Chemical Reactions and Equations – Class 10 Science Chapter 1 Notes & NCERT PDF

Class 10ScienceChapter 1NCERT book: Science - Class 10

Notes and a simple summary of Chapter 1 of the NCERT Class 10 Science book, with the official chapter PDF, flashcards, an MCQ quiz and an AI tutor for your doubts.

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Chapter 1: Chemical Reactions and Equations

Introduction

Milk turning sour in summer, an iron tawa rusting in humid air, grapes fermenting, food being cooked or digested, and the breathing we do every second — in each case the nature and identity of the starting substance changes. Whenever such a chemical change happens, we say a chemical reaction has taken place. This chapter teaches you how to recognise a reaction, how to write it as a balanced chemical equation, and how to classify reactions into types. It ends with two everyday effects of oxidation: corrosion and rancidity.

Key Concepts

1. How do we know a reaction has happened?

From Activities 1.1–1.3 (burning magnesium ribbon, mixing lead nitrate with potassium iodide, adding dilute acid to zinc granules), any of these observations signals a chemical reaction:

change in state

change in colour

evolution of a gas

change in temperature

2. Chemical equations

A word-equation writes the reaction in words: Magnesium + Oxygen → Magnesium oxide. Substances on the left are reactants; on the right are products; the arrow shows the direction of the reaction.

A skeletal equation uses formulae but may not be balanced, e.g. Mg + O₂ → MgO.

By the law of conservation of mass, mass can neither be created nor destroyed, so the number of atoms of each element must be the same on both sides. Such an equation is balanced.

3. Balancing by the hit-and-trial method (Fe + H₂O → Fe₃O₄ + H₂)

Draw boxes around each formula — never change anything inside a formula (write 4H₂O, never H₂O₄).

List the atoms of each element on both sides.

Start with the compound having the most atoms (Fe₃O₄) and its element with the most atoms (oxygen): put 4 before H₂O.

Balance hydrogen: 4H₂.

Balance iron: 3Fe.

Check every element: 3Fe + 4H₂O → Fe₃O₄ + 4H₂.

Add state symbols — (s) solid, (l) liquid, (g) gas, (aq) aqueous solution: 3Fe(s) + 4H₂O(g) → Fe₃O₄(s) + 4H₂(g). Here (g) shows that water is used as steam.

Reaction conditions such as temperature, pressure or a catalyst are written above or below the arrow (e.g. "Sunlight" and "Chlorophyll" for photosynthesis).

4. Types of chemical reactions

Combination reaction – two or more reactants form a single product. Quick lime reacts vigorously with water: CaO(s) + H₂O(l) → Ca(OH)₂(aq) + Heat. The slaked-lime solution is used for whitewashing; it slowly absorbs CO₂ to form a shiny layer of CaCO₃ (the same formula as marble) in two to three days. Burning of coal (C + O₂ → CO₂) and formation of water (2H₂ + O₂ → 2H₂O) are other examples.

Exothermic reactions release heat: burning of natural gas (CH₄ + 2O₂ → CO₂ + 2H₂O), respiration (glucose + oxygen → CO₂ + water + energy) and decomposition of vegetable matter into compost.

Decomposition reaction – a single reactant breaks into simpler products; the opposite of combination.

- Thermal decomposition (by heat): green ferrous sulphate crystals lose colour and give Fe₂O₃, SO₂ and SO₃ with a smell of burning sulphur; CaCO₃ → CaO + CO₂ (quick lime for cement); heating lead nitrate gives brown fumes of nitrogen dioxide (NO₂).

- Electrolytic decomposition (by electricity): electrolysis of water gives hydrogen and oxygen; the volume of hydrogen is double that of oxygen.

- Photolytic decomposition (by light): white silver chloride turns grey in sunlight (2AgCl → 2Ag + Cl₂); silver bromide behaves similarly — used in black-and-white photography.

- Decomposition reactions absorb energy, so they are endothermic.

Displacement reaction – a more reactive element displaces a less reactive one from its compound. An iron nail dipped in blue copper sulphate solution turns brownish and the solution fades: Fe + CuSO₄ → FeSO₄ + Cu. Zinc and lead also displace copper.

Double displacement reaction – ions are exchanged between two compounds. Sodium sulphate + barium chloride gives a white precipitate of BaSO₄. Any reaction producing an insoluble solid (a precipitate) is a precipitation reaction. Lead nitrate + potassium iodide gives yellow lead iodide.

Oxidation and reduction – heating brown copper powder forms black copper oxide (2Cu + O₂ → 2CuO); passing hydrogen over hot CuO turns it brown again (CuO + H₂ → Cu + H₂O).

- Oxidation: gain of oxygen or loss of hydrogen.

- Reduction: loss of oxygen or gain of hydrogen.

- When both happen together it is a redox reaction. In ZnO + C → Zn + CO, zinc oxide is reduced and carbon is oxidised.

5. Oxidation in everyday life

Corrosion: a metal attacked by moisture, acids or air. Iron gets a reddish-brown coating (rust), silver turns black, copper gets a green coating. Rusting damages car bodies, bridges, iron railings and ships, and wastes a lot of money every year.

Rancidity: fats and oils in food get oxidised and their smell and taste change. It is prevented by adding antioxidants, and chip bags are flushed with nitrogen, an unreactive gas, to stop oxidation.

Important Reactions (balanced)

2Mg(s) + O₂(g) → 2MgO(s)

Zn(s) + H₂SO₄(aq) → ZnSO₄(aq) + H₂(g)

3Fe(s) + 4H₂O(g) → Fe₃O₄(s) + 4H₂(g)

CaO(s) + H₂O(l) → Ca(OH)₂(aq) + Heat

2FeSO₄(s) → Fe₂O₃(s) + SO₂(g) + SO₃(g)

2Pb(NO₃)₂(s) → 2PbO(s) + 4NO₂(g) + O₂(g)

2AgCl(s) → 2Ag(s) + Cl₂(g) (sunlight)

Fe(s) + CuSO₄(aq) → FeSO₄(aq) + Cu(s)

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

CuO + H₂ → Cu + H₂O (heat)

Activities in the Chapter

Activity 1.1 – Clean a magnesium ribbon with sandpaper and burn it: it burns with a dazzling white flame to form white magnesium oxide powder. (Cleaning removes the oxide layer on its surface.)

Activity 1.4 – Add water to quick lime: the beaker becomes hot — an exothermic combination reaction.

Activity 1.5 / 1.6 – Heat ferrous sulphate crystals and lead nitrate: colour change and gases show thermal decomposition.

Activity 1.7 – Electrolysis of water: gas collected at one electrode is double the other; hydrogen burns with a pop.

Activity 1.8 – Silver chloride in sunlight turns grey.

Activity 1.9 – Iron nails in copper sulphate: nails turn brown, blue colour fades.

Activity 1.10 – Sodium sulphate + barium chloride: white precipitate.

Activity 1.11 – Heating copper powder in a china dish: surface turns black (copper oxide).

Board Exam Focus

Balancing equations and adding state symbols are asked almost every year.

Identify the type of reaction from an equation, and which substance is oxidised or reduced (e.g. in 2PbO + C → 2Pb + CO₂, lead oxide is reduced and carbon is oxidised — carbon dioxide is not the one being oxidised).

Know why decomposition is endothermic, why respiration is exothermic, why silver chloride is kept in dark bottles, and how rancidity is prevented.

Key Terms

Reactant – substance that changes in a reaction (अभिकारक)

Product – new substance formed (उत्पाद)

Balanced equation – equal atoms on both sides (संतुलित समीकरण)

Combination reaction (संयोजन अभिक्रिया)

Decomposition reaction (वियोजन अभिक्रिया)

Displacement reaction (विस्थापन अभिक्रिया)

Double displacement (द्विविस्थापन अभिक्रिया)

Precipitate – insoluble solid formed (अवक्षेप)

Exothermic – heat released (ऊष्माक्षेपी)

Endothermic – heat absorbed (ऊष्माशोषी)

Oxidation (उपचयन) / Reduction (अपचयन)

Redox reaction (रेडॉक्स अभिक्रिया)

Corrosion (संक्षारण)

Rancidity (विकृतगंधिता)

Common Mistakes

Changing a subscript to balance (writing H₂O₂ instead of 2H₂O) — only coefficients may change.

Forgetting that (g) for H₂O means steam, not liquid water.

Calling every reaction that gives off heat a combination reaction — respiration is exothermic but not a combination.

Mixing up oxidation and reduction: "gain of oxygen = oxidation", "loss of oxygen = reduction".

Thinking decomposition always needs heat — it can also use electricity or light.

💡 Key Learning Points

✓A chemical reaction is recognised by a change in state, colour or temperature, or by evolution of a gas.
✓Chemical equations must be balanced so that the number of atoms of each element is the same on both sides (law of conservation of mass).
✓Reactions are classified as combination, decomposition (thermal, electrolytic, photolytic), displacement, double displacement and redox.
✓Oxidation is gain of oxygen or loss of hydrogen; reduction is loss of oxygen or gain of hydrogen.
✓Corrosion and rancidity are harmful effects of oxidation in daily life, prevented by coatings, antioxidants and nitrogen flushing.

👨‍🏫 Teaching Tips

→Begin with the six daily-life situations (souring milk, rusting tawa, fermentation) and ask students what changed in each.
→Demonstrate balancing Fe + H₂O → Fe₃O₄ + H₂ on the board step by step, boxing each formula so students never alter subscripts.
→Perform Activity 1.4 (quick lime + water) and let students touch the beaker to feel an exothermic reaction.
→Show the silver chloride sample in sunlight vs in shade to make photolytic decomposition memorable, and link it to black-and-white photography.
→Use a sorting game: give 10 equations on cards and have groups classify them by reaction type.
→For redox, colour-code the substance gaining oxygen and the one losing oxygen in each equation.
→Bring an opened and an unopened chip packet to discuss nitrogen flushing and rancidity.

📋 Assessment Questions

Balance: BaCl₂ + Al₂(SO₄)₃ → BaSO₄ + AlCl₃ and add state symbols.
Classify CaCO₃ → CaO + CO₂ and explain why it is endothermic.
In CuO + H₂ → Cu + H₂O, name the substance oxidised and the one reduced.
Why does the colour of copper sulphate solution change when an iron nail is dipped in it?
Explain two methods of preventing rancidity.

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Next: Chapter 2 →Acids, Bases and Salts