1. Chapter 15 - Chemical Kinetics Objectives: 1.Determine rates of reactions from graphs of concentration vs. time. 2.Recall the conditions which affect the rates. 3.Recognize the order of reaction, give the rate equation, calculate the rate constant. 4.Use integrated rate laws and half-life equations in calculations. 5.Draw energy diagrams, find activation energy. 6.Identify catalysts and their properties. 7.Identify reaction intermediates. 8.Recognize the rate equation given a mechanism, and given the rate equation determine the mechanism.

2. KINETICS — the study of REACTION RATES H 2 O 2 decomposition in an insect H 2 O 2 decomposition catalyzed by MnO 2 Chemical engineering, enzymology, environmental engineering, etc.

3. Kinetics and Mechanisms KINETICS — the study of REACTION RATES and their relation to the way the reaction proceeds, i.e., its MECHANISM.KINETICS — the study of REACTION RATES and their relation to the way the reaction proceeds, i.e., its MECHANISM. The reaction mechanism is our goal!The reaction mechanism is our goal! The sequence of events at the molecular level that control the speed and outcome of a reaction.The sequence of events at the molecular level that control the speed and outcome of a reaction.

4. Br from biomass burning destroys stratospheric ozone. (See R.J. Cicerone, Science, volume 263, page 1243, 1994.) Step 1:Br + O 3 ---> BrO + O 2 Step 2:Cl + O 3 ---> ClO + O 2 Step 3:BrO + ClO + light ---> Br + Cl + O 2 NET: 2 O 3 ---> 3 O 2 Identify the intermediate Identify the catalyst(s)

5. Reaction Rates Reaction rate = change in concentration of a reactant or product with time.Reaction rate = change in concentration of a reactant or product with time. Change in distance over a period of time:  distance  time  concentration  time Rate = Three “types” of rates: Three “types” of rates: initial rate initial rate average rate average rate instantaneous rate instantaneous rate

6. Reaction Rates can be determined from a Plot Blue dye is oxidized with bleach. Its concentration decreases with time. The rate — the change in dye conc with time — can be determined from the plot. The initial rate (in this case over the first minute) is calculated from a tangent line crossing the initial concentration. Then the slope of the line is determined:

7. Reaction Rates can be determined from a Plot N 2 O 5  NO 2 + O 2 24 The average rate is calculated from a time interval. The instantaneous rate is calculated at a single point in time (or given concentration) by drawing a tangent line crossing the point. Then the slope of the line is determined. Compare average rates at the beginning and end of reaction.

8. Factors affecting the Rates ConcentrationsConcentrations Physical State of Reactants and ProductsPhysical State of Reactants and Products Surface areaSurface area TemperatureTemperature CatalystsCatalysts Review Exp. 1: Factors affecting reactions rates

9. Concentration Increasing the concentration of reactants _______________ the rate of the reaction. 0.3 M HCl6 M HCl Mg(s) + 2 HCl(aq) ---> MgCl 2 (aq) + H 2 (g)

10. Surface Area Increasing the surface area of reactants _____________ the reaction rate.

11. Temperature Increasing the temperature ____________ the rate of the reaction. Bleach at 54 ˚CBleach at 22 ˚C

12. Catalysts A _________ is present at the beginning and at the end of the reaction and it does not change; but it _______________ the rate of the reaction. 2 H 2 O 2 ---- > 2 H 2 O + O 2 MnO 2

13. Factors affecting Reaction Rates Iodine clock reaction: 1. Iodide is oxidized to iodine H 2 O 2 + 2 I - + 2 H + -----> 2 H 2 O + I 2 2. I 2 reduced to I - with vitamin C I 2 + C 6 H 8 O 6 ----> C 6 H 6 O 6 + 2 H + + 2 I - When all vitamin C is depleted, the I 2 interacts with starch to give a blue complex.

14. Factors affecting Reaction Rates

15. Concentration and Rate To postulate a mechanism we study: - The reaction rate and its - Concentration dependence. Generate a Rate Law Equation.

16. Rate Laws In general for: a A + b B --> x X Rate = k [A] m [B] n The exponents m, n are the _______________are the _______________ can be 0, 1, 2 or fractionscan be 0, 1, 2 or fractions must be determined by ____________!must be determined by ____________! With a catalyst C [C] p

17. Interpreting Rate Laws Rate = k [A] m [B] n [C] p If m = 1, rxn. is 1st order in AIf m = 1, rxn. is 1st order in A Rate = k [A] 1 If [A] doubles, then rate goes up by factor of __________ If [A] doubles, then rate goes up by factor of __________ If m = 2, rxn. is 2nd order in A.If m = 2, rxn. is 2nd order in A. Rate = k [A] 2 Rate = k [A] 2 Doubling [A] increases rate by _____________ If m = 0, rxn. is zero order.If m = 0, rxn. is zero order. Rate = k [A] 0 Rate = k [A] 0 If [A] doubles, rate ____________________

18. Deriving Rate Laws Derive rate law and k for CH 3 CHO(g) --> CH 4 (g) + CO(g) from experimental data for rate of disappearance of CH 3 CHO Expt. [CH 3 CHO] Disappear of CH 3 CHO (mol/L) (mol/Lsec) 10.100.020 20.200.081 30.300.182 40.400.318 Look at exp 1 and 2: concentration doubles; rate cuadruples Rate = k [CH 3 CHO] n 4 Rate = k [2 CH 3 CHO] n n = 2

19. Deriving Rate Laws Rate of rxn = Here the rate goes up by ________ when initial concentration doubles. Therefore, we say this reaction is __________ order. Now determine the value of k. Use any exp. Data: Using k you can calculate rate at other values of [CH 3 CHO] at same T.

20. Concentration and Time What is the concentration of reactant as a function of time? Consider First Order reactions: Integrating we get:

21. Integrated First Order Law [A] / [A] 0 =fraction remaining after time t has elapsed.

22. The decomposition of a certain insecticide in water follows first-order kinetics with a rate constant of 1.45 year -1 at 12 o C. A quantity of this insecticide is washed into a lake on June 1, leading to a concentration of 5.0 x 10 -7 g/cm 3. Assume that the average temperature of the lake is 12 o C. a) What is the concentration of the insecticide on June 1 of the following year? b) How long will it take for the concentration of the insecticide to drop to 3.0 x 10 -7 g/cm 3 ?

23. Using Integrated Rate Laws All 1st order reactions have straight line plot for ln [A] vs. time. And 2nd order gives straight line for plot of 1/[A] vs. time.

24. Using Integrated Rate Laws In an experiment for: 2 N 2 O 5 (g) ---> 4 NO 2 (g) + O 2 (g) Time (min)[N 2 O 5 ] 0 (M) 01.00 1.00.705 2.00.497 5.00.173 Data of conc. vs. time plot do not fit straight line. If it were zero order:

25. Using Integrated Rate Laws Plot of ln [N 2 O 5 ] vs. time is a straight line! ln [N 2 O 5 ] 0 0-0.35-0.70-1.75 In an experiment for: 2 N 2 O 5 (g) ---> 4 NO 2 (g) + O 2 (g) Time (min)[N 2 O 5 ] 0 (M) 01.00 1.00.705 2.00.497 5.00.173 If it were first order: Calculate the ln :

26. Using Integrated Rate Laws Plot of ln [N 2 O 5 ] vs. time is a straight line! Eqn. for straight line: y = mx + b y = mx + b Plot of ln [N 2 O 5 ] vs. time is a straight line! Eqn. for straight line: y = mx + b y = mx + b

27. The gas phase decomposition of hydrogen peroxide at 400 o C is second order in H 2 O 2. In one experiment, when the initial concentration of H 2 O 2 was 0.246 M, the concentration of H 2 O 2 dropped to 3.39 x 10 -2 M after 25.9 seconds had passed. What is the rate constant for the reaction? 2 H 2 O 2  2 H 2 O + O 2

28. Half-Life HALF-LIFE is the time it takes for 1/2 a sample is disappear. For 1st order reactions, the concept of HALF-LIFE is especially useful.

29. Half-Life Reaction is 1st order decomposition of H 2 O 2.Reaction is 1st order decomposition of H 2 O 2. Reaction after 1 half-life.Reaction after 1 half-life. 1/2 of the reactant has been consumed and 1/2 remains.1/2 of the reactant has been consumed and 1/2 remains.

30. Half-Life After 2 half- lives _____ of the reactant remains.After 2 half- lives _____ of the reactant remains. After 3 half- lives ____ of the reactant remains.After 3 half- lives ____ of the reactant remains.

31. Half-Life [A] / [A] 0 = fraction remaining when t = t 1/2 then fraction remaining = _________ ln (____) = - k t 1/2 ln [R] / [R] 0 = k t

32. In an experiment, it is determined that 75% of a sample of HCO 2 H (formic acid) has decomposed in 72 seconds following first-order kinetics. Determine t 1/2 for this reaction. HCO 2 H  CO 2 + H 2 ln [R] / [R] 0 = k tt 1/2 = 0.693 / k

33. Mechanisms Mechanism: how reactants are converted to products at the molecular level.Mechanism: how reactants are converted to products at the molecular level. RATE LAW ----> MECHANISM experiment ---->theory

34. Activation Energy Molecules need a minimum amount of energy to react. Visualized as an energy barrier - activation energy, E a. Reaction coordinate diagram

35. Activation Energy Conversion of cis to trans-2-butene requires twisting around the C=C bond.Conversion of cis to trans-2-butene requires twisting around the C=C bond. Rate = k [trans-2-butene]Rate = k [trans-2-butene]

36. Transition State Activation energy barrier CisTrans Transition state

37. Mechanisms Reaction passes thru a where there is an that has sufficient energy to become a product.Reaction passes thru a TRANSITION STATE where there is an activated complex that has sufficient energy to become a product. ACTIVATION ENERGY, E a = energy req’d to form activated complex. Here E a = ___________

38. Mechanisms Also note that trans-butene is MORE STABLE than cis-butene by about 4 kJ/mol. Therefore, cis ---> trans is _________________ This is the connection between thermo- dynamics and kinetics.

39. Effect of Temperature Reactions generally occur slower at lower T. Iodine clock reaction, book page 705. H 2 O 2 + 2 I - + 2 H + --> 2 H 2 O + I 2 Room temperature In ice at 0 o C

40. Activation Energy and Temperature Reactions are __________ at a higher T because a larger fraction of reactant molecules have enough energy to convert to product molecules. In general, differences in activation energy cause reactions to vary from fast to slow.

41. Collision Theory Molecules must collide with one another Molecules must collide with sufficient ________ to break bonds Molecules must collide in an orientation that can lead to rearrangement of the atoms.

42. Arrhenius Equation Rate constant Temp (K) 8.31 x 10 -3 kJ/Kmol Activation energy Frequency factor Frequency factor related to frequency of collisions with correct geometry. Plot ln k vs. 1/T ---> straight line. slope = -E a / R slope = -E a / R

43. Collision Theory Reactions require (a) activation energy and (b) correct geometry. O 3 (g) + NO(g) ---> O 2 (g) + NO 2 (g) O 3 (g) + NO(g) ---> O 2 (g) + NO 2 (g) 2. Activation energy and geometry 1. Activation energy

44. Mechanisms Most reactions involve a sequence of elementary steps. 2 I - + H 2 O 2 + 2 H + ---> I 2 + 2 H 2 O 2 I - + H 2 O 2 + 2 H + ---> I 2 + 2 H 2 O Rate = k [I - ] [H 2 O 2 ] Rate = k [I - ] [H 2 O 2 ]NOTE 1.Rate law comes from experiment. 2.Order and stoichiometric coefficients not necessarily the same! 3.Rate law reflects all chemistry down to and including the slowest step in multistep reaction.

45. Mechanisms Proposed Mechanism: Step 1 — slowHOOH + I - --> HOI + OH - Step 2 — fastHOI + I - --> I 2 + OH - Step 3 — fast2 OH - + 2 H + --> 2 H 2 O Rate of the reaction controlled by slow step — RATE DETERMINING STEP, Rate can be no faster than RDS! Most reactions involve a sequence of elementary steps. 2 I - + H 2 O 2 + 2 H + ---> I 2 + 2 H 2 O 2 I - + H 2 O 2 + 2 H + ---> I 2 + 2 H 2 O Rate = k [I - ] [H 2 O 2 ] Rate = k [I - ] [H 2 O 2 ]

46. Mechanisms Elementary Step 1 is bimolecular and involves I - and HOOH. Therefore, this predicts the rate law should be Rate  [I - ] [H 2 O 2 ] — as observed!! The species HOI and OH - are ________________. Proposed Mechanism: Step 1 — slowHOOH + I - --> HOI + OH - Step 2 — fastHOI + I - --> I 2 + OH - Step 3 — fast2 OH - + 2 H + --> 2 H 2 O Most reactions involve a sequence of elementary steps. 2 I - + H 2 O 2 + 2 H + ---> I 2 + 2 H 2 O 2 I - + H 2 O 2 + 2 H + ---> I 2 + 2 H 2 O Rate = k [I - ] [H 2 O 2 ] Rate = k [I - ] [H 2 O 2 ]

47. Catalysts and Activation Energy Uncatalyzed reaction Catalyzed reaction MnO 2 catalyzes decomposition of H 2 O 2 2 H 2 O 2 ---> 2 H 2 O + O 2

48. Catalysts and Activation Energy Iodine-Catalyzed Isomerization of Cis-2- Butene

49. Remember Go over all the contents of your textbook. Practice with examples and with problems at the end of the chapter. Practice with OWL tutor. Practice with the quiz on CD of Chemistry Now. Work on your OWL assignment for Chapter 15.