Chapter 13: Chemical KineticsChemical KineticsReaction RatesAverage and Instantaneous RatesSlide 5PowerPoint PresentationExampleRate LawReaction OrdersMethod of Initial RatesSlide 11Slide 12Integrated Rate Laws1st Order Integrated Rate Law2nd Order Integrated Rate LawZeroth Order Integrated Rate LawHalf-LifeHalf-Life for 2nd and 0th order reactionsDetermining the Order of a Reaction Using Integrated Rate LawsSlide 20Slide 21Slide 22Slide 23The Effect of Temperature on Reaction RateActivation EnergyFrequency FactorCH4 + CO  CH3CHOReactive CollisionUnreactive CollisionArrhenius PlotsSlide 31Rate Laws and Reaction MechanismsReaction mechanismMolecularityRate Determining StepSlide 36Reaction Energy Diagram for Previous ReactionThe Preequilibrium ApproximationSlide 39Slide 40CatalysisSlide 42Slide 43Heterogeneous CatalysisHomogenous CatalysisEnzymesSlide 47Slide 48Chapter 13 1Chapter 13: Chemical KineticsGeorgia Gwinnett CollegeChem 1212KSection 10Fall 20132•The speed of a chemical reaction is called its reaction rate.•The rate of a reaction is a measure of how fast the reaction makes products.–or uses reactants•Chemical kinetics is the study of the factors that affect reaction rates and the mechanism by which a reaction proceeds.Chemical KineticsReaction Rates•The rate of some event is measured as the change in some quantity that occurs in a given interval of time.•Reaction rates are defined as the change in concentration of reactant or product over some time interval.–typically the units are M/s -or- molecules cm-3 s-13Average and Instantaneous Rates•Average Rate: Change in concentration over some finite interval of time.•Instantaneous Rate: The rate of change of concentration at a specific instant of time.–Derivative of the concentration with respect to time.–The average rate over an infinitesimally small unit of time.–Use these most frequently. If not specified, then it is probably an instantaneous rate.4 Rate =concentration of A at time t2- concentration of A at time t1t2 - t1Reaction Rates52NO2(g)  2NO(g) + O2(g)6Concentration (mol/L)time (s)NO2NOO22NO2(g)  2NO(g) + O2(g)Relative RatesExample•What is the average rate of reaction between 0 and 200 seconds?•What is the rate of loss of NO2 during this same time period?Chapter 13 7Rate Law•The rate law tells us how the rate depends on the concentration of the reactants (and possibly products).–Consider the generic reaction:2A + 3B  2C + D–A typical rate law will have the form:Chapter 13 8 Rate=k A[ ]nB[ ]mReaction OrdersChapter 13 9A  productsMethod of Initial Rates•Initial rate: The instantaneous rate determined just after the reaction begins.–Initial concentrations of reactants hasn’t changed.–No products to complicate things.•Perform several experiments with different initial concentrations of reactants.–Compare the different initial rates to see how they depend on the concentration of each reactant.–Generally change the initial concentration of one reactant while holding the initial concentrations of the other reactants constant.•Then repeat the process for the other reactants.•Once we know the order of each reactant, the rate constant can be determined using any of the initial rates.10Example11At elevated temperatures, HI reacts according to the chemical equation2HI → H2 + I2at 443°C, the rate of reaction increases with concentration of HI, as shown in this table.Data [HI] RatePoint (mol L-1) (mol L-1 s-1)1 0.005 7.5 x 10-42 0.01 3.0 x 10-33 0.02 1.2 x 10-2a) Determine the order of the reaction with respect to HI and write the rate expressionb) Calculate the rate constant and give its unitsc) Calculate the instantaneous rate of reaction for a [HI] = 0.0020MExampleChapter 13 12Integrated Rate Laws•The rate laws we have been considering to this point are more precisely called differential rate laws (although we usually just say “rate law”).–They tell us how the rate depends on concentration.•To start our discussion of integrated rate laws we will consider reactions with just one reactant. The differential rate law for such a reaction is:•Integrated rate laws tell us how concentration depends on time.–The form of the integrated rate law depends on the order of the reaction.•We will consider 1st order, 2nd order and 0th order rate laws.131st Order Integrated Rate Law14Differential Rate Law:Integrated Rate Law: ln[X]t=- kt+ln[X]0 ln[X ]t[X]0=- kt2nd Order Integrated Rate Law15Differential Rate Law:Integrated Rate Law: 1[X]t=kt+1[X]0Zeroth Order Integrated Rate Law16Differential Rate Law:Integrated Rate Law: [X]t=- kt +[X]0Half-Life•The Half-Life (t1/2) of a reaction is the time required for the concentration of a reactant to fall to one-half of its initial value.•Determine the value of the half-life for a 1st order reaction.17Chapter 13Half-Life for 2nd and 0th order reactionsChapter 13 18 t1/ 2=1k[X]0 t1/ 2=[X]02k2nd order reaction0th order reactionDetermining the Order of a ReactionUsing Integrated Rate Laws19What is the reaction order?Chapter 13 20Chapter 13 21Chapter 13 22Example•The decomposition of SO2Cl2 is first order in SO2Cl2 and has a rate constant of 1.42 x 10−4 s−1 at a certain temperature.a) What is the half-life of this reaction?b) How long will it take for the concentration of SO2Cl2 to decrease to 25% of its initial concentration?c) If the initial concentration of SO2Cl2 is 1.00 M, how long will it take for the concentration to decrease to 0.78 M?d) If the initial concentration of SO2Cl2 is 0.150 M, what is its concentration after 200. s? Chapter 13 23The Effect of Temperature on Reaction Rate•The (differential) rate law tells us how concentration depends on temperature.•The temperature dependence is contained in the rate constant (A rate constant is only constant if the temp doesn’t change).Chapter 13 24 Rate=k[X]n k =Ae- EaRTArrhenius EquationActivation EnergyChapter 13 25Frequency Factor•Accounts for how frequently molecules collide.•Also accounts for the fact that molecules must collide in an orientation that will give rise to a reaction.Chapter 13 26 k =Ae- EaRTCH4 + CO  CH3CHO27+Reactive Collision28Unreactive Collision29Arrhenius Plots•Rearrange the Arrhenius equation to another useful form.Chapter 13 30 k =Ae- EaRTExample•The decomposition of hydrogen iodide, 2HI(g)  H2(g) + I2(g), has a rate