1 / 31

Chemical Kinetics

Chemical Kinetics. Chapter 13. A B. rate =. D [A]. D [B]. rate = -. D t. D t. Chemical Kinetics. Thermodynamics – does a reaction take place? Kinetics – how fast does a reaction proceed?.

nemo
Download Presentation

Chemical Kinetics

An Image/Link below is provided (as is) to download presentation Download Policy: Content on the Website is provided to you AS IS for your information and personal use and may not be sold / licensed / shared on other websites without getting consent from its author. Content is provided to you AS IS for your information and personal use only. Download presentation by click this link. While downloading, if for some reason you are not able to download a presentation, the publisher may have deleted the file from their server. During download, if you can't get a presentation, the file might be deleted by the publisher.

E N D

Presentation Transcript

  1. Chemical Kinetics Chapter 13

  2. A B rate = D[A] D[B] rate = - Dt Dt Chemical Kinetics Thermodynamics – does a reaction take place? Kinetics – how fast does a reaction proceed? _______________ is the change in the ___________ of a reactant or a product with time (M/s). D[A] = D[B] = Because [A] decreases with time, D[A] is_______. 13.1

  3. A B time rate = D[A] D[B] rate = - Dt Dt 13.1

  4. Br2(aq) + HCOOH (aq) 2Br-(aq) + 2H+(aq) + CO2(g) slope of tangent slope of tangent slope of tangent D[Br2] average rate = - = - Dt instantaneous rate = 13.1

  5. rate k = [Br2] rate a [Br2] rate = k [Br2] =___________ = 3.50 x 10-3 s-1 13.1

  6. Factors that Affect Reaction Rate • _____________ • __________Theory: When two chemicals react, their molecules have to collide with each other with sufficient energy for the reaction to take place. • __________Theory: _________temperature means the molecules move faster. • ___________of reactants • More reactants mean more collisions if enough energy is present • ___________ • Speed up reactions by lowering _________energy • ____________of a solid reactant • Bread and Butter theory: more area for reactants to be in contact • ____________of gaseous reactants or products • Increased number of collisions

  7. aA + bB cC + dD The Rate Law The rate law expresses the relationship of the rate of a reaction to the _________and the ____________of the reactants raised to some powers. Rate = k [A]x[B]y reaction is ______order in A reaction is ______order in B reaction is _______order overall 13.2

  8. F2(g) + 2ClO2(g) 2FClO2(g) rate = k [F2]x[ClO2]y 13.2

  9. What is the order with respect to A? What is the order with respect to B? What is the overall order of the reaction?

  10. What is the order with respect to Cl2? What is the order with respect to NO? What is the overall order of the reaction?

  11. F2(g) + 2ClO2(g) 2FClO2(g) 1 Rate Laws • Rate laws are ________determined experimentally. • Reaction order is ______defined in terms of______(not_______) concentrations. • The order of a reactant _______related to the stoichiometric coefficient of the reactant in the balanced chemical equation. rate = k [F2][ClO2] 13.2

  12. Determine the rate law and calculate the rate constant for the following reaction from the following data: S2O82-(aq) + 3I-(aq) 2SO42-(aq) + I3-(aq) = rate = k [S2O82-]x[I-]y 13.2

  13. D[A] rate = - Dt First-Order Reactions [A] = [A]0e-kt rate = k [A] [A] is the concentration of A at any time t ln[A] - ln[A]0 = - kt [A]0 is the concentration of A at time t=0 13.3

  14. Decomposition of N2O5 13.3

  15. Integrated Rate Law Problem • NH2NO2H2O + N2O • The rate law is ratio = k [NH2NO2] with a k=5.62 x 10-3 min-1 at 15 C. Starting with 0.105 M NH2NO2. • (a) at what time will [NH2NO2]= 0.0250M • (b) what is [NH2NO2] after 6.00 hr?

  16. 0.693 = = = k What is the half-life of N2O5 if it decomposes with a rate constant of 5.7 x 10-4 s-1? = First-Order Reactions The ________is the time required for the concentration of a reactant to decrease to half of its initial concentration. t = How do you know decomposition is first order? units of k (s-1) 13.3

  17. A product # of half-lives [A] = [A]0/n First-order reaction 1 2 2 4 3 8 4 16 13.3

  18. 13.3

  19. 1 1 - = kt [A] [A]0 D[A] rate = - Dt Second-Order Reactions [A] is the concentration of A at any time t rate = k [A]2 [A]0 is the concentration of A at time t=0 Half life for second order 13.3

  20. D[A] rate = - Dt Zero-Order Reactions rate = k [A]0 = k [A] is the concentration of A at any time t [A] - [A]0 = kt [A]0 is the concentration of A at time t=0 Half life for zero order 13.3

  21. Concentration-Time Equation Order Rate Law Half-Life 1 1 - = kt [A] [A]0 = [A]0 t½ = t½ t½ = Ln 2 2k k 1 k[A]0 Summary of the Kinetics of Zero-Order, First-Order and Second-Order Reactions [A] - [A]0 = - kt rate = k 0 ln[A] - ln[A]0 = - kt 1 rate = k [A] 2 rate = k [A]2 13.3

  22. A + B C + D _____thermic Reaction ____thermic Reaction The ____________________is the minimum amount of energy required to initiate a chemical reaction. 13.4

  23. Ln k = - + lnA -Ea 1 T R Temperature Dependence of the Rate Constant k = A •exp( -Ea/RT ) (Arrhenius equation) Eais the activation energy (J/mol) R is the gas constant (8.314 J/K•mol) T is the absolute temperature A is the frequency factor 13.4

  24. 2NO (g) + O2 (g) 2NO2 (g) Elementary step: NO + NO N2O2 + Elementary step: N2O2 + O2 2NO2 Overall reaction: 2NO + O2 2NO2 Reaction Mechanisms The overall progress of a chemical reaction can be represented at the molecular level by a series of simple elementary steps or elementary reactions. The sequence of elementary steps that leads to product formation is the reaction mechanism. N2O2 is detected during the reaction! 13.5

  25. Elementary step: NO + NO N2O2 + Elementary step: N2O2 + O2 2NO2 Overall reaction: 2NO + O2 2NO2 Reaction Intermediates Intermediates are species that appear in a reaction mechanism but not in the overall balanced equation. An intermediate is always formed in an early elementary step and consumed in a later elementary step. 13.5

  26. The rate-determining step is the sloweststep in the sequence of steps leading to product formation. Rate Laws and Rate Determining Steps • Writing plausible reaction mechanisms: • The sum of the elementary steps must give the overall balanced equation for the reaction. • The rate-determining step should predict the same rate law that is determined experimentally. 13.5

  27. Unimolecular reaction Bimolecular reaction Bimolecular reaction A + B products A + A products A products Rate Laws and Elementary Steps rate = k [A] rate = k [A][B] rate = k [A]2 13.5

  28. uncatalyzed catalyzed Ea k A ___________is a substance that increases the rate of a chemical reaction without itself being consumed. ratecatalyzed > rateuncatalyzed 13.6

  29. Energy Diagrams Exothermic Endothermic • Activation energy (Ea) for the forward reaction • Activation energy (Ea) for the reverse reaction • (c) Delta H

  30. The experimental rate law for the reaction between NO2 and CO to produce NO and CO2 is rate = k[NO2]2. The reaction is believed to occur via two steps: Step 1: Step 2: NO2 + NO2 NO + NO3 NO2+ CO NO + CO2 NO3 + CO NO2 + CO2 What is the equation for the overall reaction? What is the intermediate? Catalyst? NO3 NO2 What can you say about the relative rates of steps 1 and 2? rate = k[NO2]2 is the rate law for step 1 so step 1 must be slower than step 2 13.5

  31. Write the rate law for this reaction. Rate = k [HBr] [O2] List all intermediates in this reaction. HOOBr, HOBr None List all catalysts in this reaction.

More Related