Iron Oxalate Lab

Marina Horta Iron Oxalate Lab 10/22/09

Abstract; The purpose of this lab was to synthesize, isolate, and determine Fe content of a complex; and calculate the formula of the Potassium Oxalate Iron(III) Complex using chemical analyses. This lab used filtration, titration, acid base reactions, redox , synthesis, dilutions , beers law and red tide computer system to establish that the final formula for the complex was K3[Fe(C2O4)3] ∙2 H2O.

Introduction;

Part 1; The purpose of part 1 to synthesize, purify and, mass crystals of potassium oxalate ferrate (III) complex. Potassium oxalate ferrate (III) is a complex, an ion consisting of a ligand and a metal covalently bonded together. A ligand is an ion with the tendency to bond with metals because it typically donates electron pairs. This relationship between metals and ligands is only one of the key components for the synthesis of the crystals. For the precipitation of the complex ion alcohol was used to reduce solubility,

Part: 2 Part two determines the percent oxalate in the complex produced in the previous lab. Using volumetric analysis a method (similar to gravimetric analyses) in which the volume of a known substance is used to quantitatively measure the volume of an unknown substance required to react with it, or standardization a process in which the value of a potential standard is fixed by a measurement made with respect to a standard whose value is known and can be used to determine the value of unknown. Using the concentration and liters needed during the titration to react the number of moles of unknown can also be calculated. To determine the percent oxalate the end point of the titration is reached and recorded, once the solution has changed color according to the indicator in the an acid/base reaction, or a precipitant is formed in a double , single/ replacement reaction or Redox. Percent by mass can then be determined through stoichiometric calculations using the concentration, moles, liters and molar weight of the known and unknown compounds.

Part 3; The final lab is the analysis of the iron, in order to evaluate only the Fe content, calcium is used to precipitate, and vitamin C as an absorbing agent. In this third portion of the Iron oxalate experiment spectrophotometry a method of transmitted or absorbed in the electromagnetic spectrum in comparison it to the initial tests of blank and dark trails is used to measure the concentration of solutions using Beers Law. Beer’s Law describes between the relationship between absorption of energy measured and its concentration, where A is absorbance, is the molar absorptive (which depends on the chemical ...

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The samples measured by spectrophotometrics are smaller samples of serially diluted solutions, usually occurring in logarithms of -10 each time. The final procedure in the iron oxalate experiment was the empirical formula to find the values for x, y, and z.

Materials; Part 1; Hydrated salt of Fe(NH4)2SO4 * 6H2O 150 ml beaker, distilled water, weighing boat, 1M oxalic acid, 6M H2SO4, Hydrogen Peroxide, 95% ethanol, Aluminum Foil, Buchner Funnel, Hot plate, 1.8M potassium oxalate solution 1 M H2C2O4 Part 2; permanganate solution, sulfuric acid, tongs, beaker, 250 ml Erlenmeyer flask , 150 ml flask, distilled water , burette, Part 3; 25 ml of a standard solution, 1M Calcium Chloride, 100 ml beaker, Distilled water, 9M H2SO4 filtration, filter paper , Carbonic Acid, pipettes, ph 4.7 buffer solution, 125 ml Erlenmeyer flask , test tubes

Safety; Dilute Sulfuric Acid is corrosive and poisonous (labs part one and two) Hydrogen peroxide is an oxidizer and corrosive( lab one) Oxalic Acid and oxalate ions are toxic (lab one ) Ethanol is flammable (lab one) Permanganate solution is also poisonous and corrosive Take caution when handling hot glass wear (lab part two)

Calculations;

Table 1 Weight of Crystals (part 1)

Initial trey 1.49 (g)

Weight of crystal 1.3934 (g)

Weight of both 2.88 (g)

Table 2 Data (part 2)

Data mass of 1 2 3 4 ii iron oxalate 1.4974 g 1.4994 g 1.4977 g 1.4994 g

Crystals+ Trey 1.6286 g 1. 6175 g 1.6256 g 1.6247 g

Mass of Crystal .1294 g .1181 g .1288 g .1253 g

Titration ; (part 2) 1 2 3 4

Final 36.5ml 36.5ml 48.6 ml 48.4 ml

Initial 6 ml 8.5ml 18 ml 18 ml

Volume used 30.5ml 28 ml 30.6 ml 30.4 ml

Moles

Trial #1 (.01057 M)(.0305 ml) =.003224 m (5/2) = .0008059

Trial #2 (.01057M)(.028 ml) = .002959 m (5/2) =.0007398

Trail #3 (.01057M)(.0306 ml) = .0003234 m (5/2)=.0008085

Trail #4 (.01057 M)(.03 04 ml) = .0003213 m (5/2) = .00080325

Percent by Mass Oxalate

Trial# 1 (.0008059 m)(87.99) =(.07091)(1/.1294)= 54%

Trial #2 (.0007398m) (87.99) = (.06509)(1/.1181) = 55%

Trail #3 (.0008085m)(87.99) = (.0711#)(1/.1288) = 55%

Trial #4 (.00080325m )(87.99) = (.07067)(1/.1253) =56%

Empirical Formula

Fe+ 11 g 11/55.85 .1525 moles 1

C2O4-2 55g 55/82.02 .6248 moles 3

Percent Yield = 1.848/ 1.3 = 69%

Analysis; The purpose of this experiment was to determine the values of x y and z in the equation, the final (corrected) results of this lab were K3[Fe(C2O4)3] ∙2 H2O. The first experiment synthesized the green crystals through redox, precipitation and filtration and occurred largely without any significant error or deviation from the procedure all the observations were made at the correct points and reactions occurred in time. The second days trials calculated the percent by mass of oxalate in the solutions; the average value of the three trials was 55% which lies within the expected correct percent of 50-60% oxalate. An accurate result for the percent oxalate shows that the crystals in the first lab were synthesized correctly, and the titration process in the second lab occurred without error. The trials of titration used had an average of 30.2 ml their percent difference was low; enough to assume that there was no problem of over titrating (which would cause % oxalate to be off) or question whether the crystals were properly precipitated. The majority of the errors causing inaccurate values for x, y, and z occurred during the third and final days experiments. The standard curve of the Fe+ solution was first calculated using absorbance and concentration however because of random measurement errors during the serial dilutions the curve was inaccurately and imprecisely graphed too low; as a result the calculated value for r2 was not close to the correct whole number one, again an imprecise error caused by the dilutions. These dilution errors were caused by not properly using the pipettes, and a correction for any following experiments would be to more carefully A third error that took place during the experiment was the negative value for two of five concentrations of the Fe+, this error could be again because of the dilutions or it could be a result of misusing the red tide equipment however because time did not allow the negative values were kept and simply not included in the final calculations of the unknown. The actual percent iron calculated for this lab was 4.5% a value much too low to be correct or useful this lower percent Fe+ shows that the concentration of Fe+ was too low to begin or not enough of the solution was pipetted to the initial 100ml flask. The average of all the experiments 8.52 was also too low, in order to get a correct value using the empirical formula from the percent iron the percent had to be at least 11 % which shows a systematic error in the experiment. Using the empirical formula and the adjusted 11% iron the ratios of the iron and oxalate was calculated to be 3:1 instead of the original incorrect 4:1 ratio using the lower value for iron. The value for z was calculated within the reasonable 2 or 3 value, at 2.4. The final determination in the experiment was theoretical and percent yield from the moles of the third and grams of the second day. The theoretical yield was 1.848 grams while the results of the second day yielded only 1.3 resulting in a 69% yield overall. The final adjusted formula was finally determined to be K3[Fe(C2O4)3] ∙2 H2O.