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

V titre (mL)

pH

pH titration NaOH vs. HAc

½ way point

7

4

5

end point

ΔpH/ΔV

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.

The ½ way point

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

Some indicator pKa values

Indicator Use Colour change range pKa

thymol blue 0.1 % in water 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 0.1 % in water yellow to red 6.8 - 8.4 7.9

phenolphthalein 0.1 % in alc. clear to red 9.3 -10.0 9.6

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).

H+ OH- M+ X-

Predicting the pH change upon the addition of MX.

Major interactionMinor interaction

Net consumption/removal of H+ pH rises

H+ OH- M+ X-

Net consumption/removal of OH- pH falls

Solution EquilibriaReactions 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+ Ag+ + 2NH3

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)2-3 Ag+ + 2S2O3

2- (thiosulphate)Ag(CN)2

- 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,

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

1 mL of 0.1 M AgNO3

1 drop of

1 M NaXNO2

- F- Cl- Br- I-

Since Q = [Ag+][X-] an observation of: No ppt

ppt ppt ppt ppt

Means the inference KspAgNO2 > KspAgX

1 mL of 0.01 M AgNO3

1 drop of

1 M NaXNO2

- F- Cl- Br- I-

An observation of: No ppt

Noppt

ppt ppt ppt

Means the inference KspAgNO2 > KspAgF > KspAgX

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.