Chemical Kinetics Part 3: Reaction Mechanisms

1. Chemical KineticsPart 3:ReactionMechanisms Chapter 13

2. Reaction Mechanisms • The balanced chemical equation provides information about the beginning and end of reaction. • 2A + B→ C • This is an overallrxn; it doesn’t tell you how the rxn occurs. • The reaction mechanism gives the path of the reaction. (or the how) • Mechanisms provide a very detailed picture of which bonds are broken and formed during the course of a reaction. • So it tells you the paths or steps in a rxn.

3. Elementary Steps • Elementary step: any process that occurs in a single step. • Molecularity: the number of reactant molecules present in an elementary step. • Unimolecular: one molecule in the elementary step • bimolecular: two molecules in the elementary step • termolecular: three molecules in the elementary step • It is not common to see termolecular processes (statistically improbable).

4. Elementary Steps & Mechanisms • A multistep mechanism consists of a sequence of elementary steps. • The elementary steps must add to give the overall balanced chemical equation. • Many mechanisms include intermediates. • An intermediate is a species which appears in an elementary step but is not a reactant or product. • So intermediates are produced in 1 elementary step and consumed in a later step. • They do not appear in the overall equation.

5. Elementary Steps & Mechanisms • Step 1: A →B • Step 2: B →C • Overall: A →C • B is an intermediate which is produced in Step 1 and consumed in Step 2. • So it does not appear in the overall rxn.

6. Rate Laws of Elementary Steps • The rate laws of the elementary steps determine the overall rate law of the reaction. • The rate law of an elementary step is determined by its molecularity. • In other words, for an elementary step, the rate law may be written directly from the balanced equation for that step! • Remember that the rate law for a rxn can’t be written from the overall balanced equation for the rxn. (Unless it is a 1-step rxn.)

7. Rate Laws of Elementary Steps • Unimolecular processes are first order • A→B rate = k[A] • Bimolecular processes are second order • 2A→B rate = k[A]2 OR • A + B→C rate = k[A][B] • Termolecular processes are third order (very rare) • 3A→B rate = k[A]3 OR • 2A + B→C rate = k[A]2[B] OR • A + B + C→D rate = k[A][B][C] • Why are termolecular processes rare?

8. Rate Laws of Multistep Mechanisms • Most rxns occur by mechanisms with more than one elementary step. • Often, 1 step is much slower than the other steps. • This slow steplimits or governs the overall rxn rate. • The slow step is called the rate-determining step. • Therefore, the rate-determining step governs the overall rate law for the reaction.

9. Mechanisms with an Initial Slow Step • Look at the following rxn: • NO2(g) + CO(g)→NO(g) + CO2(g) • The experimentally derived rate law is: • Rate = k[NO2]2 • Because chemists want to know how the rxn occurs, they then propose mechanisms which would yield the experimentally known rate law. • They then test their proposed mechanism to see if it is correct.

10. Mechanisms with an Initial Slow Step • So for the previous rxn, the following mechanism is proposed: • Note that NO3 is an intermediate. • If k2 >> K1, then Step 1 is the rate-determining step. • So the overall rate law comes from Step 1.

11. Mechanisms with an Initial Slow Step • Does the mechanism support the known rate law of rate = k[NO2]2 ? • As Step 1 is the rate-determining step, the mechanism yields the rate law.

12. Mechanisms with an Initial Slow Step • Do the mechanism steps sum up to give the known rxn equation of: • NO2(g) + CO(g)→NO(g) + CO2(g) • Yes!

13. Mechanisms with an Initial Slow Step • This means that the proposed mechanism is consistent with the experimentally known rate law. • It does not prove that the mechanism is correct. • Many, many experiments need to done to state with reasonable confidence that a proposed mechanism is probably correct.

14. Mechanisms with an Initial Fast Step • Not all mechanisms start with the slow step. • And, it is possible for an intermediate to be in the slow step. • Consider • 2NO(g) + Br2(g)→ 2NOBr(g)

15. Mechanisms with an Initial Fast Step • 2NO(g) + Br2(g)→ 2NOBr(g) • The experimentally determined rate law is • Rate = k[NO]2[Br2] • Consider the following mechanism • What are the step rate laws?

16. Mechanisms with an Initial Fast Step • The theoretical rate law is (based on Step 2): • Rate = k2[NOBr2][NO] • But NOBr2 is an intermediate! • And we can hardly ever measure intermediates as they are unstable (don’t stick around long)! • Lastly, we want to write the rate law in terms of the reactants in the balanced equation! • How do we get rid of the intermediate NOBr2?

17. Mechanisms with an Initial Fast Step • If NOBr2 is unstable, then it doesn’t accumulate! • This means that it breaks apart as soon as it forms. • It can do this 2 ways: • It can react with NO to produce the product NOBr; but this is a slow step so only a small fraction does this. • It can fall apart and reform the reactants NO and Br2 in the reverse of Step 1: this is a fast dynamic equilibrium step, so it happens very quickly.

18. Mechanisms with an Initial Fast Step • In a dynamic equilibrium, the forward rate, rf, equals the reverse rate, rr • In our case, this means that r1 = r-1 • Therefore: • k1[NO][Br2] = k-1[NOBr2] • Rearranging, we get:

19. Mechanisms with an Initial Fast Step • Now we can get rid of [NOBr2]! • We substitute this into our original rate law from Step 2: • rate = k2[NO][NOBr2] = k-1[NOBr2] • Therefore, the overall rate law becomes • Note the final rate law is consistent with the experimentally observed rate law.

20. Catalysis • A catalyst changes the rate of a chemical reaction. • There are two types of catalyst: • homogeneous, and • heterogeneous. • Catalysts are common in the body, in the environment, and in the chemistry lab! • Homogeneous Catalysis • The catalyst and reaction is in one phase. • Hydrogen peroxide decomposes very slowly: • 2H2O2(aq)→ 2H2O(l) + O2(g).

21. Catalysis • Homogeneous Catalysis • 2H2O2(aq)→ 2H2O(l) + O2(g). • In the presence of the bromide ion, the decomposition occurs rapidly: • 2Br-(aq) + H2O2(aq) + 2H+(aq)→Br2(aq) + 2H2O(l). • Br2(aq) + H2O2(aq)→ 2Br-(aq) + 2H+(aq) + O2(g). • Note that the sum of these steps is the overall rxn. • So Br- is a catalyst because it can be recovered at the end of the reaction. • Generally, catalysts operate by lowering the activation energy for a reaction.

22. Catalysis Homogeneous Catalysis

23. Catalysis • Homogeneous Catalysis • Catalysts can operate by increasing the number of effective collisions. • That is, from the Arrhenius equation: catalysts increase k be increasing A or decreasing Ea. • A catalyst usually changes the mechanism. • Example: In the presence of Br-, Br2(aq) is generated as an intermediate in the decomposition of H2O2. • When a catalyst adds an intermediate, the activation energies for both steps must be lower than the activation energy for the uncatalyzed reaction.

24. Catalysis • Heterogeneous Catalysis • The catalyst is in a different phase than the reactants and products. • Typical example: solid catalyst, gaseous reactants and products (catalytic converters in cars). • Most industrial catalysts are heterogeneous. • First step is adsorption (the binding of reactant molecules to the catalyst surface). • Adsorbed species (atoms or ions) are very reactive. • Molecules are adsorbed onto active sites (red spheres) on the catalyst surface.

25. Catalysis • Heterogeneous Catalysis • The number of active sites on a given amount of catalyst depends on: • Type of catalyst • How the catalyst was prepared • How the catalyst was treated prior to use

26. Catalysis Heterogeneous Catalysis

27. Catalysis • Heterogeneous Catalysis • Consider the hydrogenation of ethylene: • C2H4(g) + H2(g)→C2H6(g), ΔH°= -136 kJ/mol. • The reaction is slow in the absence of a catalyst. • In the presence of a metal catalyst (Ni, Pt or Pd) the reaction occurs quickly at room temperature. • First the ethylene and hydrogen molecules are adsorbed onto active sites on the metal surface. • The H-H bond breaks and the H atoms migrate about the metal surface.

28. Catalysis • Heterogeneous Catalysis • Consider the hydrogenation of ethylene: • C2H4(g) + H2(g)→C2H6(g), ΔH°= -136 kJ/mol. • When an H atom collides with an ethylene molecule on the surface, the C-Cπbond breaks and a C-Hσbond forms. • When C2H6 forms it desorbs from the surface. • When ethylene and hydrogen are adsorbed onto a surface, less energy is required to break the bonds, the activation energy for the reaction is lowered, and the rate increases.

29. Catalysis • Enzymes • Enzymes are biological catalysts. • Most enzymes are protein molecules with large molecular masses (10,000 to 106 amu). • Enzymes have very specific shapes. • Most enzymes catalyze very specific reactions. • Substrates undergo reaction at the active site of an enzyme. • A substrate locks into an enzyme and a fast reaction occurs.

30. Catalysis Enzymes

31. Catalysis • Enzymes • The products then move away from the enzyme. • Only substrates that fit into the enzyme lock can be involved in the reaction. • If a molecule binds tightly to an enzyme so that another substrate cannot displace it, then the active site is blocked and the catalyst is inhibited (enzyme inhibitors). • The number of events (turnover number) catalyzed is large for enzymes (103 - 107 per second).

32. Catalysis • Nitrogen Fixation and Nitrogenase • Nitrogen gas cannot be used in the soil for plants or animals. • Nitrogen compounds, NO3, NO2-, and NO3- are used in the soil. • The conversion between N2 and NH3 is a process with a high activation energy (the N-N triple bond needs to be broken). • An enzyme, nitrogenase, in bacteria which live in root nodules of legumes, clover and alfalfa, catalyses the reduction of nitrogen to ammonia.

33. Catalysis Nitrogen Fixation and Nitrogenase

34. Catalysis • Nitrogen Fixation and Nitrogenase • The fixed nitrogen (NO3, NO2-, and NO3-) is consumed by plants and then eaten by animals. • Animal waste and dead plants are attacked by bacteria that break down the fixed nitrogen and produce N2 gas for the atmosphere.