Chapter Plan - Chemical reactions and equations I — lesson. Chemistry CBSE Live product, Class 10 (2026-27).

What is a chemical reaction?

A chemical reaction is a process in which one or more substances (reactants) are converted into new substances (products) with different properties.

\begin{array}{l} A + B \to C + D \\ Reactants Products \end{array}

Identification of a chemical reaction:

Change in colour
Evolution of gas
Change in temperature
Formation of precipitate
Change in state

Burning of magnesium ribbon:

Burning of magnesium ribbon

$2Mg + O_2 → 2MgO$

Type of reaction: Combination (also exothermic) reaction.

Formation of new substance white ash ($MgO$)
Type of reaction - Combination and exothermic (heat is released)
A dazzling white flame is produced

Balancing chemical equations:

Example of an unbalanced equation:

$Fe + H_2O → Fe_3O_4 + H_2$

Steps to balance (Hit-and-trial method):

Write a skeletal equation
Count the atoms of each element on both sides
Add coefficients to balance
Check for correctness
Include physical states

Balanced equation:

$3Fe({s}) + 4H_2O({g}) → Fe_3O_4({s}) + 4H_2({g})$

Types of chemical reactions:

1. Combination reaction:

When two or more reactants combine to form a single product.

$CaO + H_2O → Ca(OH)_2 + heat$

Combination reaction

Nature: Combination reactions are mostly exothermic (heat is released)

Application (Whitewashing): A shiny finish forms on walls due to calcium carbonate ($CaCO_3$)

$Ca(OH)_2 + CO_2 → CaCO_3 + H_2O$

2. Decomposition reaction:

A reaction where one reactant breaks down into two or more simpler products.

Note: Most decomposition reactions are endothermic as they require energy to break down compounds. However, the decomposition of organic or vegetable waste into compost is an exothermic process, as it releases heat.

Types of decomposition reaction:

i. Thermal decomposition (by heat):

Example 1:

2Pb (N O_{3})_{2 (s)} \overset{heat}{⟶} 2Pb O_{(s)} + 4N O_{2 (g)} + O_{2 (g)}

$Pb(NO_3)_2$: White or colourless solid
$PbO$: Yellow residue
$NO_2$: Brown fumes

Example 2:

2FeS O_{4 (s)} \overset{heat}{⟶} {Fe}_{2} O_{3 (s)} + S O_{2 (g)} + {SO}_{3 (g)}

Reddish brown residue - $Fe_2O_3$
Pungent smell - $SO_2$

ii. Electrolytic decomposition (by electricity):

Example: Electrolysis of water

2 H_{2} O_{(l)} \overset{Electricity}{⟶} {2 H}_{2 (g)} + O_{2 (g)}

Gas liberated at the cathode (hydrogen) and the anode (oxygen)
Confirmation test: Hydrogen (pop sound) and oxygen (glows brightly)
Volume ratio: $H_2 : O_2$ $=2:1$
A few drops of $H_2SO_4$ is added to increase conductivity

iii. Photolytic decomposition (by light):

Certain compounds decompose in the presence of sunlight.

2 Ag {Cl}_{(s)} \overset{Sunlight}{⟶} 2 {Ag}_{(s)} + {Cl}_{2 (g)}

2 Ag {Br}_{(s)} \overset{Sunlight}{⟶} 2 {Ag}_{(s)} + {Br}_{2 (g)}

Pale yellow silver bromide ($AgBr$) decomposes under light to greyish white silver ($Ag$) and bromine $Br_2$.
Application: Used in black and white photography

Displacement reaction:

A more reactive element displaces a less reactive element from its compound. It is also called as single displacement reaction.

Example:

Reaction between iron and copper sulphate

Blue colour of copper sulphate solution fades and a pale green colour iron sulphate solution is formed.
Brown coating forms on the nail.
Iron is more reactive and displaces copper.

Double displacement reaction:

The exchange of ions between the reactants to form new products. Double displacement reaction has the following two types:

\begin{array}{l} A + B \to C + D \\ Reactants Products \end{array}

i. Neutralisation reaction:

Acid + Base → Salt + Water

ii. Precipitation reaction:

1. $Ba {Cl}_{2 (aq)} + N a_{2} S O_{4 (aq)} \to BaS O_{4 (s)} ↓ + 2 Na {Cl}_{(aq)}$

Reactants: Both are colourless solutions

Products: White precipitate ($BaSO_4$)

2. $Pb(NO_3)_2(aq) + 2KI(aq) → PbI_2(s) + 2KNO_3(aq)$

Reactants: Both are colourless solutions

Products: Yellow precipitate $PbI_2$

Redox reaction:

Oxidation	Reduction
Gain of oxygen	Loss of oxygen
Loss of hydrogen	Gain of hydrogen
Loss of electron	Gain of electron
Increase in oxidation number	Decrease in oxidation number

Redox: Both oxidation and reduction occur simultaneously.

Example:

Redox reaction

Heat copper powder → turns black ($CuO$)
Pass hydrogen → turns brown again ($Cu$ metal)
Copper oxide gets reduced
Hydrogen gets oxidised

Agent	What is does?	What happens to it?	Example
Oxidising agent	Causes oxidation of other substance	It gets reduced	In $CuO + H_2 → Cu + H_2O$, $CuO$ is the oxidising agent
Reducing agent	Causes reduction of other substance	It gets oxidised	In $CuO + H_2 → Cu + H_2O$, $H_2$ is the reducing agent

Corrosion:

Corrosion is the progressive destruction of metals by the action of air, moisture or chemicals (such as an acid) on their surface.

Hints:

Iron oxide - Reddish brown (Rust)
Silver sulphide - Black (Tarnish)
Copper carbonate - Green (Patina)

Rancidity:

The oxidation of oils and fats present in food material resulting in bad smell and taste.

Methods to prevent rancidity: Antioxidants, air tight containers, and flushing packets with nitrogen.

Change in temperature:

i. Exothermic - Release of heat
Examples: Combustion, combination, respiration, rusting of iron, neutralisation (mostly), decomposition of vegetable waste.

ii. Endothermic - Absorption of heat

Examples: Photosynthesis, melting of ice, decomposition (mostly), dissolving ammonium chloride/nitrate.

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1. Chapter Plan - Chemical reactions and equations I

Theory: