Experiment on the preparation of potassium tris (oxalate) ferrate (ii) trihydrate
KWAME NKRUMAH UNIVERSITY OF SCIENCE AND
DEPARTMENT OF CHEMISTRY
YEAR TWO (CHEM 269)
PRACTICAL CHEMISTRY III
TITLE: PREPARATION OF POTASSIUM TRIS (OXALATE)
FERRATE (II) TRIHYDRATE
NAME: OPOKU ERNEST
DATE: 4TH NOVEMBER, 2013
TITLE: PREPARATION OF POTASSIUM TRIS(OXALATE)
AIMS AND OBJECTIVES
The aim of this experiment is to synthesize potassium tris(oxalato)ferrate (III) via the addition
of oxalic acid and potassium hydroxide to iron (III) chloride hexahydrate under gentle heating.
Various reactions will then be carried out on the product in an attempt to further understand the
characteristics of this metal complex. This experiment will help to:
1. To expose students to the use of the electronic balance.
2. To introduce titration as a useful technique in chemical analysis.
3. To prepare a pure sample of the complex K3[(Fe(C2O4)3.3H2O in good yield.
4. To use titration to determine the amount of oxalate in the product.
5. To determine percentage by mass of iron (III) produced from K3[(Fe(C2O4)3.3H2O complex.
3-dimensional structure of potassium tris(oxalato)ferrate (III)
Potassium tris(oxalato)ferrate (III) is a metal complex of iron with three oxalate ligands (C2O4
bonded to every central metal atom. These ligands are bidentate, meaning that each of them
binds to the metal atom at 2 different places. It has the chemical formula K3[(Fe(C2O4)3.3H2O,
and the three-dimensional structure proposed in Figure 1. Such complexes are often utilized in
schools and universities to introduce various concepts such as ligand strength, metal complexes,
and ligand replacement. Potassiumtris(oxalato)ferrate (III) is hygroscopic and light sensitive in
nature. In this experiment, we synthesized this fascinating compound via the addition of oxalic
acid to potassium hydroxide, forming potassium oxalate, the intermediate for this reaction
mechanism. The chemical reaction is as follows: H2C2O4(aq) + 2KOH (aq) K2C2O4·H2O
(aq) + H2O(l). Iron (III) chloride hexahydrate was then added to the reaction mixture, forming
our desired product in the following chemical reaction: K2C2O4·H2O(aq)+
FeCl3·6H2O(aq) K3[(Fe(C2O4)3.3H2O (s)+3KCl(aq)+6H2O(l).Our desired product was
produced in the form of green crystals.
The oxalic acid utilized in the first step of this reaction scheme can be synthesized by
or by oxidizing sucrose or glucose with nitric acid in the presence of a
small amount of vanadium pentoxide2
Another method of forming oxalic acid involves the oxidative carbonylation of alcohols followed
The iron (III) chloride hexahydrate used in the second step is toxic, highly corrosive and acidic.
It is usually produced by dissolving iron ore in hydrochloric acid. The representative chemical
equation is as follows:Fe3O4(s) + 8HCl (aq) FeCl2(aq) + 2 FeCl3(aq) + 4H2O.
There is another method in scientific literature that is commonly utilized to form potassium
tris(oxalato)ferrate (III). In this method, oxalic acid is added to ferrous ammonium sulphate
hexahydrate Fe(NH4)2(SO4)2 • 6H2O under acidic conditions. This forms iron (II) oxalate
(FeC2O4), a yellow precipitate. This is then added to potassium oxalate (H2C2O4) and hydrogen
peroxide, finally synthesizing our desired product. As observed, this alternative method is longer
than the method we utilized in this experiment. There are more steps, more intermediates and
more reactants required. Some of the reactants used are also rather dangerous and harmful, such
as hydrogen peroxide. It is therefore the less favoured method out of the two. After successfully
synthesizing our product, it was utilized in a variety of reactions to further understand the
chemical properties of such a metal complex.
Also, to prepare potassium tris(oxalato) iron(III) trihydrate, K3[Fe(C2O4)3] • 3H2O, ferrous
ammonium sulfate hexahydrate, Fe(NH4)2(SO4)2 • 6H2O), is dissolved in a slightly acidic
solution, excess oxalic acid is added, and the following reaction takes place:
1: Fe(NH4)2(SO4)2 • 6H2O + H2C2O4 FeC2O4(s) + H2SO4 + (NH4)2SO4 + 6H2O(l) Ferrous
oxalate, 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 ferrous oxalate precipitate that
produces a slightly basic solution for the oxidation of the ferrous ion to the ferric ion by
hydrogen peroxide. The following reaction takes place:
H2O + HO2
+ 2e- 3 OH(–)
2: H2O + HO2
+ 2 Fe(2+)
+ 3 OH(–)
Note that the ferrous oxalate is the source of the Fe(2+)
in Equation 2. The hydroxide ion
concentration of the solution is high enough so that some of the Fe(3+)
reacts with hydr- oxide
to form ferric hydroxide (brown precipitate) as follows:
With the addition of more oxalic acid, the ferric hydroxide dissolves and the soluble complex
4: 3 K2C2O4 + 2 Fe(OH)3 (s) +3 H2C2O4 2 K3[Fe(C2O4)3] • 3H2O + 3 H2O
Alcohol is added to the solution to cause the complex iron salt to precipitate since it is less
soluble in alcohol than in water.
CHEMICALS AND EQUIPMENT
1. 6% hydrogen peroxide solution
2. Oxalic acid “ AnalaR”
3. 10% oxalic acid
4. 2M H2SO4 solution
5. Di Ammonium iron (II) sulphate
6. Potassium hydrogen oxalate
7. Distilled water
8. Source of heat
10. Water bath containing ice and water
12. Electronic balance
13. 50ml measuring cylinder
14. 5ml measuring cylinder
15. Two beakers
16. Mercury thermometer
STEP OBSERVATION INFERENCE
1ml of 2M H2SO4 was weighed into a beaker
container 50ml of water and the solution
A colourless solution was
The acid dissociated totally in the aqueous
solution according to the equation.
H2SO4 + 2H2O 2H3O+
(aq) + SO2-
15g of Di-ammonium iron (II) sulphate was
weighed and added to the warm solution
The solution changed to light
Iron (II) is oxidized slightly by the acid to
75ml of a 10% w/v oxalic acid solution was
added to the green solution formed. The
solution was heated and continuously stirred.
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 oxalic acid being a reducing agent,
reduces the iron (III) formed back to iron (II).
The supernatant solution was decanted and
then 40ml of hot water was added. The
solution was filtered and the residue dried in
A dry yellow precipitate was
The precipitate formed is FeC2O4.2H2O, iron
(II) oxalate (ferrous oxalate).
5g of potassium hydrogen oxalate was added
to 15ml of water and the solution heated.
The potassium hydrogen oxalate
was insoluble in water.
3g of ferrous oxalate (the dried yellow
precipitate) was added to the warm solution of
potassium hydrogen oxalate.
A yellow solution was formed. Potassium hydrogen oxalate was added to
produce a slightly basic solution for the
oxidation of the ferrous ion to the ferric ion
13ml of 6% w/v hydrogen peroxide was added
and the solution stirred continuously.
The solution changed colour
from yellow to green with
The peroxide oxidizes the ion from Fe2+
H2O2 + 2Fe2+
25ml of 95% ethanol was added to the
solution and the solution kept in the dark for a
A green clear solution was
Ethanol was added to the solution to cause
the complex iron salt, K3[Fe(C2O4)3]. 3H2O
to precipitate since it is less soluble in alcohol
than in water.
After the week, the crystals formed were
The mass of the crystals were
found to be 2.055g.
CALCULATION AND EVALUATION OF DATA
Reaction between Di-ammonium iron (II) sulphate and 10%w/v oxalic acid solution
Mass of Di-ammonium iron (II) sulphate= 15g
volume of solvent = 50ml
M[Fe(NH4)2(SO4)2· 6H2O] = 56+2(14+4)+2[32+4(16)]+6(18)= 392g/mol
moles of [Fe(NH4)2(SO4)2· 6H2O] present in solution= 15/392= 0.0383mol
In every 100ml of solvent there was 10g of oxalic acid, hence 75ml will contain
n[H2C2O4] that was present in the reaction = 7.5/90= 0.0833mol
hence excess [H2C2O4] = 0.0833−0.0383= 0.0450mol
hence [Fe(NH4)2(SO4)2· 6H2O] was the limiting agent
mole ratio of [Fe(NH4)2(SO4)2· 6H2O]: [FeC2O4(s)]= 1:1
thus moles of [FeC2O4.2H2O(s)] produced = 0.0383mol
mass of [FeC2O4.2H2O(s)] produced= 0.0383×180= 6.894g
3g out of the 6.894 g of [FeC2O4.2H2O(s)] will contain (3×0.0383)/6.894= 0.01667mol
Reaction between FeC2O4(s) and H2O2
In every 100ml there is 6g of hydrogen peroxide
hence in 13ml there will be
M[H2O2]= 2+2(16)= 34g/mol
n[H2O2]= 0.78/34 = 0.023mol
mole ratio of H2O2 and FeC2O4.2H2O (s) = 1:2
hence 0.0383mol of FeC2O4.2H2O(s) will react with (0.5×0.01667)= 8.335×10-3
mol of the peroxide
FeC2O4.2H2O(s) is the limiting reagent in the reaction.
mole ratio between FeC2O4.2H2O(s) and K3[Fe(C2O4)3].3H2O= 1:1
mass of K3[Fe(C2O4)3].3H2O= 0.01667×491=8.18497
% yield= (2.055×100)/8.18497= 25.107%=25%
The synthesis of potassium tris(oxalato)ferrate (III) in the first part of this experiment involves a
two-step reaction scheme which first synthesizes potassium oxalate (K2C2O4·H2O), a reaction
intermediate. This compound is formed by the addition of oxalic acid to potassium hydroxide.
The chemical equation for this particular reaction is as follows:
K2C2O4·H2O(aq) + 2KOH (aq) K2C2O4·H2O (aq) + H2O(l)
Iron (III) chloride is then added to the reaction intermediate, forming our desired product
K3[Fe(C2O4)3].3H2O in the form of green crystals. The chemical equation of this is as follows:
3K2C2O4·H2O (aq)+FeCl3·6H2O (aq) K3[Fe(C2O4)3].3H2O (s)+3KCl(aq)+6H2O(l)
In the second part of this experiment, K3[Fe(C2O4)3].3H2O is put through three reactions that
provide further understanding of the properties and characteristics of this metal complex. In the
first reaction, 0.10g of solid potassium tris(oxalato) ferrate (III) is dissolved in 3.0 mL of 10%
acetic acid. This light green solution is then exposed to light. As [Fe(C2O4)3]3-
is light sensitive, Fe3+
will get reduced to Fe2+
and some oxalate ligands will get oxidized to
carbon dioxide (CO2) upon exposure to light. This phenomenon can be represented as the
following chemical equation:
(aq) + 5C2O4
This reaction forms aqueous iron (II) oxalate, which is brownish in colour, accounting for the
change in solution colour from light green to brown. The second reaction involves ligand
strength and replacement. When hydrochloric acid is added to the light green solution of
potassium tris(oxalato) ferrate (III), the solution turns yellow. This observation can be explained
by the fact that chloride ligands replaced the oxalate ligands bonded to the iron atom and formed
aqueous iron (III) chloride (FeCl3), which is yellow in solution. Although chloride anions
aren’t as strong as oxalate anions in terms of ligand strength, a great deal of chloride anions was
added. 3 drops of 6M HCl were added to a mere 2 mL of 0.20M potassium tris(oxalato)ferrate
(III). Chloride anions overwhelmed the iron (III) cations and formed yellow iron(III) chloride.
K3[Fe(C2O4)3].3H2O(aq) + 3HCl(aq) 3H2C2O4(aq) + FeCl3(aq) + 3KOH (aq)
When 3 drops of 0.5M KSCN are added to the mixture, they dissociate to form thiocyanate
anions (SCN-) which replace the chloride ligands. This move eliminates yellow iron (III) chloride
and forms dark red iron (III) thiocyanate Fe(SCN)3. This explains the change in colour from
yellow to dark red. FeCl3(aq) + 3KSCN (aq) Fe(SCN)3(aq) + 3KCl (aq)
When 10 drops of 3M KF are added, they dissociate to form potassium cations (K+
) and fluoride
). As fluoride anions are stronger ligands than thiocyanate anions, ligand replacement
occurs again, eliminating dark red iron (III) thiocyanate and forming yellow iron (III) fluoride
(FeF3) in its stead. Fe(SCN)3
(aq) + 3KF (aq) FeF3 (aq) + 3KSCN (aq)
Lastly, 15 drops of 1M oxalic acid (H2C2O4·2H2O) are added to the solution. They dissociate to
form hydrogen and oxalate ions. As oxalate ions are stronger ligands than fluoride ions, the
fluoride ligands in FeF3 get replaced, forming iron (III) oxalate.
FeF3(aq) + 3C2O4
(aq) [Fe(C2O4 )3]3-
(aq) + 3F-(aq)
After all these chemical reactions and ligand replacements, the final solution is dark yellow in
colour and contains many different ions. The last reaction involves the formation of a precipitate
after the addition of sodium hydroxide. The remaining 1 mL of 0.20M potassium
tris(oxalato)ferrate (III) solution is put in a test tube and mixed with 1 mL of 3M NaOH. This is a
precipitation reaction that forms iron (III) hydroxide (Fe(OH)3), a compound that is insoluble in
water. The chemical equation for this reaction is as follows: K3[Fe(C2O4)3].3H2O (aq) + 3NaOH
(aq)Fe(OH)3(s) + 3K+
(aq) + 3C2O4
(aq) + 3Na+
(aq) + 3H2O (l)
The precipitate filtered out and treated with 1 mL of 1M oxalic acid. This reforms the light green
solution of [Fe(C2O4 )3]3-(aq), with water as a by-product.
Fe(OH)3(s) +3H2C2O(aq) [Fe(C2O4 )3]3-
(aq) + 3H2O(l) + 3H+
Basically what happens here is this: our product in aqueous form reacts with NaOH to form a
solid(iron (III) hydroxide). When oxalic acid is added to this solid, our product gets reformed in
its aqueous state. The reaction to form Fe(OH)3 from [Fe(C2O4 )3]3-
can therefore be deemed
reversible. Such reversibility is due to the fact that these reactions are ligand replacement
reactions. In the spectrochemical series, hydroxide anions and oxalate anions are both of similar
ligand strength. Thus, the factor that determines if hydroxide anions bond to the iron atom (and
form a precipitate), or if oxalate anions bond to the iron atom (and form a light green aqueous
solution) is ion concentration. When 3M NaOH is introduced to our product (which it was),
some oxalate ions will definitely get replaced as both ligands have similar strengths, forming a
certain amount of solid iron(III) hydroxide. The reaction will gradually reach dynamic
equilibrium, where oxalate ligands and hydroxide ligands continually replace each other. When
the solid iron (III) hydroxide gets filtered out, the ligands left on the filter paper are mostly
hydroxide ligands as iron (III) hydroxide is the precipitate in this reaction. When oxalic acid gets
added, oxalate ligands get introduced, replacing some hydroxide ligands and forming a certain
amount of [Fe(C2O4 )3]3-
, which drips through the filter paper and gets collected as a light green
solution. The filtrate from the reaction of potassium tris(oxalato)ferrate (III) with sodium
hydroxide is then treated with 1 mL of 0.2 BaCl2. This filtrate includes potassium, oxalate and
sodium cations. Barium oxalate (BaC2O4), a white odourless powder, will be precipitated out.
This accounts for the white precipitate observed.
The empirical formula of our product can be determined by two methods. The first way is to
titrate a known amount of our product with potassium permanganate (KMnO4). The oxalate ion
in our product is a reducing agent that reduces KMnO4 to manganese ion (Mn2+
). The titration is
carried out by first creating a standard solution of KMnO4 with known volume and concentration.
A known mass of the product is then placed in a conical flask and diluted with excess H2SO4.
The endpoint is identified when the purple colour of the titrant remains in the beaker. MnO4-
reacts with C2O4
and sulfuric acid in the following formula:
From this titration, we can determine the concentration of the oxalate ions in the conical flask.
As we already know the concentration and volume of our product in the conical flask, we can
therefore easily determine its empirical formula. The second way in which we can determine the
empirical formula of our product is to determine the iron percentage instead of the oxalate
percentage stated above. This is also done via titration. The analyte is created by adding acid and
water to the crystals of product we obtained. 3% KMnO4 is then added and heated to near boiling
in order to get rid of the oxalate ions. This is followed by the addition of zinc powder. Finally,
the mixture is heated and filtered. The obtained filtrate is our desired analyte, which we can
titrate with known concentrations of KMnO4 in order to determine the percentage of iron present
in our product. If the iron percentage is known, we can then calculate the empirical formula of
our final product.
1. The solution was kept in a dark place to prevent solution of the ferrate oxalate complex from
being decomposed by light.
2. The solution was continuously stirred so that severe bumping cannot occur especially during the
We have successfully synthesized our desired product potassium tris(oxalato)ferrate (III) with a
yield of 25%. The product we obtained was then utilized as a reactant in various reactions that
demonstrated several concepts in chemistry such as photodecomposition, ligand strength and
2. Duncan, J. (2010). Experiment 1: synthesis and analysis of an inorganic compound.
Department of Chemistry, Plymouth State University, NewHampshire, US, United States.
InorganicCmpd_SynthAnalysis.pdf on 27.03.11
4. Coordination complex. (n.d.). Retrieved
fromhttp://en.wikipedia.org/wiki/Coordination_complex on 29.03.11
5. González , G., & Seco, M. (2004). Potassium tris(oxalato)ferrate(III): a versatile
compound toillustrate the principles of chemical equilibria. Journal of Chemical
Education, 81(8),Retrieved from http://pubs.acs.org/doi/abs/10.1021/ed081p1193
doi:10.1021/ed081p1193 on 29.03.11
6. Savelyev, G. G. (2003). The photochemistry of potassium trisoxalatoferrate(iii) trihydrate
inthe solid state. Journal of Solid State Chemistry, 12(1-2), Retrieved