Unit 5: Kinetics.

5.1 Reaction Rates.

Convert between the rate of appearance or disappearance of different species in a chemical reaction using ratios.

  1. Determine the rate of appearance of B in the reaction 2A → B, if the rate of disappearance of A is 0.35Ms-1

Determine the rate of the overall reaction from the rate of appearance or disappearance of a species.

  1. For the reaction: 2A → B. Find the overall rate of the reaction if the rate of appearance of B is 0.025 Ms-1.

Justify how a change in conditions (surface area, concentration/pressure, temperature, presence of a catalyst) will affect the rate of the reaction.

  1. Explain why increasing temperature generally increases the rate of reaction.

5.2 Introduction to Rate Law.

Write a rate law for a reaction using initial rate data from multiple trials.

  1. Determine the rate law for the reaction: X + Y → XY Using the lab data in the table below.
    Trial Initial [X] Initial [Y] Initial Rate of Appearance of XY Ms-1
    1 0.010 0.010 0.0050
    2 0.010 0.020 0.010
    3 0.020 0.010 0.020

Determine the overall order of the reaction by adding the powers of the reactant concentrations in the rate law.

  1. What is the overall order of the reaction with an experimentally determined rate law of: rate = k[A]2[B]?

Calculate the value of the rate constant (k) and express it with the correct units.

  1. The data below was used to determine the rate law of: rate = k[X][Y]. Determine the value of the rate constant (k) and its units.
    Trial Initial [X] Initial [Y] Initial Rate of Appearance of XY Ms-1
    1 0.01 0.01 0.005
    2 0.01 0.02 0.01
    3 0.02 0.01 0.01

Understand that the rate constant (k) is dependent on temperature and therefore specific to a temperature.

  1. Predict whether the following changes would affect the value of the rate constant (k)?
    1. Increasing the temperature of the reactants from 300K to 400K.
    2. Increase the concentration of the reactants from 1.0M to 2.0M.
    3. Reducing the volume of a reaction container for gas phase reactants.

5.3 Concentration Changes Over Time.

Determine the order of a reaction by examining graphs of [A], ln[A] and 1/[A] vs time.

  1. Identify the order of the reaction with respect to [X] based on the plots below.

    2. plots of concentration vs time.

Use the integrated rate law corresponding to the order of a reaction to determine the value of k with units.

  1. A student is studying a first order chemical reaction and collects the data shown in the table below. Determine the rate constant (k) using the appropriate integrated rate law.
    Time (sec) Concentration[X]
    0.0 1.0
    60.0 0.40

Use the half-life equation to solve for the rate constant (k) or the half-life (t1/2)

  1. Calculate the half-life of a reaction with k = 0.0030 s-1.

Apply the fact: only first order chemical reactions have consistent half-lives.

  1. Without graphing or using an integrated rate law, justify how the data shown below gives evidence for the reaction being first order.
    Time (sec) [X] (M)
    0 1.0
    10 0.68
    20 0.50
    30 0.34
    40 0.25
    50 0.16
    60 0.125

5.4 Elementary Reactions.

Use coefficients of elementary steps to determine the rate law.

  1. Write the rate law for the elementary step A + B → C.

Justify if a theorized elementary step is likely or not based on the number of particles colliding (its molecularity).

  1. Explain why the elementary step 2A + 3B → A2B3 is not likely part of a true reaction mechanism.

5.5 Collision Model.

Apply the fact: breaking and forming of bonds requires a collision that has sufficient energy and correct orientation.

  1. Explain why collisions between particles with very large kinetic energies as sometimes not successful.

Use the collision model to justify how reactant concentration, temperature, and surface area affects the frequency of successful collisions.

  1. Use particle level reasoning to explain how the increasing surface area of a solid reactant increases the rate of a chemical reaction.
  2. Use particle level reasoning to explain why increasing the concentration of [HCl] in the reaction between HCl(aq) and CaCO3(s) will increase the rate of the chemical reaction.

Re-draw a Maxwell Boltzmann Distribution after a temperature change has occurred.

  1. The curve below represents a Boltzmann distribution at 300K, on the same graph sketch the distribution at 400K.
    Boltzann Distribution at 300K

Apply knowledge of how Maxwell-Boltzmann distributions connect to activation energy and reaction rate

  1. In the image below shade the region under the curve that corresponds to particles with sufficient kinetic energy to overcome the activation energy
    Boltzann Distribution at 300K

5.6 Reaction Energy Profile.

Label the areas on a reaction energy profile: reactants, transition state, products, activation energy, and change in enthalpy (ΔH).

  1. Label the reactants, transition state, products, activation energy, and change in enthalpy (ΔH) on the energy profile below.
    Exothermic Energy Diagram

Determine the values of activation energy and change in enthalpy (ΔH) for a reaction using its energy profile.

  1. Determine the value of the activation energy and the change in enthalpy (ΔH) for the reaction depicted by the energy profile below.
    Exothermic Energy Diagram

Justify the relationship between temperature, activation energy, and rate of a reaction. (This relationship is explained with the Arrhenius equation but questions requiring calculating the formula are excluded from the exam.

  1. A student states: chemical reactions have a higher rate of reaction at higher temperatures because increasing the temperature decreases the activation energy of the reaction. Do you agree or disagree with the student? Justify your answer.

5.7 Introduction to Reaction Mechanisms.

Identify the: reactants, products, intermediates, catalysts from a reaction mechanism.

  1. Identify each chemical species in the reaction mechanism as either a reactant, product, intermediate, or catalyst.
    Step Speed
    NO + O3 → NO2 + O2 slow
    O + NO2 → NO + O2 fast

5.8 Reaction Mechanism and Rate Law.

Use the slow step of a mechanism to write the rate law for the overall reaction.

  1. Use the reaction mechanism below to determine the overall rate law of the reaction.
    Step Speed
    NO2+ Cl2 → NO2Cl + Cl slow
    Cl + NO2 → NO2Cl fast

Justify if a provided reaction mechanism is valid by verifying the sum of the elementary steps and that the rate law matches the rate determining step.

  1. The reaction: O3 + O → 2O2 has an experimentally determined rate law: rate=k[NO][O3]. Justify whether the proposed mechanism below is a valid mechanism.
    Step Speed
    NO2 + Cl2 → NO2Cl + Cl slow
    Cl + NO2 → NO2Cl fast

5.9 Pre-Equilibrium Approximation.

Write a rate law for the overall reaction when the slow step is not the first step by substituting species from the fast equilibrium step.

  1. Write the overall rate law based on the reaction mechanism below.
    Step Speed
    A + 2C ⇋ AC2 Fast
    A + AC2 → 2AC Slow

5.10 Multistep Reaction Energy Profile.

Draw conclusions about the relative rate of the elementary steps of a reaction mechanism by looking at the individual activation energies shown on a multistep energy profile. conditions.

  1. Sketch an energy profile for a 3 step mechanism where the first step is the rate determining step.

5.11 Catalysis.

Modify an energy profile to show the effect of a catalyst.

  1. Sketch the energy profile for the reaction depicted below after a catalyst is added.
    Exothermic Energy Diagram

Explain how a catalyst increases reaction rate by either providing an alternative reaction mechanism where the rate determining step has a lower activation energy OR improving the probability of collisions being effective.

  1. In the lungs the reaction of CO2 + H2O → H2CO3 is catalysed by the enzyme carbonic anhydrase. Suggest what carbonic anhydrase does at the particle level that increases the rate of reaction.
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