Top Posters
Since Sunday
5
o
5
4
m
4
b
4
x
4
a
4
l
4
t
4
S
4
m
3
s
3
A free membership is required to access uploaded content. Login or Register.

A Molecular Approach, 4e - Notes for Chapter (14).doc

Uploaded: 4 years ago
Contributor: Lawik
Category: Chemistry
Type: Other
Rating: N/A
Helpful
Unhelpful
Filename:   A Molecular Approach, 4e - Notes for Chapter (14).doc (127.5 kB)
Page Count: 13
Credit Cost: 1
Views: 33
Last Download: N/A
Transcript
Chapter 14. Chemical Kinetics Chapter 14. Chemical Kinetics Chapter 14. Chemical Kinetics Student Objectives 14.1 Catching Lizards Know that physiological processes (like a lizard’s metabolism) depend on temperature. Know that the rate of chemical reactions generally depends on the concentration of the reactants and the temperature of the reaction. 14.2 The Rate of a Chemical Reaction Know that a rate is the change of a particular property with respect to time. Know and understand that the rate of a chemical reaction is a change in concentration measured during a change in time. Know the difference between average and instantaneous rates, and understand how each can be measured or estimated from a plot of concentration vs. time (or from a table of data used to produce such a plot). Calculate an average rate of reaction and predict a change in concentration from a plot of concentration vs. time (or from a table of data used to produce such a plot). Know that the rate of a reaction is measured experimentally and that instruments like a spectrometer or gas chromatograph are used to measure concentrations. 14.3 The Rate Law: The Effect of Concentration on Reaction Rate Know the general definition of a rate law and the meaning and significance of the rate order and rate constant. Know that order of the reaction predicts the dependence of concentration on time. Understand that the order of a reaction must be determined experimentally and not from a balanced reaction. Know that the plot of rate (often in units of M/s) vs. concentration produces a straight horizontal line for zero-order reactions, a straight line with a positive slope for first-order reactions, and a curved line with a positive slope for higher-order reactions. Use the method of initial rates to determine the order of a reaction from a table of concentration and initial rates. Determine and identify the order of a reaction with multiple reactants by observing the effect of a concentration change on the rate of reaction for each independent reactant. 14.4 The Integrated Rate Law: The Dependence of Concentration on Time Understand the difference between rate law and integrated rate law. Know and understand the first-order integrated rate law and that a plot of the natural logarithm of the reactant concentration versus time is linear. Determine the rate constant from the slope of a plot of natural logarithm of reactant concentration versus time for a first-order reaction. Use the first-order integrated rate law to find the reactant concentration at a given time or the time elapsed for a given concentration change. Know and understand the second-order integrated rate law and that a plot of the inverse reactant concentration versus time is linear. Determine the rate constant from the slope of a plot inverse reactant concentration versus time for a second-order reaction. Use the second-order integrated rate law to find the reactant concentration at a given time or the time elapsed for a given concentration change. Know and understand the zero-order integrated rate law and that a plot of the reactant concentration versus time is linear. Determine the rate constant from the slope of a plot of reactant concentration versus time for a zero-order reaction. Use the zero-order integrated rate law to find the reactant concentration at a given time or the time elapsed for a given concentration change. Define half-life and understand how it can be identified from a plot of reactant concentration versus time. Calculate the half-life or predict the concentration of a reactant using the mathematical expressions for first- and second-order reactions. 14.5 The Effect of Temperature on Reaction Rate Know that the temperature dependence of reaction rate is modeled by the Arrhenius equation, an equation that describes the rate constant k in terms of temperature, a frequency factor, and an activation energy. Define activation energy and frequency factor. Know that a transition state or activated complex represent a transient arrangement of atoms that occurs as reactants form products. Draw and interpret a potential energy diagram, a plot of energy vs. reaction progress, including the energy levels of the reactants, transition state, and products. Know that atoms and molecules contain a distribution of energy that depends on temperature and that those exceeding the activation energy can undergo reaction. Calculate the activation energy from a plot of the natural logarithm of the rate constant versus inverse temperature, ln k vs. 1/T. Know and understand the collision model: atoms or molecules need to come in contact or collide in order to react. Know that the collision model accounts for the frequency factor in the Arrhenius equation and that effective collisions require the proper orientation between two reacting molecules. 14.6 Reaction Mechanisms Know that most reactions do not occur in a single step on a molecular level and that the series of steps is called the reaction mechanism. Know that the sum of the elementary steps must be equal to the overall chemical reaction. Know that reaction intermediates are species formed by one elementary step and consumed by another. Recognize unimolecular and bimolecular reactions. Know that the rate-determining step of a mechanism determines the rate law for the overall reaction. Draw and interpret a potential energy diagram, a plot of energy versus reaction progress, for a reaction that involves multiple steps. Understand the rate law resulting from a mechanism in which an intermediate is formed in a fast initial step. Show that a given mechanism is consistent with an observed rate law. 14.7 Catalysis Know that a catalyst increases the rate of a reaction without being consumed by the reaction. Know that catalysts lower the activation energy of a reaction, often by changing the mechanism or the nature of the transition state. Know that an automobile’s catalytic converter modifies exhaust molecules by catalyzing reactions to produce less harmful emissions. Know the definitions of homogeneous and heterogeneous catalysis and understand how they differ. Know that catalysts are involved in important atmospheric reactions including the depletion of the ozone layer and the interconversion of other pollutants. Know the general model for the behavior of enzymes and understand their role as biological catalysts. Section Summaries Lecture Outline Terms, Concepts, Relationships, Skills Figures, Tables, and Solved Examples Teaching Tips Suggestions and Examples Misconceptions and Pitfalls Lecture Outline Terms, Concepts, Relationships, Skills Figures, Tables, and Solved Examples 14.1 Catching Lizards Dependence of activity/metabolism on temperature lizards and ectotherms chemical reactions Intro figure: illustration of a lizard, glass of ice water, and the Arrhenius equation 14.2 The Rate of a Chemical Reaction Rate examples speed (change in distance per time) weight loss (change in weight per time) chemical reactions (change in amount or concentration per time) Expressions of rate mathematical relationships plots of concentration versus time average rate instantaneous rate Measuring rates spectrometer gas chromatograph Figure 14.1 The Rate of a Chemical Reaction Figure 14.2 Reactant and Product Concentrations as a Function of Time unnumbered table: table of data from Figure 14.2 Example 14.1 Expressing Reaction Rates Figure 14.3 The Spectrometer Figure 14.4 The Gas Chromatograph 14.3 The Rate Law: The Effect of Concentration on Reaction Rate Rate laws rate = k[A]n rate constant, k order, n Determining order plot of [A] vs. time for n = 0, 1, 2 plot of rate vs. [A] for n = 0, 1, 2 initial rates zero order: no change in rate first order: linear dependence second order: square dependence general: multiple reactants Figure 14.5 Reactant Concentration as a Function of Time for Different Reaction Orders Figure 14.6 Reaction Rate as a Function of Reactant Concentration for Different Reaction Orders Figure 14.7 Sublimation unnumbered tables: initial rate data for examples Example 14.2 Determining the Order and Rate Constant of a Reaction Teaching Tips Suggestions and Examples Misconceptions and Pitfalls 14.1 Catching Lizards The discussion of lizard activity and metabolism related to temperature is an interesting example of chemical kinetics, though some students may not see the connection immediately. Ask the students to suggest other examples, like the souring of milk left at room temperature. 14.2 The Rate of a Chemical Reaction There are many common examples of rates besides speed and weight loss. Solicit additional ones. Figure 14.1 shows different representations of a fast and slow A + B C reaction using symbols to represent the reactants and products. This can serve as a segue to a more mathematical representation, a plot of concentration vs. time for reactants and products (Figure 14.2). Conceptual Connection 14.1 Reaction Rates Spectrometry and gas chromatography are just two ways to measure quantities of molecules. The language of kinetics can be confusing: rate, rate law, instantaneous rate, average rate, and rate constant all have different meanings. Students often presume that energetically favorable reactions (e.g., formation of water from H2 and O2) are always fast. 14.3 The Rate Law: The Effect of Concentration on Reaction Rate The rate law for the general reaction A products is rate = k[A]n. First-order reactions are much more intuitive for most students than the other reaction orders. Reaction orders are most commonly positive integers, but they can be negative and fractional as well. Fractional orders make little sense to most students until the relevant mechanisms are explored. Conceptual Connection 14.2 Order of Reaction The method of initial rates can be used for any reaction regardless of order because the rate is relatively constant in the first 5–10% of a reaction. Conceptual Connection 14.3 Rate and Concentration The concept of reaction order is initially lost on some students. Reactions of the same order behave kinetically the same, as shown especially in Section 14.4. The order must be determined experimentally. Students are used to placing much emphasis on stoichiometry from previous studies and naturally want to extend that emphasis to reaction order. Lecture Outline Terms, Concepts, Relationships, Skills Figures, Tables, and Solved Examples 14.4 The Integrated Rate Law: The Dependence of Concentration on Time Integrated rate laws first-order plot of ln [A] vs. t slope = k intercept = ln [A]0 second-order plot of 1/[A] vs. t slope = k intercept = a/[A]0 zero-order [A]t = kt + [A]0 Half-life first-order: t1/2 = 0.693/k second-order: t1/2 = 1/k[A]0 zero-order: t1/2 = [A]0/2k Table 14.1 Atmospheric Lifetimes of Several CFCs Figure 14.8 First-Order Integrated Rate Law Example 14.3 The First-Order Integrated Rate Law: Using Graphical Analysis of Reaction Data Example 14.4 The First-Order Integrated Rate Law: Determining the Concentration of a Reactant at a Given Time Figure 14.9 Second-Order Integrated Rate Law Example 14.5 The Second-Order Integrated Rate Law: Using Graphical Analysis of Reaction Data Figure 14.10 Zero-Order Integrated Rate Law Figure 14.11 Half-Life: Concentration versus Time for a First-Order Reaction Example 14.6 Half-Life Table 14.2 Rate Law Summary Table 14.5 The Effect of Temperature on Reaction Rate Arrhenius equation Potential energy diagram activation energy, Ea transition state (activated complex) plot of ln k vs. 1/T slope = Ea/R intercept = ln A two-point equation exponential factor thermal energy distribution temperature dependence frequency factor collision model collision orientation Figure 14.12 The Activation Energy Barrier unnumbered figures: molecular models of CH3NC, CH3CN Figure 14.13 The Activated Complex unnumbered figure: activation barrier for conversion of CH3NC to CH3CN Figure 14.14 Thermal Energy Distribution Example 14.7 Using an Arrhenius Plot to Determine Kinetic Parameters Example 14.8 Using the Two-Point Form of the Arrhenius Equation Figure 14.15 The Collision Model unnumbered figures: illustrations of molecular collisions Teaching Tips Suggestions and Examples Misconceptions and Pitfalls 14.4 The Integrated Rate Law: The Dependence of Concentration on Time Integrated rate laws are provided for the first-order, second-order, and zero-order cases. In addition, plots of data are shown that yield straight lines from which the rate constant can be determined. Use of the integrated rate laws gives better understanding of the importance of reaction order. It is useful to demonstrate plotting reaction data using Excel or a similar spreadsheet program. The linearity (or lack thereof) of the data discussed in this section can also be determined by comparing the slopes between the first two points of data and the last two points. The integrated rate laws can be derived quickly using elementary calculus. Students with good math backgrounds may appreciate seeing the connection. The half-lives are derived for the three cases. Encourage students to understand these simple derivations rather than memorizing the equations. Conceptual Connection 14.4 Half-Life Part I Conceptual Connection 14.5 Half-Life Part II Conceptual Connection 14.6 Rate Law and Integrated Rate Law Students sometimes confuse the terms rate law and integrated rate law. This section in particular is very challenging to students with weak algebra skills. Since the first-order half-life equation is used so commonly (especially for nuclear processes), students tend to use the equation for all half-life cases. 14.5 The Effect of Temperature on Reaction Rate The activation barrier is shown in a plot of energy vs. reaction progress (sometimes called reaction coordinate). The thermal energy distribution is very similar to the distribution of kinetic energies from the kinetic molecular theory of gases. Conceptual Connection 14.7 Temperature Dependence of Reaction Rate The collision model gives a good conceptual basis for how simple reactions occur. The orientation factor can be modeled very well for simple molecules but not so easily for complex molecules. Conceptual Connection 14.8 Collision Theory Some of the sharper students may question why the pre-exponential factor A is independent of temperature. Strictly speaking, the collision frequency is a function of temperature, but that dependence is smaller than the temperature dependence of the exponential part of the Arrhenius equation. Lecture Outline Terms, Concepts, Relationships, Skills Figures, Tables, and Solved Examples 14.6 Reaction Mechanisms Elementary steps rate law for elementary step Reaction intermediate Molecularity unimolecular bimolecular termolecular Rate-determining step Predicting rate laws from mechanisms Table 14.3 Rate Laws for Elementary Steps unnumbered figure: illustration of traffic analogy to rate-limiting step Figure 14.16 Energy Diagram for a Two-Step Mechanism Example 14.9 Reaction Mechanisms 14.7 Catalysis Catalyst increases rate smaller Ea not consumed by the reaction example: catalytic converter Process and phase homogeneous heterogeneous hydrogenation catalyst Enzymes active site substrate enzyme–substrate binding kinetic model examples sucrase chymotrypsin Figure 14.17 Catalyzed and Uncatalyzed Decomposition of Ozone unnumbered figure: illustration of a catalytic converter Table 14.4 Change in Pollutant Levels unnumbered figure: photo of polar stratospheric cloud Figure 14.18 Homogeneous and Heterogeneous Catalysis Figure 14.19 Ozone Depletion in the Antarctic Spring Figure 14.20 Catalytic Hydrogenation of Ethene unnumbered figure: molecular models of sucrose, glucose, and fructose Figure 14.21 Enzyme–Substrate Binding Figure 14.22 An Enzyme–Catalyzed Reaction Chemistry and Medicine: Enzyme Catalysis and the Role of Chymotrypsin in Digestion Figure 14.23 The Structure of a Protein Figure 14.24 Protein Digestion Figure 14.25 Chymotrypsin, a Digestive Enzyme Figure 14.26 The Action of Chymotrypsin Teaching Tips Suggestions and Examples Misconceptions and Pitfalls 14.6 Reaction Mechanisms This section provides the second illustration of the importance of reaction order. Knowing the order of a reaction can lead to the development of a mechanism for the reaction, leading to a greater understanding of the reaction on a molecular level. The observed rate law provides evidence about the rate-determining elementary step. For most reactions, creating the actual elementary steps requires much work far beyond the scope of the class. The rate law for an overall reaction cannot be written from the balanced reaction, but the rate law for an elementary step is written from the balanced reaction. This in part defines an elementary step. A reaction intermediate is not the same as a transition state. An intermediate can be observed (e.g. carbocations) and exists as a local energy minimum on a potential energy diagram. In a multi-step mechanism, the rate-determining step has the largest Ea and the smallest k. Mechanisms can be disproven but cannot be proven to be true. 14.7 Catalysis Catalysis is a crucial process industrially (e.g. cracking and reforming gasoline molecules), environmentally (e.g. catalytic converters and pollution reduction), and biochemically (enzymes). Provide additional examples of catalysts. Examples with comparative rate constants dramatically illustrate the difference between catalyzed and uncatalyzed reactions. There are several simple demonstrations of catalysis: decomposition of hydrogen peroxide with MnO2; reaction of H2 and O2 catalyzed by Pd (Owens, G. S.; Richmond, T. G. Chemical Educator 1(4): S 1430–4171 (96), 1996, 04045–9. Catalysts decrease the activation barrier by altering the transition state or providing a different mechanism/pathway for the reaction. Catalysts are not consumed by the reaction at hand, but they can be consumed by unwanted side reactions and rendered ineffective after some amount of time. Kinetics (and catalysis) is path-dependent, whereas the thermodynamics and extent of a reaction are path-independent. Catalysis does not affect the amount of energy released or absorbed by a reaction, nor does it affect the product yield. Additional Problem for The First-Order Integrated Rate Law: Determining the Concentration of a Reactant at a Given Time (Example 14.4) Hydrogen peroxide decomposes according to the reaction: H2O2(aq) H2O(l) + 0.5 O2(g) The reaction follows first-order decay with a rate constant of 3.68 h1. If the initial concentration is 30.0 ppm, what is peroxide concentration after 2.00 hours? Sort You have been given the rate of a first-order reaction and the initial concentration of the reactant. You are asked to find the concentration at 2.00 hours. Given k = 3.68 h1 [H2O2]0 = 30.0 ppm Find [H2O2] at t = 2.00 h Strategize Use the first-order integrated rate law to get from the given information to that which you are asked to find. Conceptual Plan k, t, [A]0 [A]t Relationships Used ln [A]t = kt + ln [A]0 Solve Substitute the rate constant, the initial concentration, and the time into the integrated rate law. Solve the integrated rate law for the concentration of [H2O2]t. Solution Check The concentration has decreased a great deal. Only a very small concentration of peroxide remains. Additional Problem for Using the Two-Point Form of the Arrhenius Equation (Example 14.8) The decomposition of ozone is given by the equation: O3(g) O2(g) + O(g) The rate constant at 700 K was measured as 4.85 104 M1s1 and at 800 K was 3.58 105 M1s1. Find the activation energy in kJ/mol. Sort You are given the rate constant of a reaction at two different temperatures. You are asked to find the activation energy. Given T1 = 700 K k1 = 4.85 104 M1s1 T2 = 800 K k2 = 3.58 105 M1s1 Find Ea Strategize Use the two-point form of the Arrhenius equation, which relates activation energy to the given information and R (a constant). Conceptual Plan T1, k1, T2, k2 Ea Relationships Used Solve Substitute the two rate constants and two temperatures into the equation. Solve the equation for Ea, the activation energy, and convert to kJ/mol. Solution Check The units of the answer are correct, and the magnitude makes sense. 192 Copyright © 2017 by Education, Inc. 193 Copyright © 2017 by Education, Inc.

Related Downloads
Explore
Post your homework questions and get free online help from our incredible volunteers
  1103 People Browsing
Your Opinion
Which industry do you think artificial intelligence (AI) will impact the most?
Votes: 308

Previous poll results: Who's your favorite biologist?