Chapter 18:  Qualitative Chemical Kinetics

 

 

Section 18-1: Collision Theory and Factors That Increase Chemical Reaction Rates

Section 18-2: Reaction Energy Profiles (Reaction Progress Diagrams)

Section 18-3: Reaction Mechanisms

Section 18-4: Catalysis

Chapter 18 Practice Exercises and Review Quizzes

 

 

 

 

Section 18-1:  Collision Theory and Factors That Increase Chemical Reaction Rates

Chemical kinetics focuses on the rates of chemical reactions.  Collision theory suggests that chemical reactions occur as a result of collisions between reactant molecules.  Therefore, the rates of chemical reactions will generally increase when the number of collisions between reactant molecules is increased by increasing the concentration of aqueous or gaseous reactants or increasing the surface area of solid reactants.

 

However, for a collision between reactant molecules to actually result in a reaction, the following requirements must also be met:

 

A. The reactant molecules must collide with the proper orientation or at the proper angle.

 

B. The total kinetic energy (energy of motion) of the reactant molecules must be equal to or greater than the required minimum energy, which is known as the activation energy and will be different for each different chemical reaction.  Therefore, the rates of chemical reactions will generally increase when the percentage of reactant molecules that possess the required activation energy is increased by increasing the temperature.

 

 

Section 18-2:  Reaction Energy Profiles (Reaction Progress Diagrams)

A reaction energy profile (or reaction progress diagram) traces the changes in energy that occur as reactants are transformed into products.  Reactant molecules that collide with the required activation energy are able to form a higher-energy temporary species known as the transition state (or activated complex) on the way to becoming product molecules.  The activation energy (E_a) is shown on a reaction energy profile as the difference in energy between the reactants and the transition state.  The difference in energy between the reactants and products is represented as ΔH on a reaction energy profile.  When the products have a lower energy than the reactants, the reaction is exothermic and energy is released during the overall reaction:

 

 

 

When the products have a higher energy than the reactants, the reaction is endothermic and energy is absorbed during the overall reaction:

 

 

 

 

Note that the rate of a chemical reaction will be higher when the required activation energy is lower, but the rate of a chemical reaction does not depend on ΔH. 

 

 

 

 

 

 

 

Section 18-3:  Reaction Mechanisms

 

A reaction mechanism depicts at the molecular level the actual series of elementary steps that occur during a chemical reaction and add up to give the overall balanced equation.  If the hypothetical reaction 2 A + B → C occurred in a single step, three reactant molecules would need to collide simultaneously.  Alternatively, it is possible that the reaction occurs via a two-step mechanism wherein two A molecules collide in a first step to form an intermediate (Int.), then the intermediate collides with a B molecule in a second step to form the product C:

 

Step 1:  A + A → Int.

Step 2:  Int. + B → C

 

Note that the intermediate is formed in one step but then consumed in the second step, so the intermediate will not appear in the overall balanced equation:

 

Sample Exercise 18A:

 

The following mechanism has been proposed for a reaction:

 

Step 1:  NO + NO → N₂O₂

Step 2:  N₂O₂ + O₂ → 2 NO₂

 

Identify the intermediate and write the overall balanced equation for the reaction.

 

Solution:

 

N₂O₂ is formed in the first step but then consumed in the second step, so N₂O₂ is the intermediate that cancels out of the overall balanced equation 2 NO + O₂ → 2 NO₂.

 

 

Section 18-4:  Catalysis

 

The addition of a catalyst to a reaction mixture allows the reaction to proceed via an alternate mechanism that is faster overall than the uncatalyzed mechanism.  For example, if the hypothetical uncatalyzed reaction A + B → D proceeds via a slow one-step mechanism, addition of a catalyst (Cat.) can increase the reaction rate by providing the following faster two-step mechanism:

 

 Step 1:  A + Cat. → Int.

Step 2:  Int. + B → D + Cat.

 

Note that the catalyst is consumed in the first step but then produced in the second step, so the catalyst will not appear in the overall balanced equation.  Also, a relatively small quantity of the catalyst must be added initially because each catalyst molecule produced in the second step can be used again as a reactant in the first step.

 

The alternate mechanism provided by a catalyst results in a lower-energy transition state and, therefore, a lower activation energy for the catalyzed reaction, as shown in the reaction energy profile below for an exothermic reaction (solid curve represents the uncatalyzed reaction): 

 

 

 

 

Note that ΔH is unaffected by the addition of a catalyst because the energies of both the reactants and the products are unchanged when a catalyst is added.

 

 

 

 

 

 

 

Chapter 18 Practice Exercises and Review Quizzes:

 

 

 

 

Click for Review Quiz 1

Click for Review Quiz 1 Answers