Experiments during a recent summer on a number of fireflies (small beetles, Lampyridaes

The reaction between bromate ions and bromide ions in acidic aqueous solution is given by the equation BrO3 2 1aq2 1 5Br2 1aq2 1 6H1 1aq2 h 3Br2 1l2 1 3H2O1l2 Table 12.5 gives the results from four experiments. Using these data, determine the orders for all three reactants, the overall reaction order, and the value of the rate constant.

The decomposition of N2O5 in the gas phase was studied at constant temperature. 2N2O5 1g2 h 4NO2 1g2 1 O2 1g2 The following results were collected: [N2O5] (mol/L) Time (s) 0.1000 0 0.0707 50 0.0500 100 0.0250 200 0.0125 300 0.00625 400 Using these data, verify that the rate law is first order in [N2O5], and calculate the value of the rate constant, where the rate 5 2D[N2O5]yDt.

Using the data given in Example 12.2, calculate [N2O5] at 150 s after the start of the reaction.

A certain first-order reaction has a half-life of 20.0 minutes. a. Calculate the rate constant for this reaction. b. How much time is required for this reaction to be 75% complete?

Butadiene reacts to form its dimer according to the equation 2C4H6 1g2 h C8H12 1g2 The following data were collected for this reaction at a given temperature: [C4H6] (mol/L) Time (1 s) 0.01000 0 0.00625 1000 0.00476 1800 0.00370 2800 0.00313 3600 0.00270 4400 0.00241 5200 0.00208 6200 a. Is this reaction first order or second order? b. What is the value of the rate constant for the reaction? c. What is the half-life for the reaction under the initial conditions of this experiment?

Consider the simple reaction aA n products. You run this reaction and wish to determine its order. What if you made a graph of reaction rate versus time? Could you use this to determine the order? Sketch three plots of rate versus time for the reaction if it is zero, first, or second order. Sketch these plots on the same graph and compare them. Defend your answer.

The balanced equation for the reaction of the gases nitrogen dioxide and fluorine is 2NO2 1g2 1 F2 1g2 h 2NO2F1g2 The experimentally determined rate law is Rate 5 k3NO2 4 3F2 4 A suggested mechanism for this reaction is NO2 1 F2 h k1 NO2F 1 F Slow F 1 NO2 h k2 NO2F Fast Is this an acceptable mechanism? That is, does it satisfy the two requirements?

There are many conditions that need to be met to result in a chemical reaction between molecules. What if all collisions between molecules resulted in a chemical reaction? How would life be different?

The reaction 2N2O5 1g2 h 4NO2 1g2 1 O2 1g2 was studied at several temperatures, and the following values of k were obtained: k (s21) T (8C) 2.0 3 1025 20 7.3 3 1025 30 2.7 3 1024 40 9.1 3 1024 50 2.9 3 1023 60 Calculate the value of Ea for this reaction.

Most modern refrigerators have an internal temperature of 458F. What if refrigerators were set at 558F in the factory? How would this affect our lives?

The gas-phase reaction between methane and diatomic sulfur is given by the equation CH4 1g2 1 2S2 1g2 h CS2 1g2 1 2H2S1g2 At 5508C the rate constant for this reaction is 1.1 L/mol s, and at 6258C the rate constant is 6.4 L/mol s. Using these values, calculate Ea for this reaction.

Define reaction rate. Distinguish between the initial rate, average rate, and instantaneous rate of a chemical reaction. Which of these rates is usually fastest? The initial rate is the rate used by convention. Give a possible explanation as to why

Distinguish between the differential rate law and the integrated rate law. Which of these is often called just the rate law? What is k in a rate law, and what are orders in a rate law? Explain.

One experimental procedure that can be used to determine the rate law of a reaction is the method of initial rates. What data are gathered in the method of initial rates, and how are these data manipulated to determine k and the orders of the species in the rate law? Are the units for k, the rate constant, the same for all rate laws? Explain. If a reaction is first order in A, what happens to the rate if [A] is tripled? If the initial rate for a reaction increases by a factor of 16 when [A] is quadrupled, what is the order of n? If a reaction is third order in A and [A] is doubled, what happens to the initial rate? If a reaction is zero order, what effect does [A] have on the initial rate of a reaction?

The initial rate for a reaction is equal to the slope of the tangent line at t < 0 in a plot of [A] versus time. From calculus, initial rate 5 2d3A4 dt . Therefore, the differential rate law for a reaction is Rate 5 2d3A4 dt 5 k3A4 n . Assuming you have some calculus in your background, derive the zero-, first-, and second-order integrated rate laws using the differential rate law

Consider the zero-, first-, and second-order integrated rate laws. If you have concentration versus time data for some species in a reaction, what plots would you make to prove a reaction is either zero, first, or second order? How would the rate constant, k, be determined from such a plot? What does the y-intercept equal in each plot? When a rate law contains the concentration of two or more species, how can plots be used to determine k and the orders of the species in the rate law?

Derive expressions for the half-life of zero-, first-, and second-order reactions using the integrated rate law for each order. How does each half-life depend on concentration? If the half-life for a reaction is 20. seconds, what would be the second half-life assuming the reaction is either zero, first, or second order?

Define each of the following. a. elementary step b. molecularity c. reaction mechanism d. intermediate e. rate-determining step

What two requirements must be met to call a mechanism plausible? Why say a plausible mechanism instead of the correct mechanism? Is it true that most reactions occur by a one-step mechanism? Explain.

What is the premise underlying the collision model? How is the rate affected by each of the following? a. activation energy b. temperature c. frequency of collisions d. orientation of collisions Sketch a potential energy versus reaction progress plot for an endothermic reaction and for an exothermic reaction. Show DE and Ea in both plots. When concentrations and temperatures are equal, would you expect the rate of the forward reaction to be equal to, greater than, or less than the rate of the reverse reaction if the reaction is exothermic? Endothermic?

Give the Arrhenius equation. Take the natural log of both sides and place this equation in the form of a straight-line equation (y 5 mx 1 b). What data would you need and how would you graph those data to get a linear relationship using the Arrhenius equation? What does the slope of the straight line equal? What does the y-intercept equal? What are the units of R in the Arrhenius equation? Explain how if you know the rate constant value at two different temperatures, you can determine the activation energy for the reaction.

Why does a catalyst increase the rate of a reaction? What is the difference between a homogeneous catalyst and a heterogeneous catalyst? Would a given reaction necessarily have the same rate law for both a catalyzed and an uncatalyzed pathway? Explain.

Define stability from both a kinetic and thermodynamic perspective. Give examples to show the differences in these concepts

Describe at least two experiments you could perform to determine a rate law

Make a graph of [A] versus time for zero-, first-, and secondorder reactions. From these graphs, compare successive halflives

How does temperature affect k, the rate constant? Explain

Consider the following statements: In general, the rate of a chemical reaction increases a bit at first because it takes a while for the reaction to get warmed up. After that, however, the rate of the reaction decreases because its rate is dependent on the concentrations of the reactants, and these are decreasing. Indicate everything that is correct in these statements, and indicate everything that is incorrect. Correct the incorrect statements and explain.

For the reaction A 1 B S C, explain at least two ways in which the rate law could be zero order in chemical A.

A friend of yours states, A balanced equation tells us how chemicals interact. Therefore, we can determine the rate law directly from the balanced equation. What do you tell your friend?

Provide a conceptual rationale for the differences in the halflives of zero-, first-, and second-order reactions

The rate constant (k) depends on which of the following (there may be more than one answer)? a. the concentration of the reactants b. the nature of the reactantsc. the temperature d. the order of the reaction Explain.

Each of the statements given below is false. Explain why. a. The activation energy of a reaction depends on the overall energy change (DE) for the reaction. b. The rate law for a reaction can be deduced from examination of the overall balanced equation for the reaction. c. Most reactions occur by one-step mechanisms.

Define what is meant by unimolecular and bimolecular steps. Why are termolecular steps infrequently seen in chemical reactions?

The plot below shows the number of collisions with a particular energy for two different temperatures.a. Which is greater, T2 or T1? How can you tell? b. What does this plot tell us about the temperature of the rate of a chemical reaction? Explain your answer.

For the reaction O2 1g2 1 2NO1g2 h 2NO2 1g2 the observed rate law is Rate 5 k3NO4 2 3O2 4 Which of the changes listed below would affect the value of the rate constant k? a. increasing the partial pressure of oxygen gas b. changing the temperature c. using an appropriate catalyst

The rate law for a reaction can be determined only from experiment and not from the balanced equation. Two experimental procedures were outlined in Chapter 12. What are these two procedures? Explain how each method is used to determine rate laws.

Table 12.2 illustrates how the average rate of a reaction decreases with time. Why does the average rate decrease with time? How does the instantaneous rate of a reaction depend on time? Why are initial rates used by convention?

The type of rate law for a reaction, either the differential rate law or the integrated rate law, is usually determined by which data is easiest to collect. Explain.

The initial rate of a reaction doubles as the concentration of one of the reactants is quadrupled. What is the order of this reactant? If a reactant has a 21 order, what happens to the initial rate when the concentration of that reactant increases by a factor of two?

Hydrogen reacts explosively with oxygen. However, a mixture of H2 and O2 can exist indefinitely at room temperature. Explain why H2 and O2 do not react under these conditions.

The central idea of the collision model is that molecules must collide in order to react. Give two reasons why not all collisions of reactant molecules result in product formation.

Consider the following energy plots for a chemical reaction when answering the questions below. Reactants Products E2 Reaction progress Energy E1 a. Which plot (purple or blue) is the catalyzed pathway? How do you know? b. What does DE1 represent? c. What does DE2 represent? d. Is the reaction endothermic or exothermic?

Enzymes are kinetically important for many of the complex reactions necessary for plant and animal life to exist. However, only a tiny amount of any particular enzyme is required for these complex reactions to occur. Explain

Would the slope of a ln(k) versus 1yT plot (with temperature in kelvin) for a catalyzed reaction be more or less negative than the slope of the ln(k) versus 1yT plot for the uncatalyzed reaction? Explain. Assume both rate laws are first-order overall.

Consider the reaction 4PH3 1g2 h P4 1g2 1 6H2 1g2 If, in a certain experiment, over a specific time period, 0.0048 mole of PH3 is consumed in a 2.0-L container each second of reaction, what are the rates of production of P4 and H2 in this experiment?

In the Haber process for the production of ammonia, N2 1g2 1 3H2 1g2 h 2NH3 1g2 what is the relationship between the rate of production of ammonia and the rate of consumption of hydrogen?

At 408C, H2O2(aq) will decompose according to the following reaction: 2H2O2 1aq2 h 2H2O1l2 1 O2 1g2 The following data were collected for the concentration of H2O2 at various times. Time (s) [H2O2](mol/L) 0 1.000 2.16 3 104 0.500 4.32 3 104 0.250 a. Calculate the average rate of decomposition of H2O2 between 0 and 2.16 3 104 s. Use this rate to calculate the average rate of production of O2(g) over the same time period. b. What are these rates for the time period 2.16 3 104 s to 4.32 3 104 s?

Consider the general reaction aA 1 bB h cC and the following average rate data over some time period Dt: 2DA Dt 5 0.0080 mol/L # s 2DB Dt 5 0.0120 mol/L # s DC Dt 5 0.0160 mol/L # s Determine a set of possible coefficients to balance this general reaction.

What are the units for each of the following if the concentrations are expressed in moles per liter and the time in seconds? a. rate of a chemical reaction b. rate constant for a zero-order rate law c. rate constant for a first-order rate law d. rate constant for a second-order rate law e. rate constant for a third-order rate law

The rate law for the reaction Cl2 1g2 1 CHCl3 1g2 h HCl1g2 1 CCl4 1g2 is Rate 5 k3Cl2 4 1/2 3CHCl3 4 What are the units for k, assuming time in seconds and concentration in mol/L?

The reaction 2NO1g2 1 Cl2 1g2 h 2NOCl1g2 was studied at 2108C. The following results were obtained where Rate 5 2D3Cl2 4 Dt [NO]0 (mol/L) [Cl2]0 (mol/L) Initial Rate (mol/L min) 0.10 0.10 0.18 0.10 0.20 0.36 0.20 0.20 1.45 a. What is the rate law? b. What is the value of the rate constant?

The reaction 2I2 1aq2 1 S2O8 22 1aq2 h I2 1aq2 1 2SO4 22 1aq2 was studied at 258C. The following results were obtained where Rate 5 2D3S2O8 22 4 Dt [I2]0 (mol/L) [S2O8 22]0 (mol/L) Initial Rate (mol/L s) 0.080 0.040 12.5 3 1026 0.040 0.040 6.25 3 1026 0.080 0.020 6.25 3 1026 0.032 0.040 5.00 3 1026 0.060 0.030 7.00 3 1026 a. Determine the rate law. b. Calculate a value for the rate constant for each experiment and an average value for the rate constant

The decomposition of nitrosyl chloride was studied: 2NOCl1g2 m 2NO1g2 1 Cl2 1g2 The following data were obtained where Rate 5 2D3NOCl4 Dt [NOCl]0 (molecules/cm3) Initial Rate (molecules/cm3 s) 3.0 3 1016 5.98 3 104 2.0 3 1016 2.66 3 104 1.0 3 1016 6.64 3 103 4.0 3 1016 1.06 3 105 a. What is the rate law? b. Calculate the value of the rate constant. c. Calculate the value of the rate constant when concentrations are given in moles per liter.

The following data were obtained for the gas-phase decomposition of dinitrogen pentoxide, 2N2O5 1g2 h 4NO2 1g2 1 O2 1g2 [N2O5]0 (mol/L) Initial Rate (mol/L s) 0.0750 8.90 3 1024 0.190 2.26 3 1023 0.275 3.26 3 1023 0.410 4.85 3 1023 Defining the rate as 2D[N2O5]yDt, write the rate law and calculate the value of the rate constant.

The reaction I 2 1aq2 1 OCl2 1aq2 h IO2 1aq2 1 Cl2 1aq2 was studied, and the following data were obtained: [I2]0 (mol/L) [OCl2]0 (mol/L) Initial Rate (mol/L s) 0.12 0.18 7.91 3 1022 0.060 0.18 3.95 3 1022 0.030 0.090 9.88 3 1023 0.24 0.090 7.91 3 1022 a. What is the rate law? b. Calculate the value of the rate constant. c. Calculate the initial rate for an experiment where both I2 and OCl2 are initially present at 0.15 mol/L.

The reaction 2NO1g2 1 O2 1g2 h 2NO2 1g2 was studied, and the following data were obtained where Rate 5 2D3O2 4 Dt [NO]0 (molecules/cm3) [O2]0 (molecules/cm3) Initial Rate (molecules/cm3 s) 1.00 3 1018 1.00 3 1018 2.00 3 1016 3.00 3 1018 1.00 3 1018 1.80 3 1017 2.50 3 1018 2.50 3 1018 3.13 3 1017 What would be the initial rate for an experiment where [NO]0 5 6.21 3 1018 molecules/cm3 and [O2]0 5 7.36 3 1018 molecules/cm3 ?

The rate of the reaction between hemoglobin (Hb) and carbon monoxide (CO) was studied at 208C. The following data were collected with all concentration units in mmol/L. (A hemoglobin concentration of 2.21 mmol/L is equal to 2.21 3 1026 mol/L.) [Hb]0 (mmol/L) [CO]0 (mmol/L) Initial Rate (mmol/L s) 2.21 1.00 0.619 4.42 1.00 1.24 4.42 3.00 3.71 a. Determine the orders of this reaction with respect to Hb and CO. b. Determine the rate law. c. Calculate the value of the rate constant. d. What would be the initial rate for an experiment with [Hb]0 5 3.36 mmol/L and [CO]0 5 2.40 mmol/L?

The following data were obtained for the reaction 2ClO2 1aq2 1 2OH2 1aq2 h ClO3 2 1aq2 1 ClO2 2 1aq2 1 H2O1l2 where Rate 5 2D3ClO2 4 Dt [ClO2]0 (mol/L) [OH2]0 (mol/L) Initial Rate (mol/L s) 0.0500 0.100 5.75 3 1022 0.100 0.100 2.30 3 1021 0.100 0.0500 1.15 3 1021 a. Determine the rate law and the value of the rate constant. b. What would be the initial rate for an experiment with [ClO2]0 5 0.175 mol/L and [OH2]0 5 0.0844 mol/L?

The decomposition of hydrogen peroxide was studied, and the following data were obtained at a particular temperature: Time (s) [H2O2] (mol/L) 0 1.00 120 6 1 0.91 300 6 1 0.78 600 6 1 0.59 1200 6 1 0.37 1800 6 1 0.22 2400 6 1 0.13 3000 6 1 0.082 3600 6 1 0.050 Assuming that Rate 5 2D3H2O2 4 Dt determine the rate law, the integrated rate law, and the value of the rate constant. Calculate [H2O2] at 4000. s after the start of the reaction.

A certain reaction has the following general form: aA h bB At a particular temperature and [A]0 5 2.00 3 1022 M, concentration versus time data were collected for this reaction, and a plot of ln[A] versus time resulted in a straight line with a slope value of 22.97 3 1022 min21 . a. Determine the rate law, the integrated rate law, and the value of the rate constant for this reaction. b. Calculate the half-life for this reaction. c. How much time is required for the concentration of A to decrease to 2.50 3 1023 M?

The rate of the reaction NO2 1g2 1 CO1g2 h NO1g2 1 CO2 1g2 depends only on the concentration of nitrogen dioxide below 2258C. At a temperature below 2258C, the following data were collected: Time (s) [NO2] (mol/L) 0 0.500 1.20 3 103 0.444 3.00 3 103 0.381 4.50 3 103 0.340 9.00 3 103 0.250 1.80 3 104 0.174 Determine the rate law, the integrated rate law, and the value of the rate constant. Calculate [NO2] at 2.70 3 104 s after the start of the reaction

A certain reaction has the following general form: aA h bB At a particular temperature and [A]0 5 2.80 3 1023 M, concentration versus time data were collected for this reaction, and a plot of 1y[A] versus time resulted in a straight line with a slope value of 13.60 3 1022 L/mol s. a. Determine the rate law, the integrated rate law, and the value of the rate constant for this reaction. b. Calculate the half-life for this reaction. c. How much time is required for the concentration of A to decrease to 7.00 3 1024 M

The decomposition of ethanol (C2H5OH) on an alumina (Al2O3) surface C2H5OH1g2 h C2H4 1g2 1 H2O1g2 was studied at 600 K. Concentration versus time data were collected for this reaction, and a plot of [A] versus time resulted in a straight line with a slope of 24.00 3 1025 mol/L s. a. Determine the rate law, the integrated rate law, and the value of the rate constant for this reaction. b. If the initial concentration of C2H5OH was 1.25 3 1022 M, calculate the half-life for this reaction. c. How much time is required for all the 1.25 3 1022 M C2H5OH to decompose?

At 500 K in the presence of a copper surface, ethanol decomposes according to the equation C2H5OH1g2 h CH3CHO1g2 1 H2 1g2 The pressure of C2H5OH was measured as a function of time and the following data were obtained: Time (s) PC2H5OH (torr) 0 250. 100. 237 200. 224 300. 211 400. 198 500. 185 Since the pressure of a gas is directly proportional to the concentration of gas, we can express the rate law for a gaseous reaction in terms of partial pressures. Using the above data, deduce the rate law, the integrated rate law, and the value of the rate constant, all in terms of pressure units in atm and time in seconds. Predict the pressure of C2H5OH after 900. s from the start of the reaction. (Hint: To determine the order of the reaction with respect to C2H5OH, compare how the pressure of C2H5OH decreases with each time listing.)

The dimerization of butadiene 2C4H6 1g2 h C8H12 1g2 was studied at 500. K, and the following data were obtained: Time (s) [C4H6] (mol/L) 195 1.6 3 1022 604 1.5 3 1022 1246 1.3 3 1022 2180 1.1 3 1022 6210 0.68 3 1022 Assuming that Rate 5 2D3C4H6 4 Dt determine the form of the rate law, the integrated rate law, and the value of the rate constant for this reaction.

The rate of the reaction O1g2 1 NO2 1g2 h NO1g2 1 O2 1g2 was studied at a certain temperature. a. In one experiment, NO2 was in large excess, at a concentration of 1.0 3 1013 molecules/cm3 with the following data collected: Time (s) [O] (atoms/cm3) 0 5.0 3 109 1.0 3 1022 1.9 3 109 2.0 3 1022 6.8 3 108 3.0 3 1022 2.5 3 108 What is the order of the reaction with respect to oxygen atoms?b. The reaction is known to be first order with respect to NO2. Determine the overall rate law and the value of the rate constant

Experimental data for the reaction A h 2B 1 C have been plotted in the following three different ways (with concentration units in mol/L): 0 0.05 0.04 0.03 0.02 0.01 0 2 Time (s) [A] 4 6 0 2 4 6 3.0 3.5 4.0 4.5 5.0 Time (s) ln[A] 0 100 80 60 40 20 0 2 Time (s) 1/[A] 4 6 What is the order of the reaction with respect to A, and what is the initial concentration of A?

Consider the data plotted in Exercise 45 when answering the following questions. a. What is the concentration of A after 9 s? b. What are the first three half-lives for this experiment?

The reaction A h B 1 C is known to be zero order in A and to have a rate constant of 5.0 3 1022 mol/L s at 258C. An experiment was run at 258C where [A]0 5 1.0 3 1023 M. a. Write the integrated rate law for this reaction. b. Calculate the half-life for the reaction. c. Calculate the concentration of B after 5.0 3 1023 s has elapsed assuming [B]0 5 0

The decomposition of hydrogen iodide on finely divided gold at 1508C is zero order with respect to HI. The rate defined below is constant at 1.20 3 1024 mol/L s. 2HI1g2 hAu H2 1g2 1 I2 1g2 Rate 5 2D3HI4 Dt 5 k 5 1.20 3 1024 mol/L # s a. If the initial HI concentration was 0.250 mol/L, calculate the concentration of HI at 25 minutes after the start of the reaction. b. How long will it take for all of the 0.250 M HI to decompose?

A certain first-order reaction is 45.0% complete in 65 s. What are the values of the rate constant and the half-life for this process?

A first-order reaction is 75.0% complete in 320. s. a. What are the first and second half-lives for this reaction? b. How long does it take for 90.0% completion?

The rate law for the decomposition of phosphine (PH3) is Rate 5 2D3PH3 4 Dt 5 k3PH3 4 It takes 120. s for 1.00 M PH3 to decrease to 0.250 M. How much time is required for 2.00 M PH3 to decrease to a concentration of 0.350 M?

DDT (molar mass 5 354.49 g/mol) was a widely used insecticide that was banned from use in the United States in 1973. This ban was brought about due to the persistence of DDT in many different ecosystems, leading to high accumulations of the substance in many birds of prey. The insecticide was shown to cause a thinning of egg shells, pushing many birds toward extinction. If a 20-L drum of DDT was spilled into a pond, resulting in a DDT concentration of 8.75 3 1025 M, how long would it take for the levels of DDT to reach a concentration of 1.41 3 1027 M (a level that is generally assumed safe in mammals)? Assume the decomposition of DDT is a firstorder process with a half-life of 56.0 days.

Consider the following initial rate data for the decomposition of compound AB to give A and B: [AB]0 (mol/L) Initial Rate (mol/L s) 0.200 3.20 3 1023 0.400 1.28 3 1022 0.600 2.88 3 1022 Determine the half-life for the decomposition reaction initially having 1.00 M AB present.

The rate law for the reaction 2NOBr1g2 h 2NO1g2 1 Br2 1g2 at some temperature is Rate 5 2D3NOBr4 Dt 5 k3NOBr4 2 a. If the half-life for this reaction is 2.00 s when [NOBr]0 5 0.900 M, calculate the value of k for this reaction. b. How much time is required for the concentration of NOBr to decrease to 0.100 M?

For the reaction A n products, successive half-lives are observed to be 10.0, 20.0, and 40.0 min for an experiment in which [A]0 5 0.10 M. Calculate the concentration of A at the following times. a. 80.0 min b. 30.0 min

Theophylline is a pharmaceutical drug that is sometimes used to help with lung function. You observe a case where the initial lab results indicate that the concentration of theophylline in a patients body decreased from 2.0 3 1023 M to 1.0 3 1023 M in 24 hours. In another 12 hours the drug concentration was found to be 5.0 3 1024 M. What is the value of the rate constant for the metabolism of this drug in the body?

You and a coworker have developed a molecule that has shown potential as cobra antivenin (AV). This antivenin works by binding to the venom (V), thereby rendering it nontoxic. This reaction can be described by the rate law Rate 5 k3AV4 1 3V4 1 You have been given the following data from your coworker: 3V4 0 5 0.20 M 3AV4 0 5 1.0 3 1024 M A plot of ln[AV] versus t (s) gives a straight line with a slope of 20.32 s21 . What is the value of the rate constant (k) for this reaction?

Consider the hypothetical reaction A 1 B 1 2C h 2D 1 3E where the rate law is Rate 5 2D3A4 Dt 5 k3A4 3B4 2 An experiment is carried out where [A]0 5 1.0 3 1022 M, [B]0 5 3.0 M, and [C]0 5 2.0 M. The reaction is started, and after 8.0 seconds, the concentration of A is 3.8 3 1023 M. a. Calculate the value of k for this reaction. b. Calculate the half-life for this experiment. c. Calculate the concentration of A after 13.0 seconds. d. Calculate the concentration of C after 13.0 seconds.

Write the rate laws for the following elementary reactions. a. CH3NC1g2 S CH3CN1g2 b. O3 1g2 1 NO1g2 S O2 1g2 1 NO2 1g2 c. O3 1g2 S O2 1g2 1 O1g2 d. O3 1g2 1 O1g2 S 2O2 1g2

A possible mechanism for the decomposition of hydrogen peroxide is H2O2 h 2OH H2O2 1 OH h H2O 1 HO2 HO2 1 OH h H2O 1 O2 Using your results from Exercise 37, specify which step is the rate-determining step. What is the overall balanced equation for the reaction?

A proposed mechanism for a reaction is C4H9Br h C4H9 1 1 Br2 Slow C4H9 1 1 H2O h C4H9OH2 1 Fast C4H9OH2 1 1 H2O h C4H9OH 1 H3O1 Fast Write the rate law expected for this mechanism. What is the overall balanced equation for the reaction? What are the intermediates in the proposed mechanism?

The mechanism for the gas-phase reaction of nitrogen dioxide with carbon monoxide to form nitric oxide and carbon dioxide is thought to be NO2 1 NO2 h NO3 1 NO Slow NO3 1 CO h NO2 1 CO2 Fast Write the rate law expected for this mechanism. What is the overall balanced equation for the reaction?

For the following reaction profile, indicate a. the positions of reactants and products. b. the activation energy. c. DE for the reaction.

Draw a rough sketch of the energy profile for each of the following cases: a. DE 5 110 kJ/mol, Ea 5 25 kJ/mol b. DE 5 210 kJ/mol, Ea 5 50 kJ/mol c. DE 5 250 kJ/mol, Ea 5 50 kJ/mol

The activation energy for the reaction NO2 1g2 1 CO1g2 h NO1g2 1 CO2 1g2 is 125 kJ/mol, and DE for the reaction is 2216 kJ/mol. What is the activation energy for the reverse reaction 3NO1g2 1 CO2 1g2 h NO2 1g2 1 CO1g2 4 ?

The activation energy for some reaction X2 1g2 1 Y2 1g2 h 2XY1g2 is 167 kJ/mol, and DE for the reaction is 128 kJ/mol. What is the activation energy for the decomposition of XY?

The rate constant for the gas-phase decomposition of N2O5, N2O5 h 2NO2 1 1 2O2 has the following temperature dependence: T (K) k (s21) 338 4.9 3 1023 318 5.0 3 1024 298 3.5 3 1025 Make the appropriate graph using these data, and determine the activation energy for this reaction.

The reaction 1CH32 3CBr 1 OH2 h 1CH32 3COH 1 Br2 in a certain solvent is first order with respect to (CH3)3CBr and zero order with respect to OH2. In several experiments, the rate constant k was determined at different temperatures. A plot of ln(k) versus 1yT was constructed resulting in a straight line with a slope value of 21.10 3 104 K and y-intercept of 33.5. Assume k has units of s21 . a. Determine the activation energy for this reaction. b. Determine the value of the frequency factor A. c. Calculate the value of k at 258C

The activation energy for the decomposition of HI(g) to H2(g) and I2(g) is 186 kJ/mol. The rate constant at 555 K is 3.52 3 1027 L/mol s. What is the rate constant at 645 K?

A first-order reaction has rate constants of 4.6 3 1022 s21 and 8.1 3 1022 s21 at 08C and 20.8C, respectively. What is the value of the activation energy?

A certain reaction has an activation energy of 54.0 kJ/mol. As the temperature is increased from 228C to a higher temperature, the rate constant increases by a factor of 7.00. Calculate the higher temperature

Chemists commonly use a rule of thumb that an increase of 10 K in temperature doubles the rate of a reaction. What must the activation energy be for this statement to be true for a temperature increase from 25 to 358C?

Which of the following reactions would you expect to proceed at a faster rate at room temperature? Why? (Hint: Think about which reaction would have the lower activation energy.) 2Ce41 1aq2 1 Hg2 21 1aq2 h 2Ce31 1aq2 1 2Hg21 1aq2 H3O1 1aq2 1 OH2 1aq2 h 2H2O1l2

One reason suggested for the instability of long chains of silicon atoms is that the decomposition involves the transition state shown below: H H H SiH4 + SiH2 Si H H Si H H Si H H H H H Si The activation energy for such a process is 210 kJ/mol, which is less than either the SiOSi or the SiOH bond energy. Why would a similar mechanism not be expected to play a very important role in the decomposition of long chains of carbon atoms as seen in organic compounds?

One mechanism for the destruction of ozone in the upper atmosphere is O3 1g2 1 NO1g2 h NO2 1g2 1 O2 1g2 Slow NO2 1g2 1 O1g2 h NO1g2 1 O2 1g2 Fast Overall reaction O3 1g2 1 O1g2 h 2O2 1g2 a. Which species is a catalyst? b. Which species is an intermediate? c. Ea for the uncatalyzed reaction O3 1g2 1 O1g2 h 2O2 1g2 is 14.0 kJ. Ea for the same reaction when catalyzed is 11.9 kJ. What is the ratio of the rate constant for the catalyzed reaction to that for the uncatalyzed reaction at 258C? Assume that the frequency factor A is the same for each reaction.

One of the concerns about the use of Freons is that they will migrate to the upper atmosphere, where chlorine atoms can be generated by the following reaction: CCl2F2 1g2 hhv CF2Cl1g2 1 Cl1g2 Freon-12 Chlorine atoms can act as a catalyst for the destruction of ozone. The activation energy for the reaction Cl1g2 1 O3 1g2 h ClO1g2 1 O2 1g2 is 2.1 kJ/mol. Which is the more effective catalyst for the destruction of ozone, Cl or NO? (See Exercise 75.)

Assuming that the mechanism for the hydrogenation of C2H4 given in Section 12.7 is correct, would you predict that the product of the reaction of C2H4 with D2 would be CH2DOCH2D or CHD2OCH3? How could the reaction of C2H4 with D2 be used to confirm the mechanism for the hydrogenation of C2H4 given in Section 12.7?

The decomposition of NH3 to N2 and H2 was studied on two surfaces: Surface Ea (kJ/mol) W 163 Os 197 Without a catalyst, the activation energy is 335 kJ/mol. a. Which surface is the better heterogeneous catalyst for the decomposition of NH3? Why? b. How many times faster is the reaction at 298 K on the W surface compared with the reaction with no catalyst present? Assume that the frequency factor A is the same for each reaction. c. The decomposition reaction on the two surfaces obeys a rate law of the form Rate 5 k 3NH3 4 3H2 4 How can you explain the inverse dependence of the rate on the H2 concentration?

The decomposition of many substances on the surface of a heterogeneous catalyst shows the following behavior: Rate Concentration of reactant How do you account for the rate law changing from first order to zero order in the concentration of reactant?

For enzyme-catalyzed reactions that follow the mechanism E 1 S m E # S E # S m E 1 P a graph of the rate as a function of [S], the concentration of the substrate, has the following appearance: Rate [S] Note that at higher substrate concentrations the rate no longer changes with [S]. Suggest a reason for this

A popular chemical demonstration is the magic genie procedure, in which hydrogen peroxide decomposes to water and oxygen gas with the aid of a catalyst. The activation energy of this (uncatalyzed) reaction is 70.0 kJ/mol. When the catalyst is added, the activation energy (at 20.8C) is 42.0 kJ/mol. Theoretically, to what temperature (8C) would one have to heat the hydrogen peroxide solution so that the rate of the uncatalyzed reaction is equal to the rate of the catalyzed reaction at 20.8C? Assume the frequency factor A is constant, and assume the initial concentrations are the same

The activation energy for a reaction is changed from 184 kJ/mol to 59.0 kJ/mol at 600. K by the introduction of a catalyst. If the uncatalyzed reaction takes about 2400 years to occur, about how long will the catalyzed reaction take? Assume the frequency factor A is constant, and assume the initial concentrations are the same.

Consider the following representation of the reaction 2NO2(g) n 2NO(g) 1 O2(g). (a) time = 0 minutes (b) time = 10 minutes (c) time = ? minutes Time Determine the time for the final representation above if the reaction is a. first order b. second order c. zero order

The reaction H2SeO3 1aq2 1 6I2 1aq2 1 4H1 1aq2 h Se1s2 1 2I3 2 1aq2 1 3H2O1l2 was studied at 08C, and the following data were obtained: [H2SeO3]0 (mol/L) [H1]0 (mol/L) [I2]0 (mol/L) Initial Rate (mol/L s) 1.0 3 1024 2.0 3 1022 2.0 3 1022 1.66 3 1027 2.0 3 1024 2.0 3 1022 2.0 3 1022 3.33 3 1027 3.0 3 1024 2.0 3 1022 2.0 3 1022 4.99 3 1027 1.0 3 1024 4.0 3 1022 2.0 3 1022 6.66 3 1027 1.0 3 1024 1.0 3 1022 2.0 3 1022 0.42 3 1027 1.0 3 1024 2.0 3 1022 4.0 3 1022 13.2 3 1027 1.0 3 1024 1.0 3 1022 4.0 3 1022 3.36 3 1027 These relationships hold only if there is a very small amount of I3 2 present. What is the rate law and the value of the rate constant? aAssume that rate 5 2D3H2SeO3 4 Dt

Consider two reaction vessels, one containing A and the other containing B, with equal concentrations at t 5 0. If both substances decompose by first-order kinetics, where kA 5 4.50 3 1024 s21 kB 5 3.70 3 1023 s21 how much time must pass to reach a condition such that [A] 5 4.00[B]?

Sulfuryl chloride (SO2Cl2) decomposes to sulfur dioxide (SO2) and chlorine (Cl2) by reaction in the gas phase. The following pressure data were obtained when a sample containing 5.00 3 1022 mol sulfuryl chloride was heated to 600. K in a 5.00 3 1021 -L container. Time (hours): 0.00 1.00 2.00 4.00 8.00 16.00 PSO2Cl2 (atm): 4.93 4.26 3.52 2.53 1.30 0.34 Defining the rate as 2D3SO2Cl2 4 Dt , a. determine the value of the rate constant for the decomposition of sulfuryl chloride at 600. K. b. what is the half-life of the reaction? c. what fraction of the sulfuryl chloride remains after 20.0 h?

For the reaction 2N2O5 1g2 h 4NO2 1g2 1 O2 1g2 the following data were collected, where Rate 5 2D3N2O5 4 Dt Time (s) T 5 338 K [N2O5] T 5 318 K [N2O5] 0 1.00 3 1021 M 1.00 3 1021 M 100. 6.14 3 1022 M 9.54 3 1022 M 300. 2.33 3 1022 M 8.63 3 1022 M 600. 5.41 3 1023 M 7.43 3 1022 M 900. 1.26 3 1023 M 6.39 3 1022 M Calculate Ea for this reaction.

Experimental values for the temperature dependence of the rate constant for the gas-phase reaction NO 1 O3 h NO2 1 O2 are as follows: T (K) k (L/mol s) 195 1.08 3 109 230. 2.95 3 109 260. 5.42 3 109 298 12.0 3 109 369 35.5 3 109 Make the appropriate graph using these data, and determine the activation energy for this reaction.

Cobra venom helps the snake secure food by binding to acetylcholine receptors on the diaphragm of a bite victim, leading to the loss of function of the diaphragm muscle tissue and eventually death. In order to develop more potent antivenins, scientists have studied what happens to the toxin once it has bound the acetylcholine receptors. They have found that the toxin is released from the receptor in a process that can be described by the rate law Rate 5 k3acetylcholine receptortoxin complex4 If the activation energy of this reaction at 37.08C is 26.2 kJ/ mol and A 5 0.850 s21 , what is the rate of reaction if you have a 0.200-M solution of receptortoxin complex at 37.08C?

Iodomethane (CH3I) is a commonly used reagent in organic chemistry. When used properly, this reagent allows chemists to introduce methyl groups in many different useful applications. The chemical does pose a risk as a carcinogen, possibly owing to iodomethanes ability to react with portions of the DNA strand (if they were to come in contact). Consider the following hypothetical initial rates data: [DNA]0 (mmol/L) [CH3I]0 (mmol/L) Initial Rate (mmol/L s) 0.100 0.100 3.20 3 1024 0.100 0.200 6.40 3 1024 0.200 0.200 1.28 3 1023 Which of the following could be a possible mechanism to explain the initial rate data? Mechanism I DNA 1 CH3I h DNAiCH3 1 1 I 2 Mechanism II CH3I h CH3 1 1 I 2 Slow DNA 1 CH3 1 h DNAiCH3 1 Fast

Experiments during a recent summer on a number of fireflies (small beetles, Lampyridaes photinus) showed that the average interval between flashes of individual insects was 16.3 s at 21.08C and 13.0 s at 27.88C. a. What is the apparent activation energy of the reaction that controls the flashing? b. What would be the average interval between flashes of an individual firefly at 30.08C? c. Compare the observed intervals and the one you calculated in part b to the rule of thumb that the Celsius temperature is 54 minus twice the interval between flashes

The activation energy of a certain uncatalyzed biochemical reaction is 50.0 kJ/mol. In the presence of a catalyst at 378C, the rate constant for the reaction increases by a factor of 2.50 3 103 as compared with the uncatalyzed reaction. Assuming the frequency factor A is the same for both the catalyzed and uncatalyzed reactions, calculate the activation energy for the catalyzed reaction.

Consider the reaction 3A 1 B 1 C h D 1 E where the rate law is defined as 2D3A4 Dt 5 k3A4 2 3B4 3C4 An experiment is carried out where [B]0 5 [C]0 5 1.00 M and [A]0 5 1.00 3 1024 M. a. If after 3.00 min, [A] 5 3.26 3 1025 M, calculate the value of k.b. Calculate the half-life for this experiment. c. Calculate the concentration of B and the concentration of A after 10.0 min.

The thiosulfate ion (S2O3 22) is oxidized by iodine as follows: 2S2O3 22 1aq2 1 I2 1aq2 h S4O6 22 1aq2 1 2I2 1aq2 In a certain experiment, 7.05 3 1023 mol/L of S2O3 22 is consumed in the first 11.0 seconds of the reaction. Calculate the rate of consumption of S2O3 22. Calculate the rate of production of iodide ion

The reaction A1aq2 1 B1aq2 h products1aq2 was studied, and the following data were obtained: [A]0 (mol/L) [B]0 (mol/L) Initial Rate (mol/L s) 0.12 0.18 3.46 3 1022 0.060 0.12 1.15 3 1022 0.030 0.090 4.32 3 1023 0.24 0.090 3.46 3 1022 What is the order of the reaction with respect to A? What is the order of the reaction with respect to B? What is the value of the rate constant for the reaction?

A certain substance, initially present at 0.0800 M, decomposes by zero-order kinetics with a rate constant of 2.50 3 1022 mol/L s. Calculate the time (in seconds) required for the system to reach a concentration of 0.0210 M.

A reaction of the form aA h Products gives a plot of ln[A] versus time (in seconds), which is a straight line with a slope of 7.35 3 1023 . Assuming [A]0 5 0.0100 M, calculate the time (in seconds) required for the reaction to reach 22.9% completion

A certain reaction has the form aA h Products At a particular temperature, concentration versus time data were collected. A plot of 1y[A] versus time (in seconds) gave a straight line with a slope of 6.90 3 1022 . What is the differential rate law for this reaction? What is the integrated rate law for this reaction? What is the value of the rate constant for this reaction? If [A]0 for this reaction is 0.100 M, what is the first half-life (in seconds)? If the original concentration (at t 5 0) is 0.100 M, what is the second half-life (in seconds)?

Which of the following statement(s) is(are) true? a. The half-life for a zero-order reaction increases as the reaction proceeds. b. A catalyst does not change the value of the rate constant. c. The half-life for a reaction, aA 8n products, that is first order in A increases with increasing [A]0. d. The half-life for a second-order reaction increases as the reaction proceeds.

Consider the hypothetical reaction A2 1g2 1 B2 1g2 88n 2AB1g2 , where the rate law is: 2D3A2 4 Dt 5 k3A2 4 3B2 4 The value of the rate constant at 3028C is 2.45 3 1024 L/mol s, and at 5088C the rate constant is 0.891 L/mol s. What is the activation energy for this reaction? What is the value of the rate constant for this reaction at 3758C?

Experiments have shown that the average frequency of chirping by a snowy tree cricket (Oecanthus fultoni) depends on temperature as shown in the table. Chirping Rate (per min) Temperature (8C) 178 25.0 126 20.3 100. 17.3 What is the apparent activation energy of the process that controls the chirping? What is the rate of chirping expected at a temperature of 7.58C?

Consider a reaction of the type aA h products, in which the rate law is found to be rate 5 k[A]3 (termolecular reactions are improbable but possible). If the first half-life of the reaction is found to be 40. s, what is the time for the second half-life? Hint: Using your calculus knowledge, derive the integrated rate law from the differential rate law for a termolecular reaction: Rate 5 2d3A4 dt 5 k3A4 3

A study was made of the effect of the hydroxide concentration on the rate of the reaction I 2 1aq2 1 OCl2 1aq2 h IO2 1aq2 1 Cl2 1aq2 The following data were obtained: [I2]0 (mol/L) [OCl2]0 (mol/L) [OH2]0 (mol/L) Initial Rate (mol/L s) 0.0013 0.012 0.10 9.4 3 1023 0.0026 0.012 0.10 18.7 3 1023 0.0013 0.0060 0.10 4.7 3 1023 0.0013 0.018 0.10 14.0 3 1023 0.0013 0.012 0.050 18.7 3 1023 0.0013 0.012 0.20 4.7 3 1023 0.0013 0.018 0.20 7.0 3 1023 Determine the rate law and the value of the rate constant for this reaction

Two isomers (A and B) of a given compound dimerize as follows: 2A hk1 A2 2B hk2 B2 Both processes are known to be second order in reactant, and k1 is known to be 0.250 L/mol s at 258C. In a particular experiment A and B were placed in separate containers at 258C, where [A]0 5 1.00 3 1022 M and [B]0 5 2.50 3 1022 M. It was found that after each reaction had progressed for 3.00 min, [A] 5 3.00[B]. In this case the rate laws are defined as Rate 5 2D3A4 Dt 5 k1 3A4 2 Rate 5 2D3B4 Dt 5 k2 3B4 2 a. Calculate the concentration of A2 after 3.00 min. b. Calculate the value of k2. c. Calculate the half-life for the experiment involving A.

The reaction NO1g2 1 O3 1g2 h NO2 1g2 1 O2 1g2 was studied by performing two experiments. In the first experiment the rate of disappearance of NO was followed in the presence of a large excess of O3. The results were as follows ([O3] remains effectively constant at 1.0 3 1014 molecules/cm3 ): Time (ms) [NO] (molecules/cm3) 0 6.0 3 108 100 6 1 5.0 3 108 500 6 1 2.4 3 108 700 6 1 1.7 3 108 1000 6 1 9.9 3 107 In the second experiment [NO] was held constant at 2.0 3 1014 molecules/cm3 . The data for the disappearance of O3 are as follows: Time (ms) [O3] (molecules/cm3) 0 1.0 3 1010 50 6 1 8.4 3 109 100 6 1 7.0 3 109 200 6 1 4.9 3 109 300 6 1 3.4 3 109 a. What is the order with respect to each reactant? b. What is the overall rate law? c. What is the value of the rate constant from each set of experiments? Rate 5 kr 3NO4 x Rate 5 ks3O3 4 y d. What is the value of the rate constant for the overall rate law? Rate 5 k3NO4 x 3O3 4 y

Most reactions occur by a series of steps. The energy profile for a certain reaction that proceeds by a two-step mechanism is Reaction coordinate E On the energy profile, indicate a. the positions of reactants and products. b. the activation energy for the overall reaction. c. DE for the reaction. d. Which point on the plot represents the energy of the intermediate in the two-step reaction? e. Which step in the mechanism for this reaction is rate determining, the first or the second step? Explain

You are studying the kinetics of the reaction H2 1g2 1 F2 1g2 n 2HF1g2 and you wish to determine a mechanism for the reaction. You run the reaction twice by keeping one reactant at a much higher pressure than the other reactant (this lower-pressure reactant begins at 1.000 atm). Unfortunately, you neglect to record which reactant was at the higher pressure, and you forget which it was later. Your data for the first experiment are: Pressure of HF (atm) Time (min) 0 0 0.300 30.0 0.600 65.8 0.900 110.4 1.200 169.1 1.500 255.9 When you ran the second experiment (in which the higherpressure reactant was run at a much higher pressure), you determine the values of the apparent rate constants to be the same. It also turns out that you find data taken from another person in the lab. This individual found that the reaction proceeds 40.0 times faster at 558C than at 358C. You also know, from the energy-level diagram, that there are three steps to the mechanism, and the first step has the highest activation energy. You look up the bond energies of the species involved and they are (in kJ/mol): HOH (432), FOF (154), and HOF (565). a. Sketch an energy-level diagram (qualitative) that is consistent with the one described previously. Hint: See Exercise 106. b. Develop a reasonable mechanism for the reaction. c. Which reactant was limiting in the experiments?

The decomposition of NO2(g) occurs by the following bimolecular elementary reaction: 2NO2 1g2 h 2NO1g2 1 O2 1g2 The rate constant at 273 K is 2.3 3 10212 L/mol s, and the activation energy is 111 kJ/mol. How long will it take for the concentration of NO2(g) to decrease from an initial partial pressure of 2.5 atm to 1.5 atm at 500. K? Assume ideal gas behavior

The following data were collected in two studies of the reaction 2A 1 B h C 1 D Time (s) Experiment 1 [A] (mol/L) 3 1022 Experiment 2 [A] (mol/L) 3 1022 0 10.0 10.0 20. 6.67 5.00 40. 5.00 3.33 60. 4.00 2.50 80. 3.33 2.00 100. 2.86 1.67 120. 2.50 1.43 In Experiment 1, [B]0 5 5.0 M. In Experiment 2, [B]0 5 10.0 M. Rate 5 2D3A4 Dt a. Why is [B] much greater than [A]? b. Give the rate law and value for k for this reaction.

Consider the following hypothetical data collected in two studies of the reaction 2A 1 2B h C 1 2D Time (s) Experiment 1 [A] (mol/L) Experiment 2 [A] (mol/L) 0 1.0 3 1022 1.0 3 1022 10. 8.4 3 1023 5.0 3 1023 20. 7.1 3 1023 2.5 3 1023 30. ? 1.3 3 1023 40. 5.0 3 1023 6.3 3 1024 In Experiment 1, [B]0 5 10.0 M. In Experiment 2, [B]0 5 20.0 M. Rate 5 2D3A4 Dt a. Use the concentration versus time data to determine the rate law for the reaction. b. Solve for the value of the rate constant (k) for the reaction. Include units. c. Calculate the concentration of A in Experiment 1 at t 5 30. s.

Consider the hypothetical reaction A 1 B 1 2C h 2D 1 3E In a study of this reaction three experiments were run at the same temperature. The rate is defined as 2D[B]yDt. Experiment 1: 3A4 0 5 2.0 M 3B4 0 5 1.0 3 1023 M 3C4 0 5 1.0 M [B] (mol/L) Time (s) 2.7 3 1024 1.0 3 105 1.6 3 1024 2.0 3 105 1.1 3 1024 3.0 3 105 8.5 3 1025 4.0 3 105 6.9 3 1025 5.0 3 105 5.8 3 1025 6.0 3 105 Experiment 2: 3A4 0 5 1.0 3 1022 M 3B4 0 5 3.0 M 3C4 0 5 1.0 M [A] (mol/L) Time (s) 8.9 3 1023 1.0 7.1 3 1023 3.0 5.5 3 1023 5.0 3.8 3 1023 8.0 2.9 3 1023 10.0 2.0 3 1023 13.0 Experiment 3: 3A4 0 5 10.0 M 3B4 0 5 5.0 M 3C4 0 5 5.0 3 1021 M [C] (mol/L) Time (s) 0.43 1.0 3 1022 0.36 2.0 3 1022 0.29 3.0 3 1022 0.22 4.0 3 1022 0.15 5.0 3 1022 0.08 6.0 3 1022 Write the rate law for this reaction, and calculate the value of the rate constant.

Hydrogen peroxide and the iodide ion react in acidic solution as follows: H2O2 1aq2 1 3I2 1aq2 1 2H1 1aq2 h I3 2 1aq2 1 2H2O1l2 The kinetics of this reaction were studied by following the decay of the concentration of H2O2 and constructing plots of ln[H2O2] versus time. All the plots were linear and all solutions had [H2O2]0 5 8.0 3 1024 mol/L. The slopes of these straight lines depended on the initial concentrations of I2 and H1. The results follow: [I2]0 (mol/L) [H1]0 (mol/L) Slope (min21) 0.1000 0.0400 20.120 0.3000 0.0400 20.360 0.4000 0.0400 20.480 0.0750 0.0200 20.0760 0.0750 0.0800 20.118 0.0750 0.1600 20.174 The rate law for this reaction has the form Rate 5 2D3H2O2 4 Dt 5 1k1 1 k2 3H1 42 3I 2 4 m 3H2O2 4 n a. Specify the order of this reaction with respect to [H2O2] and [I2]. b. Calculate the values of the rate constants, k1 and k2. c. What reason could there be for the two-term dependence of the rate on [H1]?

Sulfuryl chloride undergoes first-order decomposition at 320.8C with a half-life of 8.75 h. SO2Cl2 1g2 h SO2 1g2 1 Cl2 1g2 What is the value of the rate constant, k, in s21 ? If the initial pressure of SO2Cl2 is 791 torr and the decomposition occurs in a 1.25-L container, how many molecules of SO2Cl2 remain after 12.5 h?

Upon dissolving InCl(s) in HCl, In1(aq) undergoes a disproportionation reaction according to the following unbalanced equation: In1 1aq2 h In1s2 1 In31 1aq2 This disproportionation follows first-order kinetics with a half-life of 667 s. What is the concentration of In1(aq) after 1.25 h if the initial solution of In1(aq) was prepared by dissolving 2.38 g InCl(s) in dilute HCl to make 5.00 3 102 mL of solution? What mass of In(s) is formed after 1.25 h?

The decomposition of iodoethane in the gas phase proceeds according to the following equation: C2H5I1g2 h C2H4 1g2 1 HI1g2 At 660. K, k 5 7.2 3 1024 s21 ; at 720. K, k 5 1.7 3 1022 s21 . What is the value of the rate constant for this first-order decomposition at 3258C? If the initial pressure of iodoethane is 894 torr at 2458C, what is the pressure of iodoethane after three half-lives?

Consider the following reaction: CH3X 1 Y h CH3Y 1 X At 258C, the following two experiments were run, yielding the following data: Experiment 1: [Y]0 5 3.0 M [CH3X] (mol/L) Time (h) 7.08 3 1023 1.0 4.52 3 1023 1.5 2.23 3 1023 2.3 4.76 3 1024 4.0 8.44 3 1025 5.7 2.75 3 1025 7.0 Experiment 2: [Y]0 5 4.5 M [CH3X] (mol/L) Time (h) 4.50 3 1023 0 1.70 3 1023 1.0 4.19 3 1024 2.5 1.11 3 1024 4.0 2.81 3 1025 5.5 Experiments also were run at 858C. The value of the rate constant at 858C was found to be 7.88 3 108 (with the time in units of hours), where [CH3X]0 5 1.0 3 1022 M and [Y]0 5 3.0 M. a. Determine the rate law and the value of k for this reaction at 258C. b. Determine the half-life at 858C. c. Determine Ea for the reaction. d. Given that the COX bond energy is known to be about 325 kJ/mol, suggest a mechanism that explains the results in parts a and c