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CHEMICAL KINETICS Chpt 12

CHEMICAL KINETICS Chpt 12.

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CHEMICAL KINETICS Chpt 12

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  1. CHEMICAL KINETICSChpt 12

  2. Briggs-Rauscher ReactionIO3- + 2 H2O2 + CH2(CO2H)2 + H+ --> ICH(CO2H)2 + 2 O2 + 3 H2O Reaction Mechanism:HIO3 + H2O2 HIO2 + O2 + H2OIO3- + HIO2 + H+2IO2· + H2OIO2· + Mn+2 + H2O  HIO2 + Mn(OH)+2Mn(OH)+2 + H2O2  Mn+2 + HOO· + H2O2HOO· O2 + H2O22HIO2 IO3- + H+ + HIOHIO + H2O2 I- + O2 + H+ + H2OHIO + I- + H+ I2 + H2OI2 + CH2(CO2H)2 ICH(CO2H)2 + H+ + I-IO3- + 2H2O2 + H+ HIO +2O2 + 2H2OI- + HIO + H+ I2 + H2OI2 + CH2(CO2H)2 ICH(CO2H)2 + H+ + I-HIO2 + I- + H+ 2HIOIO3- + HIO2 + H+ 2IO2· + H2OIO2· + Mn+2 + H2O  HIO2 + Mn(OH)+2I- + HIO + H+ I2 + H2OI2 + CH2(CO2H)2 ICH(CO2H)2 + H+ + I-IO3- + HIO2 + H+ 2IO2· + H2OIO2· + Mn+2 + H2O  HIO2 + Mn(OH)+2HIO2 + I- + H+ 2 HIO

  3. Thermodynamics – does a reaction take place?Kinetics – how fast does a reaction proceed?

  4. How can we MEASURE the RATE of a reaction?A  B

  5. A B rate = D[A] D[B] rate = - Dt Dt Chemical Kinetics 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

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

  7. 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

  8. 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

  9. 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

  10. What is the order with respect to A? What is the order with respect to B? What is the overall order of the reaction? 0 1 1

  11. What is the order with respect to Cl2? What is the order with respect to NO? What is the overall order of the reaction? 1 2 3

  12. 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

  13. 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

  14. 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

  15. Decomposition of N2O5 13.3

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

  17. 13.3

  18. 1 1 - = kt [A] [A]0 t½ = D[A] rate = - Dt 1 k[A]0 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 t½ = t when [A] = [A]0/2 13.3

  19. [A]0 t½ = D[A] 2k 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 t½ = t when [A] = [A]0/2 13.3

  20. 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

  21. REACTION MECHANISMS

  22. 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

  23. 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

  24. 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

  25. 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

  26. Collision Theory

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

  29. Arrhenius Equation 1899,Swedish Chemist Svante Arrhenius *High temperature and low activation energy favor larger rate constants, and thus speed up the reaction* (The larger the ratio the smaller the rate)

  30. Arrhenius Equation **Relationship of T and Rate (enough KE to satisfy Ea) **Includes Collision Frequency and Orientation (“A”)

  31. Ln k = - + lnA Ea 1 T R k = A •e( -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

  32. Ln k = - + lnA Ea 1 T R How do we experimentally find “A” and Ea for a reaction? 13.4

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

  34. Factors that Affect Reaction Rate • Temperature • Collision Theory: When two chemicals react, their molecules have to collide with each other with sufficient energy for the reaction to take place. • Kinetic Theory: Increasing temperature means the molecules move faster. • Concentrations of reactants • More reactants mean more collisions if enough energy is present • Catalysts • Speed up reactions by lowering activation energy • Surface area of a solid reactant • Bread and Butter theory: more area for reactants to be in contact • Pressure of gaseous reactants or products • Increased number of collisions

  35. 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

  36. 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.

  37. 4NH3(g) + 5O2(g) 4NO (g) + 6H2O (g) 2NO (g) + O2(g) 2NO2(g) 2NO2(g) + H2O (l) HNO2(aq) + HNO3(aq) Hot Pt wire over NH3 solution Pt-Rh catalysts used in Ostwald process Ostwald Process Pt catalyst 13.6

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

  39. Enzyme Catalysis 13.6

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