Chapter 14

1. Chapter 14 Chemical Kinetics

2. Review Section of Chapter 14 Test • Net Ionic Equations

3. Reaction Rate • The rate of a chemical reaction is measured as the decrease in the concentration of a reactant or the increase in the concentration of a product in a unit of time.

4. Reaction Rate • The rate of a chemical reaction is measured as the decrease in the concentration of a reactant or the increase in the concentration of a product in a unit of time. ∆[ ] Rate = ∆time

5. Reaction Rate • The rate of a chemical reaction is measured as the decrease in the concentration of a reactant or the increase in the concentration of a product in a unit of time. ∆[ ] What units would we use for rate? Rate = ∆time

6. Reaction Rate • The rate of a chemical reaction is measured as the decrease in the concentration of a reactant or the increase in the concentration of a product in a unit of time. ∆[ ] Rate = ∆time 2H2O2(aq) → 2H2O(l) + O2(g)

7. Reaction Rate • The rate of a chemical reaction is measured as the decrease in the concentration of a reactant or the increase in the concentration of a product in a unit of time. ∆[ ] Rate = ∆time 2H2O2(aq) → 2H2O(l) + O2(g) How could the rate be expressed for this reaction in terms of H2O2?

8. 2H2O2(aq) → 2H2O(l) + O2(g)

9. 2H2O2(aq) → 2H2O(l) + O2(g)

10. 2H2O2(aq) → 2H2O(l) + O2(g) What is the rate of the reaction from 0s to 2.16 x 104s?

11. 2H2O2(aq) → 2H2O(l) + O2(g) What is the average rate of appearance of O2 from 0s to 2.16 x 104s? 1.16 x 10-5 mol O2 L-1 s-1

12. General Rate of Reaction a A + b B → c C + d D Rate of reaction = rate of disappearance of reactants or Rate of reaction = rate of appearance (formation) of products We can use the coefficients in the equation to compare the reaction rates for all the substances in the reaction.

13. Δ[Fe2+] 0.0010 M Rate of formation of Fe2+= = = 2.6 x 10-5 M s-1 Δt 38.5 s 15-1 The Rate of a Chemical Reaction • Rate is change of concentration with time. 2 Fe3+(aq) + Sn2+(aq) → 2 Fe2+(aq) + Sn4+(aq) t = 38.5 s [Fe2+] = 0.0010 M ∆t = 38.5 s ∆[Fe2+] = (0.0010 – 0) M

14. Rates of Chemical Reaction 2 Fe3+(aq) + Sn2+(aq)→ 2 Fe2+(aq) + Sn4+(aq) Rate of formation of Fe2+ = 2.6 x 10-5 mol L-1 s-1 What is the rate of formation of Sn4+? 1.3 x 10-5 mol Sn4+ L-1 s-1 What is the rate of disappearance of Fe3+? 2.6 x 10-5 mol Fe3+ L-1 s-1

16. What does the slope of the line represent?

17. Δ[H2O2] Rate = 1.7 x 10-3 M s-1 = Δt Δ[H2O2] = (1.7 x 10-3 M s-1) (∆t) ∆[H2O2]= (1.7 x 10-3 M s-1)(100 s) = 0.17M [H2O2]100 s = 2.32 M - 0.17 M What is the concentration at 100s for the reaction: 2H2O2(aq) → 2H2O(l) + O2(g)?Given: [H2O2]i = 2.32 M = 2.15 M

18. What does it mean when the rate of a reaction reaches zero? • For a normal reaction it means that one or more of the reactants are used up and the reaction has stopped. • For a reversible reaction it means that the reaction has reached equilibrium.

19. Factors Affecting Reaction Rates • The nature of the reacting substances.

20. Factors Affecting Reaction Rates 2. The state of subdivision of the reacting substances.

21. Lycopodium Powder

22. Factors Affecting Reaction Rates 3. The temperature of the reacting substances.

23. Factors Affecting Reaction Rates 4. The concentration of the reacting substances. Air (21% oxygen) 100% oxygen

24. Factors Affecting Reaction Rates • The presence of a catalyst. Catalysts speed up reactions but are left unchanged by the reaction.

25. The Rate Law a A + b B …. → g G + h H …. Rate = k [A]m[B]n …. Rate constant = k (k is constant for a particular reaction at a specific temperature) Order of A = m Order of B = n Overall order of reaction = m + n + ….

26. Temperature and Rate • Generally, as temperature increases, so does the reaction rate. • This is because k is temperature dependent.

27. Concentration and Rate Summary • After finding the trials to compare: • A reactant is zero order if the change in concentration of that reactant produces no effect on the rate. • A reaction is first order if doubling the concentration of that reactantcauses the rate to double. • A reactant is nth order if doubling the concentration of that reactantcauses an 2n increase in rate. • Note that the rate constant does not depend on concentration.

28. Use the data provided to write the rate law and indicate the order of the reaction with respect to HgCl2 and C2O42- and also the overall order of the reaction.

29. First determine the order of HgCl2

30. Next determine the order of C2O42-

31. Now write the rate law and determine the order of the reaction.

32. Calculate the rate constant “k” and its units. Initial rate of disappearance HgCl2 mol L-1 min-1

33. What is the average rate of disappearance of C2O42- in trial 1? Initial rate of disappearance HgCl2 mol L-1 min-1

34. Use the data provided to write the rate law and indicate the order of the reaction with respect to NO2 and CO (support your answers). Also give the overall order of the reaction.

36. Calculate the rate constant “k” and its units.

37. What is the average rate of disappearance of CO in trial 2?

38. Collision Model • Key Idea: Molecules must collide to react. • However, only a small fraction of collisions produces a reaction. Why?

39. Two Factors • Collisions must have enough energy to produce the reaction (must equal or exceed the activation energy). • Orientation of reactants must allow formation of new bonds.

40. 2HI → H2 + 2I

41. Concentration and Collision Theory • Why does an increase in concentration cause an increase in reaction rate?

42. Concentration and Collision Theory • Why does an increase in concentration cause an increase in reaction rate? • Increasing the concentration increases the number of collisions and therefore there are more collisions leading to product.

43. Temperature and Collision Theory • Why does a temperature increase cause the reaction rate to increase?

44. Temperature and Collision Theory • Why does a temperature increase cause the reaction rate to increase? • At higher temperatures there are more collisions and a greater percentage of the collisions have the energy necessary to create a successful collision.

45. Activation Energy • The activation energy is the minimum amount of energy necessary for a reaction to occur.

46. Temperature and Activation Energy (Ea)Figure 14.12 (Page 432)

47. Activation Energy • The activation energy can also be thought of as the energy necessary to form an activated complex during a collision between reactants.

48. Transition State Theory • The activated complex is a hypothetical species lying between reactants and products at a point on the reaction profile called the transition state.

49. The activated complex is a transition state between reactants and products where old bonds have begun to break and new bonds have started to form. It cannot be isolated.

50. Determining the Activation Energy“The Arrhenius Equation” • Collisions must have enough energy to produce the reaction (must equal or exceed the activation energy). • Orientation of reactants must allow formation of new bonds.