# Chapter 13 Lecture Outline

## Introduction: M&M Kinetics

Introduce Kinetics with M&M's and spreadsheet M&M Kinetics
1. Zero Order Example (start with 100 on overhead, eat 10 every 5 seconds)
2. First Order Example (start with 100 on overhead, eat 1/5 every 5 seconds)
3. Second Order Example (start with 100 on overhead, eat 0.006 of number squared)
4. Experiment with General Graph to show effect of rate.Reaction Rate and Order

## The Rate of a Chemical Reaction

1. Rates as [delta]C/[delta]t
2. Overhead, and Figure 15.2 from Textbook
3. Rate for specific product or reactant
4. For the reaction 2 NO2 <-> 2 NO + O2 At 300 C
1. Data from Zumdahl

 time NO2 NO O2 0 0.0100 0 0 50 0.0079 0.0021 0.0011 100 0.0065 0.0035 0.0018 150 0.0055 0.0045 0.0023 200 0.0048 0.0052 0.0026 250 0.0043 0.0057 0.0029 300 0.0038 0.0062 0.0031 350 0.0034 0.0066 0.0033 400 0.0031 0.0069 0.0035

2. Graph concentration vs time for NO2
3. Calculate concentration of other species and graph
4. Calculate average rate
1. from 0 to 50
2. from 200 to 250
5. Completed Spreadsheet
6. Compare rates for each species
1. CD-ROM 15-2 side bar (Bleach and dye)
2. Surface Area. ( internet © Saunders, 1997)
3. Concentration. CD-ROM, 15-4 ( internet © Saunders, 1997)
1. Mg (s) + 2 HCl (aq) -> H2 (g) + Mg2+ + 2 Cl1- (aq)
2. 2 N2O5 -> 4 NO2 + O2 (Initial Rate vs concentration)

## The Rate Law: The Effect of Concentration on Reaction Rate

1. The rate law
1. Expression:
for a reaction: aA + bB -> cC + dD
rate is rate = k [A]x [B]y

2. Discuss units for kinetics problems (Handout).

3. Method of Initial Rates: Determine rate law from initial rates or rates at different concentrations. (Lecture Problems).

4. Derive Relationship

ln(Rate1/Rate2) = x ln(A1/A2)
Where:
1. Rate1 is the rate under the first conditions.
2. Rate2 is the rate under the second conditions.
3. A1 is the concentration of A for the first conditions.
4. A2 is the concentration of A for the second conditions.
5. x is the reaction order for A.

5. For the reaction:

NH4+(aq) + NO2-(aq) -> N2(g) + 2 H2O (l)

With the following experimental data.

 [NH41+] (M) [NO21-] (M) Rate (M s-1 0.100 0.005 1.35*10-7 0.100 0.010 2.75*10-7 0.200 0.010 5.4*10-7

1. Determine the order of a reaction
2. Determine k for reaction:
3. Write the rate law
4. Use the rate law to calculate the rate for new concentrations.

## The Integrated Rate Law: The Dependence of Concentration on Time

(Lecture Problems)
1. Plot of concentration vs time, and rate vs time
1. Units for axis
2. Units for area under curve
3. Integration of rate equation

2. Derive 1st order integrated rate equation

3. Use the integrated rate equation to find the concentration at time t for:

C2H5Cl -> C2H4 + HCl at 700 K, k = 2.50*10-3 min-1

1. If the initial concentration of C2H5Cl is 3.45 atm, how long does it take for the partial pressure of C2H5Cl to drop to 1.00 atm?
2. What is the partial pressure of C2H5Cl after six hours?
3. How long does it take for the C2H5Cl concentration to drop to 1/2 the initial concentration?
4. How long does it take for 99% of the C2H5Cl to react?

4. Graph Concentration vs time from integrated rate equation (overhead).

## Rate laws, integrated rate equation, and units

 Reaction Order Rate Rate Law Integrated Rate Equation Linear Plot Slope Units for k 0 [delta] concentration ----------------------- [delta] time rate = k [A]0 - [A]t = kt [A] vs time -k concentration time-1 mol liter-1 sec-1 pressure time-1 1 [delta] concentration ----------------------- [delta] time rate = k [A] ln ([A]t/[A]0) = - kt ln [A] vs time -k sec-1 2 [delta] concentration ----------------------- [delta] time rate = k [A]2rate = k [A] [B] (1/[A]t) - (1/[A]0) = kt (1/[A]) vs time k concentration-1 time-1 liter mole-1 sec-1 pressure-1 time-1

## Graphing Kinetics Data

Graphing Kinetics Data (pdf)

## The Effect of Temperature on Reaction Rate

1. Collision Theory
1. Collision Energy NO + O3 -> NO2 + O2 ( internet© Saunders, 1997)

2. Orentation NO + O3 -> NO2 + O2 ( internet© Saunders, 1997)

3. Successful Reaction NO + O3 -> NO2 + O2 ( internet© Saunders, 1997)

4. Energy Diagram (Overhead)

5. Reaction Energy Diagram F1- + CH3Cl -> CH3F + Cl1- ( internet© Saunders, 1997)

6. Temperature and reaction rate, bleach demo ( internet © Saunders, 1997)

7. Concentration and reaction (collision) rate).

2. The Arrhenius Equation

3. Graphing ln(k) vs 1/T, idea of linear relationships
1. Derive relationship
2. Show graph on Lotus

4. Calculations with the Arrhenius equiaton
1. Given A and Ea , determine k at T

2. Find change in rate at two T's from Ea

3. Determine Ea from k at two T's

4. Arrhenius Plot to find Ea

5. (Lecture Problems)

## Catalysis

1. H2O2 Decomposition ( internet© Saunders, 1997)
2. Catalysis reaction steps
1. Given the reaction steps:
1. NO + O3 -> NO3 + O (slow)
NO3 + O -> NO2 + O2 (fast)

NO + O3 -> NO2 + O2 (overall)

2. And an additional step:

3. NO2 + O -> NO + O2

4. Combines as:

5. NO + O3 -> NO2 + O2
NO2 + O -> NO + O2

O3 + O -> 2 O2

2. Which species is a catalyst?
3. Which species is an intermediate?
4. How does this catalyst effect the rate of the reaction
3. Catalysis Energetics and Reaction Rates. (Lecture Problems)

For the ovrall reaction: O3 + O -> 2 O2

1. For the catalyzed reaction Ea = 11.9 kJ
2. For the uncatalyzed reaction Ea = 14.0 kJ

3. Draw an energy level diagram
4. What is the rate constant at 298 K for each mechanism?
5. For the catalytic destruction of O3 by Cl Ea = 2.1 kJ, what would this rate constant be at 298 K?

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