Chapter 15

1. Chapter 15 Kinetics

2. Overview • Rates of Chemical Reaction • Pathways or Mechanisms of Reaction

3. Factors affecting Rate • [Reactants] recall, [ molarity] • Temp. • Surface Area • Catalyst define catalyst vs inhibitor

4. Average Rate, defined R  P rate = Δ [R] / Δ t Rate of disappearance (-) of R with respect to time or May monitor appearance (+) of product

5. Avg Rate, Example N2O5 → 2 NO2 + ½ O2 rate = - Δ [N2O5] / Δ t orrate = + ½ Δ [NO2] / Δ t or rate= +2 Δ [O2] / Δ t

6. Graphing Rate Data • You may omit graphical computations; be able to read graphs - see page 672 • Compare avg rate vs instantaneous rateex. 15.1 p.674

7. Rate Law aA + bB cC + dD • Form rate = k [A]m[B]n • Derived from experiment • m,n are coeffs of slow step • Exponent value ~ “order” (examples) • k is “rate constant”

8. Units for k, rate constant M (1- overall order) s -1 Where: M, molarity s, seconds

9. Rate Law Determination Method of “Initial Rates” Instantaneous rate at the start of the reaction(t=o) • RE:lab • Example: 2 NO + Cl2 2NOCl data p.679 We will not use graphical methods p. 687-689

10. Integrated Rate Law • Finding [R]t at a later time, t, if [R]0 is initially known ln [R]t / [R]0 = -ktfor 1st order Solve for ln [R]t : ln [R]t = -kt + ln [R]0 eqn. for a straight line graphically determined

11. Computations Integrated Rate Law We will consider 1st order only-ln [R]t / [R]0 = - kt (think algebraically!) -to find k if [R]t and [R]0 are measured -to find [R]0 if k and [R]t are known -to find the fraction, [R]t / [R]0 , if k is known

12. Half-Life of Reaction, t1/2 • Defined • Zero order: t ½ = [A]0 / 2k • Ist order: t ½ = 0.693/ k(radioactive decay is a 1st order process) • 2nd order: t ½ = 1/ k[A]0 See ex. 15.9 p.691

13. Radioactive Decaya 1st order process Recall, t ½ = 0.693/ k and ln [R]t / [R]0 = -kt If t ½ and [R]0 are given, can find k first and then use the eqn to find [R]t Alternative method- (amount remaining) = (org amount) x ( 1/2n) where n = number of half lives elapsed

14. Collision Theory • Rate depends on: # of collisions % effective (% incr, as temp incr ) • Ea defined

15. Collision Theory, contd • Transition state • Activated complex • Profiles draw & compare ΔHes and Eas for exo and endo cases • Notes: ΔH = Eafw – Eabckrate increases as T increases for both fw and bck rxns

16. Arrhenius Equation k = A e –Ea / RT Now, lnk = –Ea / RT + lnA lnk = –Ea / R (1/T) + ln A straight line; find Ea graphically from slope Or ln k2/k1 = Ea / R { 1/ T1 - 1/ T2 } where T is Kelvin (derivation in text, p. 698)

17. Mechanisms • Steps; path • Steps are reversible • “slow step” is rate determining if rate = k [A]m[B]n m,n are coeffs of slow step in mechanism • Terms: “elementary reaction” is a step unimolecular, bimolecular, etc

18. Mechanisms, contd. • Intermediates formed/ consumed • Catalysts regeneratedmechanism are proposals ex. 15.14 p. 709

19. Catalysts • Lower Ea • New mechanism • New slow step • Homogeneous catalysis • Heterogeneous catalysissee revised energy profiles

20. Enzymes • Biological catalysts, E • Lock and key model • Active site • Substrate, S

21. Lock and Key Model E active site Substrate

22. Factors Affecting Enzyme Activity • [S] • [E] • pH • T

23. Example One r = k[S] until max achieved, then rxn is zero order, independent of [S] Where [E] = constant rate max rate [S] When Ea is low, @ max rate E, sites have become saturated

24. Example tworate = k [E] Where [S] = constant and excess rate [E]

25. Inhibitors • Destroy lock and key “fit” • Active site is no longer “active” • Example, lipitor

26. End Ch. 15 • Homework • Exam # 1 Ch. 5 & 15