Acid-Base Equilibria and
Solubility Equilibria
Chemistry 3
prepared by: kayhan saied
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Prepared by:
kayhan saied
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 mixture of CH3COONa (strong electrolyte) and
CH3COOH (weak acid).
CH3COONa (s) Na+ (aq) + CH3COO- (aq)
common
CH3COOH (aq) H+ (aq) + CH3COO- (aq) ion
prepared by: kayhan saied
16.2
Consider mixture of salt NaA and weak acid HA.
NaA (s) Na+ (aq) + A- (aq) [H+][A-]
Ka =
HA (aq) H+ (aq) + A- (aq) [HA]
Ka [HA]
[H+] =
[A-] Henderson-Hasselbalch
equation
[HA]
-log [H+] = -log Ka - log
[A-] [conjugate base]
-] pH = pKa + log
[A [acid]
-log [H+] = -log Ka + log
[HA]
[A-]
pH = pKa + log pKa = -log Ka
[HA]
prepared by: kayhan saied
16.2
What is the pH of a solution containing 0.30 M HCOOH
and 0.52 M HCOOK?
Mixture of weak acid and conjugate base!
HCOOH (aq) H+ (aq) + HCOO- (aq)
Initial (M) 0.30 0.00 0.52
Change (M) -x +x +x
Equilibrium (M) 0.30 - x x 0.52 + x
[HCOO-]
Common ion effect pH = pKa + log
[HCOOH]
0.30 – x 0.30
[0.52]
0.52 + x 0.52 pH = 3.77 + log = 4.01
[0.30]
HCOOH pKa = 3.77
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16.2
A buffer solution is a solution of:
1. A weak acid or a weak base and
2. The salt of the weak acid or weak base
Both must be present!
A buffer solution has the ability to resist changes in pH upon
the addition of small amounts of either acid or base.
Consider an equal molar mixture of CH3COOH and CH3COONa
Add strong acid
H+ (aq) + CH3COO- (aq) CH3COOH (aq)
Add strong base
OH- (aq) + CH3COOH (aq) CH3COO- (aq) + H2O (l)
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16.3
HCl H+ + Cl-
HCl + CH3COO- CH3COOH + Cl-
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16.3
Which of the following are buffer systems? (a) KF/HF
(b) KBr/HBr, (c) Na2CO3/NaHCO3
(a) KF is a weak acid and F- is its conjugate base
buffer solution
(b) HBr is a strong acid
not a buffer solution
(c) CO32- is a weak base and HCO3- is its conjugate acid
buffer solution
prepared by: kayhan saied
16.3
Solubility Equilibria
AgCl (s) Ag+ (aq) + Cl- (aq)
Ksp = [Ag+][Cl-] Ksp is the solubility product constant
MgF2 (s) Mg2+ (aq) + 2F- (aq) Ksp = [Mg2+][F-]2
Ag2CO3 (s) 2Ag+ (aq) + CO32- (aq) Ksp = [Ag+]2[CO32-]
Ca3(PO4)2 (s) 3Ca2+ (aq) + 2PO43- (aq) Ksp = [Ca2+]3[PO43-]2
Dissolution of an ionic solid in aqueous solution:
Q < Ksp Unsaturated solution No precipitate
Q = Ksp Saturated solution
Q > Ksp Supersaturated solution Precipitate will form
prepared by: kayhan saied
16.6
prepared by: kayhan saied
16.6
Molar solubility (mol/L) is the number of moles of solute
dissolved in 1 L of a saturated solution.
Solubility (g/L) is the number of grams of solute dissolved in
1 L of a saturated solution.
prepared by: kayhan saied
16.6
What is the solubility of silver chloride in g/L ?
AgCl (s) Ag+ (aq) + Cl- (aq) Ksp = 1.6 x 10-10
Initial (M) 0.00 0.00 Ksp = [Ag+][Cl-]
Change (M) +s +s Ksp = s2
Equilibrium (M) s s s = Ksp
s = 1.3 x 10-5
[Ag+] = 1.3 x 10-5 M [Cl-] = 1.3 x 10-5 M
1.3 x 10-5 mol AgCl 143.35 g AgCl
Solubility of AgCl = x = 1.9 x 10-3 g/L
1 L soln 1 mol AgCl
prepared by: kayhan saied
16.6
prepared by: kayhan saied
16.6
If 2.00 mL of 0.200 M NaOH are added to 1.00 L of
0.100 M CaCl2, will a precipitate form?
The ions present in solution are Na+, OH-, Ca2+, Cl-.
Only possible precipitate is Ca(OH)2 (solubility rules).
Is Q > Ksp for Ca(OH)2?
[Ca2+]0 = 0.100 M [OH-]0 = 4.0 x 10-4 M
Q = [Ca2+]0[OH-]02 = 0.10 x (4.0 x 10-4)2 = 1.6 x 10-8
Ksp = [Ca2+][OH-]2 = 8.0 x 10-6
Q < Ksp No precipitate will form
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16.6
What concentration of Ag is required to precipitate
ONLY AgBr in a solution that contains both Br- and Cl-
at a concentration of 0.02 M?
AgBr (s) Ag+ (aq) + Br- (aq) Ksp = 7.7 x 10-13
Ksp = [Ag+][Br-]
Ksp 7.7 x 10 -13
-11 M
[Ag+] = = = 3.9 x 10
[Br-] 0.020
AgCl (s) Ag+ (aq) + Cl- (aq) Ksp = 1.6 x 10-10
Ksp = [Ag+][Cl-]
Ksp 1.6 x 10-10 -9 M
[Ag+] = = = 8.0 x 10
[Cl-] 0.020
3.9 x 10-11 M < [Ag+] < 8.0 x 10-9 M
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16.7
The Common Ion Effect and Solubility
The presence of a common ion decreases
the solubility of the salt.
What is the molar solubility of AgBr in (a) pure water
and (b) 0.0010 M NaBr?
NaBr (s) Na+ (aq) + Br- (aq)
AgBr (s) Ag+ (aq) + Br- (aq) [Br-] = 0.0010 M
Ksp = 7.7 x 10-13 AgBr (s) Ag+ (aq) + Br- (aq)
s2 = Ksp [Ag+] = s
s = 8.8 x 10-7 [Br-] = 0.0010 + s 0.0010
Ksp = 0.0010 x s
s = 7.7 x 10-10
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16.8
pH and Solubility
• The presence of a common ion decreases the solubility.
• Insoluble bases dissolve in acidic solutions
• Insoluble acids dissolve in basic solutions
remove
add
Mg(OH)2 (s) Mg2+ (aq) + 2OH- (aq)
At pH less than 10.45
Ksp = [Mg2+][OH-]2 = 1.2 x 10-11
Lower [OH-]
Ksp = (s)(2s)2 = 4s3
OH- (aq) + H+ (aq) H2O (l)
4s3 = 1.2 x 10-11
s = 1.4 x 10-4 M Increase solubility of Mg(OH)2
[OH-] = 2s = 2.8 x 10-4 M At pH greater than 10.45
pOH = 3.55 pH = 10.45 Raise [OH-]
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Decrease solubility of Mg(OH)2
16.9
Complex Ion Equilibria and Solubility
A complex ion is an ion containing a central metal cation
bonded to one or more molecules or ions.
Co2+ (aq) + 4Cl- (aq) CoCl42- (aq)
The formation constant or stability constant (Kf) is the
equilibrium constant for the complex ion formation.
[CoCl42- ]
Kf =
Co(H2O)2+ CoCl2- [Co2+][Cl-]4
6 4
stability of
Kf
complex
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16.10
prepared by: kayhan saied
16.10
prepared by: kayhan saied
16.11
Qualitative
Analysis of
Cations
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16.11
Flame Test for Cations
lithium sodium potassium copper
prepared by: kayhan saied
16.11
