Q1. Why do acidic solutions conduct electricity while glucose and alcohol solutions do not? Explain using the nail–bulb activity and suitable equations.
Answer:
In water, acids undergo ionization, producing mobile ions that carry charge. For example:
HCl(aq) → H⁺(aq) + Cl⁻(aq) (more accurately, H⁺ forms H₃O⁺(aq))
H₂SO₄(aq) → 2H⁺(aq) + SO₄²⁻(aq)
In the nail–bulb activity, a brightly glowing bulb with dilute HCl/H₂SO₄ shows strong electrical conductivity due to a high concentration of ions.
Glucose (C₆H₁₂O₆) and alcohols (e.g., C₂H₅OH) are covalent, do not dissociate into ions in water, and thus do not conduct electricity; the bulb remains off.
This proves that only ionic species in solution enable conductivity.
Key idea: Conductivity depends on the presence and mobility of ions, not merely on dissolving a substance in water.
Q2. How do bases behave in water? Describe the production of hydroxide ions, conductivity, and compare NaOH, KOH, and Mg(OH)₂.
Answer:
Bases release OH⁻ ions in water, making solutions alkaline and conductive. Examples:
NaOH and KOH are strong bases, fully dissociate, and hence their solutions are good conductors.
Mg(OH)₂ is sparingly soluble, producing fewer ions; its solutions conduct weakly, yet still show basic properties.
The OH⁻ ions are responsible for turning red litmus blue, reacting with H⁺ during neutralization, and enabling electric current flow.
Key idea: The greater the ion concentration, the stronger the conductivity and basic strength in water.
Q3. What is a neutralization reaction? Explain with molecular and ionic equations, and state its significance with examples.
Answer:
Neutralization occurs when an acid reacts with a base to form salt and water.
General reaction: Acid + Base → Salt + Water
Core ionic reaction: H⁺(aq) + OH⁻(aq) → H₂O(l)
Examples:
HCl + NaOH → NaCl + H₂O
H₂SO₄ + Ca(OH)₂ → CaSO₄ + 2H₂O
CH₃COOH + KOH → CH₃COOK + H₂O
The salt formed depends on the acid and base used.
Neutralization is often exothermic, releasing heat.
Everyday relevance:
Antacids neutralize excess stomach acid.
Lime (CaO/Ca(OH)₂) reduces soil acidity.
Wastewater treatment adjusts pH to neutral levels.
Key idea: Neutralization removes excess H⁺ or OH⁻, moving the solution toward a neutral pH.
Q4. Why must we always add acid to water and not water to acid? Explain the science and the correct lab technique.
Answer:
Mixing concentrated acids with water is highly exothermic due to the strong hydration of ions (especially H⁺ → H₃O⁺).
If water is poured into acid, the small water layer can boil instantly, causing violent splashing of hot, corrosive acid.
By adding acid to excess water slowly, the released heat is absorbed by the larger water volume, preventing splashes.
Correct technique:
Use a glass beaker or heat-resistant container.
Add acid to water in small portions, with continuous stirring.
Use PPE: goggles, lab coat, gloves; preferably work in a fume hood.
Key idea: Always remember the safety rule—A to W: Acid to Water—to control heat and avoid accidents.
Q5. Why does dry HCl gas not change the color of dry litmus paper, but moist blue litmus turns red in HCl?
Answer:
Acidity is shown only when H⁺ ions are available in aqueous solution.
Dry HCl gas contains molecules, not ions, and dry litmus has no water to allow ionization; hence, no color change occurs.
When HCl dissolves in water, it forms H⁺(aq) (actually H₃O⁺(aq)) and Cl⁻(aq); the H₃O⁺ interacts with the litmus dye, turning blue litmus red.
Therefore, moist litmus provides the water medium needed for HCl to show acidic behavior.
Key idea: The presence of water is essential for acids to ionize and exhibit their characteristic properties like indicator changes and conductivity.
High Complexity (Analytical & Scenario-Based)
Q6. You have three colorless solutions: dilute HCl, glucose, and ethanol. Design a simple experiment to identify the acid and justify your observations.
Answer:
Use the conductivity test (battery, wires, bulb, graphite/nail electrodes).
The solution that makes the bulb glow brightly is the acid (dilute HCl) due to high ion concentration: HCl → H⁺ + Cl⁻.
Glucose and ethanol are non-electrolytes; the bulb remains off.
Confirm with blue litmus:
Only dilute HCl turns blue litmus red.
Optional confirmatory test:
Add metal (e.g., Mg ribbon): HCl will produce H₂ gas (effervescence). Glucose/ethanol do not.
Justification:
Acids ionize in water to give H⁺/H₃O⁺, enabling conductivity and indicator changes.
Covalent organics (glucose, ethanol) do not furnish ions; hence, no conductivity and no acidic behavior.
Q7. Compare neutralization of strong acid–strong base (HCl + NaOH) with weak acid–strong base (CH₃COOH + NaOH). Discuss products and pH at equivalence.
Answer:
Both reactions produce salt + water via H⁺ + OH⁻ → H₂O.
Strong acid–strong base:
HCl + NaOH → NaCl + H₂O
At equivalence, the solution is approximately neutral (pH ≈ 7) because Na⁺ and Cl⁻ are spectator ions and do not hydrolyze.
Weak acid–strong base:
CH₃COOH + NaOH → CH₃COONa + H₂O
At equivalence, solution is slightly basic (pH > 7) because the acetate ion (CH₃COO⁻) undergoes hydrolysis, generating some OH⁻.
Conductivity:
During titration, conductivity decreases as ions neutralize; final conductivity depends on the salt’s ions present.
Key idea: The strength of the acid/base affects the pH at equivalence and the nature of the salt formed.
Q8. If Mg(OH)₂ is sparingly soluble, how does milk of magnesia effectively neutralize excess stomach acid?
Answer:
Milk of magnesia contains suspended Mg(OH)₂, a weak, sparingly soluble base.
Despite low solubility, it establishes an equilibrium that continuously supplies small amounts of OH⁻ to neutralize H⁺ in gastric juice.
As H⁺ from stomach acid is consumed, more Mg(OH)₂ dissolves to restore equilibrium, providing a controlled, sustained neutralization:
Mg(OH)₂ + 2HCl → MgCl₂ + 2H₂O
Benefits:
Reduced risk of over-alkalizing the stomach compared to highly soluble strong bases.
The formed MgCl₂ is relatively harmless in small medicinal doses.
Key idea: Limited solubility offers a buffer-like, gentle neutralization, making Mg(OH)₂ an effective and safer antacid.
Q9. In the bulb experiment, 1% HCl glows brightly but 1% CH₃COOH glows dimly. Explain using strong vs. weak electrolytes and ionization.
Answer:
HCl is a strong acid and strong electrolyte; it fully ionizes in water, producing a high concentration of H⁺/H₃O⁺ and Cl⁻. Hence, high conductivity and a bright bulb.
CH₃COOH (acetic acid) is a weak acid and weak electrolyte; it partially ionizes:
CH₃COOH ⇌ H⁺ + CH₃COO⁻
Fewer ions mean lower current and a dim bulb at the same concentration.
With dilution, weak acids can ionize slightly more (greater degree of ionization), but the total ion count per volume may still be lower than strong acids.
Key idea: Conductivity depends on the number of free ions, which is much higher for strong acids than weak acids at equal concentrations.
Q10. A student added water to concentrated H₂SO₄ and it splashed. Analyze why this happened, immediate first aid, and the correct preventive method.
Answer:
Cause:
Diluting concentrated H₂SO₄ is highly exothermic. Adding water to acid forms a thin water layer that heats rapidly, leading to violent boiling and splashing of corrosive acid.
Immediate actions:
Move to safety; rinse affected area with copious running water for at least 15 minutes.
Remove contaminated clothing; seek medical help. Do not apply neutralizers on skin.
Prevention:
Always add acid to water (A → W) slowly with constant stirring in a heat-resistant container; use an ice bath if needed.
Wear PPE: goggles, gloves, lab coat; work in a well-ventilated area/fume hood.
Key idea: Control the heat of dilution by ensuring a large water volume absorbs the released heat safely.
