Chemical Reactions and Equations – WBBSE Class 10 Physical Science Notes & Guide

Chemical reactions are the heart of chemistry. They show how substances transform, release or absorb energy and obey the law of conservation of mass. Understanding different kinds of reactions, how to write and balance equations, and what influences their speed will help you solve problems confidently in the board exam.

Key concepts

Types of chemical reactions

1. Combination (Synthesis) reaction – Two or more simple substances combine to form a more complex product. General form: A + B → AB
2. Decomposition reaction – A single compound breaks down into two or more simpler substances. General form: AB → A + B
3. Displacement (Single‑replacement) reaction – An element displaces another from a compound. General form: A + BC → AC + B
4. Redox (Oxidation‑reduction) reaction – Involves transfer of electrons; one species is oxidised and another reduced. Key idea: Oxidation = loss of electrons, Reduction = gain of electrons.

Balancing chemical equations

• The number of atoms of each element must be the same on both sides of the equation.
• Use the lowest whole‑number coefficients.
• Follow the order: balance metals, then non‑metals, and finally hydrogen and oxygen.

Exothermic and endothermic reactions

• Exothermic: Energy is released to the surroundings (temperature of surroundings rises). Example: combustion of methane.
• Endothermic: Energy is absorbed from the surroundings (temperature of surroundings falls). Example: thermal decomposition of calcium carbonate.

Factors affecting rate of reaction

Energy changes in reactions

• Activation energy (Eₐ): Minimum energy that reacting particles must possess to collide successfully.
• Energy profile diagram: Shows reactants, products, activation energy and overall enthalpy change (ΔH).
• ΔH < 0 → exothermic; ΔH > 0 → endothermic.

Practical: preparation of hydrogen gas

Apparatus: Test tube, rubber tubing, water trough, delivery tube, porcelain crucible, solid zinc, dilute H₂SO₄. Procedure (brief):

1. Place a small piece of zinc in a test tube.
2. Add a few drops of dilute sulphuric acid.
3. Quickly attach a rubber tube to the mouth of the test tube and pass it into a water trough upside‑down.
4. Collect the gas displaced water – the collected gas is hydrogen.

Observations: Bubbles are produced rapidly; the collected gas extinguishes a glowing splint (pop test). Safety note: Perform the experiment behind a screen; handle acid with gloves.

Important definitions and formulas

• Chemical equation: Symbolic representation of a chemical reaction.
• Mole ratio: Ratio of coefficients in a balanced equation; used for stoichiometric calculations.
• ΔH (change in enthalpy): ΔH = Σ(ΔHf products) – Σ(ΔHf reactants).
• Activation energy (Eₐ): Energy barrier shown on an energy‑profile diagram.
• Rate of reaction (average): Rate = Δ[Reactant] ÷ Δt or Δ[Product] ÷ Δt.

Solved examples

Example 1 – Balancing a synthesis reaction

Question: Balance the equation for the formation of aluminium oxide.

Unbalanced: Al + O₂ → Al₂O₃

Solution:

1. List atoms on each side.

• Reactants: Al = 1, O = 2
• Products: Al = 2, O = 3

1. Balance aluminium first: place coefficient 2 in front of Al. → 2 Al + O₂ → Al₂O₃
2. Now aluminium is balanced (2 each). Balance oxygen: need 3 O atoms on product side, so place coefficient 3/2 in front of O₂. → 2 Al + 3/2 O₂ → Al₂O₃
3. Multiply whole equation by 2 to remove fraction. → 4 Al + 3 O₂ → 2 Al₂O₃

Balanced equation: 4 Al + 3 O₂ → 2 Al₂O₃

Example 2 – Identifying reaction type & ΔH

Question: The reaction 2 H₂ + O₂ → 2 H₂O releases heat. State the type of reaction and whether it is exothermic or endothermic.

Solution:

• Two simple molecules combine to give a single product → Combination (synthesis) reaction.
• Heat is released, so the reaction is exothermic (ΔH < 0).

Example 3 – Effect of temperature on rate (quantitative)

Question: In a particular reaction, the rate at 300 K is 1.2 × 10⁻³ mol L⁻¹ s⁻¹. When the temperature is raised to 320 K, the rate becomes 3.6 × 10⁻³ mol L⁻¹ s⁻¹. By what factor does the rate increase?

Solution: Rate factor = (Rate at 320 K) ÷ (Rate at 300 K) = (3.6 × 10⁻³) ÷ (1.2 × 10⁻³) = 3. Answer: The rate triples (increases by a factor of 3) when temperature is raised from 300 K to 320 K.

Common mistakes to avoid

• Forgetting to balance atoms of hydrogen and oxygen last. Doing them early often leads to repeated changes.
• Using fractional coefficients and leaving them as is. Multiply to obtain the smallest whole numbers.
• Confusing exothermic with endothermic signs. Remember: exothermic → heat given out → ΔH negative.
• Ignoring the role of a catalyst. It changes the rate but not the ΔH of the reaction.
• Mixing up displacement with double‑replacement reactions. Single‑displacement involves an element replacing another; double‑replacement involves exchange of ions between two compounds.

Exam tips

• Write the balanced equation first; never start solving stoichiometry with an unbalanced equation.
• Memorise the general forms of the four main reaction types – it helps in quick identification.
• When a question asks for the type of reaction, look at the reactants and products rather than the word “type”.
• In rate‑related questions, remember the qualitative trends: ↑ concentration, ↑ temperature, ↑ surface area, catalyst → ↑ rate.
• For the practical on hydrogen, note the colour‑less nature of the gas, its ability to turn a glowing splint into a pop sound, and the safety precautions.

Frequently asked questions

What is the difference between a redox reaction and a displacement reaction?

A redox reaction specifically involves transfer of electrons (oxidation‑reduction). A displacement reaction may be redox (e.g., Zn + CuSO₄ → ZnSO₄ + Cu) but the classification is based on the replacement of one element by another, not on electron transfer alone.

How can I check if my balanced equation is correct?

Count the number of atoms of each element on both sides. They must be equal. Also ensure the total charge is the same on both sides for ionic equations.

Why does a catalyst not appear in the overall chemical equation?

A catalyst participates in the reaction pathway but is regenerated at the end, so it is unchanged overall. Hence it is omitted from the net equation.