Chemical Reactions and Equations – CBSE Class 10 Science Notes & Guide

Chemical reactions are the heart of chemistry. In everyday life we see rusting of iron, digestion of food, combustion of fuel and many more. Understanding how substances combine, change and release energy helps us explain natural phenomena and solve practical problems.

Key concepts

Types of chemical reactions

Chemical reactions are broadly classified into four types:

1. Combination (Synthesis) – Two or more reactants combine to form a single product. Example: 2H₂ + O₂ → 2H₂O.
2. Decomposition – A single compound breaks down into two or more simpler substances. Example: 2HgO → 2Hg + O₂.
3. Displacement (Single‑replacement) – An element displaces another from its compound. Example: Zn + 2HCl → ZnCl₂ + H₂.
4. Double‑displacement (Metathesis) – Cations and anions of two different compounds exchange places. Example: AgNO₃ + NaCl → AgCl + NaNO₃.

Each type follows a characteristic pattern that helps us predict the products.

Balancing chemical equations

A chemical equation must obey the law of conservation of mass – the number of atoms of each element is the same on both sides. To balance:

1. Write the correct formulas of reactants and products.
2. Count the atoms of each element on both sides.
3. Adjust coefficients (whole numbers placed before formulas) to make the counts equal.
4. Check again for all elements; the smallest whole‑number coefficients are the final answer.

Balancing reinforces the idea that matter is neither created nor destroyed in a chemical change.

Oxidation and reduction

Oxidation‑reduction (redox) reactions involve transfer of electrons.

• Oxidation: loss of electrons; the species becomes more positive (oxidation number increases).
• Reduction: gain of electrons; the species becomes more negative (oxidation number decreases).

The substance that gains electrons is the oxidizing agent; the one that loses electrons is the reducing agent. Identifying oxidation numbers before and after the reaction makes it easy to spot the electron transfer.

Acid–base reactions

Acids release H⁺ ions in water, while bases release OH⁻ ions. When they react, they form water and a salt – a neutralisation reaction. General form: Acid + Base → Salt + Water For strong acids and bases the reaction proceeds completely; for weak ones the extent depends on their dissociation constants.

Indicators and pH

• Indicators are substances that change colour at particular pH ranges (e.g., phenolphthalein turns pink above pH 8.2). They help us detect the endpoint of a titration.
• pH measures the acidity or alkalinity of a solution: pH = –log₁₀[H⁺].
• pH < 7 → acidic
• pH = 7 → neutral
• pH > 7 → basic

A pH meter or universal indicator paper gives a quick estimate of the solution’s nature.

Important definitions and formulas

• Chemical equation – Symbolic representation of a chemical reaction.
• Balancing coefficient – Whole number placed before a formula to satisfy the conservation of atoms.
• Oxidation number – Charge an atom would have if all bonds were ionic.
• Redox – Reaction involving transfer of electrons.
• Acid – Substance that increases H⁺ concentration in aqueous solution.
• Base – Substance that increases OH⁻ concentration in aqueous solution.
• pH – pH = –log₁₀[H⁺]; pOH = –log₁₀[OH⁻]; pH + pOH = 14 (at 25 °C).
• Neutralisation – Acid + Base → Salt + Water.

Solved examples

Example 1 – Balancing a combustion reaction

Problem: Balance the equation for the combustion of propanol, C₃H₈O. Solution:

1. Write the unbalanced equation: C₃H₈O + O₂ → CO₂ + H₂O
2. Count atoms on each side:

• C: 3 on left, 1 on right → need 3 CO₂.
• H: 8 on left, 2 on right → need 4 H₂O.
• O: left = 1 (from C₃H₈O) + 2x (from O₂), right = 3×2 (from 3CO₂) + 4×1 (from 4H₂O) = 6 + 4 = 10. Set up: 1 + 2x = 10 → 2x = 9 → x = 4.5.

1. Multiply whole equation by 2 to remove fraction: 2C₃H₈O + 9O₂ → 6CO₂ + 8H₂O.

Balanced equation: 2C₃H₈O + 9O₂ → 6CO₂ + 8H₂O.

Example 2 – Identifying oxidation and reduction

Problem: In the reaction Zn + CuSO₄ → ZnSO₄ + Cu, state the oxidising and reducing agents. Solution:

• Assign oxidation numbers: Zn (0) → Zn²⁺ (in ZnSO₄) → loss of 2 e⁻ → oxidation. Cu²⁺ (in CuSO₄) → Cu (0) → gain of 2 e⁻ → reduction.
• Therefore, Zn is the reducing agent (donates electrons) and Cu²⁺ is the oxidising agent (accepts electrons).

Example 3 – Calculating pH of a dilute acid

Problem: Find the pH of a 1.0 × 10⁻³ M HCl solution. Solution:

1. HCl is a strong acid, it dissociates completely: [H⁺] = 1.0 × 10⁻³ M.
2. pH = –log₁₀[H⁺] = –log₁₀(1.0 × 10⁻³) = 3.

Answer: pH = 3 (acidic).

Common mistakes to avoid

• Skipping the formula check – Writing wrong chemical formulas leads to an impossible balance.
• Changing subscripts – Only coefficients can be altered while balancing; subscripts represent the identity of the compound.
• Ignoring oxidation numbers – Mis‑assigning them gives wrong identification of redox partners.
• Assuming all acids are strong – Weak acids (e.g., acetic acid) do not dissociate completely; using [H⁺] = concentration gives an over‑estimated acidity.
• Using pH = –log[H⁺] without unit conversion – Concentration must be in mol L⁻¹; molarity of a solution prepared with grams must be calculated first.

Exam tips

• Memorise the four types of reactions with one example each; many MCQs ask to classify a given equation.
• Practice balancing by starting with the most complex molecule.
• Write oxidation numbers before tackling redox problems; it saves time.
• Remember the neutralisation formula: acid + base → salt + water. Check the states of the products (often aqueous).
• For pH questions, keep the log table handy and remember that pH + pOH = 14 at 25 °C.
• Use indicators wisely – know the colour range of common ones (litmus, phenolphthalein, methyl orange).

Frequently asked questions

What is the difference between a synthesis and a decomposition reaction?

In a synthesis (combination) reaction, two or more reactants form a single product (A + B → AB). In a decomposition reaction, one compound breaks into two or more simpler substances (AB → A + B).

How can I tell if a reaction is a redox reaction?

Check the oxidation numbers of each element before and after the reaction. If any element’s oxidation number increases, it is oxidised; if it decreases, it is reduced. Presence of both changes confirms a redox reaction.

Why does phenolphthalein turn pink in basic solutions but remains colourless in acidic ones?

Phenolphthalein is colourless in acidic and neutral media because its molecular structure does not absorb visible light. In basic solutions (pH > 8.2) it loses a proton, forming a coloured ion that absorbs light, giving a pink colour.