PREPARATION AND ANALYSIS OF AN IRON COORDINATION
COMPOUND1
This five lab series of experiments involves the preparation of an iron coordination compound
followed by quantitative analysis for water, oxalate, potassium and iron. You will perform quantitative
tests on the coordination compound to determine the empirical formula. The percent water is determined
gravimetrically. A solution of potassium permangate is prepared and standardized. This solution is then
used to determine the percent oxalate. Finally, percent potassium and iron is determined by ion exchange
followed by titration of the eluted solution with standard sodium hydroxide (using the standardized
solution of NaOH previously preared). After determining the empirical formula the structure of the
complex ion is drawn.
PREPARATION OF THE COMPLEX. The time required for this synthesis and the purification
is usually about two hours (including the times for cooling and recrystallization from hot water). Students
normally obtain between 30 and 50% of the theoretical yield. The yield depends on the time permitted for
the cooling and the volume of water used for the recrystallization.
DETERMINATION OF PERCENT POTASSIUM. The percentage of potassium is determined
by ion exchange followed by titration with a standard sodium hydroxide solution. The equivalence point
is determined with a pH meter.
DETERMINATION OF PERCENT IRON AFTER ION EXCHANGE. The percentage of iron is
determined by ion exchange followed by titration with a standard sodium hydroxide solution. The
equivalence point is determined with a pH meter. Each mole of iron takes three moles of sodium
hydroxide and Fe(OH)3 is formed. The potassium and iron are determined from the same solution (one
titration).
DETERMINATION OF PERCENT OXALATE. The percentage of oxalate is determined by
titration with standard potassium permanganate. The titrations are done at 60oC.
DETERMINATION OF PERCENT WATER. The percentage of water hydration is determined
by heating a weighed sample of the green hydrate in an open container in an oven until all the water of
hydration has been driven off at 1100 C.
DETERMINATION OF THE MOLECULAR FORMULA. After you have determined the mass
percentage of iron, potassium, oxalate, and water, the simplest empirical formula is calculated.
LITERATURE CITED
Aravamundan, G., J. Gopalakrishnan, M.R. Udupa, J. Chem. Educ. 1974 51, 129.
Aten, C.F. Jr., J. Chem. Educ. 1973 50, 656.
Baker, A.D., A. Cassadavell, H.D. Gafney, M. Gellender, J. Chem. Educ. 1980 57, 314- 315.
Bailar Jr., J.C. Jones E.M., "Inorganic Synthesis," McGraw Hill, Inc. New York, 1939, Vol I, p. 36.
Brooks, David W., J. Chem. Educ. 1973 50, 218.
Jekel, E and E. Jekel, Hope College, Personal Communications.
Johnson, R.C., J. Chem. Educ. 1970, 47, 702.
Olmsted, J., J. Chem. Educ. 1984, 61, 1098 -1099.
Skoog, West, Holler., "Fundamentals of Analytical Chemistry," Saunders, Philadelphia (1992) P. 882.
Young, C.G., J. Chem. Educ. 1985 62,445.
The Synthesis of an Iron Coordination Compound
Read this entire laboratory procedure before starting this lab.
Introduction:
In this experiment you will synthesize a compound that will contain the elements potassium, iron,
carbon, hydrogen, and oxygen. The final product may be given the formula KxFe(C2O4)y zH2O, where
the zH2O is called the water of hydration. This is the first experiment of a series in which you will
synthesize the compound and then determine its simplest formula (i.e., x, y, z) using a variety of analytical
techniques.
One important factor in any chemical synthesis is the actual quantity of desired product obtained
compared to the theoretical amount predicted on the basis of the stoichiometry of the reaction. The ratio
of the mass of product obtained to the theoretical quantity, expressed as a percentage, is referred to as the
"percent yield" or more simply the "yield".
There are many reasons why the actual yields are not 100 percent. Possibly the reaction reaches
equilibrium instead of going to completion. Maybe the reactants are involved in other reactions than the
one that produces the desired product. Probably some product is lost in crystallizing and separating
crystals from supernatant liquid, etc.
In this experiment you will react an aqueous solution of iron (III) chloride (FeCl3)with an aqueous
solution containing excess potassium oxalate (K2C2O4)to produce the product KxFe(C2O4)y zH2O.
x K+(aq) + Fe3+(aq) + y C2O42–(aq) + z H2O  KxFe(C2O4)y zH2O(s)
Apparatus:
Two 50 mL beakers
0.001 g balance
crucible tongs
250 mL beaker
watch glass
vacuum filtration apparatus
small vial with cap
aluminum foil
Chemicals:
FeCl3 solution (0.400 g FeCl3/mL)
K2C2O4H2O
Ice
Acetone.
Safety, Environmental, and Economic Concerns:
1. Be sure to wear safety goggles or other eye protection.
2. If any of the solutions come in contact with your skin, wash with soap and water generously.
3. Discard all waste solutions in the designated containers.
4. Avoid over heating resulting in violent boiling when heating liquids.
5. Iron compounds may stain your clothing permanently.
6. ACETONE IS A FLAMMABLE SOLVENT; extinguish all flames and turn off hot plates in your
work area while using it. Make sure that the acetone bottle is covered when not in use. If a spillage
occurs, make sure all burners and hot plates in the area are turned off and sponge or mop up the
spilled liquid using large quantities of water to rinse it down the drain.
Pre-Lab
1. Read the procedure that follows. Prepare and record in your lab notebook a data table for this experiment.
2. In this experiment, you will react an aqueous solution containing 2.95 g FeCl3 . 6H2O (270 grams/mol)
with an aqueous solution containing excess K2C2O4 to produced product KxFe(C2O4)y . zH2O.
a. Assuming that all of the Fe originally in FeCl3 . H2O ends up in the product, KxFe(C2O4)y . zH2O,
how many moles of product should be obtained?
b. What additional information would you need to calculate a percent yield?
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Notes on Experimental Procedure:
1. When a desired product is formed by crystallization from a reaction mixture containing excess
reactants and other products, the crystals are likely to be relatively impure. The crystals can be
separated from the impure solution (often called ``mother liquor'') by filtration or decantation. The
mother liquor clinging to the crystals can be removed by washing with an appropriate solvent.
However, washing will not remove impurities occluded in the crystals.
2. A standard method for purifying a crystalline product is recrystallization. For crystals that are more
soluble in hot solvent than in cold solvent, the recrystallization can be done by dissolving the crystals
in a minimum quantity of hot solvent and then cooling. The purified crystals can then be ``harvested''
by filtration. Larger, purer crystals are obtained if the hot solution is allowed to cool slowly without
being moved or disturbed. A second ``crop'' of crystals could be obtained from the filtrate by
evaporating a fraction of the solvent by heating followed by cooling the remaining solution. The
second crop of recrystallized product is generally less pure than the first.
3. The wet crystals dry very slowly. The purpose of the acetone is to wash the water off the crystals.
The acetone, which has a high vapor pressure, evaporates quickly leaving the crystals dry. Since
acetone is flammable make sure there are no flames in your work area when you do your acetone
washes.
4. The product will slowly decompose when exposed to light. Hence the crystals should be air dried in
the dark, closed drawer and eventually stored in a brown bottle.
Experimental Procedure:
Day 1
1. Obtain (in a clean 50 mL beaker) 8.00 mL of stock solution containing 0.400 g FeCl3/mL.
(Record in volume in your notebook.) Weigh 12.0 to 12.5g K2C2O4 H2O, into a clean dry 50 mL
beaker using the 0.01 g balance. (Record the mass in your notebook.) Add 20 mL of deionized
water (dH2O) to dissolve the K2C2O4 H2O. Heat on the hot plate stirring constantly until the
K2C2O4 H2O is completely dissolved.
2. Using the crucible tongs to handle the hot beaker, pour the hot solution into the beaker containing
the FeCl3 solution and stir. (Record any observations in your notebook.)
3. Cool the solution for 30-45 minutes by placing the beaker in a larger beaker containing ice and
water. Crystals should form during this time. Take care that the beaker of product does not sink
in the water as the ice melts.
4. After giving the crystals ample time to form, carefully pour off and discard the solvent without
removing any crystals, a process called decantation.
5. Add 20 mL dH2O to the crystals, heat gently with stirring to completely dissolve the crystals (note
1). If some dark residue remains undissolved, carefully decant the clear solution into another
beaker and discard the residue.
6. Cool the solution for 30 minutes by placing the beaker in a larger beaker containing ice and water.
Crystals should form during this time. If the crystals are allowed to
form slowly without being disturbed, large crystals will be obtained.
If the solution is moved or stirred while crystals are forming, smaller
crystals will result. Clean a small vial and dry it carefully with a
paper towel.
Day 2
1. Set up a vacuum filtration system as shown in the picture, then filter
the crystals by vacuum filtration using a Buchner funnel and a clean
filter flask. If your crop of crystals appears to be quite small, save
the filtrate so that you can obtain a second crop of crystals. Consult
your instructor about the procedure.
2. Wash the crystals twice with ice water. Use less than 5 mL of ice
water for each wash and work quickly to avoid dissolving the
product in the wash water. Finally wash the crystals twice with 5
mL portions of acetone (note 2).
3. Spread the crystals in the bottom of a clean, dry 250 mL beaker and
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place in your closed drawer to air dry (note 3).
4. Using the 0.001 g balance weigh the clean vial with its lid. Record the mass of the bottle plus lid.
5. Before the next lab period, when the crystals are dry and no longer have the odor of acetone, place
them in the preweighed vial, screw on the lid, and weigh the bottle plus crystals to the nearest
0.001 g. (Record the mass in your notebook.) Calculate the mass of the crystals and record it in
your notebook. Wrap the vial containing your green solid in aluminum foil and store it in your
locked drawer for use in future experiments. A minimum of 3.5 grams of product will be needed
in subsequent experiments. If your yield is less than 3.5 grams, consult your instructor.
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Name: ____________________
Date: ________________
Data Sheet for Synthesis of an Iron Oxalato Complex Salt
Volume of FeCl3 solution used
Mass of FeCl3 used
Mass of potassium oxalate used
Mass of vial plus lid
Mass of vial plus lid plus crystals
Mass of crystals
Observations:
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