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# New chm 152 unit 1 power points sp13

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### New chm 152 unit 1 power points sp13

1. 1. CHEMICAL KINETICS RATES OF REACTION Chapter 16.1-16.6,16.8 – Silberberg
2. 2. Objectives & Concepts      See the following Learning Objectives on pages 712. Understand these Concepts: 16.1-3, 16.5-15 Master these Skills: 16.2-6, 16.8
3. 3. Chemical Kinetics  The study of rates of chemical reactions and the mechanisms by which they occur.
4. 4. Reaction Rate    the increase in concentration of a product per unit time. or the decrease in concentration of a reactant per unit time.
5. 5. Reaction Mechanism  The pathway by which a reaction occurs.
6. 6. Expressing the Reaction Rate reaction rate - changes in the concentrations of reactants or products per unit time reactant concentrations decrease while product concentrations increase A for rate of reaction = - B change in concentration of A change in time - (conc A) t =- conc A2-conc A1 t2-t1
7. 7. Consider the reaction   A(g) ----> B(g) + C(g) Rate of reaction = - [A] t    = [B] = t [C] t
8. 8. Table 16.1 Concentration of O3 at Various Times in its Reaction with C2H4 at 303 K Concentration of O3 Time (s) (mol/L) 0.0 3.20x10- 10. 0 20.0 5 2.42x105 1.95x10- 30.0 5 1.63x10- 40.0 5 1.40x10- 50.0 5 1.23x10- 60.0 5 1.10x105 [C2H4] rate = t =- [O3] t
9. 9. Figure 16.5 - Three types of reaction rates for the reaction of O3 and C2H4.
10. 10. Figure 16.6A Plots of [reactant] and [product] vs. time. C2H4 + O3 → C2H4O + O2 [O2] increases just as fast as [C2H4] decreases. Rate = = [C2H4] t [C2H4O] t == [O3] t [O2] t
11. 11. Figure 16.6B Plots of [reactant] and [product] vs. time. H2 + I2 → 2HI [HI] increases twice as fast as [H2] decreases. Rate = - Rate = [H2] t [IH] t =- = -2 [I2] t [H2] t =1 2 = -2 [HI] t [I2] t The expression for the rate of a reaction and its numerical value depend on which substance serves as the reference.
12. 12. Factors That Influence Reaction Rate   Particles must collide in order to react. The higher the concentration of reactants, the greater the reaction rate.   The physical state of the reactants influences reaction rate.   A higher concentration of reactant particles allows a greater number of collisions. Substances must mix in order for particles to collide. The higher the temperature, the greater the reaction rate.  At higher temperatures particles have more energy and therefore collide more often and more effectively.
13. 13. Figure 16.3 The effect of surface area on reaction rate. A hot steel nail glows feebly when placed in O2. The same mass of steel wool bursts into flame.
14. 14. Collision Theory and Concentration The basic principle of collision theory is that particles must collide in order to react. An increase in the concentration of a reactant leads to a larger number of collisions, hence increasing reaction rate. The number of collisions depends on the product of the numbers of reactant particles, not their sum. Concentrations are multiplied in the rate law, not added.
15. 15. Figure 16.13 - The number of possible collisions is the product, not the sum, of reactant concentrations. add another 6 collisions 4 collisions add another 9 collisions
16. 16. Temperature and the Rate Constant Temperature has a dramatic effect on reaction rate. For many reactions, an increase of 10 C will double or triple the rate. Experimental data shows that k increases exponentially as T increases. This is expressed in the Arrhenius equation: k = Ae -Ea/RT Higher T larger k k = rate constant A = frequency factor Ea = activation energy increased rate
17. 17. Activation Energy In order to be effective, collisions between particles must exceed a certain energy threshold. When particles collide effectively, they reach an activated state. The energy difference between the reactants and the activated state is the activation energy (Ea) for the reaction. The lower the activation energy, the faster the reaction. Smaller Ea larger f larger k increased rate
18. 18. Figure 16.15Energy-level diagram for a reaction. Collisions must occur with sufficient energy to reach an activated state. This particular reaction is reversible and is exothermic in the forward direction.
19. 19. Temperature and Collision Energy An increase in temperature causes an increase in the kinetic energy of the particles. This leads to more frequent collisions and reaction rate increases. At a higher temperature, the fraction of collisions with sufficient energy equal to or greater than Ea increases. Reaction rate therefore increases.
20. 20. Figure 16.16 The effect of temperature on the distribution of collision energies.
21. 21. Molecular Structure and Reaction Rate For a collision between particles to be effective, it must have both sufficient energy and the appropriate relative orientation between the reacting particles. The term A in the Arrhenius equation is the frequency factor for the reaction. k = Ae -Ea/RT A = pZ p = orientation probability factor Z = collision frequency The term p is specific for each reaction and is related to the structural complexity of the reactants.
22. 22. Figure 16.18The importance of molecular orientation to an effective collision. NO(g) + NO3(g) → 2NO2(g There is only one relative orientation of these two molecules that leads to an effective collision.
23. 23. Transition State Theory An effective collision between particles leads to the formation of a transition state or activated complex. The transition state is an unstable species that contains partial bonds. It is a transitional species partway between reactants and products. Transition states cannot be isolated. The transition state exists at the point of maximum potential energy. The energy required to form the transition state is the activation energy.
24. 24. Figure 16.19 - The transition state of the reaction between BrCH3 and OH-. BrCH3 + OH- → Br- + CH3OH The transition state contains partial bonds (dashed) between C and Br and between C and O. It has a trigonal bypyramidal shape.
25. 25. Figure 16.20 - Depicting the reaction between BrCH3 and OH-.
26. 26. Figure 16.21 - Reaction energy diagrams and possible transition states for two reactions.
27. 27. Catalysis: Speeding up a Reaction A catalyst is a substance that increases the reaction rate without itself being consumed in the reaction. In general, a catalyst provides an alternative reaction pathway that has a lower total activation energy than the uncatalyzed reaction. A catalyst will speed up both the forward and the reverse reactions. A catalyst does not affect either H or the overall yield for a reaction.
28. 28. Figure 16.23 - Reaction energy diagram for a catalyzed (green) and uncatalyzed (red) process.
29. 29. Figure 16.24 - The catalyzed decomposition of H2O2. A homogeneous catalyst is in the same phase as the reaction. A small amount of NaBr is added to a solution of H2O2. Oxygen gas forms quickly as Br-(aq) catalyzes the H2O2 decomposition; the intermediate Br2 turns the solution orange.
30. 30. Figure 16.25 The metal-catalyzed hydrogenation of ethene. A heterogeneous catalyst is in a different phase than the reaction mixture.
31. 31. Rate Law Expression  The rate law expression must be determined experimentally. They cannot be written down by merely looking at the balanced chemical equation.
32. 32. Rate Law Expression    For the reaction aA + bB -----> Rate = k[A]m[B]n      cC + dD where: k = specific rate constant m = order of the reaction with respect to A n = order of the reaction with respect to B n+m = overall order of the reaction
33. 33. Reaction Orders A reaction has an individual order “with respect to” or “in” each reactant. For the simple reaction A → products: If the rate doubles when [A] doubles, the rate depends on [A]1 and the reaction is first order with respect to A. If the rate quadruples when [A] doubles, the rate depends on [A]2 and the reaction is second order with respect to [A]. If the rate does not change when [A] doubles, the rate does not depend on [A], and the reaction is zero order with respect to A.
34. 34. Figure 16.7 Plots of reactant concentration, [A], vs. time for first-, second-, and zero-order reactions.
35. 35. Figure 16.8 Plots of rate vs. reactant concentration, [A], for first-, second-, and zero-order reactions.
36. 36. Table 16.5 An Overview of Zero-Order, First-Order, and Simple Second-Order Reactions Zero Order First Order Second Order Rate law rate = k rate = k[A] rate = k[A]2 Units for k mol/L·s 1/s L/mol·s Half-life [A]0 2k Integrated rate law in straight-line form [A]t = -kt + [A]0 ln[A]t = -kt + ln[A]0 Plot for straight line [A]t vs. t ln[A]t vs. t Slope, y intercept -k, [A]0 -k, ln[A]0 ln 2 k 1 k[A]0 1 1 = kt + [A]t [A]0 1 vs. t [A]t k, 1 [A]0
37. 37. Rate Law Expression      For the reaction 2N2O5(g) -----> 4NO2(g) + O2(g) Experimentally it was determined Rate = k[N2O5] The reaction is ____order in N2O5 and _____order overall.
38. 38. Rate Law Expression      For the reaction (CH3)3CBr(aq) + OH-(aq) -----> (CH3)3COH(aq) + Br-(aq) Experimentally it was determined that the Rate = k[CH3)3CBr] The reaction is ___order in (CH3)3CBr , ____order in OH- and ___order overall.
39. 39. Rate Law Expression      For the reaction 2NO(g) + O2(g) -----> 2NO2(g) It was experimentally determined that Rate = k[NO]2[O2] The reaction is ____order in NO , ___order in O2 and _____order overall.
40. 40. Sample Problem 16.2 Determining Reaction Order from Rate Laws PROBLEM: For each of the following reactions, determine the reaction order with respect to each reactant and the overall order from the given rate law. (a) 2NO(g) + O2(g) (b) CH3CHO(g) 2NO2(g); rate = k[NO]2[O2] CH4(g) + CO(g); rate = k[CH3CHO]3/2 (c) H2O2(aq) + 3I-(aq) + 2H+(aq) I3-(aq) + 2H2O(l); rate = k[H2O2][I-] PLAN: Look at the rate law and not the coefficients of the chemical reaction. SOLUTION: (a) The reaction is 2nd order in NO, 1st order in O2, and 3rd order overall. (b) The reaction is 3/2 order in CH3CHO and 3/2 order overall. (c) The reaction is 1st order in H2O2, 1st order in I- and zero order in H+, while being 2nd order overall.
41. 41. Experimental Determination of the Rate Law  The order of the reaction for each reactant can be determined by changing individually the concentration of each reactant and observing the new rate of the reaction.
42. 42. Experimental Determination of the Rate Law       Order Rate 0 1 2 3 4 Change [A] x2 x2 x2 x2 x2 Change in 0 x2 x4(22) x8(23) x16(24)
43. 43. Example       The following rate data were obtained at 25oC for the following reaction. Determine the rate law expression and the specific rate constant for the reaction. 2A(g) + B(g) -----> 3C(g) Experiment [A]o [B]o RateM/s 1 0.10M 0.10M 2.0x10-4 2 0.20M 0.10M 4.0x10-4 3 0.10M 0.20M 2.0x10-4
44. 44. Example  The following data were obtained for the following reaction at 25oC. Determine the rate law and the specific rate constant for the reaction.
45. 45. Example(cont.)  2A(g) + B(g) + 2C(g) -----> 3D(g) + 2E(g)  Exp [A]o [B]o [C]o Rate(M/s)     1 0.20M 0.10M 0.10M 2.0x10-4 2 3 4 0.20M 0.20M 0.60M 0.30M 0.10M 0.10M 0.20M 0.30M 0.40M 6.0x10-4 2.0x10-4 6.0 x10-4
46. 46. Example        For a reaction having the rate law expression Rate = k[A][B] fill in the blanks of the following data sheet. Exp. [A] [B] Rate(Ms-1) 1 0.20M 0.050M 4.0x10-3 2 3 _____ 0.40M 0.050M _____ 1.6x10-2 3.2x10-2
47. 47. Example         For the reaction 2A(g) + 3B(g) ---> C(g) + 2D(g) fill in the blank for the following table Exp. [A]o [B]o Rate(M/s) 1 2 3 4 0.10M 0.20M 0.10M _____ 0.10M 0.10M 0.20M 0.30M 2.0x10-3 8.0x10-3 4.0x10-3 24x10-3
48. 48. The Integrated Rate Equation    The integrated rate equations will vary depending on the overall order of the reaction. For a first order, first order overall reaction the equation will be ln[A]/[A]o = -kt where time is time, [A]o is the [A] at t=0 and [A] is [A] at time t
49. 49. The Integrated Rate Equation    Use the integrated rate equation to determine the time required for one-half of the concentration on A to be consumed. This time is called the half-life. t1/2 = 0.693/k
50. 50. Reaction Halflife The half-life (t1/2) for a reaction is the time taken for the concentration of a reactant to drop to half its initial value. For a first-order reaction, t1/2 does not depend on the starting concentration. t1/2 = ln 2 k = 0.693 k The half-life for a first-order reaction is a constant. Radioactive decay is a first-order process. The half-life for a radioactive nucleus is a useful indicator of its stability.
51. 51. Examples    The first order decomposition of cyclopropane to propene is a first order, first order overall process with k=9.2s-1 at 1000oC. Determine the half-life of cyclpropane at 1000oC. Determine the amount of a 3.0 g sample of cyclopropane remaining after 0.50 seconds. Determine the time required for a 3.0 g sample to be reduced to 0.80g.
52. 52. Examples   The half-life for the following first order reaction is 688 hours at 1000oC. Determine the specific rate constant, k, at 1000oC and the amount of a 3.0 g sample of CS2 that remains after 48 hours. CS2(g) -----> CS(g) + S(g)