1. How do reactions occur?
Must have an effective collision between reacting particles for reaction to occur. “Collision Theory”
Collision must be energetic.
Collision must occur at an effective angle.

2. Particle Diagram of Collision
Activated complex or transition state.
Reactants
Products
NO O3  NO O2
Activated Complex is NOT in equation!

3. Effective vs. Ineffective Collision

4. Most collisions are NOT effective!

5. Why Do Collisions Have to be Energetic?

6. Activation Energy & Reaction

7. Energy Diagram of a Reaction
Activated Complex
Reactants
Enthalpy or Potential Energy
Products
Reaction Pathway

8. Role of Energy in Reactions
In order for a reaction to begin, energy is needed.
Energy needed to begin the reaction is the activation energy.
The activation energy comes from effective collisions.

9. Role of Energy in Reactions (con’t)
During a chemical reaction, heat may be released or absorbed.
Heat released or absorbed during a chemical reaction is called heat of reaction or enthalpy (rH).
Enthalpy is the difference between the potential energy of the products and the reactants:
rH = H (products) – H (reactants)

10. Catalysts reduce the activation energy but have no effect on the change in enthalpy (rH)

11. Role of Energy in Reactions (con’t)
Exothermic Reactions
Reactions in which energy is released
The potential energy of the products is lower than the potential energy of the reactants
The sign of rH is negative

12. Role of Energy in Reactions (con’t)
Endothermic Reactions
Reactions in which energy is absorbed
The potential energy of the products is higher than the potential energy of the reactants
The sign of rH is positive

13. Activation Energy Energy needed to initiate the reaction.
Energy needed to overcome the reaction barrier.
The difference between the top of the hill & where you start.
Difference between activated complex & reactants.

14. Activation Energy Using a match to start a fire.
The spark plug in a car engine.

15. Potential Energy Curve: Endothermic
Endothermic Reaction:
Products have more P.E. than reactants.
Start low, end high.

16. Potential Energy Curve: Exothermic
Exothermic Reaction:
Products have less P.E. than reactants.
Start high, end low.

17. Have to label 6 energies on curve.
1)Ea = Activation Energy
2)H = Hproducts – Hreactants

18. 6 Energies to Label P.E. of reactants P.E. of products
Label on both endo & exo P.E. curves.
P.E. of reactants
P.E. of products
P.E. of activated complex
Ea for the forward reaction
Ea for the reverse reaction H

19. What kind of reaction is represented?
Ea for reverse rxn
Ea for forward rxn
P.E. of activated complex
P.E. of products
P.E. of
reactants
What kind of reaction is represented?

20. H of reaction

21. What kind of reaction is represented?
forward Ea
reverse
P.E. of reactants
P.E. of activated complex
P.E. of products
What kind of reaction is represented?

22. H of reaction

23. Why does the collision have to be energetic?
The kinetic energy of the reactants is used to overcome the reaction barrier.
The kinetic energy is transformed into potential energy.

24. Factors that determine reaction rates
Nature of the reactants (ions vs. molecules)
Temperature
Concentration
Pressure (for gases)
Surface Area
The presence of a catalyst

25. Nature of the reactants: Ions or Molecules?
Ions in solution react quickly.
Covalently bonded molecules react slowly. It takes time to break all those bonds!
2 gas phase reactants tend to react more quickly than 2 liquids or 2 solids.

26. Temperature Rule of thumb:
Increasing the temperature 10oC doubles the reaction rate.

27. Temperature
A measure of the average kinetic energy of the molecules in a system.
The faster they are moving, the more often they will collide.
The faster they are moving, the more energetic the collisions.

28. Maxwell-Boltzmann Distribution

29. Increase in Temperature
Increases the frequency of collisions
Increases the percentage of collisions that lead to reaction.

30. Concentration
Increase in concentration means more particles per unit volume – so more collisions in a given amount of time.

31. Pressure For systems involving gases.
Analogous to increasing concentration.
 Pressure,  number of particles per unit volume.

32. Surface Area
Higher surface area – more particles exposed for reaction.
Higher surface area means smaller particle size.
(For heterogeneous reactions.)

33. Vocabulary Interlude
Homogeneous Reaction: all reactants are in the same phase.
Heterogeneous Reaction: reactants are in different phases.

34. Catalyst
Substance that increases the rate of reaction without itself being consumed.
Provides an alternate reaction pathway with a lower energy barrier.

36. Enzymes are catalysts!

37. Catalytic Converter in Engines