Lecture 15

Lecture 15 - Collision Theory 1.  The reac1on...

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Unformatted text preview: Collision Theory 1.  The reac1on molecules must collide with one another. 2.  The reac1on molecules must collide with sufficient energy. 3.  The molecules must collide in an orienta1on that can lead to rearrangement of the atoms. Transi1on State Theory 1.  Kine1c energy changes to poten1al energy during collision. 2. Every reac1on goes through its own transi1ons state. 3. Transi1on state corresponds to an unstable species with highest poten1al energy. Figure 16.17 Reac/on energy diagrams and possible transi/on states for two reac/ons. The ac#va#on energy (Ea) is the minimum amount of energy required to ini/ate a chemical reac/on. Sample Problem 16.9 PROBLEM: Drawing Reac/on Energy Diagrams and Transi/on States A key reac/on in the upper atmosphere is O3(g) + O(g) 2O2(g) The Ea(fwd) is 19 kJ, and the ΔHrxn for the reac/on is  ­392 kJ. Draw a reac/on energy diagram for this reac/on, postulate a transi/on state, and calculate Ea(rev). Reac/on Mechanisms Most reac/ons occurs through a sequen/al single reac/ons that are called elementary steps or elementary reac#ons. Reac#ons intermediates are formed and used up during these elementary reac/ons. Intermediates are less stable than the reactants or products. Intermediates can be isolated. Intermediates do not appear in the overall balanced reac/on. Each elementary reac/on describes a single molecular event. That is, elementary reac/ons occur in a single step. Molecularity of an elementary reac/on is the number of par/cles involved that single reac/on event. Only uni ­ and bimolecular events are seen in an elementary reac/on. In an elementary reac/on reac/on order is equal to its molecularity. Elementary reac/on rates are different. The slowest reac/on rate of an elementary reac/on is the rate determining step of the overall reac/on. REACTION MECHANISMS The Rate ­Determining Step of a Reac/on Mechanism The overall rate of a reac/on is related to the rate of the slowest step, which is the rate ­determining step. Correla/ng the Mechanism with the Rate Law The elementary steps must add up to the overall balanced equa/on. The elementary steps must be physically reasonable. The mechanism must correlate with the rate law. Reac/on Mechanisms The sequence of elementary steps that leads to product forma/on is the reac#on mechanism. 2NO (g) + O2 (g) 2NO2 (g) N2O2 is detected during the reac/on! Elementary step: NO + NO N2O2 + Elementary step: N2O2 + O2 2NO2 Overall reac/on: 2NO + O2 2NO2 Reac/on Intermediates Intermediates are species that appear in a reac/on mechanism but not in the overall balanced equa/on. An intermediate is always formed in an early elementary step and consumed in a later elementary step. Elementary step: NO + NO N2O2 + Elementary step: N2O2 + O2 2NO2 Overall reac/on: 2NO + O2 2NO2 Rate determining step •  In a mul/step process, one of the steps will be slower than all others. •  The overall reac/on cannot occur faster than this slowest, rate ­ determining step. Slow Ini/al Step NO2 + CO NO + CO2 ; Rate = k[NO2][CO] (If a single step reac/on) Rate = k[NO2]2 (Experimentally Determined Rate Law) 1) Elementary step: 2) Elementary step: Overall reac/on: NO2 + NO2 NO + NO3 Rate1= k[NO2]2 NO3 + CO NO2+ CO2 Rate2= k[NO3][CO] NO2 + CO NO + CO2 •  CO is necessary for this reac/on to occur, but the rate of the reac/on does not depend on its concentra/on sugges/ng that the reac/on occurs in two steps. •  NO3 is the intermediate; formed in the first step and consumed in the second step. •  As CO is not involved in the slow, rate ­determining step, it does not appear in the rate law. Another Example of Slow Ini/al Step 2NO2 + F2 2NO2F2 ; Rate = k[NO2][F2] (Rate Law; Experimentally determined) 1) Elementary step: 2) Elementary step: NO2 + F2 NO2F + F; Rate1= k[NO2][F2] NO2 + F NO2F; Rate2= k[NO2][F] •  Elementary reac/on1 (or step1) is the rate determining step; Overall rate law is same as the rate of the rate determining step. •  Second NO2 is not involved in the rate determining step. •  F is the intermediate Figure 16.18 Reac/on energy diagram for the two ­step reac/on of NO2 and F2. Fast Ini/al Step Elementary step: NO + O2 NO3 Elementary step: NO3 + NO 2NO2 Overall reac/on: 2NO + O2 2NO2 Rate1(f) = Rate1(r] Rate1(f)= k1[NO][O2] Rate1(r) = k ­1[NO3] Rate2 = k2[NO3[NO] k1[NO][O2] = k ­1[NO3] [NO3] =k1/k ­1[NO][O2] Rate2 = k2x k1/k ­1[NO][NO[O2] Rate2 = k [NO]2[O2] where k = k2 x k1/k ­1 Sample Problem 16.10 PROBLEM: (1) (2) Determining Molecularity and Rate Laws for Elementary Steps The following two reac/ons are proposed as elementary steps in the mechanism of an overall reac/on: NO2Cl(g) NO2(g) + Cl (g) NO2Cl(g) + Cl (g) NO2(g) + Cl2(g) (a) Write the overall balanced equa/on. (b) Determine the molecularity of each step. (c) Write the rate law for each step. CATALYSTS • Each catalyst has its own specific way of func/oning. • In general, a catalyst lowers the energy of ac/va/on. •Lowering the Ea increases the rate constant, k, and   thereby increases the rate of the reac/on. • A catalyst increases the rate of the forward and reverse reac/ons. • A catalyzed reac/on yields the products more quickly, but does not yield more product than the uncatalyzed reac/on. • A catalyst lowers Ea by providing a different mechanism, for the reac/on through a new, lower energy pathway. Figure 16.19 Reac/on energy diagram for a catalyzed and an uncatalyzed process. In heterogeneous catalysis, the reactants and the catalysts are in different phases. •  Metal ­catalyzed hydrogena/on of ethylene •  Haber synthesis of ammonia •  Ostwald process for the production of nitric acid In homogeneous catalysis, the reactants and the catalysts are dispersed in a single phase, usually liquid. •  Acid catalysis •  Base catalysis Figure 16.21 The metal ­catalyzed hydrogena/on of ethylene. Mechanism for the catalyzed hydrolysis of an organic ester Enzyme Catalysis •  Enzymes are catalysts in biological systems. •  The substrate fits into the ac/ve site of the enzyme much like a key fits into a lock. Enzyme Catalysis ...
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This note was uploaded on 02/18/2012 for the course CHEM 6C taught by Professor Hoeger during the Spring '08 term at UCSD.

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