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Chapter 16- Kinetics

by: Carly Rasmussen

Chapter 16- Kinetics Chem 1066

Carly Rasmussen
U of M
GPA 3.93
Chemical Principles II
Dr. Driessen

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About this Document

Notes on kinetics from Chapter 16 of the textbook. Let me know if you have questions and look out for the lecture notes!
Chemical Principles II
Dr. Driessen
Class Notes
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This 8 page Class Notes was uploaded by Carly Rasmussen on Sunday February 8, 2015. The Class Notes belongs to Chem 1066 at University of Minnesota taught by Dr. Driessen in Spring2015. Since its upload, it has received 175 views. For similar materials see Chemical Principles II in Chemistry at University of Minnesota.


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Date Created: 02/08/15
Chapter 16 1 Focusing on the Reaction Rate a Chemical kinetics the study of how fast reactants change into products b Reaction rate the change in the concentration of reactants or products as a function of time i Faster reaction I reactant concentration decreases quickly ii Slower reaction I reactant concentration decreases slowly iii Dependent on given set of conditions c Four factors the concentrations of the reactants their physical state the temperature of the reaction and the use of a catalyst i Concentrations of reactants molecules must collide to react 1 Rate oc collision frequency oc concentration ii Physical state molecules must mix to collide 1 more finely divided a solid or liquid reactant the greater its surface area I more contactfaster iii temperature molecules must collide with enough energy increase in temperature increases energy and of collisions 1 Rate oc collision frequency oc concentration 2 Rate oc collision energy oc concentration 2 Expressing the Reaction Rate AAl ABl a Overview Rate At or Rate At i Units moles per liter per second b Average Instantaneous and Initial Reaction Rates i Rate varies as the reactions proceeds c Three types of reaction rates i Average rate the slope of the line joining two points along the curve 1 For full time total change in concentration divided by total change in time ii Instantaneous rate the rate at a particular instant in a reaction decrease the time period of average rate 1 Slope of the line tangent to the curve at any point derivivative iii Initial rate instantaneous rate when the reactants are first mixed t0 1 Avoid complications arising from the reverse reaction rate increasing as more reactants becomes product d Expressing Rate in Terms of Reactant and Product Concentrations i Expressing the rate of a reaction and its numerical values depends on which substance is the reference pointdetermined by number of moles of each molecule that are formed ii For reaction aA bB cC dD 1AA 1 lACLAD 1 Rat a At b At c At d At 3 The Rate Law and its Components a Overview i Rate lawrate equation rate as a function of concentrations and temperature ii For reaction aA bB cC dD RatekAmB 1 krate constant dependent on reactiontemperature and doesn t change as reaction proceeds 2 Reaction orders super script letters that determine how rate is affected by reactant concentration iii Key Points 1 Balancing coefficients a and b in reaction equation not neccesarily related to reaction orders m and n 2 Components of rate law ratereaction ordersrate constant have to be found by experiment b Some Lab Methods of Determining Initial Rate i Spectrometric methods concentrations of a components that absorbs characteristic wavelengths of lightproportional to increasedecrease in intensity over time ii Conductometric methods change in electrical conductivity of reaction solution When nonionic reactants form ionic productsvice versa iii Manometric methods manometer measures pressure that attaches to reaction vessel With fixed volumetemperature When reaction causes change in gas moles c Determining Reaction Orders i MeaningTerminology reactions have an individual order with respect to each reactant and an overall order the sum of the individuals ii If reaction is Allproducts 1 First order if the rate doubles When A doubles a RatekA1kA 2 Second order rate quadruples When A doubles depends on A2 a RatekA2 3 Zero order rate doesn t change When A doubles rate doesn t depend on A a RatekA klk 4 Negative reaction number rate decreases as concentration of that component increases 5 reaction orders cannot be deduced from balanced equation need experimental data d Determining the Rate Constant ratel i39 AHBl ii Units vary depending on the order of the reaction and the time mols 1 Concentrations L mols 2 Reaction rate L X time iii Example 02g 2NOg I 2N02g 1 Ratek021NO12 N02 6 021Z ratel ratel AHBJ quot 2 k 172x10393L2molzxs 4 Integrated Rate Laws Concentration Changes Over Time a Overview i Integrated rate laws different forms of rate laws ii Include time as a variable 1 How long Will it take to use up x moles per liter of A 2 What is A after y minutes of reaction b Integrated Rate Laws for First Second and ZeroOrder Reactions i For a general first order reaction rate is the negative of the change in A divided by change in time A A 1 RateVrateklAlgi1esus AB kA A 0 Z A it 2 integrate over timel Z 6 3 ln Z a A0concentration of A at t0 b AFconcentration of A at any t ii General second order reaction With A and B 1 Simplest form rate law only contains reactant A A 22 2 Ar Z AA a Z g1ves us Rate t k6 Q A Z Z A 0 b Integrating over timel Z 1 3 iii General zero order reaction A A 0 1 RateszklA kv A AlzkAt 2 AL Ml c Determining Reaction Orders from an Integrated Rate Law i If rate data are not availableljrearrange integrated rate law in ymxb A 0 ii First order A Q kt 3 Z A at Z A 0 iii Second order 3 r 1 Z A 0 iV Zero order A Q kt Z Z d Reaction HalfLife i Halflife fl2 time it takes for the reactant concentration of a reaction to reach half its initial value ALO Ado Z A6 2 ln2kt1 t206i 11 F1rst order Z I 1 I 5 l l k 3 2 Z I mi 6 mi 1 Doesn t depend on reactant concentrations iii iv A 0 k 3 Second order 1 t 1 I T r 2 1 Inversely proportional to the initial reactant concentration shorter halflife with higher initial reactant concentration A 0 4 Zero order 6 qza E 1 Directly proportional to the initial reactant concentration 5 Theories of Chemical Kinetics a Collision Theory Basis of the Rate Law i iii iv Collision theory particles must collide to react relies on collision energy and molecular structure to explain effects of concentration and temperature on rate Why concentrations are multiplied in the rate law 1 Collision frequency provides an upper limit on reaction speed 2 Law of probabilities of collisions is product of reactant particles not sum Effect of rate temperature on rate constant and rate 1 k increase exponentially as T increases Ea 2 Arrhenius equation k AgE 3 Higher Tlarger kincreased rate Central importance of activation energy 1 Activation energy Ea energy threshold that colliding molecules must exceed to react reach activated state able to turn into products 2 AHsza rEa rev fwd Effect of temperature on collision energy 1 Rise in temperature causes a Higher collision frequency b Higher collision energy 2 Collision frequency minor factor small increase 3 Collision energy major factor Ea a f 6 RT i f fraction of collisions With energy greater than E a b Rise in temperature enlarges fraction 4 Smaller Ea higher lelarger jlarger khigher rate vi Calculating activation energy 1 2 vii Calculate from Arrhenius equation by taking natural log k E In 2 2 a i i k1 R T2 T1 Effect of molecular structure on rate 1 Effective collisions ones that lead to a product because atoms that bond to make product collide Need both enough energy and correct molecular orientation Molecular orientation represented by ApZ which is orientation probability factor p times collision frequency Z a More complex molecules mean lower p values b Transition State Theory What Activation Energy is Used For i Transition state theory focuses on high energy species that exist during moment of effective collision reactants becoming products ii Visualizing transition state theory 1 2 3 4 5 Two molecules orbit one another come closerincreasing electron cloud repulsions Enough energy to overcomemoving fast enough allows collision Correct orientation collision bonding between molecules Forms transition stateactivated complex highly unstable highest potential energy E1 needed to get to this state Doesn t necessarily mean products will form iii Every reactionstep of reaction goes through own transition state 6 Reaction Mechanisms The Steps from Reactants to Products a Reaction mechanisms sequence of single reaction steps that sum to overall reaction b Elementary reactions and molecularity i Elementary reactionssteps steps that make up reaction mechanism single molecular event ii Molecularity number of reactant particles in an elementary step 1 Unimolecular decomposition or rearrangement of a single particle 2 Bimolecular two particles interact 3 Termolecular very rare little chance of correct p and enough energy iii Reaction order equals molecularity for rate law c Rate determining step of a reaction mechanism i Each elementary step has its own rate ii Rate determininglowering step one step that is much slower limits how fast the reaction can gobecomes rate law for overall reaction iii Reaction intermediate substance formed and used up during reaction d Correlating mechanism with rate law i Can never be absolute only hypothesisthree criteria C 1 Elementary steps must add to overall balanced equation 2 Must be reasonable unimolecularbimolecular 3 Must correlate With rate laW ii Overall rate law includes all the reactants present in the rate determining step iii Only reactants involving and up to the slow step present in final rate laW Each step in mechanism has own peak in energy diagrams slow step has higher peak 7 Catalysis Speeding up a Reaction a b Catalyst substance that increases rate Without being consumed Basis of catalytic action i Generally Catalystlower Ealarger kljhigher rate ii Provides a different mechanism for the reaction Homogenous catalyst exist in solution With the reaction mixture i Example decomposition of hydrogen peroxide Heterogeneous catalyst speeds up reaction in a different phase i Mostly solids interacting with gaseousliquid ii Hydrogenation addition of Hz to CC to form CC Kinetics and functions of biological catalysts i Enzyme protein catalyst 1 Active site shaped by amino acid side chains catalyze reactions 2 Substrate reactant molecules that bind at active site via IMF ii Catalytic activity act like both homogenous interact directly With reactants and heterogeneous provide stable surface to bind With other molecules 1 Highly efficient and specific one reaction iii Induced fit active site changes along With reactants unlike old lock and key model iv Substrates and enzymes for enzymesubstrate complexes ES 1 Rate determiningratekES so adding more substrate wouldn t speed reaction V All enzymes catalyze by stabilizing reaction s transition state


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