1. Copyright©2000 by Houghton Mifflin Company. All rights reserved. 1 Chemical Kinetics The area of chemistry that concerns reaction rates.

2. Copyright©2000 by Houghton Mifflin Company. All rights reserved. 2 Reaction Rate Change in concentration (conc) of a reactant or product per unit time.

3. Copyright©2000 by Houghton Mifflin Company. All rights reserved. 3 Rate Laws Rate = k[NO 2 ] n k = rate constant n = rate order

4. Copyright©2000 by Houghton Mifflin Company. All rights reserved. 4 Types of Rate Laws Differential Rate Law: expresses how rate depends on concentration. Integrated Rate Law: expresses how concentration depends on time.

5. Copyright©2000 by Houghton Mifflin Company. All rights reserved. 5 Method of Initial Rates Initial Rate: the “instantaneous rate” just after the reaction begins. The initial rate is determined in several experiments using different initial concentrations.

6. Copyright©2000 by Houghton Mifflin Company. All rights reserved. 6 Overall Reaction Order Sum of the order of each component in the rate law. rate = k[H 2 SeO 3 ][H + ] 2 [I  ] 3 The overall reaction order is 1 + 2 + 3 = 6.

7. Copyright©2000 by Houghton Mifflin Company. All rights reserved. 7 First-Order Rate Law Integrated first-order rate law is ln[A] =  kt + ln[A] o For aA  Products in a 1st-order reaction,

8. Copyright©2000 by Houghton Mifflin Company. All rights reserved. 8 Half-Life of a First-Order Reaction t 1/2 = half-life of the reaction k = rate constant For a first-order reaction, the half-life does not depend on concentration.

9. Copyright©2000 by Houghton Mifflin Company. All rights reserved. 9 Second-Order Rate Law For aA  products in a second-order reaction, Integrated rate law is

10. Copyright©2000 by Houghton Mifflin Company. All rights reserved. 10 Half-Life of a Second-Order Reaction t 1/2 = half-life of the reaction k = rate constant A o = initial concentration of A The half-life is dependent upon the initial concentration.

11. Copyright©2000 by Houghton Mifflin Company. All rights reserved. 11 A Summary 1.Simplification: Conditions are set such that only forward reaction is important. 2.Two types: differential rate law integrated rate law 3.Which type? Depends on the type of data collected - differential and integrated forms can be interconverted.

12. Copyright©2000 by Houghton Mifflin Company. All rights reserved. 12 A Summary (continued) 4.Most common: method of initial rates. 5.Concentration v. time: used to determine integrated rate law, often graphically. 6.For several reactants: choose conditions under which only one reactant varies significantly (pseudo first-order conditions).

13. Copyright©2000 by Houghton Mifflin Company. All rights reserved. 13 Reaction Mechanism 4 The series of steps by which a chemical reaction occurs. 4 A chemical equation does not tell us how reactants become products - it is a summary of the overall process.

14. Copyright©2000 by Houghton Mifflin Company. All rights reserved. 14 Reaction Mechanism (continued) 4 The reaction has many steps in the reaction mechanism.

15. Copyright©2000 by Houghton Mifflin Company. All rights reserved. 15 Often Used Terms Intermediate: formed in one step and used up in a subsequent step and so is never seen as a product. Molecularity: the number of species that must collide to produce the reaction indicated by that step. Elementary Step: A reaction whose rate law can be written from its molecularity. uni, bi and termolecular

16. Copyright©2000 by Houghton Mifflin Company. All rights reserved. 16 Rate-Determining Step In a multistep reaction, it is the slowest step. It therefore determines the rate of reaction.

17. Copyright©2000 by Houghton Mifflin Company. All rights reserved. 17 Collision Model Key Idea: Molecules must collide to react. However, only a small fraction of collisions produces a reaction. Why? Arrhenius: An activation energy must be overcome.

18. Copyright©2000 by Houghton Mifflin Company. All rights reserved. 18 Arrhenius Equation 4 Collisions must have enough energy to produce the reaction (must equal or exceed the activation energy). 4 Orientation of reactants must allow formation of new bonds.

19. Copyright©2000 by Houghton Mifflin Company. All rights reserved. 19 Arrhenius Equation (continued) k = rate constant A = frequency factor E a = activation energy T = temperature R = gas constant

20. Copyright©2000 by Houghton Mifflin Company. All rights reserved. 20 Catalysis Catalyst: A substance that speeds up a reaction without being consumed Enzyme: A large molecule (usually a protein) that catalyzes biological reactions. Homogeneous catalyst: Present in the same phase as the reacting molecules. Heterogeneous catalyst: Present in a different phase than the reacting molecules.

21. Copyright©2000 by Houghton Mifflin Company. All rights reserved. 21 Heterogeneous Catalysis 1. Adsorption and activation of the reactants. 2. Migration of the adsorbed reactants on the surface. 3. Reaction of the adsorbed substances. 4. Escape, or desorption, of the products. Steps: