Synthesis of Potassium Iron Oxalate
From ChemEd Collaborative
Synthesis of Potassium Iron Oxalate by Laura E. Slocum October 2008
BACKGROUND
This lab comes from several of my colleagues and is a lab that I did myself about 8 or 9 years ago in a summer workshop at a conference I attended. I will be copying most of my colleagues’ work so do NOT cite me as the author of this lab. Here is the author information to the best of my knowledge –
Professor Irwin J. Brink
The Synthesis of an Iron Oxalato Complex Salt
Hope College
Holland, Michigan
I do NOT have the year that it was published, but I know that Dr. Brink retired in 1996 and it was before that. However, there have been several adaptations of this procedure since that time, I know. So, we will use 1996.
In this experiment you will synthesize a compound that will contain the elements potassium, iron, carbon, hydrogen, and oxygen. Carbon and oxygen will be present in the compound as oxalate (C2O42-) whereas hydrogen and oxygen will be present as H2O. The final product may be given the formula KxFe(C2O4)y • z nH2O, where the z nH2O 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". For example: if we react HCl with excess NaOH one of the products will be NaCl. If we assume that all of the Cl- in HCl ends up as NaCl we know that each mole of HCl consumed should produce a mole of NaCl product. Suppose that in a particular reaction 10.0 grams of HCl (36.5 g/mol) react with excess NaOH and 15.4 g of NaCl (58.5 g/mol) are isolated from the reaction mixture by crystallization. What is the percent yield of NaCl in the experiment? If you calculated 96% you did the calculation correctly.
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.
MATERIALS and EQUIPMENT
| Balance | Iron(III) chloride solution, FeCl3(aq) |
| 250-mL Beakers | Potassium Oxalate Monohydrate, K2C2O4 • H2O |
| 250-mL Beaker | Acetone |
| Crucible Tongs | Distilled Water |
| Watch Glass | Ice |
| Buchner Funnel | |
| Filter Flask | |
| Brown bottle with cap |
SAFETY
Be sure to wear safety goggles as directed by your teacher. The solutions you will be working with in this experiment are safe if handled properly. If any of the solutions come in contact with your skin, wash with soap and water generously.
Iron compounds may stain your clothing permanently. Avoid over heating resulting in violent boiling when heating liquids.
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.
Discard all waste solutions in the designated container provided by your teacher.
PROCEDURE
Notes on Experimental Procedure:
- 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.
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.
- 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.
- 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.
Synthesis Procedure:
Obtain (in a clean 50 mL beaker) 8.00 mL of stock solution containing 0.667 g FeCl3/mL.
Weigh 12.0 to 12.5g K2C2O4 • H2O, into a clean dry 50 mL beaker using the 0.01 g balance. Add 20 mL of distilled water to dissolve the K2C2O4 • H2O. Heat on the hot plate stirring constantly until the K2C2O4 • H2O is completely dissolved.
Using the crucible tongs to handle the hot beaker, pour the hot solution into the beaker containing the iron(III) chloride solution and stir.
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. If crystals do NOT start to grow after 15 minutes consult your instructor.
After giving the crystals ample time to form, carefully pour off and discard the solvent without removing any crystals, a process called decantation.
Add 20 mL distilled water 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.
Cover the beaker with a watch glass and set in on the LABELED shelf in the chemical storage area until the next laboratory period to allow crystals to form. 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 brown bottle. Allow it to drain and air dry in your desk for a week, or dry it carefully with a paper towel.
Set up a vacuum filtration system
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.
Wash the crystals twice with ice cold, distilled water. Use less than 5 mL of ice cold, distilled 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).
Spread the crystals in the bottom of a clean, dry 250 mL beaker and place on the LABELED shelf in the chemical storage area to air dry (note 3).
Using the 0.01 g balance weigh to the nearest 0.01 g the clean dry brown bottle with its lid. Record the mass of the bottle plus lid.
When the crystals are dry and no longer have the odor of acetone, place them in the pre-weighed brown bottle, screw on the lid, and weigh the bottle plus crystals to the nearest 0.01 g on the balance. Calculate the mass of the crystals. Store them in your locked drawer in the capped bottle 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. Do not put the bottle of crystals in the desiccator.
DATA TABLES
Prepare an appropriate Data Table for this lab
RESULTS
Report the mass of crystals that you obtained. Also give a brief description of their appearance.
This picture is not to be handed out for students. It is just additional information for teachers.
POST-LAB QUESTIONS
- In this experiment you reacted an aqueous solution containing 0.667 grams FeCl3 • 6H2O/mL of
solution (Molar Mass of FeCl3 • 6H2O = 270.32 g/mol) with an aqueous solution containing excess
K2C2O4 • H2O to produce your product KxFe(C2O4)y • z nH2O
x K1-(aq) + y Fe3+(aq) + z C2O42-(aq) + nH2O → KxFey(C2O4)z • nH2O (s)
Assuming that all of the Fe originally in FeCl3 • 6H2O ends up in the product, KxFe(C2O4)y • z H2O,
how many moles of product should be obtained? What additional information would you need to
calculate a "percent yield" ?
REFERENCES
Brink, Irwin J., The Synthesis of an Iron Oxalato Complex Salt, NSF Summer Project in Chemistry, Hope College, 1996.
Holley, Kathleen, The Synthesis of a Complex Iron Salt, The Oakridge School, Arlington, TX, 2006.
Sweeney-Hammond, Kathy, E-mail conversations and Handouts, Maret School, Washington, DC, 2005-2008.
