Reaction Rates (College Board AP® Chemistry): Exam Questions

Answer the following questions about the isomers fulminic acid and isocyanic acid.

Two possible Lewis electron-dot diagrams for fulminic acid , HCNO, are shown below.

Explain why the diagram on the left is the better representation for the bonding in fulminic acid. Justify your choice based on formal charges.

Fulminic acid can convert to isocyanic acid according to the equation below.

   HCNO(g)        HNCO(g)   

fulminic acid    isocyanic acid

Using the Lewis electron-dot diagrams of fulminic acid and isocyanic acid shown in the boxes above and the table of average bond enthalpies below, determine the value of ΔH° for the reaction of  HCNO (g) to form HNCO (g).

A student claims that ΔS°  for the reaction is close to zero. Explain why the student's claim is accurate.

Which species, fulminic acid  (HCNO) or isocyanic acid  (HNCO), is present in higher concentration at equilibrium at 298 K ? Justify your answer in terms of thermodynamic favorability and the equilibrium constant.

The ammonium salt of isocyanic acid is a product of the decomposition of urea, CO(NH₂)₂, represented  below.

CO(NH₂)₂ (aq) NH₄⁺ (aq) + OCN⁻ (aq)

A student studying the decomposition reaction runs the reaction at 90°C. The student collects data on the concentration of urea as a function of time, as shown by the data table and the graph below.

The student proposes that the rate law is rate = k[CO(NH₂)₂].

i) Explain how the data support the student's proposed rate law

ii) Using the proposed rate law and the student's results, determine the value of the rate constant, k.

Include units with your answer.

The student learns that the decomposition reaction was run in a solution with a pH of 13. Briefly describe an experiment, including the initial conditions that you would change and the data you would gather, to determine whether the rate of the reaction depends on the concentration of OH⁻ (aq).

Na₂C₃₇H₃₄N₂S₃O₉ + OCl⁻  →    products

            blue                                  colorless

Blue food coloring can be oxidized by household bleach (which contains OCl⁻ ) to form colorless products, as represented by the equation above. A student used a spectrophotometer set at a wavelength of 635 nm to study the absorbance of the food coloring over time during the bleaching process. In the study, bleach is present in large excess so that the concentration of OCl⁻ is essentially constant throughout the reaction. The student used data from the study to generate the graphs below. 

Based on the graphs above, what is the order of the reaction with respect to the blue food coloring?

The reaction is known to be first order with respect to bleach. In a second experiment, the student prepares solutions of food coloring and bleach with concentrations that differ from those used in the first experiment. When the solutions are combined, the student observes that the reaction mixture reaches an absorbance near zero too rapidly. In order to correct the problem, the student proposes the following three possible modifications to the experiment.

• Increasing the temperature

• Increasing the concentration of the food coloring

• Increasing the concentration of the bleach

Circle the one proposed modification above that could correct the problem, and explain how that modification increases the time for the reaction mixture to reach an absorbance near zero.

In another experiment, a student wishes to study the oxidation of red food coloring with bleach. How would  the student need to modify the original experimental procedure to determine the order of the reaction with respect to the red food coloring?

The half-life (t_½) of the catalyzed isomerization of cis-2-butene gas to produce trans-2-butene gas, represented above, was measured under various conditions, as shown in the table below.

The reaction is first order. Explain how the data in the table are consistent with a first-order reaction.

Calculate the rate constant, k, for the reaction at 350 K . Include appropriate units with your answer.

Is the initial rate of the reaction in trial 1 greater than, less than, or equal to the initial rate in trial 2? Justify your answer.

The half-life of the reaction in trial 4 is less than the half-life in trial 1. Explain why, in terms of activation energy.

2C₄H₆ (g) → C₈H₁₂ (g)

At high temperatures, the compound C₄H₆ (1,3-butadiene) reacts according to the equation above. The rate of the reaction was studied at 625 K in a rigid reaction vessel. Two different trials, each with a different starting concentration, were carried out. The data were plotted in three different ways, as shown below.

For trial 1, calculate the initial pressure, in atm, in the vessel at 625 K. Assume that initially all the gas present in the vessel is C₄H₆.

Use the data plotted in the graphs to determine the order of the reaction with respect to C₄H₆ .

The initial rate of the reaction in trial 1 is 0.0010 mol/(L·s). Calculate the rate constant, k, for the reaction at 625 K.

The following equation represents the decomposition of N₂O₅, for which the rate law is rate = k[N₂O₅]

2N₂O₅ (g) → 4NO₂ (g) + O₂ (g) 

A sample of pure N₂O₅ (g) is placed in an evacuated container and allowed to decompose at a constant temperature of 300 K. The concentration of N₂O₅ (g) in the container is measured over a period of time, and the measurements are recorded in the following table.

Determine the value of the rate constant, k, for the reaction. Include units in your answer.

The following mechanism is proposed for the decomposition of N₂O₅ (g).

Step 1: N₂O₅ (g) → NO₂ (g) + NO₃ (g)

Step 2: NO₂ (g) + NO₃ (g) → NO₂ (g) + NO (g) + O₂ (g)

Step 3: N₂O₅ (g) + NO (g) → 3NO₂ (g)

Identify which step of the proposed mechanism (1, 2, or 3) is the rate-determining step. Justify your answer in terms of the rate law given.

If this experiment was repeated at the same temperature but with twice the initial concentration of N₂O₅, would the value of k increase, decrease, or remain the same? Explain your reasoning.

A chemical reaction between gaseous hydrogen (H₂) and iodine (I₂) is represented by the following equation:

H₂ (g) + I₂ (g) → 2HI (g)

The reaction occurs in a closed container at constant temperature.

Explain how the frequency and energy of collisions affect the rate of reaction according to the collision model.

The temperature of the reaction is decreased.

Explain, using kinetic molecular theory, how this change affects the number of successful collisions per unit time.

The activation energy for this reaction is 167 kJ/mol.

Sketch a potential energy diagram for this reaction, labeling the reactants, products, activation energy (E_a), and enthalpy change (ΔH).

A powdered catalyst is added to the reaction mixture.

Explain how the catalyst affects both the activation energy and the overall rate of the reaction.

The initial concentration of hydrogen gas is 0.80 mol/L. After 4.0 seconds, the concentration decreases to 0.20 mol/L.

Calculate the average rate of reaction with respect to hydrogen over this time period. Include units.

A researcher investigates the following reaction:

A + B → C + D

The proposed mechanism is:

Step 1:

A + X ⇌ Y (fast, equilibrium)

Step 2:

Y + B → C + D (slow)

Justify why Step 2 is the rate-determining step in this mechanism.

Deduce the rate law for the overall reaction.

A student suggests that the reaction is second order with respect to reactant A. Evaluate this claim using your derived rate law.

A researcher studies the decomposition of ammonia on a platinum catalyst:

2NH₃ ​(g) ​N₂ ​(g) + 3H₂ ​(g)

Experiments show that increasing the surface area of the platinum greatly increases the reaction rate, while increasing the partial pressure of NH₃ has little effect.

Explain why increasing the surface area of platinum increases the reaction rate.

Justify why increasing the partial pressure of NH₃ has little effect on the rate.

Suggest a method to determine whether the reaction proceeds via a single step or multiple steps. Explain how the results would support the conclusion.

A chemist investigates the reaction:

A + B → Products

The initial rate of reaction is measured at varying concentrations of reactants, and the following data is recorded:

Determine the order of reaction with respect to A. Justify your answer using data from the table.

Determine the order of reaction with respect to B. Justify your answer using data from the table.

The chemist observes that at high concentrations of A, the reaction rate does not increase as expected. Suggest a plausible mechanistic explanation for this behavior and justify how it could limit the reaction rate.