This document discusses reaction kinetics and rate laws. It explains that the rate of a reaction can be defined based on the disappearance of reactants or appearance of products over time. The rate law expression relates the reaction rate to the concentrations of reactants and has the form Rate = k[A]n[B]m, where k is the rate constant and n and m are the orders of reactants A and B. The orders must be determined experimentally by measuring initial rates as concentrations are varied. A first-order reaction has a rate that depends on just one reactant concentration. Integrating the differential rate law gives an integrated rate law relating concentration to time, such as the first-order rate law ln[A] = -kt + ln

15. Figure 12.1 p. 562 Definition of Rate. 2NO 2  2NO + O 2

20. At .90 M the rate is - (.98 - .76) = -0.22 = 5.5x 10 -4 M• s -1 (0-400) -400

21. At .45 M the rate is - (.52 - .31) = -0.22 = 2.7 x 10 -4 (1000-1800) -800

41. Figure 12.4 A Plot of In (N 2 O 5 ) Versus Time (1st order)

43. Figure 12.5: A Plot of (N 2 O 5 ) vs. Time for the Decomposition Reaction of N 2 O 5

51. Z7e 546 Figure 12.7 A Plot of [A] Versus t for a Zero-Order Reaction

53. Fig. 12.8 Decomposition Reaction N 2 O  2N 2 + O 2 Even though [N 2 O] is twice as great in (b) as in (a) the reaction occurs on a saturated platinum surface, so rate is zero order

54. Zero Order Graph Time Concentration

55. Zero Order Graph Time Concentration  A]/  t = slope  t -k =  A] t 1/2 = [A 0 ]/2k

62. Use for Online HW pp Use for online HW Hint: on one problem you will have to first determine the order , then use the half-life equation (5 th line) to solve for k, then use the Integrated rate law (2 nd line) to solve for time (t).

64. Figure 12.9 A Molecular Representation of the Elementary Steps in the Reaction of NO 2 and CO Overall: NO 2 + CO  NO + CO 2 Step 1: NO 2 + NO 2  NO 3 + NO (k 1 ) Step 2: NO 3 + CO  NO 2 + CO 2 (k 2 ) NO 3 is an intermediate

80. Potential Energy Reaction Coordinate Reactants Products

81. Potential Energy Reaction Coordinate Reactants Products Activation Energy E a

82. Potential Energy Reaction Coordinate Reactants Products Activated complex

83. Potential Energy Reaction Coordinate Reactants Products  E }

84. Potential Energy Reaction Coordinate 2BrNO 2NO + Br Br---NO Br---NO 2 Transition State

88. Figure 12.13 Several Possible Orientations for a Collision Between Two BrNO Molecules. (a) & (b) can lead to a collision, (c) cannot.

89. No Reaction O N Br O N Br O N Br O N Br O N Br O N Br O N Br O N Br O N Br O N Br

91. Figure 12.14 pp Plot of In( k ) vs. 1/ T for 2N 2 O 5( g )    g ) + O 2( g ). Can get E a from slope = -E a /R Do in lab 12.

92. Figure 12.14 Plot of In( k ) vs. 1/ T for 2N 2 O 5( g )    g ) + O 2( g ). Can get E a from slope = -E a /R Do in lab 12.

93. Activation Energy and Rates The final saga

97. Second step is slow High activation energy E a

98. E a Third step is fast Low activation energy

99. Second step is rate determining

100. Intermediates are present

101. Activated Complexes or Transition States

104. Figure 12.15 Energy Plots for a Catalyzed and an Uncatalyzed Pathway for a Given Reaction. ∆ E is the same in both cases.

105. Figure 12.16 Effect of a Catalyst on the Number of Reaction-Producing Collisions (increases even though no temperature increase).

107. Heterogenous Catalysts Pt surface H H H H C H H C H H

110. Heterogenous Catalysts Pt surface H C H H C H H H H H