Introduction:
The compound to be prepared is K3[Fe(C2O4)3].3H2O.
Its correct name is potassium tris(oxalato)ferrate(III) trihydrate. The
components of this compound are K+, Fe3+, C2O42-,
and H2O. The ion C2O42- is called
the oxalate ion, and since there are three C2O42-
in the formula above, and iron is in the 3+ state (indicated by the
Roman numeral III), the complex ion in the square brackets has a 3-
charge, and thus three K+ are required to balance the charge.
Preparation of K3[Fe(C2O4)3]·
3H2O:
During the first laboratory period, K3[Fe(C2O4)3]·3H2O
will be prepared. Ferrous ammonium sulfate, Fe(NH4)2(SO4)2·6H2O,
is dissolved in a slightly acid solution, excess oxalic acid, H2C2O4,
is added, and the following reaction takes place:
|
Fe(NH4)2(SO4)2·
6H2O + H2C2O4 = FeC2O4(s)
+ H2SO4 + (NH4)2SO4
+ 6H2O(l)
|
(Eq 1)
|
FeC2O4 is a finely divided
precipitate and tends to be colloidal. However, heating the solution causes
it to coagulate and facilitates separating the precipitate from the solution.
Potassium oxalate is added to the FeC2O4
precipitate which produces a slightly basic solution for the oxidation
of the ferrous ion to the ferric ion by hydrogen peroxide, H2O2.
The following reaction takes place:
|
2Fe 2+
|
=
|
2Fe 3+ + 2e- |
(2e- "lost") |
|
|
H2O + HO2- + 2e-
|
=
|
3OH- |
(2e- "gained") |
|
|
H2O + HO2- + 2Fe 2+
|
=
|
2Fe3+ + 3OH- |
(net reaction) |
(Eq 2)
|
Note that the FeC2O4 is the source of the Fe 2+
in Equation 2.
The OH- ion concentration of the solution
is high enough so that some of the Fe3+ reacts with OH-
to form ferric hydroxide (brown precipitate) as follows:
|
Fe3+ + 3OH- = Fe(OH)3(s)
|
(Eq 3)
|
With the addition of more H2C2O4,
the Fe(OH)3 dissolves and the soluble complex K3[Fe(C2O4)3]·
3H2O is formed according to:
|
3K2C2O4 + 2Fe(OH)3(s)
+3H2C2O4 = 2K3[Fe(C2O4)3]·
3H2O + 3H2O
|
(Eq 4)
|
Alcohol is added to the solution to cause the complex
iron salt to precipitate since it is less soluble in alcohol than in water.
The complexity of the series of reactions described
in equations 1 - 4 may be greatly simplified by following the Fe2+/Fe3+
ion throughout. One discovers that for every mole of Fe(NH4)2(SO4)2].
6H2O used as starting material, one mole of K3[Fe(C2O4)3].3H2O
will be obtained as the final product.
| Fe(NH4)2(SO4)4.
6H2O + H2C2O4 |
=
|
FeC2O4
+ ......etc. |
|
(see Eq. 1) |
(Fe2+
= Fe2+) |
| FeC2O4
+
K2C2O4 + H2O2 |
=
|
Fe(OH)3(s) + .....etc |
|
(see Eq. 2 & 3) |
(Fe2+ = Fe3+) |
| Fe(OH)3 + H2O2
+ K2C2O4 |
=
|
K3[Fe(C2O4)3]·3H2O + ..etc |
|
(see Eq. 4) |
(Fe3+ = Fe3+) |
Experimental Procedure
Severe bumping can occur during the preparation of this complex, particularly
in step 2.
Preparation of K3[Fe(C2O4)3]·3H2O:
The following is a generic procedure to synthesize
the iron coordination complex. In your notebook you should
write-up a procedure using a similar format that records the exact steps
that you did in making your complex. You should record the quantities
and volumes of the reagents that you used and not those given in the generic
procedure below.
-
Using a top-loading balance, weigh about 5g of Fe(NH4)2(SO4)2·6H2O
in a 125-mL Erlenmeyer flask and record this weight on your data sheet.
Dissolve this compound in 15mL distilled water and add 4-5 drops of 3M
H2SO4.
-
Add 50mL ~0.5M H2C2O4 to this solution
and heat to boiling while stirring constantly to prevent bumping.
-
Remove the Erlenmeyer flask from the heat and allow the yellow precipitate
of FeC2O4 to settle. Decant the supernatant liquid
(pour the liquid away from the solid) and wash the precipitate using 20mL
of hot distilled water. Swirl the mixture and allow the precipitate to
settle; decant and repeat the washing once more.
-
Add 20mL of ~1M K2C2O4 to the flask containing
the precipitate, stir and heat to 40°C. While the temperature is at
40°C, immediately add 10mL of 6% H2O2
dropwise and stir continuously. Periodically check the temperature of the
solution and make sure that it is at least 40°C (but not >50°C)
during the addition of H2O2. (Some brown Fe(OH)3
may precipitate at this time.)
-
Heat the resulting solution to boiling. Obtain 20mL of ~0.5M H2C2O4
and add ~5mL of it all at once to the boiling solution. Stir continuously
and add the last few mL dropwise while maintaining the temperature near
boiling. The solution should turn clear green. If some brown residue remains,
add an additional 1mL of H2C2O4 dropwise,
again while the solution is boiling, until the solution is clear green.
However, if the residue that remains is yellow, it is probably unreacted
FeC2O4, and more H2O2 should
be added carefully. Consult your TA.
-
If the solution is cloudy, gravity filter it into a clean 125-mL Erlenmeyer.
If it is clear, no filtration is necessary. Then, while swirling constantly,
slowly add 15mL ethanol to the solution. Allow to cool, while an ice bath
is prepared in a 400 mL beaker. Immerse the bottom portion of the flask
in the ice bath and stir slowly until crystals begin to form. Stop stirring
and allow the solution to stand in the ice bath for 20 minutes. A good crop
of crystals should have formed before the solution is filtered. (Consult
your TA if you are doubtful about quantity.)
-
Prepare a vacuum filtration apparatus.
-
Decant the supernatant liquid away from the green crystals which have formed
in the flask and temporarily save the liquid in a beaker if the yield seems
low. Consult with your TA before discarding this solution. With the aid
of a glass rod, transfer the crystals to the Buchner funnel and apply suction
for about 2 minutes.
-
Stop the suction and add 10mL of a 1:1 ethanol/water solution. Wait about
30 seconds and then apply suction for 2 minutes. Repeat this washing process.
After this final wash, allow the suction to continue for a further two
minutes.
-
Transfer the crystals to a preweighed, labeled dry sample vial.
Discard the wash solutions.
Useful molar masses:
|
K2C2O4
|
|
166.2 g/mol |
|
Fe(NH4)2(SO4)2·6H2O
|
|
392.1 g/mol |
|
K3[Fe(C2O4)3]·3H2O
|
|
491.1 g/mol |
|
