Class 10 Science – Quick Revision Sheet
Chapter: Acids, Bases and Salts
1. Key Definitions
| Term | Definition (English) | Hindi (optional) |
|---|
| Acid | Substance that donates H⁺ ions (proton) in aqueous solution (Arrhenius) or accepts a pair of electrons (Lewis). | अम्ल |
| Base | Substance that produces OH⁻ ions in water (Arrhenius) or donates a pair of electrons (Lewis). | क्षार |
| Salt | Ionic compound formed when an acid reacts with a base, consisting of the cation of the base and the anion of the acid. | लवण |
| Neutralisation | Reaction of an acid with a base to give a salt + water. | निष्पक्षीकरण |
| pH | Measure of acidity: pH = –log[H⁺]. | pH |
| pOH | Measure of basicity: pOH = –log[OH⁻]. | pOH |
| Strong Acid/Base | Completely dissociates in water (≈100 % ionisation). | मजबूत अम्ल/क्षार |
| Weak Acid/Base | Partially dissociates in water (≤ 5 % ionisation). | कमजोर अम्ल/क्षार |
| Acidic Salt | Salt formed from partial neutralisation of a poly‑basic acid (contains replaceable H⁺). | अम्लीय लवण |
| Basic Salt | Salt formed from partial neutralisation of a poly‑basic base (contains OH⁻). | क्षारीय लवण |
| Normal Salt | Salt formed from complete neutralisation of a monoprotic acid and a base. | सामान्य लवण |
2. Fundamental Laws & Concepts
| Law / Concept | Statement |
|---|
| Arrhenius Theory (1887) | Acid → H⁺ in water; Base → OH⁻ in water. |
| Bronsted‑Lowry Theory (1923) | Acid = proton donor; Base = proton acceptor. |
| Lewis Theory (1923) | Acid = electron‑pair acceptor; Base = electron‑pair donor. |
| Law of Neutralisation | One mole of H⁺ reacts with one mole of OH⁻ to give one mole of H₂O. |
| pH + pOH = 14 (at 25 °C) | Relates acidity and basicity of aqueous solutions. |
| Strong Acid/Base pH | For a strong acid of concentration C (mol L⁻¹): <br> pH = –log C <br>For a strong base: <br> pOH = –log C <br>Then use pH = 14 – pOH. |
| Weak Acid Approximation | For a weak acid HA with Ka and initial concentration C: <br> [H⁺] ≈ √(Ka · C) <br> pH ≈ ½(pKa – log C). |
| Indicator Principle | Indicators change colour at a specific pH range because they are weak acids/bases themselves. |
3. Important Formulas
| Formula | Meaning |
|---|
| pH = –log[H⁺] | Calculates acidity from hydrogen‑ion concentration. |
| [H⁺] = 10^(–pH) | Finds concentration when pH is known. |
| pOH = –log[OH⁻] | Calculates basicity from hydroxide‑ion concentration. |
| [OH⁻] = 10^(–pOH) | Finds hydroxide concentration from pOH. |
| pH + pOH = 14 | Relationship at 25 °C (water ion product Kw = 1 × 10⁻¹⁴). |
| Neutralisation: Acid + Base → Salt + H₂O | General equation. |
| Strong Acid (C mol L⁻¹): pH = –log C | Complete dissociation. |
| Strong Base (C mol L⁻¹): pOH = –log C | Complete dissociation. |
| Weak Acid (HA): [H⁺] ≈ √(Ka·C) | Approximation for Ka ≪ 1. |
| Weak Base (B): [OH⁻] ≈ √(Kb·C) | Similar to weak acid. |
4. Representative Reactions
| Reaction Type | Example Equation | Notes |
|---|
| Acid + Metal | Zn + 2 HCl → ZnCl₂ + H₂↑ | Hydrogen gas liberated. |
| Acid + Carbonate | 2 HCl + CaCO₃ → CaCl₂ + H₂O + CO₂↑ | Effervescence (CO₂). |
| Acid + Base (Neutralisation) | H₂SO₄ + 2 NaOH → Na₂SO₄ + 2 H₂O | Forms normal salt (sodium sulfate). |
| Poly‑basic Acid – Partial Neutralisation → Acid Salt | H₂SO₄ + NaOH → NaHSO₄ + H₂O | Sodium bisulphate (acidic salt). |
| Poly‑basic Base – Partial Neutralisation → Basic Salt | Ca(OH)₂ + H₂SO₄ → CaSO₄ + 2 H₂O (complete) <br>Ca(OH)₂ + H₂SO₄ (1 : 0.5) → Ca(OH)SO₄ + H₂O (basic salt) | Calcium hydrogen sulphate is a basic salt. |
| Indicator Reaction | Phenolphthalein: colourless (pH < 8.2) → pink (pH > 8.2) | Used to detect endpoint of titration. |
5. Common Acids, Bases & Salts (Class‑10 Level)
| Category | Common Examples | Formula |
|---|
| Strong Acids | Hydrochloric, Sulphuric, Nitric | HCl, H₂SO₄, HNO₃ |
| Weak Acids | Acetic, Citric, Carbonic | CH₃COOH, C₆H₈O₇, H₂CO₃ |
| Strong Bases | Sodium hydroxide, Potassium hydroxide, Calcium hydroxide | NaOH, KOH, Ca(OH)₂ |
| Weak Bases | Ammonia, Aluminium hydroxide (sparingly soluble) | NH₃, Al(OH)₃ |
| Normal Salts | Sodium chloride, Potassium nitrate | NaCl, KNO₃ |
| Acid Salts | Sodium hydrogen sulphate, Potassium hydrogen phosphate | NaHSO₄, K₂HPO₄ |
| Basic Salts | Calcium carbonate, Magnesium hydroxide (as salt of weak acid) | CaCO₃, Mg(OH)₂ |
6. Quick Tips for Exams
- Identify the type of acid/base – strong ⇢ use direct pH = –log C; weak ⇢ use Ka/Kb approximation.
- Remember pH + pOH = 14 – handy for converting between acidity and basicity.
- Neutralisation stoichiometry – 1 mol H⁺ reacts with 1 mol OH⁻ → 1 mol H₂O. Use this to calculate volumes in titrations.
- Acidic vs. Basic salts – look at the number of replaceable H⁺ or OH⁻ left after partial neutralisation.
- Indicator colour range – memorize the pH range for common indicators (e.g., methyl orange 3.1–4.4, phenolphthalein 8.2–10).
7. Summary Flow‑Chart
Acid (donates H⁺) + Base (accepts H⁺) → Salt + H₂O
| |
Strong → Complete dissociation |
Weak → Use Ka/Kb for [H⁺] or [OH⁻] |
Use the above formulas and concepts to solve numerical problems, write balanced equations, and identify salts in NCERT‑style questions. Good luck!