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1 Today… Turn in: Our Plan: Homework (Write in Planner): Nothing
Test Results Videos/Notes Investigation 10 Pre-Lab Homework (Write in Planner): Be prepared for Investigation 10 next class and have report started.

2 Introduction to Chemical Kinetics
What is Kinetics? I’ll let Hank explain… Stop at 3:20

3 Kinetics in Action… Clock Reactions…

4 Unit Preview… Some Reactions
Are fast: Acid/base neutralization Sodium + water PPT reactions Are slow: Aluminum oxidation Iron oxidizing Plastic bottle decomposing

5 And some depend… On temperature: fireflies
On conditions: Iron oxide (humid or desert) On enzymes: many processes in our body On a catalyst: 2CO NO --> 2CO2 + N2 (automobile pollution and catalytic converters)

6 Chemical Kinetics: A Preview
Chemical kinetics is the study of: the rates of chemical reactions factors that affect these rates the mechanisms by which reactions occur Reaction rates vary greatly – some are very fast (burning, precipitation) and some are very slow (rusting, disintegration of a plastic bottle in sunlight).

7 Variables in Reaction Rates
Concentrations of reactants: Reaction rates generally increase as the concentrations of the reactants are increased. Temperature: Reaction rates generally increase rapidly as the temperature is increased. Surface area: For reactions that occur on a surface rather than in solution, the rate increases as the surface area is increased. Catalysts: Catalysts speed up reactions and inhibitors slow them down.

8 Think of it like this…

9 Theories of Chemical Kinetics: Collision Theory
Before atoms, molecules, or ions can react, they must first collide. An effective collision between two molecules puts enough energy into key bonds to break them. The activation energy (Ea) is the minimum energy that must be supplied by collisions for a reaction to occur. A certain fraction of all molecules in a sample will have the necessary activation energy to react; that fraction increases with increasing temperature. The spatial orientations of the colliding species may also determine whether a collision is effective.

10 Distribution of Kinetic Energies
At higher temperature (red), more molecules have the necessary activation energy.

11 Key Point Orientation of molecules at the time of their collision will determine whether they react or not!

12 Analogy - Car Crash Example
Energy and orientation of cars during a car crash can establish the change that occurs to the cars.

13 Importance of Orientation
One hydrogen atom can approach another from any direction … Effective collision; the I atom can bond to the C atom to form CH3I … and reaction will still occur; the spherical symmetry of the atoms means that orientation does not matter. Ineffective collision; orientation is important in this reaction.

14 Transition State Theory
The configuration of the atoms of the colliding species at the time of the collision is called the transition state. The transitory species having this configuration is called the activated complex. A reaction profile shows potential energy plotted as a function of a parameter called the progress of the reaction. Reactant molecules must have enough energy to surmount the energy “hill” separating products from reactants.

15 Activated complex Species formed as a result of collisions between energetic molecules that is an intermediate between the reactants and the products of a reaction. Once formed the activated complex dissociates either into products or back to the reactants

16 Reaction Profile A graphical representation of a chemical reaction in terms of the energies of the reactants, activated complexes and products

17 Reaction Profile

18 Exothermic

19 Endothermic

20 A Reaction Profile CO(g) + NO2(g)  CO2(g) + NO(g)

21 An Analogy for Reaction Profiles and Activation Energy

22 Stop! Investigation 10 - We’re going to test the effect of different variables on the rate of reaction next class. Complete steps 1 – 6 on the lab handout and be prepared to conduct the lab next class.

23 Today… Turn in: Our Plan: Homework (Write in Planner): Nothing
Investigation 10 Homework (Write in Planner): Lab Report Due Friday

24 Today… Turn in: Our Plan: Homework (Write in Planner):
Lab Report (rubric on top) Our Plan: Notes & Practice Homework (Write in Planner): Work on the Ch. 13 Homework

25 The Meaning of Rate The rate of a reaction is the change in concentration of a product per unit of time (rate of formation of product). Rate is also viewed as the negative of the change in concentration of a reactant per unit of time (rate of disappearance of reactant). The rate of reaction often has the units of moles per liter per unit time (mol∙L–1∙s–1 or M∙s–1)

26 If the rate of consumption of H2O2 is 4.6 M/h, then …
… the rate of formation of H2O must also be 4.6 M/h, and … … the rate of formation of O2 is 2.3 M/h

27 2 H2O2 --> 2 H2O + O2 1 L 0.1850 mol H2O2/L Rate = 60 s
2.960 g O2 ( mole) produced in 60 s means … … mol H2O2 reacted in 60 s. = M H2O2 s–1

28 Consider the hypothetical reaction
Example 13.1 Consider the hypothetical reaction A + 2B --> 3C + 2D Suppose that at one point in the reaction, [A] = M and 125 s later [A] = M. During this time period, what is the average (a) rate of reaction expressed in M∙s–1 and (b) rate of formation of C, expressed in M∙min–1.

29 Try It Out! EX 13.1 A: Consider the hypothetical reaction
2A + B → 2C + D Suppose that at some point during the reaction [D] = M and that 2.55 min (that is 2 min 33 sec) later [D] = M. What is the average rate of reaction during this time period, expressed in M min-1? What is the rate of formation of C expressed in M s-1?

30 Rate of Reaction Expressed as the negative of the Slope of a Tangent Line
Average rate (green dotted line) Initial Rate ( blue solid line) Instantaneous Rate (red line) Question: Over what time interval are the instantaneous rates greater than the average rate measured for the 600 sec period?

31 Average vs. Instantaneous Rate
Instantaneous rate is the slope of the tangent to the curve at a particular time. We often are interested in the initial instantane-ous rate; for the initial concentrations of reactants and products are known at this time.

32 Use data from Table 13.1 and/or Figure 13.5 to
Example 13.2 Use data from Table 13.1 and/or Figure 13.5 to determine the initial rate of reaction and calculate [H2O2] at t = 30 s.

33 Determine the instantaneous rate of reaction at t = 300 s.
Try It Out EX 13.2 A: From Figure 13.5, Determine the instantaneous rate of reaction at t = 300 s. Use the result of a) to calculate a value of [H2O2] at t = 310 s.

34 The Rate Law of a Chemical Reaction
The rate law for a chemical reaction relates the rate of reaction to the concentrations of reactants. aA + bB + cC …→ products rate = k[A]n[B]m[C]p … The exponents (m, n, p…) are determined by experiment. Exponents are not derived from the coefficients in the balanced chemical equation, though in some instances the exponents and the coefficients may be the same. The value of an exponent in a rate law is the order of the reaction with respect to the reactant in question. The proportionality constant, k, is the rate constant.

35 Rate Law Examples Rate = k[A]1 = k[A] Reaction is first order in A
Rate = k[A]2 Reaction is second order in A Rate = k[A]3 Reaction is third order in A If we triple the concentration of A in a second-order reaction, the rate increases by a factor of ________.

36 Effect of Concentration on Rate
Order of Rxn Concentration Change Effect on Rate Double/Triple… Nothing (20) 1 Double Double (21) Triple Triple (31) 2 Quadruple (22) 8x (32) Quadruple 16x (42) 3 9x (23) 27x (33)

37 More About the Rate Constant k
The rate of a reaction is the change in concentration with time, whereas the rate constant is the proportionality constant relating reaction rate to the concentrations of reactants. The rate constant remains constant throughout a reaction, regardless of the initial concentrations of the reactants. The rate and the rate constant have the same numerical values and units only in zero-order reactions. For reaction orders other than zero, the rate and rate constant are numerically equal only when the concentrations of all reactants are 1 M. Even then, their units are different.

38 To find the overall order of a reaction…
Add the orders for each compound. Example: rate = [A]2[B]1 is 3rd order overall How about rate = [A]0[B]1[C]1?

39 Overall Reaction Order
Units of k (p. 536) Overall Reaction Order Units of k Zero M s-1 First s-1 Second M-1 s-1 Third M-2 s-1

40 Method of Initial Rates
The method of initial rates is a method of establishing the rate law for a reaction—finding the values of the exponents in the rate law, and the value of k. A series of experiments is performed in which the initial concentration of one reactant is varied. Concentrations of the other reactants are held constant. When we double the concentration of a reactant A, if: there is no effect on the rate, the reaction is zero-order in A. the rate doubles, the reaction is first-order in A. the rate quadruples, the reaction is second-order in A. the rate increases eight times, the reaction is third-order in A.

41 The concentration of NO was held the same in Experiments 1 and 2 …
… while the concentration of Cl2 in Experiment 2 is twice that of Experiment 1. The rate in Experiment 2 is twice that in Experiment 1, so the reaction must be first order in Cl2. Which two experiments are used to find the order of the reaction in NO? How do we find the value of k after obtaining the order of the reaction in NO and in Cl2?

42 Example 13.3 For the reaction 2 NO(g) + Cl2(g) → 2 NOCl(g) described in the text and in Table 13.2, (a) what is the initial rate for a hypothetical Experiment 4, which has [NO] = M and [Cl2] = M? (b) What is the value of k for the reaction?

43 So, when looking at data…..
Zero order reaction: initial rate is unaffected (20 ) 1st order reaction: double the concentration, doubles the rate (21 =2) 2nd Order Reaction: Initial rate increases fourfold (22 =4) 3rd Order Reaction: Initial rate increases eightfold (23 =8)

44 Try It Out Below is some rate data for the hypothetical reaction, 2A + B --> C. What is the rate law for this reaction? Experiment [A]0 [B]0 Rate (M/s) 1 2.0 M 1.0 M 0.100 2 0.400 3 4.0 M

45 With your partner, Complete Part 1 of the Partner Review
Let’s Take a Break With your partner, Complete Part 1 of the Partner Review

46 Sum of exponents is equal to zero
Zero-Order Reactions Sum of exponents is equal to zero

47 Zero Order Rate of Reaction= k[A]0 The concentration time graph is a straight line with a negative slope

48 Zero Order Rate of the reaction: IS equal to k
remains constant throughout the reaction is the negative of the slope of the line when graph Molarity vs. time (see pg. 544)

49 Rate is independent of initial concentration
A Zero-Order Reaction rate = k[A]0 = k Rate is independent of initial concentration

50 Zero Order Integrated rate equation: [A]t = -kt + [A]0 y = mx b

51 y = [A]t = conc of A at some time x = t = time b = [A]0
Zero Order (cont.) y = [A]t = conc of A at some time x = t = time b = [A]0 m = -k (m, the slope of the straight line)

52 First-Order Reactions
In a first-order reaction, the exponent in the rate law is 1. Rate = k[A]1 = k[A] The integrated rate law describes the concentration of a reactant as a function of time. For a first-order process: ln [A]t [A]0 = –kt Look! It’s an equation for a straight line! ln [A]t – ln [A]0 = –kt ln [A]t = –kt + ln [A]0 At times, it is convenient to replace molarities in an integrated rate law by quantities that are proportional to concentration.

53 ln[A]t = -kt + ln[A]0 y = mx + b Rate of reaction= k[A]1
Integrated rate equation ln[A]t = -kt + ln[A]0 y = mx b

54 1st Order Easy test for a first order reaction is to plot the natural log of reactant conc. vs. time and see if the graph is linear.

55 Decomposition of H2O2

56 Example 13.4 For the first-order decomposition of H2O2(aq), given k = 3.66 x 10–3 s–1 and [H2O2]0 = M, determine (a) the time at which [H2O2] = M and (b) [H2O2] after 225 s.

57 Another Example H202 initially at a conc of 2.32 M, is allowed to decompose. What will the [H202] be 1200 s later? Use k = 7.3 x 10-4 s-1 for this first order decomposition.

58 NH2NO2 (aq) → H2O (l) + N2O (g)
Try It Out EX 13.4 A: The decomposition of nitramide, NH2NO2, is a first order reaction: NH2NO2 (aq) → H2O (l) + N2O (g) The rate law is rate = k[NH2NO2], with k = 5.62 x 10-3 min-1 at 15ᵒC. Starting with M NH2NO2, At what time will [NH2NO2] = M What is [NH2NO2] after 6.00 hours?

59 Half-life Time required for ½ of the reactant to be consumed Equation:
t1/2 = ln2 = 0.693 k k

60 Example 13.5 Use data from Figure 13.7 (p. 542) to evaluate the
half-life and Rate constant for the first order decomposition of N2O5 at 67 ᵒC: N2O5 (g) → 2NO2 (g) + ½ O2 (g)

61 Example 13.5 A Use the result of Example 13.5 to determine
The time required to reduce the quantity of N2O5 to 1/16 of its initial value and The mass of N2O5 remaining after a 4.80 g sample of N2O5 has decomposed for 10.0 min.

62 Second-Order Reactions
Sum of exponents = 2

63 Second-Order Reactions
A reaction that is second order in a reactant has a rate law in which the exponent for that reactant is 2. Rate = k[A]2 The integrated rate law has the form: What do we plot vs. time to get a straight line? –––– = kt + –––– [A]t [A]0 The half-life of a second-order reaction depends on the initial concentration as well as on the rate constant k: 1 t½ = ––––– k[A]0

64 Example 13.7 The second-order decomposition of HI(g) at 700 K is represented in Figure 13.9. HI(g) → ½ H2(g) + ½ I2(g) Rate = k[HI]2 What are the: (a) rate constant and (b) half-life of the decomposition of 1.00 M HI(g) at 700 K?

65 Try It Out EX 13.7 A: If, in the second order reaction A → products, it takes 55 s for the concentration of reactant A to fall to 0.40 M from an initial concentration of 0.80 M, what is the rate constant k for the reaction?

66 Summary of Kinetic Data

67 Short Cut!! [A] vs time (straight line = Zero order)
ln [A] vs. time ( straight line = first order) 1/[A] vs. time (straight line = second order)

68 Try it in your note packet!
Time, min [A] ln[A] 1/[A] 1.0 5 0.63 10 0.46 15 0.36 25 0.25

69 Crash Course Review
3:10 – 5:30

70 Complete Part 2 of the Partner Review.
Let’s Take a Break Complete Part 2 of the Partner Review.

71 Today… Turn in: Our Plan: Homework (Write in Planner):
Get out a piece of notebook paper Our Plan: Scavenger Hunt Review Investigation 11 Pre-Lab Homework (Write in Planner): Prepare Lab Report Homework Problems (Due 2/12)

72 Be prepared to experiment next class.
Pre Lab Complete steps 1 – 5 on the handout and have your formal lab report started. Be prepared to experiment next class.

73 Today… Turn in: Our Plan: Homework (Write in Planner): Nothing
Investigation 11 Homework (Write in Planner): Lab Report due Next Class (2/10)

74 Today… Turn in: Our Plan: Homework (Write in Planner):
Lab Report (rubric on top) Our Plan: Review Activity Elephant Toothpaste Demo Notes – Catalysts & Rate Determining Steps Finish Homework Problems Homework (Write in Planner): Homework Problems

75 Effect of Temperature on the Rates of Reactions
In 1889, Svante Arrhenius proposed the following expression for the effect of temperature on the rate constant, k: k = Ae–Ea/RT The constant A, called the frequency factor, is an expression of collision frequency and orientation; it represents the number of collisions per unit time that are capable of leading to reaction. The term e–Ea/RT represents the fraction of molecular collisions sufficiently energetic to produce a reaction.

76 I like this equation better…
Look at Eq on p. 551 of the text!

77 Example 13.9 Estimate a value of k at 375 K for the decomposition of dinitrogen pentoxide illustrated in Figure 13.15, given that k = 2.5 x 10–3 s–1 at 332 K.

78 Try it Out EX 13.9 B: Di-tert-butyl peroxide (DTBP) is used as a catalyst in the manufacture of polymers. In the gaseous state, DTBP decomposes to acetone and ethane by a first-order reaction. (C4H9)2O2 (g) → 2(CH3)2CO (g) + C2H6 (g) The half-life of DTBP is 17.5 h at 125ᵒC and 1.67 h at 145ᵒC. What is the activation energy, Ea, of the decomposition reaction?

79 Reaction Mechanisms Analogy: a banana split is made by steps in sequence: slice banana; three scoops ice cream; chocolate sauce; strawberries; pineapple; whipped cream; end with cherry. A chemical reaction occurs according to a reaction mechanism—a series of collisions or dissociations—that lead from initial reactants to the final products. Like making a banana split, three molecules will not collide simultaneously very often, so steps of a reaction mechanism involve only one or two reactants at a time.

80 Reaction Mechanisms An elementary reaction represents, at the molecular level, a single step in the progress of the overall reaction. A proposed mechanism must: account for the experimentally determined rate law. be consistent with the stoichiometry of the overall or net reaction.

81 Molecularity The molecularity of an elementary reaction refers to the number of free atoms, ions, or molecules that collide or dissociate in that step. Termolecular processes are unusual, for the same reason that three basketballs shot at the same time are unlikely to collide at the same instant …

82 The Rate-Determining Step
The rate-determining step is the crucial step in establishing the rate of the overall reaction. It is usually the slowest step. Some two-step mechanisms have a slow first step followed by a fast second step, while others have a fast reversible first step followed by a slow second step. Fast Mechanism for 2 NO + O2 --> 2 NO2 Slow

83 An Example Given the reaction: 2A + 2B → C + D rate = k[A]2[B]
Could take place by the following three-step mechanism: A + A ↔ X (fast) X + B → C + Y (slow) Y + B → D (fast)

84 Intermediates X & Y are called intermediates because they appear in the mechanism, but they cancel out of the balanced equation. They are products from one reaction and then a reactant in the next.

85 An Example The steps of a reaction mechanism must add up to equal the balanced equation with all intermediates cancelling out. Let’s try it with our example. A + A ↔ X (fast) X + B → C + Y (slow) Y + B → D (fast)

86 Rate-determining Step
As in any process where many steps are involved, the speed of the whole process can’t go faster than the speed of the slowest step in the process. The slowest step of a reaction is the rate-determining step. Because the slowest step is the most important step in determining the rate of a reaction, the slowest step and the steps leading up to it are used to see if the mechanism is consistent with the rate law for the overall reaction.

87 An Example Let’s look at our reaction again and show that it is consistent with the rate law: 2A + 2B → C + D rate = k[A]2[B] A + A ↔ X (fast) X + B → C + Y (slow) Y + B → D (fast)

88 An Example (p. 199 Cracking AP Chemistry Exam)
Write rate law for slow step. X is an intermediate, we need to eliminate it from the rate law. Solve for [X] in terms of [A]. Substitute for [X] in our second step. Now we have the rate law.

89 Fast step: I2 2 I Slow step: 2 I + H2 2 HI
Example For the reaction H2(g) + I2(g) --> 2 HI(g), a proposed mechanism is below. What is the net equation for the overall reaction, and what is the order of the reaction according to this mechanism? Fast step: I I Slow step: 2 I + H HI k1 k–1 k2

90 EX A The decomposition of nitrosyl choride, 2NOCl (g) → 2NO (g) + Cl2 (g) Is a first-order reaction. Propose a mechanism for this reaction consisting of one fast step and one slow step.

91 Catalyst A catalyst provides an alternative reaction pathway of lower activation energy Participates in a chemical reaction w/o undergoing permanent change Speeds up without being consumed in the reaction. It is neither a reactant nor product in a reaction.

92 Catalysis – how does it work
In general, a catalyst works by changing the mechanism of a chemical reaction. Often the catalyst is consumed in one step of the mechanism, but is regenerated in another step. The pathway of a catalyzed reaction has a lower activation energy than that of an uncatalyzed reaction, so more molecules at a fixed temperature have the necessary activation energy.

93 Effect of Catalyst on Reaction Profile and Activation Energy
A catalyst lowers the activation energy, making it easier for the reactants to “climb the energy hill” and form the products.

94 A Kinetics Pick Up Line…

95 Homogeneous Catalysis
Ozone decomposition catalyzed by chlorine atoms has a much lower activation energy and proceeds much more rapidly than the uncatalyzed reaction

96 Heterogeneous Catalysis
Many reactions are catalyzed by the surfaces of appropriate solids. A good catalyst provides a higher frequency of effective collisions. Four steps in heterogeneous catalysis: Reactant molecules are adsorbed. Reactant molecules diffuse along the surface. Reactant molecules react to form product molecules. Product molecules are desorbed (released from the surface).

97 Heterogeneous Catalysis
Hydrogen is adsorbed onto the surface of a nickel catalyst. A C=C approaches … … and is adsorbed. Hydrogen atoms attach to the carbon atoms, and the molecule is desorbed.

98 A Surface-Catalyzed Reaction

99 Enzyme Catalysis Enzymes are high-molecular-mass proteins that usually catalyze one specific reaction, or a set of similar reactions. The reactant substance, called the substrate (S), attaches itself to an area on the enzyme (E) called the active site, to form an enzyme-substrate complex (ES). The enzyme–substrate complex decomposes to form products (P), and the enzyme is regenerated.

100 Factors Influencing Enzyme Activity
The rates of enzyme-catalyzed reactions are influenced by factors such as concentration of the substrate, concentration of the enzyme, acidity of the medium, and temperature.

101 Mercury Poisoning: An Example of Enzyme Inhibition
When Hg reacts with an enzyme … … the Hg binds to sulfur atoms … … changing the shape of the active site, so that it no longer “fits” the substrate.

102 Homogeneous Catalysis
Step 1: A + catalyst → intermediate + C Step 2: B + intermediate → D + catalyst Net equation: A + B → C + D

103 Identifying Catalysts and intermediate species
O Cl∙ --> ClO∙ + O2 ClO· O∙ --> Cl∙ + O2 Catalyst? Intermediate? Net Equation?

104 Crash Course Review
7:15 - end

105 Finish the HW All homework problems are due on Wednesday!

106 Today… Before Class: Our Plan: Homework (Write in Planner):
Mark Homework Questions on the Board Our Plan: Homework Questions/Check HW Worksheet Race Study Guide Homework (Write in Planner): Test Next Class Breakfast Club 6 am on Friday

107 Study Guide Changes 7 – do 39 a only 11 – do 25 a and b only

108 Today… Turn in: Our Plan: Homework (Write in Planner): Nothing
Study Guide Questions Test Homework (Write in Planner): Complete the POGIL (Day 1 ONLY)

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