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Chapter 14 Chemical Kinetics John D. Bookstaver St. Charles Community College Cottleville, MO Chemistry, The Central Science , 11th edition Theodore L. Brown; H. Eugene LeMay, Jr.; and Bruce E. Bursten

Kinetics ,[object Object],[object Object]

Factors That Affect Reaction Rates ,[object Object],[object Object],[object Object]

Factors That Affect Reaction Rates ,[object Object],[object Object]

Reaction Rates ,[object Object]

Reaction Rates ,[object Object],C 4 H 9 Cl ( aq ) + H 2 O ( l )  C 4 H 9 OH ( aq ) + HCl ( aq )

Reaction Rates ,[object Object],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

Reaction Rates ,[object Object],[object Object],C 4 H 9 Cl ( aq ) + H 2 O ( l )  C 4 H 9 OH ( aq ) + HCl ( aq )

Reaction Rates and Stoichiometry ,[object Object],[object Object],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

Reaction Rates and Stoichiometry ,[object Object],2 HI ( g )  H 2 ( g ) + I 2 ( g ) ,[object Object],Rate = − 1 2  [HI]  t =  [I 2 ]  t

Reaction Rates and Stoichiometry ,[object Object],a A + b B c C + d D Rate = − 1 a  [A]  t = − 1 b  [B]  t = 1 c  [C]  t 1 d  [D]  t =

Concentration and Rate ,[object Object]

Concentration and Rate ,[object Object],NH 4 + ( aq ) + NO 2 − ( aq ) N 2 ( g ) + 2 H 2 O ( l )

Concentration and Rate ,[object Object],[object Object],[object Object],[object Object],[object Object],[object Object],[object Object],Therefore,

Rate Laws ,[object Object],[object Object],[object Object],[object Object],[object Object],[object Object],[object Object]

Rate Laws ,[object Object],[object Object],[object Object]

Integrated Rate Laws ,[object Object],Where [A] 0 is the initial concentration of A, and [A] t is the concentration of A at some time, t , during the course of the reaction. ln [A] t [A] 0 = − k t

Integrated Rate Laws ,[object Object],ln [A] t − ln [A] 0 = − k t ln [A] t = − k t + ln [A] 0 … which is in the form y = mx + b ln [A] t [A] 0 = − k t

First-Order Processes ,[object Object],ln [A] t = - k t + ln [A] 0

First-Order Processes ,[object Object],CH 3 NC CH 3 CN

First-Order Processes ,[object Object],[object Object],[object Object],[object Object]

Second-Order Processes ,[object Object],also in the form y = mx + b 1 [A] t = kt + 1 [A] 0

Second-Order Processes ,[object Object],1 [A] t = kt + 1 [A] 0 1 [A]

Second-Order Processes The decomposition of NO 2 at 300 ° C is described by the equation and yields data comparable to this: NO 2 ( g ) NO ( g ) + O 2 ( g ) 0.00380 300.0 0.00481 200.0 0.00649 100.0 0.00787 50.0 0.01000 0.0 [NO 2 ], M Time ( s ) 1 2

Second-Order Processes ,[object Object],[object Object],0.00380 0.00481 0.00649 0.00787 0.01000 [NO 2 ], M − 5.573 − 5.337 − 5.038 − 4.845 − 4.610 ln [NO 2 ] 300.0 200.0 100.0 50.0 0.0 Time ( s )

Second-Order Processes ,[object Object],[object Object],0.00380 0.00481 0.00649 0.00787 0.01000 [NO 2 ], M 263 208 154 127 100 1/[NO 2 ] 300.0 200.0 100.0 50.0 0.0 Time ( s ) 1 [NO 2 ]

Half-Life ,[object Object],[object Object],[object Object]

Half-Life ,[object Object],ln 0.5 = − kt 1/2 − 0.693 = − kt 1/2 NOTE: For a first-order process, then, the half-life does not depend on [A] 0 . 0.5 [A] 0 [A] 0 ln = − kt 1/2 = t 1/2 0.693 k

Half-Life ,[object Object],1 0.5 [A] 0 = kt 1/2 + 1 [A] 0 2 [A] 0 = kt 1/2 + 1 [A] 0 2 − 1 [A] 0 = kt 1/2 1 [A] 0 = = t 1/2 1 k [A] 0

Temperature and Rate ,[object Object],[object Object]

The Collision Model ,[object Object],[object Object]

The Collision Model ,[object Object]

Activation Energy ,[object Object],[object Object]

Reaction Coordinate Diagrams ,[object Object]

Reaction Coordinate Diagrams ,[object Object],[object Object],[object Object],[object Object]

Maxwell – Boltzmann Distributions ,[object Object],[object Object]

Maxwell – Boltzmann Distributions ,[object Object],[object Object],f = e - E a RT

Arrhenius Equation ,[object Object],[object Object],[object Object],- E a RT

Arrhenius Equation ,[object Object],[object Object],y = m x + b Therefore, if k is determined experimentally at several temperatures, E a can be calculated from the slope of a plot of ln k vs. . 1 T E a R 1 T

Reaction Mechanisms ,[object Object]

Reaction Mechanisms ,[object Object],[object Object]

Multistep Mechanisms ,[object Object],[object Object]

Slow Initial Step ,[object Object],[object Object],[object Object],[object Object],NO 2 ( g ) + CO ( g )  NO ( g ) + CO 2 ( g )

Slow Initial Step ,[object Object],[object Object],[object Object],[object Object],[object Object]

Fast Initial Step ,[object Object],[object Object],[object Object],2 NO ( g ) + Br 2 ( g )  2 NOBr ( g )

Fast Initial Step ,[object Object],Step 2: NOBr 2 + NO  2 NOBr (slow) Step 1 includes the forward and reverse reactions. Step 1: NO + Br 2 NOBr 2 (fast)

Fast Initial Step ,[object Object],[object Object],[object Object],[object Object]

Fast Initial Step ,[object Object],[object Object],[object Object],[object Object],[object Object],[object Object]

Fast Initial Step ,[object Object],[object Object],[object Object],k 1 k − 1 [NO] [Br 2 ] = [NOBr 2 ]

Fast Initial Step ,[object Object],= k [NO] 2 [Br 2 ] k 2 k 1 k − 1 Rate = [NO] [Br 2 ] [NO]

Catalysts ,[object Object],[object Object]

Enzymes ,[object Object],[object Object]

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