I decided to indulge in some coordination chemistry today! The oxalate ion is a strong ligand and bidentate to boot so it makes for some nice colorful salts when combined with various transition metals. I was reading about oxalic acid in a 19th century encyclopedia and came across a large section which was erroneously labeled as oxalate "double salts" (as opposed to "complex salts"). I spotted that oxalate ions coordinate with ferric ions to form emerald green salts. This sounded pretty cool and after some more research I had myself an experiment!
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the trioxalatoferrate(III) ion
Before I outline exactly what I did, I should note that in leisure I did not choose the most straightforward synthesis route. I did not particularly care about the end yield and so this allowed me to witness more chemical reactions.
Here is the specific path I took to get to Na3[Fe(C2O4)3]-H2O, a chemical with an unsurprisingly large number of synonyms which include sodium triethanedioatoferrate(III) monohydrate, sodium (tris)ethanedioatoferrate(III) monohydrate, sodium trioxalatoferrate(III) monohydrate, sodium (tris)oxalatoferrate(III) monohydrate, sodium ferric ethanedioate monohydrate, sodium iron(III) ethanedioate monohydrate, sodium iron(III) oxalate monohydrate, and sodium ferric oxalate monohydrate.
CaCl2 + FeSO4 --> CaSO4 + FeCl2
2NaHSO4 + NaOCl + NaCl --> 2Na2SO4 + H2O + Cl2
2FeCl2 + Cl2 --> 2FeCl3
FeCl3 + 3NaHCO3 --> 3NaCl + Fe(OH)3 + 3CO2
COOHCOOH + NaOH --> COOHCOONa + H2O
Fe(OH)3 + 3COOHCOONa --> Na3[Fe(C2O4)3] + 3H2O
There is little information available about the preparation of sodium trioxalatoferrate(III). Preparations of the analogous potassium salt, however, abound in literature. Most sources I read instruct to mix solutions of potassium bioxalate with ferric chloride. One preparation I read refluxed a mixture of potassium oxalate, barium oxalate, and ferric sulfate, taking advantage of the extremely low solubility of barium sulfate. Most probably, the much more common calcium oxalate cannot be substituted for the barium salt in this process because calcium sulfate is comparatively much more soluble.I decided to stray away from both of these preparations and for the heck of it, produce sodium trioxalatoferrate(III) by reacting dissolving ferric hydroxide in a solution of sodium bioxalate, which as I far as I can tell, has no problems in theory.
CaCl2 + FeSO4 --> CaSO4(s) + FeCl2
11.0g of calcium chloride and 20.0g of ferrous sulfate were each separately grounded up and mixed with enough water to completely dissolve them. The solutions were then mixed and a thick precipitate of calcium sulfate immediately formed. This was filtered and washed several times yielding a dilute yellow-green solution of ferrous chloride.
White precipitate of calcium sulfate
Filtered yellow-green solution of ferrous chloride
2NaHSO4 + NaOCl + NaCl --> 2Na2SO4 + H2O + Cl2
2FeCl2 + Cl2 --> 2FeCl3
The setup: bleach, sodium bisulfate, ferrous chloride, and sodium hydroxide from left to right
In this procedure, chlorine gas oxidizes the ferrous ion to the ferric ion. I made chlorine gas by slowly dripping 100mL of bleach from a separatory funnel into a solution containing 20g of sodium bisulfate. The chlorine generating flask was gently heated to limit chlorine’s solubility in water. It was then led into the solution of ferrous chloride and then subsequently led into a strong solution of sodium hydroxide to effectively neutralize excess chlorine gas.
The result: a deep red solution of ferric chloride!
FeCl3 + 3NaHCO3 --> 3NaCl + Fe(OH)3 + 3CO2
Next I added excess sodium bicarbonate to the solution of ferric chloride. The solution turned quickly to orange and then slowly to more of a brown, and obviously lots of frothing and foaming ensued. The stoichiometry of this product is most likely not very precise and is more aptly described as the berthollide Fe2O3-nH2O where n ranges between 2 and 3. Regardless, the important thing here is that iron is in the +3 oxidation state.
Filtering precipitated ferric hydroxide
Once filtered, I let the precipitate dry a while in the sun, but I did not worry about being able to dry it enough to accurately mass it. I was content with a ballpark estimate of 5g.
COOHCOOH + NaOH --> COOHCOONa + H2O
A solution of 5g sodium oxalate was slowly added to a 12.5g solution of oxalic acid to form a solution of sodium bioxalate. I am not sure what I was thinking here because I used rather valuable lye to neutralize the oxalic acid when I could have used the benevolent sodium bicarbonate. Oh well, hopefully I will eventually be manufacturing my own lye on a large scale from baking soda and slaked lime anyways.
"Ferric hydrate" and a solution of sodium bioxalate
The hot sodium bioxalate solution was poured on the wet ferric hydroxide and was stirred, resulting in a lime green solution.
A solution containing the trioxalatoferrate(III) ion
This 200mL solution was eventually boiled to about 40mL. I first cooled the solution to room temperature when it was at about 100mL and filtered out a white solid, which may have been excess oxalic acid and/or excess sodium bioxalate. At 40mL, I cooled the solution again and crystallized a light green solid of what is presumably sodium trioxalatoferrate(III) monohydrate. I did not bother to crystallize everything out and I was left with a medium-deep green solution.
Wet yield of sodium trioxalatoferrate(III) monohydrate