2. Chemical Kinetics  The area of chemistry that is concerned with the speeds, or rates, of reactions is called chemical kinetics.  Our goal in this chapter is to understand how to determine the rates at which reactions occur and to consider the factors that control these rates.

3. Factors that Affect Reaction Rates 1. The physical state of the reactants: The easier reactants collide, the more rapidly they react When reactants are in different phases, the reaction is limited to their area of contact Need solids with increased surface area

4. Factors that Affect Reaction Rates 2. The concentrations of the reactants: Increase the concentration of one or more of the reactants Increase in concentration leads to an increase in the frequency with which the reactant molecules collide

5. Factors that Affect Reaction Rates 3. The temperature at which the reaction occurs: Increasing temperature increases kinetic energy of molecules As molecules move more rapidly, they collide more frequently with higher energy

6. Factors that Affect Reaction Rates 4. The presence of a catalyst: Catalysts are agents that increase reaction rates without being used up They affect the kinds of collisions that occur (the reaction mechanism)

7. Reaction Rates  The speed of an event is defined as the change that occurs in a given interval of time.  Reaction rate: the change in the concentration of reactants or products per unit of time

8.  Progress of a hypothetical reaction A  B, starting with 1.00 mol A. Each red sphere represents 0.01 mol A, and each blue sphere represents 0.01 mol B. (a) At time zero, the vessel contains 1.00 mol A (100 red spheres) and 0 mol B (0 blue spheres). (b) After 20s, the vessel contains 0.54 mol A and 0.46 mol B. (c) After 40s, the vessel contains 0.30 mol A and 0.70 mol B.

9. Reaction Rates  Rate of this reaction can be expressed either as the rate of disappearance of reactant A or as the rate of appearance of product B.  Average rate of appearance of B = change in concentration of B change in time

10. Reaction Rates  Average rate of disappearance of A = - change in concentration of A change in time  Rates are always expressed as positive quantities  How do we know that A and B are in a one to one mole ratio?

11. Change of Rate with Time

12. Instantaneous Rate  Graphs showing how the concentration of a reactant or product changes with time allow us to evaluate the instantaneous rate ○ Rate at a particular moment in the reaction

13. Instantaneous Rates  From now on “rate” will mean “instantaneous rate” unless indicated otherwise  Rate = - Δ[reactant] Δt  Rate = Δ[product] Δt

14. Reaction Rates and Stoichiometry  In general for a reaction: aA + bB  cC + dD  The rate is given by:

15. Relating Rates at Which Products Appear and Reactants Disappear  (a) How is the rate at which ozone disappears related to the rate at which oxyen appears in the reaction 2O 3 (g)  3O 2 (g) ?  (b) If the rate at which O 2 appears, is 6.0x10 -5 M/s, at what rate is O 3 disappearing at this same time?

16. Relating Rates at Which Products Appear and Reactants Disappear  The decomposition of N 2 O 5 proceeds according to the following equation: 2N 2 O 5 (g)  4NO 2 (g) + O 2 (g) If the rate of decomposition of N 2 O 5 at a particular instant in a reaction vessel is 4.2x10 -7 M/s, what is the rate of appearance of (a) NO 2, (b) O 2 ?

17. The Rate Law  One way of studying the effect of concentration on reaction rate is to determine the way in which the rate at the beginning of a reaction (the initial rate) depends on the starting concentrations. Measure concentration of reactants and a function of time

18. NH 4 + (aq) + NO 2 - (aq)  N 2 (g) + 2H 2 0(l) Experiment #Initial NH 4 + Concentration (M) Initial NO 2 Concentration (M) Observed Initial Rate (M/s) 10.01000.2005.4 x10 -7 20.02000.20010.8 x10 -7 30.2000.020210.8 x10 -7 40.2000.040421.6 x10 -7 50.2000.080843.3 x10 -7

19. The Rate Law  An equation which shows how the rate depends on the concentrations of reactants, is called a rate law.  For a general reaction, aA + bB  cC + dD  The rate law takes the form:

20. Reaction Orders: The Exponents in the Rate Law  Rate laws for most reactions have the general form:  Rate = k[reactant 1] m [reactant 2] n …  The exponents m and n in a rate law are called reaction orders  The overall reaction order is the sum of the orders with respect to each reactant in the rate law.

21. Reaction Orders: The Exponents in the Rate Law  The exponents in a rate law indicate how the rate is affected by the concentration of each reactant.  Although the exponents in a rate law are sometimes the same as the coefficients in the balanced equation, this is not necessarily the case.

22. Units of Rate Constants  Units of the rate constant depend on the overall reaction order of the rate law. units of rate = (units of rate constant)* (units of concentration) Units of rate constant = (units of rate) (units of concentration)

23. Determining Reaction Orders and Units for Rate Constants What is the unit for the rate constant for the following rate law?  2N 2 O 5  4NO 2 + O 2 (a) What is the reaction order of the reactant H 2 in the following equation: H 2 + I 2  2HI? (b) What are the units of the rate constant for this equation?

24. Using Initial Rates to Determine Rate Laws  Must be determined experimentally  Observe the effect of changing the initial concentrations of the reactants on the initial rate of the reaction  The rate of reaction depends on concentration, the rate constant does not depend on concentration

25. Using Initial Rates to Determine Rate Laws  Reaction orders:  0 = changes in concentration of the reactant will have no effect on the rate because any concentration raised to the zero power = 1  1= changes in concentration of reactant will produce proportional changes in the rate  2= changes in concentration of reactant will result in changes by the factor of x 2

26. Determining Rate Law from Initial Rates  The initial rate of a reaction A + B  C  Using this data, determine (a) the rate law for the reaction, (b) the rate constant, (c) the rate of the reaction when [A] = 0.050M and [B] = 0.100M. Experiment #[A] (M)[B] (M)Initial Rate (M/s) 10.100 4.0 x 10 -5 20.1000.2004.0 x 10 -5 30.2000.10016.0 x 10 -5

27. Determining Rate Law from Initial Rates 2 NO(g) + 2H 2 (g)  N 2 (g) + 2H 2 0(g) Using this data, determine (a) the rate law for the reaction, (b) the rate constant, (c) the rate of the reaction when [NO] = 0.050M and [H 2 ] = 0.150M. Experiment #[A] (M)[B] (M)Initial Rate (M/s) 10.10 1.23 x 10 -3 20.100.202.46 x 10 -3 30.200.104.92 x 10 -3

28. The Change of Concentration with Time  Rate laws can be converted into equations that show the relationship between the concentrations of the reactants or products and time. First-order reactions Second-order reactions

29. First-Order Reactions  A first-order reaction is one whose rate depends on the concentration of a single reactant raised to the first power.  Integrated rate laws can be used to determine (1) the concentration of a reactant remaining at an time after the reaction has started, (2) the time required for a given fraction of a sample to react, or (3) the time required for a reactant concentration to fall to a certain level.

30. Using the Integrated First-Order Rate Law  The decomposition of a certain insecticide in water follows first-order kinetics with a rate constant of 1.45yr -1 at 12°C. A quantity of this insecticide is washed into a lake on June 1, leading to a concentration of 5.0x10 -7 g/cm 3. (a) What is the concentration of the insecticide on June 1 of the following year? (b) How long will it take for the concentration of the insecticide to decrease to 3.0x10 -7 g/cm 3 ?

31. Using the Integrated First-Order Rate Law  The decomposition of (CH 3 ) 2 O at 510°C is a first-order process with a rate constant of 6.8x10 -4 s -1. If the initial pressure of (CH 3 ) 2 O is 135torr, what is its pressure after 1420s?

32. Second-Order Reactions  A second-order reaction is one whose rate depends on the reactant concentration raised to the second power or on the concentrations of two different reactants, each raised to the first power.  Examples Examples

33. Half-Life  The half-life of a reaction, t 1/2, is the time required for the concentration of the reactant to reach one-half of its initial value, [A] t1/2 = ½ [A] 0.  Describes how fast a reaction occurs  A fast reaction will have a short half-life

34. Determining the Half-Life of a First-Order Reaction  At 600K, the half-life for a particular process is 3.3x10 5 s. (a) What is the rate constant at this temperature? (b) At 300K, the rate constant is 2.1x10 -5 s -1. What is the half-life at this temperature?

35. Temperature and Rate  The rates of most chemical reactions increase as the temperature rises.  Why?  The faster rate at higher temperature is due to an increase in the rate constant with increasing temperature

36. The Collision Model  We know reaction rates depend on concentration and temperature. The collision model explains why on a molecular level  Based on the kinetic molecular theory

37. The Collision Model  Molecules must collide to react  The greater the number of collisions, the greater the reaction rate 1. How does this explain rate increasing with increasing concentration? 2. How does this explain rate increasing with increasing temperature?

38. The Collision Model  For a reaction to occur, more is required than just a collision

39. The Orientation Factor

40.  Molecules must be oriented a certain way during collisions for a reaction to occur.  The relative orientations of molecules during their collisions determine whether the atoms are suitably positioned to form new bonds.

41. Activation Energy  Molecules must possess a certain minimum amount of energy to react  This energy comes from the kinetic energies of the colliding molecules  Upon collision, the kinetic energy of the molecules can be used to stretch, bend, and ultimately break bonds, leading to chemical reactions

42. Activation Energy  If molecules move too slowly, with too little kinetic energy, they merely bounce off one another without changing  The minimum amount of energy required to initiate a chemical reaction is called the activation energy, E a.

43. Activation Energy  The particular arrangement of atoms at the top of the barrier is called the activated complex or transition state.  The rate depends on the magnitude of E a  The lower E a, the faster the reaction

44. Determining Rate Law Lab -Add 2mL of HCl -Time until no longer see + -Add 2mL of Na 2 S 2 O 3 -Time until no longer see +

45. Determining Rate Law Lab