Acids, Bases and Salts
Overview
Acids, Bases and Salts form one of the most frequently tested chemistry topics in KTET Category II/III examinations. This topic bridges everyday life experiences (lemon sourness, soap slipperiness, cooking soda reactions) with fundamental chemical concepts, making it ideal for classroom teaching and exam questions alike.
For KTET aspirants, mastery requires understanding three interconnected areas: the characteristic properties of acids and bases, how indicators help identify them, and how neutralisation reactions produce salts. Questions typically test identification of common acids/bases, pH scale interpretation, indicator colour changes, and practical applications. The pedagogical component often asks how to demonstrate these concepts through safe classroom experiments.
This topic connects directly to environmental studies (acid rain, soil pH) and biology (digestion, blood pH), making it valuable for integrated science teaching at upper primary and secondary levels.
Key Concepts
- Acids release hydrogen ions (H⁺) when dissolved in water. The word "acid" comes from Latin "acidus" meaning sour. All acids contain hydrogen that can be replaced by metals.
- Bases release hydroxide ions (OH⁻) in water. Bases that dissolve in water are called alkalis. All alkalis are bases, but not all bases are alkalis (e.g., copper oxide is a base but insoluble).
- pH scale measures acidity/alkalinity from 0 to 14. pH below 7 is acidic, pH 7 is neutral, pH above 7 is basic. Lower pH means stronger acid; higher pH means stronger base.
- Indicators are substances that show different colours in acidic and basic solutions. They help identify the nature of unknown solutions without tasting (which is unsafe).
- Neutralisation is the reaction between acid and base producing salt and water: Acid + Base → Salt + Water. This reaction releases heat (exothermic).
- Salts are ionic compounds formed when the hydrogen of an acid is replaced by a metal or ammonium ion. Salts may be acidic, basic or neutral depending on the parent acid and base.
- Strong acids/bases ionise completely in water (HCl, NaOH), while weak acids/bases ionise partially (acetic acid, ammonium hydroxide).
Formulas / Key Facts
| Formula/Fact | Context |
|---|---|
| HCl → H⁺ + Cl⁻ | Hydrochloric acid ionisation in water |
| NaOH → Na⁺ + OH⁻ | Sodium hydroxide ionisation |
| HCl + NaOH → NaCl + H₂O | Neutralisation reaction |
| 2HCl + CaCO₃ → CaCl₂ + H₂O + CO₂ | Acid reacts with carbonate releasing CO₂ |
| pH = -log[H⁺] | Mathematical definition (conceptual understanding sufficient) |
| pH of pure water = 7 | Neutral reference point |
| Gastric juice pH ≈ 1.5 to 2 | Strongly acidic for digestion |
| Blood pH ≈ 7.35 to 7.45 | Slightly basic, tightly regulated |
Common Acids: Hydrochloric acid (HCl) — stomach acid; Sulphuric acid (H₂SO₄) — battery acid, "king of chemicals"; Nitric acid (HNO₃) — fertiliser production; Acetic acid (CH₃COOH) — vinegar; Citric acid — citrus fruits.
Common Bases: Sodium hydroxide (NaOH) — soap making; Calcium hydroxide Ca(OH)₂ — whitewashing; Magnesium hydroxide Mg(OH)₂ — antacid (milk of magnesia); Ammonium hydroxide (NH₄OH) — cleaning agents.
Common Indicators:
- Litmus: Red in acid, Blue in base
- Phenolphthalein: Colourless in acid, Pink in base
- Methyl orange: Red in acid, Yellow in base
- Turmeric: Yellow in acid, Reddish-brown in base
Worked Examples
Example 1: Identifying acid or base using litmus
A student tests an unknown solution with blue litmus paper. The paper turns red. What is the nature of the solution?
Step 1: Recall that blue litmus turns red in acidic solutions. Step 2: Blue litmus remains blue in neutral/basic solutions. Step 3: Since the paper turned red, the solution is acidic.
Example 2: Writing a neutralisation equation
Write the balanced equation when sulphuric acid reacts with potassium hydroxide.
Step 1: Identify reactants — H₂SO₄ (acid) and KOH (base) Step 2: Products are salt + water Step 3: Salt formed = potassium sulphate (K₂SO₄) Step 4: Balance: H₂SO₄ + 2KOH → K₂SO₄ + 2H₂O
Note: Two KOH molecules needed because sulphuric acid has two replaceable hydrogen atoms.
Example 3: pH and strength
Arrange the following in increasing order of acidic strength: Solution A (pH=3), Solution B (pH=6), Solution C (pH=1)
Step 1: Lower pH = more acidic = stronger acid Step 2: Arrange from highest pH to lowest: B (pH=6) < A (pH=3) < C (pH=1) Step 3: Answer: B < A < C (increasing acidic strength)
Common Mistakes
- Confusing "all bases are alkalis" → Only water-soluble bases are alkalis. Copper oxide and iron hydroxide are bases but not alkalis because they don't dissolve in water.
- Thinking pH 0 is neutral because zero seems "nothing" → pH 7 is neutral. pH 0 is extremely acidic (ten million times more H⁺ than neutral water).
- Assuming all salts are neutral → Salts can be acidic (NH₄Cl from weak base + strong acid), basic (Na₂CO₃ from strong base + weak acid), or neutral (NaCl from strong base + strong acid).
- Confusing acid + metal with acid + carbonate reactions → Acid + metal produces hydrogen gas (H₂) with "pop" sound. Acid + carbonate produces carbon dioxide (CO₂) which turns lime water milky.
- Forgetting that concentrated acids are not necessarily strong → Concentrated refers to amount dissolved; strong refers to degree of ionisation. Concentrated acetic acid is still a weak acid.
Quick Reference
- Acids taste sour, bases taste bitter and feel soapy (never taste in lab!)
- Blue litmus → Red = Acid; Red litmus → Blue = Base
- pH scale: 0-7 acidic, 7 neutral, 7-14 basic
- Neutralisation: Acid + Base → Salt + Water (always exothermic)
- Antacids (Mg(OH)₂) neutralise excess stomach acid (HCl)
- Baking soda (NaHCO₃) + acid releases CO₂ — used in fire extinguishers and cooking