1. 1 Reaction Mechanism The series of steps by which a chemical reaction occurs. A chemical equation does not tell us how reactants become products - it is a summary of the overall process.

2. 2 Reaction Mechanism (continued) 4 The reaction has many steps in the reaction mechanism.

3. 3 Example Consider the Reaction CH 3 Cl + Br -  CH 3 Br + Cl - What are the possible steps to this reaction?

4. 4 Example (continued) Possibility A – All in one step 1) CH 3 Cl + Br -  CH 3 Br + Cl - Possibility B – Two Steps 1) CH 3 Cl  CH 3 + + Cl - 2) CH 3 + + Br -  CH 3 Br What would the rate laws look like for each possibility?

5. 5 Often Used Terms Intermediate: formed in one step and used up in a subsequent step and so is never seen as a product. Molecularity: the number of species that must collide to produce the reaction indicated by that step. Elementary Step: A reaction whose rate law can be written from its molecularity. uni, bi and termolecular

6. 6 Example (continued) Possibility A – All in one step 1) CH 3 Cl + Br -  CH 3 Br + Cl - Possibility B – Two Steps 1) CH 3 Cl  CH 3 + + Cl - 2) CH 3 + + Br -  CH 3 Br What is the molecularity for each step? Which are unimolecular? Which are bimolecular?

7. 7 Example (continued) Data From Experiments RateCH 3 ClBr- 211 421 412 What is the rate Law? Which Mechanism is consistent with the Rate Law?

8. 8 Rate-Determining Step In a multistep reaction, the slowest step determines the rate of reaction.

9. 9 The Collision Model In a chemical reaction, bonds are broken and new bonds are formed. Molecules must collide with the correct orientation and with enough energy to cause bond breakage and formation. In other words, there is a minimum amount of energy required for reaction: the activation energy, E a.

10. 10 Reaction Coordinate Diagrams It is helpful to visualize energy changes throughout a process on a reaction coordinate diagram like this one for the rearrangement of methyl isonitrile.

11. 11 Reaction Coordinate Diagrams It 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.

12. 12 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.

13. 13 Maxwell–Boltzmann Distributions As the temperature increases, the curve flattens and broadens. Thus at higher temperatures, a larger population of molecules has higher energy.

14. 14 Maxwell–Boltzmann Distributions If the dotted line represents the activation energy, as the temperature increases, so does the fraction of molecules that can overcome the activation energy barrier. As a result, the reaction rate increases.

15. 15 Maxwell–Boltzmann Distributions This fraction of molecules can be found through the expression where R is the gas constant and T is the Kelvin temperature. f = e −E a /RT

16. 16 Arrhenius Equation Svante Arrhenius developed a mathematical relationship between k and E a : k = A e −E a /RT where A is the frequency factor, a number that represents the likelihood that collisions would occur with the proper orientation for reaction.

17. 17 Arrhenius Equation Taking the natural logarithm of both sides, the equation becomes ln k = -E a ( ) + ln A 1 RT y = mx + 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.

18. 18 Reaction Mechanisms The sequence of events that describes the actual process by which reactants become products is called the reaction mechanism.

19. 19 Reaction Mechanisms Reactions may occur all at once or through several discrete steps. Each of these processes is known as an elementary reaction or elementary process.

20. 20 Reaction Mechanisms The molecularity of a process tells how many molecules are involved in the process.

21. 21 Multistep Mechanisms In a multistep process, one of the steps will be slower than all others. The overall reaction cannot occur faster than this slowest, rate-determining step.

22. 22 Slow Initial Step The rate law for this reaction is found experimentally to be Rate = k [NO 2 ] 2 CO is necessary for this reaction to occur, but the rate of the reaction does not depend on its concentration. This suggests the reaction occurs in two steps. NO 2 (g) + CO (g)  NO (g) + CO 2 (g)

23. 23 Slow Initial Step A proposed mechanism for this reaction is Step 1: NO 2 + NO 2  NO 3 + NO (slow) Step 2: NO 3 + CO  NO 2 + CO 2 (fast) The NO 3 intermediate is consumed in the second step. As CO is not involved in the slow, rate-determining step, it does not appear in the rate law.

24. 24 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.

25. 25 Catalysts One way a catalyst can speed up a reaction is by holding the reactants together and helping bonds to break.

26. 26 Enzymes Enzymes are catalysts in biological systems. The substrate fits into the active site of the enzyme much like a key fits into a lock.

27. 27 Collision Model Key Idea: Molecules must collide to react. However, only a small fraction of collisions produce a reaction. Why? Arrhenius: An activation energy must be overcome.

28. 28 Arrhenius Equation 4 Collisions must have enough energy to produce the reaction (must equal or exceed the activation energy). 4 Orientation of reactants must allow formation of new bonds.

29. 29 Figure 12.12 Plot Showing the Number of Collisions with a Particular Energy at T 1 and T 2, where T 2  

30. 30 Arrhenius Equation (continued) k = rate constant A = frequency factor E a = activation energy T = temperature R = gas constant (8.314 J/k mole)

31. 31 Figure 12.10 A Plot Showing the Exponential Dependence of the Rate Constant on Absolute Temperature

32. 32 Take the Log Taking the natural log of both sides of the Arrhenious equation gives: Ln(k) = - Ea/R(1/T) + ln(A) Thus graphing Ln(k) vs. 1/T gives a line with a slope of –Ea/R! See figure 12.14 in book.

33. 33 Two Temperature method If only two temperatures are used, we can simplify the calculation of Ea. Ln (k 2 /k 1 ) = Ea/R (1/T 1 – 1/T 2 ) The k’s and the T’s are determined by experiment.

34. 34 Catalysis Catalyst: A substance that speeds up a reaction without being consumed Enzyme: A large molecule (usually a protein) that catalyzes biological reactions. Homogeneous catalyst: Present in the same phase as the reacting molecules. Heterogeneous catalyst: Present in a different phase than the reacting molecules.

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