1. Chapter 16

2. Thermodynamics tells us if a reaction can occur while Kinetics tells us how quickly the reaction occurs some reactions that are thermodynamically feasible are kinetically so slow they are hardly noticeable 2

3. 3 Kinetics is the study of rates of chemical reactions and the mechanisms by which they occur. Reaction rate increase in concentration of a product per unit time or decrease in concentration of a reactant per unit time Reaction mechanism series of steps by which a reaction occurs the rate of rxn is always determined by the slowest step in its mechanism (rate – determining step) Ways to measure rxn speed: for gases - pressure for concentration – color intensity

4.  Reaction rates ~ rates at which reactants disappear or products appear the change in concentration of a reactant or product. = change in concentration / (number of moles x change in time) Note: reactants are negative and products are positive

5. Ex. 1) 4A + B  2C + 3D

6.  Reaction Mechanism: a series of elementary steps that must agree with the experimentally determined rate law and the sum of the elementary steps must give the overall balanced equation for the rxn.  the individual steps in a reaction mechanism are called elementary steps. Assume they happen as written, collision wise  You don’t want to have intermediates in a rate expression. An intermediate is a chemical that is neither a reactant nor a product. It is formed and consumed in the course of the rxn. NO + Br 2 NOBr 2 NOBr 2 + NO  2NOBr NOBr 2 is the intermediate

7. Elementary StepMolecularityRate Law A  products Unimolecular – rxn involving one molecule Rate = k[A] A + A  products (2A  products) Bimolecular – rxn involving collision of two species Rate = k[A] 2 A + B  products BimolecularRate = k[A][B] A + A + B  prod. (2A + B  prod.) Termolecular – rxn involving collision of 3 species Rate = k[A] 2 [B] A + B + C  prod. TermolecularRate = k[A][B][C] Molecularity – the number of species that must collide to produce the reaction indicated by that step Examples of Elementary Steps

8. 8  Rate Law Expressions must be determined experimentally cannot be determined from balanced equations most chemical reactions are not one-step reactions  Rate law expressions are also called: rate laws rate equations rate expressions

9. By definition, the rate law for the reaction: 2 A + 3 B  5 C is: Rate = k  A  x  B  y k = specific rate constant at a given temp. x and y are called orders of rxn. x + y +…= overall order of rxn Note that rate laws are written using only the reactants. The orders of reaction,x & y, are from data, NOT the coefficients of the equation. In fact, the only time the orders equal the coefficients is when the reactions are elementary rxns.

10. 10 Orders of a reaction are expressed in terms of either: each reactant or overall reaction For example:

11. 11

12. 12  A catalyst is a substance that increases the rate of the rxn, but it remains unchanged when the rxn is complete  Catalysts change reaction rates by providing a faster alternative pathway where a different, lesser amount of activation energy is needed.

13. 13 Homogeneous catalysts exist in same phase as the reactants. Heterogeneous catalysts exist in different phases than the reactants. Often are solids Examples of catalysts include:

14. Ex. 2) Step 1: NO + Cl 2 + Pt  NOCl 2 Pt Fast Step 2: NOCl 2 Pt + NO  2NOCl + Pt Slow A. What is the overall reaction? B. What is the catalyst?

15. C. What is the intermediate? (an intermediate is a substance that is produced and then used up during a reaction) D. What is the molecularity in step 1? E. What is the rate determining step? (For a mechanism to be consistent, the steps must add up to the overall rxn, and the rate determining step must give the derived rate law.)

16. F. If this rxn is 1 st order [NOCl 2 Pt] and 1 st order [NO] because they are the reactants in the rate determining step, what is the rate law for this expression? G. What is the overall reaction order?

17. 17  Rate of a simple one-step reaction is directly proportional to the concentration of the reacting substance [X] is concentration of X in molarity or moles/L For a simple expression like R = k[A]: doubling the initial concentration of A doubles the initial rate of reaction halving the initial concentration of A halves the initial rate of reaction

18. 18  Look at the following reaction and its experimentally determined rate-law expression  because it is a second order rate-law expression doubling the [A] increases the rate of reaction by a factor of 4 2 2 = 4 halving the [A] decreases the rate of reaction by a factor of 4 (1/2) 2 = 1/4

19. Ex. 3) For the reaction 3A + B  C, the rate law is R = k [A] 2 A. What order is the reaction? B. How will these change the rate if you triple the concentration of A C. Double the concentration of B

20.  The orders of the reactions are derives by the following mathematical equation: using conc. from experiments ratios of the rates =  ratios of A  x  ratios of B  y … Or rate 1 = k[A] x 1 [B] y 1 [C] z 1 … rate 2 = k[A] x 2 [B] y 2 [C] z 2 … (we’ll solve some examples later)

21. Reactant How [M] Δ Effect on rate Order A Doubledno change 0 A Doubleddoubles 1 A Doubledquadruples 2 A Doubledeight times 3

22. Order [ Δ M] order = Δ rate 0 2 0 = 1 1 2 1 = 2 2 2 2 = 4 3 2 3 = 8

23. Or in other words: If concentration doubles (2) and the rate stays the same: ask yourself ~ 2 to what power = 1 so 2 X = 1 ( the power is 0 so the reaction is 0 order in A) Conc. doubles & rate doubles: 2 X = 2 ( x = 1 or first order) Conc. doubles & rate quadruples: 2 x = 4 ( x = 2 or second order)

24. Three basic events must occur for a reaction to occur the atoms, molecules or ions must: 1. collide 2. collide with enough energy to break and form bonds 3. collide with the proper orientation

25. 1. Temperature – if the temp. increases, then the rxn tends to speed up. Generally a raise of 10 o C doubles the rate. 2. Concentration – the more concentrated the reactants are the quicker the rxn. 3. Nature of Reactions – some chemicals react very quickly, some are very slow Na + H 2 O  Na + + OH - + H 2 very fast Al + H 2 O  Al 3+ + OH - + H 2 very slow 4. Catalyst – chemical that speeds up a rxn w/o being used up.

26. 26  For reactions that have an E a  50 kJ/mol, (E a = activation energy ) the rate approximately doubles for a 10 0 C rise in temperature, near room temperature. 2 ICl(g) + H 2 (g)  I 2 (g) + 2 HCl(g) The rate-law expression is known to be R=k[ICl][H 2 ]

27. 27 Simplified representation of effect of different numbers of molecules in the same volume. Increase in concentration, increases the chance of a rxn occurring. A(g) + B(g)  Products A B B A B 4 different possible A- B collisions 6 different possible A- B collisions 9 different possible A- B collisions

28. 28  Broad category that includes the different reacting properties of substances. For example: ~ Sodium reacts with water explosively at room temperature to liberate hydrogen and form sodium hydroxide. ~ Calcium reacts with water only slowly at room temperature to liberate hydrogen and form calcium hydroxide.

29. ~ The reaction of magnesium with water at room temperature is so slow that that the evolution of hydrogen is not perceptible to the human eye. ~ However, Mg reacts with steam rapidly to liberate H 2 and form magnesium oxide. Differences due to “nature of the reactants”

30. 30  Use experimental rate-law to postulate a mechanism.  The slowest step in a reaction mechanism is the rate determining step.  Note: Experimentally determined reaction orders indicate the number of molecules involved in: the slow step only or the slow step and the equilibrium steps preceding the slow step.

31. 31  Reaction is known to be first order in H 2 O 2, first order in I -, and second order overall.  Mechanism is thought to be: Consider the iodide ion catalyzed decomposition of hydrogen peroxide to water and oxygen.

32. 32  Important notes:  one hydrogen peroxide molecule and one iodide ion are involved in the rate determining step  the iodide ion catalyst is consumed in step 1 and produced in step 2 in equal amounts  hypoiodite ion has been detected in reaction mixture as a short-lived reaction intermediate

33. 33  Ozone, O 3, reacts very rapidly with nitrogen oxide, NO, in a reaction that is first order in each reactant and second order overall. A possible mechanism is:

34. 34  This mechanism is inconsistent with the rate-law expression b/c the slowest step doesn’t match the rate-law found in the lab:

35. 35 Integrated rate equation relates time and concentration for a chemical or nuclear reaction Why do we need this? Problem with normal rate law is that you can’t figure out the rate at some later time. First Order Reactions 1st order in reactant A & 1st order overall For example: a A  products common for many chemical reactions and all simple radioactive decays

36. 36 where: [A] 0 = mol/L of A at time t=0. [A] = mol/L of A at time t. k = specific rate constant t = time elapsed since beginning of reaction a = stoichiometric coefficient of A in balanced overall equation

37.  To find units for k, the specific rate constant, look at overall order  Copy Table 16-2, summary of orders of reactions, and Table16-3, graphing for orders of reactions, into your notes before starting on Kinetics Example Problems Overall OrderUnits 0[M] time -1 1time -1 2[M] -1 time -1 3[M] -2 time -1

38. 38  The Chernobyl nuclear reactor accident occurred in 1986. At the time that the reactor exploded some 2.4 MCi of radioactive 137 Cs into the atmosphere. The half-life of 137 Cs is 30.1 years. In what year will the amount of 137 Cs released from Chernobyl finally decrease to 100 Ci? A Ci is a unit of radioactivity called the Curie, MCi = MegaCurie