1. Introduction to Acids and Bases

Acids

• ★Taste: sour (important: do not taste chemicals in the lab).
• ★Turn blue litmus red.
• ★Examples: HCl (hydrochloric acid), H₂SO₄ (sulphuric acid), HNO₃ (nitric acid), CH₃COOH (acetic acid).

Bases

• ★Taste: bitter (again — do not taste in real life).
• ★Feel soapy or slippery to touch (for soluble bases).
• ★Turn red litmus blue.
• ★Examples: NaOH (sodium hydroxide), Ca(OH)₂ (calcium hydroxide).
• ★Alkalis are bases that dissolve in water (e.g., NaOH).

Indicators

Indicators change colour to tell us whether a solution is acidic or basic.

Natural indicators

• ★Litmus: A purple dye from lichens. Neutral → purple. Acid → red. Base → blue.
• ★Turmeric: Yellow normally; turns reddish-brown in strong bases (like soap).
• ★Other natural indicators: red cabbage, some flower petals (Hydrangea, Petunia, Geranium).

Synthetic indicators

• ★Phenolphthalein: Colourless in acid, pink in base.
• ★Methyl orange: Red in acid, yellow in base.

Olfactory indicators — some smells change in acids/bases (not used much in labs but useful to know).

2. Chemical Properties of Acids and Bases

A. Reaction with Metals

Many acids react with certain metals to produce a salt and hydrogen gas.

• Acid + Metal → Salt + H₂(g)
• Example: Zinc + dilute sulphuric acid gives zinc sulphate and hydrogen. The hydrogen gas produces a characteristic 'pop' test.
• Some strong bases (not all) also react with certain metals; for example: 2NaOH + Zn → Na₂ZnO₂ + H₂ (sodium zincate + hydrogen).
• Note: Not every metal reacts with acids; reactivity depends on the metal.

B. Reaction with Metal Carbonates and Hydrogencarbonates

Acids react with carbonates and hydrogencarbonates to give a salt, carbon dioxide and water.

• Metal carbonate/hydrogencarbonate + Acid → Salt + CO₂(g) + H₂O
• Example: Na₂CO₃ + 2HCl → 2NaCl + H₂O + CO₂↑
• Test for CO₂: Bubble the gas through lime water (Ca(OH)₂). It turns milky due to CaCO₃ precipitate. On further excess CO₂ the solution clears (forms soluble hydrogencarbonate).

C. Neutralisation (Reaction with Each Other)

Acid + Base → Salt + Water. This is called neutralisation. Acids and bases cancel each other's effects.

• Example: HCl + NaOH → NaCl + H₂O
• Ionic view: H⁺(aq) + OH⁻(aq) → H₂O(l)

D. Metallic Oxides with Acids

• ★Metallic oxides (like CuO) react with acids to give salts and water. These oxides are called basic oxides.
• Example: CuO + 2HCl → CuCl₂ + H₂O

E. Non-metallic Oxides with Bases

• ★Non-metal oxides (like CO₂) react with bases to give salts and water. These oxides are acidic oxides.
• Example: CO₂ + Ca(OH)₂ → CaCO₃ + H₂O

3. Acids and Bases in Water

Conduction of Electricity

Acids and bases conduct electricity in solution because they produce ions. Acids produce H⁺ (or H₃O⁺) ions and bases produce OH⁻ ions. Substances like sugar or alcohol do not conduct because they do not form ions in water.

Role of Water

• ★Hydrogen ions rarely exist alone; in water they form hydronium ions: HCl + H₂O → H₃O⁺ + Cl⁻.
• ★Soluble bases ionise: NaOH → Na⁺ + OH⁻ in water.

Dilution means the concentration of ions per unit volume decreases when you mix the solution with water.

4. The pH Scale

pH measures how acidic or basic a solution is. The usual scale runs from 0 to 14.

• ★pH = 7 → Neutral (pure water).
• ★pH < 7 → Acidic (higher H⁺ concentration).
• ★pH > 7 → Basic/alkaline (higher OH⁻ concentration).
• ★Universal indicator is a mixture that shows many colours across the pH range.
• ★Strong acids/bases produce more ions (e.g., HCl). Weak acids/bases produce fewer ions (e.g., acetic acid).

5. Importance of pH in Everyday Life

• ★Human body: Many body processes work best near pH 7 – 7.8 (blood ~7.4).
• ★Acid rain: Rain with pH < 5.6 harms rivers, plants and buildings.
• ★Soil pH: Plants need soil in the right pH range to absorb nutrients.
• ★Digestion: Stomach makes HCl to digest food. Excess acid causes indigestion; antacids (mild bases) help neutralise it.
• ★Tooth decay: When mouth pH < 5.5, acids produced by bacteria start dissolving tooth enamel.
• ★Nature’s defences: Bee stings (acidic) cause pain — mild bases like baking soda can give temporary relief. Nettle stings are caused by formic/methanoic acid — rubbing leaves of certain plants can soothe the pain.

Quick examples of natural acids: vinegar → acetic acid, curd → lactic acid, oranges/lemons → citric acid, tamarind → tartaric acid, tomatoes → oxalic acid.

6. Salts

A salt is formed when an acid neutralises a base. Salts with the same positive ion (cation) or negative ion (anion) are said to belong to the same family (for example, sodium salts: NaCl, Na₂SO₄).

• ★pH of salts:
  • Salt from a strong acid + strong base → neutral salt (pH ~ 7).
  • Strong acid + weak base → acidic salt (pH < 7).
  • Strong base + weak acid → basic salt (pH > 7).

Common Chemicals from Sodium Chloride (NaCl)

Sodium chloride (common salt, or "brine" when dissolved in water) is an important raw material. Some major products are:

1. Sodium hydroxide (NaOH) — Chlor-alkali process

• ★Electrolysis of brine: 2NaCl + 2H₂O → 2NaOH + Cl₂↑ + H₂↑
• ★Products and uses:
  • Cl₂ (Chlorine): Disinfectant, used to make PVC and other chemicals.
  • H₂ (Hydrogen): Fuel and used to make ammonia for fertilisers.
  • NaOH: Used in soap, paper, and textile industries.

2. Bleaching powder (CaOCl₂)

• ★Made by passing chlorine over dry slaked lime: Ca(OH)₂ + Cl₂ → CaOCl₂ + H₂O
• ★Uses: bleaching fabrics, disinfecting water, and as an oxidising agent.

3. Baking soda (NaHCO₃) — Sodium hydrogencarbonate

• ★Prepared by Solvay-type reactions or laboratory routes; mild, non-corrosive base.
• ★On heating it gives carbon dioxide which helps dough rise: 2NaHCO₃ (heat) → Na₂CO₃ + H₂O + CO₂↑
• ★Uses: baking, antacids, and soda-acid fire extinguishers.

4. Washing soda (Na₂CO₃·10H₂O)

• ★Obtained by crystallising sodium carbonate with water of crystallisation: Na₂CO₃ + 10H₂O → Na₂CO₃·10H₂O
• ★Uses: glass and soap industries, softening hard water and household cleaning.

7. Are Crystals of Salts Really Dry?

Many salts contain a fixed number of water molecules in their crystals. This is called water of crystallisation.

• ★Example: Hydrated copper sulphate is CuSO₄·5H₂O. It is blue because of the water. On heating it loses water and becomes white (anhydrous CuSO₄). Adding water restores the blue colour.
• ★Gypsum is CaSO₄·2H₂O. When heated it gives plaster of Paris (POP): CaSO₄·2H₂O (heat) → CaSO₄·½H₂O.
• ★Plaster of Paris reacts with water to set back into gypsum (hard solid): CaSO₄·½H₂O + 1½H₂O → CaSO₄·2H₂O
• ★Uses of POP: setting fractured bones, making casts, models and decorations.

Note for students: Learn the main reactions by understanding the ideas behind them — for example, neutralisation always makes a salt and water, and carbonates always give CO₂ with acids. Practice writing a few simple equations and learn common examples (HCl, NaOH, NaCl, Na₂CO₃, Ca(OH)₂).