Thermodynamics

1. Thermodynamics Tells if a reaction will occur.

2. Kinetics Tells how fast a reaction will occur.

3. Reaction Rate • Speed of the reaction. • Found experimentally. • Measure change in concentration of a reactant or a product over time. • Rate = Conc time

4. How do you measure rates? • Measure the concentration of 1 or more reactants or products over time. • Reactants disappear • Products appear • The reaction rate is the change in concentration of reactants & products in a given amount of time.

5. Concentration of Reactants, Products Appearance of products Disappearance of reactants

6. 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.

7. Particle Diagram of Collision Activated complex or transition state. Reactants Products NO + O3 NO2 + O2 Activated Complex is NOT in equation!

8. Reaction Rates depend on … • The frequency of collisions – how often they occur And • The efficiency of the collisions – what percentage are effective

9. Collision Theory • Molecules must collide in order to react. • Effective collisions lead to the formation of products. • Ineffective collisions do not lead to products.

10. Effective Collisions • Energetic • Favorable Orientation

11. Effective vs. Ineffective Collision

12. Most collisions are NOT effective!

13. Why Do Collisions Have to be Energetic?

14. Activation Energy & Reaction

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

16. 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.

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

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

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

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

21. 6 Energies to Label 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

22. They mix the arrows up! • You can’t memorize them by location – they move them around. • Have to memorize them by where they start and where they stop. • The 3 arrows for “Potential Energy of …” start at the baseline. • Ea’s start “where you are” & end at the top of the hill.

23. 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?

24. H of reaction

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

26. H of reaction

27. 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.

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

29. 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.

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

31. 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.

32. Maxwell-Boltzmann Distribution

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

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

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

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

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

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

40. Enzymes are catalysts!

41. Catalytic Converter in Engines

42. Hydrogenation & Surface Catalysis

43. Surface Science

44. Reaction Mechanism • A series of steps that leads from reactants to products. • Describes how bonds break, atoms rearrange, and bonds form in a reaction.

45. Elementary Step • Each individual step in a reaction mechanism. • The slowest elementary step is called the rate-determining step.

46. P.E. Curve for Multi-Step Rxn