Experiment on the preparation of potassium tris (oxalate) ferrate (ii) trihydrate


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Experiment on the preparation of potassium tris (oxalate) ferrate (ii) trihydrate

  2. 2. 2 TITLE: PREPARATION OF POTASSIUM TRIS(OXALATE) FERRATE(II) TRIHYDRATE 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. 3. 3 INTRODUCTION 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 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. 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.
  4. 4. 4 The oxalic acid utilized in the first step of this reaction scheme can be synthesized by hydrolyzingcyanogen1 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 by hydrolysis1 . 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.
  5. 5. 5 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: 2 Fe(2+)  2 Fe(3+) + 2e- H2O + HO2 (–) + 2e-  3 OH(–) 2: H2O + HO2 (–) + 2 Fe(2+)  2 Fe(3+) + 3 OH(–) (net reaction) 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: 3: Fe(3+) + 3OH(–)  Fe(OH)3 With the addition of more oxalic acid, the ferric hydroxide dissolves and the soluble complex is formed: 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.
  6. 6. 6 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 9. Oven 10. Water bath containing ice and water 11. Stirrer 12. Electronic balance 13. 50ml measuring cylinder 14. 5ml measuring cylinder 15. Two beakers 16. Mercury thermometer
  7. 7. 7 PROCEDURE STEP OBSERVATION INFERENCE 1ml of 2M H2SO4 was weighed into a beaker container 50ml of water and the solution heated. A colourless solution was formed. The acid dissociated totally in the aqueous solution according to the equation. H2SO4 + 2H2O 2H3O+ (aq) + SO2- 4(aq) 15g of Di-ammonium iron (II) sulphate was weighed and added to the warm solution prepared. The solution changed to light green. Iron (II) is oxidized slightly by the acid to iron (III). 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 an oven. A dry yellow precipitate was obtained. 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
  8. 8. 8 by H2O2. 13ml of 6% w/v hydrogen peroxide was added and the solution stirred continuously. The solution changed colour from yellow to green with effervescence. The peroxide oxidizes the ion from Fe2+ ion to Fe3+ ion. H2O2 + 2Fe2+ 2Fe3+ + 3OH- 25ml of 95% ethanol was added to the solution and the solution kept in the dark for a week. A green clear solution was formed. 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 weighed. 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 (10×75)/100=7.5g M[H2C2O4]= 2+2(12)+4(16)=90g/mol n[H2C2O4] that was present in the reaction = 7.5/90= 0.0833mol
  9. 9. 9 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 (13×6)/100= 0.78g 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 n[K3[Fe(C2O4)3].3H2O]= 0.01667mol M[K3[Fe(C2O4)3].3H2O]=3(39)+56+6(12)+12(16)+3(18)=491g/mol mass of K3[Fe(C2O4)3].3H2O= 0.01667×491=8.18497 % yield= (2.055×100)/8.18497= 25.107%=25% 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
  10. 10. 10 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: Fe(C2O4)3]3- (aq) 2Fe2+ (aq) + 5C2O4 2- (aq)+2CO2 (g) 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
  11. 11. 11 (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 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 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 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-(aq), with water as a by-product.
  12. 12. 12 Fe(OH)3(s) +3H2C2O(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 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. 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
  13. 13. 13 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 2- and sulfuric acid in the following formula: C2O4 2- + 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. PRECAUTION 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 oxidation reaction.
  14. 14. 14 CONCLUSION 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 ligand replacement. REFERENCES 1. http://www.scribd.com/doc/52686651/Synthesis-of-Potassium-Tris-Oxalato-Ferrate- III#fullscreen=1 2. Duncan, J. (2010). Experiment 1: synthesis and analysis of an inorganic compound. Department of Chemistry, Plymouth State University, NewHampshire, US, United States. Retrieved from 3. http://oz.plymouth.edu/~jsduncan/courses/2010_Fall/InorganicChemistry/Labs/1- 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 fromhttp://www.sciencedirect.com/science
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