(a) Consider the combustion of H2(g): 2 H2(g) + 02(g) ----> 2 H20(g). If hydrogen is

Consider the following graph of the concentration of a substance over time. (a) Is X a reactant or product of the reaction? (b) Why is the average rate of the reaction greater between points 1 and 2 than between points 2 and 3? [Section 14.2]

You study the rate of a reaction, measuring both the concentration of the reactant and the concentration of the product as a function of time, and obtain the following results: What chemical equation is consistent with this data: (a) A -----> B, (b) B -----> A, (c) A -----> 2 B, (d) B -----> 2 A. Explain your choice. [Section 14.2]

You perform a series of experiments for the reaction A -----> B + C and find that the rate law has the form rate = k [A]'. Determine the value of x in each of the following cases: (a) There is no rate change when [A] is tripled. (b) The rate increases by a factor of 9 when [A] is tripled. (c) When [A] is doubled, the rate increases by a factor of 8. [Section 14.3]

The following diagrams represent mixtures of NO(g) and 02(g). These two substances react as follows: It has been determined experimentally that the rate is second order in NO and first order in 02. Based on this fact, which of the following mixtures will have the fastest initial rate? [Section 14.3]

A friend studies a first-order reaction and obtains the following three graphs for experiments done at two different temperatures. (a) Which two lines represent experiments done at the same temperature? What accounts for the difference in these two lines? In what way are they the same? (b) Which two lines represent experiments done with the same starting concentration but at different temperatures? Which line probably represents the lower temperature? How do you know? [Section 14.4]

(a) Given the following diagrams at t = 0 and t = 30, what is the half-life of the reaction if it follows firstorder kinetics? t = Omin t = 30min (b) After four half-life periods for a first-order reaction, what fraction of reactant remains? [Section 14.4]

The following diagram shows the reaction profile of a reaction. Label the components indicated by the boxes. [Section 14.5]

You study the effect of temperature on the rate of two reactions and graph the natural logarithm of the rate constant for each reaction as a function of 1/T. How do the two graphs compare (a) if the activation energy of the second reaction is higher than the activation energy of the first reaction but the two reactions have the same frequency factor, and (b) if the frequency factor of the second reaction is ltigher than the frequency factor of the first reaction but the two reactions have the same activation energy? [Section 14.5]

Consider the diagram below, wltich represents two steps in an overall reaction. The red spheres are oxygen, the blue ones nitrogen, and the green ones fluorine. (a) Write the chemical equation for each step in the reaction. (b) Write the equation for the overall reaction. (c) Identify the intermediate in the mechanism. (d) Write the rate law for the overall reaction if the first step is the slow, rate-determining step. [Section 14.6]

Draw a possible transition state for the bimolecular reaction depicted below. (The blue spheres are nitrogen atoms, and the red ones are oxygen atoms.) Use dashed lines to represent the bonds that are in the process of being broken or made in the transition state. [Section 14.6]

The following diagram represents an imaginary twostep mechanism. Let the red spheres represent element A, the green ones element B, and the blue ones element C. (a) Write the equation for the net reaction that is occurring. (b) Identify the intermediate. (c) Identify the catalyst. [Sections 14.6 and 14.7]

(a) What is meant by the term reaction rate? (b) Name three factors that can affect the rate of a chemical reaction. (c) What information is necessary to relate the rate of disappearance of reactants to the rate of appearance of products?

(a) What are the units usually used to express the rates of reactions occurring in solution? (b) From your everyday experience, give two examples of the effects of temperature on the rates of reactions. (c) What is the difference between average rate and instantaneous rate?

Consider the following hypothetical aqueous reaction: A(aq) --> B(aq). A flask is charged with 0.065 mol of A in a total volume of 100.0 mL. The following data are collected: Time (min) Moles of A 0 10 20 0.065 0.051 0.042 30 40 0.036 0.031(a) Calculate the number of moles of B at each time in the table, assuming that there are no molecules of B at time zero. (b) Calculate the average rate of disappearance of A for each 10-min interval, in units of Mjs. (c) Between I = 10 min and I = 30 min, what is the average rate of appearance of B in units of Mjs? Assume that the volume of the solution is constant.

A flask is charged with 0.100 mol of A and allowed to react to form B according to the hypothetical gas-phase reaction A(g) ----> B(g). The following data are collected: Time (s) 0 40 80 120 160 Moles of A 0.100 0.067 0.045 0.030 0.020 (a) Calculate the number of moles of B at each time in the table. (b) Calculate the average rate of disappearance of A for each 40-s interval, in units of moljs. (c) What additional information would be needed to calculate the rate in units of concentration per time?

The isomerization of methyl isonitrile (CH3NC) to acetonitrile (CH3CN) was studied in the gas phase at 215 C, and the following data were obtained: Time (s) (CH3NC) (M) 0 0.0165 2,000 0.0110 5,000 0.00591 8,000 0.00314 12,000 0.00137 15,000 0.00074 (a) Calculate the average rate of reaction, in M/s, for the time interval between each measurement. (b) Graph [CH3NC] versus time, and determine the instantaneous rates in M/s at I = 5000 s and I = 8000 s.

The rate of disappearance of HCI was measured for the following reaction: CH30H(aq) + HC1(aq) ----> CH3Cl(aq) + HzO(/) The following data were collected: Time (min) [HCI) (M) 0.0 1 .85 54.0 1.58 107.0 1 .36 215.0 1.02 430.0 0.580 (a) Calculate the average rate of reaction, in Mjs, for the time interval between each measurement. (b) Graph [HCl] versus time, and determine the instantaneous rates in M/min and M/s at I = 75.0 min and I = 250 min.

For each of the following gas-phase reactions, indicate how the rate of disappearance of each reactant is related to the rate of appearance of each product: (a) HzOz(g) ----> Hz(g) + Oz(g) (b) 2 N20(g) ----> 2 N2(g) + 02(g) (c) Nz(g) + 3 Hz(g) ----> 2 NH3{g)

(a) Consider the combustion of H2(g): 2 H2(g) + 02(g) ----> 2 H20(g). If hydrogen is burning at the rate of 0.85 moljs, what is the rate of consumption of oxygen? What is the rate of formation of water vapor? (b) The reaction 2 NO(g) + CI2(g) ----> 2 NOCI(g) is carried out in a closed vessel. If the partial pressure of NO is decreasing at the rate of 23 torr/min, what is the rate of change of the total pressure of the vessel?

(a) Consider the combustion of ethylene, C2H4(g) + 3 02(g) ----> 2 C02(g) + 2 H20(g). If the concentration of C2H4 is decreasing at the rate of 0.025 Mjs, what are the rates of change in the concentrations of C02 and HzO? (b) The rate of decrease in N2H4 partial pressure in a closed reaction vessel from the reaction N2H4(g) + H2(g) ----> 2 NH3(g) is 63 torr/h. What are the rates of change of NH3 partial pressure and total pressure in the vessel?

A reaction A + B ----> C obeys the following rate law: Rate = k [BjZ . (a) If [A] is doubled, how will the rate change? Will the rate constant change? Explain. (b) What are the reaction orders for A and B? What is the overall reaction order? (c) What are the units of the rate constant?

Consider a hypothetical reaction between A, B, and C that is first order in A, zero order in B, and second order in C. (a) Write the rate law for the reaction. (b) How does the rate change when [A] is doubled and the other reactant concentrations are held constant? (c) How does the rate change when [B] is tripled and the other reactant concentrations are held constant? (d) How does the rate change when [C] is tripled and the other reactant concentrations are held constant? (e) By what factor ___ does the rate change when the concentrations of all three reactants are tripled?

The decomposition of N205 in carbon tetrachloride proceeds as follows: 2 N205 ----> 4 N02 + 02. The rate law is first order in N205. At 64 oc the rate constant is 4.82 X w-3 s-1. (a) Write the rate law for the reaction. (b) What is the rate of reaction when [N205] = 0.0240 M? (c) What happens to the rate when the concentration of N205 is doubled to 0.0480 M?

Consider the following reaction: 2 NO{g) + 2 Hz{g) -----> Nz(g) + 2 HzO{g) (a) The rate law for thls reaction is first order in Hz and second order in NO. Write the rate law. (b) If the rate constant for thls reaction at 1000 K is 6.0 x 104 M-z s-1 , what is the reaction rate when [NO] = 0.035 M and [Hz] = O.Q15 M? (c) What is the reaction rate at 1000 K when the concentration of NO is increased to 0.10 M, while the concentration of Hz is 0.010 M?

Consider the following reaction: CH3Br(aq) + OH-(aq) -----> CH30H(aq) + Br-(aq) The rate law for thls reaction is first order in CH3Br and first order in OH-. When [CH3Br] is 5.0 X 10- 3 M and [OH-] is 0.050 M, the reaction rate at 298 K is 0.0432 Mjs. (a) What is the value of the rate constant? (b) What are the units of the rate constant? (c) What would happen to the rate if the concentration of OH- were tripled?

The reaction between ethyl bromide (CzH5Br) and hydroxide ion in ethyl alcohol at 330 K, C2H5Br(alc) + OH-(alc) -----> CzH50H(I) + Br-(alc), is first order each in ethyl bromide and hydroxide ion. When [CzH5Br] is 0.0477 M and [OH-] is 0.100 M, the rate of disappearance of ethyl bromide is 1.7 x 10-7 Mjs. (a) What is the value of the rate constant? (b) What are the units of the rate constant? (c) How would the rate of disappearance of ethyl bromide change if the solution were diluted by adding an equal volume of pure ethyl alcohol to the solution?

The iodide ion reacts with hypochlorite ion (the active ingredient in chlorine bleaches) in the following way: OCl- + 1- -----> 01- + cr. This rapid reaction gives the following rate data: [OCn

The following data were measured for the reaction BF3{g) + NH3{g) -----> F3BNH3{g): Initial Rate Experiment [BF3] (M) [NH3 ] (M) (M/s) 0.250 0.250 0.2130 2 0.250 0.125 0.1065 3 0.200 0.100 0.0682 4 0.350 0.100 0.1193 5 0.175 0.100 0.0596 (a) What is the rate law for the reaction? (b) What is the overall order of the reaction? (c) What is the value of the rate constant for the reaction? (d) What is the rate when [BF3] = 0.100 M and [NH3] = 0.500 M?

The following data were collected for the rate of disappearance of NO in the reaction 2 NO(g) + 02{g) -----> 2 NOz(g): Initial Rate Experiment [NO] (M) [02] (M) (M/s) 0.0126 0.0125 1.41 X 10-Z 2 0.0252 0.0125 5.64 X 10-Z 3 0.0252 0.0250 1.13 X 10-1 (a) What is the rate law for the reaction? (b) What are the units of the rate constant? (c) What is the average value of the rate constant calculated from the three data sets? (d) What is the rate of disappearance of NO when [NO] = 0.0750 M and [Ozl = 0.0100 M? (e) What is the rate of disappearance of 02 at the concentrations given in part (d)?

Consider the gas-phase reaction between nitric oxide and bromine at 273 C: 2 NO{g) + Br2(g) -----> 2 NOBr{g). The following data for the initial rate of appearance of NOBr were obtained: Initial Rate Experiment [NO] (M) [Br2l (M) (M/s) 0.10 0.20 24 2 0.25 0.20 150 3 0.10 0.50 60 4 0.35 0.50 735 (a) Determine the rate law. (b) Calculate the average value of the rate constant for the appearance of NOBr from the four data sets. (c) How is the rate of appearance of NOBr related to the rate of disappearance of Br2? (d) What is the rate of disappearance of Br2 when [NO] = O.o75 M and [Br2] = 0.25 M?

Consider the reaction of peroxydisulfate ion (SzOl-) with iodide ion (I-) in aqueous solution: SzOs 2 -(aq) + 3 r-(aq) -----> 2 504 2 -(aq) + !3 -(aq) At a particular temperature the rate of disappearance of 52082 - varies with reactant concentrations in the following manner: Initial Rate Experiment [52082 -] (M) rn (a) Determine the rate law for the reaction. (b) What is the average value of the rate constant for the disappearance of 52082 - based on the four sets of data? (c) How is the rate of disappearance of S20s2 - related to the rate of disappearance of 1-? (d) What is the rate of disappearance of ,- when [S2ol-l = 0.025 M and [r-] = 0.050 M?

(a) Define the following symbols that are encountered in rate equations: [A]0, 1 1 ;2 [A]1, k. (b) What quantity, when graphed versus time, will yield a straight line for a firstorder reaction?

(a) For a second-order reaction, what quantity, when graphed versus time, will yield a straight line? (b) How do the half-lives of first-order and second-order reactions differ?

(a) The gas-phase decomposition of S02Cl2, S02Cl2(g) ----> S02(g) + Cl2(g), is first order in S02Cl2. At 600 K the half-life for this process is 2.3 X 105 s. What is the rate constant at this temperature? (b) At 320 oc the rate constant is 2.2 x 10--s s-1. What is the half-life at this temperature?

Molecular iodine, 12(g), dissociates into iodine atoms at 625 K with a first-order rate constant of 0.271 s-1. (a) What is the half-life for this reaction? (b) If you start with 0.050 M !2 at this temperature, how much will remain after 5.12 s assuming that the iodine atoms do not recombine to form 12?

As described in Exercise 14.37, the decomposition of sulfuryl chloride (S02Cl2) is a first-order process. The rate constant for the decomposition at 660 K is 4.5 x 10- 2 s-1. (a) If we begin with an initial S02Cl2 pressure of 375 torr, what is the pressure of this substance after 65 s? (b) At what time will the pressure of S02Cl2 decline to one-tenth its initial value?

The reaction S02Cl2(gl --> S02(gl + Cl2(g) is first order in S02Cl2. Using the following kinetic data, determine the magnitude of the first-order rate constant: Time (s) 0 2,500 5,000 7,500 10,000 Pressure S02C12 (atm) 1.000 0.947 0.895 0.848 0.803

From the following data for the first-order gas-phase isomerization of CH3NC at 215 C, calculate the firstorder rate constant and half-life for the reaction: Time (s) 0 2,000 5,000 8,000 12,000 15,000 Pressure CH3NC (torr) 502 335 180 95.5 41.7 22.4

Consider the data presented in Exercise 14.15. (a) By using appropriate graphs, determine whether the reaction is first order or second order. (b) What is the value of the rate constant for the reaction? (c) What is the halflife for the reaction?

Consider the data presented in Exercise 14.16. (a) Determine whether the reaction is first order or second order. (b) What is the value of the rate constant? (c) What is the half-life?

The gas-phase decomposition of N02, 2 N02(g) --> 2 NO(g) + 02(g), is studied at 383 C, giving the following data: Time (s) [N02] (M) 0.0 0.100 5.0 0.017 10.0 0.0090 15.0 0.0062 20.0 0.0047 (a) Is the reaction first order or second order with respect to the concentration of N02? (b) What is the value of the rate constant?

Sucrose (C12H22011), which is commonly known as table sugar, reacts in dilute acid solutions to form two simpler sugars, glucose and fructose, both of which have the formula C6H1206: At 23 oc and in 0.5 M HCI, the following data were obtained for the disappearance of sucrose: Time (min) [C12H220ul (M) 0 0.316 39 0.274 80 0.238 140 0.190 210 0.146 (a) Is the reaction first order or second order with respect to [C12H22011]? (b) What is the value of the rate constant?

(a) What factors determine whether a collision between two molecules will lead to a chemical reaction? (b) According to the collision model, why does temperature affect the value of the rate constant?

(a) In which of the following reactions would you expect the orientation factor to be least important in leading to reaction? NO + 0 -> N02 or H + Cl -> HCI? (b) How does the kinetic-molecular theory help us understand the temperature dependence of chemical reactions?

Calculate the fraction of atoms in a sample of argon gas at 400 K that has an energy of 10.0 kJ or greater

The gas-phase reaction Cl(g) + HBr(g) -> HCl(g) + Br(g) has an overall enthalpy change of -66 kJ. The activation energy for the reaction is 7 kj. (a) Sketch the energy profile for the reaction, and label Ea and !lE. (b) What is the activation energy for the reverse reaction?

For the elementary process N20s(g) -> N02(g) + N03(g) the activation energy (0) and overall !lE are 154 k)/mol, and 136 k)/mol, respectively. (a) Sketch the energy profile for this reaction, and label Ea and !lE. (b) What is the activation energy for the reverse reaction?

Based on their activation energies and energy changes and assuming that all collision factors are the same, which of the following reactions would be fastest and which would be slowest? Explain your answer. (a) Ea = 45 k)/mol;!lE = -25 kJ/mol (b) Ea = 35 k)/mol;!lE = -10 kJ/mol (c) Ea = 55 k)/mol;!lE = 10 kJ/mol

Which of the reactions in Exercise 14.53 will be fastest in the reverse direction? Which will be slowest? Explain.

A certain first-order reaction has a rate constant of 2.75 x 10-2 s -I at 20 oc. What is the value of k at 60 oc if (a) Ea = 75.5 k)/mol; (b) Ea = 125 k)/mol?

Understanding the high-temperature behavior of nitrogen oxides is essential for controlling pollution generated in automobile engines. The decomposition of nitric oxide (NO) to N2 and 02 is second order with a rate constant of 0.0796 M-1 s-1 at 737 oc and 0.0815 M-1 s-1 at 947 oc. Calculate the activation energy for the reaction.

The rate of the reaction CH3COOC2H5(aq) + OH-(aq) -> CH3COO-(aq) + C2H50H(aq) was measured at several temperatures, and the following data were collected: Temperature (C) 15 25 35 45 0.0521 0.101 0.184 0.332 Using these data, graph Ink versus 1/T. Using your graph, determine the value of E0

The temperature dependence of the rate constant for the reaction is tabulated as follows: Temperature (K) 600 650 700 750 800 Calculate Ea and A. 0.028 0.22 1.3 6.0 23

The activation energy of a certain reaction is 65.7 k)/mol. How many times faster will the reaction occur at 50 oc than at 0 C?

(a) What is meant by the term elementary reaction? (b) What is the difference between a unimolecular and a bimolecular elementary reaction? (c) What is a reaction mechanism?

(a) What is meant by the term molecularity? (b) Why are termolecular elementary reactions so rare? (c) What is an intermediate in a mechanism?

What is the molecularity of each of the following elementary reactions? Write the rate law for each. (a) CJ2(g) --> 2 Cl(g) (b) OCqaq) + H20(1) --> HOCl(aq) + OW(aq) (c) NO(g) + Cl2(g) --> NOCl2(gl

What is the molecularity of each of the following elementary reactions? Write the rate law for each. (a) 2 NO(g) --> N202(g) 72 (b) H2C - CH2(g) --> CH2=CH-CH3(g) (c) S03(g) __, S02(gl + O(g)

(a) Based on the following reaction profile, how many intermediates are formed in the reaction A --> D? (b) How many transition states are there? (c) Which step is the fastest? (d) Is the reaction A --> D exothermic or endothermic?

Consider the following energy profile. Reaction pathway (a) How many elementary reactions are in the reaction mechanism? (b) How many intermediates are formed in the reaction? (c) Which step is rate limiting? (d) Is the overall reaction exothermic or endothermic?

The following mechanism has been proposed for the gas-phase reaction of H2 with ICl: H#) + ICl(g) --> HI(g) + HCI(g) Hl(g) + ICl(g) --> l2(g) + HCl(g) (a) Write the balanced equation for the overall reaction. (b) Identify any intermediates in the mechanism. (c) Write rate laws for each elementary reaction in the mechanism. (d) If the first step is slow and the second one is fast, what rate law do you expect to be observed for the overall reaction?

The decomposition of hydrogen peroxide is catalyzed by iodide ion. The catalyzed reaction is thought to proceed by a two-step mechanism: (a) Write the rate law for each of the elementary reactions of the mechanism. (b) Write the chemical equation for the overall process. (c) Identify the intermediate, if any, in the mechanism. (d) Assuming that the first step of the mechanism is rate determining, predict the rate law for the overall process.

The reaction 2 NO(g) + CJ2(g) --> 2 NOCl(g) obeys the rate law, rate = k[NOj2 [CJ2l The following mechanism has been proposed for this reaction: NO(g) + Cl2(g) --> NOCl2(g) NOCI2(g) + NO(g) --> 2 NOCI(g) (a) What would the rate law be if the first step were rate determining? (b) Based on the observed rate law, what can we conclude about the relative rates of the two steps?

(a) What part of the energy profile of a reaction is affected by a catalyst? (b) What is the difference between a homogeneous and a heterogeneous catalyst?

(a) Most heterogeneous catalysts of importance are extremely finely divided solid materials. Why is particle size important? (b) What role does adsorption play in the action of a heterogeneous catalyst?

The oxidation of S02 to S03 is catalyzed by NOz. Thereaction proceeds as follows: N02(g) + S02(g) NO(g) + S03(g) 2 NO(g) + 02(g) 2 N02(g) (a) Show that the two reactions can be summed to give the overall oxidation of S02 by 02 to give S03. (Hi11t: The top reaction must be multiplied by a factor so the NO and N02 cancel out.) (b) Why do we consider N02 a catalyst and not an intermediate in this reaction? (c) Is this an example of homogeneous catalysis or heterogeneous catalysis?

NO catalyzes the decomposition of N20, possibly by the following mechanism: NO(g) + N20(g) Nz(g) + NOz(g) 2 N02(g) 2 NO(g) + 02(g) (a) What is the chemical equation for the overall reaction? Show how the two steps can be added to give the overall equation. (b) Why is NO considered a catalyst and not an intermediate? (c) If experiments show that during the decomposition of NzO, N02 does not accumulate in measurable quantities, does this rule out the proposed mechanism? If you think not, suggest what might be going on.

Many metallic catalysts, particularly the precious-metal ones, are often deposited as very thin films on a substance of high surface area per unit mass, such as alumina (Al203 ) or silica (Si02). (a) Why is this an effective way of utilizing the catalyst material? (b) How does the surface area affect the rate of reaction?

(a) If you were going to build a system to check the effectiveness of automobile catalytic converters on cars, what substances would you want to look for in the car exhaust? (b) Automobile catalytic converters have to work at high temperatures, as hot exhaust gases stream through them. In what ways could this be an advantage? In what ways a disadvantage? (c) Why is the rate of flow of exhaust gases over a catalytic converter important?

When D2 reacts with ethylene (C2H4 ) in the presence of a finely divided catalyst, ethane with two deuteriums, CHzD-CH2D, is formed. (Deuterium, D, is an isotope of hydrogen of mass 2.) Very little ethane forms in which two deuteriums are bound to one carbon (for example, CH3-CHD2). Use the sequence of steps involved in the reaction to explain why this is so.

Heterogeneous catalysts that perform hydrogenation reactions, as illustrated in Figure 14.21, are subject to poisoning, which shuts down their catalytic ability. Compounds of sulfur are often poisons. Suggest a mechanism by which such compounds might act as poisons.

(a) Explain the importance of enzymes in biological systems. (b) What chemical transformations are catalyzed (i) by the enzyme catalase, (ii) by nitrogenase?

The activation energy of an uncatalyzed reaction is 95 k)/mol. The addition of a catalyst Jowers the activation energy to 55 k)/mol. Assuming that the collision factor remains the same, by what factor will the catalyst increase the rate of the reaction at (a) 25 oc, (b) 125 C?

Suppose that a certain biologically important reaction is quite slow at physiological temperature (37 oq in the absence of a catalyst. Assuming that the collision factor remains the same, by how much must an enzyme lower the activation energy of the reaction in order to achieve a 1 X 105-fold increase in the reaction rate?

Explain why rate laws generally cannot be written from balanced equations. Under what circumstance is the rate law related directly to the balanced equation for a reaction?

Hydrogen sulfide (H2S) is a common and troublesome pollutant in industrial wastewaters. One way to remove H2S is to treat the water with chlorine, in which case the following reaction occurs: HzS(aq) + Clz(aq) S(s) + 2 H+(aq) + 2 Cqaq) The rate of this reaction is first order in each reactant. The rate constant for the disappearance of H25 at 28 C is 3.5 X 10-2 M-1 s-1 If at a given time the concentration of H2S is 2.0 X 10-4 M and that of Cl2 is 0.025 M, what is the rate of formation of cr?

The reaction 2 NO(g) + 02(g) 2 N02(g) is second order in NO and first order in 02. When [NO] = 0.040 M and [Oz] = 0.035 M, the observed rate of disappearance of NO is 9.3 X 10-5 Mfs. (a) What is the rate of disappearance of 02 at this moment? (b) What is the value of the rate constant? (c) What are the units of the rate constant? (d) What would happen to the rate if the concentration of NO were increased by a factor of 1.8?

Consider the following reaction between mercury(II) chloride and oxalate ion: 2 HgC12(aq) + C20/-(aq) 2 Cqaq) + 2 COz(g) + HgzClz(s) The initial rate of this reaction was determined for several concentrations of HgC12 and C2042 -, and the following rate data were obtained for the rate of disappearance of c2o/-: Experiment [HgC12) (M) (C20lj (M) Rate (M/s) 0.164 0.15 3.2 x 10-5 2 0.164 0.45 2.9 x 10-4 3 0.082 0.45 1.4 x 10- 4 0.246 0.15 4.8 x 10-5 (a) What is the rate law for this reaction? (b) What is the value of the rate constant? (c) What is the reaction rate when the concentration of HgCl2 is 0.100 M and that of (C20/-) is 0.25 M, if the temperature is the same as that used to obtain the data shown?

The reaction 2 N02 --+ 2 NO + 02 has the rate constant k = 0.63 M-1 s-1 Based on the units for k, is the reaction first or second order in N02? If the initial concentration of N02 is 0.100 M, how would you determine how long it would take for the concentration to decrease to 0.025 M?

Consider two reactions. Reaction (1) has a constant halflife, whereas reaction (2) has a half-life that gets longer as the reaction proceeds. What can you conclude about the rate laws of these reactions from these observations?

(a) The reaction H202(aq) --+ H20(/) + 02(g), is first order. Near room temperature, the rate constant equals 7.0 X 10-4 s-1 Calculate the half-life at this temperature. (b) At 415 C, (CH2h0 decomposes in the gas phase, (CH2h0(g) --+ CH4(g) + CO(g). If the reaction is first order with a half-life of 56.3 min at this ternperature, calculate the rate constant in s-1 .

Urea (NH2CONH2) is the end product in protein metabolism in animals. The decomposition of urea in 0.1 M HCI occurs according to the reaction NH2CONH2(aq) + H+(aq) + 2 H20(/) --+ 2 NH4 +(aq) + HC03 -(aq) The reaction is first order in urea and first order overall. When [NH2CONH2.) = 0.200 M, the rate at 61.05 oc is 8.56 X 10-5 M/s. (a) What is the value for the rate constant, k? (b) What is the concentration of urea in this solution after 4.00 x 103 s if the starting concentration is 0.500 M? (c) What is the half-life for this reaction at 61.05 C?

The rate of a first-order reaction is followed by spectroscopy, monitoring the absorption of a colored reactant. The reaction occurs in a 1.00-crn sample cell, and the only colored species in the reaction has a molar absorptivity constant of 5.60 x 103 crn-1 M-1 . (a) Calculate the initial concentration of the colored reactant if the absorbance is 0.605 at the beginning of the reaction. (b) The absorbance falls to 0.250 within 30.0 min. Calculate the rate constant in units of s-1 . (c) Calculate the half-life of the reaction. (d) How long does it take for the absorbance to fall to 0.100?

Cyclopentadiene (CgH6} reacts with itself to form dicyclopentadiene (C10Hn). A 0.0400 M solution of C5 H6 was monitored as a function of time as the reaction 2 C5H6 --+ C10H12 proceeded. The following data were collected: Time (s) ICsH6) (M) 0.0 0.0400 50.0 0.0300 100.0 0.0240 150.0 0.0200 200.0 0.0174 Plot [CsH6] versus time, ln [CsH6] versus time, and 1/[CsH6l versus time. What is the order of the reaction? What is the value of the rate constant?

(a) Two reactions have identical values for E,. Does this ensure that they will have the same rate constant if run at the same temperature? Explain. (b) Two similar reactions have the same rate constant at 25 oc, but at 35 oc one of the reactions has a higher rate constant than the other. Account for these observations.

The first-order rate constant for reaction of a particular organic compound with water varies with temperature as follows: Temperature (K) 300 320 340 355 Rate Constant (s-1) 3.2 X 10-11 1.0 X 10-'J 3.0 X 10-B 2.4 X 10-7 From these data, calculate the activation energy in units of kjfmol.

The following mechanism has been proposed for the reaction of NO with H2 to form N20 and H20: NO(g) + NO(g) --+ N202(g) N202(g) + H2(g) --+ N20(g) + H20(g) (a) Show that the elementary reactions of the proposed mechanism add to provide a balanced equation for the reaction. (b) Write a rate law for each elementary reaction in the mechanism. (c) Identify any intermediates in the mechanism. (d) The observed rate law is rate = k [NOf[H2]. If the proposed mechanism is correct, what can we conclude about the relative speeds of the first and second reactions?

Ozone in the upper atmosphere can be destroyed by the following two-step mechanism: Cl(g) + 03 (g) --+ CIO(g) + 02(g) ClO(g) + O(g) --+ Cl(g) + 02(g) (a) What is the overall equation for this process? (b) What is the catalyst in the reaction? How do you know? (c) What is the intermediate in the reaction? How do you distinguish it from the catalyst?

Using Figure 14.20 as your basis, draw the energy profile for the bromide ion-catalyzed decomposition of hydrogen peroxide. (a) Label the curve with the activation energies for reactions [14.30] and [14.31]. (b) Notice from Figure 14.19(b) that when Br-(aq) is added initially, Br2 accumulates to some extent during the reaction. What does this tell us about the relative rates of reactions [14.30] and [14.31]?

The following mechanism has been proposed for the gas-phase reaction of chloroform (CHC13) and chlorine: kl Step 1: Cl2(g) k_1 2 Cl(g) (fast) Step 2: Cl{g) + CHC13(g) HCl{g) + CC13(g) (slow) Step 3: Cl{g) + CC13(g) CC14 (fast) (a) What is the overall reaction? (b) What are the intermediates in the mechanism? (c) What is the molecularity of each of the elementary reactions? (d) What is the rate-determining step? (e) What is the rate law predicted by this mechanism? (Hint: The overall reaction order is not an integer.)

One of the many remarkable enzymes in the human body is carbonic anhydrase, which catalyzes the interconversion of carbonic acid with carbon dioxide and water. If it were not for this enzyme, the body could not rid itself rapidly enough of the C02 accumulated by cell metabolism. The enzyme catalyzes the dehydration (release to air) of up to 107 C02 molecules per second. Which components of this description correspond to the terms enzyme, substrate, and turnover number?

Enzymes are often described as following the two-step mechanism: E + S :;===' ES (fast) ES ----> E + P (slow) Where E = enzyme, S = substrate, and P = product. If an enzyme follows this mechanism, what rate law is expected for the reaction?

The enzyme invertase catalyzes the conversion of sucrose, a disaccharide, to invert sugar, a mixture of glucose and fructose. When the concentration of invertase is 4.2 x 1o-7M and the concentration of sucrose is 0.0077 M, invertsugarisformed attherate of1.5 x 10-4 Mjs. When the sucrose concentration is doubled, the rate of formation of invert sugar is doubled also. (a) Assuming that the enzyme-substrate model is operative, is the fraction of enzyme tied up as a complex large or small? Explain. (b) Addition of inositol, another sugar, decreases the rate of formation of invert sugar. Suggest a mechanism by which this occurs.

Dinitrogen pentoxide (N205) decomposes in chloroform as a solvent to yield N02 and 02. The decomposition is first order with a rate constant at 45 oc of 1.0 X 10-5 s -l Calculate the partial pressure of 02 produced from 1.00 L of 0.600 M N205 solution at 45 oc over a period of 20.0 h if the gas is collected in a 10.0-L container. (Assume that the products do not dissolve in chloroform.)

The reaction between ethyl iodide and hydroxide ion in ethanol (C2H50H) solution, C2HsJ(alc) + OH-(alc) ----> C2H50H{l) + r-(alc), has an activation energy of 86.8 kJ/mol and a frequency factor of2.10 x 1011 M-1 s-1 (a) Predict the rate constant for the reaction at35 "C. (b) A solution of KOH in ethanol is made up by dissolving 0.335 g KOH in ethanol to form 250.0 mL of solution. Similarly, 1.453 g of C2H5I is dissolved in ethanol to form 250.0 mL of solution. Equal volumes of the two solutions are mixed. Assuming the reaction is first order in each reactant, what is the initial rate at 35 C? (c) Which reagent in the reaction is limiting, assuming the reaction proceeds to completion?

Zinc metal dissolves in hydrochloric acid according to the reaction Zn(s) + 2 HCI(aq) ----> ZnC12(aq) + H2{g) Suppose you are asked to study the kinetics of this reaction by monitoring the rate of production of H2(g). (a) By using a reaction flask, a manometer, and any other common laboratory equipment, design an experimental apparatus that would allow you to monitor the partial pressure of H2{g) produced as a function of time. (b) Explain how you would use the apparatus to determine the rate law of the reaction. (c) Explain how you would use the apparatus to determine the reaction order for [H+] for the reaction. (d) How could you use the apparatus to determine the activation energy of the reaction? (e) Explain how you would use the apparatus to determine the effects of changing the form of Zn(s) from metal strips to granules.

The gas-phase reaction of NO with F2 to form NOF and F has an activation energy of Ea = 6.3 kJ/mol and a frequency factor of A = 6.0 X 108 M-1 s-1. The reaction is bel.ieved to be bimolecular: NO(g) + l':l(g) --+ NOF(g) + F(g) (a) Calculate the rate constant at 100 c. (b) Draw the Lewis structures for the NO and the NOF molecules, given that the chemical formula for NOF is misleading because the nitrogen atom is actually the central atom in the molecule. (c) Predict the structure for the NOF molecule. (d) Draw a possible transition state for the formation of NOF, using dashed l.ines to indicate the weak bonds that are beginning to form. (e) Suggest a reason for the low activation energy for the reaction.

The mechanism for the oxidation of HBr by 02 to form 2 H20 and Br2 is shown in Exercise 14.70. (a) Calculate the overall standard enthalpy change for the reaction process. (b) HBr does not react with 02 at a measurable rate at room temperature under ordinary conditions. What can you infer from this about the magnitude of the activation energy for the rate-determining step? (c) Draw a plausible Lewis structure for the intermediate HOOBr. To what familiar compound of hydrogen and oxygen does it appear similar?

Enzymes, the catalysts of biological systems, are high molecular weight protein materials. The active site of the enzyme is formed by the three-dimensional arrangement of the protein in solution. When heated in solution, proteins undergo denaturation, a process in which the three-dimensional structure of the protein unravels or at least partly does so. The accompanying graph shows the variation with temperature of the activity of a typical enzyme. The activity increases with temperature to a point above the usual operating region for the enzyme, then decl.ines rapidly with further temperature increases. What role does denaturation play in determining the shape of this curve? How does your explanation fit in with the lock-and-key model of enzyme action?

The rates of many atmospheric reactions are accelerated by the absorption of light by one of the reactants. For example, consider the reaction between methane and chlorine to produce methyl chloride and hydrogen chloride: Reaction 1: CH4(g) + Clz(g) --+ CH3Cl(g) + HCI(g) This reaction is very slow in the absence of light. However, Cl2(g) can absorb light to form Cl atoms: Reaction 2: Cl2(g) + hv --+ 2 Cl(g) Once the Cl atoms are generated, they can catalyze the reaction of CH4 and Cl2, according to the following proposed mechanism: Reaction 3: CH4 (g) + Cl(g) --+ CH3(g) + HCI(g) Reaction 4: CH 3 (g) + Cl2(g) --+ CH3 Cl(g) + Cl(g) The enthalpy changes and activation energies for these two reactions are tabulated as follows: Reaction 3 4 ll.Han (kJ/mol) +4 -109 E. (kJ/mol) 17 4 (a) By using the bond enthalpy for Cl2 (Table 8.4), determine the longest wavelength of light that is energetic enough to cause reaction 2 to occur. ln which portion of the electromagnetic spectrum is this light found? (b) By using the data tabulated here, sketch a quantitative energy profile for the catalyzed reaction represented by reactions 3 and 4. (c) By using bond enthalpies, estimate where the reactants, CH4(g) + C12(g), should be placed on your diagram in part (b). Use this result to estimate the value of E. for the reaction CH4(g) + Cl2(g) --+ CH3 (g) + HCI(g) + Cl(g). (d) The species Cl(g) and CH3(g) in reactions 3 and 4 are radicals, that is, atoms or molecules with unpaired electrons. Draw a Lewis structure of CH:v and verify that it is a radical. (e) The sequence of reactions 3 and 4 comprise a radical chain mechanism. Why do you think this is called a "chain reaction"? Propose a reaction that will terminate the chain reaction.