ICSE Class 10 Acids, Bases and Salts — Quick Revision
"Acids, Bases and Salts" is one of the chapters ICSE examiners return to again and again, because it connects definitions, ionic theory, the pH scale and a long list of preparation methods. This revision sheet puts all of it in one place — the definitions in the exact wording the paper expects, the pH calculations worked out, the salt types with real examples, and the preparation routes you must be able to name.
Definitions — know all three levels
Arrhenius definition (the one ICSE expects first): an acid is a substance which, when dissolved in water, produces hydrogen ions (H⁺) as the only positive ions. A base produces hydroxide ions (OH⁻) as the only negative ions in aqueous solution. A base soluble in water is called an alkali.
NaOH → Na⁺ + OH⁻
Bronsted–Lowry definition: an acid is a proton (H⁺) donor, a base is a proton acceptor. This explains why NH₃, which contains no OH group, is still a base: NH₃ + H₂O ⇌ NH₄⁺ + OH⁻.
Lewis definition: an acid is an electron-pair acceptor, a base an electron-pair donor. This is the broadest of the three and covers species such as AlCl₃ and BF₃.
Two facts ICSE tests directly: a substance behaves as an acid only in the presence of water (dry HCl gas does not turn dry blue litmus red), and acids conduct electricity in aqueous solution because they furnish ions.
Strong and weak, concentrated and dilute
These two pairs are not the same thing, and confusing them is a classic lost mark.
| Term | What it describes | Examples |
|---|---|---|
| Strong acid | Almost completely ionised in water | HCl, H₂SO₄, HNO₃ |
| Weak acid | Only partly ionised in water | CH₃COOH, H₂CO₃, HCOOH |
| Strong base (alkali) | Almost completely ionised | NaOH, KOH |
| Weak base | Only partly ionised | NH₄OH, Ca(OH)₂, Mg(OH)₂ |
| Concentrated | Large amount of acid in a small amount of water | Concentrated CH₃COOH is still a weak acid |
| Dilute | Small amount of acid in a large amount of water | Dilute HCl is still a strong acid |
The pH scale
pH measures how acidic or alkaline a solution is on a scale from 0 to 14.
pH + pOH = 14 (at 25 °C) [H⁺] × [OH⁻] = 1 × 10⁻¹⁴
pH below 7 is acidic, exactly 7 is neutral, above 7 is alkaline. The scale is logarithmic: a solution of pH 3 has ten times more H⁺ ions than one of pH 4, and a hundred times more than pH 5.
Example 1. A solution has [H⁺] = 1 × 10⁻³ mol L⁻¹. Find its pH.
pH = − log (10⁻³) = 3 — acidic.
Example 2. Find the pH of 0.01 M HCl.
HCl is a strong acid and gives one H⁺ per molecule, so [H⁺] = 0.01 = 10⁻² mol L⁻¹.
pH = − log (10⁻²) = 2.
Example 3. Find the pH of 0.001 M NaOH.
[OH⁻] = 10⁻³ mol L⁻¹ → pOH = 3
pH = 14 − 3 = 11 — alkaline.
Example 4. Rainwater has pH 5.6. Find [H⁺].
[H⁺] = 10−5.6 = 2.5 × 10⁻⁶ mol L⁻¹ — mildly acidic, because
dissolved CO₂ forms carbonic acid.
Neutralisation
An acid and a base react to give a salt and water. Stripped to its ions, every neutralisation is the same single reaction:
H⁺(aq) + OH⁻(aq) → H₂O(l) (the ionic equation, always exothermic)
Examples: NaOH + HCl → NaCl + H₂O; 2NaOH + H₂SO₄ → Na₂SO₄ + 2H₂O; CuO + H₂SO₄ → CuSO₄ + H₂O.
Q. 25 mL of 0.1 M NaOH is exactly neutralised by 0.05 M HCl. What volume of acid is needed?
NaOH + HCl → NaCl + H₂O, so they react 1 : 1.
moles of NaOH = 0.1 × (25 ÷ 1000) = 0.0025 mol
moles of HCl required = 0.0025 mol
volume = 0.0025 ÷ 0.05 = 0.05 L = 50 mL
The answer makes sense: the acid is half as concentrated, so twice the volume is needed.
Types of salts
| Salt type | Definition | Examples |
|---|---|---|
| Normal salt | All replaceable hydrogen of the acid is replaced by a metal or ammonium ion | NaCl, Na₂SO₄, KNO₃, CaCO₃ |
| Acid salt | Only part of the replaceable hydrogen is replaced; formed from a polybasic acid | NaHSO₄, NaHCO₃, Ca(HCO₃)₂ |
| Basic salt | Formed from partial neutralisation of a polyacidic base; still contains OH or O | Pb(OH)Cl, Zn(OH)Cl, basic copper carbonate |
| Double salt | Two simple salts crystallised together; ionises into all its ions in solution | Potash alum K₂SO₄·Al₂(SO₄)₃·24H₂O; Mohr's salt FeSO₄·(NH₄)₂SO₄·6H₂O |
| Complex salt | Contains a complex ion that does not break up in solution | Sodium argentocyanide Na[Ag(CN)₂]; potassium ferrocyanide K₄[Fe(CN)₆] |
The distinction between a double salt and a complex salt is a favourite question. A double salt gives the tests of all its component ions in solution; a complex salt does not, because the complex ion stays intact — a solution of Na[Ag(CN)₂] gives no test for ordinary Ag⁺ ions.
What kind of salt is it? — the acid–base parents rule
| Made from | Nature of the solution | Example |
|---|---|---|
| Strong acid + strong base | Neutral, pH ≈ 7 | NaCl, KNO₃ |
| Strong acid + weak base | Acidic, pH < 7 | NH₄Cl, CuSO₄ |
| Weak acid + strong base | Alkaline, pH > 7 | CH₃COONa, Na₂CO₃ |
| Weak acid + weak base | Close to neutral | CH₃COONH₄ |
Methods of preparing salts
- Direct combination of two elements — used for salts of metals with variable valency: 2Fe + 3Cl₂ → 2FeCl₃.
- Simple displacement — an active metal displaces hydrogen from a dilute acid: Zn + H₂SO₄ → ZnSO₄ + H₂.
- Precipitation (double decomposition) — used for insoluble salts: Pb(NO₃)₂ + 2KI → PbI₂↓ + 2KNO₃.
- Neutralisation of an alkali by titration — used when both the base and the salt are soluble: NaOH + HCl → NaCl + H₂O.
- Neutralisation of an insoluble base or oxide — add excess solid to the warm acid, then filter off what did not dissolve: CuO + H₂SO₄ → CuSO₄ + H₂O.
- Action of a dilute acid on a carbonate or bicarbonate: CaCO₃ + 2HCl → CaCl₂ + H₂O + CO₂; NaHCO₃ + HCl → NaCl + H₂O + CO₂.
- Oxidation of a metal by an oxidising acid — for example copper with concentrated nitric acid, where no hydrogen is released.
Choosing the right method is a standard exam question: if the salt is insoluble use precipitation; if it is soluble and the base is soluble use titration; if the salt is soluble but the base is insoluble use the excess-solid-and-filter method.
Water of crystallisation — three words to keep straight
| Term | Meaning | Example |
|---|---|---|
| Efflorescence | A hydrated salt loses its water of crystallisation to the air and becomes powdery | Na₂CO₃·10H₂O |
| Deliquescence | A solid absorbs so much moisture that it dissolves in it, forming a solution | CaCl₂, NaOH, MgCl₂ |
| Hygroscopy | A substance absorbs moisture but does not dissolve in it | Concentrated H₂SO₄, quicklime CaO |
Common mistakes that cost marks
- Saying "strong" when you mean "concentrated". Strength is about degree of ionisation; concentration is about how much water is present.
- Forgetting that acids need water. Dry HCl gas has no free H⁺ ions and shows no acidic behaviour.
- Writing pH as a unit-carrying quantity. pH is a pure number — it has no unit.
- Assuming every salt solution is neutral. Only a strong-acid + strong-base salt is neutral; NH₄Cl solution is acidic and Na₂CO₃ solution is alkaline.
- Calling every salt with hydrogen an acid salt. NaHSO₄ is an acid salt; NaOH-derived salts such as CH₃COONa contain hydrogen in the anion but are normal salts, because all the replaceable acidic hydrogen has been replaced.
- Using titration for an insoluble salt. You cannot titrate to PbSO₄ — it precipitates immediately, so the precipitation method is the right one.
Practise pH calculations until they are instant. The free pH / pOH calculator converts between [H⁺], [OH⁻], pH and pOH in either direction — enter 0.001 M NaOH and see pH 11 appear, or enter pH 5.6 and get [H⁺] back, exactly as in the worked examples above.
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