Kinetics ppt
- 1. 1 Chemical Kinetics Chapter 13 Copyright © The McGraw-Hill Companies, Inc. Permission required for reproduction or display.
- 2. 2 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). A B rate = - D[A] Dt rate = D[B] Dt 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.
- 3. 3 A B rate = - D[A] Dt rate = D[B] Dt
- 4. 4 Br2 (aq) + HCOOH (aq) 2Br- (aq) + 2H+ (aq) + CO2 (g) time 393 nm light Detector D[Br2] a D Absorption red-brown t1< t2 < t3
- 5. 5 Br2 (aq) + HCOOH (aq) 2Br- (aq) + 2H+ (aq) + CO2 (g) average rate = - D[Br2] Dt = - [Br2]final – [Br2]initial tfinal - tinitial slope of tangent slope of tangent slope of tangent instantaneous rate = rate for specific instance in time
- 6. 6 rate a [Br2] rate = k [Br2] k = rate [Br2] = rate constant = 3.50 x 10-3 s-1
- 7. 7 2H2O2 (aq) 2H2O (l) + O2 (g) PV = nRT P = RT = [O2]RT n V [O2] = P RT 1 rate = D[O2] Dt RT 1 DP Dt = measure DP over time
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- 9. 9 Reaction Rates and Stoichiometry 2A B Two moles of A disappear for each mole of B that is formed. rate = D[B] Dt rate = - D[A] Dt 1 2 aA + bB cC + dD rate = - D[A] Dt 1 a = - D[B] Dt 1 b = D[C] Dt 1 c = D[D] Dt 1 d
- 10. 10 Write the rate expression for the following reaction: CH4 (g) + 2O2 (g) CO2 (g) + 2H2O (g) rate = - D[CH4] Dt = - D[O2] Dt 1 2 = D[H2O] Dt 1 2 = D[CO2] Dt
- 11. 11 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. aA + bB cC + dD 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
- 12. 12 F2 (g) + 2ClO2 (g) 2FClO2 (g) rate = k [F2]x[ClO2]y Double [F2] with [ClO2] constant Rate doubles x = 1 Quadruple [ClO2] with [F2] constant Rate quadruples y = 1 rate = k [F2][ClO2]
- 13. 13 F2 (g) + 2ClO2 (g) 2FClO2 (g) rate = k [F2][ClO2] 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. 1
- 14. 14 Determine the rate law and calculate the rate constant for the following reaction from the following data: S2O8 2- (aq) + 3I- (aq) 2SO4 2- (aq) + I3 - (aq) Experiment [S2O8 2-] [I-] Initial Rate (M/s) 1 0.08 0.034 2.2 x 10-4 2 0.08 0.017 1.1 x 10-4 3 0.16 0.017 2.2 x 10-4 rate = k [S2O8 2-]x[I-]y Double [I-], rate doubles (experiment 1 & 2) y = 1 Double [S2O8 2-], rate doubles (experiment 2 & 3) x = 1 k = rate [S2O8 2-][I-] = 2.2 x 10-4 M/s (0.08 M)(0.034 M) = 0.08/M•s rate = k [S2O8 2-][I-]
- 15. 15 First-Order Reactions A product rate = - D[A] Dt rate = k [A] k = rate [A] = 1/s or s-1M/s M = D[A] Dt = k [A]- [A] is the concentration of A at any time t [A]0 is the concentration of A at time t=0 [A] = [A]0e−kt ln[A] = ln[A]0 - kt
- 16. 16 2N2O5 4NO2 (g) + O2 (g) Graphical Determination of k
- 17. 17 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 ? ln[A] = ln[A]0 - kt kt = ln[A]0 – ln[A] t = ln[A]0 – ln[A] k = 66 s [A]0 = 0.88 M [A] = 0.14 M ln [A]0 [A] k = ln 0.88 M 0.14 M 2.8 x 10-2 s-1 =
- 18. 18 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 ln [A]0 [A]0/2 k =t½ ln 2 k = 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? t½ ln 2 k = 0.693 5.7 x 10-4 s-1 = = 1200 s = 20 minutes How do you know decomposition is first order? units of k (s-1)
- 19. 19 A product First-order reaction # of half-lives [A] = [A]0/n 1 2 3 4 2 4 8 16
- 20. 20 Second-Order Reactions A product rate = - D[A] Dt rate = k [A]2 k = rate [A]2 = 1/M•s M/s M2= D[A] Dt = k [A]2- [A] is the concentration of A at any time t [A]0 is the concentration of A at time t=0 1 [A] = 1 [A]0 + kt t½ = t when [A] = [A]0/2 t½ = 1 k[A]0
- 21. 21 Zero-Order Reactions A product rate = - D[A] Dt rate = k [A]0 = k k = rate [A]0 = M/s D[A] Dt = k- [A] is the concentration of A at any time t [A]0 is the concentration of A at time t = 0 t½ = t when [A] = [A]0/2 t½ = [A]0 2k [A] = [A]0 - kt
- 22. 22 Summary of the Kinetics of Zero-Order, First-Order and Second-Order Reactions Order Rate Law Concentration-Time Equation Half-Life 0 1 2 rate = k rate = k [A] rate = k [A]2 ln[A] = ln[A]0 - kt 1 [A] = 1 [A]0 + kt [A] = [A]0 - kt t½ ln 2 k = t½ = [A]0 2k t½ = 1 k[A]0
- 23. 23 Exothermic Reaction Endothermic Reaction The activation energy (Ea ) is the minimum amount of energy required to initiate a chemical reaction. A + B AB C + D++
- 24. 24 Temperature Dependence of the Rate Constant Ea is 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 ln k = - Ea R 1 T + lnA (Arrhenius equation) )/( RTEa eAk Alternate format:
- 25. 25 Alternate Form of the Arrhenius Equation At two temperatures, T1 and T2 or
- 26. 26 Importance of Molecular Orientation effective collision ineffective collision
- 27. 27 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. 2NO (g) + O2 (g) 2NO2 (g) N2O2 is detected during the reaction! Elementary step: NO + NO N2O2 Elementary step: N2O2 + O2 2NO2 Overall reaction: 2NO + O2 2NO2 +
- 28. 28 2NO (g) + O2 (g) 2NO2 (g) Mechanism:
- 29. 29 Elementary step: NO + NO N2O2 Elementary step: N2O2 + O2 2NO2 Overall reaction: 2NO + O2 2NO2 + 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. The molecularity of a reaction is the number of molecules reacting in an elementary step. • Unimolecular reaction – elementary step with 1 molecule • Bimolecular reaction – elementary step with 2 molecules • Termolecular reaction – elementary step with 3 molecules
- 30. 30 Unimolecular reaction A products rate = k [A] Bimolecular reaction A + B products rate = k [A][B] Bimolecular reaction A + A products rate = k [A]2 Rate Laws and Elementary 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. The rate-determining step is the slowest step in the sequence of steps leading to product formation.
- 31. 31 Sequence of Steps in Studying a Reaction Mechanism
- 32. 32 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: NO2 + NO2 NO + NO3 Step 2: NO3 + CO NO2 + CO2 What is the equation for the overall reaction? NO2+ CO NO + CO2 What is the intermediate? NO3 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
- 33. 33 CH2 CH2 CH2 CH2 CH2 CH22 CH2 CH2 CH2 CH2 CH2 CH2 CH2 CH2 • • CH2 CH22 Chemistry In Action: Femtochemistry
- 34. 34 A catalyst is a substance that increases the rate of a chemical reaction without itself being consumed. Ea k ratecatalyzed > rateuncatalyzed Ea < Ea′ Uncatalyzed Catalyzed )/( RTEa eAk
- 35. 35 In heterogeneous catalysis, the reactants and the catalysts are in different phases. In homogeneous catalysis, the reactants and the catalysts are dispersed in a single phase, usually liquid. • Haber synthesis of ammonia • Ostwald process for the production of nitric acid • Catalytic converters • Acid catalysis • Base catalysis
- 36. 36 N2 (g) + 3H2 (g) 2NH3 (g) Fe/Al2O3/K2O catalyst Haber Process
- 37. 37 Ostwald Process Pt-Rh catalysts used in Ostwald process 4NH3 (g) + 5O2 (g) 4NO (g) + 6H2O (g) Pt catalyst 2NO (g) + O2 (g) 2NO2 (g) 2NO2 (g) + H2O (l) HNO2 (aq) + HNO3 (aq)
- 38. 38 Catalytic Converters CO + Unburned Hydrocarbons + O2 CO2 + H2O catalytic converter 2NO + 2NO2 2N2 + 3O2 catalytic converter
- 40. 40 Binding of Glucose to Hexokinase
- 41. 41 rate = D[P] Dt rate = k [ES] Enzyme Kinetics