20080220 Kinetics 6 - Intro Chemistry II 030.102 - Chapter...

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Intro Chemistry II – 030.102 - Chapter 13 Spring, 2008 1 13.6. Reaction Dynamics Compute number of collisions per unit volume. Rate of collisions of single molecule with other molecules of same type: Z 1 = 2 ! d 2 ( ) u ( N / V ) = 4 d 2 RT M ( N / V ) d = molecular diam, u = avg. speed, M = molar mass, ( N / V ) = number density For N molecules, total number of collisions per unit time = (1/2) N × Z 1 . (Note factor of 1/2.) Divide by V to get rate of collisions per unit volume: Z AA = 2 d 2 RT M ( N / V ) 2 Relate Z AA to 2nd-order rate constant for the reaction A + A products If activation energy is E a , then only fraction exp(– E a / RT ) of these collisions have sufficient energy to overcome barrier to reaction. Reaction Dynamics Each reactive collision lead to loss of 2 molecules A, so density decreases as d ( N / V ) dt = –2 Z AA e E a / RT = 2 ! 2 d 2 " RT M e E a / RT ( N / V ) 2 Convert from molecules to moles: N 0 [A] = ( N / V ) rate = 1 2 d [A] dt = 2 d 2 N 0 RT M e E a / RT [A] 2 Predicted 2nd-order rate constant is: k = 2 d 2 N 0 RT M e E a / RT Comparison with experiment: Can obtain Arrhenius parameters E a and A . For most reactions, A much less than predicted rate constant. Even for collisions with sufficient kinetic energy, reactants must have right orientation to react. Concept of steric factor .
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Intro Chemistry II – 030.102 - Chapter 13 Spring, 2008 2 13.7. Kinetics of Catalysis Catalyst = substance that takes part in chemical reaction and speeds it up but undergoes no permanent chemical change itself. They do not appear in overall balanced chemical equation but affect the rate law by modifying pathways, or more commonly providing completely new pathways for reaction. Two types:
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20080220 Kinetics 6 - Intro Chemistry II 030.102 - Chapter...

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