THE PREPARATION , ANALYSIS AND REACTION OF AN ETHANEDIOATE (OXALATE) COMPLEX OF IRON

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AUGUSTINE ADU _ KNUST
KWAME NKRUMAH UNIVERSITY O F SCIENCE AND
TECHNOLOGY


DEPARTMENT OF CHEMISTRY
YEAR TWO (CHEM 269)
PRACTICAL CHEMISTRY III
I.2.1.1 & I.2.1.2 REPORT
TITLE: THE PREPARATION, ANALYSIS AND REACTIONS OF AN
ETHANEDIOTE (OXALATE) COMPLEX IRON


NAME: AUGUSTINE ADU




TITLE: THE PREPARATION, ANALYSIS AND REACTIONS OF
AN ETHANEDIOTE (OXALATE) COMPLEX IRON

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AIMS AND OBJECTIVES
1. To prepare oxalate complex of iron from the reaction with ammonium iron (II)
sulphate
2. To deduce the oxidation states of iron in the complex
3. To determine the amount of iron and oxalate in the oxalate complex of iron by
titration.
4. To investigate the molecular formula of the complex formed by preparation, analysis
and the reaction of an ethanedioate (oxalate) complex of iron
5. To deduce the chemical properties of the complex of iron from the reactions with
dilute sodium hydroxide, ammonium thiocyanate solution, and ammonium
thiocyanate solution in the presence of dilute sulphuric acid respectively, compared to
the reaction of iron(III) chloride and the reagents above.

INTRODUCTION
Iron is nearly always determined by reduction to the dipositive state followed by titration
with manganate(VII) or dichromate(VI). However oxalate would interfere and must be
determined first by titration with permanganate. After titration, any iron present will be
Fe(III) then reduced by tin(II) chloride and hydrochloric acid, and the Fe(II) determined with
dichromate.
THE STRUCTURE OF IRON OXALATE

X-ray crystallography of simple oxalate salts show that the oxalate anion may adopt either a
planar conformation with D2h molecular symmetry, or a conformation where the O-C-C-O
dihedrals approach 90° with approximate D2d symmetry. Specifically, the oxalate moiety
adopts the planar, D2h conformation in the solid-state structures of M2C2O4 (M = Li, Na, K).

However, in structure of Cs2C2O4 the O-C-C-O dihedral angle is 81(1)°.

Therefore, Cs2C2O4
is more closely approximated by a D2d symmetry structure because the two CO2 planes are

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staggered. Interestingly, two forms of Rb2C2O4 have been structurally characterized by
single-crystal, X-ray diffraction: one contains a planar and the other a staggered oxalate.
As the preceding examples indicate that the conformation adopted by the oxalate dianion is
dependent upon the size of the alkali metal to which it is bound, some have explored the
barrier to rotation about the central C−C bond. It was determined computationally that barrier
to rotation about this bond is roughly 2–6 kcal/mole for the free dianion, C2O4
2−
. Such results
are consistent with the interpretation that the central carbon-carbon bond is best regarded as a
single bond with only minimal pi interactions between the two CO2 units.

This barrier to
rotation about the C−C bond (which formally corresponds to the difference in energy
between the planar and staggered forms) may be attributed to electrostatic interactions as
unfavorable O−O repulsion is maximized in the planar form.
It is important to note that oxalate is often encountered as a bidentate, chelating ligand, such
as in Potassium ferrioxalate. When the oxalate chelates to a single metal center, it always
adopts the planar conformation.
Oxalate occurs in many plants, where it is synthesized via the incomplete oxidation of
carbohydrates.
Oxalate-rich plants include fat hen ("lamb's quarters"), sorrel, and several Oxalis species. The
root and/or leaves of rhubarb and buckwheat are high in oxalic acid. Other edible plants that
contain significant concentrations of oxalate include—in decreasing order—star fruit
(carambola), black pepper, parsley, poppy seed, amaranth, spinach, chard, beets, cocoa,
chocolate, most nuts, most berries, fishtail palms, New Zealand spinach (Tetragonia
tetragonioides) and beans. Leaves of the tea plant (Camellia sinensis) contain among the
greatest measured concentrations of oxalic acid relative to other plants. However the
beverage derived by infusion in hot water typically contains only low to moderate amounts of
oxalic acid per serving due to the small mass of leaves used for brewing.
Transition metal ions react with charged or neutral ligands, L, (e.g. Cl
–
or H2O) to form
complex ions. Iron in the +3 oxidation state can form octahedral complexes with up to 6
unidentate ligands surrounding a central metal ion (Figure 1). The ligands act as Lewis
bases, donating at least one pair of electrons to the Fe
3+
ion. Oxalate ion, C2O4
2–
, acts as a
chelating bidentate ligand, donating 2 electron pairs from 2 oxygen atoms to the transition

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metal center, Fe
3+
. During the first week of this experiment a coordination compound with
the formula KxFey(C2O4)z·nH2O will be synthesized. A coordination compound typically
contains a complex ion (with ligands bound to a central metal cation), counter ions, and,
sometimes, waters of hydration. During the second week the empirical formula of the
coordination compound will be determined (i.e., the values of x, y, z, and n) by redox titration
and gravimetric analysis.
The general equation of the reaction is;
(NH4)2[Fe(H2O)2(SO4)2]*4H2O + H2C2O4*2H2O FeC2O4 + H2SO4 + (NH4)2SO4
+8H2O
H2C2O4*H2O + 2FeC2O4 + 3K2C2O4*H2O + H2O2 2K3[Fe (C2O4)3*3H2O + H2O

The first week synthesis of the iron complex begins with Mohr's salt: Fe(NH4)2(SO4)2
.
6H2O
The salt is dissolved in water and the solution is kept at a low pH by addition of sulfuric acid
to prevent the formation of rust coloured iron oxides and hydroxides. Oxalate ions are added
in the form of oxalic acid and potassium oxalate. The oxalate will replace some or all of the
water and sulfate ligands coordinated to the iron (II) ion and a yellow solid forms. The bright
yellow precipitate is filtered from solution, washed to remove impurities, and treated with 3%
hydrogen peroxide to oxidize the iron to the +3 state. Although the solution is heated slightly
to increase the rate of oxidation, the addition of peroxide is done slowly to prevent the heat
sensitive peroxide from decomposing before reacting with all of the iron (II) in solution. All
the Fe
2+
must be oxidized to Fe
3+
. Complex ions that form with the Fe
3+
have a different
number of oxalate groups than those that form with Fe
2+
. Empirical formula determination is
difficult with a mixture of the two complex ions.
At this point, the Fe
3+
complex ion combines with a potassium counter ion leading to the
formation of the coordination compound: KxFey(C2O4)z
.
nH2O. Since this salt is less soluble
in alcohol than in water, 95% ethanol is added to the solution and a green crystalline solid
begins to precipitate from solution within 2-3 days. The solution must be stored in the dark
during crystallization because visible light will reduce Fe
3+
to Fe
2+
.
During the second week, the crystallized salt will be analyzed to determine the mass percent
of oxalate ion. Additional data will be provided to calculate the mass percent of iron, water
and potassium so the empirical formula can be determined. The mass percent oxalate ion in
the salt will be determined by titration with a standardized KMnO4 solution according to the

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unbalanced reaction below:

(1) MnO4
–
(aq) + C2O4
2–
(aq)  Mn
2+
(aq) + CO2(g)

Since aqueous solutions of permanganate ion are not stable over a long period of time, the
exact concentration of KMnO4 must be determined by titration with a known amount of a
primary standard salt such as sodium oxalate, Na2C2O4. After KMnO4 has been
standardized, the complex iron salt can be titrated to determine its oxalate content. The
solutions containing C2O4
2–
and Mn
2+
ion are colourless; the MnO4
–
solution is a deep purple
colour. Therefore, the titrated solution will remain colourless until all the oxalate salt is
consumed in the reaction. The endpoint corresponds to the appearance of the first permanent
pink colour due to the presence of excess unreacted permanganate ion. The rate of the
reaction is very slow at room temperature so the solution must be heated to 80°C to observe
the colour change in "real time". Often, at the beginning of the titration, the purple colour of
the KMnO4 does not disappear for 30-60 seconds because the reaction has an intermediate
that must form before the reaction goes to completion.
In redox titrations, solvent impurities act as reducing or oxidizing agents requiring the
addition of more titrant. To correct for this a blank containing only the solvent must be
titrated. The "corrected volume" is equal to the volume of KMnO4
–
required to titrate oxalate
ion in solvent minus the volume required to titrate the solvent alone. (Important note: The
amount needed to titrate the blank is often only one or two drops of KMnO4
–
.
The ferric ion, Fe
3+
, is released into solution when permanganate ion reacts with oxalate ion
and destroys the complex ion. The liberated Fe
3+
ion is reddish coloured and can interfere
with observation of the faint pink titration endpoint. To eliminate the colour interference, a
small amount of concentrated phosphoric acid is added to the solution. The phosphate ion
reacts with Fe
3+
to yield a colourless complex ion, Fe(PO4)2
3–
, eliminating the reddish-brown
colour of Fe
3+
from solution.
To determine the mass percent of iron, Fe(III) must first be reduced to Fe(II) by exposure to
sunlight or by reaction with Al metal. The resulting Fe
2+
ion is then titrated with a
standardized KMnO4 solution according to the unbalanced equation below:
(2) Fe
2+
(aq) + MnO4
–
(aq)  Fe
3+
(aq) + Mn
2+
(aq)

The mass percent of water is determined by gravimetric analysis. A known mass of the
complex salt containing water is weighed, heated and reweighed. The weight of water is the

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mass difference between the hydrous and anhydrous forms of the salt.
The mass percent of potassium in the salt is determined by difference using the
experimentally determined masses of iron, oxalate and water and the mass of the complex
iron salt.

CHEMICALS AND EQUIPMENT
1. Di Ammonium iron (II) sulphate
2. 2M H2SO4 solution
3. Potassium hydrogen oxalate
4. 6% hydrogen peroxide solution
5. Distilled water
6. 0.02M permanganate solution
7. Oxalic acid “ AnalaR”
8. Mixture of ice in water
9. Electronic balance
10. Mercury thermometer
11. 2 conical flaks
12. 20ml measuring cylinder
13. burette
14. 50ml volumetric flask
15. Source of heat
16. Stirrer

PROCEDURE
STEP OBSERVATION INFERENCE
10.0g of Di Ammonium iron (II)
sulphate was weighed into a 400ml
beaker. A few drops of 2M H2SO4 and
30ml of water were added.
A light green solution was formed. The Di-ammonium iron (II)
sulphate was slightly oxidized
by H2SO4.
5.0g of oxalic acid AnalaR solution in
30ml of water was added to the green
solution.
The solution changes to yellow
solution.
The oxalic acid is a reducing
agent which reduces the iron (II)
formed back to the iron (II).
The solution was then heated
cautiously with continuous stirring to
the boiling point.
A two-layer solution was formed, a
colourless top solution and a
bottom yellow solution containing a
yellow precipitate.
The reaction that took place was
Fe(NH4)2(SO4)2· 6H2O +
H2C2O4 = FeC2O4(s) +
H2SO4 + (NH4)2SO4 + 6H2O(l)
The supernatant liquid was poured off
and the precipitate washed with hot
water.
The precipitate formed was
FeC2O4.2H2O.
7.5g of potassium hydrogen oxalate in
20ml of water was heated.
The potassium hydrogen oxalate
was insoluble in water.

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The solution was cooled to 40
o
C and
then 20ml of 6% hydrogen peroxide
was added in drops, the solution being
stirred continuously.
The yellow precipitate of iron (II)
oxalate was insoluble. The solution
however changed to green
containing green precipitate upon
the addition of H2O2 with
effervescence.
H2O2 reacted with the iron (III)
formed according to the reaction
Fe
3+
+ 3OH
-
= Fe(OH)3(s)
The potassium hydrogen oxalate
reacts with the Fe(OH)3(s)
formed to form the complex
trihydrate according to the
equation
3K2C2O4 + 2Fe(OH)3(s)
+3H2C2O4 = 2K3[Fe(C2O4)3]·
3H2O + 3H2O

The mixture now containing iron (III)
hydroxide was heated to boil and then
2.5g of oxalic acid AnalaR was added
and the mixture stirred continuously.
The green precipitate dissolved
upon heating. The addition of the
oxalic acid AnalaR also produced
some effervescence.
Both H2O2 and Fe(OH)3(s) are
unstable to heat, thus the heating
gets rid of any excess hydrogen
peroxide that may be present and
also hastens the reaction with the
potassium hydrogen oxalate by
decomposing Fe(OH)3(s).
To the clear solution, 25ml of 95%
ethanol was added and the solution
kept in the dark for a week.
There was no colour change upon
the addition of the ethanol.
Ethanol causes the complex iron
salt K3[Fe(C2O4)3]. 3H2O to
precipitate since it is less soluble
in alcohol than in water.
0.7g of the precipitate formed was
weighed and equally divided into two
separate conical flasks. They were
then dissolved in 10ml of water and
15ml of dilute H2SO4 and heated
to70
o
C and then titrated against 0.02M
permanganate.
There was a colour change from
yellow to pink.
Hot H2SO4 reduces the iron (III)
back to iron (II) which is yellow
in colour. The reaction that
occurred follows the equation
[Fe(C2O4)3]
-3
+ 6H
+

Fe
2+
+3H2C2O4
MnO
-
4+5H2C2O4+6H
+

2Mn
2+
+10CO2+8H2O

The heated mixture is titrated against KMnO4

TABLE OF RESULTS
Burette Readings/Ml A B
Final Reading (mL) 36.25 30.60
Initial Reading (mL) 0.00 0.00
Titre value (mL) 36.25 30.60

CALCULATION AND EVALUATION OF DATA
Average titre = (36.25+30.60)/2= 33.43ml
n[KMnO4] = (33.43×0.02)/1000=7.170.02)/1000=6.686×10
-4
mol

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from mole ratio 1mol of KMnO4: 5mol of H2C2O4
n[H2C2O4]= 5×6.686×10
-4
= 3.343×10
-3
mol
mole ratio of the [H2C2O4] to [K3[Fe (C2O4)3*3H2O]= 3:1
n[K3[Fe (C2O4)3*3H2O]= 3.343×10
-3
/3=1.114×10
-3
mol
M [K3[Fe (C2O4)3*3H2O]= 3(39)+56+6(12)+12(16)+3(18)=491g/mol
mass of [K3[Fe (C2O4)3*3H2O] that was produced= 491×1.114×10
-3
= 0.547138g=0.55g
this implies that the complex occupied (0.55×100)/0.7= 78.57% of the measured mass.

DISCUSSION
The H2SO4 was added to provide a slightly acidic solution in which to dissolve the Di-
ammonium iron (II) sulphate. However it slightly oxidizes the iron (II) content of the
complex. The addition of oxalic acid dihydrate to the Fe(NH4)2(SO4)2· 6H2O/H2SO4 solution
yielded a yellow solution due to the formation of iron (II) oxalate. Further addition of H2O2
aided the formation of the complex. The addition of ethanol was to precipitate any of the
formed complex present in the green solution because the complex is less soluble in alcohol
than in water. The permanganate titration oxidized the iron (II) to iron (III) this caused the
change in colour of the solution.
PRECAUTION S
1. Safety goggles, aprons, and gloves were worn in the laboratory throughout the
experiment.
2. Oxalate is very toxic via oral and inhalation routes and severe kidney damage is possible
if oxalate salts are taken internally. Oxalate compounds can be absorbed through the
skin; gloves were therefore worn and affected areas were washed with cold water.
3. 95% ethyl alcohol (ethanol) and acetone are flammable; all open flames were
extinguished in lab.
4. H2SO4 and concentrated H3PO4 are corrosive acids; all affected areas were washed
thoroughly with cold water.
5. KMnO4 is a very strong oxidizing agent; do NOT pour any permanganate solutions into
the ORGANIC collection bottles. Permanganate solutions can stain skin and clothing.
6. The solution was kept in the dark because in solution the ferrioxalate complex is
decomposed by light.
7. Severe bumping can occur especially during the oxidation reaction hence the need for

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the continuous stirring.

CONCLUSION
The oxalate complex of iron was successfully prepared from the reaction of the iron with
ammonium iron (II) sulphate. The mass of K3[Fe (C2O4)3*3H2O was found to be 0.55g out of
the total 0.7g of the sample measured. This was a percentage of 78.57% of the sample.

REFERENCES
1. Duncan, J. (2010). Experiment 1: synthesis and analysis of an inorganic compound.
Department of Chemistry, Plymouth State University, New Hampshire, US, United
States.
2. http://oz.plymouth.edu/~jsduncan/courses/2010_Fall/InorganicChemistry/Labs/1-
InorganicCmpd_SynthAnalysis.pdf on 27.03.11
3. Coordination complex. (n.d.). Retrieved from
http://en.wikipedia.org/wiki/Coordination_complex on 29.03.11
4. http://pubs.acs.org/doi/abs/10.1021/ed081p1193 doi:10.1021/ed081p1193 on 29.03.11
5. Journal of Solid State Chemistry, 12(1-2), Retrieved from
http://www.sciencedirect.com/science
6. Modern Inorganic Chemistry, Second Edition by William L. Jolly, pages 357 and 468.

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