Acids and Bases
Contents
- Theories of Acids and Bases
- Acids – Properties & Classification
- Bases – Properties & Classification
- Role of Water in Ionisation
- Neutralisation & pH Scale
- Indicators
- Salts – Formation & Types
- Industrially Important Salts
- Efflorescence, Hygroscopy & Deliquescence
Check Out- CBSE Class 9 Science Notes
Theories of Acids and Bases
Arrhenius Concept (Water Ion System)
An acid gives H⁺ ions in aqueous solution; a base gives OH⁻ ions. Neutralisation: H⁺ + OH⁻ → H₂O. H⁺ ions actually exist as H₃O⁺ (hydronium ion): H⁺ + H₂O → H₃O⁺.
Acids: HF, HCl, HBr, HI, H₂SO₄, HNO₃ | Bases: NaOH, KOH, Mg(OH)₂, Ca(OH)₂, Al(OH)₃
Brønsted–Lowry Theory (Proton Donor–Acceptor)
An acid donates a proton (H⁺); a base accepts a proton. Species that can act as both are called amphiprotic (e.g., H₂O, NH₃).
Pairs related by gain/loss of a proton are conjugate acid–base pairs. When HCl donates H⁺ to H₂O, HCl is the acid and H₂O the base; Cl⁻ is the conjugate base of HCl and H₃O⁺ is the conjugate acid of H₂O.
Lewis Concept (Electron Donor–Acceptor)
An acid accepts an electron pair to form a coordinate covalent bond; a base donates an electron pair. Neutralisation = formation of a coordinate bond (adduct).
Lewis Acids: H⁺, Cu²⁺, Fe²⁺/Fe³⁺, BF₃, AlF₃ | Lewis Bases: OH⁻, CN⁻, CH₃COO⁻, :NH₃, H₂O:
Limitations: Cannot explain relative strengths of acids/bases; neutralisation (forming salts + water) not explained; coordinate bond formation is slow, while acid–base reactions are fast.
Acids
Substances with sour taste that release one or more H⁺ ions in aqueous solution.
Types by Source
- Mineral acids: Obtained from rocks/minerals — HCl, H₂SO₄, HNO₃.
- Organic acids: Present in animals and plants — formic acid (HCOOH), acetic acid (CH₃COOH).
Methods of Preparation
- Direct combination: H₂ + Cl₂ → 2HCl
- Non-metallic oxide + water: SO₃ + H₂O → H₂SO₄
- Oxidation by oxy-acids: S + 6HNO₃ → H₂SO₄ + 2H₂O + 6NO₂
- Displacement from volatile acid salts: NaCl + H₂SO₄ → NaHSO₄ + HCl↑
Chemical Properties
Action with metals
Action with metal oxides
Action with carbonates / bicarbonates
Action with bases (neutralisation)
Classification of Acids
| Basis | Type | Description / Example |
|---|---|---|
| Strength | Strong acid | ≈100% ionisation — HCl, HNO₃, H₂SO₄ |
| Weak acid | Partial ionisation — CH₃COOH, H₂CO₃ | |
| Basicity | Monobasic | 1 H⁺ per molecule — HCl |
| Dibasic | 2 H⁺ per molecule — H₂SO₄ | |
| Tribasic | 3 H⁺ per molecule — H₃PO₄ | |
| Concentration | Concentrated | High % of acid, low % of water |
| Dilute | Low % of acid, high % of water |
Bases
Substances with bitter taste and soapy touch that release OH⁻ ions in aqueous solution. Water-soluble bases are called alkalis (e.g., NaOH, KOH). Insoluble bases include Cu(OH)₂, Fe(OH)₃, Al(OH)₃.
Methods of Preparation
- Water + metal: 2Na + 2H₂O → 2NaOH + H₂
- Metallic oxide + water: K₂O + H₂O → 2KOH
- Oxygen + metal: 2Mg + O₂ → 2MgO
- Decomposition of carbonates: ZnCO₃ → ZnO + CO₂
- Metal salt + NaOH: AlCl₃ + 3NaOH → Al(OH)₃ + 3NaCl
Chemical Properties
Action with metals (amphoteric metals)
Action with non-metallic oxides
Classification of Bases
| Basis | Type | Example |
|---|---|---|
| Strength | Strong base (100% ionised) | NaOH, KOH |
| Weak base (partial ionisation) | NH₄OH | |
| Acidity | Monoacidic (1 OH⁻) | NaOH, KOH, NH₄OH |
| Diacidic (2 OH⁻) | Ca(OH)₂, Mg(OH)₂, Cu(OH)₂ | |
| Triacidic (3 OH⁻) | Al(OH)₃, Fe(OH)₃ |
Role of Water in Ionisation
Acids and bases show their character only in the presence of water (aqueous solution). In dry (anhydrous) state, acidic/basic characters cannot be shown.
HCl gas (hydrogen chloride) cannot give H⁺ ions in dry state. In water, the polar H₂O molecules form an envelope around H and Cl atoms, separating them as hydrated ions:
These ions carry electric current. Similarly, bases release OH⁻ ions only in aqueous solution.
Neutralisation & pH Scale
Neutralisation: Reaction of acid + base → salt + water + heat.
Heat of neutralisation: Heat liberated when 1 equivalent of acid reacts with 1 equivalent of base. For strong acid–strong base reactions: constant value of 13.7 kcal/mol (both fully ionised).
pH Scale
pH = −log₁₀[H⁺]. Ranges from 0 to 14.
pH 0 → 6
Neutral
More Basic →
- pH < 7 → Acidic (lower = more acidic)
- pH = 7 → Neutral
- pH > 7 → Alkaline (higher = more alkaline)
A universal indicator shows different colours at different H⁺ concentrations: red/yellow in acid, green at pH 7, blue/violet in base.
Indicators
| Indicator | In Acid | Neutral | In Base |
|---|---|---|---|
| Litmus (from lichen) | Red | Purple | Blue |
| Phenolphthalein | Colourless | Colourless | Pink |
| Methyl orange | Red | Orange/Yellow | Orange/Yellow |
| Red cabbage juice | Red/Pink | Purple | Green |
Salts
Ionic compounds with a positive (cation, usually metal) and negative (anion, acid radical) part, electrically neutral overall.
Formation Methods
- Neutralisation of acid + base: NaOH + HCl → NaCl + H₂O
- Metal + acid: Zn + H₂SO₄ → ZnSO₄ + H₂
- Acid + carbonate: CaCO₃ + 2HCl → CaCl₂ + H₂O + CO₂
- Metal + alkali: 2NaOH + Zn → Na₂ZnO₂ + H₂
Families of Salts
- Chlorides — from HCl
- Nitrates — from HNO₃
- Sulphates — from H₂SO₄
- Carbonates — from H₂CO₃
Normal, Acidic & Basic Salts
- Normal salt: All ionisable H⁺ of polybasic acid replaced by metal ions — Na₂SO₄, K₂CO₃.
- Acidic salt: Some H⁺ remain — NaHSO₄ (sodium bisulphate), NaHCO₃ (sodium bicarbonate).
- Basic salt: Base not fully neutralised — Mg(OH)Cl (basic magnesium chloride).
Industrially Important Salts
1. Common Salt (NaCl)
Occurrence: Sea water (evaporation), rock salt, inland lakes.
- Colourless crystals; hygroscopic; reacts with conc. H₂SO₄ to give HCl gas.
- Uses: Food preservation, freezing mixture, raw material for NaHCO₃, Na₂CO₃, NaOH, Cl₂, bleaching powder.
2. Caustic Soda (NaOH)
Manufacture (Chlor-alkali process): Electrolysis of brine (aqueous NaCl):
Uses: Soap making, degreasing metals, paper/dye/rayon industry, petroleum refining, laboratory reagent.
3. Bleaching Powder (CaOCl₂)
Yellowish-white powder. Reacts with CO₂ (air) and HCl/H₂SO₄ to release Cl₂. Uses: Bleaching (textiles, paper), disinfecting water, oxidising agent, manufacture of chloroform.
4. Baking Soda (NaHCO₃)
White crystalline; alkaline solution; decomposes on heating to Na₂CO₃ + CO₂ + H₂O. Uses: Antacid, baking powder (for fluffy cakes/bread), fire extinguishers.
5. Washing Soda (Na₂CO₃·10H₂O)
Manufacture (Solvay process):
- NaCl + NH₃ + CO₂ + H₂O → NaHCO₃ + NH₄Cl
- Heat: 2NaHCO₃ → Na₂CO₃ + H₂O + CO₂ (soda ash)
- Na₂CO₃ + 10H₂O → Na₂CO₃·10H₂O (washing soda)
Uses: Laundry, removing permanent hardness of water, glass/soap/paper manufacture.
6. Plaster of Paris (CaSO₄·½H₂O)
White powder that sets hard with water (rehydration to gypsum). Uses: Setting fractured bones, making toys and statues, decorative designs, chalk, sealing lab apparatus.
Efflorescence, Hygroscopy & Deliquescence
Efflorescence
Loss of water of crystallisation by crystals exposed to dry air, turning them to powder. Examples: Na₂CO₃·10H₂O → Na₂CO₃·H₂O + 9H₂O; Na₂SO₄·10H₂O → Na₂SO₄ + 10H₂O.
Hygroscopic Substances
Absorb moisture from air without dissolving. Used as drying agents. Examples: Conc. H₂SO₄, P₂O₅, CaO, silica gel.
Deliquescent Substances
Hygroscopic substances that absorb so much moisture they dissolve in it forming a saturated solution. Deliquescence is minimised in dry conditions (where efflorescence is maximised). Examples: NaOH, KOH, MgCl₂, ZnCl₂, CaCl₂, FeCl₃, Zn(NO₃)₂, Cu(NO₃)₂.