CBSE Class 10 Science Notes: Chemical Reactions and Equations

What is a Chemical Reaction?

A chemical reaction is a process that involves the rearrangement of atoms and molecules, leading to the formation of new substances with different properties. It’s a fundamental concept in chemistry, describing how matter changes.

Everyday Examples:

• Rusting of Iron: Iron reacts with oxygen and water in the air to form iron oxide (rust).
• Burning of Fuel: Combustion, like burning wood or gasoline, involves rapid reaction with oxygen, producing heat and light.
• Respiration: The process of breaking down glucose in your body to produce energy.

Chemical Equations

A chemical equation is a symbolic representation of a chemical reaction using formulas and symbols of the reactants and products. It provides a concise way to describe what happens during a reaction.

Writing Chemical Equations:

Reactants are written on the left side, products on the right side, separated by an arrow ($\rightarrow$). Plus signs (+) are used to separate multiple reactants or products.

Example:

Formation of water from hydrogen and oxygen: $2H_2 (g) + O_2 (g) \rightarrow 2H_2O (l)$

Where: $H_2$ = Hydrogen, $O_2$ = Oxygen, $H_2O$ = Water, (g) = Gas, (l) = Liquid

Balancing Chemical Equations

Balanced chemical equations follow the law of conservation of mass, which states that matter cannot be created or destroyed. Balancing ensures the same number of atoms of each element are present on both sides of the equation.

Simple Balancing Steps:

• Write the unbalanced equation.
• Count the number of atoms of each element on both sides.
• Add coefficients (numbers in front of the formulas) to balance the number of atoms of each element. Do not change the subscripts within the formulas.
• Check the equation to ensure it’s balanced.

Example (Balancing the formation of ammonia):

Unbalanced: $N_2 + H_2 \rightarrow NH_3$

Balanced: $N_2 + 3H_2 \rightarrow 2NH_3$

Types of Chemical Reactions

1. Combination (Synthesis) Reaction:

Two or more reactants combine to form a single product.

Example: Formation of Magnesium Oxide: $2Mg (s) + O_2 (g) \rightarrow 2MgO (s)$

2. Decomposition Reaction:

A single reactant breaks down into two or more simpler products.

Example: Decomposition of Calcium Carbonate: $CaCO_3 (s) \rightarrow CaO (s) + CO_2 (g)$

3. Displacement (Single Replacement) Reaction:

A more reactive element displaces a less reactive element from a compound.

Example: Reaction of Zinc with Copper Sulfate: $Zn (s) + CuSO_4 (aq) \rightarrow ZnSO_4 (aq) + Cu (s)$

4. Double Displacement Reaction:

Exchange of ions between two compounds.

Example (Precipitation): Reaction of Barium Chloride and Sodium Sulfate: $BaCl_2 (aq) + Na_2SO_4 (aq) \rightarrow BaSO_4 (s) + 2NaCl (aq)$ (Formation of a solid precipitate, $BaSO_4$)

Endothermic vs. Exothermic Reactions

These reactions are classified based on their heat transfer.

Definitions:

• Exothermic Reactions: Release heat into the surroundings (feel warm).
• Endothermic Reactions: Absorb heat from the surroundings (feel cold).

Examples:

• Exothermic: Burning of methane gas.
• Endothermic: Melting of ice.

Oxidation and Reduction

These processes always occur together in a redox reaction.

Basic Idea:

• Oxidation: Loss of electrons or gain of oxygen.
• Reduction: Gain of electrons or loss of oxygen.

Examples:

In the reaction: $2Mg (s) + O_2 (g) \rightarrow 2MgO (s)$,

Magnesium ($Mg$) is oxidized (loses electrons) and Oxygen ($O_2$) is reduced (gains electrons).