Chapter 15 Lecture Outline

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

Measuring Reaction Rates

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

Reaction Rates

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 Cl^1- (aq)
   2. 2 N₂O₅ -> 4 NO₂ + O₂ (Initial Rate vs concentration)

Rate Equations

1. Calculating Reaction Rates: Completed Spreadsheet

   Calculate rate of NO₂ loss, NO gain, O₂ gain, and reaction for the reaction:

   2 NO₂ <-> 2 NO + O₂ At 300 C

   Data from Zumdahl

   time (s)	NO2(atm)	NO (atm)	O2 (atm)	Rate rxn (atm/s)
   0	0.0100	0	0
   50	0.0079	0.0021	0.0011	2.1*10-5
   100	0.0065	0.0035	0.0018	1.40*10-5
   150	0.0055	0.0045	0.0023	1.00*10-5
   200	0.0048	0.0052	0.0026	7.00*10-6
   250	0.0043	0.0057	0.0029	5.00*10-6
   300	0.0038	0.0062	0.0031	5.00*10-6
   350	0.0034	0.0066	0.0033	4.00*10-6
   400	0.0031	0.0069	0.0035	3.00*10-6

2. The rate law
   1. Expression:
      for a reaction: aA + bB -> cC + dD
      rate is rate = k [A]^x [B]^y
   2. For reaction of NO₂, calculate 1^st order rate constant for each rate.
   3. For reaction of NO₂, calculate 2^nd order rate constant for each rate.
   4. Fits 2^ndorder better so write rate law
      rate = k [NO₂]² Calculate average value of k, from above:
      rate = 0.277 atm^-1 sec^-1 [NO₂]²

3. Discuss units for kinetics problems (Handout). Note: Table 15.1 page 706, correct units for k all sec^-1

Method of Initial Rates

1. 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.

2. For the reaction:
   NH₄⁺(aq) + NO₂^-(aq) -> N₂(g) + 2 H₂O (l)

   With the following experimental data.

   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.

Integrated Rate Equation:

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 1^st 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

Graphing Kinetics Data

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 F^1- + CH₃Cl -> CH₃F + Cl^1- ( internet© Saunders, 1997)

Temperature and the rate constant

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

2. Concentration and reaction (collision) rate).

3. The Arrhenius Equation
   

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

5. Calculations with the Arrhenius equiaton
   1. Given A and E_a , determine k at T

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

   3. Determine E_a from k at two T's

   4. Arrhenius Plot to find E_a

6. (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:
   
   NO₂ + O -> NO + O₂
   3. Combines as:
   
   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 E_a = 11.9 kJ
   2. For the uncatalyzed reaction E_a = 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 E_a = 2.1 kJ, what would this rate constant be at 298 K?

Reaction Mechanisms

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