1. Chemical Kinetics How quickly does that chemical reaction occur?

2. Goals  Determine factors affecting reaction rates.  Experimentally measure rates of reaction.  Determine how to mathematically describe and predict rates of reaction.  Be able to use data to propose mechanisms for chemical reactions.  Understand how catalysts can change the rate of chemical reactions

3. Chemical Reaction Rates  Chemical reactions occur at different rates.  Ex. Oxidation of steel wool  http://www.youtube.com/watch?v=5MDH92VxPEQ http://www.youtube.com/watch?v=5MDH92VxPEQ  How are the rates of chemical reactions measured?

4. Factors Affecting the Rates of Reaction  The physical state of the reactants  The concentration of the reactants  The temperature at which the reaction occurs  The presence of a catalyst

5. Reaction Rates  Rate = change in concentration / change in time  What unit is used?

6. Calculating an Average Reaction Rate  Imagine a hypothetical reaction in which A  B Calculate the average rate at which A disappears over the time interval from 20 sec to 40 sec. Time (sec)[A] (M)[B] (M) 01.000 200.540.46 400.300.70

7. Rates change over time  Look at the following reaction: C 4 H 9 Cl (aq) + H 2 O (l)  C 4 H 9 OH (aq) + HCl (aq)  What happens to the rate of the chemical reaction over time?  Why?

8. Instantaneous Rates of Chemical Reaction  To find the instantaneous rate of a chemical reaction at a particular time interval, find the slope of the tangent line at the point of interest.  Calculus is often used to do this, You will sometimes see this infomation written using calculus notation.

9. Finding instantaneous rate graphically  Determine the instantaneous rate of reaction for butyl chloride at t = 0 sec.

10. Instantaneous Reaction Rates  In chemical kinetics we will always assume rate to mean instantaneous reaction rate.

11. Stoichiometric Relationships in Kinetics  Imagine the following reaction C 4 H 9 Cl (aq) + H 2 O (l)  C 4 H 9 OH (aq) + HCl (aq)  How does the rate of decrease of butyl chloride (C 4 H 9 Cl) compare to the rate of increase of butyl alcohol (C 4 H 9 OH)?

12. Stoichiometric Relationships in Kinetics  Imagine the following reaction 2 HI (g)  H 2(g) + I 2(g)  How is the rate of formation of iodine related to the rate of HI decrease

13. General Stoichiometric Relationships  For the general reaction: aA + bB  cC + dD

14. Sample Problem  The decomposition reaction of N 2 O 5 is represented in the following equation. 2 N 2 O 5 (g)  4 NO 2 (g) + O 2 (g) If the rate of disappearance of N 2 O 5 = 4.2x10 -7 M/sec, what is the rate of appearance of each of the products?

15. Rate Laws Mathematical Expressions to Predict Reaction Rates

16. Rate of Reaction is Dependent on Concentration  Examine the reaction between nitrogen monoxide and hydrogen 2 NO (g) + 2 H 2 (g)  N 2 (g) + 2 H 2 O (g) The initial rates of reaction under varying concentrations: Exp. #[NO] (M)[H 2 ] (M)Initial Rate (M/sec) 10.10 1.23x10 -3 20.100.202.46x10 -3 30.200.104.92x10 -3

17. Determining rate law  What happens to the rate when [NO] doubles?  What happens to the rate when [H 2 ] doubles?

18. Using the information give, determine k Exp. #[NO] (M)[H 2 ] (M)Initial Rate (M/sec) 10.10 1.23x10 -3 20.100.202.46x10 -3 30.200.104.92x10 -3

19. Using the information we have calculated so far …  Determine the initial rate of reaction when [NO] = 0.050 M and [H 2 ] = 0.150 M.

20. A note about rate laws  You cannot determine a rate law by looking at an equation.  You must use experimental data to determine a rate law.  The most common exponents are 1 and 2. It also possible to have exponents of 0, ½, and 3.  We will discuss the physical reason for exponents later in the unit.

21. Rate Laws and Reaction Order  First order reactions: Reactions whose rate depends on the concentration of only one reactant raised to the first power.  Imagine the reaction A  B where the rate of the reaction is dependent on [A] so that The rate law in this form is called the differential rate law.

22. Integrated rate law  Integrating the rate law gives us:  This can be manipulated to fit the form of y = mx + b  Or expressed as:

23. 1 st Order Rate Law Practice Problem  The decomposition of an insecticide in water follows first order reaction kinetics with a k = 1.45/yr at 12 o C. A quantity of insecticide washes into the lake so that the [insecticide] in the lake = 5.00x10 -7 g/cm 3. Assume that the average temp of the lake is 12 o C.  A. What is the [insecticide] after one year?  B. How long will it take for [insecticide] to decrease to 3.0x10 -7 g/cm 3 ?

24. Second order reactions  Reactions in which two reactants react with and each reactant reacts has an exponent of one in the rate law. OR  Reactions with one reactant reacts and has an exponent of 2 in the rate law.

25. Second Order Reactions  Consider the reaction: A  B Where the reaction order is 2 so that the differential rate law is:

26. Integrated Rate Law: 2 nd order reaction  Integration of the previous equation yields  Notice the y = mx + b format of the resulting equation.

27. Sample Problem  In the following reaction NO 2 (g)  NO (g) + ½ O 2 (g) the decomposition of NO 2 is second order with k = 0.543/M sec. If the initial [NO 2 ] = 0.0500 M, what is the remaining concentration after 0.500 hour?

28. Using Kinetic Data to Determine Reaction Order  Cyclopentadiene (C 5 H 6 ) reacts with itself to form dicyclopentadiene (C 10 H 12 ). A 0.0400 M solution of C 5 H 6 was monitored as a function of time as the reaction proceeded and the following data was collected.  From this data determine the order of the reaction. Time (s)[C 5 H 6 ] (M) 0.00.0400 50.00.0300 100.00.0240 150.00.0200 200.00.0174

29. How to solve this kind of problem? Examine the integrated rate law equations 1 st order 2 nd order If it is first order, what will a graph of ln [C 5 H 6 ] v. time look like? If it is second order, what will a graph of 1/[C 5 H 6 ] v. time look like?

32. Determine the value of the rate constant (k) from this data.  For the slope of the line is equal to k

33. What happens in a chemical reaction?  Molecules involved must collide with one another.  For the reaction to occur, the molecules must  Collide at the with the correct orientation  Collide with sufficient energy. http://www.youtube.com/watch?v=pTp0R6WuSks&feature=relate d

34. Remember the Reaction Profile

35. Effects of temperature  Increasing temperature increases the energy at which collisions occur.  This means that there will be a greater chance that molecules will collide with sufficient activation energy.

36. Relating E a to temperature  The fraction of molecules with sufficient energy to react is given by the equation  Note: If all other factors are equal, the primary determinant of reaction rate is activation energy.

37. Arrhenius Equation  Arrhenius noted this relationship and noted the effect of collision frequency.  Modified equation to find the rate constant (k)  A is the frequency factor and is nearly constant as temperature varries.

38. Arrhenius Equation in Point-Slope format  With this equation, a graph of ln k vs. 1/T will give a line with slope = -E a /R. From this E a can be easily determined.

39. Sample Problem  The following values of k were determined at different temperatures for a chemical reaction.  Find the activation energy. Temperature (OC)K (s -1 ) 189.72.52x10 -5 198.95.25x10 -5 230.36.30x10 -4 251.23.16x10 -3

40. Solving  Arrange so that you can graph. T (K)1/T (K -1 )ln k 462.92.160x10 -3 -10.598 472.12.118x10 -3 -9.855 503.51.986x10 -3 -7.370 524.41.907x10 -3 -5.757

42.  Slope is equal to –Ea/R.  For R, use 8.314 J/mol K  Find Ea Ea = -slope R = -(-19105K)(8.314J/mol K) = 1.6 kJ/mol

43. Reaction Mechanisms  See notes and handouts from class.