1. CHEMICAL KINETICS Chapter 13

2. Objectives  Define key terms and concepts.  Predict how temperature, catalysts, concentration, and surface area affect reaction rates.  Calculate the rate of a reaction.  Distinguish between a zero, first, and second order reactions.  Identify the rate determining step of a reaction.

3. Rates of Reaction  Chemical Kinetics  Area of chemistry that studies the rate of reactions.  Reaction Rate  The speed a chemical reaction occurs.  Rate of a reaction may change over time.  May also have a back-reaction occurring later in the reaction that wasn’t occurring at the start.

4. Rates of Reaction  Rate Constant (k)  A constant the relates the reaction rate and the concentration of the reactants.  Holds true for the entire reaction.  Rate Law  Expression relating the rate of a reaction to the rate constant and the concentrations of the reactants.

5. Rate of Reaction  Lead Iodide http://www.youtube.com/watch?feature=endscreen&NR=1&v=RE-dFN7U91M http://www.youtube.com/watch?feature=endscreen&NR=1&v=RE-dFN7U91M Pb(NO 3 ) 2(aq) + 2KI (aq)  2KNO 3(aq) + PbI 2(s)  Iodine Clock http://www.youtube.com/watch?v=LcHKWChwP58 http://www.youtube.com/watch?v=LcHKWChwP58  Reaction of hydrogen peroxide and sulfuric acid with potassium iodide, sodium thiosulfate, and starch.

6. Affects on Rates of Reaction  Collision Theory of Chemical Kinetics  The rate of reaction is directly proportionate to the number of collisions per second.  Temperature http://www.youtube.com/watch?v=bk_GIddjR6c  Collision Orientation  Surface Area http://www.youtube.com/watch?v=FJtwkum_QAY&feature=results_main&playnext=1&list=PLA8CB0B1A42A2EA34

7. Rates of Reaction  Concentration  Increasing/decreasing the concentration of either of the reactants will also increase the reaction rate proportionately.  If the reaction order is 0 for a reactant, the rate is independent of the concentration of that reactant.  In the lab, zinc granules react fairly slowly with dilute hydrochloric acid, but much faster if the acid is concentrated. Zn (s) + HCl (aq)  ZnCl 2(aq) + H 2(g)

8. Affects on Rates of Reaction  Activation Energy (E a )  The minimum amount of energy required to initiate a chemical reaction.  If the products of a reaction are more stable than the reactants, the reaction is exothermic.  If the products of a reaction are less stable than the reactants, the reaction is endothermic.

9. Affects on Rates of Reaction

10. Rates of Reaction  The Reaction Rate is expressed in terms of the decrease in the concentration of reactants as compared to the increase in the concentration of products over time.  For the general reaction aA  bB  Where [A] is the change in concentration of the reactants  [B] is the change in concentration of the products  t is time in seconds Rate = - ∆[A] ∆t Rate = ∆[B] ∆t 1a1a 1b1b

11. Rates of Reaction  For the general reaction aA + bB  cC + dD Rate = - = - == ∆[A] ∆t 1a1a ∆[B] ∆t 1b1b ∆[C] ∆t 1c1c ∆[D] ∆t 1d1d

12. If aluminum is reacting at the rate of 0.015M/second with chlorine gas as illustrated in the following reaction, Al + Cl 2  AlCl 3 at what rate is aluminum chloride being formed?

13. If aluminum is reacting at the rate of 0.015M/second with chlorine gas as illustrated in the following reaction, Al + Cl 2  AlCl 3 at what rate is chlorine gas being consumed?

14. If silver nitrate reacts at a rate of 0.094M/sec in the following reaction AgNO 3 + CaCl 2  Ca(NO 3 ) 2 + AgCl at what rate is calcium chloride being consumed? At what rate are calcium nitrate and silver chloride being produced?

15. What Are Your Questions?

16. Rates of Reaction  For the general reaction aA + bB  cC + dD  Reaction Order  The sum of the powers to which all reactant concentrations appear in the rate law are raised.  Given by adding together x and y in the above reaction.  If for a reaction, x = 1 and y = 2, then overall the reaction order is 3. Rate = k[A] x [B] y

17. Rates of Reactions

18. Rates of Reaction – First Order  The reaction rate and concentration of reactants can be calculated for a reaction at any time during that reaction.  First Order Reactions  Reaction whose rate depends on the reactant concentration raised to the first power.  Rate=k[A] [A] t [A] 0 In = -kt In[A] t = -kt + ln[A] 0

19. Rates of Reaction – First Order

20. The following reaction is first order with a rate constant of 4.80x10 -4 /sec at 45°C. N 2 O 5  NO 2 + O 2 If the initial concentration is 1.65x10 -2 M, what is the concentration after 825sec? How long would it take for the concentration of N 2 O 5 to decrease to 1.00x10 -2 M from the initial value provided?

21. The following reaction is first order with a rate constant of 4.80x10 -4 /sec at 45°C. N 2 O 5  NO 2 + O 2 If the initial concentration is 1.65x10 -2 M, what is the concentration after 600 seconds? How long would it take for the concentration of dinitrogen pentoxide to decrease to 10% of its initial value?

22. Rates of Reaction – Half Life  Half Life  The amount of time it takes for the concentration of a reactant to decrease by half.  Radioactivity t 1/2 = 0.693 k

23. Sulfuryl chloride, SO 2 Cl 2, is a colorless, corrosive liquid whose vapor decomposes in a first- order reaction to sulfur dioxide and chlorine. SO 2 Cl 2(g)  SO 2(g) + Cl 2(g) At 320°C, the rate constant is 2.20x10 -5 /sec. What is the half-life of SO 2 Cl 2 at this temperature? How long would it take for 50% of the SO 2 Cl 2 to decompose? How long for 75%?

24. Rates of Reaction – Second Order  Second Order Reaction  A reaction in which the rate depends on the concentration of 1 reactant raised to the 2 nd power or 2 different reactants, raised to the 1 st power.  Rate=k[A] 2 k[A] 0 t 1/2 = 1 1 = [A] t 1 [A] 0 + kt

25. Rates of Reaction – Second Order

26. Rates of Reaction – Zero Order  Zero Order Reaction  The concentration of the reactants does not affect the rate of the reaction.  Rate = k 2k t 1/2 = [A] 0 [A] t = - kt + [A] 0

27. Rates of Reaction – Zero Order

28. Hydrogen sulfide is oxidized by chlorine in aqueous solution according to the following reaction H 2 S (aq) + Cl 2(aq )  S (s) + 2HCl (aq) If the rate law is experimentally determined as Rate=k[H 2 S][Cl 2 ] What is the reaction order with respect to each of the reactants? What is the overall reaction order?

29. Oxalic acid, H 2 C 2 O 4, is oxidized by permanganate in aqueous solution according to the following reaction 2MnO 4 - (aq) + 5H 2 C 2 O 4(aq) + 6H + (aq)  2Mn 2+ (aq) + 10CO 2(aq) + 8H 2 O (l) If the rate law is experimentally determined as Rate=k[MnO 4 - ][H 2 C 2 O 4 ] What is the reaction order with respect to each of the reactants? What is the overall reaction order?

30. A reaction is zero order with a rate constant of 5.4x10 8 /M sec. If the amount of the reactant is increased, what is the rate of reaction? What is the half-life of this reaction if the concentration of the products is 0.25M?

31. Using the following reaction and initial rates of reaction provided, determine the reaction order for the reactants and the rate constant for the reaction. 2NO 2(g)  2NO (g) + O 2(g) Experimental Trial Initial NO 2 Concentration (M) Initial Rate of Formation of O 2 (mole/L*sec) Exp 10.0107.1x10 -5 Exp 20.02028x10 -5

32. Using the following reaction and initial rates of reaction provided, determine the reaction order for the reactants and the rate constant for the reaction. H 2 O 2(aq) + 3I - (aq) + 2H + (aq)  I 3 - (aq) + 2H 2 O (l) Experimental Trial Initial Concentration (M) Initial Rate (mole/L*sec) H2O2H2O2 I-I- H+H+ Exp 10.010 0.000501.15x10 -6 Exp 20.0200.0100.000502.30x10 -6 Exp 30.0100.0200.000502.30x10 -6 Exp 40.010 0.001001.15x10 -6

33. HO 2 is a highly reactive chemical that can form in the atmosphere. It can decompose into hydrogen peroxide and oxygen according to the following reaction: HO 2(g)  H 2 O 2(g) + O 2(g) Determine the reaction order and rate constant at 25  C using the following data: Trial[HO 2 ] 0 (M)Initial Rate (M/s) 11.1x10 -8 1.7x10 -7 22.5x10 -8 8.8x10 -7 33.4x10 -8 1.6x10 -6 45.0x10 -8 3.5x10 -6

34. What Are Your Questions?

35. The Arrhenius Equation  Arrhenius Equation  Expresses how temperature effects the rate constant of a reaction Where Ea is the activation energy in kJ/moles R is 8.314 J/Kmole T is the temperature in kelvins k = Ae -E a /RT lnk = lnA - E a /RT

36. The rate constant for the formation of hydrogen iodide is 2.7x10 -4 L/mole sec at 600K and 3.5x10 -3 L/mole sec at 650K. What is the activation energy for this reaction? What is the rate constant at 700K? H 2(g) + I 2(g)  2HI (g)

37. Acetaldehyde decomposes when heated as illustrated in the following reaction: CH 3 CHO (g)  CH 4(g) + CO (g) If the rate constant for this decomposition reaction is 1.05x10 -3 L/mole sec at 760K and 2.14x10 -2 L/mole sec at 830K. What is the activation energy for this reaction? What is the rate constant at 880K?

38. What Are Your Questions?

39. Reaction Mechanisms  A balanced chemical reaction only explains the overall reaction process, but not all the individual steps it required to make the reaction happen.  Elementary Steps  Series of simple reactions that represent the progress of an overall reaction at the molecular level.  Reaction Mechanism  The sequence of elementary steps that lead to the products of a reaction.

40. Reaction Mechanisms CH 2 =CH 2 + HCl  CH 3 CH 3 Cl  Intermediates  A compound that appears in the elementary steps of the reaction, but not the overall reaction.

41. Reaction Mechanisms  Need to be aware of the number of molecules reacting in an elementary step to determine the rate order.  Bimolecular Reaction  Involves 2 molecules  Unimolecular Reaction  Involves 1 molecule  Decomposition or dehydration reactions  Termolecular Reactions  Involves 3 molecules

42. Reaction Mechanisms  Rate-Determining Step  The slowest in the sequence of steps in a chemical reaction that leads to the formation of the products.  Back to the Iodine Clock In the first, slow reaction, the triiodide ion is produced H 2 O 2(aq) + 3 I − (aq) + 2 H + (aq) → I 3 − (aq) + 2 H 2 O (l) In the second, fast reaction, triiodide is reconverted to iodide by the thiosulfate. I 3 − (aq) + 2 S 2 O 3 2 − (aq) → 3 I − (aq) + S 4 O 6 2- (aq)

43. Reaction Mechanisms Overall Balanced Reaction 2HC 2 H 3 O 2(aq) + Na 2 CO 3(aq)  2C 2 H 3 O 2 Na (aq) + H 2 O (l) + CO 2(g) Step 1 2HC 2 H 3 O 2(aq) + Na 2 CO 3(aq)  2C 2 H 3 O 2 Na (aq) + H 2 CO 3(aq) Step 2 H 2 CO 3(aq)  H 2 O (l) + CO 2(g)

44. Reaction Mechanisms  When determining the Rate Law  The sum of the elementary steps must give the overall balanced equation for the reaction.  The rate-determining step should show the same rate law as predicted experimentally.

45. Carbon tetrachloride is produced by chlorinating methane or a chlorinated methane molecule. The mechanism for the overall reaction is provided below: Cl 2 2Cl Cl + CHCl 3 HCl + CCl 3 Cl + CCl 3 CCl 4 What is the overall, balanced net reaction? Identify the intermediates.

46. Ozone reacts with nitrogen dioxide according to the following equation: O 3(g) + 2NO 2(g)  O 2(g) + N 2 O 5(g) The proposed steps for this reaction are O 3(g) + NO 2(g)  NO 3(g) + O 2(g) (slow step) NO 3(g) + NO 2(g)  N 2 O 5(g) (fast step) What is the rate law predicted by this mechanism?

47. The iodide ion-catalyzed decomposition of hydrogen peroxide occurs by the following mechanism: H 2 O 2(aq) + I - (aq)  H 2 O (l) + IO - (aq) (slow step) H 2 O 2(aq) + IO - (aq)  H 2 O (l) + O 2(g) + I - (aq) (fast step) What is the rate law predicted by this mechanism?

48. Catalysis  Catalyst  Something that speeds up the rate of a reaction by lower the activation energy required for the reaction.  Is not used up in the reaction.

49. Catalysis  Heterogeneous Catalysis  The reactants and the catalyst are in different phases. Synthesis of Ammonia https://www.youtube.com/watch?v=uMkzxV_y7tY https://www.youtube.com/watch?v=uMkzxV_y7tY Fritz Haber https://www.youtube.com/watch?v=uMkzxV_y7tY https://www.youtube.com/watch?v=uMkzxV_y7tY Catalytic Converters https://www.youtube.com/watch?v=1e9EvrThk1Y https://www.youtube.com/watch?v=1e9EvrThk1Y  Homogeneous Catalysis  When the reactants and catalyst are in the same phase, typically liquid. https://www.youtube.com/watch?v=MrUv1KjaNXY https://www.youtube.com/watch?v=MrUv1KjaNXY

50. Catalysis

51. What Are Your Questions?