MG SCHOOL SCIENCE

Friday, 6 November 2026

Chemical Reactions And Equations Notes

MG SCHOOL SCIENCE • CLASS 10

Chemical Reactions and Equations

Chapter 1 learning notes • Understand reactions, balance equations and revise the important concepts.

Chapter 01Study NotesQuick RevisionInteractive Quiz
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01 What is a Chemical Reaction?

In our daily lives, physical changes involve alterations in appearance or state without creating new substances, like melting ice. Chemical changes result in the formation of entirely new substances, such as cooking food or rusting iron.

Definition: A chemical reaction is a process in which one or more substances, called reactants, are converted into one or more different substances, called products.
REACTANTS

Starting substances

The substances that undergo chemical change in a reaction.

PRODUCTS

New substances

The different substances formed as a result of the reaction.

02 Chemical Equations

A chemical equation is a symbolic representation of a chemical reaction using the symbols and formulae of the substances involved. For example, magnesium burns in air and combines with oxygen to form magnesium oxide.

Example • Magnesium + Oxygen2Mg + O2 → 2MgO

Writing a chemical equation

The reactants are written on the left-hand side (LHS) of the arrow and the products on the right-hand side (RHS). A plus sign (+) links different reactants and products.

Physical state symbols

  • (s) – solid
  • (l) – liquid
  • (g) – gas
  • (aq) – aqueous solution

Other symbols

  • ↑ – gas produced
  • ↓ – precipitate formed
  • Conditions such as temperature, pressure and catalyst may be written near the arrow.

Balanced chemical equations

Law of Conservation of Mass: Atoms can neither be created nor destroyed in a chemical reaction. Therefore, the number of atoms of each element must be the same on both sides of a balanced equation.

A chemical equation is balanced when the number of atoms of each element on the reactant side is identical to the number on the product side.

Steps for balancing an equation

  1. Write the number of atoms of each element on the reactant and product sides.
  2. Balance the element that has the maximum number of atoms first.
  3. Adjust coefficients to make the number of each type of atom equal on both sides.
  4. Check the final number of atoms on both sides.
Example: Balance Fe + H2O → Fe3O4 + H2.
Fe + 4H2O → Fe3O4 + 4H2
3Fe(s) + 4H2O(g) → Fe3O4(s) + 4H2(g)

After balancing, the same number of atoms of each element is present on both sides. Physical states make the equation more informative.

Photosynthesis example: (aq) indicates substances dissolved in water and (l) indicates liquid water. Sunlight supplies energy absorbed by chlorophyll.
6CO2 + 6H2O + Sunlight → C6H12O6 + 6O2

03 Types of Chemical Reactions

Chemical reactions are classified into several categories considering different factors.

1. Combination Reaction
A reaction in which a single product is formed from two or more reactants is known as a combination reaction.

Example – Formation of slaked lime: Calcium oxide reacts vigorously with water and releases a large amount of heat.

CaO(s) + H2O(l) → Ca(OH)2(aq) + Heat

The calcium hydroxide solution is used for whitewashing walls. It reacts with carbon dioxide in air to form calcium carbonate, which gives the walls a shiny finish after two to three days. Marble also has the formula CaCO3.

Ca(OH)2 + CO2 → CaCO3 + H2O

Other examples of exothermic combination reactions

C(s) + O2(g) → CO2(g)
2H2(g) + O2(g) → 2H2O(l)
CH4(g) + 2O2(g) → CO2(g) + 2H2O(g)
Revision point: Calcium oxide reacts with water to form calcium hydroxide, which is used for whitewashing. Calcium hydroxide later reacts with carbon dioxide in air to form calcium carbonate on the walls.
2. Decomposition Reaction
A single reactant breaks down when exposed to heat, light or electricity, resulting in the formation of two or more products.

(a) Thermal decomposition

A decomposition reaction carried out by heating is called thermal decomposition.

  • Ferrous sulphate: Ferrous sulphate crystals (FeSO4·7H2O) lose water when heated and their green colour fades. They then decompose to ferric oxide, sulphur dioxide and sulphur trioxide.
  • Calcium carbonate: Heating calcium carbonate breaks it down into calcium oxide (lime) and carbon dioxide. Calcium oxide is used in cement production.
  • Lead nitrate: Heating lead nitrate powder releases brown fumes of nitrogen dioxide (NO2).
2FeSO4 → Fe2O3 + SO2 + SO3
CaCO3 → CaO + CO2
2Pb(NO3)2 → 2PbO + 4NO2 + O2

(b) Photolytic decomposition

Photolytic decomposition is a decomposition reaction caused by light.

  • Hydrogen peroxide breaks down into water and oxygen gas when exposed to sunlight.
  • White silver chloride turns grey in sunlight due to decomposition into silver and chlorine. Silver bromide behaves in the same way. These reactions are used in black-and-white photography.
2H2O2 → 2H2O + O2
2AgCl → 2Ag + Cl2
2AgBr → 2Ag + Br2

(c) Electrolytic decomposition

Electrolytic decomposition is a decomposition reaction caused by an electric current. In electrolysis of water, water breaks down into oxygen and hydrogen gases when electric current is passed through it.

2H2O(l) → 2H2(g) + O2(g)
Revision: Decomposition reactions are endothermic because they require an energy input (heat, light or electricity) to break down compounds.
3. Displacement Reaction
A displacement reaction is a chemical reaction in which a more reactive element displaces a less reactive element from its compound.

Iron nails in copper sulphate solution: The iron nail becomes brownish and the blue colour of copper sulphate solution fades. Iron displaces copper from copper sulphate solution.

Fe(s) + CuSO4(aq) → FeSO4(aq) + Cu(s)

Zinc and lead are also more reactive than copper and can displace copper from its compounds.

Zn + CuSO4 → ZnSO4 + Cu
4. Double Displacement Reaction
A double displacement reaction is a reaction in which two compounds react and their positive ions (cations) and negative ions (anions) switch places, forming two new compounds.

When barium chloride solution reacts with sodium sulphate solution, a white precipitate of barium sulphate is formed along with sodium chloride.

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

04 Endothermic and Exothermic Reactions

ENERGY RELEASED

Exothermic reaction

Energy is released from the reaction system into the surroundings, usually in the form of heat.

C(s) + O2(g) → CO2(g) + Heat

Example: Burning of carbon.

ENERGY ABSORBED

Endothermic reaction

Energy is absorbed from the surroundings in the form of heat, light or electricity.

6CO2 + 6H2O + Sunlight → C6H12O6 + 6O2

Example: Photosynthesis.

05 Oxidation and Reduction (Redox)

In a redox reaction, one substance is oxidised and another substance is reduced.

Oxidation

Oxidation is the gain of oxygen or loss of hydrogen/electrons by a substance.

Reduction

Reduction is the loss of oxygen or gain of hydrogen/electrons by a substance.

Oxidation of copper to copper oxide

On heating, copper powder becomes coated with black copper(II) oxide because oxygen is added to copper.

2Cu + O2 → 2CuO

When hydrogen gas is passed over heated copper(II) oxide, the black coating turns brown as copper is obtained. Copper(II) oxide loses oxygen and is reduced, while hydrogen gains oxygen and is oxidised.

CuO + H2 → Cu + H2O

Other examples of redox reactions

ZnO + C → Zn + CO
MnO2 + 4HCl → MnCl2 + 2H2O + Cl2

06 Effects of Oxidation Reactions in Everyday Life

1. Corrosion

Numerous metals react with moisture, oxygen and acids. Iron, for example, starts as shiny but eventually rusts, forming a reddish-brown powder due to oxidation.

The process of metal deterioration due to environmental substances is called corrosion.
  • Iron rusts; its oxides do not firmly attach, causing flaking, structural weakening and breakdown.
  • Copper develops a green coating.
  • Silver acquires a black coating.

Rusting of iron

Iron reacts with oxygen and moisture to form rust (hydrated ferric oxide).

Corrosion of copper

Copper reacts with substances in moist air to form a green coating of basic copper carbonate.

Corrosion of silver

Silver reacts with sulphur compounds in air to form a black coating of silver sulphide.

2. Rancidity

The taste and odour of food materials containing fats and oils change when they are left exposed to air for a long time. This process is called rancidity, and it is caused by oxidation of fats and oils present in food materials.

Methods to prevent rancidity

  • Seal food in air-tight containers.
  • Use nitrogen packaging (flush bags with nitrogen gas).
  • Store food in refrigeration.
  • Add antioxidants or preservatives.
Revision answer: The primary cause of rancidity in food containing fats and oils is oxidation of fats and oils when food materials are exposed to air.

07 Quick Revision – Important Points

Key definitions
  • Chemical reaction: Reactants are converted into products.
  • Chemical equation: Symbolic representation of a chemical reaction.
  • Combination: Two or more reactants form a single product.
  • Decomposition: One reactant breaks down into two or more products.
  • Displacement: A more reactive element displaces a less reactive element.
  • Double displacement: Ions of two compounds exchange places.
  • Corrosion: Deterioration of metals due to environmental substances.
  • Rancidity: Oxidation of fats and oils that changes food taste and odour.
Important equations at a glance
2Mg + O2 → 2MgO
3Fe + 4H2O → Fe3O4 + 4H2
CaO + H2O → Ca(OH)2 + Heat
CaCO3 → CaO + CO2
Fe + CuSO4 → FeSO4 + Cu
Na2SO4 + BaCl2 → BaSO4 + 2NaCl
CuO + H2 → Cu + H2O

08 Interactive MCQ Practice

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