Reaction Rates

1. Reaction Rates • Different reactions take place at different rates. • Spontaneous vs. Non-spontaneous (reactants vs. products) • Nature tends towards stability

2. Reaction Rates • Thermodynamically stable: a reaction which has an overall positive energy change (non-spontaneous, endothermic)

3. Reaction Rates

4. Reaction Rates • Kinetically stable: a reaction that proceeds so slowly that no change is detectable.

5. Reaction Rates • Some reactions go to completion. That is, it proceeds until one of the reactants is used up and then stops.Most do not.

6. Reaction Rates • Many reactions are reversible. • Reactants produce products, while products may decompose back into reactants. A + B AB

7. Reaction Rates • When the rate at which the products form, equals the rate at which the products decompose, the reaction has reached dynamic equilibrium.

8. Reaction Rates • NOTE: chemical equilibrium says nothing about the amount of products produced (some reactions only produce a small amount of product). Why?

9. Reaction Rates • Reversibilty implies that there is always some reactants present, even after products have formed and equilibrium has been established.Why?

10. Reaction Rates • Reaction rates are measured in terms of the appearance of the products, or the disappearance of the reactants.Which do you choose?

11. Reaction Rates • As a reaction proceeds, samples can be taken and the concentrations of the reactants and products determined. Rates are then computed.

12. Reaction Rates • What is concentration? • How is it measured? Which units? • What would be the units for reaction rates?

13. Reaction Rate • Meaning: rate = conc. of species/time • aA + bB --> cC + dD • Rate = C]/ct = -A]/at • Why is rate, computed by species]/t also divided by the molar coefficient?

14. Reaction Rate • The minus sign is used because concentrations of reactants decrease with time. The concentration will be expressed in molarity; time may be in seconds, years, etc.

15. Reaction Rates Concentration vs Time

16. Reaction Rates Concentration vs Time

17. Reaction Rates • Factors: Nature of the reactants, concentration of the reactants, temperature, amount of reactant surface area, and the presence of a catalysis.

18. Nature of Reactants • Reactions with bond rearrangements take the longest. Covalent? • Ionic reactions proceed almost instantaneously.

19. Nature of Reactants • Collisions must have enough energy to rearrange bonds to form products.

20. Nature of Reactants • Molecules sometimes have to form activated complexes. Activation energy is the energy required to do this.

21. Concentration • Concentrations: increasing the concentration of a reactant increases the reaction rate. Why? • How does pressure effect a reaction?

22. Temperature • How does temperature effect reaction rate?

23. Phases • Homogeneous reaction: a reaction where all the reactants are in the same phase. • Heterogeneous reactions: reactants are in differentphases (interface).

24. Heterogeneous rx • The reaction must take place on the surface of the solid or liquid. How does the amount of exposed surface area affect the rate of the reaction? How can it be increased?

25. Catalyst • Catalysts: a chemical substance that lowers the activation energy and speeds up a reaction but is not consumed in the reaction.

26. Catalyst • Homogeneous vs heterogeneous catalysts. • Activated complex vs surface catalysts. • Inhibitors

27. Catalyst • Heterogeneous catalysts: uses adsorption to the surface to speed the production of a product.

28. Catalyst • Homogeneous catalysts: 1. reduce activation energy, 2. assist in bringing together reactants.

29. Catalyst • 3. Low activation energy = more successful collisions = faster reaction.

30. Catalyst • Homogeneous catalyst = activated complex • Heterogeneous catalysts = surface catalysts.

31. Catalyst • Inhibitors = do not slow reactions, rather they compete for and tie-up reactants.

32. Rate Expressions • Rate expression / rate constants • Reaction mechanisms • Rate determining step

33. Rate Expressions • The reaction rate varies directly as the product of the concentrations of the reactants.

34. Rate Expressions • Rate expression: • rate = k[reactants]n • The rate expression is always determined experimentally.

35. Rate Expressions • The specific rate constant is dependent on the size, speed, and kind of molecules involved in the reaction.

36. Rate Expressions • There is one k for each reaction at a given temperature.

37. Rate Expressions • The exponents of the concentrations determines the ‘order’ of the reaction. • Ie. A squared concentration in the expression is considered second order.

38. Rate and Concentration • Rate dependence on concentration: single reactant, A: rate = k[A]m; k = rate constant, m = order • Two reactants, A, B; rate = k[A]m[B]n, overall order = m+n

39. Rate and Concentration • In general, m and n are usually positive integers. However, they can be 0 or a fraction.

40. Rate/Concentration Determine m and k from the following rate-conc. data: CH3CHO(g) --> CH4(g) + CO(g);

41. Rate/Concentration • Sample problem

42. Rate/Concentration Determine m and k from the following rate-conc. data: CH3CHO(g) --> CH4(g) + CO(g) m = 2, k = 2.0 (M.s)-1

43. First Order Kinetics • First order reactions: • ln ([A]o / [A]) = kt • [A]o = original conc. of reactant A • [A] = conc. of A at any time t.

44. First Order Kinetics Suppose k = 0.250/s, [A]o = 1.00 M. What is the conc. of A after 10.0s? [A] = 0.0819 M

45. First Order Kinetics How long does it take to drop to one half its original value? t1/2 = 2.77s

46. First Order Kinetics • Note that, for a first order reaction: • t1/2 is independent of the initial conc. It takes just as long to drop from 2.0M to 1.0M as from 1.0M to 0.50M.

47. First Order Kinetics • Note that, for a first order reaction: • t1/2 is inversely related to k. If t1/2 is small, then k is large and vice versa.

48. Zero Order Kinetics • Rate = k; [A] = [A]o - kt. The plot of [A] versus time is linear.

49. Second Order Kinetics • Rate = k[A]2; 1/[A] - 1/[A]o = kt. The plot of 1/[A] versus time is linear.

50. Reaction Mechanisms • A series of reactions may have to take place for a reaction to go to completion. The steps are considered the reaction mechanism.