1. Chemical Kinetics In kinetics we study 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). © 2012 Pearson Education, Inc.

2. Chemical Kinetics The Collision Theory In a chemical reaction, bonds are broken(endothermic) and new bonds are formed (exothermic). There are three conditions that need to be met in order for a reaction to take place © 2012 Pearson Education, Inc.

3. Chemical Kinetics The Three Conditions of the Collision Theory 1. The reactants need to come into contact (collide) © 2012 Pearson Education, Inc.

4. Chemical Kinetics The Three Conditions of the Collision Theory 2. The collision must occur with a certain minimum energy to form an activated complex (activation energy E act ) Must have enough energy to overcome the electron/electron repulsion of the valence shell electrons of the reacting species © 2012 Pearson Education, Inc.

5. Chemical Kinetics Activation Energy In other words, there is a minimum amount of energy required for reaction: the activation energy, E a. Just as a ball cannot get over a hill if it does not roll up the hill with enough energy, a reaction cannot occur unless the molecules possess sufficient energy to get over the activation-energy barrier. © 2012 Pearson Education, Inc.

6. Chemical Kinetics Reaction Coordinate Diagrams The diagram shows the energy of the reactants and products (and, therefore,  E). The high point on the diagram is the transition state. The species present at the transition state is called the activated complex. The energy gap between the reactants and the activated complex is the activation-energy barrier. © 2012 Pearson Education, Inc.

7. Chemical Kinetics The Three Conditions of the Collision Theory 3. The molecules that are colliding must have the correct orientation © 2012 Pearson Education, Inc.

8. Chemical Kinetics Proper orientation

9. Chemical Kinetics The Collision Model Furthermore, molecules must collide with the correct orientation and with enough energy to cause bond breakage and formation. © 2012 Pearson Education, Inc.

10. Chemical Kinetics The Collision Model Furthermore, molecules must collide with the correct orientation and with enough energy to cause bond breakage and formation. © 2012 Pearson Education, Inc.

11. Chemical Kinetics 5 Factors that affect Reaction Rate

12. Chemical Kinetics 5 Factors That Affect Reaction Rates 1. Concentration of reactants. –As the concentration of reactants increases, so does the likelihood that reactant molecules will collide. © 2012 Pearson Education, Inc.

13. Chemical Kinetics 5 Factors That Affect Reaction Rates 2. Nature of the reactants –Some reactant molecules react in a hurry, others very slowly. –Physical state– gasoline (l) vs gasoline (g) or K 2 SO 4 (s) + Ba(NO 3)(S)  no rxn, but they will react when aqueous –Chemical identity: What is reacting? Oppositely charged ions react very rapidly. Or metallic sodium reacts more rapidly than metallic magnesium. Or nature of bonds in reactants © 2012 Pearson Education, Inc.

14. Chemical Kinetics 4 Factors That Affect Reaction Rates 3. Temperature –At higher temperatures, reactant molecules have more kinetic energy, move faster, and collide more often and with greater energy. –As a guideline in many reactions a 10 C rise in temp will double reaction rate © 2012 Pearson Education, Inc.

15. Chemical Kinetics Maxwell–Boltzmann Distributions Temperature is defined as a measure of the average kinetic energy of the molecules in a sample. At any temperature there is a wide distribution of kinetic energies. © 2012 Pearson Education, Inc.

16. Chemical Kinetics Maxwell–Boltzmann Distributions As the temperature increases, the curve flattens and broadens. Thus, at higher temperatures, a larger population of molecules has higher energy., and the particles have the activation energy on collision This allows for a greater number of successful collisions. © 2012 Pearson Education, Inc.

17. Chemical Kinetics 5 Factors That Affect Reaction Rates 4. Presence of a catalyst. –Catalysts speed up reactions by changing the pathway (place reactants in proper orientation—”plays matchmaker”) of the reaction. –Catalysts are not consumed during the course of the reaction. –lowers the activation energy of a reaction © 2012 Pearson Education, Inc.

18. Chemical Kinetics Catalysts Catalysts increase the rate of a reaction by decreasing the activation energy of the reaction. Catalysts change the mechanism by which the process occurs. The  H Remains unchanged © 2012 Pearson Education, Inc.

19. Chemical Kinetics

20. Chemical Kinetics Catalysts One way a catalyst can speed up a reaction is by holding the reactants together and helping bonds to break. © 2012 Pearson Education, Inc.

21. Chemical Kinetics 5 Factors That Affect Reaction Rates 5. Surface Area of the reactants- Greater surface area, increases the chance of collision—more speed

22. 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. © 2012 Pearson Education, Inc.

23. Chemical Kinetics Rate= change in concentration of a species Time interval  [A]  t Rate of  reactant is always negative Rate of  product is always positive

24. Chemical Kinetics Reaction Rates In this reaction, the concentration of butyl chloride, C 4 H 9 Cl, was measured at various times. The speed of a reaction is expressed in terms of “its rate”, some measurable quantity changes with time © 2012 Pearson Education, Inc.

25. Chemical Kinetics Reaction Rates The average rate of the reaction over each interval is the change in concentration [Molarity] 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 =  [C 4 H 9 Cl]  t © 2012 Pearson Education, Inc.

26. Chemical Kinetics Reaction Rates C 4 H 9 Cl(aq) + H 2 O(l)  C 4 H 9 OH(aq) + HCl(aq) Is average rate constant?? : © 2012 Pearson Education, Inc.

27. 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) © 2012 Pearson Education, Inc.

28. Chemical Kinetics Reaction Rates A plot of [C 4 H 9 Cl] 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) © 2012 Pearson Education, Inc.

29. Chemical Kinetics Reaction Rates All reactions slow down over time. Therefore, the best indicator of the rate of a reaction is the instantaneous rate near the beginning of the reaction. C 4 H 9 Cl(aq) + H 2 O(l)  C 4 H 9 OH(aq) + HCl(aq) © 2012 Pearson Education, Inc.

30. Chemical Kinetics Reaction Rates 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 © 2012 Pearson Education, Inc.

31. Chemical Kinetics Graphing data of an experiment will show average reaction rate. Instantaneous rate = slope (rise/run) of line tangent of to the curve at that point!!

32. Chemical Kinetics Example Problem Examine the graph above and calculate the AVERAGE rate at which [NO 2 ] changes in the first 50 seconds What is the instantaneous rate and 600 seconds?

33. Chemical Kinetics Relative Reaction Rates We can consider the appearance of product along with the disappearance of reactants. What if the ratio is not 1:1?? The reactants concentration is declining while the products’ are increasing 2 HI(g)  H 2 (g) + I 2 (g) © 2012 Pearson Education, Inc.

34. Chemical Kinetics Reaction Rates and Stoichiometry We look at the Relative Rates of Reaction 2 HI(g)  H 2 (g) + I 2 (g) =  1212  [HI] = 1111  [H 2 ]  t 1111   = © 2012 Pearson Education, Inc. In such a case tt tt

35. Chemical Kinetics Relative Reaction Rates and Stoichiometry To generalize, then, for the reaction Make sure reactants have neg sign and watch stoich aA + bBcC + dD Rate =  1a1a  [A]  t =  1b1b  [B]  t = 1c1c  [C]  t 1d1d  [D]  t = © 2012 Pearson Education, Inc.

36. Chemical Kinetics Example: What are the relative rates of change in concentration of products and reactant in the decomposition of nitrosyl chloride? 2NOCl (g)  2 NO (g) + Cl 2 (g)

37. Chemical Kinetics Differential Rate Law Considers only the relation between the reaction rate and the concentration of reactants given by a mathematical equation

38. Chemical Kinetics Differential Rate Law To find the exact relation between rate and concentration, we must conduct experiments and collect information. © 2012 Pearson Education, Inc.

39. Chemical Kinetics Concentration and Rate If we compare Experiments 1 and 2, we see that when [NH 4 + ] doubles, the initial rate doubles. NH 4 + (aq) + NO 2  (aq) N 2 (g) + 2 H 2 O(l) © 2012 Pearson Education, Inc.

40. Chemical Kinetics Concentration and Rate Likewise, when we compare Experiments 5 and 6, we see that when [NO 2  ] doubles, the initial rate doubles. NH 4 + (aq) + NO 2  (aq) N 2 (g) + 2 H 2 O(l) © 2012 Pearson Education, Inc.

41. Chemical Kinetics Concentration and Rate This means Rate  [NH 4 + ] Rate  [NO 2  ] Rate  [NH 4 + ] [NO 2  ] which, when written as an equation, becomes Rate = k [NH 4 + ] [NO 2  ] This equation is called the rate law, and k is the rate constant. Therefore, © 2012 Pearson Education, Inc.

42. Chemical Kinetics Rate Laws A rate law shows the relationship between the reaction rate and the concentrations of reactants. The exponents tell the order of the reaction with respect to each reactant. Since the rate law is Rate = k[NH 4 + ] [NO 2  ] the reaction is First-order in [NH 4 + ] and First-order in [NO 2  ] © 2012 Pearson Education, Inc.

43. Chemical Kinetics Rate Laws Rate = k[NH 4 + ] [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. © 2012 Pearson Education, Inc.

44. Chemical Kinetics The rate constant k Is temperature dependent Must be determined by experimentation

45. Chemical Kinetics The formula for the differential rate law C aA + bB  xX Where C is the catalyst Initial rxn rate= k[A] m [B] n [C] p Where m=order of reaction for reactant A, n=order for reactant B,p=order of catalyst Exponents can be 0,whole numbers or fractions and must be determined experimentally. The exponent is 0 it has no effect on the rate.

46. Chemical Kinetics BrO 3 - (aq) + 5 Br - (aq) + 6 H + (aq)  3 Br 2 (aq) + 3 H 2 O Initial Concentrations Rate in M per unit time Mixture[BrO 3 ‑ ]/ M[Br - ] / M[H + ] /M A0.00500.0250.03010 B0.0100.0250.03020 C0.0100.0500.03040 D0.0100.0500.060160

47. Chemical Kinetics Another example Initial Concentrations Rate in mol L -1 hr -1 Experiment[A][B] 10.50 0.20 0.50 x 10 -2 20.75 0.20 0.50 x 10 -2 31.00 0.20 0.50 x 10 -2 40.50 0.40 1.00 x 10 -2 50.50 0.60 1.50 x 10 -2

48. Chemical Kinetics Experiment Starting [X] /M Starting [Y]/M Rate/ M min -1 10.005000.02501.00 20.01000.02504.00 30.005000.01251.00 Finding order of a reaction with algebra