33-1. Energy Profiles. Energy Profiles Reactions will:- 1.Break or weaken bonds in reactants then form bonds in products.

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1 Energy Profiles Energy Profiles Reactions will:- 1.Break or weaken bonds in reactants then form bonds in products. 2. Reactants pass over a Potential Energy Barrier on way to Products. Plot: NO 33-1

2 Energy Profile Diagrams so: example: water can be formed from the reaction of hydrogen and oxygen or can be decomposed into hydrogen and oxygen 33-2

3 in most chemical reactions there are a series of elementary reactions E TS = transition state I = intermediate TS 2 E a Reactants TS 1 I 2 TS 3 Ea (reverse) I 1 )E rxn Products reaction coordinate 33-3

4 important points Each hump represents an elementary step e.g., A + B C Rate = k[a][b] 2A products Rate = k [A] 2 sum of elementary steps must add up to give overall balanced equation (also see examples coming up) maxima are transition states: cannot be isolated valley are intermediates : can be trapped and isolated Exo or endo- depends on sign of )E rxn Ea refers to highest energy transition state : TS 2 here the reaction cannot go faster than the slowest step 33-4

5 example Consider the reaction: CR (g) + CH 4 (g) 6 CR-H (g) + CH 3 (g) An Elementary Reaction for which E a = 16 kj mol -1 and )EE = )HE = +8 kj mol -1 Plot Energy Profile: E 16 kj mol -1 CR (g) + CH 4 (g) CR-H (g) + CH 3 (g) +8 kj mol -1 reaction coordinate 33-5

6 Rate determining step the experimentally observed rate (and rate law) tell us only about the slowest step in the reaction sequence this turns out to be very useful... concept first 7 /hr observed rate observed rate is governed by this step 33-6

7 Rate limiting reactions. The reaction mechanism is the dissection of a complicated reaction sequence into its elementary steps. The overall rate of a complex reaction is the rate of the slowest of these elementary steps, which is called the rate determining or the rate limiting step. In a chemical reaction, the reaction rate is determined by the events leading up to and including the rate determining step, but not by the (fast) reactions which occur afterwards. lets see how this works

8 example : then a bit more on catalysts for the upcoming quiz The following mechanism is suggested for the reaction of with stoichiometry: CO(g) + NO 2 (g) v NO(g) + CO 2 (g) rate law: rate = k[no 2 ] 2 1, slow: 2 NO 2 (g) v NO(g) + NO 3 (g) 2, fast: NO 3 (g) + CO(g) v NO 2 (g) + CO 2 (g) rate = k 1 = k[no 2 ] 2 which is what we observe. Thus the rate equation includes 2 mol NO 2, which were involved at the slow step of the reaction, but does not include CO, which did not become involved in the reaction until after the slow step. Since the rate of the reaction does not depend on [CO], the reaction is said to be zero order in CO. We can then conclude: "The proposed mechanism is consistent with the Experimental data." 33-8

9 F 2 (g) + 2NO(g) 6 2 ONF(g) from experiment: rate = k[no][f 2 ] Possible Mechanism: 1. NO + F 2 6 ONF + F (slow) 2. F + NO 6 ONF (fast) If this were true, then it would give rise to the rate equation: rate = k 1 [NO][F 2 ] which is what we OBSERVE. We can then conclude: "The proposed mechanism is consistent with the Experimental data. more examples soon 33-9

10 More on catalysts as we saw in lecture 31 catalysts speed up reactions by lowering the Ea E with catalyst R P there are several other properties of catalysts we need to know r.c 33-10

11 Catalysis Homogeneous and Heterogeneous A Catalyst increases the rate of a chemical reaction without itself being used up in the process. It does this by providing an alternative path with a lower activation energy for the conversion of reactants to products. Homogeneous: Reactants and Catalyst in the same phase: Example: the conversion of 2O 3 (g) 6 3O 2 (g) is catalyzed by Freons (molecules like CF 2 CR 2 which are used in airconditioners etc.). Heterogeneous Reactant and catalyst in different phase. Usually involve a solid as the catalyst important point: catalyst is regenerated 33-11

12 example: Ozone decomposition uncatalyzed mechanism: (two steps) O 3 º O 2 + O rapid equilibrium, K 1 O 3 + O 6 2O 2, slow, k 2 rate determining step Overall: 2O 3 (g) 6 3O 2 (g): Rate = -0.5d[O 3 ]/dt = k 2 [O 3 ][O] see page 27-6 (get rid of the [O] (an intermediate) method: slow step: rate = k 2 [O 3 ] [O] = k 2 [O 3 ] K 1 [O 3 ]/[O 2 ] new idea: intermediates do not appear in rate law --only reactants (but from equil step K 1 =[O 2 ][O]/[O 3 ] so: [O] = K 1 [O 3 ]/[O 2 ] = k 2 K 1 [O 3 ) 2 /[O 2 ] = k [O 3 ) 2 /[O 2 ] same as observed by expt. k is exptl rate constant 33-12

13 Catalyzed mechanism CF 2 CR 2 6 CR@ 2 CR CR@ + O 3 + O + O + O 2 CR@ 2 CR 6 CF 2 CR 2 Net effect is O 3 + O 6 2O 2, but now no longer "slow" it is important that you can identify intermediates and catalysts from a scheme like this intermediates: 2 catalyst: CF 2 CR 2 (regenerated) 33-13

14 Note that with the catalyst resulting in a lower Activation Energy than the uncatalyzed reaction, a larger fraction of molecules will have sufficient energy to exceed E a and hence participate in the reaction. recall: 33-14

15 another one what is the catalyst Cl or ClO?? 33-15

16 homework One pathway for destruction of ozone in the upper atmosphere is: O 3 º O 2 + O (rapid equilibrium) O 3 + NO 6 NO 2 + O 2 (slow) NO 2 + O 6 NO + O 2 (fast) 2O 3 6 3O 2 (overall) The catalyst in this reaction is?? (a) O (b) NO 2 (c) O 3 (d) O 2 (e) NO solution is on bottom LHS of p