Chemical Kinetics

Chemical kinetics is the branch of chemistry that deals with the rate of chemical reactions, the factors affecting it and the mechanism by which reactions occur. In this chapter you will learn to express average and instantaneous rates in terms of the change in concentration of reactants or products, and how rate depends on concentration, temperature and catalysts. You will distinguish between the rate law and the law of mass action, and between the order of a reaction, which is determined experimentally, and its molecularity, which applies only to elementary reactions, and you will derive the units of rate constants. The chapter then develops the integrated rate equations for zero order reactions, [R] = [R]0 - kt, and first order reactions, k = (2.303/t) log([R]0/[R]), along with the half-life of each and the idea of pseudo first order reactions such as the acid hydrolysis of ethyl acetate. Finally, you will study the temperature dependence of rate through the Arrhenius equation, k = A e^(-Ea/RT), activation energy, the effect of a catalyst, and the basic ideas of collision theory.

Detailed Worksheet: Chemical Kinetics Section A - Definitions (10 marks) 1. Distinguish between average rate and instantaneous rate of a reaction. (2 marks) 2. Distinguish between order and molecularity of a reaction on two bases. (2 marks) 3. Define half-life of a reaction. Write its expression for a first order reaction. (2 marks) 4. What is a pseudo first order reaction? Give one example. (2 marks) 5. Define activation energy. How does a catalyst affect it? (2 marks) Section B - Calculations and Applications (15 marks) 6. A first order reaction has a rate constant of 6.93 x 10^-3 s^-1. Calculate its half-life, and the time required for 75% completion of the reaction. (3 marks) 7. The rate constant of a reaction doubles when the temperature is raised from 298 K to 308 K. Calculate the activation energy (R = 8.314 J K^-1 mol^-1, log 2 = 0.3010). (3 marks) 8. For the reaction A + B -> products, the following data were obtained: experiment 1, [A] = 0.1 M, [B] = 0.1 M, initial rate 2.0 x 10^-3 M/s; experiment 2, [A] = 0.2 M, [B] = 0.1 M, rate 4.0 x 10^-3 M/s; experiment 3, [A] = 0.1 M, [B] = 0.2 M, rate 8.0 x 10^-3 M/s. Find the order with respect to A and B, the rate law and the rate constant with units. (3 marks) 9. For the reaction 2N2O5 -> 4NO2 + O2, the rate of disappearance of N2O5 is 0.02 mol L^-1 s^-1. Calculate the rate of formation of NO2 and of O2, and the rate of reaction. (3 marks) 10. A zero order reaction has k = 0.002 mol L^-1 s^-1 and initial concentration 0.40 mol L^-1. Calculate the concentration after 100 s, the half-life, and the time for complete reaction. (3 marks) Section C - Diagrams (10 marks) 11. Draw the energy profile diagram for an exothermic reaction with and without a catalyst, labelling reactants, products, activated complex, activation energy and enthalpy change. (4 marks) 12. Draw graphs of (i) concentration versus time and (ii) rate versus concentration for a zero order reaction, and of log[R] versus time for a first order reaction, stating what the slope represents in each straight line graph. (3 marks) 13. Draw the Maxwell-Boltzmann distribution curves at temperatures T and T + 10 K, and mark the fraction of molecules with energy equal to or greater than the activation energy. (3 marks) Section D - Analysis and Higher-order Thinking (15 marks) 14. Show that for a first order reaction the time required for 99.9% completion is about ten times its half-life. Explain why the half-life of a first order reaction is independent of initial concentration, while that of a zero order reaction is proportional to it. (5 marks) 15. The rate constant of a first order reaction at 500 K is 2 x 10^-4 s^-1 and at 700 K is 8 x 10^-3 s^-1 (log 40 = 1.602). Calculate the activation energy, and explain on the basis of the Arrhenius equation why rate increases so rapidly with temperature. (5 marks) 16. A reaction has the rate law rate = k[NO]^2[O2]. Explain how the rate changes when (i) the concentration of NO is doubled (ii) the concentration of O2 is tripled (iii) the volume of the vessel is halved. Analyse why the order of a reaction may differ from the stoichiometric coefficients. (5 marks) Instructions: Time allowed 2 hours. Attempt all sections. Show all calculations with formulae and units. Use log values given in the questions.