Acids, Bases and Salts – CBSE Class 10 Science Notes & Guide

Acids, bases and salts are essential parts of everyday chemistry. They influence the taste of food, the cleanliness of our homes, the health of our bodies and many industrial processes. This guide summarises the main ideas, definitions, formulas and problem‑solving techniques you need to master the chapter for CBSE Class 10.

Key concepts

Properties of acids and bases

• Acids taste sour, turn blue litmus red, react with metals to release hydrogen gas, conduct electricity (electrolytes), and neutralise bases. Common examples: HCl, H₂SO₄, CH₃COOH.
• Bases feel slippery, taste bitter, turn red litmus blue, conduct electricity, and neutralise acids. Common examples: NaOH, KOH, Ca(OH)₂.
• Both acids and bases dissociate in water to give ions: acid → H⁺ + conjugate base, base → OH⁻ + conjugate acid.

pH scale and indicators

• pH = –log₁₀[H⁺]. The scale runs from 0 (strong acid) to 14 (strong base); 7 is neutral.
• Indicators are substances that change colour at particular pH ranges. Example: phenolphthalein (colourless in acid, pink in base), litmus paper, universal indicator.
• To calculate pH, use the concentration of hydrogen ions. Example: if [H⁺] = 1×10⁻⁴ M, pH = –log(1×10⁻⁴) = 4.

Neutralisation reactions

• A neutralisation reaction occurs when an acid and a base react to form water and a salt.
• General equation: acid + base → salt + water.
• The reaction is exothermic; heat is released.
• Example: HCl + NaOH → NaCl + H₂O.

Preparation of salts

• Direct combination: metal oxide or metal reacts with acid (e.g., CuO + 2H₂SO₄ → CuSO₄ + 2H₂O).
• Reaction of acid with metal: acid + metal → salt + hydrogen gas (e.g., Zn + 2HCl → ZnCl₂ + H₂).
• Reaction of base with acid (neutralisation) as above.
• Reaction of acid with carbonate: acid + carbonate → salt + water + CO₂ (e.g., HCl + Na₂CO₃ → 2NaCl + H₂O + CO₂).

Importance of salts in daily life

• Food: table salt (NaCl) enhances flavour and preserves food.
• Health: electrolytes like KCl and Na⁺/Cl⁻ maintain fluid balance.
• Cleaning: washing soda (Na₂CO₃) softens water; bleach (NaClO) disinfects.
• Industry: fertilizers (NH₄NO₃, K₂SO₄) provide essential nutrients for crops.
• Water treatment: calcium carbonate removes hardness.

Important definitions and formulas

Formulas

• pH = –log₁₀[H⁺]
• [H⁺] = 10⁻ᵖᴴ
• For strong acid/base, concentration of H⁺ (or OH⁻) ≈ molarity of the solution.
• Neutralisation stoichiometry: moles of H⁺ = moles of OH⁻.

Solved examples

Example 1

Question: 25 mL of 0.1 M HCl is mixed with 25 mL of 0.1 M NaOH. Calculate the pH of the resulting solution. Solution:

1. Find moles of H⁺: 0.1 M × 0.025 L = 0.0025 mol.
2. Find moles of OH⁻: 0.1 M × 0.025 L = 0.0025 mol.
3. Since moles are equal, complete neutralisation occurs; only water and NaCl remain.
4. The solution is neutral, so pH = 7.

Answer: pH = 7 (neutral).

Example 2

Question: 10 mL of 0.2 M H₂SO₄ is diluted to 250 mL. What is the pH of the diluted solution? (Assume complete dissociation of the first H⁺ only.) Solution:

1. Moles of H₂SO₄ = 0.2 M × 0.010 L = 0.002 mol.
2. First proton gives equal moles of H⁺: 0.002 mol.
3. Total volume after dilution = 0.250 L.
4. [H⁺] = 0.002 mol ÷ 0.250 L = 0.008 M.
5. pH = –log₁₀(0.008) ≈ –log₁₀(8×10⁻³) = 2.10.

Answer: pH ≈ 2.1 (acidic).

Example 3

Question: Write the balanced equation for the preparation of copper(II) sulfate by reacting copper(II) oxide with sulfuric acid, and calculate the mass of CuSO₄ formed when 5 g of CuO reacts with excess H₂SO₄. Solution:

1. Balanced equation: CuO + H₂SO₄ → CuSO₄ + H₂O.
2. Molar mass CuO = 79.5 g mol⁻¹; moles of CuO = 5 g ÷ 79.5 g mol⁻¹ = 0.0629 mol.
3. From the equation, 1 mol CuO gives 1 mol CuSO₄.
4. Molar mass CuSO₄ = 159.6 g mol⁻¹.
5. Mass of CuSO₄ = 0.0629 mol × 159.6 g mol⁻¹ = 10.04 g (≈10 g).

Answer: 10 g of CuSO₄ is produced.

Common mistakes to avoid

• Confusing concentration with amount: Always convert volume to litres before using molarity.
• Ignoring dilution factor: When a solution is diluted, the number of moles stays same, but concentration changes.
• Using wrong stoichiometric coefficients: Write balanced equations first; coefficients give the mole ratio.
• Assuming strong acids are completely dissociated for poly‑acidic compounds: For H₂SO₄, the second proton is weak; treat only the first H⁺ as fully ionised unless otherwise specified.
• Mixing up pH and pOH: pOH = –log[OH⁻]; pH + pOH = 14 at 25 °C.

Exam tips

• Memorise the pH ranges for common indicators; a quick table helps.
• Practice converting between pH, [H⁺] and [OH⁻] using the relation pH + pOH = 14.
• In neutralisation problems, always equate moles of H⁺ to moles of OH⁻.
• Write the complete balanced equation before calculating masses or volumes.
• For preparation of salts, identify the type of reaction (acid‑metal, acid‑carbonate, etc.) and apply the correct stoichiometry.

Frequently asked questions

What is the difference between a strong acid and a weak acid?

A strong acid dissociates completely in water, giving a high concentration of H⁺ ions (e.g., HCl). A weak acid only partially dissociates, so the [H⁺] is less than the initial concentration (e.g., CH₃COOH).

How does a universal indicator work?

Universal indicator is a mixture of several dyes that change colour over a wide pH range. Each component reacts at a specific pH, producing a smooth colour transition from red (acidic) to violet (basic).

Why does neutralisation release heat?

When H⁺ ions combine with OH⁻ ions to form water, strong O–H bonds are formed. The formation of these bonds releases energy, making the reaction exothermic.