1. Chemical Kinetics 1 Chemical kinetics Plan 1. The subject of a chemical kinetics. 2. Classification of chemical reactions. 3. Determination methods of the rates of chemical reaction. 4. Reaction order and Molecularity of the chemical reaction. 5. Determination methods of the reaction order. Assistant Kozachok S.S. prepared

2. Chemical Kinetics Studies the rate at which a chemical process occurs. Besides information about the speed at which reactions occur, kinetics also sheds light on the reaction mechanism (exactly how the reaction occurs).

3. Chemical Kinetics Outline: Kinetics Reaction Rates How we measure rates. Rate Laws How the rate depends on amounts of reactants. Integrated Rate Laws How to calc amount left or time to reach a given amount. Half-life How long it takes to react 50% of reactants. Arrhenius Equation How rate constant changes with T. Mechanisms Link between rate and molecular scale processes.

4. Chemical Kinetics Factors That Affect Reaction Rates Concentration of Reactants  As the concentration of reactants increases, so does the likelihood that reactant molecules will collide. Temperature  At higher temperatures, reactant molecules have more kinetic energy, move faster, and collide more often and with greater energy. Catalysts  Speed rxn by changing mechanism.

5. Chemical Kinetics Reaction Rates Rates of reactions can be determined by monitoring the change in concentration of either reactants or products as a function of time.  [A] vs  t, unints mol/L s Rxn Movie

6. Chemical Kinetics Reaction Rates In this reaction, the concentration of butyl chloride, C 4 H 9 Cl, was measured at various times, t. C 4 H 9 Cl (aq) + H 2 O (l)  C 4 H 9 OH (aq) + HCl (aq) [C 4 H 9 Cl] M

7. Chemical Kinetics Reaction Rates The average rate of the reaction over each interval is the change in concentration divided by the change in time: C 4 H 9 Cl (aq) + H 2 O (l)  C 4 H 9 OH (aq) + HCl (aq) Average Rate, M/s

8. Chemical Kinetics Reaction Rates Note that the average rate decreases as the reaction proceeds. This is because as the reaction goes forward, there are fewer collisions between reactant molecules. C 4 H 9 Cl (aq) + H 2 O (l)  C 4 H 9 OH (aq) + HCl (aq)

9. Chemical Kinetics Reaction Rates A plot of concentration vs. (versus) time for this reaction yields a curve like this. The slope of a line tangent to the curve at any point is the instantaneous rate at that time. C 4 H 9 Cl (aq) + H 2 O (l)  C 4 H 9 OH (aq) + HCl (aq)

10. Chemical Kinetics Reaction Rates The reaction slows down with time because the concentration of the reactants decreases. C 4 H 9 Cl (aq) + H 2 O (l)  C 4 H 9 OH (aq) + HCl (aq)

11. Chemical Kinetics Reaction Rates and Stoichiometry In this reaction, the ratio of C 4 H 9 Cl to C 4 H 9 OH is 1:1. Thus, the rate of disappearance of C 4 H 9 Cl is the same as the rate of appearance of C 4 H 9 OH. C 4 H 9 Cl (aq) + H 2 O (l)  C 4 H 9 OH (aq) + HCl (aq) Rate = -  [C 4 H 9 Cl]  t =  [C 4 H 9 OH]  t

12. Chemical Kinetics Reaction Rates and Stoichiometry What if the ratio is not 1:1? H 2 (g) + I 2 (g)  2 HI (g) Only 1/2 HI is made for each H 2 used.

13. Chemical Kinetics Reaction Rates and Stoichiometry To generalize, for the reaction aA + bBcC + dD Reactants (decrease) Products (increase)

14. Chemical Kinetics Concentration and Rate Each reaction has its own equation that gives its rate as a function of reactant concentrations.  this is called its Rate Law To determine the rate law we measure the rate at different starting concentrations.

15. Chemical Kinetics Concentration and Rate Compare Experiments 1 and 2: when [NH 4 + ] doubles, the initial rate doubles.

16. Chemical Kinetics Concentration and Rate Likewise, compare Experiments 5 and 6: when [NO 2 - ] doubles, the initial rate doubles.

17. Chemical Kinetics Concentration and Rate This equation is called the rate law, and k is the rate constant.

18. Chemical Kinetics Rate Laws A rate law shows the relationship between the reaction rate and the concentrations of reactants. The rate of a reaction is directly proportional to the concentration of the reactants.  For gas-phase reactants use P A instead of [A]. k is a constant that has a specific value for each reaction. The value of k is determined experimentally. “Constant” is relative here - k is unique for each rxn k changes with T

19. Chemical Kinetics The dependence of the reaction rates on concentration may be expressed in terms of order of a reaction. To calculate the order of a reaction, we must carry out the experiments to find out the dependence of reaction rate on the concentration of each individual reactant at a given temperature. Moreover, the concentration dependence of various reactants, as we have seen, is given by the rate law. Thus, the order of a reaction is defined as: the sum of the power to which the concentration terms are raised in the rate law expression.

20. Chemical Kinetics Rate Laws Exponents tell the order of the reaction with respect to each reactant. This reaction is First-order in [NH 4 + ] First-order in [NO 2 − ] The overall reaction order can be found by adding the exponents on the reactants in the rate law. This reaction is second-order overall.

21. Chemical Kinetics Integrated Rate Laws Consider a simple 1st order rxn: A  B How much A is left after time t? Integrate: Differential form:

22. Chemical Kinetics Integrated Rate Laws The integrated form of first order rate law: Can be rearranged to give: [A] 0 is the initial concentration of A (t=0). [A] t is the concentration of A at some time, t, during the course of the reaction.

23. Chemical Kinetics Integrated Rate Laws Manipulating this equation produces… …which is in the form y = mx + b

24. Chemical Kinetics First-Order Processes If a reaction is first-order, a plot of ln [A] t vs. t will yield a straight line with a slope of -k. So, use graphs to determine rxn order.

25. Chemical Kinetics First-Order Processes Consider the process in which methyl isonitrile is converted to acetonitrile. CH 3 NCCH 3 CN How do we know this is a first order rxn?

26. Chemical Kinetics First-Order Processes This data was collected for this reaction at 198.9°C. CH 3 NCCH 3 CN Does rate=k[CH 3 NC] for all time intervals?

27. Chemical Kinetics First-Order Processes When ln P is plotted as a function of time, a straight line results.  The process is first-order.  k is the negative slope: 5.1  10 -5 s -1.

28. Chemical Kinetics Second-Order Processes Similarly, integrating the rate law for a process that is second-order in reactant A: also in the form y = mx + b Rearrange, integrate:

29. Chemical Kinetics Second-Order Processes So if a process is second-order in A, a plot of 1/[A] vs. t will yield a straight line with a slope of k. If a reaction is first-order, a plot of ln [A] t vs. t will yield a straight line with a slope of -k. First order:

30. Chemical Kinetics Determining rxn order The decomposition of NO 2 at 300°C is described by the equation NO 2 (g) NO (g) + 1/2 O 2 (g) and yields these data: Time (s)[NO 2 ], M 0.00.01000 50.00.00787 100.00.00649 200.00.00481 300.00.00380

31. Chemical Kinetics Graphing ln [NO 2 ] vs. t yields: Time (s)[NO 2 ], Mln [NO 2 ] 0.00.01000-4.610 50.00.00787-4.845 100.00.00649-5.038 200.00.00481-5.337 300.00.00380-5.573 The plot is not a straight line, so the process is not first-order in [A]. Determining rxn order Does not fit:

32. Chemical Kinetics Second-Order Processes A graph of 1/[NO 2 ] vs. t gives this plot. Time (s)[NO 2 ], M1/[NO 2 ] 0.00.01000100 50.00.00787127 100.00.00649154 200.00.00481208 300.00.00380263 This is a straight line. Therefore, the process is second-order in [NO 2 ].

33. Chemical Kinetics Determination of the reaction order by a graphics method ln [A] t The first order 1/[A] t The second order Zero-order t [A] t 1/[A] 2 The third order

34. Chemical Kinetics Determination methods of the reaction order. The substitution method of kinetic equalizations C 0 or [A] 0 – initial (first) values of the concentration of the initial substance A C or [A]– the concentration of initial substance A in a specific time

35. Chemical Kinetics Half-Life Half-life is defined as the time required for one-half of a reactant to react. Because [A] at t 1/2 is one-half of the original [A], [A] t = 0.5 [A] 0.

36. Chemical Kinetics Half-Life For a first-order process, set [A] t =0.5 [A] 0 in integrated rate equation: NOTE: For a first-order process, the half-life does not depend on [A] 0.

37. Chemical Kinetics Half-Life- 2nd order For a second-order process, set [A] t =0.5 [A] 0 in 2nd order equation.

38. Chemical Kinetics Determination of the reaction order according to the half-lives The first order: The 2-d order: The 3-dorder: Zero order:

39. Chemical Kinetics Outline: Kinetics First orderSecond order Rate Laws Integrate d Rate Laws complicated Half-life complicated

40. Chemical Kinetics Reaction Mechanisms Reactions may occur all at once or through several discrete steps. Each of these processes is known as an elementary reaction or elementary process.

41. Chemical Kinetics 41 Classification of chemical reactions according to the quantity of stages (phases). Simple reactions go in a one elementary chemical act Compositum reactions go in several stages Double-sided : А В Parallel: В A C In consecutive order: А→В→С Conjugating:А D С В Е

42. Chemical Kinetics Classification of the chemical reaction according to the quantity of the reacting phases Homogeneous: N 2 (g) + H 2 (g) → NH 3 (g) Heterogeneous: Mg (s) + HCl (l) → MgCl 2 (l) + H 2 g) Topochemical ( in the hard phase)

43. Chemical Kinetics Reaction Mechanisms The molecularity of a process tells how many molecules are involved in the process. The rate law for an elementary step is written directly from that step.

44. Chemical Kinetics Unimolecular: Н 2 СО 3 → Н 2 О + СО 2 Bimolecular: CuO + CO → Cu + CO 2 Termolecular: 2 NO + O 2 = 2 NO 2

45. Chemical Kinetics Multistep Mechanisms In a multistep process, one of the steps will be slower than all others. The overall reaction cannot occur faster than this slowest, rate-determining step.

46. Chemical Kinetics Slow Initial Step The rate law for this reaction is found experimentally to be Rate = k [NO 2 ] 2 CO is necessary for this reaction to occur, but the rate of the reaction does not depend on its concentration. This suggests the reaction occurs in two steps. NO 2 (g) + CO (g)  NO (g) + CO 2 (g)

47. Chemical Kinetics Slow Initial Step A proposed mechanism for this reaction is Step 1: NO 2 + NO 2  NO 3 + NO (slow) Step 2: NO 3 + CO  NO 2 + CO 2 (fast) The NO 3 intermediate is consumed in the second step. As CO is not involved in the slow, rate-determining step, it does not appear in the rate law.

48. Chemical Kinetics Fast Initial Step The rate law for this reaction is found (experimentally) to be Because termolecular (= trimolecular) processes are rare, this rate law suggests a two-step mechanism.

49. Chemical Kinetics Fast Initial Step A proposed mechanism is Step 1 is an equilibrium- it includes the forward and reverse reactions.

50. Chemical Kinetics Fast Initial Step The rate of the overall reaction depends upon the rate of the slow step. The rate law for that step would be But how can we find [NOBr 2 ]?

51. Chemical Kinetics Fast Initial Step NOBr 2 can react two ways:  With NO to form NOBr  By decomposition to reform NO and Br 2 The reactants and products of the first step are in equilibrium with each other. Therefore, Rate f = Rate r

52. Chemical Kinetics Fast Initial Step Because Rate f = Rate r, k 1 [NO] [Br 2 ] = k −1 [NOBr 2 ] Solving for [NOBr 2 ] gives us k1k−1k1k−1 [NO] [Br 2 ] = [NOBr 2 ]

53. Chemical Kinetics Fast Initial Step Substituting this expression for [NOBr 2 ] in the rate law for the rate-determining step gives

54. Chemical Kinetics