15 February 2012 1 Objective You will be

15 February 2012 1 Objective: You will be able to: define “kinetics” and identify factors that affect the rate of a reaction. write rate expressions for balanced chemical reactions.

Agenda Do now II. Kinetics notes III. Reaction Rates Demonstrations IV. Rate constant and reaction rates problems. Homework: p. 602 #2, 3, 5, 7, 12, 13, 15, 16, 18: Thurs. I.

3 Chemical Kinetics

Aspects of Chemistry 4 How can we predict whether or not a reaction will take place? Thermodynamics Once started, how fast does the reaction proceed? Chemical kinetics: this unit! How far will the reaction go before it stops? Equilibrium: next unit

Chemical Kinetics The area of chemistry concerned with the speeds, or rates, at which a chemical reaction occurs. reaction rate: the change in the concentration of a reactant or product with time (M/s) Why do reactions have such very different rates? Steps in vision: 10 -12 to 10 -6 seconds! Graphite to diamonds: millions of years! In chemical industry, often more important to maximize the speed of a reaction, not necessarily yield.

A D[A] rate = Dt D[B] rate = Dt 6 B

Chemical Kinetics Reaction rate is the change in the concentration of a reactant or a product with time (M/s). A B D[A] rate = Dt D[A] = change in concentration of A over time period Dt D[B] rate = Dt D[B] = change in concentration of B over time period Dt Because [A] decreases with time, D[A] is negative.

red-brown Br 2 (aq) + HCOOH (aq) 2 Br- (aq) + 2 H+ (aq) + CO 2 (g) time t 1< t 2 < t 3 393 nm light 8 Detector D[Br 2] a D Absorption

Br 2 (aq) + HCOOH (aq) 2 Br- (aq) + 2 H+ (aq) + CO 2 (g) slope of tangent [Br 2]final – [Br 2]initial D[Br 2] average rate = =Dt tfinal - tinitial 9 instantaneous rate = rate for specific instance in time

Factors that Affect Reaction Rates 10 Concentration of reactants: higher concentrations = faster reactions as concentration increases, the frequency of collisions increases, increasing reaction rate Temperature: increasing temperature increases reaction rate because of increased KE Physical state of reactants: homogeneous mixtures of either liquids or gases react faster than heterogeneous mixtures Presence of a catalyst: affects the kinds of collisions that lead to a reaction.

Question and Demo 11 Mine explosions from the ignition of powdered coal dust are relatively common, yet lumps of coal burn without exploding. Explain.

Reaction Rates and Stoichiometry 2 A B Two moles of A disappear for each mole of B that is formed. 1 D[A] rate = 2 Dt a. A + b. B D[B] rate = Dt c. C + d. D 1 D[A] 1 D[C] 1 D[D] 1 D[B] rate = =a Dt b Dt c Dt d Dt 12

Example Write the rate expression for the following reaction: CH 4 (g) + 2 O 2 (g) CO 2 (g) + 2 H 2 O (g)

Write the rate expression for the following reaction: CH 4 (g) + 2 O 2 (g) CO 2 (g) + 2 H 2 O (g) D[CH 4] D[CO 2] 1 D[H 2 O] rate = = == Dt Dt Dt 2 14

Practice Problems Write the rate expressions for the following reactions in terms of the disappearance of the reactants and appearance of products. a. I-(aq) + OCl-(aq) + OI-(aq) b. 4 NH 3(g) + 5 O 2(g) 4 NO(g) + 6 H 2 O(g)

rate a [Br 2] rate = k [Br 2] rate = rate constant k= [Br 2] = 3. 50 x 10 -3 s-1

Using Rate Expressions Consider the reaction: 4 NO 2(g) + O 2(g) 2 N 2 O 5(g) Suppose that, at a particular moment during the reaction, molecular oxygen is reacting at the rate of 0. 024 M/s. a. At what rate is N 2 O 5 being formed? b. At what rate is NO 2 reacting?

16 February 2012 Objective: You will be able to: solve rate expressions. determine the order of a reaction from experimental data Homework Quiz: N 2(g) + 3 H 2(g) → 2 NH 3(g) Suppose that at a particular moment during the reaction, hydrogen is reacting at the rate of 0. 074 M/s. a. At what rate is NH 3 being formed? b. At what rate is nitrogen reacting? 18

Agenda Do now II. Iodine clock reaction. III. Solving rate equations IV. Determining order of reactions Homework: p. 602 #15, 16, 18, 19, 20: Mon after break Hint: Use pressure just like concentration. Diagnostic test (Tues after break) I.

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Example Consider the reaction: 4 PH 3(g) P 4(g) + 6 H 2(g) Suppose that, at a particular moment during the reaction, molecular hydrogen is being formed at the rate of 0. 078 M/s. a. At what rate is P 4 being formed? b. At what rate is PH 3 reacting?

Problem 22 Consider the reaction between gaseous hydrogen and gaseous nitrogen to produce ammonia gas. At a particular time during the reaction, H 2(g) disappears at the rate of 3. 0 M/s. a. What is the rate of disappearance of N 2(g)? b. What is the rate of appearance of NH 3(g)?

23 If ammonia appears at 2. 6 M/s, how fast does hydrogen disappear?

The Rate Law The rate law is a mathematical relationship that shows how rate of reaction depends on the concentrations of reactants a. A + b. B c. C + d. D Rate = k [A]x[B]y x and y are small whole numbers that relate to the number of molecules of A and B that collide and are determined experimentally!

The Rate Law a. A + b. B c. C + d. D Rate = k [A]x[B]y Reaction is xth order in A Reaction is yth order in B Reaction is (x +y)th order overall Rate = k [A]1[B]2

Example 26 Experime nt 1 2 [A](M) [B](M) Rate = −d[A]/dt (M/s) 0. 10 0. 20 0. 04 0. 08 3 0. 20 0. 32 What is the numerical value of the rate constant for the reaction described in the table above? Specify units.

F 2 (g) + 2 Cl. O 2 (g) 2 FCl. O 2 (g) rate = k [F 2]x[Cl. O 2]y Double [F 2] with [Cl. O 2] constant Rate doubles rate = k [F 2][Cl. O 2] x=1 Quadruple [Cl. O 2] with [F 2] constant Rate quadruples y=1

Write the reaction rate expressions for the following in terms of the disappearance of the reactants and the appearance of products: a) 2 H 2(g) + O 2(g) 2 H 2 O(g) b) 4 NH 3(g) + 5 O 2(g) 4 NO(g) + 6 H 2 O(g)

Consider the reaction N 2(g) + 3 H 2(g) 2 NH 3(g) Suppose that at a particular moment during the reaction molecular hydrogen is reacting at a rate of 0. 074 M/s. a) At what rate is ammonia being formed? b) At what rate is molecular nitrogen reacting?

27 February 2012 Take Out: p. 602 #15, 16, 18, 19, 20 Objective: You will be able to determine the rate of a reaction given experimental data and reactant concentrations. Homework Quiz: What is the rate law for the reaction shown below? What is the rate when [A]=1. 50 M and [B]=0. 50 M? 30 Run # Initial [A] ([A]0) Initial [B] ([B]0) Initial Rate (v 0) 1 1. 00 M 1. 25 x 10 -2 M/s 2 1. 00 M 2. 5 x 10 -2 M/s 3 2. 00 M 2. 5 x 10 -2 M/s

Agenda 31 Homework Quiz II. Homework answers III. Determining and solving rate laws IV. Hand back tests and assignments Homework: Diagnostic test revisit/correct p. 603 #15, 16, 18 I.

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 not related to the stoichiometric coefficient of the reactant in the balanced chemical equation. F 2 (g) + 2 Cl. O 2 (g) 2 FCl. O 2 (g) rate = k [F 2][Cl. O 2] 1 32

Determine the rate law and calculate the rate constant for the following reaction from the following data: S 2 O 82 - (aq) + 3 I- (aq) 2 SO 42 - (aq) + I 3 - (aq) Experiment [S 2 O 82 -] [I-] Initial Rate (M/s) 1 0. 08 0. 034 2. 2 x 10 -4 2 0. 08 0. 017 1. 1 x 10 -4 3 0. 16 0. 017 2. 2 x 10 -4

Determine the rate law and calculate the rate constant for the following reaction from the following data: S 2 O 82 - (aq) + 3 I- (aq) 2 SO 42 - (aq) + I 3 - (aq) Experiment [S 2 O 82 -] [I-] Initial Rate (M/s) 1 0. 08 0. 034 2. 2 x 10 -4 2 0. 08 0. 017 1. 1 x 10 -4 3 0. 16 0. 017 2. 2 x 10 -4 rate = k [S 2 O 82 -]x[I-]y y=1 x=1 rate = k [S 2 O 82 -][I-] Double [I-], rate doubles (experiment 1 & 2) Double [S 2 O 82 -], rate doubles (experiment 2 & 3) 34 2. 2 x 10 -4 M/s rate k= = = 0. 08/M • s 2[S 2 O 8 ][I ] (0. 08 M)(0. 034 M)

Practice Problems The reaction of nitric oxide with hydrogen at 1280 o. C: 2 NO(g) + 2 H 2(g) N 2(g) + 2 H 2 O(g) From the following data collected at this temperature, determine (a) the rate law, (b) the rate constant and (c) the rate of the reaction when [NO] = 12. 0 x 10 -3 M and [H 2] = 6. 0 x 10 -3 M 35 Experiment [NO] M [H 2] M Initial Rate (M/s) 1 5. 0 x 10 -3 2. 0 x 10 -3 1. 3 x 10 -5 2 10. 0 x 10 -3 2. 0 x 10 -3 5. 0 x 10 -5 3 10. 0 x 10 -3 4. 0 x 10 -3 10. 0 x 10 -5

Calculate the rate of the reaction at the time when [F 2] = 0. 010 M and [Cl. O 2] = 0. 020 M. F 2(g) + 2 Cl. O 2(g) 2 FCl. O 2(g) [F 2] (M) [Cl. O 2] (M) Initial Rate (M/s) 0. 10 0. 010 1. 2 x 10 -3 0. 10 0. 040 4. 8 x 10 -3 0. 20 0. 010 2. 4 x 10 -3

Consider the reaction X + Y Z From the following data, obtained at 360 K, a) determine the order of the reaction b) determine the initial rate of disappearance of X when the concentration of X is 0. 30 M and that of Y is 0. 40 M 37 Initial Rate of Disappearance of X (M/s) [X] (M) [Y] (M) 0. 053 0. 10 0. 50 0. 127 0. 20 0. 30 1. 02 0. 40 0. 60 0. 254 0. 20 0. 60 0. 509 0. 40 0. 30

38 Consider the reaction A B. The rate of the reaction is 1. 6 x 10 -2 M/s when the concentration of A is 0. 35 M. Calculate the rate constant if the reaction is a. first order in A b. second order in A

The rate laws can be used to determine the concentrations of any reactants at any time during the course of a reaction.

29 Nov. 2010 Take Out Homework p. 603 #19, 21, 22, 23, 2529 Objective: SWBAT compare 1 st order, 2 nd order, and zero order reactions, and describe how temperature and activation energy effect the rate constant. Do now: Calculate the half life of the reaction F 2(g) 40 + 2 Cl. O 2(g) 2 FCl. O 2(g), with rate data shown below: [F 2] (M) [Cl. O 2] (M) Initial Rate (M/s) 0. 10 0. 010 1. 2 x 10 -3 0. 10 0. 040 4. 8 x 10 -3 0. 20 0. 010 2. 4 x 10 -3

28 February 2012 Take Out: Diagnostic test Objective: You will be able to determine order of a reaction and k graphically. Homework Quiz: What is the rate law for the reaction shown below? What is the rate when [A]=1. 50 M and [B]=0. 50 M? 41 Run # Initial [A] ([A]0) Initial [B] ([B]0) Initial Rate (v 0) 1 1. 00 M 1. 25 x 10 -2 M/s 2 1. 00 M 2. 5 x 10 -2 M/s 3 2. 00 M 2. 5 x 10 -2 M/s

Agenda 42 I. III. Homework Quiz 1 st order reactions graphically Half life calculations Homework: p. 603 #19, 20 (use Excel!), 24, 26

First Order (Overall) Reactions 43 rate depends on the concentration of a single reactant raised to the first power. rate = k[A] = Using calculus, this rate law is transformed into an equation for a line: ln[A] = ln[A]0 - kt

First-Order Reactions A k= product D[A] rate = Dt rate M/s = = 1/s or s-1 M [A] = [A]0 e−kt rate = k [A] D[A] = k [A] Dt ln[A] = ln[A]0 - kt

Graphical Determination of k 2 N 2 O 5 4 NO 2 (g) + O 2 (g)

A non-graphical example 46 The reaction 2 A B is first order in A with a rate constant of 2. 8 x 10 -2 s-1 at 800 C. How long will it take for A to decrease from 0. 88 M to 0. 14 M ?

The reaction 2 A B is first order in A with a rate constant of 2. 8 x 10 -2 s-1 at 800 C. How long will it take for A to decrease from 0. 88 M to 0. 14 M ? [A]0 = 0. 88 M ln[A] = ln[A]0 - kt [A] = 0. 14 M kt = ln[A]0 – ln[A] = t= k 47 ln [A]0 [A] k ln = 0. 88 M 0. 14 M 2. 8 x 10 -2 s-1 = 66 s

The conversion of cyclopropane to propene in the gas phase is a first order reaction with a rate constant of 6. 7 x 10 -4 s-1 at 500 o. C. a) b) c) If the initial concentration of cyclopropane was 0. 25 M, what is the concentration after 8. 8 minutes? How long, in minutes, will it take for the concentration of cyclopropane to decrease from 0. 25 M to 0. 15 M? How long, in minutes, will it take to convert 74% of the starting material?

29 February 2012 Objective: You will be able to: calculate the half-life of a first order reaction explore the relationship between time and concentration of a second order reaction Homework Quiz: The conversion of cyclopropane to propene in the gas phase is a first order reaction with a rate constant of 6. 7 x 10 -4 s-1 at 500 o. C. If the initial concentration of cyclopropane was 0. 25 M, what is the concentration after 8. 8 minutes? 49

The rate of decomposition of azomethane (C 2 H 6 N 2) is studied by monitoring partial pressure of the reactant as a function of time: CH 3 -N=N-CH 3(g) → N 2(g) + C 2 H 6(g) The data obtained at 300 o. C are shown here: Time (s) Partial Pressure of Azomethane (mm. Hg) 0 284 100 220 150 193 200 170 250 150 300 132 Are these values consistent with first-order kinetics? If so, determine the rate constant.

51 The following gas-phase reaction was studied at 290 o. C by observing the change in pressure as a function of time in a constant-volume vessel: Cl. CO 2 CCl 3(g) 2 COCl 2(g) Determine the order of the reaction and the rate constant based on the following data, where P is the total pressure Time (s) P (mm. Hg) 0 400 2, 000 316 4, 000 248 6, 000 196 8, 000 155 10, 000 122

Ethyl iodide (C 2 H 5 I) decomposes at a certain temperature in the gas phase as follows: C 2 H 5 I(g) → C 2 H 4(g) + HI(g) From the following data, determine the order of the reaction and the rate constant: Time (min) [C 2 H 5 I] (M) 0 0. 36 15 0. 30 30 0. 25 48 0. 19 75 0. 13

First-Order Reactions The half-life, t½, is the time required for the concentration of a reactant to decrease to half of its initial concentration. t½ = t when [A] = [A]0/2 ln t½ = [A]0/2 k ln 2 0. 693 = = k k What is the half-life of N 2 O 5 if it decomposes with a rate constant of 5. 7 x 10 -4 s-1? How do you know decomposition is first order?

First-Order Reactions The half-life, t½, is the time required for the concentration of a reactant to decrease to half of its initial concentration. t½ = t when [A] = [A]0/2 ln t½ = [A]0/2 k ln 2 0. 693 = = k k What is the half-life of N 2 O 5 if it decomposes with a rate constant of 5. 7 x 10 -4 s-1? 0. 693 t½ = ln 2 = = 1200 s = 20 minutes -4 -1 k 5. 7 x 10 s How do you know decomposition is first order? 54 units of k (s-1)

First-order reaction A product # of half-lives [A] = [A]0/n 1 2 55 2 4 3 8 4 16

56 The decomposition of ethane (C 2 H 6) to methyl radicals is a first-order reaction with a rate constant of 5. 36 x 10 -4 s-1 at 700 o. C: C 2 H 6(g) 2 CH 3(g) Calculate the half-life of the reaction in minutes.

57 Calculate the half-life of the decomposition of N 2 O 5: 2 N 2 O 5 4 NO 2(g) + O 2(g) t (s) [N 2 O 5] (M) ln [N 2 O 5] 0 0. 91 -0. 094 300 0. 75 -0. 29 600 0. 64 -0. 45 1200 0. 44 -0. 82 3000 0. 16 -1. 83

Second-Order Reactions A product D[A] rate = Dt rate M/s = = 1/M • s k= 2 2 M [A] 1 1 = + kt [A]0 58 D[A] = k [A]2 Dt [A] is the concentration of A at any time t [A]0 is the concentration of A at time t=0 t½ = t when [A] = [A]0/2 1 t½ = k[A]0 rate = k [A]2

Iodine atoms combine to form molecular iodine in the gas phase: I(g) + I(g) I 2(g) This reaction follows second-order kinetics and has the high rate constant 7. 0 x 109/M·s at 23 o. C. a. If the initial concentration of I was 0. 086 M, calculate the concentration after 2. 0 minutes. b. Calculate the half-life of the reaction if the initial concentration of I is 0. 60 M and if it is 0. 42 M.

The reaction 2 A → B is second order with a rate constant of 51/M·min at 24 o. C. a. b. Starting with [A]o = 0. 0092 M, how long will it take for [A]t = 3. 7 x 10 -3 M? Calculate the half-life of the reaction.

1 March 2012 61 Objective: You will be able to: determine the activation energy for a reaction Homework Quiz: The reaction 2 A → B is second order with a rate constant of 51/M·min at 24 o. C. a. b. Starting with [A]o = 0. 0092 M, how long will it take for [A]t = 3. 7 x 10 -3 M? Calculate the half-life of the reaction.

Agenda 62 Homework Quiz II. Questions? III. Kinetics Quiz IV. Activation Energy Homework: p. I.

Zero-Order Reactions A product D[A] rate = Dt D[A] =k Dt rate = M/s k= 0 [A] = [A]0 - kt [A] is the concentration of A at any time t [A]0 is the concentration of A at time t = 0 t½ = t when [A] = [A]0/2 [A]0 t½ = 2 k 63 rate = k [A]0 = k

Summary of the Kinetics of Zero-Order, First-Order and Second-Order Reactions Order 0 64 Rate Law rate = k 1 rate = k [A] 2 [A]2 rate = k Concentration-Time Equation [A] = [A]0 - kt Half-Life t½ = [A]0 2 k ln[A] = ln[A]0 - kt t½ = ln 2 k 1 1 = + kt [A]0 1 t½ = k[A]0

65 Activation Energy and Temperature Dependence of Rate Constants Reaction rates increase with increasing temperature Ex: Hard boiling an egg Ex: Storing food How do reactions get started in the first place?

Collision Theory 66 Chemical reactions occur as a result of collisions between reacting molecules. reaction rate depends on concentration But, the relationship is more complicated than you might expect! Not all collisions result in reaction

Question 67 Explain in terms of collision theory why temperature affects rate of reaction.

So, when does the reaction happen? 68 In order to react, colliding molecules must have a total KE ≥ activation energy (Ea) Ea: minimum amount of energy required to initiate a chemical reaction activated complex (transition state): a temporary species formed by the reactant molecules as a result of the collision before they form the product.

A+B Exothermic Reaction + + AB C+D Endothermic Reaction The activation energy (Ea ) is the minimum amount of energy required to initiate a chemical reaction. =a barrier that prevents less energetic molecules from reacting

Rate Constant is Temp. Dependent 70 Arrhenius equation Ea is the activation energy (J/mol) R is the gas constant (8. 314 J/K • mol) T is the absolute temperature A is the frequency factor

Alternate Arrhenius Equation 71 To relate k at two temperatures, T 1 and T 2:

The rate constants for the decomposition of acetaldehyde: CH 3 CHO(g) → CH 4(g) + CO(g) were measured at five different temperatures. The data are shown below. Plot lnk versus 1/T, and determine the activation energy (in k. J/mol) for the reaction. (Note: the reaction is order in CH 3 CHO, so k has the units of ) k 0. 011 0. 035 0. 105 T (K) 700 730 760 0. 343 0. 789 790 810

Determining Graphically 0. 001185 0. 001235 0. 001285 0. 001335 0. 001385 0. 001435 0. 001485 -0. 50 -1. 00 -1. 50 ln K -2. 00 -2. 50 -3. 00 -3. 50 -4. 00 -4. 50 -5. 00 1/T (K-1) slope = -2. 19 x 104 slope =

Determining activation energy 74 The second order rate constant for the decomposition of nitrous oxide (N 2 O) into nitrogen molecule and oxygen atom has been measured at different temperatures. Determine graphically the activation energy for the reaction. o k 1. 87 x 10 -3 0. 0113 0. 0569 T ( C) 600 650 700 0. 244 750

5 March 2012 75 Objective: You will be able to: review and correct answers to the multiple choice questions on the diagnostic test. Homework Quiz: Please use the same sheet of paper as last week!

Agenda 76 Homework Quiz II. Homework answers III. Correct and explain answers to diagnostic test multiple choice questions. Homework: Finish correcting and explaining answers to multiple choice: due Weds. I.

With one partner: 77 Check your answers to the multiple choice against my answers on the board. For each question you answered incorrectly, or skipped, or guessed and happened to get it right: Write 1 to 2 sentences to explain why the correct answer is correct. Use resources! Textbook, notes, internet…

7 March 2012 78 Objective: You will be able to: review, correct and explain answers to the free response questions on the diagnostic test. Do now: Look at your free response 1 -6 and decide on your first three preferences for creating a poster and explaining your answers. Write them down on your slip of paper.

Agenda 79 I. II. Objective and agenda Correct and explain answers to diagnostic test free response questions

With your group… 80 1. 2. 3. 4. Check your answers with the answer key. Make notes about how to solve the problem/answer the question. Design and create a poster that shows the work and answers, as well as additional explanations of how to solve the problem or answer the question. Post your poster in the room! Then, go look at other groups posters and correct your work.

30 Nov. 2010 Take Out Homework p. 605# 31, 32, 35, 37, 39 Objective: SWBAT use the Arrhenius equation to solve for rate constants and temperatures, and solve practice problems on kinetics. Do now: Match 81 Order 2 1 0 Rate Law rate = k[A]2 rate = k Conc-Time Eq. [A]=[A]0 -kt 1/[A]=1/[A]0 + kt ln[A]=ln[A]0 –kt Half Life Eq. t 1/2=1/k[A]o t 1/2=ln 2/k t 1/2=[A]0/2 k

Agenda 82 Homework solutions II. Using the Arrhenius equation part 2 III. Molecular orientation IV. Problem Set work time Homework: Complete problem set and p. 605 #40, 42 Quiz tomorrow I.

8 March 2012 83 Objective: You will be able to: review, correct and explain answers to the free response questions on the diagnostic test. describe the reaction mechanism of a reaction Do now: Finish and hang up your poster. (10 min. )

Agenda 84 Objective and agenda II. Gallery Walk: Correct and explain answers to diagnostic test free response questions III. Using the Alternate Arrhenius Equation IV. Hand back quizzes Homework p. 605 #44, 45, 49, 51, 52, 54: Mon. I.

Gallery Walk 85 Walk with your group Spend 5 minutes at each station Correct/complete your work and make notes of how/why each problem is solved.

Using the alternate Arrhenius Equation 86 The rate constant of a first order reaction is 3. 46 x 10 -2 /s at 298 K. What is the rate constant at 350 K if the activation energy for the reaction is 50. 2 k. J/mol?

Using the Arrhenius Equation 87 The first order rate constant for the reaction of methyl chloride (CH 3 Cl) with water to produce methanol (CH 3 OH) and hydrochloric acid (HCl) is 3. 32 x 10 -10/s at 25 o. C. Calculate the rate constant at 40 o. C if the activation energy is 116 k. J/mol.

88 Frequency of Collisions and Orientation Factor For simple reactions (between atoms, for example) the frequency factor (A) is proportional to the frequency of collision between the reacting species. “Orientation factor” is also important.

Importance of Molecular Orientation effective collision ineffective collision 89

Reaction Mechanisms 90 A balanced chemical equation doesn’t tell us much about how the reaction actually takes place. It may represent the sum of elementary steps Reaction mechanism: the sequence of elementary steps that leads to product formation.

Reaction Mechanisms The overall progress of a chemical reaction can be represented at the molecular level by a series of simple elementary steps or elementary reactions. The sequence of elementary steps that leads to product formation is the reaction mechanism. 2 NO (g) + O 2 (g) 2 NO 2 (g) N 2 O 2 is detected during the reaction! 91 Elementary step: NO + NO N 2 O 2 + Elementary step: N 2 O 2 + O 2 2 NO 2 Overall reaction: 2 NO + O 2 2 NO 2

2 NO (g) + O 2 (g) Mechanism: 92 2 NO 2 (g)

13 March 2012 93 Objective: You will be able to identify overall reactions, intermediates and rate laws for reaction mechanisms.

Agenda 94 Objectives and Agenda II. Review: Reaction mechanisms III. Elementary step examples IV. Catalysts Homework: p. 605 #44, 45, 49, 51, 52, 54, 55, 56, 61: Tues. I.

Intermediates are species that appear in a reaction mechanism but not in the overall balanced equation. An intermediate is always formed in an early elementary step and consumed in a later elementary step. Elementary step: NO + NO N 2 O 2 + Elementary step: N 2 O 2 + O 2 2 NO 2 Overall reaction: 2 NO + O 2 2 NO 2 The molecularity of a reaction is the number of molecules reacting in an elementary step. • Unimolecular reaction – elementary step with 1 molecule • Bimolecular reaction – elementary step with 2 molecules • Termolecular reaction – elementary step with 3 molecules 95

Rate Laws and Elementary Steps Unimolecular reaction A products rate = k [A] Bimolecular reaction A+B products rate = k [A][B] Bimolecular reaction A+A products rate = k [A]2 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. The rate-determining step is the slowest step in the sequence of steps leading to product formation. 96

The experimental rate law for the reaction between NO 2 and CO to produce NO and CO 2 is rate = k[NO 2]2. The reaction is believed to occur via two steps: Step 1: NO 2 + NO 2 NO + NO 3 Step 2: NO 3 + CO NO 2 + CO 2 What is the equation for the overall reaction? What is the intermediate? What can you say about the relative rates of steps 1 and 2?

The experimental rate law for the reaction between NO 2 and CO to produce NO and CO 2 is rate = k[NO 2]2. The reaction is believed to occur via two steps: Step 1: NO 2 + NO 2 NO + NO 3 Step 2: NO 3 + CO NO 2 + CO 2 What is the equation for the overall reaction? NO 2+ CO NO + CO 2 What is the intermediate? NO 3 What can you say about the relative rates of steps 1 and 2? 98 rate = k[NO 2]2 is the rate law for step 1 so step 1 must be slower than step 2

Rate Determining Step 99 rate determining step: the slowest step in the sequence of steps leading to product formation.

Problem 100 Propose a mechanism for the overall reaction: 2 A + 2 B → A 2 B 2

Example The gas-phase decomposition of nitrous oxide (N 2 O) is believed to occur via two elementary steps: Step 1: N 2 O N 2 + O Step 2 N 2 O + O N 2 + O 2 Experimentally the rate law is found to be rate = k[N 2 O]. a) Write the equation for the overall reaction. b) Identify the intermediates. c) What can you say about the relative rates of steps 1 and 2? 101

102 NO 2 + F 2 → NO 2 F + F NO 2 + F → NO 2 F a. b. c. d. Write the overall reaction. What is the intermediate? If the rate law is rate = k[NO 2][F 2], which step is the rate determining step? Which step proceeds at the fastest rate?

103 Hydrogen and iodine monochloride react as follows: H 2(g) + 2 ICl(g) → 2 HCl(g) + I 2(g) The rate law for the reaction is rate = k[H 2][ICl]. Suggest a possible mechanism for the reaction.

Decomposition of Hydrogen Peroxide 104 2 H 2 O 2(aq) 2 H 2 O(l) + O 2(g) Can be catalyzed using iodide ions (I-) rate = k[H 2 O 2][I-] Why? ! Determined experimentally. Step 1: H 2 O 2 + I- H 2 O + IOStep 2: H 2 O 2 + IO- H 2 O + O 2 + I-

105 For the decomposition for H 2 O 2, the reaction rate depends on the concentration of I- ions, even though Idoesn’t appear in the overall equation. I- is a catalyst for the reaction.

A catalyst is a substance that increases the rate of a chemical reaction without itself being consumed. Ea Catalyzed k Uncatalyzed ratecatalyzed > rateuncatalyzed 106 Ea′ < Ea

Catalysts 107 forms an alternative reaction pathway lowers overall activation energy for example, it might form an intermediate with the reactant. Ex: 2 KCl. O 3(s) 2 KCl(s) + 3 O 2(g) Very slow, until you add Mn. O 2, a catalyst. The Mn. O 2 can be recovered at the end of the reaction!

Week of March 12 108 Step 1: HBr + O 2 → HOOBr Step 2: HOOBr + HBr → 2 HOBr Step 3: HOBr + HBr → H 2 O + Br 2 Step 4: HOBr + HBr → H 2 O + Br 2 a. Write the equation for the overall reaction. b. Identify the intermediate(s). c. What can you say about the relative rate of each step if the rate law is rate =

13 March 2012 109 Objective: You will be able to identify and describe the effect of catalysts in a reaction mechanism. Agenda: I. Homework Quiz II. Homework Answers III. Catalysts IV. Problem Set Homework: Problem Set: Monday

Catalyst Example: Ozone Cycle 110 Step 1: O 2(g) + hv → O(g) + O(g) Step 2: O(g) + O 2(g) → O 3(g) Step 3: O 3(g) + hv → O 2(g) + O(g) Step 4: O(g) + O(g) → O 2(g) Overall: O 3(g) + O 2(g) → O 2(g) + O 3(g) This cycle continually repeats, producing and destroying ozone at the same rate while absorbing harmful ultraviolet light from the sun.

Chlorofluorocarbons and Ozone Chlorine atoms from CFCs released into the atmosphere catalyze the O 3(g) → O 2(g) reaction. Net result: ozone is depleted faster that is generated by the natural cycle. Cl atoms from CFCs deplete the ozone layer! Step 1: 2 Cl(g) + 2 O 3(g) → 2 Cl. O(g) + 2 O 2(g) Step 2: Cl. O(g) + Cl. O(g) → O 2(g) + 2 Cl(g) Overall: 2 O 3(g) → 3 O 2(g) 111

In heterogeneous catalysis, the reactants and the catalysts are in different phases (usually, catalyst is a solid, reactants are gases or liquids). • Haber synthesis of ammonia • Ostwald process for the production of nitric acid • Catalytic converters In homogeneous catalysis, the reactants and the catalysts are dispersed in a single phase, usually liquid. 112 • Acid catalysis • Base catalysis

Haber Process Synthesis of Ammonia Extremely slow at room temperature. Must be fast and high yield! Process occurs on the surface of the Fe/Al 2 O 3/K 2 O catalyst, which weakens the covalent N-N and H-H bonds. Fe/Al 2 O 3/K 2 O N 2 (g) + 3 H 2 (g) catalyst 113 2 NH 3 (g)

Ostwald Process 4 NH 3 (g) + 5 O 2 (g) Pt catalyst 2 NO (g) + O 2 (g) 2 NO 2 (g) + H 2 O (l) 114 4 NO (g) + 6 H 2 O (g) 2 NO 2 (g) HNO 2 (aq) + HNO 3 (aq) Pt-Rh catalysts used in Ostwald process

Catalytic Converters CO + Unburned Hydrocarbons + 2 NO 2 115 catalytic O 2 converter catalytic converter CO 2 + H 2 O 2 N 2 + 3 O 2

Enzyme Catalysis biological catalysts 116

Binding of Glucose to Hexokinase 117

14 March 2012 118 Objective: You will be able to: demonstrate your knowledge of chemical kinetics on a problem set and a lab. Agenda: I. Objectives and Agenda II. Work time: I. Problem Set II. Kinetics Pre-Lab

AP Exam 119 Monday, May 7 If you have a year average >80%, you pay $13 (full cost = $87!) This is due, in CASH (no coins), by next Friday. If your average is <80%, I’ll chat with you privately today about your options.

Homework 120 Pre-lab: due tomorrow Lab procedure: read by tomorrow Problem set: due Monday Kinetics test: Tuesday

Expectations 121 Choose ONE person to work with. Work either on the problem set or the prelab questions (or split your time…) Stay at your table. Use a professional tone and volume of voice. Use this time wisely!

15 March 2012 122 Sit at a lab table with your group. Take Out: Lab notebook and lab packet Objective: You will be able to: determine the rate law and the activation energy for the oxidation of iodide ions by bromate ions in the presence of an acid.

Homework 123 Problem Set due Monday Kinetics Unit Test Tuesday Gas Unit revisions due tomorrow

Logistics 124 Half of the groups will do Part 1 on page 5 while the other half does steps 1 -3 on page 6. Then, we’ll switch!

Changes to the Procedure 125 Instead of reaction strips, you’ll be using spot plates. Instead of inverting one reaction strip over the other and shaking down to mix, you’ll be adding the drops of KBr. O 3, starting the stopwatch, and stirring with a toothpick to mix. You must do this at the same way, in the same order, and at the same speed each time!

126 Put the reagents for reaction strip 1 in one well plate. If more than 2 drops of KBr. O 3, place the drops in a second well plate. Transfer them with a separate pipette so you can dispense them all at once into the first well plate. Start timing and stir.

Precision and Consistency 127 Be very precise in your work, or your results won’t be meaningful. Be very consistent in the way your carry out the procedure: the way you hold the pipette to drop solutions, the way you add the KBr. O 3 (from “reaction strip 2”), the rate at which you stir, when you start and stop timing, etc.

Reagents and Equipment 128 Leave reagents at the front table. Bring your labeled pipettes to the table to fill them.

Data 129 Record your data immediately and carefully in tables in your lab notebook.

19 March 2012 130 Objective: You will be able to: determine the reaction order, rate law, and activation energy for an iodine clock reaction. Reminder: $13 (cash) due by Friday for AP Exam

Homework 131 Problem Set due today Kinetics Test tomorrow 10 MC 1 -2 FRQ

What’s the purpose? 132

22 March 2012 133 Objective: You will be able to: determine the rate law, reaction constant and activation energy for the iodine clock reaction.

Agenda 134 Finish lab II. Clean up/return materials III. Work on lab calculations, analysis and conclusions in your lab notebook Note: all data, etc. must also be in your lab notebook! Homework: Lab notebook due Monday $13 for AP Exam due by 8: 00 am TOMORROW!!! I.

Water baths 135 Warm water bath (40 o. C) on the side bench. If it’s too cool, remove some water, and add some hot water from the beaker on the hot plate. It should be shallow! Don’t swamp your spot plate. Record the actual temp. Ice bath (Oo. C): create one using ice and water in your metal pan. Use a little thermometer to record the temperature.

Safety 136 Keep your goggles on your eyes! One warning Then you’re out. Label your reagents and store them carefully. Use a professional voice and stay at your table unless you need to get something.

Cleanup 137 Keep your labeled pipettes in the cassette case. Rinse transfer pipettes in water and squirt out water to dry. Return equipment to the cart. Make sure your lab table is clean and neat.