The document discusses corrections made to tables in a textbook on chemical kinetics. Table 1.1 corrected the concentration of Mg2+ and noted analytical errors in data for the Mississippi River sample that will be replaced. Table 1.3a corrected the coefficient in the Davies equation from 0.2 to 0.3, as suggested by Davies. The rest of the document covers an overview of chemical kinetics including expressions for irreversible reactions, the effects of temperature, and reversible reactions.

1. Corrections to the text
 Table 1.1 Mg2+ = 1280 mg/L
 Table 1.1 electroneutrality is usually ±15%,
which indicates analytical errors and/or missing
solute. The Mississippi River sample has an
error of about 50%; those data will be replaced.
 Table 1.3a Davies equation
 The final term has a coefficient of 0.3
 The example has a coefficient of 0.2
 Davies suggested a value from 0.2 to 0.3
1

2. Chemical Kinetics
Material from the handout

3. Two views of equilibrium
 Based on the minimum energy
 Thermodynamic perspective
 Approach used after tonight
 Based on kinetics
 Forward reaction rate
 Reverse reaction rate
 Equilibrium occurs when the net rate = 0

4. An overview
 Expressions/analysis of irreversible reactions
 0, 1st, and 2nd order
 Elementary and non-elementary reactions
 Series and parallel reactions
 Reversible reactions
 Effect of temperature on kinetics

5. Kinetic expressions for…
 Basic irreversible reactions:
 Zero order: A→P
 1st order : A→P
 2nd order: A+A→P
 2nd order : A+B→P

6. An example: Given these batch
experimental data, estimate
the rate constant
Time (d) 0 5 10 15 20 25 30
C (mg/L) 25 15 9 6 3 2 1

7. Maybe it’s a 1 order reaction
st

8. Integrate…

9. Convenient linear form

10. Is a 1 order model appropriate?
st
 Plot ln(C) = f(t)
 Look for straight line
 Slope = k
t (d) 0 5 10 15 20 25 30
C (mg/L) 25 15 9 6 3 2 1
ln C 3.22 2.71 2.20 1.79 1.10 0.69 0

11. 1 order assessment: k ≈ -0.11 d
st -1
3.5
3 ln C = -0.1057 t + 3.26
R2 = 0.9961
2.5
2
Ln C
1.5
1
0.5
0
0 5 10 15 20 25 30 35
Time (d)

12. Kinetic expressions for
elementary reactions

13. Kinetic expressions for
elementary reactions - II

14. Differential rate analysis

15. Example
Apply differential approach
t (min) [A] (mg/L)
0 10
2 5.8
4 3.7
6 2.6
8 1.9
10 1.5

16. Example
Data manipulation
t (min) [A] (mg/L) ∆[A]/∆t [A]avg Ln(-∆[A]/∆t)
0 10 - -
2 5.8 -2.1 7.9 0.74
4 3.7 -1.05 4.75 0.049
6 2.6 -0.55 3.15 -0.598
8 1.9 -0.35 2.25 -1.05
10 1.5 -0.02 1.7 -1.61
[A]avg = average concentration during the period ∆t.

17. Example
Integral analysis
n = 1.51
kn = exp(-2.34) = 0.096
1
0.5
y = 1.5102x - 2.3399
R2 = 0.9964
0
ln(-∆A/∆t)
-0.5
-1
-1.5
-2
0 0.5 1 1.5 2 2.5
ln(Aavg)

18. Example
2.5
2
What about 1st and 1.5
y = -0.1884x + 2.1705
R2 = 0.9794
2nd order kinetics?
ln (A)
1
0.5
0
0 2 4 6 8 10
Time (min)
0.7
0.6
y = 0.0573x + 0.067
0.5 2
R = 0.9861
0.4
1/A
0.3
0.2
0.1
0
0 2 4 6 8 10
Time (min)

19. Which model is correct?
n = 1.51 provides the best fit;
n = 1 or 2 is not too bad

20. Elementary and
non-elementary reactions

21. Elementary reactions

22. An enzyme-catalyzed reaction

23. Pseudo 1 order reactions
st

24. Experimental design

25. Multiple step reactions
are common
 For example, reactions in series:
A→B→C
 Reactions in parallel:
A→B
A→C

26. Reactions in series example
 Reaction A → B is a 0 order reaction
k0 = 10-5 mg.L-1.min-1
 Reaction B → C is a 1st order reaction
k1 = 0.01 min-1
 Assume all concentrations can be expressed as
mg of carbon per L
 Initial concentrations
A0 = 10-3 mg/L B0 = C0 = 0 mg/L
 Determine A, B, and C = f(t)

27. Rate expressions
27

28. Integrate for B
28

29. Rearrange to solve for B
29

30. 1.E-03
A
B
1.E-03 C
Conentration (mg C / L)
8.E-04
6.E-04
4.E-04
2.E-04
0.E+00
0 50 100 150 200 250 300
Time (min)

31. Parallel reactions example

32. Quiz 1
 Imagine you can follow a molecule of copper
from the source to the mouth of the Mississippi
River. Assume that the average concentration of
dissolved copper remains constant at 5 µg/L.
The total dissolved solids, however, increases
significantly from source to mouth.
 Would you expect the activity to increase,
decrease, or remain the same? Why?
 From the perspective of a fish, what is more
important, the activity or the concentration of
copper? Briefly explain your answer.
32

33. Reversible reactions

34. Relating A and B

35. Substitute to simplify

36. Integrate for A

37. Rearrange to solve for A

38. Some special conditions

39. Similar analysis for B

40. Irreversible reactions?
 If the reverse reaction is very slow relative to
the forward reaction…
 Examples of irreversible reactions:
 Precipitation
 Life

41. Temperature effects

42. Assessing experimental data
 Plot ln(k) versus 1/T
 Intercept = ln(kAr)
 Slope = - EAr/R

43. Example for the Q-10 rule
 Suppose a biochemical rate doubles…
 For a system increase from T = 10°C to 20°C.
 Estimate the activation energy for the reaction

44. Arrhenius equation at any T

45. Solve for EAr

46. Summary of key points
 Models for simple kinetics
 Basic elementary reactions
 0, 1st, or 2nd order models
 Evaluation of experimental data
 Complex reaction mechanisms are likely
 Parallel or series combinations
 Temperature effects

47. How does water temperature
affect pH?
 Equilibrium expression:
{H+} {OH-} = Kw = 10-14.0 (T = 25 C)
 Charge balance in pure water:
[H+] = [OH-]
What is the pH at T = 1 °C?

48. van’t Hoff equation

49. Substitute
49

50. Neutral pH at T = 1° C