Download Preparation of an Inorganic Complex Potassium

Chem 109
Bowdoin College Department of Chemistry
Spring 2003
Preparation of an Inorganic Complex
Potassium trisoxalatoferrate(III) trihydrate
PURPOSE
The purpose of this experiment is to synthesize a transition metal coordination compound and
determine the percent yield for the synthesis. This introduces an important class of compounds.
GENERAL BACKGROUND
Transition metal complexes are ions or compounds formed between transition metal ions(Lewis
acids) and a variety of organic and inorganic anions or neutral molecules, called ligands(Lewis
bases). Some neutral metal atoms also form such compounds. Examples of metal ions and
ligands are given below:
Transition metal ions
Cu2+, Fe2+, Fe3+, Ni2+
Mn2+, Cr3+, Co2+, Co3+
Ligands
Cl-, Br-, I-, SCN-, CO32F-, C2O42-, H2O, NH3
Such compounds are studied for, among other reasons, their roles in living systems, their potential
in solar power conversion applications and as anticancer drugs.
A well-known coordination ion is light blue hexaaquacopper(II), Cu(H2O)62+, in which six water
molecules (ligands) are arranged (coordinated) at the six corners of an octahedron about a central
Cu2+ ion (on the 6 arms of an x,y,z set of axes with Cu2+ at the origin). The nature of the bonds
between the central atom and the coordinating ligands varies with the nature of the species
involved. They generally involve unshared electron pairs on the ligand. In some coordination
compounds, the bonding may be described as largely covalent in character, whereas in others,
electron sharing is relatively unimportant and the interaction between the central ion and the ligands
is essentially ionic. The electrons are considered to "belong" to the ligand. For example, in
Cu(H2O)62+, the Cu atom is still Cu2+ and the H2O ligands are still neutral.
Some ligands can bond to the metal at more than one bonding site if the spacing of the unbonded
electron pairs and the geometry of the ion are correct. The ligand must be able to wrap itself
around the metal ion to bring more than one electron pair into orientation toward the metal, so that
the ligand is actually bonded two, three or even six times to the same central metal ion. These are
referred to as polydentate (many-toothed) ligands. Consider the structure and the unbonded
electron pairs of the ethylenediamine (H2N-CH2-CH2-NH2) molecule - how many bonding sites
do you think it would have? One hexadentate ligand, EDTA4- is used to treat systemic lead
poisoning and as a preservative to bind metal ions and prevent oxidation of foods. Oxalate,
C2O42-, the ligand used in this synthesis is the dianion of oxalic acid, H2C2O4, a toxic natural
compound found in the leaves of rhubarb and many other plants.
BACKGROUND SPECIFIC TO THE SYNTHESIS OF K3 Fe(C 2 O 4 )3 •3H 2 O
The compound you will synthesize has iron(III), Fe3+, as the central metal ion with three bidentate
oxalate ions, C2O4-2, as the ligands. The compound also contains potassium ions (to form the
neutral salt) and waters of hydration. The molecular formula is K3Fe(C2O4)3•3H2O. You will
synthesize the iron compound from the compound you synthesized earlier, iron(II) oxalate
dihydrate, FeC2O4•2H2O, plus oxalic acid (H2C2O4), hydrogen peroxide, (H2O2), and
potassium oxalate hydrate, (K2C2O4•H2O).
The overall reaction for the synthesis, omitting spectator ions is:
3 K+ + Fe2+ + 3 C2 O 4 2- + 3 H2 O + 1/2 H2 O 2 K 3 Fe(C 2 O 4 )3 •3H 2 O + OH-
Note that the reagents listed above do not appear as such on the reagent side. For example, all
the compounds that contain oxalate are combined into one oxalate reagent term. The
stoichiometric coefficient of 3 for oxalate means only that 3 moles of oxalate (from the combined
sources) are required for each Fe2+.
This reaction can be further broken down into the following subsidiary reactions.
The iron(II) oxalate dihydrate, Fe(C2O4)•2H2O is mixed with basic potassium oxalate solution
and an orange intermediate complex of Fe(II) and oxalate forms, K2Fe(C2O4)2•2H2O.
FeC2O4•2H2O + 2 K+ +C2O42-
K2Fe(C2O4)2•2H2O (orange solid)
K2Fe(C2O4)2•2H2O is oxidized to Fe(III) with hydrogen peroxide (H2O2), forming a brown
precipitate of iron(III) hydroxide (a base).
K2Fe(C2O4)2•2H2O + 1/2 H2O2 2 K+ + 2 C2O42- + Fe3+ + OH- and
Fe3+ + 3OH- Fe(OH)3 (brown solid)
Upon the addition of more oxalic acid, H2C2O4, the iron(III) hydroxide dissolves. The Fe(III) in
solution with oxalate undergoes complex ion formation, yielding Fe(C2O4)33-, which gives the
solution a clear, lime green color.
3 H+ + Fe(OH)3 3 H2O + Fe3+ and
Fe3+ + 3 C2O42-
Fe(C2O4)33-, (clear, lime green solution)
The ionic product, the K+ salt of the complex ion, is very water soluble but only slightly soluble in
ethanol. Addition of ethanol, CH3CH2OH, followed by cooling in an ice bath allows the product to
crystallize.
3 K+ + Fe(C2O4)33- + 3 H2O K3Fe(C2O4)3•3 H2O(s)
BEFORE YOU COME TO LAB
Review as necessary
Limiting reagents and % yield in pp. 87-95 in Chang and in Lab Guide (pp. 20 - 21).
If you are interested in the subject of transition metal complexes, see Chang, pp. 883-902 where
these complexes are described in greater detail, particularly why they are often colored (891-894).
Prelab Assignment
Complete the prelab assignment work sheet that is on the last page of this handout. Detach the
worksheet and pass in before the prelab discussion begins.
Prepare in your notebook.
Complete the Heading, Purpose, and Reference Sections. Read the procedure carefully and
note all the chemicals you will use. The calculated values for question 1 of the prelab must be
reproduced in your lab notebook (Calculations Section) for this experiment before coming to lab.
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PROCEDURE:
Write down what you do and see as you go. Be sure to record all observations and
measurements in your notebook. Include colors and whether reagents and solutions are cloudy or
clear. Careful recording of observations made during the synthesis allows confirmation that the
correct reactions occurred.
CAUTION: While heating solutions with precipitates, stir continuously to avoid
bumping (splattering).
1. Retrieve your sample of FeC2O4•2H2O from your drawer. Weigh the solid and compare to
the mass you obtained earlier. (Make a note if there is a significant difference in mass.) Put the
solid in a 100 mL beaker and wet it with a few drops of deionized water.
2. In a 30-mL beaker, dissolve 0.90 g (record to ± 0.0001 grams) of potassium oxalate hydrate,
K2C2O4•H2O, in 4 mL of warm deionized water and add to the yellow solid.
3. Warm to about 40 °C(check with a thermometer), remove from heat, and add 5.0 mL of 3%
hydrogen peroxide, H2O2, dropwise while stirring continuously.
4. Heat to boiling and add 3.0 mL of 1.0 M oxalic acid solution (H2C2O4) and the solution should
turn to a clear lime-green color. Be sure to record the observations you make and not simply
copy the previous sentence.
5. Add a few more drops of 1.0 M oxalic acid (H2C2O4) if the solution is not clear lime-green.
Was this necessary in your case? Indicate what you actually did in your notebook.
6. Cool the solution to ~ 50 °C and add 5 mL of 95% ethanol (CH3CH2OH).
7. Cover with a watch glass and cool in an ice-water bath (for 15 min. after reaching 5°C).
8. Vacuum filter the product and wash twice with 5 mL of 50-50 ethanol-water. Describe the
appearance of the product and the filtrate. Filtrate goes to the waste container in the hood not
down the drain.
9. Have the instructor check your product and sign your notebook. Leave it on a watch glass in
your locker to dry.
10. Return in two days or more, weigh* your product, record the mass in the procedure/data
section of your notebook along with the date you weighed it, place the solid in a shell vial, label
with your name and the formula of product and save in your locker. Turn in the labeled vial to
your instructor at the beginning of your next lab session.
*The lab will be open Monday, Tuesday, Wednesday, and Thursday afternoons from
1:00 - 3:30 p.m. Contact your instructor if you are unable to come during these times.
Important! If your final mass is greater than 1.2 grams, see your lab instructor before
completing your notebook entry for the post-experiment notebook sections.
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POST-EXPERIMENT NOTEBOOK SECTIONS
NOTE: This assignment covers part 3 of your notebook entry.
Analysis & Results Section:
State the actual yield, the theoretical yield(calculated in the worksheet on page 5) and percent
yield. Show calculation of percent yield.
Chemistry:(Hint: Review page 2 of this handout)
1. Include the overall balanced reaction for the synthesis.
2. Two intermediate solid iron compounds are formed during the synthesis. What are
they(formulas please) and what colors are they? What is the oxidation state of iron in each?
3. What two things did you do to the lime-green, dissolved product to cause it to crystallize?
4. What does the color of your green filtrate tell you about the completeness of precipitation of the
product? How does the completeness of precipitation affect your yield?
Conclusion:
Comment on your percent yield(Refer to page 16 of your Chem 109 Lab Guide for a frame of
reference). Find the compound in the CRC Handbook of Chemistry and Physics (you must use
the 75th or earlier edition of the CRC) in the Table of Physical Constants of Inorganic Compounds
(listed alphabetically under Potassium). Compare the color of your product(state the color you
used to describe your solid in your procedure/data section) to the color(state the exact wording
used by the CRC ) listed for this compound in the CRC. Provide a full citation for the CRC.
Worksheet:
Attach the LIMITING REAGENT / THEORETICAL YIELD worksheet to your yellow notebook
pages (page 5 of this handout).
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NAME _____________________
LIMITING REAGENT / THEORETICAL YIELD WORKSHEET
Attach this worksheet to your yellow notebook pages. Show all calculations with units &
correct significant figures for full credit.
The amount (in moles) of Fe2+ available is the moles of iron(II) oxalate dihydrate (use the mass
you weighed out).
Amount (in moles) of Fe2+______________
The K+ comes from K2C2O4•H2O (note that there are 2 K+ in the formula). From the mass used
(use the mass you weighed out), calculate the amount (in moles) of K+ available in Step Two.
Amount (in moles) of K+_____________
There are 3 sources of oxalate. Calculate the amount (in moles) of each and add to obtain the total
amount (in moles) of oxalate available.
Oxalate from iron(II) oxalate dihydrate (use the mass you weighed out).
More oxalate comes from the added K2C2O4•H2O (one C2O42- in the formula, use your
mass).
The third source is the oxalic acid solution (H2C2O4).
Total amount (in moles) of C2O42- ______________
Fill in the total moles of each reagent below and from the overall reaction stoichiometry, calculate the
number of moles of product possible from each. Circle the limiting reagent and calculate the
theoretical yield (in grams) using the molar mass of the product. Show your calculation of theoretical
yield.
moles used
moles product
theoretical
possible
yield
Fe2+
______________
______________
K+
______________
______________
C 2 O 4 2-
______________
______________
5
______________ g
Name_____________________________
CHEM 109
Inorganic Synthesis Prelab Assignment
Spring 2003
The values for question 1 must be reproduced in your lab notebook in the calculation section
before coming to lab. You may use your Molar Mass spreadsheet to calculate.
1. Calculate the molar masses of the solids used in the synthesis. Record in your notebook.
potassium oxalate hydrate, K2C2O4•H2O
Potassium trisoxalatoferrate(III) trihydrate (product), K3Fe(C2O4)3•3H2O
2. Suppose a student synthesizes potassium trisoxalatoferrate(III) trihydrate and uses the
following amounts (moles) of each of the reactants. Calculate the moles of product possible from
each reactant using the stoichiometry of the reaction, then state which species is the limiting reagent.
Calculate the theoretical yield of the product in grams. Show all calculations for full credit.
Fe2+
0.00572 moles
K+
0.0232 moles
C 2 O 4 2-
0.0306 moles
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