7 Reaction Kinetics

VICTORIA JUNIOR COLLEGE CHEMISTRY DEPARTMENT Dr WU Jiang (wu.jiang@vjc.sg) REACTION KINETICS Lecture 1 2 3 4 5 Pages 3-8 8-14 15-20 20-25 26-31 Completed by T2W6 T2W7 T2W8 T2W9 Tutorial Q Q1 – Q3 Q4, Q5 Q6, Q7 Q8-Q11 Lecture Outline 1 Introduction 2 Concept of Reaction Rate 3 Experimental methods of finding reaction rates 3.1 Titrimetric Analysis 3.2 Colorimetric Analysis 3.3 Volume measurements 3.4 Pressure measurements 3.5 Conductometric measurements 4 The Rate Equation 4.1 The Rate Constant 4.2 The Order of Reaction 4.2.1 Zero order reaction 4.2.2 First order reaction 4.2.3 Second order reaction 4.2.4 Pseudo order reaction 5 Methods to determine order of reaction 5.1 Graphical method / Continuous method 5.2 Initial rate method 6 Theories of reaction rate 6.1 The Collision Theory 6.2 The Transition State Theory 7 Factors affecting rate of reaction 7.1 Concentration 7.2 Temperature 7.3 Light 7.4 Surface area 7.5 Catalysis 8 Reaction Mechanism Reference Texts 1. Chemistry for Advanced Level by Peter Cahn / Hughes 2. Advanced A Level Chemistry by Ramsden 3. Chemistry in Context by Hill & Holman

Victoria Junior College JC1 2021 Reaction Kinetics Page 2 Learning outcomes (9729 H2 Chemistry) Candidates should be able to: (a) explain and use the terms: rate of reaction; rate equation; order of reaction; rate constant; half –life of a reaction; rate–determining step; activation energy; catalysis (b) construct and use rate equations of the form rate = k[A] m[B]n (limited to simple cases of single –step reactions and of multi –step processes with a rate –determining step, for which m an d n are 0, 1 or 2), including: (i) deducing the order of a reaction by the initial rates method (ii) justifying, for zero– and first–order reactions, the order of reaction from concentration–time graphs (iii) verifying that a suggested reaction mechanism is consistent with the observed kinetics (iv) predicting the order that would result from a given reaction mechanism (v) calculating an initial rate using concentration data [integrated forms of rate equations are not required] (c) (i) show understanding that the half–life of a first–order reaction is independent of concentration (ii) use the half–life of a first–order reaction in calculations (d) calculate a rate constant using the initial rates method (e) devise a suitable experimental technique for studying the rate of a reaction, from given information (f) explain qualitatively, in terms of collisions, the effect of concentration changes on the rate of a reaction (g) show understanding, including reference to the Boltzmann distribution, of what is meant by the term activation energy (h) explain qualitatively, in terms both of the Boltzmann distribution and of collision frequency, the effect of temperature change on a rate constant (and hence, on the rate) of a reaction (i) (i) explain that, in the presence of a catalyst, a reaction has a different mechanism, i.e. one of lower activation energy, giving a larger rate constant (ii) interpret this catalytic effect on a rate constant in terms of the Boltzmann distribution (j) outline the different modes of action of homogeneous and heterogeneous catalysis, including: (i) the Haber process (ii) the catalytic removal of oxides of nitrogen in the exhaust gases from car engines (iii) the catalytic role of atmospheric oxides of nitrogen in the oxidation of atmospheric sulfur dioxide (iv) catalytic role of Fe2+ in the Ι–/S2O82– reaction (k) describe enzymes as biological catalysts which may have specific activity (l) explain the relationship between substrate concentration and the rate of an enzyme–catalysed reaction in biochemical systems Learning outcomes (8873 H1 Chemistry) Candidates should be able to: (a) explain and use the terms: rate of reaction; rate equation; order of reaction; rate constant; half–life of a reaction; activation energy; catalysis; (b) construct and use rate equations of the form rate = k[A] m[B]n (limited to simple cases of single –step reactions, for which m and n are 0, 1 or 2), including: (i) deducing the order of a reaction by the initial rates method; (ii) justifying, for zero– and first–order reactions, the order of reaction from concentration–time graphs; (iii) calculating an initial rate using concentration data; [integrated forms of rate equations are not required] (c) show understanding that the half–life of a first–order reaction is independent of concentration; (d) explain qualitatively, in terms of collisions, the effect of concentration changes on the rate of a reaction; (e) show understanding, including reference to the Boltzmann dist ribution, of what is meant by the term activation energy; (f) explain qualitatively, in terms both of the Boltzmann distribution and of collision frequency, the effect of temperature change on a rate constant (and, hence, on the rate) of a reaction; (g) (i) explain that, in the presence of a catalyst, a reaction follows a different pathway, i.e. one of lower activation energy, giving a larger rate constant; (ii) interpret this catalytic effect in terms of the Boltzmann distribution; (h) Outline the mode of action of heterogeneous catalysis, as exemplified by the catalytic removal of oxides of nitrogen in the exhaust gases from car engines. (i) describe enzymes as biological catalysts which may have specific activity.

Victoria Junior College JC1 2021 Reaction Kinetics Page 3 1 Introduction Industrially, chemists want to know the rate of a reaction as they want the most economical process of making a product. This would not only save time but makes it profitable as well. Archaeologists can estimate the age of fossils of prehistoric remains by a process known as radioactive decay. This involves measuring the concentration of a decaying radioactive isotope 14C in the object under study. Reaction kinetics is the study of the rates of chemical reactions, which include the factors that affect them and the mechanisms by which the reactions occur. 2 Concept of Reaction Rate The rate of reaction can be defined as the change in the concentration of a particular reactant or product per unit time. Mathematically, it can be calculated as: Rate of reaction = The rate can show how fast a product is formed or how fast a reactant is used up in a given time. It has the unit of mol dm–3 s–1. (sometimes, mol dm–3 min–1 may be used.) For a hypothetical reaction, aA + bB → cC + dD Worked Example 1 For reaction, 2NO2(g) → 2NO(g) + O2 (g), the concentration-time graph is given below. NO2 NO O2 rate =  a 1 dt ]A[d =  b 1 dt ]B[d = c 1 dt ]C[d = d 1 dt ]D[d Calculate the rate of reaction at  t = 100 s  t = 250 s

Victoria Junior College JC1 2021 Reaction Kinetics Page 4 Making Thinking Visible 1. How do we get the rate consumption of a reactant or formation of a product from the concentration-time graph? 2. Is the rate of consumption of a reactant or formation of a product always the same as the rate of reaction? 3. Why is it necessary to divide by the stoichiometric coefficient when calculating the rate of reaction? 4. Why is there a negative sign when determining the rate of reaction based on the concentration of a reactant? Exercise 1: (with reference to the graph on page 3) 1. Without drawing a tangent, suggest the rate of formation of O2 at 100 s. 2. Determine from the graph the rate of reaction at t = 150 s. Instantaneous rate v.s. Average rate In the earlier discussion, we have used the instantaneous rate to represent the rate of reactio