1. Laboratory Experiments 3 and 4
Based on solution chemistry.
Experiment 3:
Aqueous acid – base chemistry
Experiment 4:
Aqueous complex formation and solubility equilibria

2. 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

3. pH titration NaOH vs. HAc
end point pH 7
½ way point 5
ΔpH/ΔV 4
V titre (mL)

4. 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.

5. Coloured End Point Indicators
Some weak acids have different colours for their two forms
HIn H In-
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.
In HIn

6. Some indicator pKa values
Indicator Use Colour change range pKa
thymol blue % in water red to yellow
methyl orange 0.1 % in water red to yellow
phenol red 0.1 % in water yellow to red
phenolphthalein 0.1 % in alc. clear to red

7. 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).

8. Predicting the pH change upon the addition of MX.
Major interaction
Minor interaction
H+ OH-
M X-
Net consumption/removal of H pH rises
H+ OH-
M X-
Net consumption/removal of OH pH falls

9. 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) Ag NH3

10. Complex equilibrium constants are constructed in the
same way as other equilibria.
Ag(NH3) Ag NH3
Kinstab = [Ag+][NH3]2/[Ag(NH3)2+]
Kstab = 1/Kinstab
Most metal form complexes with negative ions in solution. These may result in a soluble or insoluble product.
Ag(S2O3) Ag S2O32- (thiosulphate)
Ag(CN) Ag CN- (cyanide)

11. 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,
AgX (s) Ag 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

12. Since Q = [Ag+][X-] an observation of:
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
ppt
ppt
Means the inference KspAgNO2 > KspAgX

13. Means the inference KspAgNO2 > KspAgF > KspAgX
1 mL of 0.01 M AgNO3
1 drop of
1 M NaX
NO2- F-
Cl-
Br- I-
An observation of:
No ppt No
ppt
ppt
ppt
ppt
Means the inference KspAgNO2 > KspAgF > KspAgX

14. Lowering the [Ag+] further can be achieved using complex formation.
Ag(NH3) Ag NH3
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.