1. 1 MECHANISMS A Microscopic View of Reactions Sections 15.5 and 15.6 How are reactants converted to products at the molecular level? Want to connect the RATE LAW ----> MECHANISM experiment ---->theory How are reactants converted to products at the molecular level? Want to connect the RATE LAW ----> MECHANISM experiment ---->theory

2. 2 MECHANISMSMECHANISMS For example Rate = k [trans-2-butene] Conversion requires twisting around the C=C bond. For example Rate = k [trans-2-butene] Conversion requires twisting around the C=C bond.

3. 3 MECHANISMSMECHANISMS Conversion of trans to cis butene

4. 4 Energy involved in conversion of trans to cis butene MECHANISMSMECHANISMS See Figure 15.15

5. 5 MECHANISMSMECHANISMS Reaction passes thru a TRANSITION STATE where there is an activated complex that has sufficient energy to become a product. ACTIVATION ENERGY, E a = energy req’d to form activated complex. Here E a = 233 kJ/mol Reaction passes thru a TRANSITION STATE where there is an activated complex that has sufficient energy to become a product. ACTIVATION ENERGY, E a = energy req’d to form activated complex. Here E a = 233 kJ/mol

6. 6 Also note that trans-butene is MORE STABLE than cis-butene by about 4 kJ/mol. Therefore, trans ---> cis is ENDOTHERMIC. This is the connection between thermo- dynamics and kinetics. Also note that trans-butene is MORE STABLE than cis-butene by about 4 kJ/mol. Therefore, trans ---> cis is ENDOTHERMIC. This is the connection between thermo- dynamics and kinetics. MECHANISMSMECHANISMS

7. 7 Activation Energy A flask full of trans-butene is stable because only a tiny fraction of trans molecules have enough energy to convert to cis. In general, differences in activation energy are the reason reactions vary from fast to slow. A flask full of trans-butene is stable because only a tiny fraction of trans molecules have enough energy to convert to cis. In general, differences in activation energy are the reason reactions vary from fast to slow.

8. 8 MECHANISMSMECHANISMS 1.Why is reaction observed to be 1st order? As [trans] doubles, number of molecules with enough E also doubles. As [trans] doubles, number of molecules with enough E also doubles. 2.Why is the reaction faster at higher temperature? Fraction of molecules with sufficient activation energy increases with T. Fraction of molecules with sufficient activation energy increases with T. 1.Why is reaction observed to be 1st order? As [trans] doubles, number of molecules with enough E also doubles. As [trans] doubles, number of molecules with enough E also doubles. 2.Why is the reaction faster at higher temperature? Fraction of molecules with sufficient activation energy increases with T. Fraction of molecules with sufficient activation energy increases with T.

9. 9 MECHANISMSMECHANISMS Reaction of trans --> cis is UNIMOLECULAR- only one reactant is involved.

10. 10 Reaction of trans --> cis is UNIMOLECULAR- only one reactant is involved. BIMOLECULAR — two different molecules must collide --> products Reaction of trans --> cis is UNIMOLECULAR- only one reactant is involved. BIMOLECULAR — two different molecules must collide --> products MECHANISMSMECHANISMS

11. 11 MECHANISMS Reaction of trans --> cis is UNIMOLECULAR - only one reactant is involved. BIMOLECULAR — two different molecules must collide --> products Reaction of trans --> cis is UNIMOLECULAR - only one reactant is involved. BIMOLECULAR — two different molecules must collide --> products A bimolecular reaction Exo- or endothermic?

12. 12 Collision Theory Reactions require (a) activation energy and (b) correct geometry. O 3 (g) + NO(g) ---> O 2 (g) + NO 2 (g) O 3 (g) + NO(g) ---> O 2 (g) + NO 2 (g) Reactions require (a) activation energy and (b) correct geometry. O 3 (g) + NO(g) ---> O 2 (g) + NO 2 (g) O 3 (g) + NO(g) ---> O 2 (g) + NO 2 (g)

13. 13 Collision Theory Reactions require (a) activation energy and (b) correct geometry. O 3 (g) + NO(g) ---> O 2 (g) + NO 2 (g) O 3 (g) + NO(g) ---> O 2 (g) + NO 2 (g) Reactions require (a) activation energy and (b) correct geometry. O 3 (g) + NO(g) ---> O 2 (g) + NO 2 (g) O 3 (g) + NO(g) ---> O 2 (g) + NO 2 (g) 2. Activation energy and geometry 1. Activation energy

14. 14 MECHANISMSMECHANISMS O 3 + NO reaction occurs in a single ELEMENTARY step. Most others involve a sequence of elementary steps. Adding elementary steps gives NET reaction. O 3 + NO reaction occurs in a single ELEMENTARY step. Most others involve a sequence of elementary steps. Adding elementary steps gives NET reaction.

15. 15 MECHANISMSMECHANISMS O 3 + NO reaction occurs in a single ELEMENTARY step. Most others involve a sequence of elementary steps. Adding elementary steps gives NET reaction. O 3 + NO reaction occurs in a single ELEMENTARY step. Most others involve a sequence of elementary steps. Adding elementary steps gives NET reaction.

16. 16 MECHANISMSMECHANISMS Most rxns. involve a sequence of elementary steps. 2 I - + H 2 O 2 + 2 H + ---> I 2 + 2 H 2 O 2 I - + H 2 O 2 + 2 H + ---> I 2 + 2 H 2 O Rate = k [I - ] [H 2 O 2 ] Rate = k [I - ] [H 2 O 2 ] Step 1 — slowHOOH + I - --> HOI + OH - Step 2 — fastHOI + I - --> I 2 + OH - Step 3 — fast2 OH - + 2 H + --> 2 H 2 O Rate of the reaction controlled by slow step — RATE DETERMINING STEP, rds. RATE DETERMINING STEP, rds. Rate can be no faster than rds! Most rxns. involve a sequence of elementary steps. 2 I - + H 2 O 2 + 2 H + ---> I 2 + 2 H 2 O 2 I - + H 2 O 2 + 2 H + ---> I 2 + 2 H 2 O Rate = k [I - ] [H 2 O 2 ] Rate = k [I - ] [H 2 O 2 ] Step 1 — slowHOOH + I - --> HOI + OH - Step 2 — fastHOI + I - --> I 2 + OH - Step 3 — fast2 OH - + 2 H + --> 2 H 2 O Rate of the reaction controlled by slow step — RATE DETERMINING STEP, rds. RATE DETERMINING STEP, rds. Rate can be no faster than rds!

17. 17 MECHANISMSMECHANISMS Step 1 is bimolecular and involves I - and HOOH. Therefore, this predicts the rate law should be Rate  [I - ] [H 2 O 2 ] — as observed!! The species HOI and OH - are reaction intermediates. Step 1 is bimolecular and involves I - and HOOH. Therefore, this predicts the rate law should be Rate  [I - ] [H 2 O 2 ] — as observed!! The species HOI and OH - are reaction intermediates. 2 I - + H 2 O 2 + 2 H + ---> I 2 + 2 H 2 O Rate = k [I - ] [H 2 O 2 ] Rate = k [I - ] [H 2 O 2 ] Step 1 — slowHOOH + I - --> HOI + OH - Step 2 — fastHOI + I - --> I 2 + OH - Step 3 — fast2 OH - + 2 H + --> 2 H 2 O

18. 18 Arrhenius Equation Reaction rates depend on energy, frequency of collisions, temperature, and geometry of molecules given by: A = frequency of collisions with correct geometry at concentration of 1M (L/mol*s) R = gas constant (8.314 x 10 -3 kJ/K*mol) e -Ea/RT is fraction of molecules having the minimum energy required for reaction

19. 19 Arrhenius Equation Calculate the value of the activation energy from the temp. dependence of the rate constant Calculate the rate constant for a given temp. (if activation energy and A are known)

20. 20 Arrhenius Equation Taking the natural log and rearranging: Straight line plot of ln k vs 1/T Slope of –E a /R

21. 21 CATALYSISCATALYSIS Catalysts speed up reactions by altering the mechanism to lower the activation energy barrier.

22. 22 CATALYSISCATALYSIS Catalysts speed up reactions by altering the mechanism to lower the activation energy barrier. What is a catalyst?Catalysts and society Dr. James Cusumano, Catalytica Inc.

23. 23 CATALYSISCATALYSIS In auto exhaust systems — Pt, NiO 2 CO + O 2 ---> 2 CO 2 2 NO ---> N 2 + O 2 In auto exhaust systems — Pt, NiO 2 CO + O 2 ---> 2 CO 2 2 NO ---> N 2 + O 2

24. 24 CATALYSISCATALYSIS 2.Polymers: H 2 C=CH 2 ---> polyethylene 3.Acetic acid: CH 3 OH + CO --> CH 3 CO 2 H 4. Enzymes — biological catalysts 2.Polymers: H 2 C=CH 2 ---> polyethylene 3.Acetic acid: CH 3 OH + CO --> CH 3 CO 2 H 4. Enzymes — biological catalysts

25. 25 CATALYSISCATALYSIS Catalysis and activation energy Uncatalyzed reaction Catalyzed reaction MnO 2 catalyzes decomposition of H 2 O 2 2 H 2 O 2 ---> 2 H 2 O + O 2

26. 26 Iodine-Catalyzed Isomerization of cis-2-Butene Figure 15.18

27. 27 Iodine-Catalyzed Isomerization of cis-2-Butene Figure 15.19