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Reaction Rates

Reaction Rates. --insert obligatory “slow children” joke. Definition. A reaction rate is the rate at which a reaction takes place. Definition. A reaction rate is the rate at which a reaction takes place. Helpful, huh?.

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Reaction Rates

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  1. Reaction Rates --insert obligatory “slow children” joke

  2. Definition • A reaction rate is the rate at which a reaction takes place.

  3. Definition • A reaction rate is the rate at which a reaction takes place. • Helpful, huh?

  4. Usually, a rate is expressed in moles/sec, or more commonly, M/s (mol/L/s, or mol L-1 s-1) • Q: Moles of what? • A: Anything convenient.

  5. Rates decrease with time

  6. But over short times, you can measure an initial rate. Reactant 1 is used Reactant 2 is used faster Product is produced

  7. For the general reaction aA + bBcC + dD --at a constant volume • Rate = -D[A] = -D [B] =D [C] =D [D] aDt bDt cDt dDt (notice that the reactants are disappearing, while products are being formed)

  8. For the general reaction aA + bBcC + dD --at a constant volume • Rate = -D[A] = -D [B] =D [C] =D [D] aDt bDt cDt dDt (notice, also, that the rate is independent of which substance is used to measure it.)

  9. N2 + 3H22NH3 • If, under some conditions, 2.40 moles of NH3 are produced each second, what is the rate of use of N2 and H2?

  10. N2 + 3H22NH3 • If, under some conditions, 2.40 moles of NH3 are produced each second, what is the rate of use of N2 and H2? • 2.40 mol NH3/s x -1 N2/2 NH3 =-1.20 mol N2/s

  11. N2 + 3H22NH3 • If, under some conditions, 2.40 moles of NH3 are produced each second, what is the rate of use of N2 and H2? • 2.40 mol NH3/s x -1 N2/2 NH3 =-1.20 mol N2/s • 2.40 mol NH3/s x -3 H2/2 NH3 =-3.60 mol H2/s

  12. N2 + 3H22NH3 • If, under some conditions, 2.40 moles of NH3 are produced each second, what is the rate of use of N2 and H2? • You might say that the rate of the reaction is 1.20 mole/s as written, but we are usually interested in the rate of change of a substance

  13. How can you speed up a reaction?

  14. How can you speed up a reaction? • --Heat it up. • --Crush, grind or powder a solid reactant. • --Increase pressure of a gaseous reactant • --Increase concentrations of aqueous reactants • --Add a catalyst (if known) • (Stir or shake to bring reactants together.)

  15. Now, for convenience, measure the easiest rate: Changes in… • Color -- pressure • Conductivity --Mass • Gas volume --pH Whatever changes can be measured

  16. How would you speed up… • Hydrochloric acid acts on tin metal to form hydrogen gas and aqueous tin (II) chloride

  17. How would you speed up… • Hydrochloric acid acts on tin metal to form hydrogen gas and aqueous tin (II) chloride • Increase concentration of HCl • Powder the tin • Heat the reactants

  18. How would you measure the rate of… • Hydrochloric acid acts on tin metal to form hydrogen gas and aqueous tin (II) chloride

  19. How would you measure the rate of… • Hydrochloric acid acts on tin metal to form hydrogen gas and aqueous tin (II) chloride • Rate of change in H2 volume • loss of mass of tin, appearance of tin (II), increase in pH, and loss of conductivity are all harder

  20. What happens to the rate after (mumble, mumble) seconds? • Hydrochloric acid acts on tin metal to form hydrogen gas and aqueous tin (II) chloride

  21. What happens to the rate after (mumble, mumble) seconds? • Hydrochloric acid acts on tin metal to form hydrogen gas and aqueous tin (II) chloride • It slows down. • Why?

  22. What happens to the rate after (mumble, mumble) seconds? • Hydrochloric acid acts on tin metal to form hydrogen gas and aqueous tin (II) chloride • It slows down. • Why? • --[H+] decreases as reaction proceeds

  23. Rates decrease with time

  24. We measure the initial rates of reactions because —there in the first instant— that is the only time when we know all of the concentrations!

  25. Reaction rates are affected by • -Concentrations • -Temperature • -etc.

  26. Production of product Product appearance in a reaction Product appearance in a reaction at a higher temperature Product appearance in a reaction at a lower concentration

  27. So– we choose a temperature, control that temperature and vary concentrations to discover the effect that each concentration has on the rate.

  28. Rate Laws • Rate= k [A]x[B]y • x and y are not related to the balancing of the reaction! (They are experimentally determined.)

  29. You will be asked to… • Write rate laws from empirical data. • Multiple trials show the exponents • Calculate rate constants from the laws you write. • Watch your units! Units vary for different laws • Calculate a rate from a known rate law, constant, and concentrations. • Use M/s or mol/s (mol L-1s-1 or mol s-1 )

  30. The rate of decay of 3H (tritium) • [3H] (M) Rate of Decay (nM/sec) • .01 .018 • .02 .036 • .03 .054 • .04 .072

  31. The rate is directly related to [3H]

  32. This is a first order reaction. • The rate law is Rate=k[3H] where k=1.78 x 10-9 /s

  33. A first order reaction has a rate of .024 M/s when [A] = .20 M • What is the value of the rate constant, k?

  34. A first order reaction has a rate of .024 M/s when [A] = .20 M • What is the value of the rate constant, k? • What are the units on the rate constant, k?

  35. A first order reaction has a rate of .024 M/s when [A] = .20 M • What is the value of the rate constant, k? • What are the units on the rate constant, k? • What is the rate of the reaction when [A]=.30M?

  36. A first order reaction has k=1.6 x 10-4 /s • What is the rate when [A]= .050M?

  37. A first order reaction has k=1.6 x 10-4 /s • What is the rate when [A]= .050M? • At what [A] will the rate be 5.4 x 10-6 M/s?

  38. For example, the rate of IO3- reduction by HSO3- • [IO3-](mM) Rate of Reaction (M/s) • .3 .015 • .6 .060 • .9 .135 • 1.2 .240

  39. When you double the concentration—the rate quadruples • When you triple the original concentration, the rate increases by a factor of 9  4=22  9=32

  40. This is a second order reaction Rate = k [IO3-]2 where k=167/Ms

  41. Concentrations of reactants fall • In a first order reaction, the concentration of reactant decreases with time • The decrease is faster with a larger k

  42. Decrease of reactant Decrease of reactant when k is larger

  43. For a first order reaction… • The amount of reactant left, [A], is [A]= [Ao]e-kt —where [Ao] is the original conc. (at time 0) e is the base of the natural logs k is the rate constant and t is time

  44. For a first order reaction… • This equation is usually used as ln[A]- ln[Ao]= -kt —where [Ao] is the original conc. (at time 0) e is the base of the natural logs k is the rate constant and t is time

  45. A first order reaction takes [A] from .080M to .055M in 85 s. • Write the rate law.

  46. A first order reaction takes [A] from .080M to .055M in 85 s. • Write the rate law. • What is k?

  47. A first order reaction takes [A] from .080M to .055M in 85 s. • Write the rate law. • What is k? • What is [A] after 120s?

  48. A first order reaction takes [A] from .080M to .055M in 85 s. • Write the rate law. • What is k? • What is [A] after 120s? • At what time will [A]=.020M?

  49. A first order reaction takes [A] from .080M to .055M in 85 s. • Write the rate law. • What is k? • What is [A] after 120s? • At what time will [A]=.020M? • What is the half life of this reaction?

  50. For a second order reaction: • The relationship between concentrations, k and t is 1 - 1 = kt [A]t [A]o

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