1. CHEMICAL KINETICS CHAPTER 17, Kinetics Fall 2009, CHEM 1310 1

2. KeKin Kinetics vs Thermodynamics A: Reactants B: Transition state C: products E: Forward Activation Free Energy F: Reverse Activation Free Energy

3. Most reactions occur through several steps and are not single step reactions. Each step in a multi-step reaction is called an elementary reaction. Types of elementary reactions 1.Unimolecular (a single reactant) 2.Bimolecular 3.Termolecular (very unlikely) 3 Reaction Mechanisms

4. Each step of this reaction is an “elementary step”. Each elementary step has reactant(s), a transition state, and product(s). Products that are consumed in subsequent elementary reaction are called intermediates.

5. 5 Reaction Rates: To measure a reaction rate we could monitor the disappearance of reactants or appearance of products. e.g., 2NO 2 + F 2 → 2NO 2 F

6. 6 Gen. Rxn:aA + bB → cC + dD NO 2 + CO → NO + CO 2

7. 7 Order of a Reaction The power (n) to which the concentration of A is raised in the rate expression describes the order of the reaction with respect to A. Do not confuse the order (n) with the stoichiometric coefficient (a).

8. 8 m th order in [A] n th order in [B]

9. 9 [A] (mol L -1 ) [B] (mol L -1 )Rate (mol L -1 s -1 ) 11.0x10 -4 1.0x10 -4 2.8x10 -6 21.0x10 -4 3.0x10 -4 8.4x10 -6 32.0x10 -4 3.0x10 -4 3.4x10 -5

10. 10 [A] (mol L -1 ) [B] (mol L -1 )Rate (mol L -1 s -1 ) 11.0x10 -4 1.0x10 -4 2.8x10 -6 21.0x10 -4 3.0x10 -4 8.4x10 -6 32.0x10 -4 3.0x10 -4 3.4x10 -5

11. Example: At elevated temperatures, HI reacts according to the chemical equation 2HI → H 2 + I 2 The rate of reaction increases with concentration of HI, as shown in this table. Data [HI]Rate Point (mol L -1 ) (mol L -1 s -1 ) 1 0.0057.5 x 10 -4 2 0.0103.0 x 10 -3 3 0.0201.2 x 10 -2 a) Determine the order of the reaction with respect to HI and write the rate expression b) Calculate the rate constant and give its units c) Calculate the instantaneous rate of reaction for a [HI] = 0.0020M 11

12. INTEGRATED RATE LAWS Single Reactant (three cases) – Zero-Order Rate Law (n = 0) – First-Order Rate Law (n = 1) – Second-Order Rate Law (n = 2) More than one Reactant – Must state the order of the reaction with respect to each reactant (rate = k[A] n [B] m [C] p ) 12

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14. 14 INTEGRATED RATE LAWS n=0,1,2

15. 15 In the real world, if we do not know the order of the reaction we can use experimental plots to estimate the order. If a plot of [A] vs t is a straight line, then the reaction is zero order. If a plot of ln[A] vs t is a straight line, then the reaction is 1 st order. If a plot of 1/ [A] vs t is a straight line, then the reaction is 2 nd order.

16. INTEGRATED RATE LAWS 16 Zero Order Reactions [A] - [Ao] = -kt Graph [A] vs t Slope = -k, intercept = [Ao]

17. INTEGRATED RATE LAWS 17 First Order Reactions ln[A]-ln[Ao] = -kt Graph ln[A] vs t Slope = -k, intercept = [Ao]

18. 18 In[N 2 0 5 ] versus time. Slope = - k 2 N 2 O 5 (g) → 4 NO 2 (g) + O 2 (g) This graph gives a straight line, and so is First order with respect to the decomposition of N 2 O 5 If a plot of ln[A] vs t is not a straight line, the reaction is not first order!

19. 19 Dimerization Data set provided [C 4 H 6 ] vs time 2 C 4 H 6 (g) → C 8 H 12 (g) [C 4 H 6 ] ˚ = 0.01M

21. Most reactions proceed not through a single step but through a series of steps Each Step is called an elementary reaction Types of elementary reactions 1.Unimolecular (a single reactant) E.g., A → B + C (a decomposition) 2.Bimolecular (most common type) E.g., A + B → products 3.Termolecular (less likely event) E.g., A + B + C → products 21 Reaction Mechanisms

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23. 23 Notice that NO 3 is formed and consumed. This is called a __________________________________. Notice also that Step 1 is bimolecular and Step 2 is bimolecular

24. CHEMICAL EQUILIBRIUM 24 A direct connection exists between the equilibrium constant of a reaction and the rate constants. a) at equilibrium: forward reaction rate = reverse reaction rate. b) K eq = k f / k r (same as K = k 1 /k -1 ) A ⇌ B kfkf krkr

25. REACTION MECHANISM & RATE LAWS 25 Typically with a reaction one of several elementary step reaction is the slowest step. This is called the Rate Determining Step (RDS) Case #1: When the RDS occurs first, the first step is slow and determines the rate of the overall reaction. Example 15.6

26. 26 Reaction Progress EnergyEnergy F + NO 2 F NO 2 + F 2 slow fast NO 2 F

27. Chem 1310 Spring 2009 stop here

28. 28 Reaction Progress EnergyEnergy

29. 29 Need to express [intermediates] in terms of other reactants

30. 30 Substituting for [N 2 O 2 ] in the rate expression above

31. 31 Reaction Progress EnergyEnergy N 2 O 2 + O 2 slow fast 2NO 2NO 2

32. 32 Reaction Progress EnergyEnergy Reaction Mechanism Intermediates Transition states

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34. A MODEL FOR CHEMICAL KINETICS 34

35. 35 Chapter 5: The Kinetic Molecular Theory of Gases The Meaning of Temperature: temperature is a measure of the average kinetic energy of the gas particles. The Kelvin temperature of a gas is a measure of the random motions of the particles of gas. With higher temperature, greater motion.

36. 36 Chapter 5: Speed Distribution Curves Maxwell-Boltzmann speed distribution Temperature is a measure of the average kinetic energy of molecules when their speeds have Maxwell Boltzmann distribution. i.e., the molecules come to thermal equilibrium.

37. 37 Transition State, also called Activated Complex Two requirements must be satisfied for reactants to collide successfully to rearrange to form products 1. 2.

38. 38 Consider two different temperatures. 1.) Collisions must have enough energy to produce a reaction. Not all collisions have enough energy to make product E collision > E act Number of collisions Distribution of velocities

39. 39 2 BrNO (g) → 2 NO (g) + Br 2 (g) 2.) Molecular Orientation Relative orientations of the reactants must allow formation of any new bonds to produce products. Orientation a or b lead to product, c does not.

42. 42 Find the rate constant k at several temperatures. Plot of In(k) versus 1/T for the reaction y = mx + b Slope =

43. 43 Reaction Progress EnergyEnergy EarEar EafEaf Transition State The Activation Energy (E a ) is the minimum collision energy that reactants must have in order to form products

44. 44 Reaction Progress EnergyEnergy ΔE = E a f - E a r EarEar EafEaf Transition State The Activation Energy (E a ) is the minimum collision energy that reactants must have in order to form products

48. CHEMICAL KINETICS Catalyst Inhibitor 48

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51. KINETICS OF CATALYSIS A catalyst has no effect on the thermodynamics of the overall reaction It only provides a lower energy path Examples – Pt and Pd are typical catalysts for hydrogenation reactions (e.g., ethylene to ethane conversion) – Enzymes act as catalysts Phases – Homogenous catalysis – the reactants and catalyst are in the same catalyst (gas or liquid phase) – Heterogeneous catalysis – reaction occurs at the boundary of two different phases (a gas or liquid at the surface of a solid) 51

52. 52 Effect of a catalyst on the number of reaction-producing collisions.