(Paper) Chemistry MCQ: Chemical kinetics

Paper : Chemistry : Chemistry MCQ: Chemical kinetics

  1. Consider the following reaction in aqueous solution,

5Br-(aq) + BrO3-(aq) + 6H+(aq) --> 3Br2(aq) + 3H2O(l)

If the rate of appearance of Br2 at a particular moment during the reaction is 0.025 M s-1, what is the rate of disappearance (in M s-1) of Br- at that moment?

  1. Consider the following reaction at 25oC,

(CH3)3COH(l) + HCl(aq) --> (CH3)3CCl(l) + H2O(l)

The experimentally determined rate law for this reaction indicates that the reaction is first-order in (CH3)3COH and that the reaction is first-order overall. Which of the following would produce an increase in the rate of this reaction?

    1. increasing the concentration of (CH3)3COH
    2. increasing the concentration of HCl
    3. decreasing the concentration of HCl
    4. decreasing the concentration of (CH3)3CCl
    5. It is impossible to tell.

  1. A certain first-order reaction has a rate constant, k, equal to 2.1 x 10-5 s-1 at 355 K. If the activation energy for this reaction is 135 kJ/mol, calculate the value of the rate constant (in s-1) at 550 K.

  2. Which of the following influences the rate of a chemical reaction performed in solution?
    1. temperature
    2. activation energy
    3. presence of a catalyst
    4. concentrations of reactants
    5. All of the above influence the rate.

  3. Laughing gas, N2O, can be prepared (ha, ha!) from H2 and NO:

H2(g) + 2 NO(g) --> N2O(g) + H2O(g)

A study of initial concentration (ha, ha!) versus initial rate at a certain temperature yields the following data for this reaction (ha, ha!):

[H2], M

[NO], M

initial rate, M s-1

0.1000

0.5000

2.560 x 10-6

0.2000

0.3000

1.843 x 10-6

0.1000

0.3000

9.216 x 10-7

0.2000

0.6000

7.373 x 10-6

Which of the following is the correct rate law for this reaction (ha, ha!)?

    1. Rate = k[H2][NO]2
    2. Rate = k[H2][NO]
    3. Rate = k[NO]2
    4. Rate = k[H2]2
    5. Rate = k

  1. Iodine-131, a radioactive isotope of iodine, is used medicinally as a radiotracer for the diagnosis and treatment of illnesses associated with the thyroid gland. The half-life of iodine-131 is 7.0 x 105 seconds. If a patient is given 0.45 g of iodine-131, calculate how long it would take (in seconds) for 90.0% of the iodine-131 to decay. Recall: radioactive decay is a first-order process.

  2. Consider a reaction which is first-order in A and first-order in B,

  3. The complex ion, [Cr(NH3)5Cl]2+, reacts with OH- ion in aqueous solution,

[Cr(NH3)5Cl]2+(aq) + OH-(aq) --> [Cr(NH3)5(OH)]2+(aq) + Cl-(aq)

The following data were obtained for this reaction at 25oC,

time, min

[Cr(NH3)5Cl]2+, M

0

1.00

6

0.657

12

0.432

18

0.284

24

0.186

30

0.122

36

0.0805

The order of the reaction with respect to the [Cr(NH3)5Cl]2+ ion is:

    1. zero order
    2. first order
    3. second order
    4. third order
    5. fourth order

  1. A student determined the value of the rate constant, k, for a chemical reaction at several different temperatures. Which of the following graphs of the student's data would give a straight line?
    1. k versus T
    2. k versus (1/T)
    3. ln k versus (1/T)
    4. ln k versus T
    5. ln k versus Ea

  2. In the experiment, "How Can Spectrophotometric Methods Be Used to Determine the Order of a Chemical Reaction", it is necessary to remove invalid data points towards the end of the reaction. Which of the following statements best explains why this is necessary?
    1. The Spectronic 20 becomes unstable towards the end of the reaction.
    2. Towards the end of the reaction, the temperature of the solution is significantly different than the initial temperature of the solution.
    3. Towards the end of the reaction, the concentrations of the reactants are so high that it is difficult to measure them accurately.
    4. Towards the end of the reaction, the concentrations of the products are sufficiently high that the reverse reaction competes with the forward reaction.
    5. None of these.

The next two questions are about this reaction:

2N2O5 (g) <==> 4NO2 (g) + O2 (g)

  1. The rate law for the above reaction is:
    1. rate = k [N2O5]2
    2. rate = [N2O5]2
    3. rate = k [N2O5]2 / [NO2]4 [O2]1
    4. rate = k [N2O5]x
    5. rate = [N2O5]x

  2. If the instantaneous rate of appearance of NO2 (g) is 0.0400 M/s at some moment in time, what is the rate of disappearance of N2O5 (g) in M/s ?

  3. The rate laws for certain enzyme-activated reactions in your body have a specific rate constant k, with units of M/s. What is the overall order of these reactions?
    1. 0
    2. 1
    3. 2
    4. 3
    5. Cannot be determined.

The next two questions are about this reaction:

2NO(g) + Cl2(g) <==> 2NOCl(g)

  1. The rate law for the above reaction has been determined to be

rate = k[NO][Cl2].

What is the overall order of the reaction?

    1. 0
    2. 1
    3. 2
    4. 3
    5. Cannot be determined.

  1. A mechanism involving the following steps has been proposed for the above reaction:

(1)

NO(g) + Cl2(g) --> NOCl2(g)

(2)

NOCl2(g) + NO(g) --> 2NOCl(g)

  1. Based on the rate law given in the preceding problem, which step is the rate-limiting step?
    1. Step (1)
    2. Step (2)
    3. Both Steps (1) & (2)
    4. Either Steps (1) or (2)

  2. Consider the following reaction in aqueous solution,

I- (aq) + OCl- (aq) -> IO- (aq) + Cl- (aq)

and the following initial concentration and initial rate data for this reaction:

[I-], M

[OCl-], M

initial rate, M s-1

0.1000

0.0500

3.05 x 10-4

0.2000

0.0500

6.10 x 10-4

0.3000

0.0100

1.83 x 10-4

0.3000

0.0200

3.66 x 10-4

Which of the following is the correct rate law for this reaction?

    1. Rate = k[I-]
    2. Rate = k[OCl-]
    3. Rate = k[I-]2
    4. Rate = k[I-][OCl-]
    5. Rate = k[I-]2[OCl-]

  1. Which of the following statements best describes how the "Method of Initial Rates" is used to measure the initial rate of an equilibrium reaction?
    1. The rate of the reaction is measured when the reaction is very close to equilibrium.
    2. The rate of the reaction is measured immediately after the reaction is started.
    3. The rate of the reaction is measured when the reaction is about one-half complete.
    4. The rate of the reaction is measured after five half-lives.
    5. The rate of an equilibrium reaction cannot be measured using this method.

  2. Which of the following are the correct units for the rate constant, k, for a zero-order reaction?
    1. M s-1
    2. M-1 s-1
    3. M-2 s-1
    4. M-3 s-1
    5. M

  3. Which of the following statements is TRUE?
    1. The existence of certain intermediates in a reaction mechanism can sometimes be proven because intermediates can sometimes be trapped and identified.
    2. Intermediates in a reaction mechanism cannot be isolated because they do not have finite lifetimes.
    3. Reaction mechanisms cannot have any more than one intermediate.
    4. Intermediates in a reaction mechanism appear in the overall, balanced equation for the reaction.
    5. None of the above statements is TRUE.

  4. Which of the following statements best describes how a catalyst works?
    1. A catalyst changes the potential energies of the reactants and products.
    2. A catalyst decreases the temperature of the reaction which leads to a faster rate.
    3. A catalyst lowers the activation energy for the reaction by providing a different reaction mechanism.
    4. A catalyst destroys some of the reactants, which lowers the concentration of the reactants.
    5. A catalyst raises the activation energy for the reaction which produces a faster rate.

  5. In terms of the "Collision Theory of Chemical Kinetics", the rate of a chemical reaction is proportional to:
    1. the change in free energy per second.
    2. the change in temperature per second.
    3. the number of collisions per second.
    4. the number of product molecules.
    5. none of the above.

  6. Nitrogen monoxide, NO, reacts with hydrogen, H2, according to the following equation:

2 NO (g) + 2 H2 (g) -> N2 (g) + 2 H2O (g)

If the mechanism for this reaction were,

2NO(g) + H2(g) -> N2(g) + H2O2(g)

slow

H2O2(g) + H2(g) -> 2H2O(g)

fast

which of the following rate laws would we expect to obtain experimentally?

    1. Rate = k[H2O2][H2]
    2. Rate = k[NO]2[H2]
    3. Rate = k[NO]2[H2]2
    4. Rate = k[NO][H2]
    5. Rate = k[N2][H2O2]

  1. Consider the following reaction in aqueous solution,

5 Br- (aq) + BrO3- (aq) + 6 H+ (aq) -> 3 Br2 (aq) + 3 H2O (l)

If the rate of disappearance of Br- (aq) at a particular moment during the reaction is 3.5 x 10-4 M s-1, what is the rate of appearance of Br2 (aq) at that moment?

  1. Consider the following gas-phase reaction,

2 HI (g) -> H2 (g) + I2 (g)

and the following experimental data obtained at 555 K,

[HI], M

rate, M s-1

0.0500

8.80 x 10-10

0.1000

3.52 x 10-9

0.1500

7.92 x 10-9

What is the order of the reaction with respect to HI (g)?

  1. Radioactive phosphorus is used in the study of biochemical reaction mechanisms. The isotope phosphorus-33 decays by first-order kinetics with a half-life of 14.3 days. If a chemist initially has a 7.5 M solution of pure phosphorus-33, calculate the concentration (in M) of phosphorus-33 in the solution after 2.4 days.

  2. Hydrogen iodide, HI, decomposes in the gas phase to produce hydrogen, H2, and iodine, I2:

2 HI (g) -> H2 (g) + I2 (g)

The value of the rate constant, k, for this reaction was measured at several different temperatures and the data are shown below:

temperature, K

k, M-1 s-1

555

6.23 x 10-7

575

2.42 x 10-6

645

1.44 x 10-4

700

2.01 x 10-3

Calculate the value of the activation energy (in kJ/mol) for this reaction.

  1. Listed below are some statements that pertain to chemical kinetics. For each statement, first decide whether the statement is TRUE or FALSE. If the statement is FALSE, briefly explain why the statement is FALSE. If the statement is TRUE, you do not need to provide any additional explanation.
    1. Rate laws for chemical reactions can be determined from the stoichiometry of the overall balanced equation.
    2. The rate of a chemical reaction that occurs in solution depends on the concentration, the temperature and the viscosity of the solvent.
    3. The "Method of Initial Rates" cannot be used for equilibrium reactions.
    4. Experimentally, transition states can sometimes be trapped or isolated whereas intermediates cannot be either trapped or isolated.

  2. Consider the following reaction,

2 ClO2 (aq) + 2 OH- (aq) -> ClO3- (aq) + ClO2- (aq) + H2O (l)

Write an expression that describes the relationship between the rates of disappearance of ClO2 and OH- and the rates of appearance of ClO3-, ClO2- and H2O.

  1. A Chemistry 116 student was charged with the task of determining the activation energy, Ea, for a particular first-order reaction. The student measured the value of the rate constant for this reaction at several temperatures and obtained the following data:

k, s-1

T, K

3.94 x 10-4

384

1.17 x 10-3

397

5.26 x 10-2

447

4.63 x 10-1

481

  1. Nitrogen monoxide, NO, reacts with ozone, O3, to produce nitrogen dioxide, NO2, and oxygen, O2:

NO (g) + O3 (g) -> NO2 (g) + O2 (g)

The experimentally determined rate law for this reaction is: Rate = k[O3]. Consider the following proposed mechanism for this reaction:

Step 1:

O3 (g) -> O2 (g) + O (g)

fast

Step 2:

NO (g) + O (g) -> NO2 (g)

slow

This proposed mechanism is NOT an acceptable possibility for this reaction. Briefly explain why this mechanism is not acceptable. ALSO, identify the rate-determining step and any intermediates in this proposed mechanism.

  1. Using the information in the previous question, and your knowledge about reaction mechanisms, write an acceptable mechanism for the reaction described in the previous question.

  2. In many kinetics studies, the rate of the reaction is determined almost immediately after the reactants are mixed. The BEST reason for doing this is:
    1. The reaction proceeds faster as time increases.
    2. Intermediates are easier to detect.
    3. The reaction proceeds slower as time increases.
    4. Reverse reactions are avoided.
    5. The concentrations of the reactions hasn't changed much.

USE THE FOLLOWING DATA FOR THE NEXT THREE (3) QUESTIONS.

Nitric oxide gas reacts with chlorine gas according to the equation,

2 NO + Cl2 -> 2 NOCl

The following data were obtained for this reaction:

initial [NO]

initial [Cl2]

initial rate

0.50 M

0.50 M

1.19 mol/L hr

1.00

0.50

4.79

1.00

1.00

9.59

1.50

1.50

32.27