Transcript Document 7587084

Atkins & de Paula:
Elements of Physical Chemistry: 5e
Chapter 8:
Chemical Equilibrium:
Equilibria in Solution
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
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!
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 aH3O+
Brønsted-Lowry theory
base
acid
conjugate
acid
conjugate
base
Protonation and deprotonation
• Strong & weak electrolytes, acids, bases
• Acidity constants & basicity constants
• HA + H2O  H3O+ + A[H3O+][A-]
Ka =
[HA]
=
aH O+ aA3
aHA
pKa = -log Ka
• B +H2O  BH+ + OH[BH+][OH-]
=
Kb =
[B]
aBH+ a OH-
aB
pKb = -log Kb
Protonation and deprotonation
•
•
•
•
•
Autoionization (autoprotolysis) of water
H2O + H2O  H3O+ + OHKw = Ka x Kb & pKw = pKa + pKb
pKw = pH + pOH
pH = -log aH3O+ and
pOH = -log aOH-
Protonation and deprotonation
• The fraction (f) of molecules of a weak
acid (HA) that has donated a proton
(deprotonated):
• f=
• f=
equilibrium molar concentration of conjugate base
molar concentration of acid as prepared
[A-]equilibrium
[HA]as prepared
Protonation and deprotonation
• The fraction (f) of molecules of a weak
base (B) that has accepted a proton
(been protonated):
• f=
• f=
equilibrium molar concentration of conjugate acid
molar concentration of base as prepared
[BH+]equilibrium
[B]as prepared
Table 8.1 Acidity and basicity constants* at 298.15 K (1)
Acid/base
Kb
pKb
Ka
pKa
Strongest weak acids
Trichloroacetic acid, CCl3COOH
3.3  10
14
13.48
3.0  10
Benzenesulfonic acid, C6H5SO3H
5.0  10
14
13.30
2  10
Iodic acid, HIO3
5.9  10
14
13.23
1.7  10
1
0.77
Sulfurous acid, H2SO3
6.3  10
13
12.19
1.6  10
2
1.81
Chlorous acid, HClO2
1.0  10
12
12.00
1.0  10
2
2.00
Phosphoric acid, H3PO4
1.3  10
12
11.88
7.6  10
3
2.12
Chloroacetic acid, CH2ClCOOH
7.1  10
12
11.15
1.4  10
3
2.85
Lactic acid, CH3CH(OH)COOH
1.2  10
11
10.92
8.4  10
4
3.08
Nitrous acid, HNO2
2.3  10
11
10.63
4.3  10
4
3.37
Hydrofluoric acid, HF
2.9  10
11
10.55
3.5  10
4
3.45
Formic acid, HCOOH
5.6  10
11
10.25
1.8  10
4
3.75
Benzoic acid, C6H5COOH
1.5  10
10
9.81
6.5  10
5
4.19
Acetic acid, CH3COOH
5.6  10
10
9.25
5.6  10
5
4.75
Carbonic acid, H2CO3
2.3  10
8
7.63
4.3  10
7
6.37
Hypochlorous acid, HClO
3.3  10
7
6.47
3.0  10
8
7.53
Hypobromous acid, HBrO
5.0  10
6
5.31
2.0  10
9
8.69
Boric acid, B(OH)3†
1.4  10
5
4.86
7.2  10
10
9.14
Hydrocyanic acid, HCN
2.0  10
5
4.69
4.9  10
10
9.31
Phenol, C6H5OH
7.7  10
5
4.11
1.3  10
10
9.89
Hypoiodous acid, HIO
4.3  10
4
3.36
2.3  10
11
10.64
Weakest weak acids
Weakest weak acids
1
0.52
1
0.70
p.182
Table 8.1 Acidity and basicity constants* at 298.15 K (2)
Acid/base
Kb
pKb
Ka
pKa
Weakest weak acids
Acid/base
Kb
pKb
Ka
pKa
13.90
7.7  10
9.37
13.90
8.75
9.37
7.97
8.75
5.98
7.97
5.79
5.98
5.77
5.79
4.75
5.77
4.19
4.75
3.44
4.19
3.27
3.44
3.19
3.27
2.99
3.19
2.3  101
7.7  10 6
5.6  105
2.3  10 7
9.1  106
5.6  10 8
1.0  107
9.1  10 9
6.3  108
1.0  10 9
5.9  109
6.3  10 10
5.6  109
5.9  10
10
1.5  1010
5.6  10
11
2.8  1010
1.5  10
11
1.9  1011
2.8  10
11
1.5  1011
1.9  10
11
1.0  1011
1.5  10
5
4.63
0.10
5.35
4.63
6.03
5.35
8.02
6.03
8.21
8.02
8.23
8.21
9.25
8.23
9.81
9.25
10.56
9.81
10.73
10.56
10.81
10.73
11.01
10.81
1.0  10
11
11.01
Weakest
Weakest weak
weak bases
acids
Urea, CO(NH2)2
Weakest weak bases
Aniline, C6H5NH2
Urea, CO(NH2)2
Pyridine, C5H5N
Aniline, C6H5NH2
Hydroxylamine, NH2OH
Pyridine, C5H5N
Nicotine, C10H11N2
Hydroxylamine, NH2OH
Morphine, C17H19O3N
Nicotine, C10H11N2
Hydrazine, NH2NH2
Morphine, C17H19O3N
Ammonia, NH3
Hydrazine, NH2NH2
Trimethylamine, (CH3)3N
Ammonia, NH3
Methylamine, CH3NH2
Trimethylamine, (CH3)3N
Dimethylamine, (CH3)2NH
Methylamine, CH3NH2
Ethylamine, C2H5NH2
Dimethylamine, (CH3)2NH
Triethylamine, (C2H5)3N
Ethylamine, C2H5NH2
Strongest weak bases
Triethylamine, (C2H5)3N
1.3  10
14
4.3  1014
1.3  10 9
1.8  1010
4.3  10 8
1.1  109
1.8  10 6
1.0  108
1.1  10 6
1.6  106
1.0  10 6
1.7  106
1.6  10 5
1.8  106
1.7  10
5
6.5  105
1.8  10
4
3.6  105
6.5  10
4
5.4  104
3.6  10
4
6.5  104
5.4  10
3
1.0  104
6.5  10
10
1.0  10
3
2.99
1
0.10
* Values for polyprotic acids—those capable of donating more than one proton—refer to the first deprotonation.
Strongest weak bases
† The proton transfer equilibrium is B(OH)3(aq)  2 H2O(l) e H3O(aq)  B(OH)4– (aq).
* Values for polyprotic acids—those capable of donating more than one proton—refer to the first deprotonation.

–
p.182
Work through these
• Example 8.1, p.176
• Example 8.2, p.176f
• Do any of you need help working
I
C
E
problems?
Polyprotic acids
• H2A(aq) + H2O(l)  H3O+(aq) + HA-(aq)
Ka1 =
aH O+ aHA3
aH2A
• HA-(aq) + H2O(l)  H3O+(aq) + A2-(aq)
Ka2 =
aH O+ aA23
aHA-
• Work Example 8.3 & Example 8.4 (p.178f)
Polyprotic acids
• Work Example 8.3 (p.178)
• Work Example 8.4 (p.178f)
• Understand the graphs in Fig 8.1 & 8.2
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
Table 8.2 Successive acidity constants
of polyprotic acids at 298.15 K
Acid
Ka1
pKa1
Ka2
Carbonic acid, H2CO3
4.3  10
Hydrosulfuric acid, H2S
pKa2
7
6.37
5.6  10
11
10.25
1.3  10
7
6.88
7.1  10
15
14.15
Oxalic acid, (COOH)2
5.9  10
2
1.23
6.5  10
5
4.19
Phosphoric acid, H3PO4
7.6  10
3
2.12
6.2  10
8
7.21
Phosphorous acid, H2PO3
1.0  10
2
2.00
2.6  10
7
6.59
Sulfuric acid, H2SO4
Strong
1.2  10
2
1.92
Sulfurous acid, H2SO3
1.5  10
2
1.81
1.2  10
7
6.91
Tartaric acid, C2H4O2(COOH)2
6.0  10
4
3.22
1.5  10
5
4.82
Ka3
pKa3
2.1  10
13
12.67
p.177
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”?
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)
pH curve
acid
Fig 8.3 (181)
Titration of a
strong acid with a
strong base
base
pH curve
acid
Fig 8.4 (181)
Titration of a weak
acid with a strong
base
base
pH curve
base
Fig 8.5 (183)
Titration of a weak
base with a strong
acid
acid
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
• 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
pH curve of a buffer solution
base
Buffer action
Consider an aqueous solution of CH3COOH
CH3COOH(aq) + H2O(l)
acid
CH3COO-(aq) + H3O+
conjugate
base
Consider an equimolar mixture of CH3COOH and CH3COO-Na+
Add strong acid
H+ (aq) + CH3COO- (aq)
Add strong base
OH- (aq) + CH3COOH (aq)
CH3COOH (aq)
CH3COO- (aq) + H2O (l)
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
Acid-base indicators
HIn (aq)
H+ (aq) + In- (aq)
[HIn]
 10 Color of acid (HIn) predominates
[In ]
If [HIn] 
[In-],
combination color
[HIn]
1
[In ]
[HIn]
-) predominates
Color
of
conjugate
base
(In

10
[In-]
Table 8.3 Indicator color changes
Indicator
Acid colour
pH range of
colour change
pKIn
Base colour
Thymol blue
Red
1.2–2.8
1.7
Yellow
Methyl orange
Red
3.2–4.4
3.4
Yellow
Bromophenol blue
Yellow
3.0–4.6
3.9
Blue
Bromocresol green
Yellow
4.0–5.6
4.7
Blue
Methyl red
Red
4.8–6.0
5.0
Yellow
Bromothymol blue
Yellow
6.0–7.6
7.1
Blue
Litmus
Red
5.0–8.0
6.5
Blue
Phenol red
Yellow
6.6–8.0
7.9
Red
Thymol blue
Yellow
9.0–9.6
8.9
Blue
Phenolphthalein
Colourless
8.2–10.0
9.4
Pink
Alizarin yellow
Yellow
10.1–12.0
11.2
Red
Alizarin
Red
11.0–12.4
11.7
Purple
• Fig 8.7 (194)
• Strong acid, strong
base titration
• Which indicators
work well here?
• Why will several
indicators be
adequate here?
acid
pH curve
base
• Fig 8.8 (194)
• Weak acid, strong
base titration
• Which indicators
work well here?
• Why will bromothymol blue not
work well here?
acid
pH curve
base
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 = aM+ a X-
or
Ks = [M+][X-]
The solubility constant
Ks varies according to the formula of the salt
Ks = [Ag+][Cl-]
AgCl (s)
Ag+ (aq) + Cl- (aq)
MgF2 (s)
Mg2+ (aq) + 2F- (aq)
Ks = [Mg2+][F-]2
2Ag+ (aq) + CO32- (aq)
Ks = [Ag+]2[CO32-]
3Ca2+ (aq) + 2PO43- (aq)
Ks = [Ca2+]3[PO33-]2
Ag2CO3 (s)
Ca3(PO4)2 (s)
Relationship between Ks and Molar Solubility (s)
Compound
Ks expression
Cation
Anion Relation between Ks & s
Table 8.4 Solubility constants at 298.15 K (1)
Compound
Formula
Ks
Aluminium hydroxide
Al(OH)3
1.0  10
33
Antimony sulfide
Sb2S3
1.7  10
93
Barium carbonate
BaCO3
8.1  10
9
fluoride
BaF2
1.7  10
6
sulfate
BaSO4
1.1  10
10
Bismuth sulfide
Bi2S3
1.0  10
97
Calcium carbonate
CaCO3
8.7  10
9
fluoride
CaF2
4.0  10
11
hydroxide
Ca(OH)2
5.5  10
6
sulfate
CaSO4
2.4  10
5
CuBr
4.2  10
8
chloride
CuCl
1.0  10
6
iodide
CuI
5.1  10
12
sulfide
Cu2S
2.0  10
47
Cu(IO3)2
1.4  10
7
oxalate
CuC2O4
2.9  10
8
sulfide
CuS
8.5  10
45
Copper(I) bromide
Copper(II) iodate
Table 8.4 Solubility constants at 298.15 K (2)
Compound
Formula
Ks
Iron(II) hydroxide
Fe(OH)2
1.6  10
14
FeS
6.3  10
18
Iron(III) hydroxide
Fe(OH)3
2.0  10
39
Lead(II) bromide
PbBr2
7.9  10
5
chloride
PbCl2
1.6  10
5
fluoride
PbF2
3.7  10
8
iodate
Pb(IO3)2
2.6  10
13
iodide
PbI2
1.4  10
8
sulfate
PbSO4
1.6  10
8
sulfide
PbS
3.4  10
28
ammonium phosphate
MgNH4PO4
2.5  10
13
carbonate
MgCO3
1.0  10
5
fluoride
MgF2
6.4  10
9
hydroxide
Mg(OH)2
1.1  10
11
sulfide
Magnesium
Table 8.4 Solubility constants at 298.15 K (3)
Compound
Formula
Ks
Mercury(I) chloride
Hg2Cl2
1.3  10
18
Hg2I2
1.2  10
28
HgS black:
1.6  10
52
red:
1.4  10
53
iodide
Mercury(II) sulfide
Nickel(II) hydroxide
Ni(OH)2
6.5  10
18
Silver bromide
AgBr
7.7  10
13
carbonate
Ag2CO3
6.2  10
12
chloride
AgCl
1.6  10
10
hydroxide
AgOH
1.5  10
8
iodide
AgI
1.5  10
16
sulfide
Ag2S
6.3  10
51
Zn(OH)2
2.0  10
17
ZnS
1.6  10
24
Zinc hydroxide
sulfide
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)
CH3COOH (aq)
Na+ (aq) + CH3COO- (aq)
H+ (aq) + CH3COO- (aq)
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)
Solubility Rules for Common Ionic Compounds in water at 25°C
Soluble Compounds
Exceptions
Compounds containing
alkali metal ions & NH4+
NO3-, HCO3-, ClO3Cl-, Br-, ISO4
2-
Halides of Ag+, Hg22+, Pb2+
Sulfates of Ag+, Ca2+, Sr2+, Ba2+,
Hg2+, Pb2+
Insoluble Compounds
Exceptions
CO32-, PO43-, CrO42-, S2-
Compounds containing alkali
metal ions and NH4+
OH-
Compounds containing alkali
metal ions and Ba2+
You should know these!
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
Key
Ideas
Key
Ideas
The End
…of this chapter…”