Chemical Kinetics CHAPTER 14 Chemistry: The Molecular Nature of Matter, 6 th edition

1. Chemical Kinetics CHAPTER 14 Chemistry: The Molecular Nature of Matter, 6th edition By Jesperson, Brady, & Hyslop

2. CHAPTER 14 Chemical Kinetics • Learning Objectives: • Factors Affecting Reaction Rate: • Concentration • State • Surface Area • Temperature • Catalyst • Collision Theory of Reactions and Effective Collisions • Determining Reaction Order and Rate Law from Data • Integrated Rate Laws • Rate Law  Concentration vs Rate • Integrated Rate Law  Concentration vs Time • Units of Rate Constant and Overall Reaction Order • Half Life vs Rate Constant (1st Order) • Arrhenius Equation • Mechanisms and Rate Laws • Catalysts Jesperson, Brady, Hyslop. Chemistry: The Molecular Nature of Matter, 6E

3. CHAPTER 14 Chemical Kinetics Lecture Road Map: Factors that affect reaction rates Measuring rates of reactions Rate Laws Collision Theory Transition State Theory& Activation Energies Mechanisms Catalysts Jesperson, Brady, Hyslop. Chemistry: The Molecular Nature of Matter, 6E

4. CHAPTER 14 Chemical Kinetics Mechanisms of Reactions Jesperson, Brady, Hyslop. Chemistry: The Molecular Nature of Matter, 6E

5. Mechanisms Overall vs Individual Steps Sometimes rate law has simple form • N2O5 NO2 + NO3 • NO2 + NO3 N2O5 But others are complex • H2 + Br2 2 HBr Jesperson, Brady, Hyslop. Chemistry: The Molecular Nature of Matter, 6E

6. Mechanisms Overall vs Individual Steps Some reactions occur in a single step, as written Others involve a sequence of steps • Reaction Mechanism • Entire sequence of steps • Elementary Process • Each individual step in mechanism • Single step that occurs as written Jesperson, Brady, Hyslop. Chemistry: The Molecular Nature of Matter, 6E

7. Mechanisms Overall vs Individual Steps • Exponents in rate law for elementary process are equal to coefficients of reactants in balanced chemical equation for that elementary process • Rate laws for elementary processes are directly related to stoichiometry • Number of molecules that participate in elementary process defines molecularity of step Jesperson, Brady, Hyslop. Chemistry: The Molecular Nature of Matter, 6E

8. Mechanisms Unimolecular Process • Only one molecule as reactant • H3C—NC  H3C—CN • Rate = k[CH3NC] • 1st order overall • As number of molecules increases, number that rearrange in given time interval increases proportionally Jesperson, Brady, Hyslop. Chemistry: The Molecular Nature of Matter, 6E

9. Mechanisms Bimolecular Process • Elementary step with two reactants • NO(g) + O3(g)  NO2(g) + O2(g) • Rate = k[NO][O3] • 2ndorder overall • From collision theory: • If[A]doubles, number of collisions between A and B will double • If [B] doubles, number of collisions between A and B will double • Thus, process is 1storder in A, 1storder in B, and 2ndorder overall Jesperson, Brady, Hyslop. Chemistry: The Molecular Nature of Matter, 6E

10. Mechanisms Termolecular Process • Elementary reaction with three molecules • Extremely rare • Why? • Very low probability that three molecules will collide simultaneously • 3rd order overall Jesperson, Brady, Hyslop. Chemistry: The Molecular Nature of Matter, 6E

11. Mechanisms Elementary Processes • Significance of elementary steps: • If we know that reaction is elementary step • Then we know its rate law Jesperson, Brady, Hyslop. Chemistry: The Molecular Nature of Matter, 6E

12. Mechanisms Multi-step Mechanisms • Contains two or more steps to yield net reaction • Elementary processes in multi-step mechanism must always add up to give chemical equation of overall process • Any mechanism we proposemust be consistent with experimentally observed rate law • Intermediate = species which are formed in one step and used up in subsequent steps • Species which are neither reactant nor product in overall reaction • Mechanisms may involve one or more intermediates Jesperson, Brady, Hyslop. Chemistry: The Molecular Nature of Matter, 6E

13. Mechanisms Example The net reaction is: NO2(g) + CO(g)  NO(g) + CO2(g) The proposed mechanism is: NO2(g) + NO2(g)  NO3(g) + NO(g) NO3(g) + CO(g)  NO2(g) + CO2(g) 2NO2(g) + NO3(g) + CO(g)  NO2(g) + NO3(g) + NO(g) + CO2(g) or NO2(g) + CO(g)  NO(g) + CO2(g) 1 Jesperson, Brady, Hyslop. Chemistry: The Molecular Nature of Matter, 6E

14. Mechanisms Rate Determining Step • If process follows sequence of steps, slow step determines rate = rate determining step. • Think of an assembly line • Fast earlier steps may cause intermediates to pile up • Fast later steps may have to wait for slower initial steps • Rate-determining step governs rate law for overall reaction • Can only measure rate up to rate determining step Jesperson, Brady, Hyslop. Chemistry: The Molecular Nature of Matter, 6E

15. Mechanisms Example: Rate Determining Step (CH3)3CCl(aq) + OH–(aq)  (CH3)3COH(aq) + Cl–(aq) chlorotrimethylmethanetrimethylmethanol • Observed rate = k[(CH3)3CCl] • If reaction was elementary • Rate would depend on both reactants • Frequency of collisions depends on both concentrations • Mechanism is more complex than single step • What is mechanism? • Evidence that it is a two step process Jesperson, Brady, Hyslop. Chemistry: The Molecular Nature of Matter, 6E

16. Mechanisms Rate Determining Step as Initial Step Step 1: (CH3)3CCl(aq)  (CH3)3C+(aq) + Cl–(aq) (slow) Step 2: (CH3)3C+(aq) + OH–(aq)  (CH3)3COH(aq) (fast) • Two steps each at different rates • Each step in multiple step mechanism is elementary process, so • Has its own rate constant and its own rate law • Hence only for each step can we write rate law directly • Observed rate law says that step 1 is very slow compared to step 2 • In this case step 1 is rate determining • Overall rate = k1[(CH3)3CCl] Jesperson, Brady, Hyslop. Chemistry: The Molecular Nature of Matter, 6E

17. Mechanisms Mechanisms with Fast Initial Step 1st step involves fast, reversible reaction Ex.Decomposition of ozone (No catalysts) Net reaction: 2O3(g)  3O2(g) Proposed mechanism: O3(g)  O2(g) + O(g) (fast) O(g) + O3(g)  2O2(g) (slow) Jesperson, Brady, Hyslop. Chemistry: The Molecular Nature of Matter, 6E

18. Mechanisms Is the Mechanism Rate Law Consistent? • Rate of formation of O2 = Rate of reaction 2 = k2[O][O3] • But O is intermediate • Need rate law in terms of reactants and products • and possibly catalysts • Rate (forward) = kf[O3] • Rate (reverse) = kr[O2][O] • When step 1 comes to equilibrium • Rate (forward) = Rate (reverse) • kf[O3] = kr[O2][O] Jesperson, Brady, Hyslop. Chemistry: The Molecular Nature of Matter, 6E

19. Mechanisms Is the Mechanism Rate Law Consistent? • Solving this for intermediate O gives: • Substitution into rate law for step 2 gives: • Rate of reaction 2 = k2[O][O3] = • where • This is observed rate law • Yes, mechanism consistent Jesperson, Brady, Hyslop. Chemistry: The Molecular Nature of Matter, 6E

20. Group Problem The reaction mechanism that has been proposed for the decomposition of H2O2 is • H2O2 + I–→ H2O + IO– (slow) • H2O2 + IO–→ H2O + O2 + I– (fast) What is the expected rate law? First step is slow so the rate determining step defines the rate law rate=k [H2O2][I–] Jesperson, Brady, Hyslop. Chemistry: The Molecular Nature of Matter, 6E

21. Group Problem The reaction: A + 3B → D + F was studied and the following mechanism was finally determined: • A + BC (fast) • C + B→D + E (slow) • E + B → F (very fast) What is the expected rate law? Rate Step 2=k2[C][B] Rate forward = kf[A][B] Rate reverse = kr[C] kf[A][B] = kr[C] [C]= kf[A][B]/kr Rate = kobs[A][B]2 Jesperson, Brady, Hyslop. Chemistry: The Molecular Nature of Matter, 6E

22. CHAPTER 14 Chemical Kinetics Catalysts Jesperson, Brady, Hyslop. Chemistry: The Molecular Nature of Matter, 6E

23. Catalyst Definition • Substance that changes rate of chemical reaction without itself being used up • Speeds up reaction, but not consumed by reaction • Appears in mechanism, but not in overall reaction • Does not undergo permanent chemical change • Regenerated at end of reaction mechanism • May appear in rate law • May be heterogeneous or homogeneous Jesperson, Brady, Hyslop. Chemistry: The Molecular Nature of Matter, 6E

24. Catalyst Activation Energy • By providing alternate mechanism • One with lower Ea • Because Ea lower, more reactants and collisions have minimum KE, so reaction proceeds faster Jesperson, Brady, Hyslop. Chemistry: The Molecular Nature of Matter, 6E

25. Catalyst Activation Energy Jesperson, Brady, Hyslop. Chemistry: The Molecular Nature of Matter, 6E

26. Catalyst Homogeneous Catalyst • Same phase as reactants Consider : S(g) + O2(g) + H2O(g)  H2SO4(g) S(g) + O2(g)  SO2(g) NO2(g) + SO2(g)  NO(g) + SO3(g) Catalytic pathway SO3(g) + H2O(g)  H2SO4(g) NO(g) + ½O2(g) NO2(g) Regeneration of catalyst  Net: S(g) + O2(g) + H2O(g)  H2SO4(g) • What is Catalyst? • Reactant (used up) in early step • Product (regenerated) in later step • Which are Intermediates? Jesperson, Brady, Hyslop. Chemistry: The Molecular Nature of Matter, 6E

27. Catalyst Heterogeneous Catalyst • Exists in separate phase from reactants • Usually a solid • Many industrial catalysts are heterogeneous • Reaction takes place on solid catalyst Ex.3H2(g) + N2(g) 2NH3(g) Jesperson, Brady, Hyslop. Chemistry: The Molecular Nature of Matter, 6E

28. Catalyst Heterogeneous Catalyst H2 and N2 approach Fe catalyst H2 and N2 bind to Fe & bonds break N—H bonds forming N—H bonds forming NH3 formation complete NH3 dissociates Jesperson, Brady, Hyslop. Chemistry: The Molecular Nature of Matter, 6E

29. Catalyst Enzymes: Superoxide Dismutase Miller, Anne-Frances. “Fe Superoxide Dismutase” Handbook of Metalloproteins. John Wiley & Sons, Ltd, Chinchester, 2001 Rodrigues, J. V; Abreu, I. A.; Cabelli, D; Teixeira, M. Biochemistry 2006, 45, 9266-9278.