Chapter in a nutshell: Acids give H⁺ (turn blue litmus red), bases/alkalis give OH⁻ (turn red litmus blue); the pH scale measures acidity (0–14). They neutralise to form salts + water. Salts are normal, acid or basic, prepared by several standard methods, and may carry water of crystallisation — leading to efflorescence, deliquescence and hygroscopy.
1. Acids
An acid is a substance that produces hydrogen ions (H⁺ / H₃O⁺) in aqueous solution. Properties: sour taste; turn blue litmus red; turn methyl orange red, phenolphthalein colourless; conduct electricity (electrolytes); corrosive. Chemical reactions:- With active metals → salt + hydrogen: $\mathrm{Zn + 2HCl \rightarrow ZnCl_2 + H_2}$
- With carbonates/bicarbonates → salt + water + CO₂: $\mathrm{Na_2CO_3 + 2HCl \rightarrow 2NaCl + H_2O + CO_2}$
- With bases (neutralisation) → salt + water: $\mathrm{HCl + NaOH \rightarrow NaCl + H_2O}$
Classification of acids:
| Basis | Types |
|---|---|
| Source | mineral (HCl, H₂SO₄, HNO₃) / organic (acetic, citric) |
| Strength | strong (fully ionised: HCl, H₂SO₄, HNO₃) / weak (partly ionised: CH₃COOH, H₂CO₃) |
| Concentration | concentrated / dilute |
| Basicity | number of replaceable H⁺: mono- (HCl), di- (H₂SO₄), tri-basic (H₃PO₄) |
2. Bases and Alkalis
A base reacts with an acid to give salt + water; an alkali is a water-soluble base that gives OH⁻ ions. Properties: bitter taste, soapy feel; turn red litmus blue; methyl orange yellow, phenolphthalein pink.- Strong alkalis: NaOH, KOH (fully ionised). Weak base: NH₄OH.
- Acidity of a base = number of OH⁻ it provides (NaOH mono-acidic, Ca(OH)₂ di-acidic).
- Amphoteric oxides/hydroxides (Al₂O₃, ZnO, PbO) react with both acids and alkalis.
3. Indicators & the pH Scale
| Indicator | In acid | In base |
|---|---|---|
| Litmus | red | blue |
| Methyl orange | red/pink | yellow |
| Phenolphthalein | colourless | pink |
| Turmeric | yellow | red/brown |
4. Neutralisation
Acid + Base → Salt + Water (an exothermic reaction): $$\mathrm{HCl + NaOH \rightarrow NaCl + H_2O} \qquad \mathrm{H^+ + OH^- \rightarrow H_2O}$$ Uses: treating acidity (antacids), soil treatment (lime), insect-sting relief.5. Salts — Types
A salt is formed when the replaceable H⁺ of an acid is replaced by a metal/ammonium ion.| Type | Definition | Example |
|---|---|---|
| Normal salt | all replaceable H⁺ replaced | NaCl, Na₂SO₄, K₂CO₃ |
| Acid salt | only part of the H⁺ replaced | NaHCO₃, NaHSO₄ |
| Basic salt | contains OH⁻ along with the salt | Pb(OH)Cl, basic copper carbonate |
6. Methods of Preparing Salts
| Method | Reaction | Use |
|---|---|---|
| Direct combination | $\mathrm{Fe + S \rightarrow FeS}$ | element + element |
| Acid + metal | $\mathrm{Zn + H_2SO_4 \rightarrow ZnSO_4 + H_2}$ | active metals |
| Acid + base (neutralisation) | $\mathrm{HCl + NaOH \rightarrow NaCl + H_2O}$ | soluble salts (titration) |
| Acid + carbonate/oxide | $\mathrm{CuO + H_2SO_4 \rightarrow CuSO_4 + H_2O}$ | insoluble base + acid |
| Precipitation (double decomposition) | $\mathrm{AgNO_3 + NaCl \rightarrow AgCl{\downarrow} + NaNO_3}$ | insoluble salts |
- Insoluble salts (AgCl, BaSO₄, PbI₂, most carbonates) → made by precipitation.
7. Water of Crystallisation
The fixed number of water molecules chemically combined in one formula unit of a crystalline salt.| Hydrated salt | Formula | Water molecules |
|---|---|---|
| Blue vitriol (copper sulphate) | CuSO₄·5H₂O | 5 |
| Washing soda | Na₂CO₃·10H₂O | 10 |
| Gypsum | CaSO₄·2H₂O | 2 |
| Epsom salt | MgSO₄·7H₂O | 7 |
8. Efflorescence, Deliquescence, Hygroscopy
| Term | Meaning | Example |
|---|---|---|
| Efflorescence | a hydrated salt loses water to air, becoming powdery | Na₂CO₃·10H₂O |
| Deliquescence | a substance absorbs moisture and dissolves in it | NaOH, CaCl₂, MgCl₂ |
| Hygroscopy | absorbs moisture but does not dissolve | conc. H₂SO₄, CaO, silica gel |
9. Worked / Structured Examples (ICSE pattern)
Q1. Name the gas evolved when (i) zinc reacts with dilute H₂SO₄, (ii) sodium carbonate reacts with dilute HCl. Solution: (i) hydrogen, (ii) carbon dioxide.Q2. Classify: (i) NaHSO₄, (ii) Na₂SO₄, (iii) Pb(OH)Cl. Solution: (i) acid salt, (ii) normal salt, (iii) basic salt.
Q3. A salt is insoluble in water — which method prepares it? Give an example. Solution: Precipitation (double decomposition), e.g. $\mathrm{BaCl_2 + Na_2SO_4 \rightarrow BaSO_4{\downarrow} + 2NaCl}$.
Q4. What happens when blue copper sulphate crystals are heated? Solution: They lose water of crystallisation and turn into white anhydrous CuSO₄; adding water restores the blue colour.
Q5. State the pH nature of (i) lemon juice, (ii) NaOH solution, (iii) pure water. Solution: (i) acidic (pH < 7), (ii) alkaline (pH > 7), (iii) neutral (pH = 7).
10. Key Terms — Quick Glossary
| Term | One-line definition |
|---|---|
| Acid | gives H⁺ ions in water; turns blue litmus red. |
| Base | reacts with acid to give salt + water. |
| Alkali | water-soluble base giving OH⁻ ions. |
| Basicity | number of replaceable H⁺ in an acid. |
| pH | measure of acidity (0–14). |
| Neutralisation | acid + base → salt + water. |
| Normal/acid/basic salt | all / part / OH-containing salt. |
| Water of crystallisation | fixed water molecules in a crystal. |
| Deliquescence | absorbs moisture and dissolves. |
| Efflorescence | loses water of crystallisation to air. |
11. Common Mistakes to Avoid
- Saying all bases are alkalis — only water-soluble bases are alkalis.
- Confusing strength (degree of ionisation) with concentration (amount in water).
- Mixing up basicity (acids) with acidity (bases).
- Confusing deliquescence (dissolves) with hygroscopy (doesn't dissolve) and efflorescence (loses water).
- Choosing neutralisation for an insoluble salt — those need precipitation.
12. Likely Exam Questions (with crisp answers)
- Define an acid in terms of ions. → A substance that gives H⁺ (H₃O⁺) ions in aqueous solution.
- What is an alkali? → A water-soluble base that provides OH⁻ ions.
- Define basicity of an acid; give a dibasic acid. → Number of replaceable H⁺; H₂SO₄.
- State the colour change of methyl orange in acid and base. → Red/pink in acid, yellow in base.
- What is the pH of a neutral solution? → 7.
- Name the gas evolved when an acid reacts with a metal carbonate. → Carbon dioxide.
- Distinguish a normal salt from an acid salt with examples. → Normal: all H⁺ replaced (NaCl); acid: part replaced (NaHCO₃).
- Which method prepares an insoluble salt? → Precipitation (double decomposition).
- Define water of crystallisation. → Fixed water molecules chemically combined in a crystalline salt.
- Define deliquescence with an example. → Absorbing moisture and dissolving in it; NaOH.
- Define efflorescence with an example. → Losing water of crystallisation to air; washing soda.
- Name a hygroscopic drying agent. → Concentrated sulphuric acid (or CaO).
- What is neutralisation? Write its ionic equation. → Acid + base → salt + water; $\mathrm{H^+ + OH^- \rightarrow H_2O}$.
- Why does anhydrous CuSO₄ test for water? → It turns from white to blue in the presence of water.
13. Action of Acids — more reactions
- Acid + metal oxide (basic oxide) → salt + water: $\mathrm{CuO + 2HCl \rightarrow CuCl_2 + H_2O}$.
- Acid + hydroxide → salt + water: $\mathrm{2NaOH + H_2SO_4 \rightarrow Na_2SO_4 + 2H_2O}$.
- Acid + sulphite → salt + water + SO₂; acid + sulphide → salt + H₂S.
- Dilute vs concentrated: dilute acids supply H⁺ for typical reactions; concentrated H₂SO₄ also acts as a dehydrating/oxidising agent.
14. Solubility Rules (for choosing the prep method)
| Soluble | Insoluble |
|---|---|
| All nitrates | — |
| All Na⁺, K⁺, NH₄⁺ salts | — |
| Most chlorides | AgCl, PbCl₂ |
| Most sulphates | BaSO₄, PbSO₄, CaSO₄(slightly) |
| Na/K/NH₄ carbonates | most other carbonates |
15. Action of Heat on Some Salts
- Hydrated salts lose water: $\mathrm{CuSO_4{\cdot}5H_2O \xrightarrow{\Delta} CuSO_4 + 5H_2O}$ (blue → white).
- Carbonates (except Na/K) → oxide + CO₂: $\mathrm{CaCO_3 \xrightarrow{\Delta} CaO + CO_2}$.
- Nitrates: $\mathrm{2Pb(NO_3)_2 \xrightarrow{\Delta} 2PbO + 4NO_2 + O_2}$ (brown fumes of NO₂).
- Ammonium salts sublime/decompose: $\mathrm{NH_4Cl \xrightarrow{\Delta} NH_3 + HCl}$.
16. Simple Identification Tests
- Hydrogen (H₂): burns with a 'pop' sound with a lighted splint.
- Carbon dioxide (CO₂): turns lime water milky.
- Ammonia (NH₃): turns moist red litmus blue; dense white fumes with HCl.
- Sulphur dioxide (SO₂): turns acidified potassium dichromate paper green.
17. More Worked Examples
Q6. Write the equation for the action of dilute sulphuric acid on copper oxide. Solution: $\mathrm{CuO + H_2SO_4 \rightarrow CuSO_4 + H_2O}$. Q7. How would you prepare lead sulphate (insoluble)? Solution: Precipitation: $\mathrm{Pb(NO_3)_2 + Na_2SO_4 \rightarrow PbSO_4{\downarrow} + 2NaNO_3}$. Q8. Name the brown gas evolved when lead nitrate is heated. Solution: Nitrogen dioxide (NO₂). Q9. A solution turns phenolphthalein pink. Acidic or basic? Solution: Basic (alkaline).18. More Exam Questions (with crisp answers)
- Why is concentrated sulphuric acid stored in airtight bottles? → It is hygroscopic (absorbs moisture from air).
- Give the test for carbon dioxide. → It turns lime water milky.
- Name a salt that effloresces and one that deliquesces. → Washing soda (effloresces); NaOH/CaCl₂ (deliquesces).
- Why is a salt of a strong acid and weak base acidic? → On hydrolysis it releases excess H⁺.
- Write the action of heat on calcium carbonate. → $\mathrm{CaCO_3 \rightarrow CaO + CO_2}$.
- Name the salt and gas formed when zinc reacts with dilute HCl. → Zinc chloride and hydrogen.
- Define a drying (desiccating) agent; give one. → A substance that absorbs moisture; conc. H₂SO₄ / silica gel.
- State one industrial use of neutralisation. → Treating acidic soil with lime (or antacids for acidity).