1. 13.6-13.7 Reaction Mechanisms -Catalysis

2. Reaction Mechanisms
Most reactions do not occur in a single step, but rather, in a series of steps. However, when we usually see a reaction written, we are usually seeing the overall reaction.
The overall reaction = the sum of the intermediate steps.
Each individual step is called an elementary step. These represent the exact species that are colliding in a chemical reaction.
The reaction mechanism: a series of steps by which a chemical reaction can occur.

3. Reaction Mechanisms (continued)
Let’s consider the over reaction below: H2(g) + 2ICl(g)  2HCl(g) + I2(g) This reaction actually contains two steps: Step 1: H2(g) + ICl(g)  HI(g) + HCl(g) Step 2: HI(g) + ICl(g)  HCl(g) + I2(g) Now look to see if any molecules appear on both side of the reaction. HI is both a product in step 1 and a reactant in step 2 and is not part of the overall reaction. Therefore, HI is known as a reaction intermediate. Reaction intermediates are formed in one elementary step, and consumed in another.

4. Rate Laws for Elementary Steps
Elementary steps are categorized by their molecularity, the number of reactant particles involved in the reaction.
Examples:
A  product unimolecular (first order reactions)
A + B  product bimolecular (second order)
A + B + C  product termolecular  these are rare!
We can’t deduce the rate law from the overall reaction, but we can for the elementary step from its equation.

5. The Rate Determining Step
Rate determining step: the elementary step in a reaction that is slower than all the others.
The rate determining step in a reaction mechanism limits the overall rate of the reaction.
For a proposed reaction mechanism to be valid, two conditions must be met:
1.) The elementary steps in the mechanism must sum to
the overall reaction.
2.) The rate law predicted by the mechanism must be
consistent with the experimentally observed rate law.

6. Example
In a reaction with two elementary steps ,where the first is slower (RDS) than the second, the experimental rate law was determined to be: rate = k[NO2]2 NO2(g) + NO2(g)  NO3(g) + NO(g) Slow NO3(g) + CO(g)  NO2(g) + CO2(g) Fast Overall reaction: NO2(g) + CO(g)  NO(g) + CO2(g) The rate = k[NO2]2 which is second order and agrees with the experimental rate law. So this reaction mechanism is valid.

7. Overall Reaction Rate Laws
If the initial step in the reaction is the rate limiting step, then the rate order of the overall reaction should be equal to the rate law of the initial step.
However, if the initial step is fast, and another elementary step that follows is limiting, the RDS may contain a reaction intermediate. If this occurs, the reactant that the reaction intermediate depends on must be used in the rate law, in place of the intermediate.
There is an example on the following slide 

8. Let’s Try a Practice Problem!
Predict the overall reaction and rate law that result from the following two-step mechanism: 2A  A2 Slow A2 + B  A2B Fast Overall reaction: 2A + B  A2B Rate = k[A]2

9. Example
In a reaction with three elementary steps, where the second is rate limiting, and an equilibrium is reached in the first reaction, the experimental rate law was determined to be Rate= k[H2][NO]2 k1 Step 1: 2NO(g) N2O2(g) Fast k-1 Step 2: H2(g) + N2O2(g)  H2O(g) + N2O(g) Slow k2 Step 3: N2O(g) + H2(g)  N2(g) + H2O(g) Fast k3 The overall reaction: 2H2(g) + 2NO(g)  2H2O(g) + N2(g) Rate= k[H2][NO]2 , The reaction mechanism is valid!

10. Catalysis
So far, we have seen how both the increase in concentration and temperature can increase the reaction rate, but they may not always be the best way to do so.
A catalyst: a substance that increases the rate of a chemical reaction, but is not consumed by the reaction.
Works by providing an alternate mechanism for the reaction, one in which the RDS has a lower activation energy.

11. Homogeneous and Heterogeneous Catalysts
Homogeneous catalysts: exists in the same phase as the reactants. (i.e. If both reactants are gases, and so is the catalyst, the catalyst is said to be homogeneous)
Heterogeneous catalyst: the catalyst exists in a different phase than the reactants.
Directly from the college board: (above) Talked about 

12. Biological Catalysts
Living organisms rely on enzymes, biological catalysts that increase the rate of biochemical reactions, to survive.
Enzymes: usually large protein molecules with complex three dimensional structures. Enzymes are extremely specific, only catalyzing a single reaction.
Active site: a specific area within the enzyme that a reactant molecule (usually called a substrate) can bind to.
The substrate fits into the active site just as a key fits in a lock (this binding occurs through IMFs or covalent bonds). When this happens, the activation energy is of the reaction is lowered greatly.
An example of this is the enzyme sucrase, which catalyzes the breaking up of sucrose (table sugar) into glucose and fructose.

13. pgs #’s 74, 76, and 78
Study for quiz: chapter 13