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Chapter 12

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

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Chapter 12

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  1. Chapter 12 Kinetics

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

  3. Reaction Rate • Rate = Conc. of A at t2 -Conc. of A at t1t2- t1 • Rate =D[A]Dt • Change in _____________________________per unit time • For this reaction • N2 + 3H2 2NH3

  4. As the reaction progresses the concentration H2 goes _________________ Concentration [H2] Time

  5. As the reaction progresses the concentration N2 goes down ________________________________ Concentration [N2] [H2] Time

  6. As the reaction progresses the concentration NH3 goes ___________________. Concentration [N2] [H2] [NH3] Time

  7. Calculating Rates • Average rates are taken over ___________ intervals • ___________________________ 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. Concentration D[H2] Dt Time

  9. Concentration D[H2] D t Time

  10. __________________________ • 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] DtDtDt

  11. Rate Laws • Reactions are ________________________. • 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 ________________. • This is called the _______________________________________.

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

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

  14. 2 NO2 2NO + O2 • The rate of _________________________ 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 • ________________ Rate law - describes how rate depends on concentration. • _______________ 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 __________________________.

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

  17. [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 Now graph the data

  18. 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 (.98 - .76) = 0.22 =- 5.5x 10 -4 (0-400) -400

  20. At .40 M the rate is (.52 - .31) = 0.22 =- 2.7 x 10 -4 (1000-1800) -800

  21. 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 _______________ 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 ______________________ times. • The ____________________ concentrations of the reactants are varied. • The reaction rate is measured 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 __________________________

  24. Initial concentrations (M) Rate (M/s) Now we have to see how the rate changes with concentration 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

  25. Integrated Rate Law • Expresses the reaction concentration as a function of ______________________. • Form of the equation depends on the order of the rate law (differential). • Changes Rate = D[A]nDt • We will only work with ________________________

  26. First Order • For the reaction 2N2O5 4NO2 + O2 • We found the Rate = k[N2O5]1 • If concentration doubles _____________ doubles. • If we integrate this equation with respect to time we get the Integrated Rate Law • ln[N2O5] = - kt + ln[N2O5]0 • [N2O5]0 is the ______________________ concentration.

  27. First Order • General form Rate = D[A] / Dt = k[A] • ln[A] = - kt + ln[A]0 • In the form y = mx + b • y = ln[A] m = -k • x = t b = ln[A]0 • A graph of ln[A] vs time is a _____________________________.

  28. First Order • By getting the straight line you can ______________________ it is first order • Often expressed in a ratio

  29. Half Life • The time required to reach______________the original concentration. • If the reaction is first order • [A] = [A]0/2 when t = t1/2

  30. Half Life • The time required to reach half the original concentration. • If the reaction is first order • [A] = [A]0/2 when t = t1/2 • ln(2) = kt1/2

  31. Second Order • Rate = -D[A] / Dt = k[A]2 • integrated rate law • 1/[A] = kt + 1/[A]0 • y= 1/[A] m = k • x= t b = 1/[A]0 • A ____________________________if 1/[A] vs t is graphed • Knowing k and [A]0 you can calculate [A] at any time t

  32. Second Order Half Life • [A] = [A]0 /2 at t = t1/2

  33. Zero Order Rate Law • Rate = k[A]0 = k • Rate does ________________________ change with concentration. • Integrated [A] = -kt + [A]0 • When [A] = [A]0 /2 t = t1/2 • t1/2 = [A]0 /2k

  34. Zero Order Rate Law • Most often when reaction happens on a __________________ because the surface area stays constant. • Also applies to ____________________ chemistry.

  35. Concentration Time

  36. Concentration k = D[A]/Dt D[A] Dt Time

  37. Reaction Mechanisms • The series of _________________ that actually occur in a chemical reaction. • _________________________ can tell us something about the mechanism • A _____________________ equation does not tell us how the reactants become products.

  38. Example of a Reaction Mechanism – H2+ Cl2 2HCl

  39. Reaction Mechanisms • 2NO2 + F2 2NO2F • Rate = k[NO2][F2] • The proposed mechanism is • NO2 + F2 NO2F + F (slow) • F + NO2 NO2F (fast) • F is called an ___________________It is formed then consumed in the reaction

  40. Reaction Mechanisms • Each of the two reactions is called an _________________________________. • The rate for a reaction can be written from its ______________________________. • Molecularity is the number of ___________ that must come together.

  41. Unimolecular step involves ___________ molecule - Rate is first order. • Bimolecular step - requires ___________ molecules - Rate is second order • Trimolecular step- requires ___________ molecules - Rate is third order • Termolecular steps are almost never heard of because the chances of three molecules coming into contact at the same time are miniscule.

  42. A products Rate = k[A] • A+A products Rate= k[A]2 • 2A products Rate= k[A]2 • A+B products Rate= k[A][B] • A+A+B Products Rate= k[A]2[B] • 2A+B Products Rate= k[A]2[B] • A+B+C Products Rate= k[A][B][C]

  43. How to get rid of intermediates • Use the reactions that form them • If the reactions are fast and irreversible - the concentration of the intermediate is based on ________________________. • If it is formed by a reversible reaction set the rates ______________________ to each other.

  44. Formed in reversible reactions • 2 NO + O2 2 NO2 • Mechanism • 2 NO N2O2 (fast) • N2O2 + O2 2 NO2 (slow) • rate = k2[N2O2][O2] • k1[NO]2 = k-1[N2O2] • rate = k2 (k1/ k-1)[NO]2[O2]=k[NO]2[O2]

  45. Formed in fast reactions • 2 IBr I2+ Br2 • Mechanism • IBr I + Br (fast) • IBr + Br I + Br2 (slow) • I + I I2 (fast) • Rate = k[IBr][Br] but [Br]= [IBr] • Rate = k[IBr][IBr] = k[IBr]2

  46. Collision theory • Molecules must collide to react. • Concentration affects rates because collisions are more likely. • Must collide hard enough. • Temperature and rate are related. • Only a small number of collisions produce reactions.

  47. Potential Energy Reactants Products Reaction Coordinate

  48. Potential Energy Activation Energy Ea Reactants Products Reaction Coordinate

  49. Activated complex Potential Energy Reactants Products Reaction Coordinate

  50. Potential Energy } Reactants DE Products Reaction Coordinate

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