AP Chemistry Kinetics Quiz MyAP

H3AsO4 + 3 I- + 2 H3O+ → H3AsO3 + I3- + H2O The oxidation of iodide ions by arsenic acid in acidic aqueous solution occurs according to the stoichiometry shown above. The experimental rate law of the reaction is: Rate = k[H3AsO4] [I-] [H3O+] According to the rate law for the reaction, an increase in the concentration of hydronium ion has what effect on this reaction? A The rate of reaction increases. B The rate of reaction decreases. C The value of the equilibrium constant increases. D The value of the equilibrium constant decreases. E Neither the rate nor the value of the equilibrium constant is changed.

A

The proposed rate-determining step for a reaction is 2 NO2(g)→NO3(g)+NO(g). The graph above shows the distribution of energies for NO2(g) molecules at two temperatures. Based on the graph, which of the following statements best explains why the rates of disappearance of NO2(g) are different at temperature 2 and temperature 1 ? A NO2(g) is consumed at a faster rate at temperature 2 because more molecules possess energies at or above the minimum energy required for a collision to lead to a reaction compared to temperature 1. B NO2(g) is consumed at a faster rate at temperature 2 because the molecules have a wider range of energies allowing for a better orientation during a collision compared to temperature 1. C Fewer NO2(g) molecules have a relatively high energy at temperature 1, which favors collisions between molecules rather than between the molecules and the container, leading to a faster rate of disappearance compared to temperature 2. D More NO2(g) molecules have a relatively low energy at temperature 1, which increases the number of effective collisions taking place and the rate of disappearance compared to temperature 2.

A

Which graph best represents the changes in concentration of O2(g), and why? A Graph 1, because the rate of O2 consumption is half the rate at which NO is consumed; two molecules of NO react for each molecule of O2 that reacts. B Graph 1, because O2 molecules are consumed at a slower rate at the beginning of the reaction when there are not as many molecules of NO2 produced. C Graph 2, because there is a large excess of O2 molecules and its concentration will not change drastically over time. D Graph 2, because a collision between three molecules to form a product has a low probability and the concentration of O2 will remain relatively constant.

A

2 N2O5(g) → 4 NO2(g) + O2(g) A sample of N2O5 was placed in an evacuated container, and the reaction represented above occurred. The value of PN2O5 , the partial pressure of N2O5(g), was measured during the reaction and recorded in the table below. Which of the following correctly describes the reaction? A The decomposition of N2O5 is a zero-order reaction. B The decomposition of N2O5 is a first-order reaction. C The decomposition of N2O5 is a second-order reaction. D The overall reaction order is 3.

B

If 87.5 percent of a sample of pure 131I decays in 24 days, what is the half-life of 131I? A 6 days B 8 days C 12 days D 14 days E 21 days

B

NO(g) + NO3(g) → 2 NO2(g) rate = k[NO][NO3] The reaction represented above occurs in a single step that involves the collision between a particle of NO and a particle of NO3. A scientist correctly calculates the rate of collisions between NO and NO3 that have sufficient energy to overcome the activation energy. The observed reaction rate is only a small fraction of the calculated collision rate. Which of the following best explains the discrepancy? A The energy of collisions between two reactant particles is frequently absorbed by collision with a third particle. B The two reactant particles must collide with a particular orientation in order to react. C The activation energy for a reaction is dependent on the concentrations of the reactant particles. D The activation energy for a reaction is dependent on the temperature.

B

S2O82−(aq)+3I−(aq)→2SO42−(aq)+I3−(aq) In aqueous solution, the reaction represented by the balanced equation shown above has the experimentally determined rate law: rate=k[S2O82−][I−]. If the concentration of [S2O82−] is doubled while keeping [I−] constant, which of the following experimental results is predicted based on the rate law, and why? A The rate of reaction will remain the same, because k will decrease by half. B The rate of reaction will double, because the rate is directly proportional to [S2O82−]. C The rate of reaction will increase by a factor of four, because two moles of SO42− are produced for each mole of S2O82− consumed. D The rate of reaction will increase by a factor of four, because the reaction is second order overall.

B

Which of the following statements best explains why an increase in temperature of 5-10 Celsius degrees can substantially increase the rate of a chemical reaction? A The activation energy for the reaction is lowered. B The number of effective collisions between reactant particles is increased. C The rate of the reverse reaction is increased. D ∆H for the reaction is lowered. E ∆G for the reaction becomes mor

B

Which of the following statements is a correct interpretation of the data regarding how the order of the reaction can be determined? A The reaction must be first order because there is only one reactant species. B The reaction is first order if the plot of ln [H2O2] versus time is a straight line. C The reaction is first order if the plot of 1/[H2O2] versus time is a straight line. D The reaction is second order because 2 is the coefficient of H2O2 in the chemical equation.

B

Which of the following will most likely increase the rate of the reaction represented above? A Decreasing the temperature of the reaction system B Adding a heterogeneous catalyst to the reaction system C Increasing the volume of the reaction vessel using a piston D Removing some H2(g) from the reaction system

B

Cl−(aq) + ClO−(aq) + 2 H+(aq) → Cl2(g) + H2O(l) What effect will increasing [H+] at constant temperature have on the reaction represented above? A The activation energy of the reaction will increase. B The activation energy of the reaction will decrease. C The frequency of collisions between H+(aq) ions and ClO−(aq) ions will increase. D The value of the rate constant will increase.

C

Cu(s) + 4 HNO3(aq) → Cu(NO3)2(aq) + 2 NO2(g) + 2 H2O(l) Each student in a class placed a 2.00 g sample of a mixture of Cu and Al in a beaker and placed the beaker in a fume hood. The students slowly poured 15.0 mL of 15.8 M HNO3(aq) into their beakers. The reaction between the copper in the mixture and the HNO3(aq) is represented by the equation above. The students observed that a brown gas was released from the beakers and that the solutions turned blue, indicating the formation of Cu2+(aq). The solutions were then diluted with distilled water to known volumes. In one student's experiment the reaction proceeded at a much slower rate than it did in the other students' experiments. Which of the following could explain the slower reaction rate? A In the student's sample the metal pieces were much smaller than those in the other students' samples. B The student heated the reaction mixture as the HNO3(aq) was added. C The student used a 1.5 M solution of HNO3(aq) instead of a 15.8 M solution of HNO3(aq). D The student used a 3.00 g sample of the mixture instead of the 2.00 g sample that was used by the other students.

C

Gaseous cyclobutene undergoes a first-order reaction to form gaseous butadiene. At a particular temperature, the partial pressure of cyclobutene in the reaction vessel drops to one-eighth its original value in 124 seconds. What is the half-life for this reaction at this temperature? A 15.5 sec B 31.0 sec C 41.3 sec D 62.0 sec E 124 sec

C

The gas-phase reaction A2(g)+B2(g)→2 AB(g) is assumed to occur in a single step. Two experiments were done at the same temperature inside rigid containers. The initial partial pressures of A2 and B2 used in experiment 1 were twice the initial pressures used in experiment 2. Which statement provides the best comparison of the initial rate of formation of AB in experiments 1 and 2 ? A The initial rate of formation of AB is the same in both experiments because they were done at the same temperature and the frequency and energy of the collisions between A2 and B2would have been about the same. B The initial rate of formation of AB is slower in experiment 1 than in with experiment 2 because at the same temperature, a higher pressure would reduce the volume available for A2 and B2molecules to achieve the proper orientation for a successful collision. C The initial rate of formation of AB is faster in experiment 1 than in experiment 2 because at a higher pressure the collisions between A2 and B2 molecules would have been more frequent, increasing the probability of a successful collision. D The initial rate of formation of AB is faster in experiment 1 than in experiment 2 because at a higher pressure a larger fraction of the A2 and B2 molecules would have the minimum energy required to overcome the activation energy barrier.

C

The initial rate of formation of CO2(g) from the chemical reaction represented by the equation above was studied in two separate experiments. The table above provides the experimental conditions used. If both experiments are carried out with finely powdered samples of the solid and 50.0mL of HCl(aq), which experiment, if any, will have the faster initial rate of formation of CO2(g) and why? A The rate of formation of CO2(g) will be the same because the mass of CaCO3(s) and the volume of HCl(aq) used will be the same in both experiments. B The rate of formation of CO2(g) will be the same because the surface area of the solid and the average kinetic energy of the particles will be the same in both experiments. C CO2(g) will be formed at a faster rate in experiment 2 because more H+ particles can react per unit time. D CO2(g) will be formed at a faster rate in experiment 1 because the proportion of CaCO3(s)particles to H+ particles will be greater.

C

Two samples of Mg(s) of equal mass were placed in equal amounts of HCI(aq) contained in two separate reaction vessels. Particle representations of the mixing of Mg(s) and HCI(aq) in the two reaction vessels are shown in Figure 1 and Figure 2 above. Water molecules are not included in the particle representations. Which of the reactions will initially proceed faster, and why? A The reaction in Figure 1, because the atoms of Mg are more concentrated than those in Figure 2 B The reaction in Figure 1, because the Mg(s) in Figure 1 has a larger mass than the Mg(s) in Figure 2 C The reaction in Figure 2, because more Mg atoms are exposed to HCI(aq) in Figure 2 than in Figure 1 D The reaction in Figure 2, because the Mg(s) in Figure 2 has less surface area than the Mg(s) in Figure 1

C

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.

C

Approximately how long did it take for 75 percent of the initial amount of C25H30N3+ (aq) to react? A 75 s B 225 s C 300 s D 600 s

C. 300 s

Relatively slow rates of chemical reaction are associated with which of the following? A The presence of a catalyst B High temperature C High concentration of reactants D Strong bonds in reactant molecules E Low activation energy

D

The initial rates of the reaction represented by the equation shown above were measured for different initial concentrations of NO(g) and Cl2(g). Based on the data given in the table above, which of the following is the rate law expression for the reaction, and why? A Rate=k[NO]2, because the initial rate quadrupled when [NO] was doubled but remained constant when [Cl2] was doubled. B Rate=k[NO][Cl2], because the initial rate doubled when either [NO] or [Cl2] was doubled. C Rate=k[NO][Cl2]2, because the initial rate doubled when [NO] was doubled and quadrupled when [Cl2] was doubled. D Rate=k[NO]2[Cl2], because the initial rate quadrupled when [NO] was doubled and doubled when [Cl2] was doubled.

D

The table above shows the results from a rate study of the reaction X + Y → Z. Starting with known concentrations of X and Y in experiment 1, the rate of formation of Z was measured. If the reaction was first order with respect to X and second order with respect to Y, the initial rate of formation of Z in experiment 2 would be A R4 B R2 C R D 2R E 4R

D

A kinetics experiment is set up to collect the gas that is generated when a sample of chalk, consisting primarily of solid CaCO3, is added to a solution of ethanoic acid, CH3COOH. The rate of reaction between CaCO3 and CH3COOH is determined by measuring the volume of gas generated at 25oC and 1 atm as a function of time. Which of the following experimental conditions is most likely to increase the rate of gas production? A Decreasing the volume of ethanoic acid solution used in the experiment B Decreasing the concentration of the ethanoic acid solution used in the experiment C Decreasing the temperature at which the experiment is performed D Decreasing the particle size of the CaCO3 by grinding it into a fine powder

D. Decreasing the particle size of the CaCO3 by grinding it into a fine powder

Factors that affect the rate of a chemical reaction include which of the following? I. Frequency of collisions of reactant particles II. Kinetic energy of collisions of reactant particles III. Orientation of reactant particles during collisions A II only B I and II only C I and III only D II and III only E I, II, and III

E