Kinetics

1. Kinetics Reaction Rates

2. Reaction Rates Collision theory Potential energy diagrams Factors affecting reaction rate Activated complex catalysts temperature Activation energy concentration Surface area

3. Reaction Rates • A measure of how quickly a reaction occurs • An experimental, measurable quantity • Rate = change in property/change in time • Example: Speed = miles/hour • Speed = Ddistance/Dtime • Chemical kinetics: the study of reaction rates and the factors that affect them

4. What could we measure for a reaction? Easily measured properties include: • Change in mass of a solid • Change in concentration • Temperature changes • pH changes • Gas volume changes • Color changes • We must also measure changes in TIME!

5. Writing Rate Expressions • For a general reaction aA + bB  cC + dD • General form: • We need to modify the rate expression to compensate for stoichiometry • The reaction has only one rate for a given set of conditions • Convention: all reaction rates are positive

6. Measuring Rates • Average Rate • Initial Rate • Calculate average rate for early part of data when plot is nearly linear

7. What happens to the rate over time? • Compare average rate at beginning vs. average rate at end • Reaction rates typically slow down over time • Why? • There are fewer moles of reactants left, and therefore fewer collisions.

8. Collision Theory • Molecules must collide in order to react. • They must collide with the correct orientation. • “Effective collision” • Has appropriate orientation; molecules may react. • “Ineffective collision” • Doesn’t have needed orientation; particles will separate.

9. Collision Theory, cont. • Molecules must collide in order to react. • They must have enough energy to react. • Activation Energy, Ea • The minimum energy that reactants must have for the reaction to occur

10. Potential Energy Diagrams • Activation Energy: from reactants to top of “hill” • Transition State • Aka Activated Complex • High energy state, where bonds are broken and new bonds are formed • DHrxn = energy of products – energy of reactants

11. Potential Energy Diagrams • Reactions with a smaller activation energy will occur more quickly than reactions with a larger Ea. Which reaction would you expect to be fastest? Slowest? Why?

12. Collision Theory • Basic premise: • More collisions = faster reaction rate • More collisions = greater likelihood for effective collisions

13. How can we speed up the rate of a reaction? • Increase temperature • Particles move more quickly, so more possible collisions • More particles are likely to have enough energy to overcome activation energy barrier • Increase concentration • More particles, so more possible collisions • Increase surface area • More particles are exposed, so more collisions are possible

14. Catalysts • Speed up reaction rates, without being consumed • Homogeneous vs. heterogeneous catalysts • Enzymes • Catalytic RNA • Catalytic antibodies • Catalytic converter in car engine • Effectively lower the activation energy of the reaction • May even change the mechanism of the reaction