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.
FAQs for CBSE Class 9 Science Notes Acids and Bases Chapter
Acids are substances that produce hydrogen ions in solution, while bases produce hydroxide ions. Acids usually taste sour, whereas bases are bitter and slippery. The chapter introduces their properties, examples, and uses in daily life.
Indicators are substances that change color in acidic or basic solutions. Common indicators include litmus, turmeric, and phenolphthalein. They help identify whether a substance is acidic, basic, or neutral, making them useful in laboratory experiments.
A neutralization reaction occurs when an acid reacts with a base to form salt and water. This reaction reduces the acidic and basic properties of the substances involved. The chapter explains its practical applications in medicine, agriculture, and industry.