Materials Included In Kit

Additional Materials Required

Prelab Preparation

Oxalic acid solution, H₂C₂O₄, 10%: Add 100 g of oxalic acid to a 2-L beaker containing 900 mL of deionized water. Stir to dissolve. Transfer to a clean 1-L capped container and label as 10% oxalic acid.

Potassium oxalate solution, 25%: Add 75 g of potassium oxalate to a 400-mL beaker containing 225 mL of distilled water. Stir to dissolve. Transfer to a clean 500-mL capped container and label as 25% potassium oxalate.

Safety Precautions

The sulfuric acid solution is corrosive to eyes, skin and other tissue. Always add acid to water, never the reverse. Do not heat sulfuric acid. The oxalic acid solution is a skin and eye irritant and moderately toxic by ingestion. The 6% hydrogen peroxide solution is an oxidizer and a skin and eye irritant. The 95% ethyl alcohol solution and the acetone are both flammable and dangerous fire risks. Keep both away from open flames and other sources of ignition. The addition of denaturants makes the 95% ethyl alcohol solution and the 50% ethyl alcohol solution poisonous; they cannot be made nonpoisonous. The acetone is slightly toxic by ingestion and inhalation. Wear chemical splash goggles and chemical-resistant gloves and a chemical-resistant apron. Wash hands thoroughly with soap and water before leaving the laboratory. Please review current Safety Data Sheets for additional safety information.

Disposal

Please consult your current Flinn Scientific Catalog/Reference Manual for general guidelines and specific procedures, and review all federal, state and local regulations that may apply, before proceeding. The sulfuric acid solution and the oxalic acid solution may be disposed of according to Flinn Suggested Disposal Method #24b. The filtrate containing ethyl alcohol, acetone, and water may be disposed of according to Flinn Suggested Disposal Method #26b. The potassium trioxalateferrate(III) trihydrate product may be disposed of according to Flinn Suggested Disposal Method #26a, its solution according to Flinn Suggested Disposal Method #26b.

Lab Hints

• Caution the students to be sure to stir the hot mixtures containing solids to prevent “bumping.” Stirring should continue until the solution has cooled significantly. Also caution students not to use thermometers to stir the solutions.
• Review with students how to properly set up and use a Büchner funnel. Review proper transfer techniques.
• An alternative procedure for identifying the product involves the redox titration of a solution of the product with a standard solution of potassium permanganate.
• The absorption curve for a solution of K₃[Fe(C₂O₄)₃]•3H₂O is shown. 0.050 g of K₃[Fe(C₂O₄)₃]•3H₂O were dissolved in 100 mL of deionized water. The spectrum was taken every 10 nm from 340 nm to 800 nm using a Flinn Scientific Spectrophotometer and a standard ½"-diameter cuvet.
  {12653_Hints_Figure_2}
• If the absorbance value of the solution at 360 nm is greater than 1.5, have the student dilute the solution (1:1) until the absorbance value is <1.5.

Answers to Prelab Questions

1. Define the terms Lewis acid and Lewis base.

   A Lewis acid is a substance that can accept a pair of electrons from another atom to form a new bond.
   A Lewis base is a substance that contains an atom with a lone pair of electrons that can be donated to another atom to form a new bond.

2. Define the terms ligand and coordination number.

   A ligand is an atom, ion or molecule that acts as a Lewis base to bond to the central metal atom or ion in a complex ion. The coordination number of the central metal atom of a complex ion refers to the number of ligand atoms bonded to it.

3. What are the oxidation numbers of the metal atoms in each of the following coordination compounds?
   1. [Ni(NH₃)₆](NO₃)₂

      NO₃ has –1 charge; therefore the Ni(NH₃)₆ complex ion has a +2 charge. Since each NH₃ molecule is neutral, the Ni atom must have a +2 oxidation number.

   2. K₃[Co(CN)₆]

      K has a +1 charge; therefore the complex ion Co(CN)₆ has a –3 charge. Each CN ligand has a –1 charge, giving the Co atom an oxidation number of +3.

   3. [Pt(NH₃)3Br]Cl

      Cl has a –1 charge; therefore the complex ion Pt(NH₃)₃Br has a +1 charge. Each NH₃ ligand is neutral and the Br ligand has a –1 charge. The Pt atom has, therefore, a +2 oxidation number.

4. For each of the following ligands, draw the Lewis structures and indicate the atom that donates an electron pair for complex ion formation.
   {12653_Answers_Figure_3}
5. What is the coordination number of the metal in each of the following compounds?
   1. [FeCO(CN)₅](NO₃)₃

      The coordination number of the Fe atom equals the number of ligands in the complex ion. Coordination number of Fe = 6.

   2. [Ag(CN)₂]Cl

      Coordination number of Ag = 2.

   3. [Cr(H₂O)₂Cl₂]Br

      Coordination number of Cr = 4.

6. Suppose a student synthesizes potassium trioxalatoferrate(III) trihydrate, K₃[FeC₂O₄)₃]•3H₂O, by starting with 11.356 g of ferrous ammonium sulfate, Fe(NH₄)₂SO₄•6H₂O.
   1. What is the theoretical yield, in grams, for K₃[FeC₂O₄)₃]•3H₂O?

      Mass of Fe(NH₄)₂SO₄•6H₂O = 11.356 g

      {12653_PreLabAnswers_Equation_9}

      The theoretical moles of K₃[FeC₂O₄)₃]•3H₂O are the same, 0.02896 moles.
      The theoretical yield, in grams, is: Yield = 0.02896 moles x 491.16 g/mole = 14.22 g

   2. If 9.376 g of K₃[FeC₂O₄)₃]•3H₂O were actually synthesized, what is the percent yield?
      {12653_PreLabAnswers_Equation_10}

Sample Data

Data Table

1. Mass of Fe(NH₄)₂(SO₄)₂•6H₂O (g): ___9.879 g___
2. Mass of K₃[Fe(C₂O₄)₃]•3H₂O (g): ___9.347 g___

Answers to Questions

Results Table
Theoretical yield of K₃[Fe(C₂O₄)₃]•3H₂O (g): ___12.37 g___
Percent yield of K₃[Fe(C₂O₄)₃]•3H₂O (%): ___75.56%___

1. Calculate the theoretical yield of K₃[Fe(C₂O₄)₃]•3H₂O, based on the sample weight of Fe(NH₄)₂(SO₄)₂•6H₂O. Enter this value in the results table.

   Mass of Fe(NH₄)₂(SO₄)₂•6H₂O = 9.879 g

   {12653_Answers_Equation_11}

   The theoretical moles of K₃[Fe(C₂O₄)₃]•3H₂O are the same as Fe(NH₄)₂(SO₄)₂•6H₂O, or 0.02519 moles.
   Theoretical yield of K₃[Fe(C₂O₄)₃]•3H₂O:
   Yield = 0.02519 moles x 491.16 g/mole = 12.37 g

2. Calculate the percent yield for the K₃[Fe(C₂O₄)₃]•3H₂O product. Enter this value in the results table.
   {12653_Answers_Equation_12}
3. Calculate the absorbance ratios of the product solution. Calculate the 360/370 nm, the 370/380 nm, and the 360/380 nm absorbance ratios and enter these values in the results table. Was the product K₃[Fe(C₂O₄)₃]•3H₂O?
   {12653_Answers_Equation_13}
   Yes, the product was K₃[Fe(C₂O₄)₃]•3H₂O.

Introduction

Coordination compounds are interesting substances, usually highly colored and containing a complex ion in their structures. In this lab, green crystals of the coordination compound potassium trioxalatoferrate(III) trihydrate are synthesized. These crystals are then analyzed to determine their identity and the percent yield of the experiment.

Concepts

• Coordination compound
• Ligands
• Lewis acids and bases
• Coordination number
• Complex ions

Background

Coordination compounds are compounds that contain a metal atom or ion bonded to a group of molecules or ions. While these compounds may be neutral molecules, most are ionic compounds consisting of a complex ion (a metal ion with its attached molecules or ions) and a counter ion or ions to balance the charge. The molecules or ions attached to the central metal atom are called ligands.

FeCl₂•6H₂O is a typical coordination compound. The formula as written reflects the compound makeup, but not its structure. When writing the formula for coordination compounds, brackets are put around the metal and its ligands, indicating the actual structure and bonds in the compound. For FeCl₂•6H₂O the formula becomes [Fe(H₂O)₆]Cl₂.

The number of ligands atoms bonded to the metal ion is called the coordination number of the metal ion. This value is usually, but not always, 2, 4 or 6 and depends on the particular metal ion.

Ligands have lone pairs of electrons that can be used to form a bond with the metal ion. The ligand acts as a Lewis base and the metal ion acts as a Lewis acid.

The bonding in coordination compounds involves the overlap of the metal d-orbitals and the ligand lone pair orbitals. If two atoms in a ligand donate lone pair electrons to form separate single bonds with the metal ion, the ligand is said to be a bidentate ligand. As many as six atoms in an individual ligand can be bonded with the metal ion.

In Part 1, ferrous ammonium sulfate, Fe(NH₄)₂(SO₄)₂•6H₂O, is reacted with oxalic acid to form an intermediate compound, iron(II) oxalate dihydrate, FeC₂O₄•2H₂O. This compound is isolated and then converted in Part 2 to potassium trioxalatoferrate(III) trihydrate, K₃[Fe(C₂O₄)₃]•3H₂O, by oxidation with hydrogen peroxide in the presence of potassium oxalate and oxalic acid.

In Part 1, the reaction to form the intermediate compound is: In Part 2, the iron(II) oxalate dihydrate is first reacted with potassium oxalate, forming an orange complex of iron(II) and oxalates. It is this iron(II) compound that is oxidized by the hydrogen peroxide. Initially, the K₂[Fe(C₂O₄)₂]•2H₂O is oxidized to iron(III) hydroxide, Fe(OH)₃, a brown precipitate.

and

When more oxalic acid is added, the iron(III) hydroxide dissolves. The iron(III) ion in solution forms a complex ion with the oxalate ligands, yielding a clear green solution of the complex ion, Fe(C₂O₄)₃^3–. The potassium salt of this ion is soluble in water, but only slightly soluble in ethanol. The addition of ethanol forces the precipitation of the product salt.

Experiment Overview

The purpose of this experiment is to synthesize the coordination compound potassium trioxalatoferrate(III) trihydrate, K₃[Fe(C₂O₄)₃]•3H₂O. Once produced, the product will be identified by colorimetric methods and the percent yield of the product will be determined. Many coordination compounds absorb visible light. The absorbance of a solution of the synthesized product will be taken at three wavelengths. The ratios of these absorbances will then be used to identify the product as K₃[Fe(C₂O₄)₃]•3H₂O.

Materials

Prelab Questions

See Student PDF.

Safety Precautions

The sulfuric acid solution is corrosive to eyes, skin and other tissue. Always add acid to water, never the reverse. Do not heat sulfuric acid. The oxalic acid solution is a skin and eye irritant and moderately toxic by ingestion. The 6% hydrogen peroxide solution is an oxidizer and a skin and eye irritant. The 95% ethyl alcohol solution and the acetone are both flammable and dangerous fire risks. Keep both away from open flames and other sources of ignition. The addition of denaturates makes the 95% ethyl alcohol solution and the 50% ethyl alcohol solution poisonous; they cannot be made nonpoisonous. The acetone is slightly toxic by ingestion and inhalation. Wear chemical splash goggles, chemical-resistant gloves and a chemical-resistant apron. Wash hands thoroughly with soap and water before leaving the laboratory.

Procedure

Part 1. Preparation of Iron(II) Oxalate Intermediate

1. On an analytical balance, accurately mass between 10.0 and 10.5 g of ferrous ammonium sulfate, Fe(NH₄)₂(SO₄)₂•6H₂O. Record the exact mass in thedata table.
2. Obtain 30 mL of deionized water in a clean 50-mL graduated cylinder and transfer the water to a clean 150-mL beaker.
3. Add 10 g of ferrous ammonium sulfate to the 150-mL beaker and stir to dissolve.
4. Using a Beral-type pipet, transfer 6 drops (1 mL) of 2 M H₂SO₄ to the solution of ferrous ammonium sulfate. Stir to mix.
5. Obtain 50 mL of the 10% oxalic acid solution in a clean 50-mL graduated cylinder. Add this solution slowly, while stirring, to the 150-mL beaker. Yellow iron(II) oxalate now precipitates.
6. Place the 150-mL beaker on a hot plate and carefully heat the mixture to boiling. Be sure to continuously stir the contents to prevent bumping.
7. Turn off the hot plate. Carefully remove the 150-mL beaker using heat-resistant gloves and place the beaker on a ceramic fiber square to cool.
8. Continue stirring the beaker contents as they cool to prevent “bumping.”
9. Obtain a 400-mL beaker and decant the supernatant liquid in the 150-mL beaker into the 400-mL beaker. The supernatant liquid is the liquid above the precipitate.
10. Add approximately 100 mL of deionized water to a clean 250-mL Erlenmeyer flask. Place the Erlenmeyer flask on the hot plate and heat water to approximately 40 °C.
11. Wash the yellow precipitate remaining in the 150-mL beaker with 3 successive portions of approximately 30 mL of hot deionized water. Decant the liquid into the 400-mL beaker between each washing.

Part 2. Synthesis of K₃[Fe(C₂O₄)₃]•3H₂O

1. Obtain 18 mL of the 25% potassium oxalate solution in a clean 50-mL graduated cylinder. Add the 18 mL of 25% potassium oxalate solution to the 150-mL beaker containing the yellow iron(II) oxalate solid.
2. Heat the iron(II) oxalate/potassium oxalate mixture to 40 °C on a hot plate.
3. Add 17 mL of 6% hydrogen peroxide solution to a clean 50-mL beaker. Use a clean graduated cylinder to transfer.
4. Obtain 15 mL of the 10% oxalic acid solution in a clean 50-mL beaker. Transfer 8 mL of this solution to a clean 10-mL graduated cylinder.
5. Monitor the temperature of the iron(II) oxalate/potassium oxalate solution. When the mixture is around 40 °C, slowly andcarefully, while stirring, add the 17 mL of 6% hydrogen peroxide solution drop-wise using a Beral-type pipet. Continuously stir the solution while adding the 6% hydrogen peroxide. A reddish brown precipitate forms.
6. Once all the 6% hydrogen peroxide solution has been added to the oxalate solution, heat the solution in the beaker to its boiling point. Be sure to continuously stir the solution as it heats.
7. When the solution is boiling, quickly add the 8 mL of 10% oxalic acid solution. Add the remaining 7 mL of 10% oxalic acid in the 50-mL beaker drop-wise to the solution using a Beral-type pipet. The composition and color of the mixture will change from brick red precipitate to an olive or brown solution with suspended brown solid.
8. Turn off the hot plate. Remove the 150-mL beaker using heat-resistant gloves and place the beaker on a ceramic fiber square to cool. Continue stirring the beaker contents as they cool to prevent “bumping.”
9. Once the beaker contents have cooled, filter the brown solid by gravity using a filter flask. The filtrate should be a bright green solution due to the presence of dissolved K₃[Fe(C₂O₄)₃].
10. Obtain 20 mL of 95% ethyl alcohol in a clean 50-mL graduated cylinder. Add the 95% ethyl alcohol to the bright green solution of K₃[Fe(C₂O₄)₃]. Swirl to mix—this will produce a green precipitate of K₃[Fe(C₂O₄)₃]•3H₂O.
11. Cover the beaker with a watch glass and store the mixture as directed by the instructor.
12. Set up the filtration assembly as shown in Figure 1.
    {12653_Procedure_Figure_1}
13. Turn on the water to start the suction in the Büchner funnel. Add the quantitative filter paper to the Büchner funnel and wet the paper with deionized water from a wash bottle.
14. Transfer the contents of the 150-mL beaker to the Büchner funnel. Wash the beaker and green precipitate with 20 mL of 50% ethyl alcohol, then 20 mL of acetone. Transfer each washing to the Büchner funnel.
15. Allow the precipitate to air dry under suction for 5–10 minutes.
16. Disconnect the hosing from the aspirator, then turn off the faucet.
17. Carefully remove the filter paper and precipitate. On an analytical balance, tare a piece of weighing paper or a weigh boat.
18. Transfer the precipitate to the weighing paper or weight boat. Record the mass of the precipitate in the Data Table as mass of K₃[Fe(C₂O₄)₃]•3H₂O.

Part 3. Colorimetric Identification of K₃[Fe(C₂O₄)₃]•3H₂O

1. On a tared piece of weighing paper or dish, mass approximately 0.050 g of the green precipitate on an analytical balance.
2. Transfer the massed precipitate to a clean 250-mL beaker.
3. Obtain 100 mL of distilled water in a clean 100-mL graduated cylinder and add the distilled water to the 250-mL beaker. Stir to dissolve the green precipitate.
4. Follow the procedure for taking colorimetric measurements of the solution as directed by the instructor. Generally, spectrophotometers are used as follows: Turn the instrument on and allow it to warm up for 15 minutes. Set the wavelength at 360 nm. With no light passing through the instrument to the photo tube, set the percent transmittance to zero with the “zero” control. Handle cuvets at the top so no fingerprints are in the light path. Polish cuvets with a tissue. Place a cuvet that is about ⅔ full of distilled water into the sample holder and set the percent transmittance to 100% with the appropriate control (not the zero control). Fill a cuvet about ⅔ full of a test solution, place it in the spectrophotometer and read the absorbance. Consult the instrument manual for the details on its use.
5. Repeat step 4 for the wavelengths of 370 nm and 380 nm. Record all absorbances in the data table.

Student Worksheet PDF

12653_Student1.pdf