Presentation Transcript

1. Kinetics

2. Brainstorm: What are some factors that affect the rate of a reaction?
3. Reaction Rate Change in [reactants] or [products] per unit time (M/s) Always expressed as a positive quantity Can be expressed as appearance of B OR disappearance of A A  B
4. Changes in reaction rate Typical for the rate to decrease as the reaction proceeds Instantaneous rate = tangent to graph Initial rate = instantaneous rate at t = 0
5. Reaction Rate Problem Consider the hypothetical aqueous reaction: A  B. Calculate the average rate of disappearance of A for each 10-min interval, in units of M/s.
6. Calculating Instantaneous Rate of Reaction Using Figure 14.4 from the book, calculate the instantaneous rate of disappearance of C4H9Cl at t = 0 (the initial rate)
7. Reaction Rates and Stoichiometry What if stoichiometric relationships are not one to one? 2 HI (g)  H2 (g) + I2 (g)
8. Summary of Stoichiometry and Rates For the reaction a A + b B c C + d D Rate = - 1 [A] = - 1 [B] = 1 [C] = 1 [D] a t b t c t d t
9. Reaction Rates How is the rate at which ozone disappears related to the rate at which oxygen appears in the reaction 2 O3 (g)  3 O2 (g)? If the rate at which O2 appears is 6.0 x 10-5 M/s at a particular instant, at what rate is O3 disappearing at this same time?
10. The Rate Law An equation that shows how the rate depends on the reactants k = the rate constant (changes only with temperature aA + bB  cC + dD Rate = k[A]m[B]n m and n are NOT stoichiometric coefficients) They are determined experimentally
11. Units on Rate Constants Depends on overall reaction order What are the units of the rate constant for the following rate laws? Rate = k[N2O5] Rate = k[H2][I2]
12. Reaction order The exponents m and n in the rate law Overall reaction order = sum of exponents in the rate law DETERMINED EXPERIMENTALLY Rate = k[H2][I2] Reaction = 1st order in H2 1st order in I2 2nd order overall (1 + 1)
13. Determining a Rate Law What is the rate law for the following reaction: 2 NO (g) + 2 H2 (g)  N2 (g) + 2 H2O (g)
14. First Order Reactions Rate depends on the concentration of a single reactant raised to the first power A  products Rate = - [A] = k[A] t
15. Integrated Rate Law for 1st order reactions ln [A]t – ln[A]0 = -kt ln[A]t = -kt [A]0
16. Integrated rate law in y = mx + b format ln [A]t = -kt + ln [A]0
17. Using the Integrated First-Order Rate Law The decomposition of a certain insecticide in water follows first order kinetics with a rate constant of 1.45 yr-1 at 12ºC. A quantity of this insecticide is washed into a lake on June 1, leading to a concentration of 5.0 x 10-7 g/cm3. Assume that the average temperature of the lake is 12ºC. What is the concentration of the insecticide on June 1 of the following year? How long will it take for the concentration to decrease to 3.0 x 10-7 g/cm3?
18. Solving the Problem
19. How can you tell if a reaction is first order? Graph time on the x-axis, and ln [A]t on the y-axis If the graph is a straight line, the reaction is first order in A slope = -k y-intercept = ln [A]0
20. Second Order Reactions Rate depends on: the reactant concentration raised to the second power OR on the concentration of two different reactants, each raised to the first power Rate = -[A] = k[A]2 t
21. Integrated Rate Law for Second Order Reactions = y = mx + b format 1 kt + 1 [A]t [A]0
22. Determining Reaction Order from the Integrated Rate Law
23. Determining Reaction Order from the Integrated Rate Law
24. Determining Reaction Order from the Integrated Rate Law
25. Half Life The time required for the concentration of a reactant to reach one-half of its initial value Half life of a first order reaction: t1/2 = 0.693 k
26. Determining the half-life of a first –order reaction Use Figure 14.4 in the book to estimate the half-life for that reaction. Use the half-life to calculate the rate constant.
27. Temperature and Rate Most reaction rates increase with temperature The rate constant increases Collision model – molecules must collide to react More collisions/sec = higher reaction rate Higher temperature = higher speed = more forceful collision = more energy in collision
28. Orientation of molecules Molecules must be oriented in a certain way Allows certain bonds to break/form Example:
29. Activation Energy (Ea) Minimum amount of energy needed for molecules to react and initiate a chemical reaction Kinetic energy of collision is used to change the potential energy of the molecule At the top of the energy barrier, an activated complex forms (also called the transition state) Lower Ea = faster reaction rate
30. Arrhenius Equation k = Ae-Ea/RT Ea = activation energy (Joules) R = gas constant = 8.31 J/mol-K T = temperature (Kelvin) A = frequency factor (nearly constant) Increase in rate with increasing temperature is generally nonlinear Three factors affect reaction rates The fraction of molecules possessing an energy of Ea or greater The number of collisions occurring per second The fraction of collisions that have the correct orientation
31. What does the Arrhenius Equation really mean? k = Ae-Ea/RT As activation energy increases, reaction rate decreases
32. Analyzing the Arrhenius Equation by graph ln k = - Ea + ln A RT
33. Another way of looking at the Arrhenius Equation Subtract one from the other: lnk1= - Ea+ ln A RT1 lnk2= - Ea+ ln A RT2 Rearranged: ln k1 = Ea 1 - 1 k2 R T2 T1
34. Determining the energy of activation From these data, calculate the activation energy for the reaction What is the value of the rate constant at 430.0 K?
35. Reaction Mechanisms The process by which a reaction occurs Elementary reaction – occurs in a single event (step) Molecularity – the number of molecules that participate as reactants in an elementary reaction Unimolecular Bimolecular Termolecular ???
36. Example of an Elementary Reaction
37. Multistep Mechanisms Most reactions involve multiple steps Each step is an elementary reaction Sum of these elementary reactions is the overall process
38. Multistep Mechanism Potential Energy C A B D Reaction Pathway How many intermediates are formed in the reaction A – D? How many transition states are there? Which step is the fastest? Is the reaction endothermic or exothermic?
39. Rate Laws for Elementary Reactions
40. Rate-Determining Step for Multistep Mechanisms One step is usually much slower than the others – this is called the rate-determining step The overall rate of reaction cannot exceed this rate This determines the rate law for the overall reaction
41. Slow Initial Step The rate of the overall reaction depends on the initial step These molecules in the initial step end up in the overall rate law A fast initial step makes it more difficult to determine the rate law
42. Rate Law for Multistep Mechanism The following mechanism has been proposed for the gas-phase reaction of H2 with Icl: H2 (g) + ICl (g)  HI (g) + HCl (g) HI (g) + ICl (g)  I2 (g) + HCl (g) Write the balanced equation for the overall reaction. Identify intermediates. Write rate laws for each elementary reaction. If the first setup is slow and the second one is fast, what rate law do you expect to be observed for the overall reaction?
43. Rate Law for Multistep Mechanism The reaction 2 NO (g) + Cl2 (g)  2 NOCl (g) obeys the rate law, rate = k[NO]2[Cl2]. The following mechanism has been proposed: NO (g) + Cl2 (g)  NOCl2 (g) NOCl2 (g) + NO (g)  2 NOCl (g) What would be the rate law if the first step were rate determining? What can we conclude about the relative rates of the two steps?