VI. How Reactions Occur

1. VI. How Reactions Occur • Most chemical reactions occur through several small steps, not one big step. • A chemical equation typically shows the overall reaction, not the intermediate steps. • e.g. H2(g) + 2ICl(g) 2HCl(g) + I2(g) only shows what’s at the beginning and what you end up with.

2. VI. Reaction Mechanisms • A reaction mechanism is a series of individual chemical steps through which an overall chemical reaction occurs. • A proposed mechanism for the reaction H2(g) + 2ICl(g) 2HCl(g) + I2(g) is: Step 1 H2(g) + ICl(g) HI(g) + HCl(g) Step 2 HI(g) + ICl(g)  HCl(g) + I2(g)

3. VI. Elementary Steps • The reactions in a mechanism are called elementary steps; what’s implied in these steps is exactly what happens. • Proposed reaction mechanisms must add up to the overall reaction! • Does the previous mechanism add up? • Species that are formed in one step and then consumed in another are known as intermediates. What is/are the intermediate(s) in the previous mechanism?

4. VI. Elementary Step Rate Laws • Elementary steps are characterized by their molecularity, i.e. the # of reactant particles involved in the step. • Rate laws for elementary steps can be written directly from their stoichiometry! • e.g. If A + B  C + D is an elementary step, then the rate law for this step is: Rate = k[A][B].

5. VI. Energy Diagram, 2-Step Mechanism

6. VI. The Rate-Determining Step • The slow step in the mechanism will determine the overall rate of reaction. • This step is known as the rate-determining step. • It’s the bottleneck of the reaction.

7. VI. Valid Mechanisms • Valid mechanisms satisfy 2 criteria: • Elementary steps add up to overall reaction. • Rate law predicted by mechanism must be consistent with experimental rate law. • Note that a valid mechanism is not a proven mechanism.

8. VI. Example • Consider the reaction: NO2(g) + CO(g) NO(g) + CO2(g). • Experimentally, Rate = k[NO2]2. This implies it’s not a single-step reaction. Why? • Is the mechanism below valid? NO2(g) + NO2(g) NO3(g) + NO(g) Slow NO3(g) + CO(g)  NO2(g) + CO2(g) Fast

9. VI. Rate Laws w/ Intermediates • Rate laws must always be written from the rate-determining step. • However, rate laws cannot contain intermediates. • Rate laws from other steps can be used to substitute for intermediates. • We look at fast first steps.

10. VI. Fast 1st Steps • When the 1st step is fast, its products will build up and reverse reaction starts. • Eventually, an equilibrium is set up. • Thus, for A + B  C + D (Fast), we can write A + B  C + D. • Rate = k[A][B] and Rate = k-1[C][D]. • At equilibrium, k[A][B] = k-1[C][D]. • This can be used to rewrite rate laws.

11. VI. Sample Problem • What is the overall reaction and rate law for the mechanism below? Identify the intermediates as well. Cl2(g) 2Cl(g) Fast Cl(g) + CHCl3(g)  HCl(g) + CCl3(g) Slow CCl3(g) + Cl(g)  CCl4(g) Fast

12. VII. Catalysts • We know we can change reaction rates by changing the temperature or changing reactant concentrations. • However, there are limits to these tactics. • If available, can use catalysts, substances that increase reaction rate, but aren’t used up in the reaction.

13. VII. Catalytic Destruction of O3 Catalyzed: Cl(g) + O3(g) ClO(g) + O2(g) ClO(g) + O(g)  Cl(g) + O2(g) Uncatalyzed: O3(g) + O(g) 2O2(g) O3(g) + O(g)  2O2(g) Atomic chlorine from photodissociated CFC’s is the catalyst.

14. VII. How Do Catalysts Work? • Catalysts provide a lower-energy mechanism for the reaction.

15. VII. Types of Catalysts • There are homogeneous and heterogeneous catalysts.

16. VII. Biological Catalysts • A biological catalyst is called and enzyme. • An enzymes has an active site into which a specific substrate fits – like a lock and key.