Transcript Slide 1

Laboratory Experiments 3 and 4
Based on solution chemistry.
Experiment 3:
Aqueous acid – base chemistry
Experiment 4:
Aqueous complex formation and solubility equilibria
Titration Curves
strong base OH- (burette) vs. weak acid HAc (aliquot)
Ka = [H+][Ac-]/[HAc] and
[H+] = Ka[HAc]/[Ac-]
after each addition of OH- the equilibrium readjusts
as [HAc] approaches 0
[H+] also approaches 0
and pH changes rapidly
pH titration NaOH vs. HAc
end point
pH
7
½ way point
5
ΔpH/ΔV
4
V titre (mL)
The ½ way point
At the ½ way point of the titration
[HAc] = [Ac-] and [HAc]/[Ac-] = 1
since log (1) = 0
pH = pKa - log{[HAc]/[Ac-]} the H-H equation
pH = pKa at the ½ way point
The Ka values for weak acids and the Kb for their
conjugates can be determined by a ‘simple’ titration.
Coloured End Point Indicators
Some weak acids have different colours for their two
forms
HIn
H+
+ In-
In-
HIn
pH = pKa + log {[In-]/[HIn]}
The solution colour depends on the pH.
As the pH changes rapidly at the end point the solution
colour changes rapidly.
The Indicator is chosen so that it’s pKa value is close to
the pH at the end point.
Some indicator pKa values
Indicator
Use
thymol blue
0.1 % in water
Colour change
range
pKa
red to yellow
1.2 - 2.8
1.7
methyl orange 0.1 % in water red to yellow
3.1 - 4.4
3.7
phenol red
6.8 - 8.4
7.9
9.3 -10.0
9.6
0.1 % in water yellow to red
phenolphthalein 0.1 % in alc. clear to red
Effect of Ions in Water
The pH value of a solution is set by the position of the
equilibrium.
Kw = [H+][OH- ] = K[H2O] = 1 x 10-14
Ions that remove H+ or OH- will lower or raise the
solution pH.
i.e. Adding NaAc (sodium acetate) consumes H+ by
forming HAc. The Na+ ion does not consume OH- by
forming a complex. The pH rises ( > 7).
Predicting the pH change upon the addition of MX.
H+
OH-
M+
X-
H+
OH-
M+
X-
Major interaction
Minor interaction
Net consumption/removal of H+
pH rises
Net consumption/removal of OH-
pH falls
Solution Equilibria
Reactions controlled by equilibrium occur in solution.
i.e. the precipitation of salts
AgCl
Ag+
+ Cl-
Other major reactions are called
COMPLEX FORMATION
These are homogeneous reactions (all in one phase).
Consider the case where ammonia (NH3) is added to Ag+
ions in solution.
Ag(NH3)2+
Ag+
+ 2NH3
Complex equilibrium constants are constructed in the
same way as other equilibria.
Ag(NH3)2+
Kinstab
Kstab
Ag+
=
=
+ 2NH3
[Ag+][NH3]2/[Ag(NH3)2+]
1/Kinstab
Most metal form complexes with negative ions in solution.
These may result in a soluble or insoluble product.
Ag(S2O3)2-3
Ag(CN)2-
Ag+ + 2S2O32- (thiosulphate)
Ag+ + 2CN- (cyanide)
The absolute values of Ksp and Kinstab are difficult to
determine the relative values are not.
Consider the salts
Ag NO2, AgF, AgCl, AgBr, AgI,
Ag+
AgX (s)
+ X- Q = [Ag+][X-]
A solution of AgNO3 (soluble) Ag+ (0.1 M) mixed
with a small volume of 1 M NaX, if:
Q > KspAgX
ppt
Q < KspAgX
no ppt
1 mL of 0.1
M AgNO3
1 drop of
1 M NaX
NO2-
F-
Cl-
Br-
I-
Since Q = [Ag+][X-] an observation of:
No
ppt
ppt
ppt
Means the inference KspAgNO2 > KspAgX
ppt
ppt
1 mL of 0.01 M
AgNO3
1 drop of
1 M NaX
NO2-
F-
Cl-
No
ppt
ppt
Br-
I-
An observation of:
No
ppt
ppt
Means the inference KspAgNO2 > KspAgF > KspAgX
ppt
Lowering the [Ag+] further can be achieved using complex
formation.
Ag(NH3)2+
Ag+ + 2NH3
4.0 M ammonia leaves low concentration of ‘free’ [Ag+]
16 M leaves even less.
This logic can be used to test the relative strengths of
complexes.