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Chapter 16: Kinetics Rates and Mechanisms of Chemical Reactions

Chapter 16: Kinetics Rates and Mechanisms of Chemical Reactions. 16.1 Factors That Influence Reaction Rate 16.2 Expressing the Reaction Rate 16.3 The Rate Law and Its Components 16.4 Integrated Rate Laws: Concentration Changes over Time 16.5 The Effect of Temperature on Reaction Rate

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Chapter 16: Kinetics Rates and Mechanisms of Chemical Reactions

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  1. Chapter 16: Kinetics Rates and Mechanisms of Chemical Reactions 16.1 Factors That Influence Reaction Rate 16.2 Expressing the Reaction Rate 16.3 The Rate Law and Its Components 16.4 Integrated Rate Laws: Concentration Changes over Time 16.5 The Effect of Temperature on Reaction Rate 16.6 Explaining the Effects of Concentration and Temperature 16.7 Reaction Mechanisms: Steps in the Overall Reaction 16.8 Catalysis: Speeding Up a Chemical Reaction

  2. A  B Reaction Progress Red  Blue t (min) 100 molecules 54 Red : 46 Blue 30 Red : 70 Blue

  3. A  B Moles A & B vs. Time

  4. A  B

  5. C4H9Cl(aq) + H2O(l)  C4H9OH(aq) + HCl(aq)

  6. Rate of Reaction C4H9Cl(aq) + H2O(l)  C4H9OH(aq) + HCl(aq) • The average rate can be expressed in terms of the disappearance of C4H9Cl. • The units for average rate are mol/L.s or M/s. • The average rate decreases over time. • Plot [C4H9Cl] versus time. • The rate at any instant in time (instantaneous rate) is the slope of the tangent to the curve. • Instantaneous rate is different from average rate. • The instantaneous rate is usually the rate of reaction.

  7. Plot of [C4H9Cl] vs. Time

  8. Important Factors that Effect Reaction Rates 1) Concentration: molecules must collide in order to react. The higher the concentration, the higher number of collisions. Rate = k (collision frequency) = k (concentration) k = rate constant 2) Physical state: molecules must physically mix in order to collide. The physical state (solid, liquid, gas) will affect frequency of collisions, as well as the physical size of droplets (liquid) or particles in the case of solids.(heterogeneous vs. homogeneous) 3) Temperature: molecules must collide with enough energy to react. (Activation Energy) Raising the temperature increases the K.E. of the molecules, the number of collisions and the energy of the collisions.

  9. Fig. 16.13

  10. Fig. 16.14

  11. Fig 16.15

  12. Fig. 16.16

  13. Rate Constant vs. Temperature

  14. Reaction Rate Expression / Rate Law For a Chemical Reaction:aA + bB cC + dD A Rate Expression for this reaction could be written as: Rate = k [A]m[B]n . . . k = the reaction rate constant m & n are called reaction orders. They define how the rate is effected by the concentration of each reactant, for example, if the rate doubles when the concentration of A doubles, the rate depends on [A]1, so a = 1; if the rate quadruples when the concentration of B doubles, the rate depends on [B]2, so b = 2. Can a reaction order equal zero?

  15. First Order Reaction Plots

  16. Second-Order Reaction Plots

  17. Integrated Rate Laws and Reaction Order Fig. 16.7

  18. First or Second Order? Fig. 16.8

  19. Fig. 16.9

  20. (p. 684)

  21. The Arrhenius Equation • k is the rate constant, Eais the activation energy, R is the gas constant (8.314 J/K-mol) and T is the temperature in K. • A is called the frequency factor. • A is a measure of the probability of a favorable collision. • Both A and Ea are specific to a given reaction.

  22. The Arrhenius Equation • With sufficient data, Ea and A can be determined graphically by rearranging the Arrhenius equation: • With only two sets of data:

  23. Fig. 16.11

  24. Fig. 16.12

  25. Fig. 16.17

  26. Rate Laws for General Elementary Steps Elementary Step Molecularity Rate Law A product Unimolecular Rate = k[A] 2A product Bimolecular Rate = k[A]2 A + B product Bimolecular Rate = k[A][B] Least likely of the elementary steps. Why? 2A + B product Termolecular Rate = k[A]2[B] Table 16.6 (p. 693)

  27. Bimolecular (2nd Order) Reaction

  28. Nature of the Transition State in the Reaction Between CH3Br and OH- Fig. 16.18

  29. Fig. 16.19

  30. Overall Reaction Mechanism

  31. Fig. 16.22

  32. Heterogeneous Catalyst

  33. Ozone Kinetics Stratospheric ozone is very important. It absorbs short-wavelength (~ 3x10-7m) ultraviolet (UV) radiation from the sun which is one of the strongest mutagens. O2 (g) + O(g) O3 (g) [formation] O3 (g) + O(g) 2 O2 (g) [breakdown] UV photon O2 (g) 2 O(g) UVA UV photon O3 (g) O2 (g) + O(g) UVB

  34. Ozone Kinetics: Catalysis UV photon . . CF2Cl2 (g) CF2Cl(g) + Cl(g) UVA

  35. O3 (g) + Cl(g) ClO(g) + O2 (g) ClO(g) + O(g) Cl(g) + O2 (g) CFCs & Depletion of Ozone . . UV photon Freon -12 CF2Cl2 (g) CF2Cl(g) + Cl(g) UVA . . . . O3 (g) + O(g) 2 O2 (g) The • represents an unpaired electron, resulting from homolytic bond breaking, and the resultant molecules are respectively called a “free radical”. They are very reactive.

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