Acids, Bases & Salts

🔹 1. Definition of Acid

• An acid is a substance that produces H⁺ ions (protons) in aqueous solution.
• According to the Arrhenius definition: Acid increases H⁺ in water.
• According to the Brønsted-Lowry definition: Acid is a proton donor.

🔸 Examples:

• Hydrochloric acid (HCl)
• Sulfuric acid (H₂SO₄)
• Nitric acid (HNO₃)
• Ethanoic acid (CH₃COOH)

🔹 2. Properties of Acids

• Sour taste (e.g., lemon – citric acid)
• Turns blue litmus red
• Conducts electricity in solution (due to mobile H⁺ ions)
• Corrosive in concentrated form
• Reacts with metals, bases, and carbonates

🔹 3. Indicators in Acids

🔹 4. Reactions of Acids

1. Acid + Metal → Salt + Hydrogen
   • E.g., 2HCl + Zn → ZnCl₂ + H₂↑
2. Acid + Base (alkali) → Salt + Water(Neutralization)
   • E.g., HCl + NaOH → NaCl + H₂O
3. Acid + Carbonate → Salt + Water + Carbon Dioxide
   • E.g., 2HCl + CaCO₃ → CaCl₂ + H₂O + CO₂↑

🔹 5. Types of Acids

• Strong Acids: Fully ionize in water
  • E.g., HCl, HNO₃, H₂SO₄
• Weak Acids: Partially ionize
  • E.g., CH₃COOH (ethanoic acid), citric acid

🔹 6. pH of Acids

• pH < 7 (The lower the pH, the stronger the acid)
• Measured using pH paper, universal indicator, or pH meter
• Strong acid: pH 1–3 | Weak acid: pH 4–6

🔹 7. Definition of Base

• A base is a substance that reacts with an acid to form salt and water.
• Alkali is a soluble base that produces OH⁻ ions in aqueous solution.
• Brønsted-Lowry: Base is a proton acceptor.

Examples:

• Sodium hydroxide (NaOH)
• Ammonia (NH₃)
• Copper(II) oxide (CuO)

🔹 8. Properties of Bases

• Bitter taste
• Soapy to touch
• Turns red litmus blue
• Conducts electricity in solution (for alkalis)
• Reacts with acids in neutralization

🔹 9. Indicators of Bases

🔹 10. Reactions of Bases

1. Base + Acid → Salt + Water (Neutralization)
   • E.g., NaOH + HNO₃ → NaNO₃ + H₂O
2. Base + Ammonium Salt → Salt + Water + Ammonia (on heating)
   • E.g., NH₄Cl + NaOH → NaCl + H₂O + NH₃

🔹 11. Types of Bases

• Strong Bases: Completely dissociate in water (e.g., NaOH, KOH)
• Weak Bases: Partially dissociate (e.g., NH₃)
• Metal Oxides and Hydroxides as insoluble bases (e.g., CuO)

🔹 12. Neutralization Reaction (with Ionic Equation)

Example:
HCl (aq) + NaOH (aq) → NaCl (aq) + H₂O (l)

Ionic Equation:

H⁺ (aq) + OH⁻ (aq) → H₂O (l)

This shows that neutralization involves hydrogen ions combining with hydroxide ions to form water.

🔹 13. Salts – Definitions and Examples

• A salt is a compound formed when the hydrogen of an acid is fully or partially replaced by a metal or ammonium ion.

Types of Salts:

• Normal salt: All H⁺ replaced (e.g., NaCl)
• Acid salt: Some H⁺ remains (e.g., NaHSO₄)
• Double salt: Two different cations (e.g., alum – KAl(SO₄)₂·12H₂O)

Examples:

🔹 14. Solubility of Salts

Soluble Salts

• All sodium, potassium, and ammonium salts
• All nitrates
• Most chlorides, bromides, iodides (except Ag⁺ and Pb²⁺)
• Most sulfates (except Ba²⁺, Pb²⁺, Ca²⁺ slightly)

Insoluble Salts

• Most carbonates (except Group 1 and NH₄⁺)
• Most hydroxides (except Group 1 and NH₄⁺, Ca²⁺ slightly soluble)
• Silver chloride, Lead(II) sulfate

🔹 15. Preparation of Salts

A. Preparation of Soluble Salts

Method depends on solubility of base/metal:

(i) Acid + Excess Base (Insoluble Base Method)

• Used when base is insoluble (e.g., CuO, ZnCO₃)
• E.g., CuO + H₂SO₄ → CuSO₄ + H₂O
• Steps:
  1. Add base to warm acid until no more reacts
  2. Filter excess base
  3. Heat to concentrate
  4. Crystallize

(ii) Titration (Acid + Alkali)

• When both acid and base are soluble (e.g., NaOH + HCl)
• Steps:
  1. Use indicator to find exact volume needed
  2. Repeat without indicator
  3. Evaporate and crystallize

B. Preparation of Insoluble Salts

By Precipitation Reaction

E.g., AgNO₃ (aq) + NaCl (aq) → AgCl (s) + NaNO₃ (aq)

Steps:

1. Mix two soluble salt solutions
2. Filter the precipitate
3. Wash and dry

✅ Summary Table

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