Today…

1. Today… • Turn in: • Nothing • Our Plan: • Test Results • Videos/Notes • Investigation 10 Pre-Lab • Homework (Write in Planner): • Be prepared for Investigation 10 next class and have report started.

2. Introduction to Chemical Kinetics • What is Kinetics? I’ll let Hank explain… http://www.youtube.com/watch?v=7qOFtL3VEBc • Stop at 3:20

3. Kinetics in Action… • Clock Reactions… • http://www.youtube.com/watch?v=BqeWpywDuiY

4. Unit Preview…Some Reactions • Are fast: • Acid/base neutralization • Sodium + water • PPT reactions • Are slow: • Aluminum oxidation • Iron oxidizing • Plastic bottle decomposing

5. And some depend… • On temperature: fireflies • On conditions: Iron oxide (humid or desert) • On enzymes: many processes in our body • On a catalyst: 2CO + 2NO --> 2CO2 + N2 (automobile pollution and catalytic converters)

6. Chemical Kinetics: A Preview • Chemical kinetics is the study of: • the rates of chemical reactions • factors that affect these rates • the mechanisms by which reactions occur • Reaction rates vary greatly – some are very fast (burning, precipitation) and some are very slow (rusting, disintegration of a plastic bottle in sunlight).

7. Variables in Reaction Rates • Concentrations of reactants: Reaction rates generally increase as the concentrations of the reactants are increased. • Temperature: Reaction rates generally increase rapidly as the temperature is increased. • Surface area: For reactions that occur on a surface rather than in solution, the rate increases as the surface area is increased. • Catalysts: Catalysts speed up reactions and inhibitors slow them down.

8. Think of it like this… • http://ed.ted.com/lessons/how-to-speed-up-chemical-reactions-and-get-a-date

9. Theories of Chemical Kinetics: Collision Theory • Before atoms, molecules, or ions can react, they must first collide. • An effective collision between two molecules puts enough energy into key bonds to break them. • The activation energy (Ea) is the minimum energy that must be supplied by collisions for a reaction to occur. • A certain fraction of all molecules in a sample will have the necessary activation energy to react; that fraction increases with increasing temperature. • The spatial orientationsof the colliding species may also determine whether a collision is effective.

10. Distribution of Kinetic Energies At higher temperature (red), more molecules have the necessary activation energy.

11. Key Point Orientation of molecules at the time of their collision will determine whether they react or not!

12. Analogy - Car Crash Example Energy and orientation of cars during a car crash can establish the change that occurs to the cars.

13. Importance of Orientation One hydrogen atom can approach another from any direction … Effective collision; the I atom can bond to the C atom to form CH3I … and reaction will still occur; the spherical symmetry of the atoms means that orientation does not matter. Ineffective collision; orientation is important in this reaction.

14. Transition State Theory • The configuration of the atoms of the colliding species at the time of the collision is called the transition state. • The transitory species having this configuration is called the activated complex. • A reaction profile shows potential energy plotted as a function of a parameter called the progress of the reaction. • Reactant molecules must have enough energy to surmount the energy “hill” separating products from reactants.

15. Activated complex • Species formed as a result of collisions between energetic molecules that is an intermediate between the reactants and the products of a reaction. Once formed the activated complex dissociates either into products or back to the reactants

16. Reaction Profile • A graphical representation of a chemical reaction in terms of the energies of the reactants, activated complexes and products

17. Reaction Profile

18. Exothermic

19. Endothermic

20. A Reaction Profile CO(g) + NO2(g)  CO2(g) + NO(g)

21. An Analogy for Reaction Profiles and Activation Energy

22. Stop! • Investigation 10 - We’re going to test the effect of different variables on the rate of reaction next class. • Complete steps 1 – 6 on the lab handout and be prepared to conduct the lab next class.

23. Today… • Turn in: • Nothing • Our Plan: • Investigation 10 • Homework (Write in Planner): • Lab Report Due Friday

24. Today… • Turn in: • Lab Report (rubric on top) • Our Plan: • Notes & Practice • Homework (Write in Planner): • Work on the Ch. 13 Homework

25. The Meaning of Rate • The rate of a reaction is the change in concentration of a product per unit of time (rate of formation of product). • Rate is also viewed as the negative of the change in concentration of a reactant per unit of time (rate of disappearance of reactant). • The rate of reaction often has the units of moles per liter per unit time (mol∙L–1∙s–1or M∙s–1)

26. If the rate of consumption of H2O2 is 4.6 M/h, then … … the rate of formation of H2O must also be 4.6 M/h, and … … the rate of formation of O2 is 2.3 M/h

27. 1 L 0.1850 mol H2O2/L Rate = 60 s 2 H2O2 --> 2 H2O+ O2 2.960 g O2 (0.09250 mole) produced in 60 s means … … 0.1850 mol H2O2reacted in 60 s. = 0.00131 M H2O2 s–1

28. Example 13.1 Consider the hypothetical reaction A + 2B --> 3C + 2D Suppose that at one point in the reaction, [A] = 0.4658 M and 125 s later [A] = 0.4282 M. During this time period, what is the average (a) rate of reaction expressed in M∙s–1and (b) rate of formation of C, expressed in M∙min–1.

29. Try It Out! EX 13.1 A: Consider the hypothetical reaction 2A + B → 2C + D Suppose that at some point during the reaction [D] = 0.2885 M and that 2.55 min (that is 2 min 33 sec) later [D] = 0.3546 M. • What is the average rate of reaction during this time period, expressed in M min-1? • What is the rate of formation of C expressed in M s-1?

30. Rate of Reaction Expressed as the negative of the Slope of a Tangent Line • Average rate (green dotted line) • Initial Rate ( blue solid line) • Instantaneous Rate (red line) • Question: Over what time interval are the instantaneous rates greater than the average rate measured for the 600 sec period?

31. Average vs. Instantaneous Rate Instantaneous rate is the slope of the tangent to the curve at a particular time. We often are interested in the initial instantane-ous rate; for the initial concentrations of reactants and products are known at this time.

32. Example 13.2 Use data from Table 13.1 and/or Figure 13.5 to • determine the initial rate of reaction and • calculate [H2O2] at t = 30 s.

33. Try It Out EX 13.2 A: From Figure 13.5, • Determine the instantaneous rate of reaction at t = 300 s. • Use the result of a) to calculate a value of [H2O2] at t = 310 s.

34. The Rate Law of a Chemical Reaction • The rate law for a chemical reaction relates the rate of reaction to the concentrations of reactants. aA + bB + cC…→ products rate = k[A]n[B]m[C]p … • The exponents (m, n, p…) are determined by experiment. • Exponents are not derived from the coefficients in the balanced chemical equation, though in some instances the exponents and the coefficients may be the same. • The value of an exponent in a rate law is the order of the reaction with respect to the reactant in question. • The proportionality constant, k, is the rate constant.

35. Rate Law Examples Rate = k[A]1 = k[A] Reaction is first order in A Rate = k[A]2 Reaction is second order in A Rate = k[A]3 Reaction is third order in A If we triple the concentration of A in a second-order reaction, the rate increases by a factor of ________.

36. Effect of Concentration on Rate

37. More About the Rate Constant k • The rate of a reaction is the change in concentration with time, whereas the rate constant is the proportionality constant relating reaction rate to the concentrations of reactants. • The rate constant remains constant throughout a reaction, regardless of the initial concentrations of the reactants. • The rate and the rate constant have the same numerical values and units only in zero-order reactions. • For reaction orders other than zero, the rate and rate constant are numerically equal only when the concentrations of all reactants are 1 M. Even then, their units are different.

38. To find the overall order of a reaction… • Add the orders for each compound. • Example: • rate = [A]2[B]1 is 3rd order overall • How about rate = [A]0[B]1[C]1?

39. Units of k (p. 536)

40. Method of Initial Rates • The method of initial rates is a method of establishing the rate law for a reaction—finding the values of the exponents in the rate law, and the value of k. • A series of experiments is performed in which the initial concentration of one reactant is varied. Concentrations of the other reactants are held constant. • When we double the concentration of a reactant A, if: • there is no effect on the rate, the reaction is zero-order in A. • the rate doubles, the reaction is first-order in A. • the rate quadruples, the reaction is second-order in A. • the rate increases eight times, the reaction is third-order in A.

41. The concentration of NO was held the same in Experiments 1 and 2 … … while the concentration of Cl2 in Experiment 2 is twice that of Experiment 1. The rate in Experiment 2 is twice that in Experiment 1, so the reaction must be first order in Cl2. Which two experiments are used to find the order of the reaction in NO? How do we find the value of k after obtaining the order of the reaction in NO and in Cl2?

42. Example 13.3 For the reaction 2 NO(g) + Cl2(g) →2 NOCl(g) described in the text and in Table 13.2, (a) what is the initial rate for a hypothetical Experiment 4, which has [NO] = 0.0500 M and [Cl2] = 0.0255 M? (b) What is the value of k for the reaction?

43. So, when looking at data….. • Zero order reaction: initial rate is unaffected (20 ) • 1st order reaction: double the concentration, doubles the rate (21 =2) • 2nd Order Reaction: Initial rate increases fourfold (22 =4) • 3rd Order Reaction: Initial rate increases eightfold (23 =8)

44. Try It Out • Below is some rate data for the hypothetical reaction, 2A + B --> C. What is the rate law for this reaction?

45. Let’s Take a Break • With your partner, Complete Part 1 of the Partner Review

46. Zero-Order Reactions • Sum of exponents is equal to zero

47. Zero Order Rate of Reaction= k[A]0 • The concentration time graph is a straight line with a negative slope

48. Zero Order • Rate of the reaction: • IS equal to k • remains constant throughout the reaction • is the negative of the slope of the line when graph Molarity vs. time (see pg. 544)

49. A Zero-Order Reaction rate = k[A]0 = k Rate is independent of initial concentration

50. Zero Order Integrated rate equation: [A]t = -kt + [A]0 y = mx + b