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Applications of Aqueous Equilibria
log fun with Ka:Henderson-Hasselbalch Equation
+ A-H+(aq) (aq)HA (aq)
Ka =
[H+] [A-][HA]
log Ka = log [H+] + log[A-]
[HA]
= [A-][HA]- log [H+] - log Ka log+
= [A-][HA]pH pKa log+
Ka = [H+] x [A-]
[HA]
= [conj.base ][acid]
pH pKa log+
Buffer Solutions
• contain a weak acid (or base) and its conjugate in a salt form.
• resist changes in pH upon the addition of small amounts of either acid or base
Example
Which of the following are buffer systems?
(b) KBr/ HBr
(a) KF / HF
(c) Na2CO3/ NaHCO3
yes; HF is a weak acid and F- is its conjugate base
no; HBr is a strong acid
yes; HCO3- is a weak acid and
CO32- is its conjugate base
(d) 0.5 mol HCl added to 1 mol KF solution
yes H+ + F- → HFI 0.5 1 0E 0 0.5 0.5
How a sodium acetate/ acetic acid buffer buffers
OAcH AcO-
the chemistry
AcO- OAcH
OH- H+
Weak acidIts conjugate base
in salt formbuffers have ~equal amounts of these to
start with
add baseadd acid
the mathematics
Ka =
[H+] [AcO-][HOAc]
How a sodium acetate/ acetic acid buffer buffers
= [AcO- ]
[HOAc]pH pKa log+
What pH’s do you get for this buffer when[AcO-]/[HOAc] = 1:1, 2:1, 3:1, 4:1, 5:1, 1:2, 1:3, 1:4? Group # 1 2 3 4 5 6 7 8
Rewriting in the form of Henderson-Hasselbalch equationWhen buffered (initial amounts of acid and its
conjugate base are present) this term stays pretty
constant even when acids or bases are added.
= 1.8 x 10-5
Calculate the pH of two 1.00-L beakers of water after adding a) 0.100 mol HCl to one and b) 0.100 mol
NaOH to the other. (both are strong so use stoichy!)
a) pH = 1
UNBUFFERED SOLUTIONS
+ Cl-H+ (aq) (aq)0.100 M 00
HClI
pH = -log[H+] = -log(0.100)
pOH = 1
+ OH-Na+(aq) (aq)0.100 M 00
NaOHI
pOH = -log[OH-] = -log(0.100)
E 0.100 M 0.100 M0
E 0.100 M 0.100 M0
pH + pOH = 14 b) pH = 13
BUFFERED SOLUTIONS
1. HOAc/OAc- buffer alone
2. #1 with 0.10 mol HCl added
3. #1 with 0.10 mol NaOH added
3 Scenarios to contrast with our unbuffered examples:
Calculate the pH of a buffer containing 1.0 M acetic acid and 1.0 M sodium acetate.
+ AcO-H+ (aq) (aq)OAc (aq)H
pH = 4.74
1.0 M 1.0 M+x-x + x
BUFFERED SOLUTIONS
Ka = [H+] [AcO-]
[HOAc]1.8 x 10-5 =
= [conj.base ][acid]
pH pKa log+
=(1.0 M + x)
pH -log(1.8 x 10-5) log+ (1.0 M - x)
1.0 M 1.0 M0IE
pH = 4.74What happens when 0.10 mol of gaseous HCl is added
to 1.0 L of this solution?
Calculate the pH of a buffer containing 1.0 M acetic acid and 1.0 M sodium acetate.
0.1 mol 1.0 mol 1.0 mol
≅ 0.0 mol 0.9 mol 1.1 mol
Key 1st step: Strong acid + weak base (OAc-) goes to completion. Use stoichiometry with acid/base
neutralization reaction.+ AcO-H+ (aq) (aq) OAc (aq)H
I
E
+H+ (aq) AcO-(aq) OAc (aq)H≅ 0.0 mol 0.9 mol 1.1 mol
Ka = [H+] [AcO-]
[HOAc] 1. 8 x 10-5 = x (0.9 + x)(1.1 - x)
x = [H+] = 2.2 x 10-5 M pH = 4.66
OAc (aq)H AcO-(aq) H+ (aq)+
E
E
I 1.1 mol 0.9 mol1L 1L
0 M
1.1 M - x 0.9 M + x +x
Key 2nd step: Now let system reach equilibrium using these as initial conditions in the acid dissociation (Ka) reaction.
Same calculation using the Henderson-Hasselbalch equation
= [c.base ][acid]pH pKa log+
= (0.9 + x)(1.1 - x)pH - log Ka log+
pH = 4.74 - 0.09 = 4.65
OAc (aq)H AcO- (aq) H+ (aq)+1.1 mol - x 0.9 mol + xE +x
Calculate the pH of a buffer containing 1.0 M acetic acid and 1.0 M sodium acetate.
pH = 4.74
What happens when 0.10 mol of gaseous HCl is added to 1.0 L of this solution?
pH = 4.66
What happens when 0.10 mol of solid NaOH is added to 1.0 L of this solution?
add 0.10 mol of NaOH(s) to 1.0 L of a solution that is 1.0 M in NaOAc and HOAc
+ AcO-OH- OAcH
0.1 mol 1.0 mol 1.0 mol
≅ 0.0 mol 0.9 mol 1.1 mol
+H2O
Key 1st Step: Strong base + weak acid goes to completion. Use stoichiometry with acid/base neutralization reaction.
Ka = [H+][AcO- ]
[HOAc] 1.8 x 10-5 = [H+](0.9 - x)
(1.1 + x)
[H+] = 1.47 x 10-5 M pH = 4.83
+ AcO-OH- OAcH≅ 0.0 mol 0.9 mol 1.1 mol
+H2O
AcO-OAcH +H+
E 0.9 mol 1.1 molx1L 1L
- x + x
Key 2nd step: Now let system reach equilibrium using these as initial conditions for the acid dissociation (Ka) reaction.
Same calculation using the Henderson-Hasselbalch equation
= [c.base ][acid]pH pKa log+
= (0.9 - x)(1.1 + x)pH - log (1.8 x 10-5) log+
pH = 4.74 + 0.09 = 4.83
0.9 mol - x 1.1 mol + xAcO-OAcH +H+
+xE
Calculate the pH of a buffer containing 1.0 M acetic acid and 1.0 M sodium acetate.
pH = 4.74
What happens when 0.10 mol of gaseous HCl is added to 1.0 L of this solution?
pH = 4.66
What happens when 0.10 mol of solid NaOH is added to 1.0 L of this solution?
pH = 4.83Compare to pH = 1 and pH = 13 for unbuffered solution!
Show the fraction of each species present as a function of pH
Distribution Curves
acid and conjugate base
base and conjugate acid
Buffer range = pKa ± 1.00
CH3COOH
100
75
25
50
0 1 2 3 4 5 6 7 8 9 10 11 12 13 14distribution curves for acetic acid and
acetate ion as a function of pH
pH = pKa = 4.74
pH
%
CH3COO -
It is important because it contributes to controlling the pH of blood at 7.4.
The bicarbonate/carbonate acid buffer is more complicated because carbonic acid is
a diprotic acid.
pH
100
75
25
50
0 1 2 3 4 5 6 7 8 9 10 11 12 13 14
pKa =
pH =
6.83 10.32pKa =pH =
distribution curves for carbonic acid and bicarbonate ion as a function of
pH
CO32-HCO3
-H2CO3
%
Preparing a Buffer Solution with a Specific pH
= [c.base ][acid]
pH pKa log+
• center of buffer range: [acid] = [conj. base]What’s true about pH then?
• choose a weak acid/conjugate base combination where the pKa of the acid
corresponds to the pH at which you want the solution buffered
pH = pKa
How would you prepare a liter of “carbonate buffer” at a pH of 10.10 ?
you have available:
H2CO3
NaHCO3
Na2CO3
pKa 6.38
pKa 10.32
Pick NaHCO3 and its conjugate base.
How would you prepare a liter of “carbonate buffer” at a pH of 10.10 ?
=[c.base ]
[acid]pH pKa log+
=[Na2CO3
]
[NaHCO3]10.10 10.32 log+
[Na2CO3 ]
[NaHCO3]= 10-0.22 = 0.60
Acid-Base Titrations
Titration
a solution of known concentration, called a standard solution is added gradually to another solution of
unknown concentration until the chemical reaction between the two solutions is complete
equivalence point• the exact number of moles required for the
reaction
• detected by a color change in an acid-base indicator or electronically
Acid-Base Titrations
How does the pH change during the course of the following kinds of titration?
strong acid-strong base
weak acid-strong base
weak base-strong acid
Titrations Involving aStrong Acid and a
Strong Base
Titrations Involving a Strong Acid and a Strong Base
titrate 25.0 mL of 0.100 M HCl with 0.100 M NaOH
mmol molNote:mL and L
are equivalent
And for strong w/ strong, simply use stoichy! No Keq.
Titrations Involving a Strong Acid and a Strong Base
What is the pH after the addition of 10.0 mL of 0.100 M NaOH to 25.0 mL of 0.100 M HCl ?
0.1mmol NaOH10ml x
mL= 1.0mmol OH-
0.1mmol HCl25ml x
mL= 2.5mmol H+
H2O (l )H+ (aq) + OH- (aq)
1.0mmol2.5mmol H+I 0
1.0mmol1.5mmol H+E 0
HCl (aq) NaCl (aq)H2O (l ) ++ NaOH(aq)
Titrations Involving a Strong Acid and a Strong Base
What is the pH after the addition of 10.0 mL of 0.100 M NaOH to 25.0 mL of 0.100 M HCl ?
H2O (l )H+ (aq) + OH- (aq)
1.5mmol H+
35.0 mL (solution)[H+] = = 4.28 x 10-2 M
pH = 1.37
Titrations Involving a Strong Acid and a Strong Base
What is the pH after the addition of 25.0 mL of 0.100 M NaOH to 25.0 mL of 0.100 M HCl ?
0.1mmol NaOH25ml x
mL= 2.5mmol OH-
0.1mmol HCl25ml x
mL= 2.5mmol H+
H2O (l )H+ (aq) + OH- (aq)
2.5 mmol2.5 mmol H+I 0
2.5 mmol0E 0
Titrations Involving a Strong Acid and a Strong Base
What is the pH after the addition of 25.0 mL of 0.100 M NaOH to 25.0 mL of 0.100 M HCl ?
[H+] = 1.0 x 10-7 M
pH = 7.0
H2O (l ) + H3O+ (aq)+ OH- (aq)H2O (l )
pH versus ml 0.10 M NaOH added
(NaOH)volume
(total)volume
(HCl)mmol
M[H+] pH
0 25.0ml 2.5 0.10 1.0010.0 ml 35.0 ml 1.5 0.043 1.37
20.0 ml 45.0 ml 0.5 0.011 1.95
24.0 ml24.5 ml25.0 ml
49.0 ml49.5 ml50.0 ml
0.1
0.050
0.0020.00110-7
2.693.007.00
_
_
_
_
15
10
5
0
25 500
_ __ __
(NaOH)volume
pH7.0 at
equivalence point
Titration of 25.00 ml of
0.100 M HCl with 0.100 M
NaOH
Titrations Involving a Strong Acid and a Strong Base
What is the pH after the addition of 35.0 mL of 0.100 M NaOH to 25.0 mL of 0.100 M HCl ?
0.1mmol NaOH35ml x
mL= 3.5mmol OH-
0.1mmol HCl25ml x
mL= 2.5mmol H+
H2O (l )H+ (aq) + OH- (aq)
3.5 mmol2.5 mmolstart 0
2.5 mmol0end 1.0 mmol
Titrations Involving a Strong Acid and a Strong Base
H2O (l )H+ (aq) + OH- (aq)
1.0 mmol OH-
60.0 mL (solution)[OH- ] = = 1.66 x 10-2 M
pH = 12.22
pOH = 1.78
What is the pH after the addition of 35.0 mL of 0.100 M NaOH to 25.0 mL of 0.100 M HCl ?
pH versus ml 0.10 M NaOH added
(NaOH)volume
(total)volume
(HCl)mmol
M[H+] pH
25.0 ml 50.0ml 0 10-7 7.00
25.5 ml 50.5 ml 0.05 0.001 11.0026.0 ml 51.5 ml 0.10 0.002 11.2930.0 ml40.5 ml50.0 ml
55.0 ml65.0 ml75.0 ml
0.501.50
2.50
0.0090.0230.033
11.9612.3612.52
(NaOH)mmol
M[OH-]
_
_
_
_
15
10
5
0
25 500
_ __ __
(NaOH)volume
pH
Titration of 25.00 ml of
0.100 M HCl with 0.100 M
NaOH
7.0 at equivalence
point
Sketch the titration curve if strong base is in the beaker and strong acid is in the buret. _
_
_
_
13
10
5
0
25 500
_ __ __
(HCl)volume
pH7.0 at
equivalence point
Titration of 25.00 ml of
0.100 M HCl with 0.100 M
NaOH
Titrations Involving a weak Acid and a Strong Base
Calculate the pH in the titration of 25.0 mL of 0.100 M acetic acid with 0.100 M NaOH
after adding (a) 10.0 mL of 0.100 M NaOH(b) 25.0 mL of 0.100 M NaOH(c) 35.0 mL of 0.100 M NaOH
Titrations Involving a weak Acid and a Strong Base
CH3COOH H2OCH3COONa ++ NaOH
CH3COOH H2OCH3COO - ++ OH -
Molecular Eqn
Net Ionic Eqn
Titrations Involving a weak Acid and a Strong Base
CH3COOH H2OCH3COO - ++ OH -
1.0mmol2.5mmol 0 0
1.0mmol1.5mmol 0 1.0mmol
Calculate the pH in the titration of 25.0 mL of 0.100 M acetic acid with 0.100 M NaOH after adding (a) 10.00 mL of 0.100 M NaOH. Key 1st step: Strong base (sodium hydroxide) + weak acid
(acetic) goes to completion. Use stoichiometry with neutralization reaction.
Key 2nd step: Now let system reach equilibrium using these as initial conditions for the acid dissociation (Ka) reaction.
1.0mmol + x1.5mmol - x +x
I
E
CH3COOH CH3COO - + H+
Ka = 1.8 x 10-5
E
Calculate the pH in the titration of 25.0 mL of 0.100 M acetic acid with 0.100 M NaOH after adding (a) 10.00 mL of 0.100 M NaOH.
Titrations Involving a weak Acid and a Strong Base
1.0mmol1.5mmol
35.0 mL 35.0 mL= 0.0043 M 0.0028 M
pH = 4.57
Plan: Assume x is small since Ka is small and turn mmol into M .
= (0.0043 - x)(0.0028 + x)
pH - log (1.8 x 10-5) log+
1.0mmol + x1.5mmol - x +xCH3COOH CH3COO - + H+
Ka = 1.8 x 10-5
E
Calculate the pH in the titration of 25.0 mL of 0.100 M acetic acid with 0.100 M NaOH after adding (b) 25.0 mL of 0.100 M NaOH.
Titrations Involving a weak Acid and a Strong Base
CH3COOH H2OCH3COO - ++ OH -
2.5mmol2.5mmol 0 02.5mmol0 0 2.5mmol
50.0 mL= 5.0 x 10-2 M
2.5mmol
H2OCH3COO - + CH3COOH + OH - 5.6 x 10-10Kb =
Key 1st step: Strong base (sodium hydroxide) + weak acid (acetic) goes to completion. Use stoichiometry with neutralization reaction.
Key 2nd step: Now let system reach equilibrium using these as initial conditions for the base dissociation (Kb) reaction.
IE
Calculate the pH in the titration of 25.0 mL of 0.100 M acetic acid with 0.100 M NaOH after adding
(b) 25.0 mL of 0.100 M NaOH.
Titrations Involving a weak Acid and a Strong Base
H2OCH3COO - + CH3COOH + OH -
5.0 x 10-2 M - x x x
(5.0 x 10-2 - x)5.6 x 10-10 =
x2
x = 5.29 x 10-6 [OH -] =
pOH = 5.27
pH = 8.72
E
Calculate the pH in the titration of 25.0 mL of 0.100 M acetic acid with 0.100 M NaOH after adding
(c) 35.0mL of 0.100 M NaOH.
Titrations Involving a weak Acid and a Strong Base
CH3COOH H2OCH3COO - ++ OH -
3.5mmol2.5mmol 0 02.5mmol0 1.0mmol 2.5mmol
1.0mmol OH -
60.0 mL (solution)
= 1.67 x 10-2 MpH = 12.22
pOH = 1.78
strong base!
Strong base dominates, so we can ignore weak base and directly determine pH.
IE
pH versus ml 0.10 M NaOH added
(NaOH)volume
(total)volume
(NaOAc)mmol [H+] pH
0 ml 25.0ml 0 1.3 x 10-3 2.87
(HOAc)mmol
2.5
10.0 ml 35.5 ml 1 2.7 x 10-5 4.571.5
(NaOH)mmol [OH-]
25.0 ml 50.0 ml 2.5 8.720 1.9 x 10-9
35.0 ml 60.0 ml 1 12.220 1.7 x 10-2
_
_
_
_
15
10
5
2.5
25 500
_ __ __
(0.10 M NaOH)volume
pH
Titration of 25.0 ml of
0.10 M acetic acid with 0.10
M NaOH
_
_7.5
12.5
_
++ +
++
+++++++
+
+
++++
++++++
+4.75
8.72
Maximum bufferWhy?
Why?
7
Titrations Involving a Strong Acid and a Weak Base
Calculate the pH in the titration of 25.0 mL of 0.100 M HCl with 0.100 M NH3 after adding
(a) 10.0 mL of 0.100 M NH3(b) 25.0 mL of 0.100 M NH3(c) 35.0 mL of 0.100 M NH3
Titrations Involving a Strong Acid and a Weak Base
HCl Cl -NH4+ ++ NH3
H + NH4++ NH3
Calculate the pH in the titration of 25.0 mL of 0.100 M HCl with 0.100 M NH3 after adding (a) 10.00 mL of 0.100 M NH3.
Titrations Involving a Strong Acid and a Weak Base
H + NH4++ NH3
1.0mmol2.5mmol 0
1.0mmol1.5mmol 0
1.5 mmol H + pH = 1.3735.0 mL
= 4.29 x 10-2M[H +] =
E
I
Key 1st step: Strong acid (HCl) + weak base (ammonia) goes to completion. Use stoichiometry with neutralization reaction.
strong acid!
Strong acid dominates, so we can ignore weak acid NH4+ and directly determine pH.
2.5mmol
50.0 mL= 5.0 x 10-2M
Calculate the pH in the titration of 25.0 mL of 0.100 M HCl with 0.100 M NH3 after adding (b) 25.0 mL of 0.100 M NH3.
Titrations Involving a Strong Acid and a Weak Base
H + NH4++ NH3
2.5 mmol2.5mmol 02.5 mmol00E
I
Key 1st step: Strong acid (HCl) + weak base (ammonia) goes to completion. Use stoichiometry with neutralization reaction.
Key 2nd step: Now let system reach equilibrium using these as initial conditions for the acid dissociation (Ka) reaction.
NH3 H+Ka = 5.6 x 10-10
NH4+ +
E 2.5mmol - x +x +x
Calculate the pH in the titration of 25.0 mL of 0.100 M HCl with 0.100 M NH3 after adding (b) 25.0 mL of 0.100 M NH3.
Titrations Involving a Strong Acid and a Weak Base
(5.0 x 10-2 - x)5.6 x 10-10 =
x2
x = 5.29 x 10-6 [H+] =pH = 5.28
NH3 H+Ka = 5.6 x 10-10
NH4+ +
E 2.5mmol - x +x +x
Calculate the pH in the titration of 25.0 mL of 0.100 M HCl with 0.100 M NH3 after adding (c) 35.0mL of 0.100 M NH3.
Titrations Involving a Strong Acid and a Weak Base
3.5 mmol2.5mmol 0
H + + NH3 NH4+
2.5 mmol0 1.0 mmol
NH3 H+
Ka = 5.6 x 10-10
NH4+ +
2.5mmol 1.0mmol
60.0 mL= 4.1 x 10-2M
60.0 mL= 1.66 x 10-2M
Calculate the pH in the titration of 25.0 mL of 0.100 M HCl with 0.100 M NH3 after adding (c) 35.0mL of 0.100 M NH3.
Titrations Involving a Strong Acid and a Weak Base
NH3 H+
Ka = 5.6 x 10-10
NH4+ +
4.16 x 10-2 1.66 x 10-2
(4.16 x 10-2)=
1.66 x 10-2
(5.6 x 10-10 )= 1.4 x 10-9
pH = 8.85
[NH4+]
[H+] =[NH3]
Ka
pH versus ml 0.10 M NH3 added
(NH3)
volume (total)volume [H+] pH
(HCl)mmol
0 ml 25.0ml 0.1 1.02.5
10.0 ml 35.0 ml 0.0429 1.371.5
25.0 ml 50.0 ml 5.282.5 0
35.0 ml 60.0 ml 8.8512.5
(NH4+ )
mmol (NH3
)mmol
+
_
_
_
_
10
5
2.5
20 400
_ __ _
(0.10 M NH3)volume
pH
Titration of 25.0 ml of
0.10 M HCl acid with 0.10
M NH3
_
_
7.5
12.5
_
+ + ++
+
++++
+++
60 80
+++++
++++
5.36 at equivalence point
+
_
_
_
_
10
5
2.5
20 400
_ __ _
(0.10 M NH3)volume
pH
_
_
7.5
12.5
_
+ + ++
+
++++
+++
9.25 when [NH3] = [NH4+]
60 80
+++++
++++
Titration of 25.0 ml of
0.10 M HCl acid with 0.10
M NH3
Solubility Equilibria
Applies the principles of equilibrium to ionic solids of low solubility in water
• Milk of Magnesia demo
Solubility product (Ksp)
in a saturated solution, the species in solution is in equilibrium with undissolved
material
is characterized by an equilibrium constant Ksp called the solubility product
Ag+AgCl (s) (aq) + Cl- (aq)
Examples:
= [Cl-]Ksp [Ag+] = 1.6 x 10-10
2Ag+Ag2CO3 (s) (aq) + CO32- (aq)
= [CO32-]Ksp [Ag+]2 = 8.1 x 10-12
3Ca2+Ca3(PO4)2 (s) (aq) + 2PO43- (aq)
= [PO43-]2 Ksp [Ca2+] 3 = 1.2 x 10-26
When ions in solution reform solid they do so on the surface of the solid.
NaCl ( s )
Na+Cl- Cl-Na+
The amount of excess solid present has no effect on the position of solubility equilibrium.
NaCl ( s )
Na+Cl- Cl-Na+
Doubling the surface area of the solid doubles both the rate of reforming and dissolving. position of equilibrium unchanged
Molar Solubility and Solubility
Molar solubility: the number of moles solute in one liter of a saturated solution (mol/L)
Solubility: the number of grams of solute in one liter of a solution (g/L)
Depending on the text being used
Solubility v.s. Solubility Product
• solubility product is an equilibrium constant and thus has only one value at a certain temperature
• solubility is an equilibrium position and has an infinite number of possible values at a given
temperature depending on the conditions
ie: the common ion effect
given Ksp you should be able to calculate
molar solubility
concentration of anion in a saturated solution
concentration of cation in a saturated solution
molar solubilities when common-ion effect comes into play
Practice Exercise
The solubility of PbCrO4 is 4.5 x 10-5 g/L. What is its solubility product ?
Pb2+PbCrO4 (s) + CrO42- (aq)(aq)
4.5 x 10-5g
Lx
323.2 g
1 mol= 1.4 x 10-7 mol/L
First convert to mol/L
Practice Exercise
The solubility of PbCrO4 is 4.5 x 10-5 g/L. What is its solubility product ?
Pb2+PbCrO4 (s) + CrO42- (aq)(aq)
1.4 x 10-7 M 1.4 x 10-7 M
[CrO42 -] = Ksp [Pb2+]
(1.4 x 10-7) (1.4 x 10-7)= Ksp
= 1.9 x 10-14Ksp
Practice Exercise
Calculate the molar solubility of PbCO3. (Ksp = 3.3 x 10-14 )
Pb2+PbCO3 (s) + CO32- (aq)(aq)
E +x +x
I 0.00 M 0.00 M
[CO32 -] = Ksp [Pb2+]
(x) (x)= Ksp = (x)2
x = 1.8 x 10-7mo l /L
3.3 x 10-14=
Practice Exercise
The Ksp value for Cu(OH)2 = 2.2 x 10-20 at 25ºC. Calculate its solubility in g/L.
Cu2+Cu(OH)2 (s) + 2OH- (aq)(aq)
E +x +2xI 0.00 M 0.00 M
[OH-]2 = Ksp [Cu2+]
(x)(2x)2= Ksp = 4x3 2.2 x 10-20== (x)(4x2)
4x3 2.2 x 10-20=
2.2 x 10-20
4( (1/3
x =
x = 1.77 x 10-7 molL
x97.55 g
1 mol
x = 1.72 x 10-5 g/L
Relative Solubilities
CaSO4
SaltAgCl
CuI
Ksp
8.3 x 10-17
5.1 x 10-12
6.1 x 10-5
all Ksp expressions are of the form[Y ] = Ksp [X ]
Therefore, solubility decreases in the order
CaSO4 AgClCuI> >
Relative Solubilities
Bi2S3
SaltCuSAg2S
Ksp
8.3 x 10-45
1.6 x 10-49
1.1 x 10-73
Therefore, need to do a calculation
Ksp expressions are of the form[Cu2+][S2-], [Ag+]2[S2-], [Bi3+]2[S2-]3
Relative Solubilities
CuS
8.3 x 10-45 [Cu2+][S2-]
Cu2+ + S2 - (aq)(aq)
= = x2
Ag2S
1.6 x 10-49 [Ag+]2 [S2-]
2Ag+ + S2 - (aq)(aq)
= = 4x3
Bi2S3
1.1 x 10-73 [Bi3+]2 [S2-]3
2Bi3+ + 3S2 - (aq)(aq)
= = 108x5
Relative Solubilities
Bi2S3
Salt
CuS
Ag2S
Ksp
8.3 x 10-45
1.6 x 10-49
1.1 x 10-73
Molar solubility
9.2 x 10-23
3.4 x 10-17
1.0 x 10-15
Predicting Precipitation Reactions
Ion Product
Qc > Ksp: precipitate will form
Qc = Ksp: saturated solution
Qc < Ksp: no precipitate will form; substance dissolves
Q = [products] x [reactants] y
• analogous to reaction quotient
• tells us whether a precipitate will form under a given set of reaction conditions
Practice Exercise
If 2.00 ml of 0.200 M NaOH are added to 1.0 L of 0.100 M CaCl2, will a
precipitate form?
Practice Exercise
If 2.00 ml of 0.200 M NaOH are added to 1.0 L of 0.100 M CaCl2, will a
precipitate form?= 8.0 x 10-6 Ksp
Ca2+(aq)Ca(OH)2 (s) + 2OH- (aq)
Ca2+(aq)Ca(OH)2 (s) + 2OH- (aq)
Ca2+ = 1.0 L x0.10 mol
L= .1 mol
OH- = 2.0 x 10-3 L x0.20 mol
L= 4.0 x 10-4 mol
.1 mol
1.002 L
4.0 x 10-4 mol
1.002 L
[Ca2+] = 0.100 M [OH ] = 0.0004 M
Ca2+(aq)Ca(OH)2 (s) + 2OH- (aq).100 M .0004 M
[OH-]2 = Q [Ca2+]
< KspQ
(.0004 )2 (.100 )= Q
= 8.0 x 10-6 Ksp
= 1.6 x 10-8 Q
; precipitate will not form
Separation of Ions by Fractional Precipitation
When a solution contains several ions can one be removed by precipitation, leaving
all the others in solution?
Practice Exercise
(a) AgCl
Calculate the concentration of Ag+
necessary to cause precipitation of
(Ksp = 1.6 x 10-10 )
(b) Ag3PO4 (Ksp = 1.8 x 10-18 )
from a solution in which Cl- and PO43-
are each present at a concentration of 0.10 M.
Practice Exercise
Ag+AgCl (s) + Cl- (aq)(aq)
Ksp 1.6 x 10-10=[Cl-] = 0.1 M ;
3Ag+Ag3PO4 (s) + PO43- (aq) (aq)
Ksp 1.8 x 10-18=[PO43- ] = 0.1 M ;
Practice Exercise
calculate the concentration of Ag+at which AgCl begins to precipitate
Ag+AgCl (s) + Cl- (aq)(aq)
Ksp 1.6 x 10-10=[Cl-] = 0.1 M ;
1.6 x 10-10 = [Cl-] [Ag+]
1.6 x 10-10 = (0.1) [Ag+]
[Ag+] = 1.6 x 10-9
Practice Exercise calculate the concentration of Ag+ at which
Ag3PO4 begins to precipitate
1.8 x 10-18 = [Ag+] 3 [PO43- ]
1.8 x 10-18 = (0.10) [Ag+]3
[Ag+] = 2.6 x 10-6
3Ag+Ag3PO4 (s) + PO43- (aq) (aq)
Ksp 1.8 x 10-18=[PO43- ] = 0.1 M ;
[Ag+]3 = 1.8 x 10-17
Practice Exercise
[Ag+] to begin precipitation of:
[PO43-] = 2.6 x 10-6 M
[Cl-] = 1.6 x 10-9 M
Who precipitates first?AgCl precipitates before Ag3PO4
Practice Exercise
What is [Cl-] when Ag3PO4 begin precipitate ?
PO43- begins to precipitate when [Ag+]
reaches 2.6 x 10-6 M.
Ksp = 1.6 x 10-10 = [Ag+] [Cl-]
1.6 x 10-10 = (2.6 x 10-6) [Cl-]
[Cl-] = 6.2 x 10-5
The Common Ion Effect and Solubility
A common ion suppresses the solubility of an ionic substance
Practice Exercise Calculate the solubility of AgBr in g/L in:
Ksp = 7.7 x 10-13
Ag+AgBr (s) + Br- (aq)(aq)
(a) pure water
(b) 0.0010 M NaBr
Practice Exercise
Ag+AgBr (s) + Br- (aq)(aq)
(a) pure water
(b) 0.0010 M NaBr
Le Chatelier’s principle: by adding NaBr, the equilibrium should shift to the _____and the
solubility of AgBr is ______.
Calculate the solubility of AgBr in g/L in: Ksp = 7.7 x 10-13
leftless
Practice Exercise
Ag+AgBr (s) + Br- (aq)(aq)
(a) pure water
x2 7.7 x 10-13=x 8.8 x 10-7 M=
[Br-] = Ksp [Ag+] 7.7 x 10-13=
= 1.7 x 10-4 g/Lsolubility of
AgBr
Calculate the solubility of AgBr in g/L in: Ksp = 7.7 x 10-13
[Br-] = Ksp [Ag+] 7.7 x 10-13=
Practice Exercise
Ag+AgBr (s) + Br- (aq)(aq)
(b) 0.0010 M NaBr
[Ag+] 7.7 x 10-10 M=
7.7 x 10-13=[Ag+] ( 0.0010)
= 1.4 x 10-7 g/L solubility of AgBr
Calculate the solubility of AgBr in g/L in: Ksp = 7.7 x 10-13
pH and Solubility
pH and Solubility
Mg2+Mg(OH)2 + 2OH-
Ksp = 1.2 x 10-11 = [Mg2+] [OH-]2
consider Mg(OH)2
solubility decreases in basic solution because increase in [HO-] requires decrease in
[Mg2+]
pH and Solubility
Mg2+Mg(OH)2 + 2 HO-
Ksp = 1.2 x 10-11 = [Mg2+] [HO-]2
consider Mg(OH)2
solubility increases in acidic solution because [HO-] is
decreased
H+
H2O
Practice Exercise
Calculate its molar solubility of Cr(OH)3 (Ksp = 3.0 x 10-29 ) in a pH 10 buffer.
Cr3+Cr(OH)3 (s) + 3OH - (aq)(aq)
= Ksp 3.0 x 10-29 = [OH -]3 [Cr3+]
3.0 x 10-29 = [10-4]3 [Cr3+]
= 3 x 10-17 mol/L[Cr3+]
pOH = 4
In General
salts of weak acids will be more soluble in acid solution than in pure water
because the anion of a weak acid (conjugate base) is basic
pH and Solubility
Ba2+BaF2 + 2F -Ksp = 1.7 x 10-6 = [Ba2+] [F -]2
BaF2 is a salt of a weak acid; its anion (F-) is basic
solubility increases in acid because F - concentration is
decreased
H+
HF
Complex Ion Equilibria and Solubility
“A complex ion is an ion containing a central metal cation bonded to one or more molecules or ions.”
demos - write the equilibria for each stepCuCl2 + NH3 + HCl
AgNO3 + NaCl + NH3 + HCl
kind of like Ksp except....
Formation constant
�
K f =[metal cation][neutral or anion ligand]
[complex ion]
the product is a dissolved ion complex, not a solid precipitate
Cu2+(aq) Cu(NH3)42+(aq)
[Cu(NH3)42+]
[Cu2+] [NH3]4Kf = = 5.0 x 1013
NH3 (aq)
2+
• a complex ion is formed by Lewis acid-Lewis base reaction;
• the bond between the Lewis acid and Lewis base is covalent
• a complex ion is characterized by the formation constant ( Kf )
metal ion is the Lewis acid
the neutral molecule or ion that acts as the Lewis base is called a Ligand
• the number of ligands attached to the metal ion is called the coordination number
Some typical complex ions
complex ion Kf
Ag(NH3)2+ 1.5 x 107Ag+ 2NH3+
Cu2+ 4CN -+
Co3+ 6NH3+
Cu(CN)42- 1.0 x 1025
Co(NH3)63+ 5.0 x 1031
The effect of complex ion formation generally is to increase the solubility of a
substance.
AgCl (s) Ag(NH3)2+ (aq)
[Ag(NH3)2+]
[Ag+] [NH3]2Kf = = 1.5 x 107
NH3 (aq)
Practice exercise
Ag+(aq)AgBr (s) + Br - (aq)
Ksp = 7.7 x 10-13
What is the molar solubility of AgBr in a 1.0 M solution of NH3?
(aq)Ag+ 2NH3 Ag(NH3)2++(aq) (aq)
1.5 x 107Kf =
AgBr (s) 2NH3+ (aq) + Br - (aq)Ag(NH3)2
+ (aq)Kf Ksp = K
AgBr (s) 2NH3+ (aq)
+ Br - (aq)Ag(NH3)2+ (aq)
Practice exercise
What is the molar solubility of AgBr in a 1.0 M solution of NH3?
Kf Ksp = K [Ag(NH3)2
+] [Br-]
[NH3]2 = 12.3 x 10-6 =
12.3 x 10-6 =x2
1.0x = 0.0035 M
the solubility of AgBr in pure water
is 9 x 10-7 M
Some typical coordination numbersCoordination
numbers
Ag+
Cu+
Co3+
Au+
Mn2+
Fe2+
Co2+
Ni2+
Cu2+
Zn2+
Cr3+
Au3+
Sc3+
Coordination numbers
Coordination numbers
2,4
2
2,4
4,6
4,6
4,6
4,6
4,6
6
6
6
6
4
Some common ligands
CN-
NH3
H2O
CH3NH2
CO
NO
F-
Cl-
Br-
I-
SCN-
Indicators