Chemistry 102(01) Spring 2012

1. Chemistry 102(01) Spring 2012 CTH 328 9:30-10:45 am Instructor: Dr. UpaliSiriwardane e-mail: upali@latech.edu Office: CTH 311 Phone 257-4941 Office Hours: M,W 8:00-9:00 & 11:00-12:00 am; Tu,Th,F8:00 - 10:00 am.. Exams: 9:30-10:45 am, CTH 328. March 21 ,2012 (Test 1): Chapter 13 April 18 , 2012 (Test 2): Chapter 14 &15 May 14 , 2012 (Test 3):Chapter 16 &18 Optional Comprehensive Final Exam: May 17, 2012 : Chapters 13, 14, 15, 16, 17, and 18

2. Chapter 13. Chemical Kinetics 13.1 Reaction Rate 13.2 Effect of Concentration on Reaction Rate 13.3 Rate Law and Order of Reaction 13.4 A Nanoscale View: Elementary Reactions 13.5 Temperature and Reaction Rate: The Arrhenius Equation 13.6 Rate Laws for Elementary Reactions 13.7 Reaction Mechanisms 13.8 Catalysts and Reaction Rate 13.9 Enzymes: Biological Catalysts 13‑10 Catalysis in Industry

3. How do you measure rates? Rates are related to the time it required to decay reactants or form products. The rate reaction = change in concentration of reactants/products per unit time Average rate rate of reaction = –D[reactant]/Dt Instantaneous rate rate of reaction = – d[reactant]/dt

4. Rate of Appearance & Disappearance • 2 N2O5(g) -----> 4 NO2 (g) + O2 (g) • Disappearance is based on reactants • rate = -(D[N2O5]/ D t • Appearance is based on products • rate = D[NO2]/ D t • rate = D[O2]/ D t • Converting rates of Appearance. • rate = (D[NO2]/ D t = - 4/2 D[N2O5]/ D t • D[O2]/ D t = - 1/2 D[N2O5]/ D t

5. Measuring Rate a A --> b B Based on reactants rate = -(1/a) D[A]/ D t Based on products rate = +(1/b) D[B]/ D t D[A]= [A]f - [A]I Change in A D t= tf - ti Change in t

6. Reaction of cis-platin with Water

7. Disappearance of Color

8. An example reaction where gas is produced Gas buret Constant temperature bath

9. Time vs. volume of gas • Time (s) Volume STP O2, mL • 0 0 • 300 1.15 • 600 2.18 • 900 3.11 • 1200 3.95 • 1800 5.36 • 2400 6.50 • 3000 7.42 • 4200 8.75 • 5400 9.62 • 6600 10.17 • 7800 10.53 Here are the results for our experiment.

10. 2 N2O5(g) -----> 4 NO2 (g) + O2 (g)

11. Graphof 2 N2O5(g) ---> 4 NO2 (g) + O2 (g)

12. Graph

13. Factors that affect rates of chemical reactions • a) Temperature • b) Concentration • c) Catalysts • d) Particle size of solid reactants

14. Effect of Particle Size on Rate

15. Chemical Kinetics Definitions and Concepts a) rate law b) rate constant c) order d) differential rate law c) integral rate law

16. Every chemical reaction has a Rate Law The rate law is an expression that relates the rate of a chemical reaction to a constant (rate constant-k) and concentration of reactants raised to a power. The power of a concentration is called the order with respect to a particular reactant. Rate Law

17. Rate Law E.g. A + B -----> C rate a [A]l[B]m rate = k [A]l[B]m; k = rate constant [A] = concentration of A [B] = concentration of B l = order with respect to A m = order with respect to B l & m have nothing to do with stoichiometric coefficients

18. Rate Constant E.g. A + B -----> C rate a [A]l[B]m rate = k [A]l[B]m; k = rate constant proportionality constant of the rate law Larger the k faster the reaction It is related inversely to t½

19. Decomposition Reaction

20. Rate Law E.g. 2 N2O5(g) -----> 4 NO2 (g) + O2 (g) rate a [N2O5]1 rate = k [N2O5]1 ;k = rate constant [N2O5] = concentration of N2O5 1 = order with respect to N2O5 Rate and the order are obtained by experiments

21. Order The power of the concentrations is the order with respect to the reactant. E.g. A + B -----> C If rate law: rate = k [A]1[B]2 The order of the reaction with respect to A is one (1). The order of the reaction with respect to B is two (2). Overall order of a chemical reaction is equal to the sum of all orders (3).

22. Finding rate laws • Method of initial rates • The order for each reactant is found by: • Changing the initial concentration of that reactant. • Holding all other initial concentrations and conditions constant. • Measuring the initial rates of reaction • The change in rate is used to determine the order for thatspecific reactant. The process is repeated for each reactant.

23. Initial rate

24. How do you find order? A + B -----> C rate = k [A]l[B]m; Hold concentration of other reactants constant If [A] doubled, rate doubled 1st order, [2A]1 = 2 1 x [A]1 , 2 1 = 2 b) If [A] doubled, rate quadrupled 2nd order, [2A]2 = 2 2 x [A]2 , 2 2 = 4 c) If [A] doubled, rate increased 8 times 3rd order, [2A]3 = 2 3 x [A]3 , 2 3 = 8

25. Rate data

26. Determining order

27. DeterminingK, Rate Constant

28. Overall order

29. Units of the Rate Constant (k) 1 first order: k =───=s-1 s L second order k =─── mol s L2 third order k =─── mol2 s

30. First order reactions

31. Rate Law Differential Rate Law Integral Rate rate = k [A]0 -D [A]/Dt =k ; ([A]0=1) [A]f-[A]i = -kt rate = k [A]1-D [A]/Dt = k [A] ln [A]o/[A]t = kt rate = k [A]2 -D [A]/Dt = k [A]2 1/ [A]f = kt + 1/[A]i Differential and Integral Rate Law

32. Integrated Rate Laws

33. 1 [A]0 Graphical method 1 [A]t 1 [A]t

34. Graphical Ways to get Order

35. First-order, Second-order,and Zeroth-order Plots

36. Finding rate laws 0 order plot 2nd order plot [N2O5] 1/[N2O5] Time (s) Time (s) As you can see from these plots of the N2O5 data, only a first order plot results in a straight line. Time (s) ln[N2O5] 1st order plot

37. Comparing graphs This plot of ln[cis-platin] vs. time produces a straight line, suggesting that the reaction is first-order.

38. First Order Reactions A ----> B

39. t1/2 equation 0.693 = k t1/2 0.693 t1/2 = ---- k

40. Half-life The half-life and the rate constant are related. t1/2 = Half-life can be used to calculate the first order rate constant. For our N2O5 example, the reaction took 1900 seconds to react half way so: k = = = 3.65 x 10-4 s-1 0.693 k 0.693 t1/2 0.693 1900 s

41. A Nanoscale View:Elementary Reactions Most reactions occur through a series of simple steps or elementary reactions. Elementary reactions could be unimolecular - rearrangement of a molecule bimolecular - reaction involving the collision of two molecules or particles termolecular - reaction involving the collision of three molecules or particles

42. Elementary Reactions and Mechanism 2NO2 (g) + F2(g) 2NO2F (g) If the reaction took place in a single step the rate law would be: rate = k [NO2]2 [F2] Observed: rate = k1 [NO2] [ F2] If the observed rate law is not the same as if the reaction took place in a single step that more than one step must be involved

43. Elementary Reactions A possible reaction mechanism might be: Step one NO2 + F2 NO2F + F (slow) Step two NO2 + F NO2F (fast) Overall 2NO2 + F2 2NO2F slowest step in a multi-step mechanism the step which determines the overall rate of the reaction rate = k1 [NO2] [ F2] Rate Determining Step

44. Reaction profile of rate determining step This type of plot shows the energy changes during a reaction. Potential Energy activation energy DH Reaction coordinate

45. What Potential Energy Curves Show Exothermic Reactions Endothermic Reactions Activation Energy (Ea) of reactant or the minimum energy required to start a reaction Effect of catalysts Effect of temperature

46. Examples of reaction profiles Exothermic reaction Endothermic reaction

47. Examples of reaction profiles High activation energy (kinetic) Low heat of reaction (thermodynamic) Low activation energy (kinetic) High heat of reaction (thermodynamic)

48. Unimolecular Reaction cis-2-butene trans-2-butrne

49. Bimolecular Reaction I- + CH3Br ICH3 + Br-

50. Orientation Probability: Some Unsuccessful Collisions I- + CH3Br ICH3 + Br-