1 / 60

CHEMICAL KINETICS CHAPTER 17, Kinetics Fall 2009, CHEM 1310

CHEMICAL KINETICS CHAPTER 17, Kinetics Fall 2009, CHEM 1310. KeKin. Kinetics vs Thermodynamics A: Reactants B: Transition state C: products E: Forward Activation Free Energy F: Reverse Activation Free Energy. Reaction Mechanisms.

june
Download Presentation

CHEMICAL KINETICS CHAPTER 17, Kinetics Fall 2009, CHEM 1310

An Image/Link below is provided (as is) to download presentation Download Policy: Content on the Website is provided to you AS IS for your information and personal use and may not be sold / licensed / shared on other websites without getting consent from its author. Content is provided to you AS IS for your information and personal use only. Download presentation by click this link. While downloading, if for some reason you are not able to download a presentation, the publisher may have deleted the file from their server. During download, if you can't get a presentation, the file might be deleted by the publisher.

E N D

Presentation Transcript

  1. CHEMICAL KINETICSCHAPTER 17, KineticsFall 2009, CHEM 1310

  2. KeKin Kinetics vs Thermodynamics A: Reactants B: Transition state C: products E: Forward Activation Free Energy F: Reverse Activation Free Energy

  3. Reaction Mechanisms • Most reactions occur through several steps and are not single step reactions. • Each step in a multi-step reaction is called an elementary reaction. Types of elementary reactions 1. Unimolecular (a single reactant) 2. Bimolecular 3. Termolecular(very unlikely)

  4. Each step of this reaction is an “elementary step”. Each elementary step has reactant(s), a transition state, and product(s). Products that are consumed in subsequent elementary reaction are called intermediates.

  5. Reaction Rates: To measure a reaction rate we could monitor the disappearance of reactants or appearance of products. e.g., 2NO2 + F2→ 2NO2F

  6. Gen. Rxn: aA + bB → cC + dD NO2 + CO → NO + CO2

  7. Order of a ReactionThe power (n) to which the concentration of A is raised in the rate expression describes the order of the reaction with respect to A. Do not confuse the order (n) with the stoichiometric coefficient (a).

  8. mthorder in [A] nth order in [B]

  9. [A] (mol L-1)[B] (mol L-1) Rate (mol L-1 s-1) 1 1.0x10-4 1.0x10 -4 2.8x10 -6 2 1.0x10-4 3.0x10 -4 8.4x10 -6 3 2.0x10-4 3.0x10 -4 3.4x10 -5

  10. [A] (mol L-1)[B] (mol L-1) Rate (mol L-1 s-1) 1 1.0x10-4 1.0x10 -4 2.8x10 -6 2 1.0x10-4 3.0x10 -4 8.4x10 -6 3 2.0x10-4 3.0x10 -4 3.4x10 -5

  11. Example: At elevated temperatures, HI reacts according to the chemical equation 2HI → H2 + I2 The rate of reaction increases with concentration of HI, as shown in this table. Data [HI] Rate Point (mol L-1) (mol L-1 s-1) • 0.005 7.5 x 10-4 • 0.010 3.0 x 10-3 • 0.020 1.2 x 10-2 a) Determine the order of the reaction with respect to HI and write the rate expression b) Calculate the rate constant and give its units c) Calculate the instantaneous rate of reaction for a [HI] = 0.0020M

  12. INTEGRATED RATE LAWS • Single Reactant (three cases) • Zero-Order Rate Law (n = 0) • First-Order Rate Law (n = 1) • Second-Order Rate Law (n = 2) • More than one Reactant • Must state the order of the reaction with respect to each reactant (rate = k[A]n[B]m[C]p)

  13. INTEGRATED RATE LAWS n=0,1,2

  14. In the real world, if we do not know the order of the reaction we can use experimental plots to estimate the order. If a plot of [A] vst is a straight line, then the reaction is zero order. If a plot of ln[A] vst is a straight line, then the reaction is 1st order. If a plot of 1/ [A] vst is a straight line, then the reaction is 2nd order.

  15. INTEGRATED RATE LAWS Zero Order Reactions [A] - [Ao] = -kt Graph [A] vst Slope = -k, intercept = [Ao]

  16. INTEGRATED RATE LAWS First Order Reactions ln[A]-ln[Ao] = -kt Graph ln[A] vst Slope = -k, intercept = [Ao]

  17. 2 N2O5 (g) → 4 NO2 (g) + O2 (g) Slope = - k This graph gives a straight line, and so is First order with respect to the decomposition of N2O5 If a plot of ln[A] vs t is not a straight line, the reaction is not first order! In[N205] versus time.

  18. Dimerization Data set provided [C4H6] vs time 2 C4H6 (g) → C8H12 (g) [C4H6]˚ = 0.01M

  19. Reaction Mechanisms • Most reactions proceed not through a single step but through a series of steps • Each Step is called an elementary reaction Types of elementary reactions • Unimolecular (a single reactant) E.g., A → B + C (a decomposition) • Bimolecular(most common type) E.g., A + B → products • Termolecular (less likely event) E.g., A + B + C → products

  20. Notice that NO3 is formed and consumed. This is called a __________________________________. Notice also that Step 1 is bimolecular and Step 2 is bimolecular

  21. CHEMICAL EQUILIBRIUM A direct connection exists between the equilibrium constant of a reaction and the rate constants. a) at equilibrium: forward reaction rate = reverse reaction rate. b) Keq = kf/ kr(same as K = k1/k-1) kf A⇌ B kr

  22. REACTION MECHANISM & RATE LAWS Typically with a reaction one of several elementary step reaction is the slowest step. This is called the Rate Determining Step (RDS) Example 15.6 Case #1:When the RDS occurs first, the first step is slow and determines the rate of the overall reaction.

  23. Energy Reaction Progress fast slow F + NO2F NO2F NO2+ F2

  24. Chem 1310 Spring 2009 stop here

  25. Energy Reaction Progress

  26. Need to express [intermediates] in terms of other reactants

  27. Substituting for [N2O2] in the rate expression above

  28. slow Energy fast N2O2 + O2 2NO 2NO2 Reaction Progress

  29. Energy Reaction Progress • Reaction Mechanism • Intermediates • Transition states

  30. A MODEL FOR CHEMICAL KINETICS

  31. Chapter 5: The Kinetic Molecular Theory of Gases The Meaning of Temperature: temperature is a measure of the average kinetic energy of the gas particles. The Kelvin temperature of a gas is a measure of the random motions of the particles of gas. With higher temperature, greater motion.

  32. Chapter 5: Speed Distribution Curves Maxwell-Boltzmann speed distribution Temperature is a measure of the average kinetic energy of molecules when their speeds have Maxwell Boltzmann distribution. i.e., the molecules come to thermal equilibrium.

  33. Transition State, also called Activated Complex Two requirements must be satisfied for reactants to collide successfully to rearrange to form products

  34. Number of collisions Consider two different temperatures.1.) Collisions must have enough energy to produce a reaction. Not all collisions have enough energy to make productEcollision > Eact Distribution of velocities

  35. 2.) Molecular OrientationRelative orientations of the reactants must allow formation of any new bonds to produce products. 2 BrNO (g)→ 2 NO (g) + Br2(g) Orientation a or b lead to product, c does not.

  36. y = mx + b Find the rate constant k at several temperatures. Plot of In(k) versus 1/Tfor the reaction Slope =

  37. Eaf Ear Energy Reaction Progress Transition State The Activation Energy (Ea) is the minimum collision energy that reactants must have in order to form products

  38. Eaf Energy Reaction Progress Transition State Ear ΔE = Eaf - Ear The Activation Energy (Ea) is the minimum collision energy that reactants must have in order to form products

  39. CHEMICAL KINETICS Catalyst Inhibitor

More Related