Acid-Base and Solubility Equilibria

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Acid-Base and Solubility Equilibria

• Common-ion effect
• Buffer solutions
• Acid-base titration
• Solubility equilibria
• Complex ion formation
• Qualitative analysis

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I. Common ion effect

• Common ion an ion that is produced from more
  than one solute
• Common ion effect shift in equilibrium caused by
  the addition of a compound having an ion in
  common with the dissolved substance.

Example Consider CH3COOH (weak acid, partial
ionizes)
Add CH3COONa (strong electrolyte, completely
ionizes)
Le Chatelier's principle equilibrium () shift
This shift to the left will also result in a
decrease of the H
The ionization of a weak electrolyte is decreased
by adding a strong electrolyte (i.e. a salt) that
produces an common ion
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Ex
What is the the pH of a solution containing 0.20
M CH3COOH and 0.30 M CH3COONa at 25oC? (Ka1.8
x10-5).
Initial (M)
0.20
0.00
0.30
Change (M)
-x
x
x
Equilibrium (M)
0.20 - x
x
0.30 x
Assume 0.20 x ? 0.20 0.30 x ? 0.30
x 1.2 x 10-5 M
H CH3COO 1.2 x 10-5 M
pH -log H 4.92

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II. Buffer solution

• A. A solution that resist pH change when acids
  or base are added.
• B. In general, a buffer is made of
• A weak acid and its salt
• A weak base and its salt

HFNaF
NH3NH4Cl
Consider an equal molar mixture of HCOOH HCOONa
Add strong acid
Add strong base
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Ex.
Which of the following can be used to prepare for
buffer solutions? (a) NaF/HF (b) NaBr/HBr,
(c)Na2SO3/NaHSO3, and (d) CH3COOH and NaOH.
(a) HF is a weak acid and F- is its conjugate
base
buffer solution
(b) HBr is a strong acid
not a buffer solution
(c) SO32- is a weak base and HSO3- is it
conjugate acid
buffer solution
(d) CH3COOH is a weak base and NaOH is a strong
base
CH3COOH NaOH --gt CH3COONa H2O
buffer solution
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C. pH of a Buffer solution

• Consider a mixture of a weak acid (HA) and its
  salt (XA)

Henderson-Hasselbalch equation
pKa -log Ka
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Ex
What is the the pH of a solution containing 0.20
M CH3COOH and 0.30 M CH3COONa at 25oC? (Ka1.8
x10-5)? What is the pH after the addition of (a)
20.0 mL of 0.050 M NaOH or (b) 20.0 mL of 0.050 M
HCl to 80.0 mL of this buffer solution?
Henderson-Hasselbalch eq.
4.74 0.18 4.92
(a) Addition of 20.0 mL of 0.050 M NaOH (0.001
mol NaOH)
NaOH will react with the acid in the buffer
start (moles)
0.016
0.001
0.024
end (moles)
0.015
0.00
0.025
final volume 80.0 mL 20.0 mL100.0 mL
4.96
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Continued
Henderson-Hasselbalch eq.
pKa 4.74
(b) Addition of 20.0 mL of 0.050 M HCl (0.001 mol
HCl)
HCl will react with the base in the buffer
start (moles)
0.024
0.001
0.016
end (moles)
0.023
0.0
0.017
final volume 80.0 mL 20.0 mL100.0 mL
4.87
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D. Preparation of buffer

• Select acid/base pair with pKa that is closest to
  the desired pH.
• Adjust ratio of base/acid to have the desired pH.

Ex. How to prepare a buffer with pH 3.85?
H 10-3.851.4 x10-4
HCOOH seems to be a good choice, Ka1.7 x10-4
The buffer can be prepared by mixing 1.0 M HCOOH
and 1.2 M HCOONa
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E. Buffer capacity

• Quantity of acid or base that can be added before
  pH changes significantly
• Buffers work best when the ratio of base/acid
  is 1.
• The greater the concentrations of both base and
  acid, the ________ buffering capacity

greater
Ex. Which of the following has the greater
buffering capacity? (a) 1.0 L of 1.0 M NaF
1.0 M HF solution, or (b) 1.0 L of 0.1 M
NaF 0.1M HF solution
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III. Titration

• Acid-base titration to determine the
  concentration of unknown solution
• Standard solution (titrant) a solution of known
  concentration
• Equivalence point the point at which the
  reaction is complete

• Indicator a substance changes color at/near the
  equivalence point.

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A. Strong acid-strong base titration
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Q.
A 40.0 mL sample of 0.200 M HCl is titrated with
0.200 M NaOH. Calculate the pH of the soution
after the following volume of NaOH have been
added. (a) 35.0 mL, (b) 39.0 mL, (c) 41.0 mL, and
(d) 45.0 mL.
Initial pH 0.699

• mol HCl 0.0400 L x 0.200 M 0.00800 mol
• mol NaOH 0.0350 L x 0.200 M 0.00700 mol
• excess HCl 0.00800 mol - 0.00700 mol 0.00100
  mol
• Final volume (40.0 35.0)mL 75.0 mL
• H0.00100 mol/ 0.0750 L 0.0133 M
• pH -logH 1.875

(b) mol HCl 0.0400 L x 0.200 M 0.00800 mol
mol NaOH 0.0390 L x 0.200 M 0.00780
mol excess HCl 0.00800 mol - 0.00780 mol
0.00020 mol Final volume (40.0 39.0)mL 79.0
mL H0.00020 mol/ 0.0790 L 0.0025 M pH
-logH 2.60
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continued
(c) mol HCl 0.0400 L x 0.200 M 0.00800 mol
mol NaOH 0.0410 L x 0.200 M 0.00820
mol excess NaOH 0.00820 mol - 0.00800 mol
0.00020 mol Final volume (40.0 41.0)mL 81.0
mL OH0.00020 mol/ 0.0810 L 0.0025 M pOH
-logOH 2.61 pH14.0 -pOH 11.39
(d) mol HCl 0.0400 L x 0.200 M 0.00800 mol
mol NaOH 0.0450 L x 0.200 M 0.00900
mol excess NaOH 0.00900 mol - 0.00800 mol
0.00100 mol Final volume (40.0 45.0)mL 85.0
mL OH0.00100 mol/ 0.0850 L 0.0118 M pOH
-logOH 1.929 pH14.0 -pOH 12.071
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B. Weak acid-strong base titration
At equivalence point (pH gt 7)
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Q.
Exactly 100 mL of 0.10 M HNO2 are titrated with a
0.10 M NaOH solution. What is the pH at the
equivalence point ?
start (moles)
0.01
0.01
end (moles)
0.0
0.0
0.01
Final volume 200 mL
0.05
0.00
0.00
-x
x
x
0.05 - x
x
x
pOH 5.98
2.2 x 10-11
pH 14 pOH 8.02
0.05 x ? 0.05
x ? 1.05 x 10-6
OH-
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B. Weak acid-strong base titration curve
1. Initial pH of a weak acid 2. Before
equivalence point buffering region 3. At the
equivalence point a weak base 4. After the
equivalence point excess strong base raises pH
50 mL of 0.1 M weak acid is titrated with 0.1 M
NaOH
pHpKa
equivalence point
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C. Strong acid-weak base titration
At equivalence point (pH lt 7)
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IV. Acid-base indicator
Color of acid (HIn) predominates
Color of conjugate base (In-) predominates
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The titration curve of a strong acid with a
strong base
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Q Which indicator(s) would you use for a
titration of HCOOH with NaOH ?
Weak acid titrated with strong base.
At equivalence point, will have conjugate base of
weak acid.
At equivalence point, pH gt 7
Use cresol red or phenolphthalein
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V. Solubility equilibria
A. Solubility product constant, Ksp
Ksp AgCl-
Ksp Ca2F-2
Ksp Ag2SO42-
Ksp Ca23PO33-2
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B. Solubility and KSP
a. Dissolution of an ionic solid in aqueous
solution
Q lt Ksp
Unsaturated solution
No precipitate
Q Ksp
Saturated solution
Precipitate will form
Q gt Ksp
Supersaturated solution
b. Determine Ksp from solubility or solubility
from Ksp
Molar solubility (mol/L) moles of solute in 1 L
of a saturated solution
Solubility (g/L) grams of solute dissolved in 1
L of a saturated solution.
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Q.
What is the solubility of silver chloride in g/L
? Ksp1.6x10-10
0.00
0.00
s
s
s
s
Ksp AgCl-
Ksp s2
s 1.3 x 10-5
Ag 1.3 x 10-5 M
Cl- 1.3 x 10-5 M
Solubility of AgCl
1.9 x 10-3 g/L
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Q.
The solubility of SrF2 is found to be 1.1 x 10-2
g in 100 mL of aqueous solution. Determine the
value of Ksp of SrF2.
0.00
0.00
s
2s
-s
s
2s
s Solubility of SrF2 in M
8.74 x 104 M
Ksp Sr2F2
s(2s)2
4 s3
2.7 x 10-9
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Q.
If 2.00 mL of 0.200 M NaOH are added to 1.00 L of
0.100 M CaCl2, will a precipitate form? (Ksp
8.0 x10-6)
Which ions are present in solution?
Na, OH-, Ca2, Cl-.
Only possible precipitate is _________(solubility
rules).
Ca(OH)2
Is Q gt Ksp for Ca(OH)2?
Ca20 0.100 M x 1.00 L/(1.00 L 0.00200
L)0.100 M
OH-0 0.00200 L x 0.200 M /(1.00L 0.00200
L)4.0 x 10-4 M
0.10 x (4.0 x 10-4)2 1.6 x 10-8
Ksp Ca2OH-2 8.0 x 10-6
Q lt Ksp
No precipitate will form
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C. Factors affecting solubility

1. Common ion effect common ion decreases the
   solubility.

left
If NaF is added, equilibrium shifts __________.
Ex. What is the molar solubility of CaF2 in (a)
pure water and (b) 0.010 M NaF? (Ksp3.9
x10-11)
s 2.1 x 10-4
Ksp 3.9 x 10-11
s(2s)2 4s3
(b)
F- 0.010 M
F- 0.010 2s ? 0.010
0.00 0.010
Ksp (0.010)2 x s
-s s 2s
s 3.9 x 10-7
s 0.0102s
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2. Effect of pH

• Insoluble bases dissolve in acidic solutions
• Insoluble acids dissolve in basic solutions

increases
In acidic solution, solubility_________.
Why?
Ex. Which salt will be more soluble in acidic
solution than in pure water (a) ZnCO3, (b) AgCN,
and (c) BiI3.
(a) and (b) because CO32 and CN react with H
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Q.
What concentration of Ag is required to
precipitate ONLY AgBr in a solution that contains
both Br- and Cl- at a concentration of 0.010 M?
Ksp 7.7 x 10-13
Ksp AgBr-
Ksp 1.6 x 10-10
Ksp AgCl-
7.7 x 10-11 M lt Ag lt 1.6 x 10-8 M
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VI. Complex ion equilibria and solubility

• Complex ion an ion containing a central metal
  cation (Lewis acid) bonded to one or more
  molecules or ions (Lewis base).
• Solubility of slightly soluble salts can be
  increased by complex ion formation

By adding NH3, solubility of AgCl increases due to
3. The formation constant or stability constant
(Kf) is the equilibrium constant for the complex
ion formation.
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VII. Qualitative analysis
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Flow chart for separation of cations
Solution containing ions of all cation groups
Group 1 precipitates AgCl, Hg2Cl2, PbCl2
HCl
Filtration
Solution containing ions of remaining groups
H2S
Group 2 precipitates CuS, CdS, SnS, Bi2S3
Filtration
Solution containing ions of remaining groups
Group 3 precipitates CoS, FeS, MnS, NiS ZnS,
Al(OH)3,Cr(OH)3
NaOH
Filtration
Solution containing ions of remaining groups
Na2CO3
Group 4 precipitates BaCO3, CaCO3, SrCO3
Filtration
Solution contains Na, K, NH ions

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Flame test of cations
lithium
sodium
potassium
copper