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CHAPTER 12

CHAPTER 12. Chemical Kinetics. Kinetics. The study of reaction rates. Spontaneous reactions are reactions that will happen - but we can’t tell how fast. Diamond will spontaneously turn to graphite – eventually. Reaction mechanism- the steps by which a reaction takes place. Reaction Rate.

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CHAPTER 12

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  1. CHAPTER 12 Chemical Kinetics

  2. Kinetics • The study of reaction rates. • Spontaneous reactions are reactions that will happen - but we can’t tell how fast. • Diamond will spontaneously turn to graphite – eventually. • Reaction mechanism- the steps by which a reaction takes place.

  3. Reaction Rate • Reaction Rate • Rate = Conc. of A at t2 -Conc. of A at t1 t2- t1 • Rate =D[A] Dt • Change in concentration per unit time • For this reaction • N2 + 3H2 2NH3

  4. REACTION RATE GRAPH • As Time Increases, Concentration deceases

  5. H2 VS N2 • As the reaction progresses the concentration N2 goes down 1/3 as fast

  6. PRODUCT CONCENTRATION • As the reaction progresses the concentration NH3 goes up.

  7. Calculating Rates Average rates are taken over long intervals • Instantaneous rates are determined by finding the slope of a line tangent to the curve at any given point because the rate can change over time • Derivative.

  8. Average slope method SLOPE

  9. Instantaneous slope method • Instantaneous slope method

  10. Defining Rate We can define rate in terms of the disappearance of the reactant or in terms of the rate of appearance of the product. • In our example • N2 + 3H2  2NH3 • -D[N2] = -3D[H2] = 2D[NH3] Dt Dt Dt

  11. Rate Laws Reactions are reversible. • As products accumulate they can begin to turn back into reactants. • Early on the rate will depend on only the amount of reactants present. • We want to measure the reactants as soon as they are mixed. • This is called the Initial rate method.

  12. Rate Laws • Two key points • The concentration of the products do not appear in the rate law because this is an initial rate. • The order must be determined experimentally, • can’t be obtained from the equation

  13. 2 NO2  2 NO + O2 • You will find that the rate will only depend on the concentration of the reactants. • Rate = k[NO2]n • This is called a rate law expression. • k is called the rate constant. • n is the order of the reactant -usually a positive integer.

  14. 2 NO2  2 NO + O2 • The rate of appearance of O2 can be said to be. • Rate' = D[O2] = k'[NO2] Dt • Because there are 2 NO2 for each O2 • Rate = 2 x Rate' • So k[NO2]n = 2 x k'[NO2]n • So k = 2 x k'

  15. Types of Rate Laws • Differential Rate law - describes how rate depends on concentration. • Integrated Rate Law - Describes how concentration depends on time. • For each type of differential rate law there is an integrated rate law and vice versa. • Rate laws can help us better understand reaction mechanisms.

  16. Determining Rate Laws The first step is to determine the form of the rate law (especially its order). • Must be determined from experimental data. • For this reaction • 2 N2O5 (aq)  4NO2 (aq) + O2(g)The reverse reaction won’t play a role

  17. Now graph the data • [N2O5] (mol/L) Time (s) • 1.00 0 • 0.88 200 • 0.78 400 • 0.69 600 • 0.61 800 • 0.54 1000 • 0.48 1200 • 0.43 1400 • 0.38 1600 • 0.34 1800 • 0.30 2000

  18. To find rate • To find rate we have to find the slope at two points We will use the tangent method.

  19. At .90 M the rate is (0.98 M - .76 M) = 0.22 = -5.5 x10-4 (0 – 400) -400

  20. At 0.40M the rate is:(.52 - .31) = 0.22 = -2.7 x 10-4(1000-800) -800

  21. WHAT IT MEANS • Since the rate at twice the concentration is twice as fast the rate law must be.. • Rate = -D[N2O5] = k[N2O5]1 = k[N2O5] Dt • We say this reaction is first order in N2O5 • The only way to determine order is to run the experiment.

  22. The method of Initial Rates • This method requires that a reaction be run several times. • The initial concentrations of the reactants are varied. • The reaction rate is measured bust after the reactants are mixed. • Eliminates the effect of the reverse reaction.

  23. An example • For the reaction BrO3- + 5 Br- + 6H+ 3Br2 + 3 H2O • The general form of the Rate Law is Rate = k[BrO3-]n[Br-]m[H+]p • We use experimental data to determine the values of n,m,and p

  24. Initial concentrations (M) Rate (M/s) BrO3- Br- H+ 0.10 0.10 0.10 8.0 x 10-4 0.20 0.10 0.10 1.6 x 10-3 0.20 0.20 0.10 3.2 x 10-3 0.10 0.10 0.20 3.2 x 10-3 Now we have to see how the rate changes with concentration

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