Equilibrium

When a reaction takes place in a closed vessel, the rates of the forward and backward reactions eventually become equal, and the concentrations of reactants and products stop changing. This dynamic equilibrium occurs in physical processes like ice melting at 273 K and in reactions like H2 + I2 <-> 2HI. This chapter explains the law of chemical equilibrium and the equilibrium constants Kc and Kp, related by Kp = Kc(RT)^(delta n). It covers homogeneous and heterogeneous equilibria, the use of the reaction quotient Q to predict the direction of reaction, and the relation delta G zero = -RT ln K. Le Chatelier's principle explains how concentration, pressure, temperature, catalysts and inert gases affect equilibrium, as in the Haber process. The second part covers ionic equilibrium: Arrhenius, Bronsted-Lowry and Lewis concepts of acids and bases, the ionic product of water (Kw = 1.0 x 10^-14 at 298 K), the pH scale, the ionisation constants Ka and Kb, the common ion effect, hydrolysis of salts, buffer solutions and the solubility product of sparingly soluble salts.

Detailed Worksheet: Equilibrium Section A - Definitions (10 marks) 1. What is dynamic equilibrium? Give one example from a physical process. (2 marks) 2. State the law of chemical equilibrium. (2 marks) 3. State Le Chatelier's principle. (2 marks) 4. What is a Bronsted-Lowry acid? Identify the conjugate base of H2O and of HCO3-. (2 marks) 5. What is a buffer solution? Give one example. (2 marks) Section B - Calculations and Applications (15 marks) 6. For the reaction H2(g) + I2(g) <-> 2HI(g), the equilibrium concentrations at 731 K are [H2] = 0.10 M, [I2] = 0.10 M and [HI] = 0.70 M. Calculate Kc. If the reaction is written as HI(g) <-> 1/2 H2(g) + 1/2 I2(g), what is the new equilibrium constant? (3 marks) 7. For N2(g) + 3H2(g) <-> 2NH3(g), Kc = 0.50 at 700 K. Calculate Kp (R = 0.0831 bar L per mol per K). State whether high or low pressure favours the formation of ammonia. (3 marks) 8. Calculate the pH of: (i) 0.001 M HCl (ii) 0.01 M NaOH (iii) a 0.10 M acetic acid solution with Ka = 1.8 x 10^-5. (3 marks) 9. The solubility product of AgCl at 298 K is 1.8 x 10^-10. Calculate its molar solubility in water. How will its solubility change in 0.10 M NaCl? Calculate the new solubility. (3 marks) 10. For a reaction at 298 K, K = 10. Calculate delta G zero using delta G zero = -2.303 RT log K (R = 8.314 J per mol per K). What does the sign of delta G zero indicate about the reaction? (3 marks) Section C - Diagrams (10 marks) 11. Draw a graph showing how the concentrations of H2, I2 and HI change with time until equilibrium is reached. (4 marks) 12. Draw a graph showing how the rates of the forward and backward reactions change with time and become equal. (3 marks) 13. Draw a pH scale from 0 to 14, marking acidic, neutral and basic ranges, and place on it gastric juice, blood, pure water and sodium hydroxide solution. (3 marks) Section D - Analysis and Higher-order Thinking (15 marks) 14. Using Le Chatelier's principle, explain the conditions used in the Haber process for ammonia. Why does a catalyst not change the equilibrium composition? (5 marks) 15. Explain how the reaction quotient Q can be compared with K to predict the direction of a reaction. For PCl5 <-> PCl3 + Cl2 with Kc = 8.3 x 10^-3, predict the direction if [PCl5] = 0.20 M, [PCl3] = 0.10 M and [Cl2] = 0.10 M. (5 marks) 16. Explain the common ion effect and the action of an acetic acid-sodium acetate buffer when a small amount of acid or base is added. Calculate the pH of a buffer that is 0.10 M in both acetic acid and sodium acetate (pKa = 4.74). (5 marks) Instructions: Time allowed 2 hours. Attempt all sections. Show all calculations with units. Use '<->' for reversible reactions. Draw graphs neatly with labelled axes.