Ch. 12-13: Kinetics and Equilibrium AP Review Questions **Know rate constant units: - (Order – 1) = power on M, power on time is always -1. Rate = k [A]2[B]1 units of k = M-(3-1)min-1 = M-2 min-1

1) Which of the following must be true for a reaction that proceeds spontaneously from initial standard state conditions?

a) ∆Go > 0 and Keq > 1 b) ∆Go > 0 and Keq < 1 c) ∆Go < 0 and Keq > 1 d) ∆Go < 0 and Keq < 1 e) ∆Go = 0 and Keq = 1

2) If ∆Go > 0 for a reaction, which of the

following statements about the equilibrium constant, K, is true?

a) K = 0 b) K > 1 c) K < 1 d) K = 1 e) K = Go

Questions 3-6 a) VSEPR theory b) Kinetic molecular theory c) Transition state theory d) Quantum theory e) Atomic Theory

3) Theory used to determine, predict, or explain

molecular geometry. 4) Theory that explains paramagnetism the best. 5) Theory that incorporates the activation energy. 6) Theory used to explain the effect of

temperature on reaction kinetics.

Questions 7-9 a) ideal gas law b) law of conservation of matter c) equilibrium law d) rate law e) Hess’ law

7) This law describes the order of a reaction. 8) This law can be used to determine molar

masses. 9) This law is fundamental to all stoichiometry

calculations.

10) Which of the following is an integral part of collision theory?

1. collision frequency 2. collision energy 3. collision orientation a) only 1 and 2 b) only 1 and 3 c) only 1 d) 1, 2, and 3 e) only 2 and 3

11) A reaction has the following rate law: Rate =

k [A][B]2. The units for the rate constant are a) s-1 b) mol L-1 s-1 c) L2 mol-2 s-1 d) L mol-1 s-1 e) mol s-1

12) The kinetic molecular theory is used to explain a) reaction rates b) bond vibrations c) gas behavior d) catalysts e) activated complexes

13) A certain chemical reaction is described by a

first-order rate law. What is the rate constant, in units of s-1, for the reaction if the half-life is determined to be 228 min?

a) ln (228 X 60) b) ln2 / [228 X 60] c) ln 2 d) ln ([At]/[Ao]) e) 228 s-1

14) The rate law for a reaction is found to be rate = k [A]2 [B]. Which of the following is true about this system?

a) A plot of log rate versus time is a straight line.

b) The units for the rate constant are mol2 L-2 s-1.

c) This reaction is unlikely since it implies the simultaneous collision of two atoms of A and one of B.

d) It is unlikely that the first step of the mechanism is the rate-limiting step.

e) All third order reactions are endothermic.

15) The addition of a catalyst to a chemical reaction will bring about a change in which of the following characteristics of the reaction?

I. the activation energy II. the enthalpy change III. the value of the equilibrium constant a) I only b) II only c) I and II only d) I and III only e) II and III only

16) All of the following affect reaction rates when

all reactants are in solution EXCEPT a) nature of reaction substances b) pressure c) temperature of reactants d) concentration of reactants e) presence of catalyst

17) The rate constants for the decomposition of

acetaldehyde were measured at five different temperatures. A plot of ln k versus 1/T was found to give a straight line. The slope of this plot is equal to

a) –k b) k c) –Ea / R d) Ea e) A

Questions 18-19 F2(g) + 2 ClO2(g) → 2 FClO2(g) Experiment [F2] [ClO2] initial rate of (M) (M) disappearance of F2 (M/sec) 1 0.10 0.010 1.2 X 10-3 2 0.20 0.010 2.4 X 10-3 3 0.40 0.020 9.6 X 10-3 18) Based on the data given in the above table,

which of the following expressions is equal to the rate law for the reaction given above.

a) Rate = k [F2] b) Rate = k [ClO2] c) Rate = k [F2] [ClO2] d) Rate = k [F2]2 [ClO2] e) Rate = k [F2] [ClO2]2

19) What is the initial rate of disappearance of

ClO2 in experiment 2? a) 1.2 X 10-3 M/sec b) 2.4 X 10-3 M/sec c) 4.8 X 10-3 M/sec d) 7.0 X 10-3 M/sec e) 9.6 X 10-3 M/sec

20) Dinitrogen pentoxide decomposes according to the following balanced equation:

N2O5(g) → 2 NO2(g) + ½ O2(g) The rate of decomposition was found to be 0.80 mole liter-1 sec-1 at a given concentration and temperature. What would the rate be for the formation of oxygen gas under the same conditions? a) 0.20 mole liter-1 sec-1 b) 0.40 mole liter-1 sec-1 c) 0.80 mole liter-1 sec-1 d) 1.60 mole liter-1 sec-1 e) 3.20 mole liter-1 sec-1

21) Time(min) [A] M [B] M 0 0.50 6.00 10 0.36 6.00 20 0.25 6.00 30 0.18 6.00 40 0.13 6.00 A reaction occurred in which A reacted with a large excess of B to form C. The concentrations of the reactants were measured periodically and recorded in the chart above. Based on the data in the chart, which of the following statements is NOT true?

a) The reaction is first order in [A]. b) The reaction is first order overall. c) The rate of reaction is constant over time. d) The half-life of reactant A is 20 minutes. e) The graph of ln [A] will be a straight line.

22) For the reaction: 2 NO(g) + O2(g) → 2 NO2(g) which two of the following possible intermediate mechanisms would support this reaction?

1. 2 NO(g) → N2O2(g) 2. NO(g) + O2(g) → NO3(g) 3. 2 NO2(g) → N2O2(g) + O2(g) 4. NO3(g) + NO(g) → 2 NO2(g) 5. NO(g) → NO(g) + O2(g)

a) 1 and 2 b) 2 and 3 c) 3 and 4 d) 2 and 4 e) 1 and 4

Questions 23-26

a) activation energy b) free energy c) ionization energy d) kinetic energy e) lattice energy 23) The energy required to convert a ground state atom in the gas phase to a gaseous positive ion. 24) The energy change that occurs in the conversion of an ionic solid to widely separated gaseous ions. 25) The energy in a chemical or physical change that is available to do useful work. 26) The energy required to form the transition state in a chemical reaction.

27) Which of the following best describes the role of the spark from the spark plug in an automobile engine?

a) The spark decreases the energy of activation for the slow step.

b) The spark increases the concentration of the volatile reactant.

c) The spark supplies some of the energy of activation for the combustion reaction.

d) The spark provides a more favorable activated complex for the combustion reaction.

e) The spark provides the heat of vaporization for the volatile hydrocarbon.

Questions 28-32

a) activation energy, Eact b) standard entropy of formation, Sf

o c) enthalpy of reaction, ∆Hrxn

o d) total entropy of change for the universe,

∆Souniverse

e) free energy of formation, ∆Gfo

28) Equals the energy difference between reactants and the transition state 29) Is always greater than or equal to zero 30) Is defined as zero for pure elements in their standard state 31) Indicates the amount of disorder in a pure substance at the standard state 32) Is always negative for an exothermic reaction

33) Activation energy and enthalpy of reaction data are listed, and one of these is wrong. Which of these combinations is impossible?

f) Activation energy 100 kJ and an enthalpy of reaction of +200 kJ

g) Activation energy 250 kJ and an enthalpy of reaction of +200 kJ

h) Activation energy 200 kJ and an enthalpy of reaction of -200 kJ

i) Activation energy 50 kJ and an enthalpy of reaction of -200 kJ

j) Activation energy 250 kJ and an enthalpy of reaction of +100 kJ

34) Refer to the diagram:

Which will be affected by the addition of a catalyst?

I. A II. B III. C a) I only b) II only c) III only d) I and II e) II and III

35) The activation energy for the reverse of the

reaction in the following diagram is a) A—B b) B—C c) C—B d) C—A e) A—C

36) In the Arrhenius equation, -Ea/RT, the exponential factor represents a) collision frequency b) activation energy c) activated complex d) fraction of molecules having energy of activation e) energy of activated complex

37) Based on the data below for a reaction in which A and B react to form C, what is the rate law for the reaction?

Initial Rate of Formation of

[A] (mol L-1) [B] (mol L-1) C (mol L-1 s-1) 0.2 0.2 0.50 0.4 0.2 2.00 0.8 0.2 8.00 0.2 0.4 1.00 0.2 0.8 2.00

a) rate = k [A] [B] b) rate = k [A]2 [B] c) rate = k [A] [B]2 d) rate = k [A]2 [B]2 e) rate = k [A]3

38) Rate = k [NO]2 [Cl2] What is the order of the reaction with respect to nitric oxide, NO? a) 0 b) 1 c) 2 d) 3 e) 4

39) CO(g) + NO2(g) → CO2(g) + NO(g) For the above reaction, the experimental rate law is given as follows: Rate = k [NO2]2 If additional CO gas is added to the reaction vessel, while temperature remains constant, which of the following is true? a) Both the reaction rate and k increase. b) Both the reaction rate and k decrease. c) Both the reaction rate and k remain the same. d) The reaction rate increases, but k remains the

same. e) The reaction rate decreases, but k remains the

same.

**40) 2 NO ↔ N2O2 (fast equilibrium) k1 for forward rate constant k-1 for reverse rate constant N2O2 + H2 → N2O + H2O (slow) k2 for forward rate constant N2O + H2 → N2 + H2O (fast) k3 for forward rate constant Nitric oxide, NO, can be reduced by hydrogen gas to yield nitrogen gas and water vapor. The decomposition is believed to occur according to the reaction mechanism shown above. The rate law for the reaction that is consistent with this mechanism is given by which of the following? a) Rate = k [NO]2 b) Rate = k [NO]2 [N2O2] c) Rate = k [N2O2] [H2] d) Rate = k [NO]2 [H2] e) Rate = k [N2O] [H2]

41) 2 NO(g) + O2(g) → 2 NO2(g) The reaction above occurs by the following two step process: Step I: NO(g) + O2(g) → NO3(g) Step II. NO3(g) + NO(g) → 2 NO2(g) Which of the following is true of Step II if it is the rate-limiting step? a) Step II has a lower activation energy and

occurs more slowly than Step I. b) Step II has a higher activation energy and

occurs more slowly than Step I. c) Step II has a lower activation energy and

occurs more quickly than Step I. d) Step II has a higher activation energy and

occurs more quickly than Step I. e) Step II has the same activation energy and the

same speed as Step I.

42) A plot of logarithm of reactant concentration versus time is linear for

a) a zero-order reaction b) a first-order reaction c) a second-order reaction d) a third-order reaction e) none of the above

43) For the reaction A + 3 B → 4 C, the rate of disappearance of B can be expressed as a) 1/3 ∆[B]/∆t b) –1/3 ∆[B]/∆t c) –3/4 ∆[B]/∆t d) –4/3 ∆[B]/∆t e) –1/4 ∆[B]/∆t

44) Acetaldehyde decomposes when heated to yield methane, CH4, and carbon monoxide, CO, according to the following equation: CH3CHO(g) → CH4(g) + CO(g) Experimental data show that the rate increases by a factor of four when the concentration of acetaldehyde is doubled. The rate law for the reaction a) cannot be determined from the given

information b) is rate = k [CH3CHO] c) is rate = k [CH3CHO]2 d) is rate = k [CH3CHO]4 e) is rate = k4

45) Which of the following is true for a zero-order reaction?

a) The rate constant equals zero. b) The activation energy equals zero. c) The half-life time equals 0.693/k. d) The rate is independent of

concentration. e) The rate is independent of

temperature.

46) For the reaction 2 NO2 + O3 → N2O5 + O2, the proposed mechanism is Step 1 NO2 + O3 → NO3 + O2 (slow) Step 2 NO3 + NO2 → N2O5 (fast) The rate law for the reaction is a) k [NO2] [O3] b) k [NO3] [NO2] c) k [NO3] [O2] d) k [NO2]2 [O3] e) more information is needed to determine the

rate law

47) The following graph represents the disappearance of a reactant in a kinetics experiment. What is the initial rate of disappearance of this reactant?

a) 4 X10-3 mol L-1 s-1 b) 2.00 mol c) 45 mol L-1 s-1 d) 0.022 mol L-1 s-1 e) 90

48) A concentration-versus-time plot is curved for all of the following except

f) a zero-order reaction g) a first-order reaction h) a second-order reaction i) a third-order reaction j) none of the above

49) If a reactant concentration is doubled, and the reaction rate increases by a factor of 8, the exponent for that reactant in the rate law should be a) ¼ b) ½ c) 2 d) 3 e) 4

50) Which of the following is not commonly used as a catatlyst? a) NaCl b) Pt c) Au d) enzymes e) MnO2

51) The reaction 2 HgCl2 + C2O4

2- → 2Cl- + 2 CO2 + Hg2Cl2(s) has a rate law of Rate = k [HgCl2] [C2O4

2-]2

Overall this is a ________ -order reaction. a) first b) second c) zero d) third e) fourth

52) Solid calcium carbonate decomposes to produce solid calcium oxide and carbon dioxide gas. The value of ∆Go for this reaction is 130.24 kJ/mole. Calculate ∆G at 100 oC for this reaction if the pressure of the carbon dioxide gas is 1.00 atm. a) –998.56 kJ/mole b) –604.2 kJ/mole c) 56.31 kJ/mole d) 130.24 kJ/mole e) 256.24 kJ/mole

53) 2 NO(g) + 2 H2(g) → N2(g) + 2 H2O(g) Which of the following is true regarding the relative molar rates of disappearance of the reactants and appearance of the products? I. N2 appears at the same rate that H2 disappears II. H2O appears at the same rate that NO disappears III. NO disappears at the same rate that H2 disappears. a) I only b) I and II only c) I and III only d) II and III only e) I, II, and III

54) If the temperature at which a reaction takes place is increased, the rate of the reaction will

a) increase if the reaction is endothermic and decrease if the reaction is exothermic

b) decrease if the reaction is endothermic and increase if the reaction is exothermic

c) increase if the reaction is endothermic and increase if the reaction is exothermic

d) decrease if the reaction is endothermic and decrease if the reaction is exothermic

e) remain the same for both an endothermic and an exothermic reaction

55) The reaction of elemental chlorine with ozone in the atmosphere occurs by the two-step process shown below.

I. Cl + O3 → ClO + O2 II. ClO + O → Cl + O2

Which of the statements below is true regarding this process? a) Cl is a catalyst b) O3 is a catalyst c) ClO is a catalyst d) O2 is an intermediate e) O is an intermediate

56) A set of reactants can follow two reaction paths. One path produces “A” and the other produces “I”. At 30 oC, both “A” and “I” are produced at the same rate. At 50 oC, “A” is produced faster than “I”. Which of the following statements is true?

a) The two reaction paths have different values for Ea.

b) Both reactions have the same Ea. c) Ea for the reaction path for formation of I is

smaller than Ea for formation of A. d) Differences in activation energy have no role

in the rate changes. e) Both reaction rates should go up by a factor of

4 because T rose by 20 degrees.

Questions 57-60 A(g) + 2 B(g) + 3 C(g) → 4 D(g) + 5 E(g) Rate of formation of E = d[E]/dt = k [A]2 [B] 57) If one were to double the concentration of B, the rate of the reaction shown above would increase by a factor of a) ½ b) 1 c) 2 d) 4 e) 8 58) –d[B]/dt is equal to a) –d[A]/dt b) –d[C]/dt c) +d ½ [D]/dt d) +d 1/5 [E]/dt e) none of these 59) To decrease the rate constant k, one could a) increase [E] b) decrease [B] c) decrease the temperature d) increase the volume e) increase the pressure 60) If one were to reduce the volume of a container by 1/3, the rate of the reaction would increase by a factor of a) 3 b) 9 c) 16 d) 27 e) Reducing the volume of the container has no effect on the rate.

61) Sulfur trioxide gas dissociates into sulfur dioxide gas and oxygen gas at 1250 oC. In an experiment, 3.60 moles of sulfur trioxide were placed into an evacuated 3.0 liter flask. The concentration of sulfur dioxide gas measured at equilibrium was found to be 0.20 M. What is the equilibrium constant, Kc, for the reaction? a) 1.6 X10-4 b) 1.0 X10-3 c) 2.0 X10-3 d) 4.0 X10-3 e) 8.0 X10-3

62) The graph above shows the results of a study of the reaction of X with a large excess of Y to yield Z. The concentrations of X and Y were measured over a period of time. According to the results, which of the following can be concluded about the rate law for the reaction under the conditions studied?

a) It is zero order in [X]. b) It is first order in [X]. c) It is second order in [X]. d) It is first order in [Y]. e) The overall order of the reaction is 2.

63) N2(g) + O2(g) ↔ 2 NO In the formation of NO (shown above), what effect will the addition of a catalyst have on the equilibrium constant, Keq, for the reaction. Assume temperature and pressure remain constant.

a) A catalyst will increase Keq. b) A catalyst will decrease Keq. c) A catalyst will have no effect on Keq. d) A catalyst will first increase Keq, until the

reaction slows, then Keq will decrease. e) A catalyst will cause a dramatic increase in

Keq.

64) All of the following are true about a boiling liquid except a) ∆Go must be zero b) ∆So must be positive c) ∆Ho must be positive d) Keq must be 1.00 e) ∆So must equal ∆Ho

65) The equilibrium law for the following reaction is Na2CO3(s) + 2 HCl(g) ↔ H2O(l) + CO2(g)

a) [HCl]2 / [CO2] b) [Na2CO3] [HCl]2 / {[H2O][CO2]} c) [CO2] / [H2O] d) [H2O][CO2] / {[Na2CO3][HCl]2} e) [CO2] / [HCl]2

66) A reaction vessel contains N2(g), H2(g) and NH3(g) at a certain temperature. If the equation is represented in the following two ways, N2(g) + 3 H2(g) ↔ 2 NH3(g) ½ N2(g) + 3/2 H2(g) ↔ NH3(g), then the value of Kc

a) will be the same using either equation b) will be larger when equation (ii) is used c) will be larger when equation (i) is used d) for (i) is equal to square of Kc value for (ii) e) Both (b) and (d)

67) At 500 Kelvin, the Kp = 3.33 X10-17 for the equilibrium reaction N2(g) + 2 O2(g) → N2O4(g). What is the value for Kp for the reaction ½ N2O4(g) → ½ N2(g) + O2(g)? a) √3.33 X10-17 b) 3.19 X1019 c) (1/2) (3.33 X 10-17) d) √(1/3.33 X10-17) e) √3.33 X 10-17

68) PCl5(g) ↔ PCl3(g) + Cl2(g) When PCl5 is placed in an evacuated container at 250

oC, the above reaction takes place. The pressure in the container before the reaction takes place is 6 atm and is entirely due to the PCl5 gas. After the reaction comes to equilibrium, the partial pressure due to Cl2 is found to be 2 atm. What is the value of the equilibrium constant, Kp, for this reaction? a) 0.5 b) 1 c) 2 d) 5 e) 10

69) For which of the following reactions will the equilibrium constants Kc and Kp have the same value?

a) 2 N2O5(g) ↔ 2 NO2(g) + O2(g) b) 2 CO2(g) ↔ 2 CO(g) + O2(g) c) H2O(g) + CO(g) ↔ H2(g) + CO2(g) d) 3 O2(g) ↔ 2 O3(g) e) CO(g) + Cl2(g) ↔ COCl2(g)

70) N2(g) + 3 H2(g) ↔ 2 NH3(g) ∆Ho = -92 kJ Which of the following changes would cause a decrease in the equilibrium constant, Keq, for the reaction shown above?

a) decrease the temperature b) increase the temperature c) increase the pressure of the reaction vessel by

decreasing the volume d) the addition of nitrogen gas to the reaction

vessel e) the addition of argon gas to the reaction vessel

71) N2(g) + 3 H2(g) ↔ 2 NH3(g) If 2.00 mol H2, 1.00 mol N2 and 2.00 mol NH3 are placed into a 1.00 L evacuated flask. How will the reaction above proceed if Kc = 0.105 at this temperature?

a) The reaction must proceed from right to left. b) The reaction must proceed from left to right. c) The reaction is already at equilibrium, so no

net movement will occur. d) The reaction cannot establish equilibrium at

this temperature. e) The reaction will proceed from left to right

until 2.23 mol NH3 have been produced.

72) C(s) + CO2(g) ↔ 2 CO(g) In the reaction shown above, the value of Kp is 167.5 at a temperature of 1273 K. If the PCO2 at this temperature is 0.10 atm (at equilibrium), what will the PCO be? a) 0.1 atm b) 4.1 atm c) 12.9 atm d) 16.7 atm e) 53.0 atm

73) A reaction has an equilibrium constant of 3.8 X103. This constant will change if

a) a catalyst is added to the reaction mixture b) additional reactant is added c) the temperature is changed d) the pressure is decreased e) a precipitate is formed

74) Given the following reactions with their equilibrium constants, calculate the equilibrium constant for the overall reaction. Cd2+ + 4 CN- ↔ Cd(CN)4

2- Kf = 7.7 X10+16 CdCO3 ↔ Cd2+ + CO3

2- Ksp = 1.8 X10-14

CdCO3 + 4 CN- ↔ Cd(CN)42- + CO3

2- Koverall=? a) 1.4 X103 b) 4.3 X1030 c) 2.2 X106 d) 9.5 X102 e) 2.3 X10-31

75) 2 CO2(g) ↔ 2 CO(g) + O2(g) ∆Ho = -514 kJ The equilibrium constant will be the highest when the reaction above is carried out at a) low temperature and low pressure b) low temperature and high pressure c) high temperature and high pressure d) high temperature and low pressure e) any temperature or pressure; neither affect the

reaction

76) H2(g) + I2(g) ↔ 2 HI(g) At 450 oC the equilibrium constant, Kc, for the reaction shown above has a value of 50. Which of the following sets of initial conditions at 450 oC will cause the reaction above to produce more H2? I. [HI]= 5-molar, [H2]= 1-molar, [I2]= 1-molar II. [HI]= 10-molar, [H2]= 1-molar, [I2]= 1-molar III. [HI]= 10-molar, [H2]= 2-molar, [I2]= 2-molar a) I only b) II only c) I and II only d) II and III only e) I, II, and III

77) 4 FeS(s) + 7 O2(g) → 2 Fe2O3(s) + 4 SO2(g) ∆Ho = -2432 kJ mol-1 When the above reaction is at equilibrium for any given pressure, P, and temperature, T, which of the following will shift the position of equilibrium so that more product is formed?

a) increase the temperature of the system without changing the pressure

b) add an inert gas to increase the pressure of the system

c) add a catalyst specific for the forward reaction d) remove the Fe2O3 as it is formed e) remove the SO2 as it is formed

78) 2 CO(g) + O2(g) ↔ 2 CO2(g) ∆Ho = -514.2 kJ In the above reaction, which factors will cause the equilibrium to shift to the right? I. An increase in volume II. An increase in temperature III. Removal of CO2 a) I only b) II only c) III only d) I and II only e) II and III only

79) Consider the reaction 2SO2(g) + O2(g) ↔ 2 SO3(g) ∆H = -198.2 kJ Which of the following statements is NOT true?

a) Equilibrium will shift to the right if pressure is increased.

b) Equilibrium will shift to the right if temperature is decreased.

c) More SO3(g) can be produced by carrying the reaction at high pressure and low temperature.

d) Equilibrium will shift to the left if temperature is increased.

e) More SO3(g) can be produced if the reaction is carried out at low pressure and high temperature.

80) Which of the following does not necessarily produce more product in a chemical reaction?

a) increasing the temperature b) increasing the amount of reactants c) removing product as it is formed d) decreasing the volume of a reaction where ∆n

is negative e) increasing the volume of a reaction where ∆n

is positive

Questions 81-82 2 SO2(g) + O2(g) ↔ 2 SO3(g) ∆H = -197 kJ 81) Starting with a system at equilibrium, which of the following operations will not increase the amount of SO3(g)?

a) decreasing the temperature of the reaction vessel

b) decreasing the volume of the reaction vessel c) adding N2(g) to increase the pressure in the

reaction vessel d) adding O2(g) to the reaction vessel e) adding SO2(g) to the reaction vessel

82) SO2 and O2 are mixed in an insulated vessel and sealed so that there is no heat exchange with the surroundings. When the reaction comes to equilibrium, what best describes what has happened to the system? Use the information from the previous question if needed.

a) Total energy remains constant and the entropy increases.

b) Total energy increases and the entropy remains constant.

c) Total energy remains constant and the temperature of the system increases.

d) Total energy decreases and the temperature increases.

e) Total energy remains constant and the temperature decreases.

83) 2 NO(g) + O2(g) ↔ 2 NO2(g) ∆H < 0 Which of the following changes alone would cause a decrease in the value of Keq for the reaction represented above?

a) decreasing the temperature b) increasing the temperature c) decreasing the volume of the reaction vessel d) increasing the volume of the reaction vessel e) adding a catalyst

84) 2 SO3(g) ↔ 2 SO2(g) + O2(g) After the equilibrium represented above is established, some pure O2(g) is injected into the reaction vessel at constant temperature. After equilibrium is reestablished, which of the following has a lower value compared to its value at the original equilibrium?

a) Keq for the reaction b) The total pressure in the reaction vessel c) The amount of SO3(g) in the reaction vessel d) The amount of O2(g) in the reaction vessel e) The amount of SO2(g) in the reaction vessel

Written Questions:

1) 2 NO(g) + Cl2(g) → 2 NOCl(g) The following data were collected for the reaction above. All of the measurements were taken at a temperature of 263 K. Experiment Initial [NO] Initial [Cl2] Initial rate of disappearance (M) (M) of Cl2 (M/min) 1 0.15 0.15 0.60

2 0.15 0.30 1.2 3 0.30 0.15 2.4 4 0.25 0.25 ?

a) Write the expression for the rate law for the reaction above. b) Calculate the value of the rate constant for the above reaction and specify the units. c) What is the initial rate of appearance of NOCl in experiment 2? d) What is the initial rate of disappearance of Cl2 in experiment 4?

2) The reaction between NO and H2 is believed to occur in the following three-step process. NO + NO ↔ N2O2 (fast) N2O2 + H2 → N2O + H2O (slow) N2O + H2 → N2 + H2O (fast) a) Write a balanced equation for the overall reaction. b) Identify the intermediates in the reaction. Explain your reasoning. c) From the mechanism represented above, a student correctly deduces that the rate law for the

reaction is rate = k [NO]2[H2]. The student then concludes that (1) the reaction is third order and (2) the mechanism involves the simultaneous collision of two NO molecules and an H2 molecule. Are conclusions (1) and (2) correct? Explain.

d) Explain why an increase in temperature increases the rate constant, k, given the rate law in (c).

3) 2 A + B → C + D The following results were obtained when the reaction represented above was studied at 25 oC. Experiment Initial Initial Initial rate of formation [A] [B] of C (mol L-1 min-1) 1 0.25 0.75 4.3 X10-4 2 0.75 0.75 1.3 X10-3 3 1.50 1.50 5.3 X10-3 4 1.75 ? 8.0 X10-3 a) Determine the order of the reaction with respect to A and to B. Justify your answer. b) Write the rate law for the reaction. Calculate the value of the rate constant, specifying units. c) Determine the initial rate of change of [A] in Experiment 3. d) Determine the initial value of [B] in Experiment 4. e) Identify which of the reaction mechanisms represented below is consistent with the rate law

developed in part (b). Justify your choice Choices:

1. A + B → C + M fast M + A → D slow 2. B ↔ M fast equilibrium M + A → C + X slow A + X → D fast 3. A + B ↔ M fast equilibrium M + A → C + X slow X → D fast

4) Answer the following questions regarding the kinetics of chemical reactions. a) The diagram below shows the energy pathway for the reaction O3 + NO → NO2 + O2.

Clearly label the following directly on the diagram. (i) The activation energy (Ea) for the forward reaction. (ii) The enthalpy change (∆H) for the reaction.

b) The reaction 2 N2O5 → 4 NO2 + O2 is first order with respect to N2O5.

(i) Using the axes below, complete the graph that represents the change in [N2O5] over time as the reaction proceeds.

(ii) Describe how the graph in (i) could be used to find the reaction rate at a given time, t. (iii) Considering the rate law and the graph in (i), describe how the value of the rate constant, k, could be determined. (iv) If more N2O5 were added to the reaction mixture at constant temperature, what would be the effect on the rate constant, k? Explain.

c) Data for the chemical reaction 2 A → B + C were collected by measuring the concentrations of A at 10-minute intervals for 80 minutes. The following graphs were generated from the analysis of the data.

Use the information in the graphs above to answer the following. (i) Write the rate-law expression for the reaction. Justify your answer. (ii) Describe how to determine the value of the rate constant for the reaction.

5) (I) A2 + B2 → 2 AB

(II) X2 + Y2 → 2 XY Two reactions are represented above. The potential-energy diagram for reaction I is shown below. The potential energy of the reactants in reaction II is also indicated on the diagram. Reaction II is endothermic, and the activation energy of reaction I is greater than that of reaction II.

a) Complete the potential-energy diagram for reaction II on the graph above. b) For reaction I, predict how each of the following is affected as the temperature is increased

by 20 oC. Explain the basis for each prediction. (i) Rate of Reaction (ii) Heat of Reaction

c) For reaction II, the form of the rate law is rate = k [X2]m[Y2]n. Briefly describe an experiment that can be conducted in order to determine the values of m and n in the rate law for the reaction.

d) From the information given, determine which reaction initially proceeds at the faster rate under the same conditions of concentration and temperature. Justify your answer.

6) Hydrogen peroxide decomposes according to the following reaction: H2O2(aq) → 2 H2O(l) + O2(g) A graph of concentration versus time for a 1.00 M solution of H2O2 is shown below:

a) How can you determine the instantaneous rate

at 4,000 sec? b) What is the approximate value of half-life as it

is read from the graph? c) Assuming that decomposition of H2O2 is a

first-order reaction, how can you determine the rate constant (k) from the half-life?

7) C(s) + H2O(g) ↔ CO(g) + H2(g) ∆Ho = +131 kJ A rigid container holds a mixture of graphite pellets (C(s)), H2O(g), CO(g) and H2(g) at equilibrium. State whether the number of moles of CO(g) in the container will increase, decrease, or remain the same after each of the following disturbances is applied to the original mixture. For each case, assume that all other variables remain constant except for the given disturbance. Explain each answer with a short statement.

a) Additional H2(g) is added to the equilibrium mixture at constant volume. b) The temperature of the equilibrium mixture is increased at constant volume. c) The volume of the container is decreased at constant temperature. d) The graphite pellets are pulverized.

8) Le Chatlier’s principle is central to many qualitative and quantitative aspects of chemical equilibrium. Use the following equation to answer the questions.

N2(g) + 3 H2(g) ↔ 2 NH3(g) (∆Ho = -46.0 kJ mol-1) a) Briefly but completely summarize Le Chatlier’s principle. b) Which changes in concentration will increase the amount of NH3 produced? c) What changes in pressure will not increase the amount of NH3 produced? d) How should the temperature be changed to increase the amount of NH3 produced? e) Will a catalyst increase the amount of NH3 produced? f) What will happen if liquid water is present in the reaction system?

9) Hydrogen bromide decomposes according to the equation 2 HBr(g) ↔ H2(g) + Br2(g) a) Write the equilibrium law in terms of concentrations, Kc, and partial pressures, Kp. b) A 55.5 g sample of HBr (molar mass = 80.9) is transferred to an evacuated 22.0 L flask at 142 oC.

(1) What is the initial concentration in moles per liter of HBr? (2) What is the initial pressure of HBr in mm Hg?

c) When the system comes to equilibrium, 22.4 g of Br2 is found to be in the flask. Calculate the value of Kc and Kp. Be sure to indicate whether you are calculating Kc or Kp.

d) At another temperature, Kc = 17.7. 1.37 mol HBr(g), 2.31 mol H2(g), and 0.551 mol Br2(g) are introduced into an evacuated 15.0 L flask. Determine whether the reaction proceeds in the forward or in the reverse direction.

10) H2(g) + CO2(g) ↔ H2O(g) + CO(g) When H2(g) is mixed with CO2(g) at 2000 K, equilibrium is achieved according to the equation above. In one experiment, the following equilibrium concentrations were measured. [H2] = 0.20 mol/L [CO2] = 0.30 mol/L [H2O] = [CO] = 0.55 mol/L

a) What is the mole fraction of CO(g) in the equilibrium mixture? b) Using the equilibrium concentrations given above, calculate the value of Kc, the equilibrium

constant for the reaction. c) Determine Kp in terms of Kc for this system. d) When the system is cooled from 2000K to a lower temperature, 30.0 percent of the CO(g) is

converted back to CO2(g). Calculate the value of Kc at this lower temperature. e) In a different experiment, 0.50 mole of H2(g) is mixed with 0.50 mole of CO2(g) in a 3.0 liter reaction

vessel at 2000 K. Calculate the equilibrium concentration, in moles per liter, of CO(g) at this temperature.

11) The rate of decomposition of di-tert-butyl peroxide (DTBP) to acetone and ethane: (CH3)3COOC(CH3)3(g) → 2 (CH3)2CO(g) + CH3CH3(g) DTBP acetone ethane The progress of the reaction can be followed by measuring the total pressure in a closed system at constant temperature. The following data were collected in one experiment. Time, s Ptotal, mm Hg 0 700 0.50 X104 1420 1.00 X104 1770 1.50 X104 1940 2.00 X104 2020 2.50 X104 2062

a) What is the partial pressure PDTBP at the beginning, t = 0, of the experiment before any decomposition?

b) What is PDTBP at t = ∞? (Assume the reaction goes to completion.) c) What is Ptotal at t = ∞? d) What is PDTBP at t = 0.50 X104? e) What is the half-life for the decomposition reaction? f) What is the order for the reaction?

12) Consider the dimerization reaction of NO2: 2 NO2(g) ↔ N2O4(g)

a) Using Lewis structures, suggest why NO2 dimerizes and CO2 does not. b) What is the sign of the entropy change, ∆So, for this reaction in the forward direction? c) At 100 oC the value of Kc is approximately 3.3. What is the sign for ∆Ho? d) Can the reverse reaction be made spontaneous? If so how?

13) For the gaseous equilibrium represented below, it is observed that greater amounts of PCl3 and Cl2 are produced as the temperature is increased. PCl5(g) ↔ PCl3(g) + Cl2(g)

a) What is the sign of ∆So for the reaction? Explain. b) What change, if any, will occur in ∆Go for the reaction as the temperature is increased? Explain

your reasoning in terms of thermodynamic principles. c) If He gas is added to the original reaction mixture at constant volume and temperature to half the

original volume, what will happen to the partial pressure of Cl2? Explain. d) If the volume of the reaction mixture is decreased at constant temperature to half the original

volume, what will happen to the number of moles of Cl2 in the reaction vessel? Explain.

14) C2H2(g) + 2 H2(g) → C2H6(g) Information about the substances involved in the reaction represented above is summarized in the following tables. Substance So (J/mol•K) ∆Hf

o (kJ/mol) Bond Bond Energy (kJ/mol) C2H2(g) 200.9 226.7 C—C 347 H2(g) 130.7 0 C==C 611 C2H6(g) ----- -84.7 C—H 414 H—H 436

a) If the value of the standard entropy change, ∆So, for the reaction is –232.7 joules per mole• Kelvin, calculate the standard molar entropy, So, of C2H6 gas.

b) Calculate the value of the standard free-energy change, ∆Go, for the reaction. What does the sign of ∆Go indicate about the reaction above?

c) Calculate the value of the equilibrium constant, K for the reaction at 298 K. d) Calculate the value of the C==C (triple) bond energy in C2H2 in kilojoules per mole.