1. Reaction Mechanisms

2. A balanced equation for a chemical reaction indicates the substances present at the start of the reaction and those produced as the reaction takes place. It provides no information about how the reaction occurs. The process by which a reaction occurs is called the reaction mechanism.

3. Reactions take place because of collisions between reacting molecules. The reaction of NO and O 3 to form NO 2 and O 2 occurs as a result of a single collision. NO(g) + O 3 (g)  NO 2 (g) + O 2 (g) Processes that occur in a single step are called elementary reactions.

4. Multistep Mechanisms The net change represented by a balanced chemical equation often occurs by a multistep mechanism, which consists of two or more elementary reactions.

5. Consider this reaction: NO 2 (g) + CO(g)  NO (g) +CO 2 (g) This reaction appears to have two elementary reactions(or steps). NO 2 (g) + NO 2 (g)  NO 3 (g) + NO (g) NO 3 (g) + CO (g)  NO 2 (g) + CO 2 (g) The chemical equations for the elementary reactions must always add to give the chemical equations of the overall process. 2NO 2 (g) +NO 3 (g) + CO (g)  NO 2 (g) +NO 3 (g) + NO (g) + CO 2 (g) When we simplify the equation by eliminating substances that appear on both sides of the arrow, we get NO 2 (g) + CO(g)  NO (g) +CO 2 (g) Because NO 3 is not a reactant or product in the overall reaction– it is formed in one elementary reaction and consumed in the next– it is called an intermediate.

6. Intermediates Intermediates are species that appear in some steps but not in the net equation.

7. Study Check The conversion of ozone into O 2 has a two step mechanism: O 3 (g)  O 2 (g) +O(g) O 3 (g) +O(g)  2O 2 (g) Write the equation for the overall reaction. Identify the intermediate.

8. The Rate Determining Step for a Multistep Reaction When you have a multistep reaction, the slowest step limits the reaction rate. The rate law for the overall reaction is based on the rate law for the slowest step in that reaction.

9. Consider the following reaction: NO 2 + CO  NO + CO 2 The reaction is believed to be a two-step process: NO 2 + NO 2  NO 3 + NO slow NO 3 + CO  NO 2 + CO 2 fast

10. NO 2 + NO 2  NO 3 + NO slow NO 3 + CO  NO 2 + CO 2 fast In the first step, two molecules of NO 2 collide, forming the intermediate species NO 3. This molecule then collides with one molecule of CO and reacts quickly to produce one molecule each of NO 2 and CO 2. The first step is the slower of the two, so it is the rate determining step, and we use it to write the rate law for the overall reaction. The reaction is 2 nd order, so the rate law is Rate = k[NO 2 ] 2

11. Study Check 2NO 2 + F 2  2NO 2 F This reaction has a two step mechanism: NO 2 + F 2  NO 2 F + F slow F + NO 2  NO 2 F fast Identify the rate determining step, and write the rate law if the reaction is first order with respect to each reactant.