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Chemical Kinetics Ch 13

Chemical Kinetics Ch 13. We have learned that enthalpy is the sum of the internal energy plus the energy associated with the work done by the system (PV) on the atmosphere In addition, entropy and Gibb’s Free Energy predict whether a reaction will occur or not

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Chemical Kinetics Ch 13

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  1. Chemical Kinetics Ch 13 • We have learned that enthalpy is the sum of the internal energy plus the energy associated with the work done by the system (PV) on the atmosphere • In addition, entropy and Gibb’s Free Energy predict whether a reaction will occur or not • Now we will examine the rate at which a reaction will proceed

  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? Reaction rate is the change in the concentration of a reactant or a product with time (M/s). D[A] = change in concentration of A over time period Dt D[B] = change in concentration of B over time period Dt Because [A] decreases with time, D[A] is negative. 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) time Br2(aq) 393 nm 393 nm Detector light D[Br2] aDAbsorption 13.1

  5. Br2(aq) + HCOOH (aq) 2Br-(aq) + 2H+(aq) + CO2(g) slope of tangent slope of tangent slope of tangent [Br2]final – [Br2]initial D[Br2] average rate = - = - Dt tfinal - tinitial instantaneous rate = rate for specific instance in time 13.1

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

  7. 2H2O2 (aq) 2H2O (l) + O2 (g) [O2] = P n V 1 1 D[O2] P = RT = [O2]RT RT RT DP rate = = Dt Dt measure DP over time PV = nRT 13.1

  8. 2H2O2 (aq) 2H2O (l) + O2 (g) 13.1

  9. 2A B aA + bB cC + dD rate = - = = rate = - = - D[D] D[B] D[A] D[B] D[C] D[A] rate = 1 1 1 1 1 Dt Dt Dt Dt Dt Dt c d a 2 b Reaction Rates and Stoichiometry Two moles of A disappear for each mole of B that is formed. 13.1

  10. D[CO2] = Dt D[CH4] rate = - Dt Write the rate expression for the following reaction: CH4(g) + 2O2(g) CO2(g) + 2H2O (g) D[H2O] = Dt D[O2] = - 1 1 Dt 2 2 13.1

  11. aA + bB cC + dD The Rate Law The rate law expresses the relationship of the rate of a reaction to the rate constant and the concentrations of the reactants raised to some powers. Rate = k [A]x[B]y reaction is xth order in A reaction is yth order in B reaction is (x +y)th order overall 13.2

  12. F2(g) + 2ClO2(g) 2FClO2(g) rate = k [F2]x[ClO2]y Double [F2] with [ClO2] constant Rate doubles x = 1 rate = k [F2][ClO2] Quadruple [ClO2] with [F2] constant Rate quadruples y = 1 13.2

  13. F2(g) + 2ClO2(g) 2FClO2(g) 1 Rate Laws • Rate laws are always determined experimentally. • Reaction order is always defined in terms of reactant (not product) concentrations. • The order of a reactant is not related to the stoichiometric coefficient of the reactant in the balanced chemical equation. rate = k [F2][ClO2] 13.2

  14. 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 = 2.2 x 10-4 M/s = [S2O82-][I-] (0.08 M)(0.034 M) rate = k [S2O82-]x[I-]y y = 1 x = 1 rate = k [S2O82-][I-] Double [I-], rate doubles (experiment 1 & 2) Double [S2O82-], rate doubles (experiment 2 & 3) = 0.08/M•s 13.2

  15. Definitions: 1st Order Reaction • First Order Reaction: The reaction that occurs first, not always the one desired. For example, the formation of brown gunk in an organic prep.

  16. A product rate = [A] M/s D[A] - M = k [A] Dt [A] = [A]0exp(-kt) ln[A] = ln[A]0 - kt D[A] rate = - Dt First-Order Reactions rate = k [A] = 1/s or s-1 k = [A] is the concentration of A at any time t [A]0 is the concentration of A at time t=0 13.3

  17. Definitions: Physical Chemistry • The pitiful attempt to apply y=mx+b to everything in the universe.

  18. Decomposition of N2O5 13.3

  19. Learn this equation & how to use

  20. An Example • A reaction is 25% complete in 42 seconds. Calculate the half life. Ln 1/.75=-.287, Div by 42 and k=6.84x10-3 T1/2 =0.693/k=101 s

  21. 0.88 M ln The reaction 2A B is first order in A with a rate constant of 2.8 x 10-2 s-1 at 800C. How long will it take for A to decrease from 0.88 M to 0.14 M ? 0.14 M = 2.8 x 10-2 s-1 ln ln[A]0 – ln[A] = k k [A]0 [A] [A]0 = 0.88 M ln[A] = ln[A]0 - kt [A] = 0.14 M kt = ln[A]0 – ln[A] = 66 s t = 13.3

  22. [A]0 What is the half-life of N2O5 if it decomposes with a rate constant of 5.7 x 10-4 s-1? ln t½ [A]0/2 0.693 = = = = k k t½ ln2 ln2 0.693 = k k 5.7 x 10-4 s-1 First-Order Reactions The half-life, t½, is the time required for the concentration of a reactant to decrease to half of its initial concentration. t½ = t when [A] = [A]0/2 = 1200 s = 20 minutes How do you know decomposition is first order? units of k (s-1) 13.3

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

  24. 13.3

  25. Second-Order Reactions- ignore 13.3

  26. Concentration-Time Equation Order Rate Law Half-Life 1 1 = + kt [A] [A]0 = [A]0 t½ = t½ t½ = ln2 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

  27. A + B C + D Endothermic Reaction Exothermic Reaction The activation energy (Ea) is the minimum amount of energy required to initiate a chemical reaction. 13.4

  28. lnk = - + lnA Ea 1 T R Temperature Dependence of the Rate Constant also known as the Arrhenius equation 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

  29. Activation Energy • The useful quantity of energy available in one cup of coffee

  30. A product rate [A]0 D[A] - = k Dt [A]0 t½ = D[A] 2k rate = - Dt Zero-Order Reactions rate = k [A]0 = k = M/s 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 t½ = t when [A] = [A]0/2 13.3

  31. lnk = - + lnA Ea 1 T R 13.4

  32. 13.4

  33. uncatalyzed catalyzed Ea k ‘ Ea< Ea A catalyst is a substance that increases the rate of a chemical reaction without itself being consumed. k = A •exp( -Ea/RT ) ratecatalyzed > rateuncatalyzed 13.6

  34. In heterogeneous catalysis, the reactants and the catalysts are in different phases. • Haber synthesis of ammonia • Ostwald process for the production of nitric acid • Catalytic converters In homogeneous catalysis, the reactants and the catalysts are dispersed in a single phase, usually liquid. • Acid catalysis • Base catalysis 13.6

  35. Fe/Al2O3/K2O N2 (g) + 3H2 (g) 2NH3 (g) catalyst Haber Process 13.6

  36. 4NH3(g) + 5O2(g) 4NO (g) + 6H2O (g) 2NO (g) + O2(g) 2NO2(g) 2NO2(g) + H2O (l) HNO2(aq) + HNO3(aq) Pt-Rh catalysts used in Ostwald process Hot Pt wire over NH3 solution Ostwald Process Pt catalyst 13.6

  37. catalytic CO + Unburned Hydrocarbons + O2 CO2 + H2O converter catalytic 2NO + 2NO2 2N2 + 3O2 converter Catalytic Converters 13.6

  38. Enzyme Catalysis 13.6

  39. enzyme catalyzed uncatalyzed D[P] rate = Dt rate = k [ES] 13.6

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