ethanedioate complex of iron
Title: The preparation, analysis, and reactions of an ethanedioate complex of iron
Objective:
* To prepare oxalate complex of iron from the reaction with ammonium iron (II) sulphate
* To deduce the oxidation states of iron in the complex
* To determine the amount of iron and oxalate in the oxalate complex of iron by
titration.
* To investigate the molecular formula of the complex formed by preparation, analysis and the reaction of an ethanedioate (oxalate) complex of iron
* 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.
Materials and Methods
Procedure for prepartion
10.53g of ammonium iron(II) sulphate, (NH4)2Fe(SO4)2?6H2O , was weighed using a rough balance and put into a 400cm3 beaker, followed by the addition of few drops of dilute sulphuric acid and 30cm3 of water, and heated until a clear solution was obtained. Then a solution of 5.08g oxalic acid dihydrate in 30cm3 of water was added. The mixture was heated cautiously, with continuous stirring to boil. Then the yellow precipitate of hydrated iron(II) oxalate was allow to settle. The supernatant liquid was poured off and the precipitate was washed once with hot water by decantation. A hot solution of 7.49g of potassium oxalate monohydrate in 20cm3 of water was added and the mixture was cooled to 40?. 20cm3 of 6% hydrogen peroxide solution was added a few drops at a time with continuous stirring. The temperature must be kept below 50?. Then the mixture was heat to the boiling point and 2.503g of oxalic acid dihydrate was added. After stirring for 2-3 minutes, a clear green solution was obtained. Then, it was filtered using a filter paper and a long-stemmed funnel. 10cm3 of 95% ethanol was added to the clear solution to redissolve any precipitated crystals by warming on a steam bath. The solution bottle was then covered by aluminum foil and was set-aside in a dark cupboard for crystallization to occur. The crystals were separate by filtration using a sintered-glass funnel. Then it were washed twice with a mixture of equal volumes of ethanol and water, and followed by twice with acetone. The solid was allowed to air dry by suction. The weight of the dry solid was weighed and the yield was calculated.
Procedure for Analysis
(a)Oxidation with permanganate
Two portions of 0.38g of the sample complexes were added into two conical flasks separately. Each sample was dissolved in 10cm3 of water and 15cm3 of dilute sulphuric acid. When the solution was heated to 70?, the thermometer was removed and washed down, and was titrated with the standard, approximately 0.02M solution of permanganate. A pink color was persisting for 30 seconds at the end-point. The resulting solution was kept for the determination of iron in (b).
(b)Percentage of iron
15cm3 of concentrated hydrochloric acid was added to ...
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Procedure for Analysis
(a)Oxidation with permanganate
Two portions of 0.38g of the sample complexes were added into two conical flasks separately. Each sample was dissolved in 10cm3 of water and 15cm3 of dilute sulphuric acid. When the solution was heated to 70?, the thermometer was removed and washed down, and was titrated with the standard, approximately 0.02M solution of permanganate. A pink color was persisting for 30 seconds at the end-point. The resulting solution was kept for the determination of iron in (b).
(b)Percentage of iron
15cm3 of concentrated hydrochloric acid was added to the solution obtained in (a). When the solution was heated to the boiling point, tin(II) chloride was added a few drops at a time until the yellow colour of the complex was discharged. Then few more drops of tin(II) chloride were added . Once the solution was cooled to room temperature, one portion 10cm3 of mercury(II) chloride solution was added quickly. 10cm3 of 40% phosphoric acid and 10 drops of barium diphenylamine sulphonate indicator solution were added to the reduced solution. The solution was then titrated with the standard (approx. 0.01M) dichromate solution. Just before the end-point, the color of the solution was changed from colorless to grey-green; at this point the dichromate solution was added dropwise until the violet color persisted throughout the solution and this was taken as the end-point.
Procedure for the reactions of the complex Oxalate
Three tests tubes with our product and three with iron(III) chloride were prepared and marked correspondingly. Dilute sodium hydroxide, ammonium thiocyanate, and ammonium thiocyanate with dilute sulphuric acid were added separately into each of the three test tubes containing our product and the three containing iron(III) chloride. The interaction of aqueous solution of iron(III) chloride and the complex with the above added reagents was compared.
Results:
Mass of ammonium iron(II) sulphate used= 10.53g
Mass of diluted sulphuric acid used=5.08g
Mass of potassium oxalate monohydrate used = 7.49g
Mass of oxalic acid dihydrate used= 2.503g
Mass of our product obtained= 5.925g
Titration with permanganate titration
Mass of complex (g)
0.38
0.38
Initial reading (cm3)
1st 0.8, 2nd 2.6
1st 5.7, 2nd 18.3
Final reading (cm3)
1st 48.9, 2nd 5.7
1st 48.2, 2nd 18.8
Volume of 0.02M permanganate used (cm3)
48.1+3.1=51.2
42.5+0.5=43
The average of the weight of complex = (0.38+0.38)/2 = 0.38g
The average volume of permanganate used = (51.2+43)/2 = 47.1 cm3
Titration with dichromate
Weight if complex (g)
0.38
0.38
Final reading (cm3)
23.2
35.5
Volume of 0.01M dichromate used (cm3)
11.2
12.3
The average of the weight of complex = (0.38+0.38)/2 = 0.38g
The average volume of dichromate used = (11.2+12.3)/2 = 11.75cm3
Observation of the reactions of the complex oxalate
Iron(III) chloride
Complex
Dilute sodium hydroxide solution
The solution changed from yellow to pale yellow solution with reddish brown precipitate.
The solution changed from pale green to pale yellow solution with reddish brown precipitate.
Ammonium thiocyanate solution
The solution changed from yellow to clear dark red solution.
No observable change
Ammonium thiocyanate solution + dilute sulphuric acid
The solution changed form yellow to dark red solution.
The solution change form pale green to red solution.
Discussion
Question
)Deduction for the qualitative test
In the qualitative test, both the complex and iron (III) chloride were reacted with dilute sodium hydroxide solution respectively and brown precipitates were formed that indicated the presence of Fe(OH)3(s) :
Fe3+ + 3OH- › Fe(OH)3
For the reaction between iron (III) chloride reacted and ammonium thiocyanate, it showed a similar reaction:
Fe3+ + SCN- › Fe(SCN)2+
The dark red solution formed indicated the formation of Fe(SCN)2+ .However,
as Fe3+ form stable complex with the oxalate, the Fe3+ ions cannot react with CNS- ions to form red complex. While in the presence of dilute sulphuric acid, more Fe3+ ions can be released to form red complex with CNS- ions. Thus, the resulting solution is dark red.
As the complex prepared had similar reactivity as FeCl3, we can assume that the oxidation state of iron in the complex is also +3.
2)Calculations
a) From the dichromate titration,
Cr2O72-+6Fe2++14H+ › 2Cr3++6Fe3++7H2O
Volume of dichromate used = 11.75cm3
Molarity of dichromate = 0.01M
no. Of moles of dichromate ions = 0.01×11.75×10-3
= 1.175×10-4mol
Mole ratio of Cr2O72-: Fe3+ = 1: 6
No. of moles of iron(III) ions = (1.175×10-4)× 6
= 7.05×10-4mol
The mass of complex containing 1 mole of iron = 0.38/ 7.05×10-4mol
= 539.01g
b) From the permanganate titration,
2MnO4-+5C2O42-+6H+ › 2Mn2++10CO2+8H2O
Volume of permanganate used = 47.1cm3
Molarity of permanganate used = 0.02M
no. of moles of permanganate ions
= 0.02×47.1×10-3
= 9.42×10-4mol
Mole ratio of MnO4-: C2O42- = 2:5
no. of moles of oxalate ions = (9.42×10-4×5)/2
= 2.355×10-3mol
Mole ratio of C2O42-:Fe3+ =2.355×10-3: 7.05×10-4
= 3.34:1
~3:1
As the sample complex contain iron(III) ions, oxalate ions, potassium ions and water molecules, the formula of the complex is KxFe(C2O4)3.yH2O
Since the complex is electrically neutral, so the no. of moles of potassium:
[(+1)(x)]+(+3)+(-2)(3) = 0
x = 3
?the formula of the complex is K3[Fe(C2O4)3].yH2O
3×39.10+55.85+3×(2×12.01+4×16.00)+y(1.01×2+16.00)=539.01g
437.21+ 18.02y = 539.01g
18.02y = 101.8g
y = 5.649
y~6
?the formula of the hydrated complex is K3[Fe(C2O4)3]?6H2O
c) no. of moles of ammonium iron(II) sulphate
= 10.53g/[2(14.01)+2(32.07)+14(16.00)+20(1.01) +55.85]
= 0.0268mol
no. of moles of complex= 5.925/539.01
= 0.011mol
Percentage yield of the complex= _______no. of mole of iron in the complex______
no. of mole of iron in ammonium iron(III) sulphate
= 0.011/ 0.0268 ×100%
= 41.04%
3)The structure for the complex anion present in the product:
The complex anion lacks a centre and a plane of symmetry so optical isomers exists so the compound could in principle exhibit optical isomerism.
Source of errors and improvement
The percentage yield of the complex was only 41.04%, it might due to the loss of product and the presence of impurities. For instance, too much acetone was used to wash the crystals as a result some of the crystals are dissolved and lost. Moreover, there was loss of sample during the transfer process as some solid might adhere on the flasks or Buchner funnel. For improvement, proper amount of acetone should be used to wash the crystals to minimize the loss of product. Slowly pour off the supernatant liquid and try to avoid the loss of solids. More sets of titration can be done in order to improve the accuracy. When adding hydrogen peroxide solution into the mixture during the preparation of the oxalate, the process should be taken slow enough to ensure that the temperature does not rise above 50?. Otherwise, side products may come out which affect the yield of our final product. Also, the oxalate should be equally placed in the sintered-glass funnel to ensure best filtration.
According to our calculation, the formula of the hydrated complex is K3[Fe(C2O4)3].6H2O. However, the formula of the hydrated complex should have 3 water molecules theoretically. This may due to the inaccurate mass of the complex due to the experimental error such as the presence of extra water or impurities which increases the final mass of the complex.
Some important notes and safety precaution
Aluminium foil should be used to wrap the conical flask containing the product before putting it in a dark place as the product is photosensitive.
Sintered-glass was used instead of filter paper as filter paper is easier to retain the product causing lost in yield. After the suction filtration, remove the rubber tube first instead of the sintered-glass funnel to prevent sucking back of the filtrate.
As concentrated hydrochloric acid was highly corrosive, it can attack the respiratory system and skins, the adding process should be preceded in the fume cupboard and should take more care when pouring it. Moreover, most of the reactants used in this experiment are highly toxic, we should avoid handling all these solid and avoid spillage of solutions. When handling highly flammable ethanol, no naked flames in the vicinity should exist.
Conclusion:
In this experiment, the mass of the complex prepared was 5.925g and the percentage yield is 41.04%. The oxidation state of the iron in the complex is +3 and the mass of complex containing 1 mole of iron is 539.01g. The formula of the hydrated complex is K3Fe(C2O4)3?6H2O and the complex is optically active according to its structure.