Transcript Slide 1

Exam 1 Stats
• Average 51.8 ± 22.1
• High 92.5 (2 of you)
• Low 14
A
100-83
4
B
82-64
11
C
63-41
15
D
40-22
8
F
21-0
6
Chapter 8
Activity
The “true” nature of
ionic species in solution
• Ions are charged molecules
• As a result, they tend to attract polar
solvent molecules (like water, for instance)
and other ions
– Hydrated radius
– Ionic atmosphere
• The thickness of the ionic atmosphere is a
function of the ionic strength of the
solution “the concentration of charge”
Hydrated radius
_
+
water
_
+
Chloride
Sodium ion
Ionic Atmosphere and Shielding
Low ionic strength
+ +
+ - +
+ + +
High ionic strength
+ +
+ + +
+ + + + +
+
- - + - - -
+
_
Calcium
ion
Sulfate
ion
+
+ +
+ +
+
+
-
Sodium
ion
Chloride
ion
- - - -- +
- - - --
- - - -
Activity
• All ionic species in any equilibrium
expression are more accurately expressed
as activities.
A Ca2+ = [Ca2+]gCa2+
or
A + H2O ↔ A- + H3O+
Ka = A A- A H3O+ / [A] = [A-]gA-[H3O+]gH3O+ / [A]
where g is the activity coefficient and is a
function of the ionic radius of the ion and
the ionic strength of the solution.
Activity - continued
• The higher the ionic strength of the
solution, the larger the smaller the activity
coefficient.
• Rationalization: At higher ionic strengths
the ion cloud around any ion is thicker,
which weakens the attractive forces
between the ion and its counterpart,
inhibiting recombination.
Ionic strength
• A measure of the concentration of ions in
solution
m = ½ ∑ c iz i2
0.010 M Ca(NO3)2
m = ½ ([NO3-](-1)2 + [Ca2+](2+)2) =
½ {(0.02)(-1)2 + (0.01)(2)2 } = 0.03 M
gCa2+@m=0.03
=?
Use extended Debye-Huckle eq or Table 8.1 and
extrapolation
Take home message
• At high ionic strengths, solubility increases
slightly (by a factor of 2-10).
• pH is influenced by ionic strength
• Weak acid and base dissociation is
influenced a little by ionic strength
• Significant figures?
Example 8-12
• Solubility of Hg2Br2
– in pure water
– in 0.00100 M KNO3
– in 0.0100 M KNO3
– in 0.100 M KNO3
In pure water
Hg2Br2  Hg22+ + 2BrKsp = [Hg22+]gHg22+ [Br-]2gBr-2
2[Hg22+] = [Br-] and let [Hg22+] = x
Ionic strength is very low.
gHg22+ and gBr- are close to 1.00
so,
Ksp = 4 [x]3 = 5.6E-23
x = 2.4E-8 M
Activity coefficient as a function of
ionic strength Table 8 -1
m = 0.001
m = 0.01
m = 0.1
Hg22+
0.867
0.660
0.335
Br-
0.964
0.898
0.75
in 0.00100 M KNO3
m = ½ ((.001)(+1)2 + (0.001)(-1)2 = 0.001 M
gHg22+ @ m = 0.001 = 0.867
gBr-2 @ m = 0.001 = 0.964
Ksp = 4 gHg22+ gBr-2 [x]3 = 5.6E-23
x = 2.6E-8M
in 0.0100 M KNO3
m = ½ ((.01)(+1)2 + (0.01)(-1)2 = 0.01 M
gHg22+ @ m = 0.001 = 0.660
gBr-2 @ m = 0.001 = 0.898
Ksp = 4 gHg22+ gBr-2 [x]3 = 5.6E-23
x = 3.0E-8M
in 0.100 M KNO3
m = ½ ((0.1)(+1)2 + (0.1)(-1)2 = 0.1 M
gHg22+ @ m = 0.001 = 0.335
gBr-2 @ m = 0.001 = 0.75
Ksp = 4 gHg22+ gBr-2 [x]3 = 5.6E-23
x = 4.2E-8M
solubility vs log(m)
solubility of Hg2Br2
4.7E-08
4.2E-08
3.7E-08
3.2E-08
2.7E-08
2.2E-08
-8
-6
-4
log(m)
-2
0
pH and ionic strength
• True definition of pH
• pH = -logAH+ = -log {[H+]gH+}
• pH of a 0.00100 M HCl solution
– Ionic strength, m, = 0.001; gH+ = 0.967
– pH = -log(.001*.967) = 3.01
• pH of a 0.100 M HCl solution
– Ionic strength, m, = 0.1; gH+ = 0.83
– pH = -log(0.1*0.83) = 1.08
• pH of a 0.00100 M HCl/0.100 M NaCl solution
– Ionic strength, m, = 0.1; gH+ = 0.83
– pH = -log(0.001*0.83) = 3.08
What is the concentration H+ an
OH- in a 0.100 M NaCl solution?
Kw = [OH-]gOH-[H+]gH+ = 1.01E-14
At m = 0.100, gOH- = 0.76 and gH+ = 0.83
H2O is the only source of H+ and OH- so let
x = [H+] = [OH-]
Kw = (0.76)(0.83) x2
x = 1.27E-7 M
EDTA Prob 12-31
• A 50 mL sample containing Ni2+ is treated
with 25.00 mL of 0.0500 M EDTA to
complex all of the Ni2+. The excess EDTA
is back-titrated, requiring 5.00 mL of
0.0500 M Zn2+. What is the concentration
of Ni2+ IN THE ORIGINAL SAMPLE?