositions are endothermic (but compost formation is exothermic)
- • | Combination | Decomposition |
- • | Two or more reactants form a single product | A single reactant gives two or more products |
- • | Usually releases energy (exothermic) | Usually absorbs energy (endothermic) |
- • | CaO + H₂O → Ca(OH)₂ | CaCO₃ →[heat] CaO + CO₂ |
8. Displacement reactions
The figure, Activity 1.9. Two test tubes, A and B, have blue copper sulphate solution.
- • Activity 1.9: iron nails in copper sulphate solution
- • Observation: a brown coating (copper) on the nails; blue solution turns pale green
- • Fe(s) + CuSO₄(aq) → FeSO₄(aq) + Cu(s)
- • Solution: blue to pale green; brown copper on the nail
- • Displacement reaction: a more reactive element displaces a less reactive one from its compound
- • Zn(s) + CuSO₄(aq) → ZnSO₄(aq) + Cu(s)
- • Pb(s) + CuCl₂(aq) → PbCl₂(aq) + Cu(s)
9. Double displacement and precipitation reactions
The figure, Activity 1.10.
- • Activity 1.10: sodium sulphate + barium chloride
- • Observation: a white precipitate (barium sulphate) forms
- • Na₂SO₄(aq) + BaCl₂(aq) → BaSO₄(s) + 2NaCl(aq)
- • Ba²⁺ + SO₄²⁻ → BaSO₄ (white precipitate)
- • Na⁺ and Cl⁻ stay in the solution (NaCl)
- • Na: 2 = 2, Cl: 2 = 2, Ba: 1 = 1, S: 1 = 1, O: 4 = 4 ✓
- • Double displacement: exchange of ions between the reactants
- • Precipitation: a reaction that produces a precipitate (an insoluble substance)
- • Pb(NO₃)₂(aq) + 2KI(aq) → PbI₂(s) + 2KNO₃(aq) (yellow precipitate)
- • | Displacement | Double displacement |
- • | An element + a compound | Two compounds |
- • | A more reactive element removes another | The two exchange their ions |
- • | Fe + CuSO₄ → FeSO₄ + Cu | Na₂SO₄ + BaCl₂ → BaSO₄ + 2NaCl |
10. Oxidation and reduction
The figure, Activity 1.11.
- • Activity 1.11: heating copper powder
- • Observation: a black coating (copper oxide) on the brown powder
- • 2Cu + O₂ →[heat] 2CuO
- • CuO + H₂ →[heat] Cu + H₂O
- • CuO: loses oxygen H₂: gains oxygen
- • Oxidation: gain of oxygen or loss of hydrogen
- • Reduction: loss of oxygen or gain of hydrogen
- • CuO: reduced (oxidising agent) H₂: oxidised (reducing agent)
- • Redox: oxidation and reduction in the same reaction
- • ZnO + C → Zn + CO (ZnO reduced, C oxidised)
- • MnO₂ + 4HCl → MnCl₂ + 2H₂O + Cl₂
- • MnO₂: loss of oxygen (reduced) HCl: loss of hydrogen (oxidised)
11. Oxidation in everyday life: corrosion and rancidity
Oxidation shows up in our daily life too.
- • Corrosion: a metal attacked by moisture, acids, etc. around it
- • Iron: reddish brown coating (rust)
- • Silver: black coating Copper: green coating
- • Damage: car bodies, bridges, railings, ships and other iron objects
- • Rancidity: fats and oils oxidised, so their smell and taste change
- • Prevention: antioxidants, airtight containers
- • Chips packets flushed with nitrogen: to prevent oxidation
Summary
Very good! First, we saw four signs that tell us a chemical reaction has taken place, a change in state, colour, gas and temperature. Then we wrote a chemical equation from a word-equation, and because of the law of conservation of mass, we balanced it by the hit-and-trial method, with the state symbols too. Then the types of reactions. In combination, several join into one. In decomposition, one breaks into several, by heat, light or electricity. In displacement, a more reactive element removes another. In double displacement, the ions are exchanged, and often a precipitate forms. Exothermic releases heat, endothermic absorbs energy. And finally oxidation, reduction, redox, and their effects, corrosion and rancidity. Write all of today's balanced equations once in your notebook without looking, and check each one by counting the atoms on both sides. Your homework is on your screen. In the next class we will study acids, bases and salts, where these equations will be used again. If anything was not clear, press the Ask button and ask me. Done, you worked very well today!
AskMedha · Medha Ma’am — CBSE Class 10 Science, Chapter 1. Written from the textbook, checked twice.
Questions parents and students ask
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