1. Rate Law & Reaction Order 02
Reaction Order: The sum of the powers to which all reactant concentrations appearing in the rate law are raised.
Reaction order is determined experimentally:
By inspection.
From slope of the line, using the appropriote plot: [A] vs. t, ln[A] vs. t, 1/[A] vs. t

2. First-Order Reactions 02
Using calculus we obtain the integrated rate equation:
ln [A]t = – kt + ln [A]0
y = m x b
Plotting ln[A]t against t will give:

3. Determining Reaction Order
Reaction: H2O2 → H2O ½ O2
Plot [H2O2] vs. time
time
[H2O2]
ln[H2O2] 1
0.0000
0.705
0.3496 2
0.497
0.6992 3
0.349
1.0527 4
0.246
1.4024 5
0.173
1.7545

4. Determining Reaction Order
Reaction: H2O2 → H2O ½ O2
Plot ln [H2O2] vs. time
time
[H2O2]
ln[H2O2] 1
0.000
0.705
-0.349 2
0.497
-0.699 3
0.349
-1.053 4
0.246
-1.402 5
0.173
-1.754

5. Determining Reaction Order
Reaction: H2O2 → H2O ½ O2
Plot ln [H2O2] vs. time
time
[H2O2]
ln[H2O2] 1
0.0000
0.705
0.3496 2
0.497
0.6992 3
0.349
1.0527 4
0.246
1.4024 5
0.173
1.7545

6. First-Order Reactions 04
Show that the decomposition of N2O5 is first order and calculate the rate constant.
Plot: ln[N2O5] vs. t Is it linear?

7. First-Order Reactions 06
Half-Life: Time for reactant concentration to decrease by half its original value.

8. Second-Order Reactions 01
A  Products A + B  Products
Rate = k[A] Rate = k[A][B]
These equations can then be integrated to give:
y = m x b
Does this form look familiar?

9. Second-Order Reactions 01
Rate = k[A]2
y = m x b
Plot 1/[A] vs. t

10. Second-Order Reactions 02
Half-Life: Time for reactant concentration to decrease by half its original value.

11. Second-Order Reactions 03
Iodine atoms combine, form molecular iodine in gas phase:
I(g) + I(g)  I2(g)
The reaction follows second-order kinetics. Rate constant: k = 7.0 x 10–1 M–1s– at 23°C.
1. If initial concentration of I is M, calculate the concentration after 2.0 min.
2. Calculate the half-life at the start of the reaction if the initial concentration of I is 0.60 M.

12. Second-Order Reactions 03
Iodine atoms combine, form molecular iodine in gas phase:
I(g) + I(g)  I2(g)
2nd Order, k = 7.0 x 10–1 M–1s–1 [I]0 = M, Find [I]120

13. Second-Order Reactions
Iodine atoms combine, form molecular iodine in gas phase:
I(g) + I(g)  I2(g)
2nd Order, k = 7.0 x 10–1 M–1s–1 [I]0 = 0.60 M, Find t1/2

14. Review – 1st and 2nd Order Reactions
(integrated form)

15. Review – Rate Constants
Units for “Rate” will always be M/sec
Molarity units will vary, depending on reaction order

16. Review – Equations for Kinetics Problems
General Rate Law:
rate = k [A]x [B]y
First, determine the order of the reaction (find exponents x and y) by inspection or by graphing.
Then this equation can be used to calculate the rate constant k.

17. Review – Equations for Kinetics Problems
Integrated forms of Rate Laws:
1st Order Reaction:
2nd Order Reaction:
Zero Order Reaction:
rate = k [A]x [B]y
ln[A]t = –kt + ln[A]0
1/[A]t = kt + 1/[A]0
[A]t = -kt + [A]0