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Question Bank Ionic Equilibrium - Chemistry Class 11 - NEET PDF Download
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FAQs on Question Bank: Ionic Equilibrium
| 1. What is the difference between strong and weak acids in ionic equilibrium? |  |
Ans. Strong acids completely dissociate into ions in solution, while weak acids only partially dissociate, establishing equilibrium between molecules and ions. Strong acids like HCl ionise 100%, whereas weak acids like acetic acid ionise only a small percentage. This distinction directly affects pH calculations and equilibrium constant expressions in ionic equilibrium problems.
| 2. How do I calculate the pH of a weak acid using the ionisation constant Ka? | |
Ans. Use the Ka expression: Ka = [H⁺][A⁻]/[HA]. For weak acids, assume initial concentration minus x gives equilibrium concentrations, where x is the H⁺ produced. Solve for x using the quadratic equation or approximation method, then calculate pH = -log[H⁺]. Understanding Ka values helps predict how much a weak acid dissociates in solution.
| 3. Why does adding a common ion suppress the dissociation of a weak electrolyte? | |
Ans. The common ion effect occurs because adding ions shifts the equilibrium leftward according to Le Chatelier's principle. When acetate ions are added to acetic acid solution, the equilibrium HA ⇌ H⁺ + A⁻ shifts left, reducing H⁺ concentration and increasing pH. This suppression is crucial for buffer solutions in CBSE Class 11 chemistry.
| 4. What's the relationship between Kh, Kb, and Ka for salt hydrolysis problems? | |
Ans. The hydrolysis constant Kh relates to ionisation constants through: Kh = Kw/(Ka × Kb). Salts of weak acids undergo hydrolysis, producing acidic solutions, while salts of weak bases produce basic solutions. This relationship explains why ammonium chloride solutions are acidic and sodium acetate solutions are basic in ionic equilibrium.
| 5. How do buffer solutions work, and why don't they change pH when small amounts of acid or base are added? | |
Ans. Buffer solutions contain a weak acid and its conjugate base (or weak base and its conjugate acid), maintaining relatively constant pH. When acid is added, the conjugate base neutralises it; when base is added, the weak acid neutralises it. The Henderson-Hasselbalch equation describes this equilibrium: pH = pKa + log([A⁻]/[HA]), making buffers essential in ionic equilibrium applications.
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