Synthesis K3[Fe(C2O4)3]

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Experiment 6 
Synthesis of Potassium tris(oxalato)ferrate (III) 
 
Timothy Tan Xin Zhong 
M11605 
 
 
 
 
 
Experiment 6 
Synthesis of Potassium tris(oxalato)ferrate (III) 
 
Page 1 
 
 
 
Experiment 6 – Synthesis of Potassium tris(oxalato)ferrate (III) 
 
Aim 
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. 
 
Introduction 
Potassium tris(oxalato)ferrate (III) is a metal complex of iron with 
three oxalate ligands (C2O4
2-) 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. Potassium 
tris(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: 
3K2C2O4·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 with a yield of 55.83%. 
 
The oxalic acid utilized in the first step of this reaction scheme can be synthesized by hydrolyzing 
cyanogen1 or by oxidizing sucrose or glucose with nitric acid in the presence of a small amount of 
vanadium pentoxide.2 Another method of forming oxalic acid involves the oxidative carbonylation 
of alcohols followed by hydrolysis.1 
 
 
 
Figure 1: 3-dimensional structure of 
potassium tris(oxalato)ferrate (III) 
Experiment 6 
Synthesis of Potassium tris(oxalato)ferrate (III) 
 
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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 sulfate 
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 (K2C2O4) 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 favored 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. 
 
Experimental Procedure 
The first part of our experiment involved synthesizing our desired product, potassium 
tris(oxalato)ferrate (III). This was done by mixing 13 mmol of oxalic acid (H2C2O4·2H2O) with 24 
mmol of potassium hydroxide (KOH) in a 25 mL conical flask. 7 mL of distilled water was added to 
the mixture. The flask was then gently heated until complete dissolution took place. 4 mmol of iron 
(III) chloride hexahydrate (FeCl3·6H2O) was then added to the mixture. 
 
The solution was filtered into another 25 mL conical flask. This flask was then wrapped thoroughly 
in aluminum foil and placed in ice for 30 minutes. As we had no problems producing visible crystals, 
we did not have to scratch the inner walls of the flask. The green crystals produced were then 
collected via suction filtration. Recrystallization was carried out on the crystals with minimal 
amounts of hot water. The weight of the crystals produced was then recorded. 
 
In the second part of our experiment, we put the synthesized crystals through three different 
reactions. The first reaction involved the photodecomposition of potassium tris(oxalato)ferrate (III). 
This was done by dissolving 0.10g of the product in a test tube with 3.0 mL of 10% acetic acid 
(CH3COOH). The solution was then exposed to light for 30 minutes, turning brown after a period of 
time. 
 
As the other two reactions needed 0.20M potassium tris(oxalato)ferrate (III) solution, 2.0 mL of this 
solution was prepared. 1.0 mL of this solution was placed in a test tube for the second reaction. 
First, 3 drops of 6M HCl was added. This was followed by 3 drops of 0.5M KSCN, 10 drops of 3M KF, 
and 15 drops of 1M H2C2O4·2H2O. Upon addition of HCl, the light green potassium 
tris(oxalato)ferrate (III) solution turned yellow. Adding KSCN turned the solution dark red. Finally, 
Experiment 6 
Synthesis of Potassium tris(oxalato)ferrate (III) 
 
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the addition of KF turned the solution yellow, while the addition of oxalic acid turned the solution 
dark yellow. These observations were duly recorded. 
 
For the last reaction, 1 mL of 3M NaOH was added to the remaining 1.0 mL of 0.2M stock solution. 
A reddish brown precipitate was formed. This precipitate was separated from the solution via 
filtration and treated with 1 mL of 1M H2C2O4·2H2O, reforming Fe(C2O4)3
3-. The filtrate was also put 
through a chemical reaction with 1 mL of 0.2M BaCl2. This formed a white precipitate. These 
observed changes were also properly logged down. 
 
Results and Observations 
Molecular weight of K3[Fe(C2O4)3]·3H2O = 491.25g/mol 
Expected moles of product produced = 4 mmol 
 
Expected mass of product produced = 491.25g/mol X 4 mmol = 1.965g 
Actual mass of product produced = 1.097g 
Yield = 
 
 
 
 
 
 
 
Discussion 
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: 
H2C2O4 (aq) + 2KOH (aq)  K2C2O4·H2O (aq) + H2O (l) 
 
Iron (III) chloride is then added to the reaction intermediate, forming our desired product 
K3Fe(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, K3Fe(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. Thisphenomenon can be represented as the following chemical equation: 
2[Fe(C2O4)3]
3- (aq)  2Fe2+ (aq) + 5C2O4
2- (aq) + 2CO2 (g) 
 
This reaction forms aqueous iron (II) oxalate, which is brownish in color, accounting for the change 
in solution color 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 
Experiment 6 
Synthesis of Potassium tris(oxalato)ferrate (III) 
 
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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 color 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 
anions (F-). 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
2- (aq)  [Fe(C2O4)3]
3- (aq) + 3F- (aq) 
 
After all these chemical reactions and ligand replacements, the final solution is dark yellow in color 
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
2- (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-, with water as a byproduct. 
Fe(OH)3 (s) + 3H2C2O4 (aq)  [Fe(C2O4)3]
3- (aq) + 3H2O (l) + 3H
+ (aq) 
 
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 
Experiment 6 
Synthesis of Potassium tris(oxalato)ferrate (III) 
 
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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 odorless powder, will be precipitated out. This accounts for the white 
precipitate observed. 
Ba2+ (aq) + C2O4
2- (aq)  BaC2O4 (s) 
 
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 (Mn
2+). 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 color of the titrant remains in the beaker. MnO4
- reacts with C2O4
2- and 
sulfuric acid in the following formula: 
5C2O4
2- + 2MnO4
- + 16H+  10CO2 + 2Mn
2+ + 8H2O 
 
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. 
 
 
Experiment 6 
Synthesis of Potassium tris(oxalato)ferrate (III) 
 
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Conclusion 
We have successfully synthesized our desired product potassium tris(oxalato)ferrate (III) with a 
yield of 55.83%. 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 
ligand replacement. 
 
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