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Chapter 14

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  1. Chapter 14 Chemical Kinetics

  2. Reaction Rate • The rate of a chemical reaction is measured as the decrease in the concentration of a reactant or the increase in the concentration of a product in a unit of time.

  3. Reaction Rate • The rate of a chemical reaction is measured as the decrease in the concentration of a reactant or the increase in the concentration of a product in a unit of time. ∆[ ] Rate = ∆time

  4. Reaction Rate • The rate of a chemical reaction is measured as the decrease in the concentration of a reactant or the increase in the concentration of a product in a unit of time. ∆[ ] What units would we use for rate? Rate = ∆time

  5. Reaction Rate • The rate of a chemical reaction is measured as the decrease in the concentration of a reactant or the increase in the concentration of a product in a unit of time. ∆[ ] Rate = ∆time 2H2O2(aq) → 2H2O(l) + O2(g)

  6. Reaction Rate • The rate of a chemical reaction is measured as the decrease in the concentration of a reactant or the increase in the concentration of a product in a unit of time. ∆[ ] Rate = ∆time 2H2O2(aq) → 2H2O(l) + O2(g) How could the rate be expressed for this reaction in terms of H2O2?

  7. 2H2O2(aq) → 2H2O(l) + O2(g)

  8. 2H2O2(aq) → 2H2O(l) + O2(g)

  9. 2H2O2(aq) → 2H2O(l) + O2(g) What is the rate of the reaction from 0s to 2.16 x 104s?

  10. 2H2O2(aq) → 2H2O(l) + O2(g) What is the average rate of appearance of O2 from 0s to 2.16 x 104s? 1.16 x 10-5 mol O2 L-1 s-1

  11. General Rate of Reaction a A + b B → c C + d D Rate of reaction = rate of disappearance of reactants or Rate of reaction = rate of appearance (formation) of products We can use the coefficients in the equation to compare the reaction rates for all the substances in the reaction.

  12. Δ[Fe2+] 0.0010 M Rate of formation of Fe2+= = = 2.6 x 10-5 M s-1 Δt 38.5 s 15-1 The Rate of a Chemical Reaction • Rate is change of concentration with time. 2 Fe3+(aq) + Sn2+(aq) → 2 Fe2+(aq) + Sn4+(aq) t = 38.5 s [Fe2+] = 0.0010 M ∆t = 38.5 s ∆[Fe2+] = (0.0010 – 0) M

  13. Rates of Chemical Reaction 2 Fe3+(aq) + Sn2+(aq)→ 2 Fe2+(aq) + Sn4+(aq) Rate of formation of Fe2+ = 2.6 x 10-5 mol L-1 s-1 What is the rate of formation of Sn4+? 1.3 x 10-5 mol Sn4+ L-1 s-1 What is the rate of disappearance of Fe3+? 2.6 x 10-5 mol Fe3+ L-1 s-1

  14. What does the slope of the line represent?

  15. Δ[H2O2] Rate = 1.7 x 10-3 M s-1 = Δt Δ[H2O2] = (1.7 x 10-3 M s-1) (∆t) ∆[H2O2]= (1.7 x 10-3 M s-1)(100 s) = 0.17M [H2O2]100 s = 2.32 M - 0.17 M What is the concentration at 100s for the reaction: 2H2O2(aq) → 2H2O(l) + O2(g)?Given: [H2O2]i = 2.32 M = 2.15 M

  16. What does it mean when the rate of a reaction reaches zero? • For a normal reaction it means that one or more of the reactants are used up and the reaction has stopped. • For a reversible reaction it means that the reaction has reached equilibrium.

  17. Factors Affecting Reaction Rates • The nature of the reacting substances.

  18. Factors Affecting Reaction Rates 2. The state of subdivision of the reacting substances (surface area).

  19. Lycopodium Powder

  20. Factors Affecting Reaction Rates 3. The temperature of the reacting substances.

  21. Factors Affecting Reaction Rates 4. The concentration of the reacting substances. (Except in zero order reactions) Air (21% oxygen) 100% oxygen

  22. Factors Affecting Reaction Rates • The presence of a catalyst. Catalysts speed up reactions but are left unchanged by the reaction.

  23. The Rate Law a A + b B …. → g G + h H …. Rate = k [A]m[B]n …. Rate constant = k (k is constant for a particular reaction at a specific temperature) Order of A = m Order of B = n Overall order of reaction = m + n + ….

  24. Temperature and Rate • Generally, as temperature increases, so does the reaction rate. • This is because k is temperature dependent. • Therefore the temperature dependence of reaction rates is contained in the temperature dependence of the rate constant.

  25. Temperature dependence of “k” . . . . .

  26. Concentration and Rate Summary • After finding the trials to compare: • A reactant is zero order if the change in concentration of that reactant produces no effect on the rate. • A reaction is first order if doubling the concentration of that reactantcauses the rate to double. • A reactant is nth order if doubling the concentration of that reactantcauses an 2n increase in rate. • Note that the rate constant does not depend on concentration.

  27. Use the data provided to write the rate law and indicate the order of the reaction with respect to HgCl2 and C2O42- and also the overall order of the reaction.

  28. First determine the order of HgCl2

  29. Next determine the order of C2O42-

  30. Now write the rate law and determine the order of the reaction.

  31. Calculate the rate constant “k” and its units. Initial rate of disappearance HgCl2 mol L-1 min-1

  32. What is the average rate of disappearance of C2O42- in trial 1? Initial rate of disappearance HgCl2 mol L-1 min-1

  33. Use the data provided to write the rate law and indicate the order of the reaction with respect to NO2 and CO (support your answers). Also give the overall order of the reaction.

  34. Calculate the rate constant “k” and its units.

  35. What is the average rate of disappearance of CO in trial 2?

  36. How do we make these charts? Initial rate of disappearance HgCl2 mol L-1 min-1 • Rates can be measured experimentally using a variety techniques: • moniter pH changes • Titrations • Change in volume or mass (gas production) • Basically we can use any method to follow a reaction that produces a measurable change.

  37. How do we make these charts? Initial rate of disappearance HgCl2 mol L-1 min-1 One important method involves the spectroscopic determination of concentration through Beer’s Law.

  38. Using Beer’s Law to Determine [ ] vs. time. • For each trial, the reactants are mixed and the reaction mixture is transferred into a test tube or cuvette. • Without any delay, the reaction vessel is placed into a spectrophotometer. The absorbance data is then collected at the wavelength of maximum absorbance as a function of time. • This absorbance data is then converted to concentration data using Beer’s Law: A = ɛ l c

  39. Fe(s)+CuSO4(aq)→Fe2SO4(aq)+Cu(s) • The solution gradually gets paler as the concentration of copper sulfate decreases and the concentration of iron sulfate increases. Concentration of copper sulfate solution

  40. Using Beer’s Law to Determine [ ] vs. time. • A graph of concentration vs. time can be prepared and then used to experimentally determine the rate.

  41. What does this tangent allow us to measure?

  42. Half Life of a First Order Reaction • Half-life is the time required to convert one half of a reactant to product. • For first-order reactions, half-life is often used as a representation for the rate constant. • This is because the half-life of a first-order reaction and the rate constant are inversely proportional, and the half-life is independent of concentration.

  43. Radioactivity Radioactive decay is the spontaneous breakdown of unstable atoms into more stable atoms with the simultaneous emission of particles and rays. Radioactive decay occurs at a constant rate that is a first order process.

  44. Radioactivity and Half - Life The half-life of carbon-14 is 5730 years. How old is a bone that has about 12.5% of the carbon-14 that a living organism would have in it?

  45. Carbon Dating

  46. Big Question • How can we experimentally determine the order of a reaction?

  47. Make “3” Graphs • In order to determine order of reactant, A. We must collect data consisting of concentration versus time. • One common way to determine concentration vs. time data is through the use of a spectrophotometer.

  48. Make “3” Graphs • We then use the data to make three graphs. • [A] versus t • ln [A] versus t • 1 / [A] versus t • By examining these graphs we can determine the order of the reaction with respect to a particular reactant and determine the rate constant.