Ch15_Kinetics_2.pdf

Page 1 1MECHANISMS 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 2MECHANISMS For example H3C C C CH3 H H H3C C C H H CH3 trans-2-butene cis-2-butene Rate = k [trans-2-butene] Conversion requires twisting around the C=C bond. 3MECHANISMS Conversion of trans to cis buteneConversion of trans to cis butene 4 Energy involved in conversion of trans to cis butene Energy involved in conversion of trans to cis trans cis energy ActivatedComplex 27 kJ/mol 31 kJ/mol 266 kJ -262 kJ MECHANISMS See Figure 15.15 4 kJ/mol 5Mechanisms ? Reaction passes thru a TRANSITION STATE where there is an activated complex that has sufficient energy to become a product. ? Reaction passes thru a ACTIVATION ENERGY, Ea = energy req?d to form activated complex. Here Ea = 266 kJ/mol trans cis energy Activated Complex 27 kJ/mol 31 kJ/mol 4 kJ/mol 266 kJ -262 kJ 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- MECHANISMS trans cis energy Activated Complex 27 kJ/mol 31 kJ/mol 4 kJ/mol 266 kJ -262 kJ Page 2 7Activation Energy A flask full of trans-butene is stable because only a tiny fraction of trans molecules have enough energy to convert to cis. 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 cause reactions to vary from fast to slow.from fast to slow. 8Mechanisms 1. Why is cis <--> trans reaction observed to be 1st order? As [trans] doubles, number of molecules with enough E also doubles. 2. Why is the cis <--> trans reaction faster at higher temperature? Fraction of molecules with sufficient activation energy increases with T. As [trans] doubles, number of <--> reaction 9Mechanisms Reaction of trans --> cis is UNIMOLECULAR - only one reactant is involved. BIMOLECULAR ? two different molecules must collide --> products Reaction of trans --> cis - A bimolecular reaction Exo- or endothermic? 10Collision Theory Reactions require (a) activation energy and (b) correct geometry. O3(g) + NO(g) ---> O2(g) + NO2(g) 2. Activation energy and geometry 1. Activation energy 11 More About Activation Energy k = Ae -E a /RT ln k = -( Ea R )( 1 T ) + ln A Arrhenius equation ?Arrhenius equation ? Rate constant Temp (K) 8.31 x 10-3 kJ/K?molActivation energyFrequency factor Frequency factor = frequency of collisions with correct geometry. Plot ln k vs. 1/T ---> straight line. slope = -Ea/R 12Mechanisms O3 + NO reaction occurs in a single ELEMENTARY step. Most others involve a sequence of elementary steps. Adding elementary steps gives NET reaction. ELEMENTARY step. Most others involve a Page 3 13Mechanisms Most rxns. involve a sequence of elementary steps. 2 I- + H2O2 + 2 H+ ---> I2 + 2 H2O Rate = k [I-] [H2O2] NOTE 1. Rate law comes from experiment 2. Order and stoichiometric coefficients not necessarily the same! 3. Rate law reflects all chemistry down to and including the slowest step in multistep reaction. Most rxns. involve a sequence of elementary - - Order and stoichiometric coefficients not and including the slowest step in multistep 14Mechanisms Proposed Mechanism Step 1 ? slow HOOH + I- --> HOI + OH- Step 2 ? fast HOI + I- --> I2 + OH- Step 3 ? fast 2 OH- + 2 H+ --> 2 H2O Rate of the reaction controlled by slow step ? RATE DETERMINING STEP, rds. Rate can be no faster than rds! - - - - - , rds. Rate can be no faster than rds! Most rxns. involve a sequence of elementary steps. 2 I- + H2O2 + 2 H+ ---> I2 + 2 H2O Rate = k [I-] [H2O2] 15Mechanisms Elementary Step 1 is bimolecular and involves I- and HOOH. Therefore, this predicts the rate law should be Rate [I-] [H2O2] ? as observed!! The species HOI and OH- are reaction intermediates. Elementary Step 1 is - [I- - 2 I- + H2O2 + 2 H+ ---> I2 + 2 H2O Rate = k [I-] [H2O2] Step 1 ? slow HOOH + I- --> HOI + OH- Step 2 ? fast HOI + I- --> I2 + OH- Step 3 ? fast 2 OH- + 2 H+ --> 2 H2O 16 Ozone Decomposition Mechanism Proposed mechanism Step 1: fast O3 (g) ¸ O2 (g) + O (g) Step 2: slow O3 (g) + O (g) ---> 2 O2 (g) 2 O3 (g) ---> 3 O2 (g) Rate = k [O3] 2 [O2] 17CATALYSIS Catalysts speed up reactions by altering the mechanism to lower the activation energy barrier. Dr. James Cusumano, Catalytica Inc. What is a catalyst? Catalysts and society 18CATALYSIS In auto exhaust systems ? Pt, NiO In auto exhaust systems ? Pt, NiO 2 CO + O 2 ---> 2 CO2 2 NO ---> N2 + O2 Page 4 19CATALYSIS 2. Polymers: H2C=CH2 ---> polyethylene 3. Acetic acid: CH3OH + CO --> CH3CO2H 4. Enzymes ? biological catalysts 20CATALYSIS Catalysis and activation energy Uncatalyzed reaction Catalyzed reaction MnO2 catalyzes decomposition of H2O2 2 H2O2 ---> 2 H2O + O2 21 Iodine-Catalyzed Isomerization of cis-2-Butene Figure 15.18 22 Iodine-Catalyzed Isomerization of cis-2-Butene Figure 15.19 John Kotz kinetics/2