Ch 2Acids, Bases and Salts

Chapter 2: Acids, Bases and Salts

1. Acids and Bases — Basic Definitions

Acids are substances that are sour in taste and turn blue litmus red. In aqueous solution they produce hydrogen ions, H⁺(aq). Examples: HCl (hydrochloric acid), H₂SO₄ (sulphuric acid), HNO₃ (nitric acid), CH₃COOH (acetic acid, in vinegar), citric acid (lemon), tartaric acid (tamarind).

Bases are bitter in taste, soapy to touch and turn red litmus blue. In aqueous solution they produce hydroxide ions, OH⁻(aq). Examples: NaOH (sodium hydroxide), KOH (potassium hydroxide), Ca(OH)₂ (calcium hydroxide), Mg(OH)₂ (milk of magnesia). Bases soluble in water are called alkalis (e.g. NaOH, KOH). All alkalis are bases but all bases are not alkalis.

2. Indicators

• Natural indicators:
  • Litmus (extracted from lichens): acid turns it red, base turns it blue; purple when neutral.
  • Turmeric: remains yellow in acid, turns reddish-brown in base (a soap stain on a turmeric curry stain turns red-brown).
  • Red cabbage extract: red/pink in acid, greenish in base.
• Synthetic indicators: Methyl orange — red in acid, yellow in base. Phenolphthalein — colourless in acid, pink in base.
• Olfactory indicators change their smell in acidic or basic medium: onion and vanilla essence lose their characteristic smell in a basic solution but retain it in acid. Useful for visually impaired students.

3. Chemical Properties of Acids and Bases

(a) Reaction with metals

Acid + active metal → salt + hydrogen gas. H₂ gas burns with a pop sound when a burning candle is brought near it.

Some metals also react with strong bases to give hydrogen:

(b) Reaction with metal carbonates and metal hydrogencarbonates

Both give salt + water + carbon dioxide. CO₂ turns lime water milky due to formation of insoluble CaCO₃; with excess CO₂ the milkiness disappears because soluble Ca(HCO₃)₂ forms.

(c) Neutralisation: acid + base

In general: Base + Acid → Salt + Water. Essentially H⁺ + OH⁻ → H₂O.

(d) Reaction of acids with metal oxides

Metal oxides are basic; they neutralise acids to give salt and water.

(The solution turns blue-green due to copper(II) chloride.)

(e) Reaction of bases with non-metal oxides

Non-metal oxides are acidic; they react with bases to give salt and water.

4. What Do All Acids Have in Common?

All acids produce H⁺(aq) ions in water — this is responsible for their acidic behaviour. H⁺ cannot exist alone; it combines with water to form the hydronium ion:

Dry HCl experiment: Dry HCl gas (from NaCl + conc. H₂SO₄) does not change the colour of dry blue litmus paper, but turns moist litmus paper red. Conclusion: acids show acidic character only in the presence of water, because ionisation (formation of H⁺ ions) needs water. Solutions of acids and bases conduct electricity because they contain ions; solutions of glucose and alcohol do not conduct and are not acids even though they contain hydrogen.

Bases give OH⁻(aq) in water:

Dilution

Mixing an acid or base with water is highly exothermic. Always add acid to water slowly with constant stirring, never water to concentrated acid — otherwise the heat generated can splash the acid out and cause burns, and the glass container may crack. Dilution decreases the concentration of ions (H₃O⁺/OH⁻) per unit volume.

5. Strength of Acids and Bases — the pH Scale

The pH scale (p = potenz, German for power) measures hydrogen ion concentration; it runs from 0 to 14.

• pH < 7 → acidic (lower pH = stronger acid = more H⁺ concentration)
• pH = 7 → neutral (pure water)
• pH > 7 → basic (higher pH = stronger base = more OH⁻ concentration)

Strong acids ionise completely in water (HCl, H₂SO₄, HNO₃); weak acids ionise only partially (CH₃COOH, citric acid, carbonic acid). Strong bases: NaOH, KOH; weak base: NH₄OH. A universal indicator (mixture of indicators) shows different colours at different pH values and is used with pH paper.

6. Importance of pH in Everyday Life

• Our body: works within the pH range 7.0 to 7.8. Acid rain (rain with pH below 5.6) makes river water acidic and harms aquatic life.
• Digestion: the stomach produces HCl which helps digestion. Excess acid causes indigestion and pain; antacids such as Mg(OH)₂ (milk of magnesia), a mild base, neutralise the excess acid.
• Tooth decay: starts when the pH in the mouth falls below 5.5 — bacteria degrade sugar and food particles to acids which corrode tooth enamel (calcium hydroxyapatite, the hardest substance in the body). Basic toothpastes neutralise the acid.
• Soil: plants require a specific pH range for healthy growth; soil that is too acidic is treated with quicklime (CaO), slaked lime [Ca(OH)₂] or chalk (CaCO₃).
• Bee sting / self-defence in plants and animals: a bee sting injects methanoic (formic) acid — rubbing a mild base like baking soda gives relief. Nettle leaf hairs also inject methanoic acid; the traditional remedy is rubbing with the leaf of the dock plant, which often grows nearby.

7. Salts — Family and pH

A salt is formed by the neutralisation of an acid and a base. Salts having the same positive or negative radicals belong to a family, e.g. NaCl and Na₂SO₄ (sodium family); NaCl and KCl (chloride family).

• Strong acid + strong base → neutral salt, pH = 7 (e.g. NaCl, Na₂SO₄).
• Strong acid + weak base → acidic salt, pH < 7 (e.g. NH₄Cl).
• Weak acid + strong base → basic salt, pH > 7 (e.g. Na₂CO₃, CH₃COONa).

8. Chemicals from Common Salt (NaCl)

Common salt (from sea water or as rock salt, deposits formed when ancient seas dried up) is the raw material for several important chemicals.

(a) Sodium hydroxide — the chlor-alkali process

Electrolysis of brine (concentrated aqueous NaCl solution). It is called chlor-alkali because the products are chlor (chlorine) and alkali (NaOH).

Cl₂ is given off at the anode, H₂ at the cathode, NaOH solution forms near the cathode. Uses: Cl₂ — water treatment, PVC, disinfectants, CFCs; H₂ — fuels, margarine, ammonia for fertilisers; NaOH — soaps and detergents, paper making, artificial fibres; combined products — HCl and bleach.

(b) Bleaching powder, CaOCl₂

Produced by the action of chlorine on dry slaked lime:

Uses: bleaching cotton and linen in textiles and wood pulp in paper industry; oxidising agent in chemical industries; disinfecting drinking water.

(c) Baking soda, NaHCO₃ (sodium hydrogencarbonate)

Produced using sodium chloride as one of the raw materials:

It is a mild, non-corrosive base. On heating during cooking:

Uses: in baking powder (baking soda + a mild edible acid such as tartaric acid — on heating/mixing with water, CO₂ is released and makes cakes soft and spongy); as an antacid; in soda-acid fire extinguishers.

(d) Washing soda, Na₂CO₃·10H₂O

Obtained by heating baking soda to get sodium carbonate, then recrystallising it:

Uses: in glass, soap and paper industries; manufacture of borax; as a cleaning agent for domestic purposes; removing permanent hardness of water.

(e) Plaster of Paris, CaSO₄·½H₂O

Made by heating gypsum at exactly 373 K (above this temperature the anhydrous "dead burnt plaster" forms):

On mixing with water it sets into a hard solid mass (gypsum) — this is why it is used for setting fractured bones:

Uses: supporting fractured bones, making toys, materials for decoration, making smooth surfaces. Note: the formula CaSO₄·½H₂O means that two formula units of CaSO₄ share one molecule of water.

9. Water of Crystallisation

The fixed number of water molecules present in one formula unit of a salt is its water of crystallisation. Such salts are called hydrated salts.

• CuSO₄·5H₂O (copper sulphate crystals, blue vitriol) — on heating, blue crystals turn white (anhydrous CuSO₄) and water droplets appear on the test tube; adding water restores the blue colour. The change is reversible.
• Na₂CO₃·10H₂O — washing soda.
• CaSO₄·½H₂O — Plaster of Paris; gypsum is CaSO₄·2H₂O.