1. Kinetics Chapter 12

2. Reaction Rates  Kinetics is concerned with studying the reaction mechanism of a reaction.  An average reaction rate describes how fast (not how spontaneous) a given reaction is overall.  Only positive rates are considered. In the event of a decrease in concentration, we insert a negative sign into the equation.

3. Reaction Rates on a Graph  Graph concentration and time of reaction.  Slope of a line tangent to a point gives instantaneous rate. 2NO 2  2NO + O 2

4. Reaction Rate Summary  Average rate is not always equal to instantaneous rate.  Concentration of the products and reactants change over time, producing logarithmic curves.  Because concentration changes, rate is changing.  Coefficients in the balanced reaction indicate the relationship between the rates of reactions AND the concentrations.

5. Rate Laws Intro  All chemical reactions are reversible.  When the rate of the forward and reverse reactions are equal, it is called equilibrium.  Having the reverse reaction occur simultaneously with the forward messes with concentration.  We manipulate conditions to inhibit the progress of the reverse reaction when measuring rates.

6. Rate Law Equation  There are two types of rate laws.  Integrated  Regular Integrated Rate Law Compares concentration and time Rate =  [A]/  t Regular Rate Law Compares concentration with rate of reaction Rate = k[A] n  Things to watch for:  How is the rate defined (in reference to which reactant).  Only reactants are included in rate law.

7. Determining the Order of Each Reactant in the Rate Law  The order is the superscript on the rate law.  These values are determined from data that is derived from several experiments that measure initial concentration and initial rate of reaction.  You must scrutinize the relationship between concentration and rate to determine the order of a given reactant.  Each reactant can have a different order.

8. Types of Orders  Zero Order  Change in concentration elicits no change in rate.  This will eliminate the reactant from the rate law.  First Order  Direct, equally proportional relationship between concentration and rate.  ½ the concentration, ½ the rate  3x the concentration, 3x the rate, etc.  Second Order  Exponential relationship  ½ concentration, ¼ rate.  3x concentration, 9x rate

9. Calculating orders Only compare one concentration variable at a time…look for reactions in which one variable is held constant.

10. Reaction Mechanisms  A reaction mechanism is a series of smaller reactions by which the overall reaction occurs.  The smaller reactions must add together to give the complete balanced overall reaction.  The experimentally derived rate law must agree with the rate law given in the slowest step of the reaction mechanism.

11. Types of elementary steps

12. Rate Determining Step  A reaction is only as fast as its slowest step. NO 2 + NO 2  NO 3 + NO NO 3 + CO  NO 2 + CO 2 We have to hypothesize the slowest step based upon an experimentally known rate. NO 2 + CO  NO + CO 2 has a rate law: Rate = k[NO2] 2 Which is the rate determining step?

13. Try Me  The balanced equation for the reaction of gases nitrogen dioxide and fluorine is 2NO 2(g) + F 2(g)  2NO 2 F (g) The experimentally determined rate law is Rate = k[NO 2 ][F 2 ] A mechanism for this reaction is: NO 2 + F 2  NO 2 F + F F + NO 2  NO 2 F Which part of the mechanism do you suppose is the most likely Rate Determining Step?

14. More Practice  A proposed mechanism for a reaction is slow C 4 H 9 Br  C 4 H 9 + + Br - slow C 4 H 9 Br  C 4 H 9 + + Br - fast C 4 H 9 + + H 2 O  C 4 H 9 OH 2 + fast C 4 H 9 + + H 2 O  C 4 H 9 OH 2 + fast C 4 H 9 OH 2 + + H 2 O  C 4 H 9 OH + H 3 O + fast C 4 H 9 OH 2 + + H 2 O  C 4 H 9 OH + H 3 O + Write the rate law for this mechanism. What is the overall balanced equation for the reaction? What are the intermediates in the proposed mechanism?

15. Collision Theory and Kinetics  Molecules must collide to react.  Only some collisions yield a reaction.  Activation energy must be overcome to break and form necessary bonds.

16. Increase Productivity!  Temperature  Increase temperature, increase KE  # effective collisions = (total collisions)*e -E a /RT  Molecular Orientation  What touches in a collision will determine if something happens or not.

17. Calculating Activation Energy  Arrhenius Equation:  Natural log of Arrhenius Equation: k = [A]e -E a /RT ln(k) = -(E a /R)(1/T)+ ln[A]

18. Try Me!!  The Reaction: 2 N 2 O 5  4NO 2 + O 2 yielded the following data. Calculate E a. k (s -1 ) T ( o C) 2.0 x 10 -5 20 7.3 x 10 -5 30 2.7 x 10 -4 40 9.1 x 10 -4 50 2.9 x 10 -3 60

19. Derive This Equation: ln(k 2 /k 1 ) = (E a /R)(1/T 1 – 1//T 2 ) From These Equations: ln(k 1 ) = -(E a /R)(1/T 1 )+ ln(A) ln(k 2 ) = -(E a /R)(1/T 2 )+ ln(A)

20. Catalysis  Catalyst: a substance that speeds up a reaction without being consumed during the process.  How?  The presence of a catalyst causes the reaction to take a different reaction mechanism.  The new reaction mechanism has a different slow step.  The new slow step has a lower activation energy. What do you already know about how a catalyst works?

22. Heterogeneous Catalysts  Reactants are ADSORBED on the surface of the catalyst.  Reactants migrate on the surface of the catalyst.  Reactants form bonds and are desorbed from the surface of the catalyst. Homogeneous catalysts exist in the same phase as the reactants…ie freon gas Heterogeneous catalysts are in a phase…ie Pt metal