Experiment on the preparation, analysis and reactions of an ethanedioate (oxalate) complex iron


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Experiment on the preparation, analysis and reactions of an ethanedioate (oxalate) complex iron

  2. 2. TITLE: THE PREPARATION, ANALYSIS AND REACTIONS OF AN ETHANEDIOTE (OXALATE) COMPLEX IRON 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
  3. 3. 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 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
  4. 4. unidentate ligands surrounding a central metal ion (Figure 1). The ligands act as Lewis bases, donating at least one pair of electrons to the Fe3+ ion. Oxalate ion, C2O4 2– , acts as a chelating bidentate ligand, donating 2 electron pairs from 2 oxygen atoms to the transition metal center, Fe3+ . 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 Fe2+ must be oxidized to Fe3+ . Complex ions that form with the Fe3+ have a different number of oxalate groups than those that form with Fe2+ . Empirical formula determination is difficult with a mixture of the two complex ions. At this point, the Fe3+ 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 Fe3+ to Fe2+ . During the second week, the crystallized salt will be analyzed to determine the mass percent
  5. 5. 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 unbalanced reaction below: (1) MnO4 – (aq) + C2O4 2– (aq)  Mn2+ (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 Mn2+ 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, Fe3+ , is released into solution when permanganate ion reacts with oxalate ion and destroys the complex ion. The liberated Fe3+ 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 Fe3+ to yield a colourless complex ion, Fe(PO4)2 3– , eliminating the reddish-brown colour of Fe3+ 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 Fe2+ ion is then titrated with a standardized KMnO4 solution according to the unbalanced equation below:
  6. 6. (2) Fe2+ (aq) + MnO4 – (aq)  Fe3+ (aq) + Mn2+ (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 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)
  7. 7. 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. The solution was cooled to 40o 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 Fe3+ + 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 to70o 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+ Fe2+ +3H2C2O4 MnO- 4+5H2C2O4+6H+ 2Mn2+ +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
  8. 8. 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 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. PRECAUTIONS 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
  9. 9. 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 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.