1. Atkins & de Paula: Elements of Physical Chemistry: 5e
Chapter 8:
Chemical Equilibrium:
Equilibria in Solution

2. End of chapter 8 assignments
Discussion questions:
3, 5, 6
Exercises:
1, 5, 7, 8, 9, 10, 11, 12, 13, 24, 28, 29, 34, 35, 36
Use Excel if data needs to be graphed

3. Homework assignments Did you:
Read the chapter?
Work through the example problems?
Connect to the publisher’s website & access the “Living Graphs”?
Examine the “Checklist of Key Ideas”?
Work any end-of-chapter exercises?
Review terms and concepts that you should recall from previous courses
Study the introductory pages closely!

4. Brønsted-Lowry theory
Note the use of activity, a, for [J]
Activity is technically correct; ions interact over distances, making [J] “hazardous”
Conjugate acid=accepted the H+
Conjugate base=anion remains from acid
So… pH = -log[H3O+] is pH = -log a
H3O+

5. Brønsted-Lowry theory
conjugate acid
conjugate base
base
acid

6. Protonation and deprotonation
Strong & weak electrolytes, acids, bases
Acidity constants & basicity constants
HA + H2O  H3O+ + A-
B +H2O  BH+ + OH-
aHA a
H3O+ A-
Ka =
[H3O+][A-]
[HA] =
pKa = -log Ka aB a
BH+
OH-
Kb =
[BH+][OH-]
[B] =
pKb = -log Kb

7. Protonation and deprotonation
Autoionization (autoprotolysis) of water
H2O + H2O  H3O+ + OH-
Kw = Ka x Kb & pKw = pKa + pKb
pKw = pH + pOH
pH = -log a and pOH = -log a
H3O+
OH-

8. Protonation and deprotonation
The fraction (f) of molecules of a weak acid (HA) that has donated a proton (deprotonated):
f =
equilibrium molar concentration of conjugate base
molar concentration of acid as prepared
[A-]equilibrium
[HA]as prepared

9. Protonation and deprotonation
The fraction (f) of molecules of a weak base (B) that has accepted a proton (been protonated):
f =
equilibrium molar concentration of conjugate acid
molar concentration of base as prepared
[BH+]equilibrium
[B]as prepared

10. Table 8.1 Acidity and basicity constants* at 298.15 K (1)
Weakest weak acids
p.182
Weakest weak acids

11. Table 8.1 Acidity and basicity constants* at 298.15 K (2)
p.182
Table 8.1 Acidity and basicity constants* at K (2)

12. Work through these Example 8.1, p.176 Example 8.2, p.176f
Do any of you need help working I C E
problems?

13. Polyprotic acids H2A(aq) + H2O(l)  H3O+(aq) + HA-(aq)
HA-(aq) + H2O(l)  H3O+(aq) + A2-(aq)
Work Example 8.3 & Example 8.4 (p.178f)
Ka1 =
aH2A a
H3O+
HA-
Ka2 =
aHA- a
H3O+
A2-

14. Polyprotic acids Work Example 8.3 (p.178) Work Example 8.4 (p.178f)
Understand the graphs in Fig 8.1 & 8.2

15. Polyprotic acids For Example 8.4
For pH < pKa, the acid form dominates
For pH = pKa, the conjugate pair have equal concentrations
For pH > pKa, the base form dominates

16. Table 8.2 Successive acidity constants of polyprotic acids at 298.15 K

17. Amphiprotic systems Amphoteric is the old term
Can act as an acid or a base
Examples: HCO3- and amino acids
Is NaHCO3(aq) acidic or basic?
pH = ½ (pKa1 + pKa2) [See Derivation 8.1 p.180]
(Back up one slide and calculate the pH)
Does this “make sense”?

18. Acid-base titrations Terms you must understand:
Stoichiometric point (equivalence point)
Analyte
Titrant
pH curve
Work through Example 8.5 (p.182) and Illustration 8.3 (p.183)

19. pH curve Fig 8.3 (181) Titration of a strong acid with a strong base

20. pH curve Fig 8.4 (181) Titration of a weak acid with a strong base

21. pH curve Fig 8.5 (183) Titration of a weak base with a strong acid

22. Buffer action
Buffer solutions change pH very little when a small amount of acid or base is added to the solution
Acid buffer stabilizes pH < 7 – a weak acid and a salt that contains its conjugate base
Base buffer stabilizes pH > 7 – a weak base and a salt that contains its conjugate acid

23. pH curve of a buffer solution
Fig 8.6 (191)
pH of a solution changes slowly in the region halfway to the stoichiometric point
Here, the solution is buffered
pH ~ pKa
acid
base

24. Buffer action Consider an aqueous solution of CH3COOH
CH3COOH(aq) + H2O(l) CH3COO-(aq) + H3O+
conjugate base
acid
Consider an equimolar mixture of CH3COOH and CH3COO-Na+
Add strong acid
H+ (aq) + CH3COO- (aq) CH3COOH (aq)
Add strong base
OH- (aq) + CH3COOH (aq) CH3COO- (aq) + H2O (l)

25. Acid-base indicators
A water-soluble organic molecule with acid (HIn) and conjugate base (In-) forms with different colors.
Examples:
Phenolphthalein
Cresol red
Methyl red
Thymol blue
Litmus

26. Acid-base indicators HIn (aq) H+ (aq) + In- (aq) [HIn]  10
Color of acid (HIn) predominates
 1
[HIn]
[In-]
If [HIn]  [In-], combination color
 10
[HIn]
[In-]
Color of conjugate base (In-) predominates

27. Table 8.3 Indicator color changes

28. pH curve Fig 8.7 (194) Strong acid, strong base titration
Which indicators work well here?
Why will several indicators be adequate here?
acid
base

29. pH curve Fig 8.8 (194) Weak acid, strong base titration
Which indicators work well here?
Why will bromo-thymol blue not work well here?
acid
base

30. Solubility
Solubility: the maximum quantity of solute that will dissolve in a solvent (saturation)
Does T or p have to be specified? Explain
The solubility constant, A/K/A “solubility product constant.” Why?
MX  M+ + X-
Ks = a a or Ks = [M+][X-] M+ X-

31. The solubility constant
Ks varies according to the formula of the salt
AgCl (s) Ag+ (aq) + Cl- (aq)
Ks = [Ag+][Cl-]
MgF2 (s) Mg2+ (aq) + 2F- (aq)
Ks = [Mg2+][F-]2
Ag2CO3 (s) Ag+ (aq) + CO32- (aq)
Ks = [Ag+]2[CO32-]
Ca3(PO4)2 (s) Ca2+ (aq) + 2PO43- (aq)
Ks = [Ca2+]3[PO33-]2

32. Relationship between Ks and Molar Solubility (s)
Ks expression
Compound
Cation
Anion
Relation between Ks & s

33. Table 8.4 Solubility constants at 298.15 K (1)

34. Table 8.4 Solubility constants at 298.15 K (2)

35. Table 8.4 Solubility constants at 298.15 K (3)

36. The common ion effect
The common ion effect is the shift in equilibrium caused by the addition of a compound having an ion in common with the dissolved substance.
The presence of a common ion suppresses the ionization of a weak acid or a weak base.
Consider a mixture of CH3COO-Na+ (strong electrolyte) and CH3COOH (weak acid).
CH3COO-Na+ (s) Na+ (aq) + CH3COO- (aq)
CH3COOH (aq) H+ (aq) + CH3COO- (aq)

37. The common ion effect
The common ion effect is the shift in equilibrium caused by the addition of a compound having an ion in common with the dissolved substance.
The addition of a common ion decreases the solubility of a salt.
“It is very dangerous to neglect deviations from ideal behavior in ionic solutions…. [Q]uantitative calculations are unreliable.” (p.196)

38. Solubility Rules for Common Ionic Compounds in water at 25°C
Soluble Compounds
Exceptions
Compounds containing alkali metal ions & NH4+
NO3-, HCO3-, ClO3-
Cl-, Br-, I-
Halides of Ag+, Hg22+, Pb2+
SO42-
Sulfates of Ag+, Ca2+, Sr2+, Ba2+, Hg2+, Pb2+
Insoluble Compounds
CO32-, PO43-, CrO42-, S2-
Compounds containing alkali metal ions and NH4+
OH-
Compounds containing alkali metal ions and Ba2+
You should know these!

39. Solubility Equilibria
Dissolution of an ionic solid in aqueous solution:
Q < Ksp
Unsaturated solution
No precipitate
Q = Ksp
Saturated solution
No precipitate
Q > Ksp
Supersaturated solution
Precipitate will form

40. Key Ideas

41. Key Ideas

42. The End
…of this chapter…”