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Chemical Kinetics PowerPoint Presentation
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Chemical Kinetics

Chemical Kinetics

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Chemical Kinetics

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  1. Chemical Kinetics

  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 rateis 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.

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

  4. 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

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

  6. 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

  7. aA + bB cC + dD The Rate Law The rate lawexpresses 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 orderin A reaction is yth orderin B reaction is (x +y)th order overall

  8. 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

  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? 0 1 1

  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? 1 2 3

  11. 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 notrelated to the stoichiometric coefficient of the reactant in the balanced chemical equation. rate = k [F2][ClO2]

  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 = 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. The order of a reaction can be found out from a graph of conc against time. The gradient of the graph at any time is the rate of the reaction. If the rate of the reaction (gradient) is a constant, it is zero order reaction. This is because the change in concentration does not affect the reaction rate. If the rate (gradient) of reaction is halved, when the concentration is halved, it is first order reaction. If the halving the concentration makes the rate (gradient) decreased by a factor of 4, the reaction is second order. If one of the reagent has a high concentration and at the end of the reaction, the other reagent disappeared, the reaction is virtually proportional to the that reagent disappeared.

  14. Half life • This is the time required for half of the reagent to disappear. • If t½ halves when the concentration is halved, it is zero order. • If t½ remains same when the concentration halved, it is first order reaction. • If t½ doubles when concentration halved, it is second order. • Radio active decay is first order.

  15. Rate Law Order Half-Life = [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 rate = k 0 rate = k [A] 1 rate = k [A]2 2 K= rate constant; [A]0 = initial conc; [A]= conc at any time;

  16. 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 stepsor elementary reactions. The sequence of elementary stepsthat leads to product formation is the reaction mechanism. N2O2 is detected during the reaction!

  17. 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 notin the overall balanced equation. Anintermediateis always formed in an early elementary step and consumed in a later elementary step.

  18. The rate-determining stepis 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.

  19. 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

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

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

  22. 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

  23. 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.

  24. THE END