1. Kinetics Cartoon courtesy of NearingZero.net ©2011 University of Illinois Board of Trustees http://islcs.ncsa.illinois.edu/copyright

2. Chemical Kinetics The area of chemistry that concerns reaction rates. However, only a small fraction of collisions produces a reaction. Why? Key Idea: Molecules must collide to react. ©2011 University of Illinois Board of Trustees http://islcs.ncsa.illinois.edu/ copyright

3. Collision Theory Orientation of reactants must allow formation of new bonds. Collisions must have enough energy to produce the reaction (must equal or exceed the activation energy). ©2011 University of Illinois Board of Trustees http://islcs.ncsa.illinois.edu/copyright

5. Factors Affecting Rate  Temperature Increasing temperature always increases the rate of a reaction.  Surface Area Surface Area Increasing surface area increases the rate of a reaction  Concentration Increasing concentration USUALLY increases the rate of a reaction  Presence of Catalysts  All increase number of effective collisions ©2011 University of Illinois Board of Trustees http://islcs.ncsa.illinois.edu/copyright

6. Catalysis Catalyst: A substance that speeds up a reaction without being consumedCatalyst: A substance that speeds up a reaction without being consumed Enzyme: A large molecule (usually a protein) that catalyzes biological reactions.Enzyme: A large molecule (usually a protein) that catalyzes biological reactions. Homogeneous catalyst: Present in the same phase as the reacting molecules.Homogeneous catalyst: Present in the same phase as the reacting molecules. Heterogeneous catalyst: Present in a different phase than the reacting molecules.Heterogeneous catalyst: Present in a different phase than the reacting molecules. ©2011 University of Illinois Board of Trustees http://islcs.ncsa.illinois.edu/copyright

7. Reaction Rate The change in concentration of a reactant or product per unit of time ©2011 University of Illinois Board of Trustees http://islcs.ncsa.illinois.edu/copyright

8. Activation Energy The minimum energy required to transform reactants into the activated complex (The minimum energy required to produce an effective collision) Flame, spark, high temperature, radiation are all sources of activation energy ©2011 University of Illinois Board of Trustees http://islcs.ncsa.illinois.edu/copyright

9. Endothermic Reactions ©2011 University of Illinois Board of Trustees http://islcs.ncsa.illinois.edu/copyright

10. Exothermic Reactions ©2011 University of Illinois Board of Trustees http://islcs.ncsa.illinois.edu/copyright

11. Endothermic Reaction with a Catalyst ©2011 University of Illinois Board of Trustees http://islcs.ncsa.illinois.edu/copyright

12. Exothermic Reaction with a Catalyst ©2011 University of Illinois Board of Trustees http://islcs.ncsa.illinois.edu/copyright

13. Decomposition of H 2 O 2 by Various Catalysts ©2011 University of Illinois Board of Trustees http://islcs.ncsa.illinois.edu/copyright

14. Catalysts Increase the Number of Effective Collisions ©2011 University of Illinois Board of Trustees http://islcs.ncsa.illinois.edu/copyright

15. Rate Laws for Reactions Rate Laws are determined experimentally Relationship between concentration of one reactant and the rate ©2011 University of Illinois Board of Trustees http://islcs.ncsa.illinois.edu/copyright

16. 2NO 2 (g)  2NO(g) + O 2 (g) Reaction Rates: 2. Can measure appearance of products 1. Can measure disappearance of reactants 3. Are proportional stoichiometrically ©2011 University of Illinois Board of Trustees http://islcs.ncsa.illinois.edu/copyright

17. Rate Laws Consider the system If the concentration of NO is held constant and the concentration of H 2 is varied, the experiment determines how the rate of the reaction varies with respect to the concentration of H 2 ©2011 University of Illinois Board of Trustees http://islcs.ncsa.illinois.edu/copyright

18. The following results were found experimentally… doubling the concentration of H 2 doubled the rate and tripling the concentration of H 2 tripled the rate Therefore the concentration of H 2 is directly related to the rate of the reaction and can be written R is proportional to [H 2 ] The reaction is first order with respect to hydrogen (the exponent is 1) ©2011 University of Illinois Board of Trustees http://islcs.ncsa.illinois.edu/copyright

19. When the concentration of H2 was held constant… Doubling the concentration of NO increased the rate of the reaction four times Tripling the concentration of NO increased the rate of the reaction nine times Therefore the rate of the reaction is directly related to the square of the concentration of NO R is proportional to [NO] 2 The reaction is second order with respect to NO (the exponent is 2) ©2011 University of Illinois Board of Trustees http://islcs.ncsa.illinois.edu/copyright

20. Combining these results Introducing a proportionality constant k gives an overall rate law R = k [H 2 ][NO] 2 ©2011 University of Illinois Board of Trustees http://islcs.ncsa.illinois.edu/copyright

21. Reaction Mechanism The series of steps by which a chemical reaction occurs. A chemical equation does not tell us how reactants become products It is a summary of the overall process. Example: has many steps in the reaction mechanism ©2011 University of Illinois Board of Trustees http://islcs.ncsa.illinois.edu/copyright

22. Rate-Determining Step In a multi-step reaction, the slowest step is the rate-determining step. It therefore determines the rate of reaction. slowest step is the rate-determining step slowest step is the rate-determining step ©2011 University of Illinois Board of Trustees http://islcs.ncsa.illinois.edu/copyright