Science and Technology Class 9 Notes Chapter 5 Maharashtra Board

Introduction to Acids, Bases, and Salts

What Are Acids, Bases, and Salts?

Acids: Sour taste (e.g., lemon, vinegar).

Bases: Bitter taste, soapy feel (e.g., baking soda, milk of magnesia).

Salts: Formed when acids react with bases.

Litmus Test:

• Acids turn blue litmus red (e.g., lemon, tamarind, vinegar, buttermilk).
• Bases turn red litmus blue (e.g., baking soda, milk of magnesia, lime).
• Neutral: No change (e.g., water, milk).

Ionic Compounds Recap:

Ionic compounds have two parts:

• Cation (positive ion, basic radical): e.g., Na⁺, Ca²⁺.
• Anion (negative ion, acidic radical): e.g., Cl⁻, SO₄²⁻.

Types:

• Acids: Have H⁺ as cation (e.g., HCl → H⁺ + Cl⁻).
• Bases: Have OH⁻ as anion (e.g., NaOH → Na⁺ + OH⁻).
• Salts: Have other cations and anions (e.g., NaCl → Na⁺ + Cl⁻).

Arrhenius Theory of Acids and Bases

Proposed by Svante Arrhenius (1887):

Acid: Produces H⁺ ions in water.

• Example: HCl → H⁺ + Cl⁻; H₂SO₄ → H⁺ + HSO₄⁻ → 2H⁺ + SO₄²⁻.

Base: Produces OH⁻ ions in water.

• Example: NaOH → Na⁺ + OH⁻; Ca(OH)₂ → Ca²⁺ + 2OH⁻.

Classification of Acids and Bases:

• Strong Acid: Fully dissociates in water (e.g., HCl, HNO₃ → mostly H⁺ and anions).
• Weak Acid: Partially dissociates (e.g., CH₃COOH → some H⁺, mostly undissociated).
• Strong Base: Fully dissociates (e.g., NaOH, KOH → mostly OH⁻ and cations).
• Weak Base: Partially dissociates (e.g., NH₃ → some OH⁻, mostly undissociated).
• Alkali: Bases highly soluble in water (e.g., NaOH, KOH – strong; NH₃ – weak).

Basicity and Acidity:

Basicity of Acid: Number of H⁺ ions produced per molecule.

• Monobasic: 1 H⁺ (e.g., HCl).
• Dibasic: 2 H⁺ (e.g., H₂SO₄).
• Tribasic: 3 H⁺ (e.g., H₃PO₄).

Acidity of Base: Number of OH⁻ ions produced per molecule.

• Monoacidic: 1 OH⁻ (e.g., NaOH).
• Diacidic: 2 OH⁻ (e.g., Ca(OH)₂).
• Triacidic: 3 OH⁻ (e.g., Al(OH)₃).

Concentration of Acids and Bases

What is Concentration?

• The amount of solute (acid/base) in a solution.
• Concentrated: High amount of solute (more sour/bitter).
• Dilute: Low amount of solute (less sour/bitter).

Activity (Lemon Juice):

• Take juice from half a lemon in two beakers.
• Beaker A: Add 10 ml water. Beaker B: Add 20 ml water.
• Result: Beaker A tastes more sour (higher concentration of lemon juice).
• Conclusion: Concentration affects the intensity of properties (taste, reactivity).

Units of Concentration:

• Grams per Liter (g/L): Mass of solute in grams in 1 liter of solution.
• Molarity (M): Moles of solute in 1 liter of solution.
• Example: [NaCl] = 1 M means 1 mole of NaCl in 1 liter.

Example: 1.5 mol of NaOH in 2 L →

• Molarity = 1.5 / 2 = 0.75 M.
• Mass = 1.5 × 40 = 60 g → Concentration = 60 / 2 = 30 g/L.

pH of a Solution

What is pH?

• A scale (0 to 14) to measure H⁺ ion concentration, introduced by Sorensen (1909).
• pH 7: Neutral (e.g., pure water, [H⁺] = 10⁻⁷ mol/L).
• pH < 7: Acidic (e.g., HCl, [H⁺] = 10⁰ mol/L → pH 0).
• pH > 7: Basic (e.g., NaOH, [H⁺] = 10⁻¹⁴ mol/L → pH 14).

pH of Common Solutions:

• Lemon Juice: 2.2 (acidic).
• Vinegar: 3 (acidic).
• Pure Water: 7 (neutral).
• Milk of Magnesia: 10 (basic).
• NaOH (1M): 14 (basic).

Measuring pH:

Universal Indicator: Shows different colors for different pH values.

• Acidic (pH 0-6): Red to yellow.
• Neutral (pH 7): Green.
• Basic (pH 8-14): Blue to purple.

pH Meter: Most accurate; uses electrodes to measure pH.

Other Indicators:

• Litmus: Red in acid, blue in base.
• Phenolphthalein: Colorless in acid, pink in base.
• Methyl Orange: Red in acid, yellow in base.

Reactions of Acids and Bases

1. Neutralization Reaction:

• Acid + Base → Salt + Water.
• Example: HCl + NaOH → NaCl + H₂O.
• Ionic Reaction: H⁺ + OH⁻ → H₂O.

Activity:

• • Add NaOH drop by drop to 10 ml HCl.
  • Check pH with pH paper after each addition.
  • pH increases from <7 to 7 (green color, neutral).
  • Reason: H⁺ from HCl reacts with OH⁻ from NaOH to form water.

Examples:

• • H₂SO₄ + 2KOH → K₂SO₄ + 2H₂O.
  • HNO₃ + KOH → KNO₃ + H₂O.

2. Reaction of Acids with Metals:

• Acid + Metal → Salt + Hydrogen Gas.
• Example: Mg + 2HCl → MgCl₂ + H₂ (hydrogen gas burns with a pop sound).
• Activity:
  • Add dilute HCl to Mg ribbon in a test tube.
  • Hydrogen gas is released, confirmed by a burning candle (pop sound).

Examples:

• • Zn + H₂SO₄ → ZnSO₄ + H₂.
  • Cu + 2HNO₃ → Cu(NO₃)₂ + H₂.

3. Reaction of Acids with Metal Oxides:

• Metal Oxide + Acid → Salt + Water.
• Example: Fe₂O₃ + 6HCl → 2FeCl₃ + 3H₂O (red oxide dissolves, solution turns yellow).
• Metal oxides are basic (neutralize acids).

Examples:

• • CaO + 2HCl → CaCl₂ + H₂O.
  • MgO + 2HCl → MgCl₂ + H₂O.

4. Reaction of Bases with Non-Metal Oxides:

• Non-Metal Oxide + Base → Salt + Water.
• Example: CO₂ + 2NaOH → Na₂CO₃ + H₂O.
• Non-metal oxides are acidic (neutralize bases).

Examples:

• • CO₂ + 2KOH → K₂CO₃ + H₂O.
  • SO₃ + 2NaOH → Na₂SO₄ + H₂O.
• Amphoteric Oxides: React with both acids and bases (e.g., ZnO, Al₂O₃).
  • ZnO + 2NaOH → Na₂ZnO₂ + H₂O.
  • Al₂O₃ + 2NaOH → 2NaAlO₂ + H₂O.

5. Reaction of Acids with Carbonates/Bicarbonates:

• Acid + Carbonate/Bicarbonate → Salt + Water + CO₂.
• Example: NaHCO₃ + HCl → NaCl + H₂O + CO₂ (CO₂ turns limewater milky).

Activity:

• • Add lemon juice (acid) to baking soda (NaHCO₃).
  • CO₂ gas is released, turns limewater [Ca(OH)₂] milky: Ca(OH)₂ + CO₂ → CaCO₃ + H₂O.

Examples:

• • Na₂CO₃ + 2HCl → 2NaCl + H₂O + CO₂.
  • CaCO₃ + 2HCl → CaCl₂ + H₂O + CO₂.

Salts

Types of Salts (Based on pH):

• Neutral Salt: pH = 7 (strong acid + strong base, e.g., NaCl, KCl).
• Acidic Salt: pH < 7 (strong acid + weak base, e.g., NH₄NO₃).
• Basic Salt: pH > 7 (weak acid + strong base, e.g., Na₂CO₃, NaHCO₃).

Activity:

• Test pH of NaCl (neutral, pH 7), NH₄Cl (acidic, pH < 7), NaHCO₃ (basic, pH > 7).

Classification:

• Neutral: NaCl, KCl, Na₂SO₄.
• Acidic: NH₄NO₃.
• Basic: Na₂CO₃, NaHCO₃, NaCH₃COO.

Water of Crystallization:

• Water molecules in the crystal structure of some salts.

Examples:

• • Blue Vitriol: CuSO₄·5H₂O (blue).
  • Green Vitriol: FeSO₄·7H₂O.
  • Washing Soda: Na₂CO₃·10H₂O.
  • Alum: K₂SO₄·Al₂(SO₄)₃·24H₂O.

Activity:

• • Heat blue vitriol (CuSO₄·5H₂O) → turns white (CuSO₄) + water vapor.
  • Add water to white powder → turns blue again (reversible physical change).
  • Equation: CuSO₄·5H₂O → CuSO₄ + 5H₂O.

Similar for FeSO₄·7H₂O and Na₂CO₃·10H₂O.

Ionic Compounds and Electrical Conductivity

Conductivity Test:

• Activity: Test solutions of NaCl, CuSO₄, H₂SO₄, NaOH, glucose, urea.
• Result: NaCl, CuSO₄, H₂SO₄, NaOH conduct electricity (bulb glows); glucose, urea do not.
• Reason: Conductors have dissociated ions (cations to cathode, anions to anode).
• Electrodes:
  • Cathode: Negative electrode (connected to battery’s negative terminal).
  • Anode: Positive electrode (connected to battery’s positive terminal).

Electrolysis:

Electrolysis of CuSO₄:

• Cathode: Cu²⁺ + 2e⁻ → Cu (copper deposits).
• Anode: Cu → Cu²⁺ + 2e⁻ (copper dissolves).

Electrolysis of Water (with salt):

• Add salt to water to make it conductive.
• Cathode: 2H₂O + 2e⁻ → H₂ + 2OH⁻ (hydrogen gas, volume double).
• Anode: 2H₂O → O₂ + 4H⁺ + 4e⁻ (oxygen gas).
• Test: H₂ at cathode (neutral), OH⁻ in solution (basic); H⁺ near anode (acidic).

Electrolytes:

• Strong Electrolytes: High conductivity (e.g., NaCl, H₂SO₄, NaOH).
• Weak Electrolytes: Low conductivity (e.g., CH₃COOH, NH₃).
• Pure Water: Bad conductor (low H⁺/OH⁻, 10⁻⁷ mol/L).