Chem 120 - Lab 9 Synthesis of Potassium trisoxolatoferrate (III) trihydrate
Labs 9 and 10 will be typed and illustrated with line drawings. This will be
your final lab report and it will count as a double lab. Part of the marks will be
for answering the theory questions and making a presentable job of your
report. Part of the marks will be on the yield and purity of your green crystal
product. Lab 9 is the synthesis. Lab 10 is the analysis of your complex and use
of it as an actinometer (light meter, photon counter).
1. Introduction lecture in Lab on: Chelation complexes – (See Brown and web
references.)
Ligands – the atom (molecule or ion) that donates the lone pair of electrons to
form a coordinate bond. Ligands are Lewis bases (electron pair donors) because of
this. Metals have unfilled d orbitals and can accept electron pairs and thus act as
Lewis acids (Electron pair acceptors).
Together they make coordination complexes or coordinate bonds.
Enantiomers - non-superposeable stereo isomers. Optical isomers. Right and left
handed versions of molecules. Often in biochemistry only one isomer is useful.
Synthesis of organometallic crystals. Metals in biochemical systems are complexed
by organic molecules like the Fe in hemoglobin used to bind and move Oxygen
gas (O2) or the Zn in carbonic anhydrase used to remove carbon dioxide (CO2)
from the blood or various chlorophylls & porphyrins. In geochemistry, original
bio-porphyrins become transmetallated and are selectively replaced by Ni or V
so these become progressively more dominant in kerogen, coal and petroleum
(geoporphyrins). The remainder preserve the transition metal abundances of the
ground water the plants drank for geological amounts of time.
Crystal structures are long range, highly symmetric arrangement of molecules or
ions which achieve a favoured low energy state by their packing geometry and
maximizing bonds or coulomb force interactions between the essential chemical
structure (stoichiometry). There are only 6 crystal systems ( rhombohedral is a
subclass of hexagonal) and 32 point groups or space groups for fundamental
symmetry. It usually takes several formula units to make a suitably symmetric
building block. Halite, common table salt or rock salt has a formula of NaCl but
because it crystallizes in a face-centered cubic lattice, it requires 4 formula units
to build the unit cell, hence the molecular formula is Na4Cl4 or (NaCl)4. These are
highly symmetric cubic crystals with 4/m 3 2/m as their symmetry elements. Due
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to the large number of elements and the asymmetric shapes of some of the
molecules, your compound has Monoclinic symmetry and like a great number of
chiral compounds, it belongs to Space Group P21/c. These kinds of crystals
contain only a single 2 fold rotational axis and can never contain a mirror plane
due to the chiral centre. Monoclinic unit cells have 3 unequal axes a<b<c and one
angle > 90° between a & c and the other 2 are right angles.
Do not use all overhead lights especially after step 6 as crystals & complex are
photosensitive!
Hazards: Hot water and strong sulphuric acid both burn skin. Oxalic acid is toxic
and will pass through your skin. It blocks Calcium uptake and will kill you
painfully in many biochemical ways. Wash off any spills immediately. Wear
goggles and do not get anything in your eyes.
Procedural steps:
1. Make ~200 mL of hot deionized water (~40° not boiling) in a 250 mL beaker
(for dissolving and rinsing complex at various stages). (hotplate)
2. Pump 25 mL of 0.8 M Oxalic Acid (e.g. 0.020 mol) into a clean, rinsed, 100 mL
beaker. This is to isolate Iron as a yellow ppt. Fe2+Ox2 , preserve it in reduced
form (ferrous ion). (n.b. More is not better here as excess oxalic acid is a
contaminant for your complex.)
3. Weigh 5.0xx g of Mohr’s salt (Ferrous ammonium sulphate hexahydrate,
392.139 f.w.) (e.g. 0.0130 mol of salt and of Ferrous iron) into a weighing boat.
Transfer this analytically to the beaker with the oxalic acid. Rinse tray as needed
to include all of the iron. Dissolve by stirring with a clean glass rod. The Mohr’s
salt provides Fe2+ for complex, and as K.F. Mohr found out in the 1800’s this salt
makes a ferrous complex which resists oxidation in air, especially in acidic
solutions. (n.b. If you forget to add the Mohr’s salt you’ll not have any green
colour for your crystals later on!
4. Dispense 1 mL of strong aqueous sulphuric acid from the dropper bottle in the
fume hood into a clean plastic graduated cylinder. Go back to your bench and
add this to acidify your oxalic acid solution and help to keep the pH low.
5. Measure out 17 mL of dionized water into a clean graduated cylinder and add
this to your oxalic acid solution. Solutions depend critically on ionic strength and
their structure depends not only on the ions and molecules present but on their
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dilution and temperature. Solubility is basically a very finicky structural thing
right down to the molecular level and how many waters can pack around the
solutes.
6. Gradually heat the Oxalic acid solution to boiling while stirring constantly
with a clean glass rod to avoid bumping (bursting steam bubbles which spit hot
solution around and lose your ingredients). Once this comes to boiling, remove it
from the hotplate and set it aside to allow the yellow precipitate to settle.
Heating dehydrates the solution and isolates all of the iron as ferrous oxalate.
7. Once your solution has settled into a clear superlatant liquid and a solid
yellow precipitate, carefully decant the clear liquid into a waste beaker. This
rinsing removes more oxalic acid and improves purity of precipitate. Do not lose
the yellow precipitate as this contains all of your iron needed to form the core of
the molecule for the crystals. Do not filter this solution as you will lose too much of
your iron oxalate solid to the filter paper.
8. Take 20 mL of your warm DI water and wash off your yellow precipitate by
suspending it and allowing it to resettle. For the second time, allow your yellow
precipitate to settle then decant off the rinse water into your waste beaker. This
rinses off any excess oxalic acid and improves purity of precipitate.
9. Weigh out 3.0xx g of potassium oxalate hydrate (K2C2O4 – H2O, 166.216 f.w.) to
provide 0.016 mol of salt and 2 x this of K+ ions to allow your complex to
crystallize during the final step of the synthesis. If you miss this step you’ll never
get a green crystal! Dissolve this in 10 mL of your hot DI water in a small beaker.
10. Add the potassium oxalate hydrate solution to your yellow precipitate in the
small beaker. Make a visual note of how much precipitate you have formed, it
should be about 1 teaspoonful. This gives you a reality check on how much iron
you have retained and a minimum of how much green crystal you can form.
Add a magnetic stir bar.
11. At the sink, rinse and clean a burette then charge it with ~15 mL of 0.88 mol/L
hydrogen peroxide, H2O2 (aq) , place it in a burette clamp on a ring stand to ready
it for the oxidation step (13).
12. Gently warm your beaker with the clear potassium oxalate solution and egg
yolk yellow ferrous oxalate precipitate on a hotplate on setting 3 to about 40°C
while stirring continuously with the magnetic stirring bar. This is not boiling but
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like a good hot tub temperature! This step is to prepare a well mixed uniform
suspension for a complete oxidation in the next step.
13. Slowly add while stirring 8.5 mL of the 0.88 M = 3.0% H2O2 from burette into
your hot suspension to oxidize your iron in the suspended particles from ferrous
to ferric keeping solution at 40°C throughout this step. Since they are solid tiny
crystals they only react through their surface area. Stirring and several minutes
cooking time are essential here.
14. Heat your beaker to boiling to re-dissolve the yellowish precipitate and
ensure complete oxidation to ferric iron. This step should completely dissolve
your yellow precipitate. Sometimes a brown precipitate of Fe (OH)3 remains and
is reluctant to dissolve. This will reduce your yield as not all of the ferric iron is
available to make the ferric oxalate molecular anion for the core of your green
crystals.
15. Drop the heat to just below the boiling point for your solution. Dispense 7 mL
of 0.8 Mol/L of oxalic acid into a small graduated cylinder. At nearly boiling
slowly add this to your mixture. This will provide enough oxalic acid (0.0056 mol)
to complete chelation of the ferric iron. Do not add any excess oxalic acid as this
will drop the purity of your product and encourage the solid solution of extra
oxalate and cause poor quality, small, impure crystals to form.
16. To your nearly boiling solution, while continually stirring, gradually add
another 1.5 mL of the H2O2 (aq) solution from your burette. This ensures the
solution is no longer too acidic and encourages the iron to remain oxidized as
ferric oxalate.
17. By now you should have a transparent, dark chrismas-tree green solution of
potassium ferric oxalate which you will need to form the green crystals. If so pat
yourself on the back for successfully forming the complex and a synthesis well
done. (If colour is not bright green, instructor to fix colour at this point . First try up to
10 mL 3.0% H2O2 otherwise add 1.0 mL of sulfuric acid.) If you have any remaining
brown precipitate (this should be way less than the yellow egg yolk you started
with in step 10), you must remove it by filtration at this stage using a flat filter
paper and a flat, perforated, ceramic Buchner funnel. Gravity filtration is rather
slow and will not recover as much green solution so aspiration using a sidearm
Buchner flask as per the Copper bromide decomposition in Lab 3. Ensure you
get tight fitting hoses to pull a vacuum. Also ensure that you can recover the
filtered liquid, uncontaminated as this will be used to form your crystals. Weigh
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a clean dry beaker and record its weight, then transfer in your green solution.
Record whether you needed to use filtration to purify your solution.
18. Obtain 10 mL of ethanol in a small clean (no residual chemicals from prior
steps) graduated cylinder. This alcohol solution is water soluble and it has a
greater affinity for water, you’re your green complex does. Place your beaker of
green liquid on the bench top to cool and loosely cover it with aluminum foil to
prevent light from entering. When your beaker of green liquid is cool to the
touch, gradually add 10 mL of ethanol of ethanol while swirling the beaker with
a flick of the wrist. This should dehydrate your potassium trisoxalato ferrate
complex and allow the solution to separate into 2 components, one with water
and alcohol and one with water and green complex. If your solution goes cloudy
or crystallizes too fast, heat it gently to redissolve the complex again and cover it
to let it slowly crystallize in a labeled beaker, covered entirely by aluminum foil,
in the dark.
Washing instructions:
In dark lab (especially if you will do lab 10 to analyze the compound and make
the actinometer), set up a sidearm buchner flask (250 ml), get a buchner funnel
from the cabinets under the counter and a rubber hose and set up an aspirator
like with the bromine vapour. Weigh dry and then put a flat filter paper in the
funnel. Scrape out your green crystals and slurry them out with a rubber
policeman and 50% ethanol. Rinse once with 50% Ethanol (or a small amount of
95%). Get a red acetone squeeze bottle from under the hoods and rinse down
while in the aspirating funnel to dry crystals. When not using the acetone bottle
crack open the cap so it doesn’t spontaneously siphon dry all over the counter.
Transfer filter paper and green crystals to a tared weighing boat and weigh final
product to compare to theoretical yield. You should have: dark green, large, flat
bladed monoclinic crystals.
19. Cover your beaker and place it in your top glassware drawer gently and leave
it there for 15 minutes to crystallize. This step is to permit enough time for
dehydration of the complex, nucleation and crystallization. Check after 15
minutes and ensure that green crystals have started to form. If no crystals have
formed and your solution remains green, ask to buy a green seed crystal from
your instructor for 5 points out of 100. Allow this to crystallize until you have
green crystal and clear superlatant solution.
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20. Once your crystals have completely precipitated leaving a clear solution, label
a filter paper with a pencil and your names (Justin Bieber and Hannah Montana)
and lab number (Chem 120-666). Weigh the dry filter paper and record its
weight. Transfer your crystals to the filter paper and rinse them off with enough
50% ethanol solution to recover all of your crystals and to further dehydrate the
crystals and rinse of any excess oxalic acid. Aspirate this dry with a sidearm
Buchner funnel and suction flask as above and rinse with a final 10 mL of
acetone to promote dehydration of your sample.
21. Weigh a piece of weighing paper or a clean sample boat. Carefully transfer
your dried crystals to it and calculate their weight and obtain your yield. You
will compare this to the theoretical yields you expect from your initial weights
and limiting reagents. You will have to calculate % yield and reflect on this and
explain why your yield is different in the lab report. Calculate a theoretical yield
for the experiment from the ingredients you used and the limiting reagent.
Compare this to your final product weight and discuss any discrepancy. If less
where were weights of materials lost. If more, what else could be in your green
crystals?
22. Save your green crystals in a glass vial, labeled with lab tape until the
analysis lab 10. Be sure to label them with your 2 names, the lab date and your
course and section number.
23. Wrap your vial in aluminum foil and keep your crystals them in the dark, as
they are photosensitive and this is part of their use in Lab 10 as a photon counter
detector.
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Lab 9 and Lab 10 write up (analysis next week) This counts 20 points (2 labs).
Provide a title page with a photograph of your green crystals (from your cell
phone) .
I. Procedure of Synthesis:
1) Give the stoichiometry of your ideal green crystals and explain where each
component came from in your synthesis.
2) Give a point form, step by step recipe for the synthesis with a 1 line
explanation for each step in the procedure.
II. Calculations, Questions/Problems and Discussion: Do these point form as
per list below:
1. For your green crystal formation, given by:
3K+ Fe3+ + 3 (C2O4)2- + 3H2O  K3[Fe(C2O4)3] – 3H2O
Calculate the total number of moles of:
a) K+ ions (from potassium oxalate, K2C2O4-H2O ) solution used
b) Oxalate ions (C2O4) from potassium oxalate and oxalic acid used and
c) Fe+3 ions available from Mohr’s salt ((NH4)2Fe2+(SO4)2·6H2O ) used.
2. Calculate the limiting reagent, from the stoichiometry of your complex:
(K3Fe(C2O4)3 – 3H2O ) and the available components from calculations 1a-c above.
3. Using your limiting reagent as determined in 2 above, calculate the theoretical
yield in grams of your green crystals.
4. Calculate your percentage yield from the weight of your dried green crystals
and the theoretical yield from part 3 above. Explain why your yield is high, low
or reasonable depending on what happened in your particular synthesis. (note:
Different groups may have different responses for this question.) If low, explain
why only part of the complex formed and where the rest of your reagents went
to. If high account for what the other stuff is that added to your crystal weight
but not to its theoretical stoichiometry by counting the oxalates you reacted with
the permanganate ions in your lab 10 titrations.
5. Explain why your neutral green crystals need the 3 positive potassium ions to
form a solid?
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6. Give a balanced chemical formula for potassium bis-oxalato ferrate (II)
dihydrate and explain why this form (with more reduced iron) should have this
particular stoichiometry?
7. Write a balance chemical reaction for the oxidation of Fe2+ ions to Fe3+ ions in
acidic aqueous solution using hydrogen peroxide (H2O2). For this step look up
the half cell reactions in a table of standard reduction potentials. Electrons lost by
the oxidized element, ion or molecule must equal the electrons gained by the
reductant.
8. Draw & label the green crystal complex showing where K,+, polar waters of
crystallization, Fe3+ ion & oxalate ions sit. Label each type of chemical bond that
holds this molecular ion and salt together: covalent, ligand, ionic, ion-dipole, etc.
III. Provide peer reviewed academic references for this synthesis, analysis and
related questions as per those below, to demonstrate where the science came
from for you to do this lab. See below:
References:
Dallinger, Richard, F. (1995). Synthesis and Characterization of Potassium
Tris(oxalato)ferrate(III) Trihydrate: A Spectrophotometric Method of Iron
Analysis. J. Chem. Educ., 72 (10), p. 936. DOI: 10.1021/ed072p936.
Junk, Peter, C. (2004). Supramolecular interactions in the X-ray crystal structure
of potassium tris(oxalato)ferrate(III) trihydrate. Journal of Coordination
Chemistry
Volume 58, Issue 4, p.355-361.
http://www.tandfonline.com/doi/abs/10.1080/00958970512331334250#.VFPkocrY
udw
Ladriere, J., (1992). Mössbauer study on the thermal decomposition of potassium
tris (oxalato) ferrate(III) trihydrate and bis (oxalato) ferrate(II) dehydrate.
Hyperfine Interactions, April 1992, Volume 70, Issue 1-4, pp. 1095-1098.
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