Synthesis of Potassium Tris(Oxalato)Ferrate (III)

Experiment 6 Synthesis of Potassium tris(oxalato)ferrate (III) Timothy Tan Xin Zhong M11605

Experiment 6 Synthesis of Potassium tris(oxalato)ferrate (III)

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 (C2O42-) 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.

Figure 1: 3-dimensional structure of potassium tris(oxalato)ferrate (III)

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

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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,

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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)33-. 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. This phenomenon can be represented as the following chemical equation: 2[Fe(C2O4)3]3- (aq)  2Fe2+ (aq) + 5C2O42- (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

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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 (FeF 3) 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) + 3C2O42- (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) + 3C2O42- (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

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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) + C2O42- (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 (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 color of the titrant remains in the beaker. MnO4- reacts with C2O42- and sulfuric acid in the following formula: 5C2O42- + 2MnO4- + 16H+  10CO2 + 2Mn2+ + 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.

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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. References 1. Riemenschneider , W., & Tanifuji, M. (2000). Oxalic acid. doi: 10.1002/14356007.a18_247. Retrieved on 27.03.11 2. Eiichi, Y., Tomiya, I., Tsuyoshi, S., Yukio, Y. (1972). US Patent No. 3,678,107. Washington, DC: US Patent and Trademark Office. Retrieved on 27.03.11 3. Duncan, J. (2010). Experiment 1: synthesis and analysis of an inorganic compound. Informally published manuscript, Department of Chemistry, Plymouth State University, New Hampshire, US, United States. Retrieved from http://oz.plymouth.edu/~jsduncan/courses/2010_Fall/InorganicChemistry/Labs/1InorganicCmpd_SynthAnalysis.pdf on 27.03.11 4. Coordination complex. (n.d.). Retrieved from http://en.wikipedia.org/wiki/Coordination_complex on 29.03.11 5. González , G., & Seco, M. (2004). Potassium tris(oxalato)ferrate(iii): a versatile compound to illustrate 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 in the solid state. Journal of Solid State Chemistry, 12(1-2), Retrieved from http://www.sciencedirect.com/science?_ob=ArticleURL&_udi=B6WM2-4B6NPPXH5&_user=10&_coverDate=01%2F01%2F1975&_rdoc=1&_fmt=high&_orig=gateway&_origi n=gateway&_sort=d&_docanchor=&view=c&_acct=C000050221&_version=1&_urlVersion= 0&_userid=10&md5=c2e314ad54e5453fd871881144379b59&searchtype=a doi: 10.1016/0022-4596(75)90183-8 on 29.03.11

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Synthesis of Potassium Tris(Oxalato)Ferrate (III)

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