Sciencemadness Discussion Board

Oxalate complexes

sasan - 15-6-2014 at 01:14


Hi guys,I know that oxalate complexes had discussed in madness so much,but they are not complete and discussed in different topics with different dates,so I decided to open new topic to discuss about oxalate complexes

potassium trisoxalatoferrate(III):Fe2(SO4)3 + 3 BaC2O4 + 3 K2C2O4 → 2 K3[Fe(C2O4)3] + 3 BaSO4 (mixing in hot water and stiring and filtering the barium sulfate and drying the complex)

second method:2FeC2O4.2H2O + H2C2O4 + H2O2(aq) + 3K2C2O4 ==) 2K3[Fe(C2O4)3]+...... avaporating and drying the
product

potassium tirsoxalatochromate(III):K2Cr2O7 + 7H2C2O4 + K2C2O4==) K3[Cr(C2O4)3] + 6CO2 +H2O precipitating by alcohol and cooling and drying

potassium bisoxalatochromate(III):K2Cr2O7 + 7H2C2O4 + ??H2O ==)2K[Cr(C2O4)2]·4H2O+ ??CO2 drying the bluish violet solution

potassium bis(oxalato)cuprate(II)dihydrate:2K2C2O4 + CuSO4.5H2O ==)K[Cu(C2O4)2].2H2O + ......mixing stochiometric amounts of copper sulfate and potassium oxalate solutions at 60 C.precipitated by cooling,it has easiest pocedure in the oxalate stuff.
this is the picture of copper complex,very beautiful blue color

There is aluminium complex too,but it is boring and white to make that

Now my questions:is there other complex of oxalate with transition metals?and have you guys ever work on oxalate complexes?

I know there is titanium(III) complex too,but nothing on googling.

DSC_0244.jpg - 393kB DSC_0244.jpg - 393kB

[Edited on 15-6-2014 by sasan]

Brain&Force - 15-6-2014 at 11:10

Nice work sasan! Have you checked the tris(oxalato)ferrate complex for fluorescence?

Are there vanadium, manganese, nickel, and cobalt complexes as well?

unionised - 15-6-2014 at 11:13

Does this help?
http://scripts.iucr.org/cgi-bin/paper?S0567740877009819

blogfast25 - 15-6-2014 at 13:11

Quote: Originally posted by Brain&Force  
Nice work sasan! Have you checked the tris(oxalato)ferrate complex for fluorescence?



I made that complex and it fluoresces mildly in visible light. Didn't check in UV.

It's also a bit unusual to have a green ferric compound.

[Edited on 15-6-2014 by blogfast25]

numos - 15-6-2014 at 15:45

Ooo... something new to try, do you know if sodium oxalate could be used instead of potassium? Would Na3[Fe(C2O4)3] have the same properties?

Otherwise... I have sodium oxalate, but I have no Potassium bases to make potassium oxalate out of.

Also, this isn't just directed at you, but when you use H2O2 and other aqueous substances (eg. Ammonia) , can you give the conc.? 3%? 30%? 95%?

Justin Blaise - 15-6-2014 at 19:31

I once neutralized a small amount of CoCO3 with aqueous oxalic acid. I boiled the mixture until it stopped producing CO2 and I was left with a salmon colored precipitate of what I believed to be cobalt (II) oxalate. The color was not much different than that of the original carbonate, so it was a little disappointing.

numos - 15-6-2014 at 22:19

I tried the Iron and Copper complexes with Sodium oxalate, and it definitely does work, the copper looks identical to the picture.

sasan - 15-6-2014 at 22:35


B&F oxalatoferrates have fluorescence property,but not strong and beautiful like sodium flourscein and pyranine and they will decomposes under UV light:(

There is manganese oxalate complex too and it is more pretty(deep red maroon color)but the production is difficult and highly depends on the concentration and temprature.I go for this complex today but I dont think if it will work or not

I tried nickel complex too,but I didn't get a good result,maybe I used wrong amounts.I will retry this later

I don't know about the vanadium complex,but I think there is cobalt (III) complex exists(maybe treating cobalt salts with pottasium oxalate and hydrogen peroxide solution with stochiometric amounts.)

justin blaise use my procedure for cobalt complex,right now I don't have enough cobalt salt,make potassium oxalate by treating your oxalic acid with potassium carbonate(not potassium hydroxide)and make cobalt salt with your cobalt carbonate.

numos I have sodium oxalate too,but it is very less soluble comparing to potassium oxalate,and for making oxalate complexes you need concentrated solution and you can't make concentrated solution of oxalate salt with your sodium oxalate,treat your sodium oxalate with hydrochloric acid to obtain oxalic acid precipitate(I know oxalic acid is soluble in water,but I think it is insoluble in HCl(aq) and in sodium chloride solution),then combine your oxalic acid with potassium carbonate in right amounts to gain potassium oxalate solution and dry it to form K2C2O4.2H2O

about the concentration of the hydrogen peroxide,for every reaction I use excess hydrogen peroxide solution with concentration of ~30%

I hope I answered your questions

papaya - 16-6-2014 at 01:24

I must say that lately I prepared copper (II) oxalate (which must be regarded as a complex), though unwillingly, with a rather unexpected route and without using any oxalic acid.
Take copper metal, dissolve in 50% nitric acid - there must be enough excess of acid left in solution. Then add ethylene glycol and let it stay for a day. I tried this on a test tube scale - don't scale up, it may be not safe, since after HOURS an exothermic runaway with NOx evolution occurs after which light blue precipitate starts to form (supposedly copper oxalate dihydrate). After 24 hours I filtered that, washed with water and dried at RT. It's a beautiful light blue fine powder, the color seems very interesting under daylight - must be a little flourescence.
But the most striking thing to me was how it decomposes when heated on the foil - RED fumes are evolved with NO residue left at all - the fume consists of probable nano-sized copper/copper oxide particles(not NO2).. If the fumes hit the flame - it gets colored intense green! Just try!

[Edited on 16-6-2014 by papaya]

woelen - 16-6-2014 at 01:41

Some oxalates and oxalato complexes are explosive!

Copper(II) is borderline. It very easily loses CO2 and leaves behind a very fine powder of copper. Ferrous oxalate is somewhat more difficult to decompose, but it can be used to make pyrophoric iron powder. Silver oxalate explodes on heating, expelling CO2 with great violence and leaving fine silver powder behind.

papaya - 16-6-2014 at 01:47

Quote: Originally posted by woelen  
Some oxalates and oxalato complexes are explosive!

Copper(II) is borderline. It very easily loses CO2 and leaves behind a very fine powder of copper. Ferrous oxalate is somewhat more difficult to decompose, but it can be used to make pyrophoric iron powder. Silver oxalate explodes on heating, expelling CO2 with great violence and leaving fine silver powder behind.


Sure, but copper one, at least the dihydrate didn't show any energetic properties, even when heating confined in a foil it decomposed the same way.

woelen - 16-6-2014 at 02:08

As I said, copper(II) oxalate is borderline. It can decompose quickly, but not explosively. I can imagine that if it is really confined, that it can lead to an explosion, due to pressure buildup. Confining it in a little ball of Al-foil is not sufficient. The produced gas is not formed quickly enough to rip the foil apart, it quickly escapes through the folds in the foil.

bismuthate - 16-6-2014 at 02:49

I just made some and found that it decomposes almost instantly under flame. I've been wondering (and experimenting with to find out) if complexes would have energetic properties.
Would an ethylenediamine complex be possible? That might be explosive.

[Edited on 16-6-2014 by bismuthate]

PHILOU Zrealone - 16-6-2014 at 05:23

Quote: Originally posted by woelen  
Some oxalates and oxalato complexes are explosive!

Copper(II) is borderline. It very easily loses CO2 and leaves behind a very fine powder of copper. Ferrous oxalate is somewhat more difficult to decompose, but it can be used to make pyrophoric iron powder. Silver oxalate explodes on heating, expelling CO2 with great violence and leaving fine silver powder behind.

Mercury oxalate is even better for the explosive purpose...at least it crushes more sand in sand test than silver oxalate.

PHILOU Zrealone - 16-6-2014 at 05:38

Quote: Originally posted by bismuthate  
I just made some and found that it decomposes almost instantly under flame. I've been wondering (and experimenting with to find out) if complexes would have energetic properties.
Would an ethylenediamine complex be possible? That might be explosive.

[Edited on 16-6-2014 by bismuthate]

Oxalate is a weak explosophoric group. The driving force is the oxydoredox potential of:
-the oxydation of the oxalate dianion (reducer) into 2 gaseous CO2 molecules:
(-)O2C-CO2(-) --> 2 CO2(g) + 2e(-)
-the reduction of the oxydising metalic cation into metal
Cu(+) + 1e(-) --> Cu
Cu(2+) +2e(-) --> Cu
Ag(+) + 1e(-) --> Ag
Hg(+) + 1e(-) --> Hg
Hg(2+) + 2e(-) --> Hg

The heat of reaction (explosion) is not that much with comparison with other energetic materials...mostly because the carbon in oxalate is already almost fully oxydised.

Thus adding highly defficient OB complexants (like ethylene diamine) will temper the already poor explosive properties.

sasan - 16-6-2014 at 06:43


please don't go far from the topic,it is about the oxalate complexes not explosives or etc;)

papaya" potassium oxalato copprate is not fluorescent,I have UV tube I tried that under UV,but is has no fluorescence property,it is just very shiny and crystaline under lights,the ferrioxalate is somewhat fluorescent.

The cobalt complex is too unstable to make by amateur chemists,It is very photosensitive and very airsensitive(unlike cobalt (II) hydroxides that will oxidizes to (III) in presence of air:o)and reduced to cobalt (II) and destroys the potassium trisoxalatocobaltate(III)

The manganese complex is similar to cobalt,its complex exists just for Mn(3+),and in presence of air will reduced to Mn(2+)(Unlike manganese (II) hydroxide that will oxidize to Manganese(III)oxide:o) and destroys the potassium trisoxalatomanganate(III) compelx,but it is more easier to make manganese complex comparing too cobalt

any ideas?

Woelen,in your website I saw the potassium bis&trisoxalatochromates(III) pictures,how did you make that?I tried my procedure but I had a very low yield.



Praxichys - 16-6-2014 at 07:53

Anyone try tetraaminecopper II oxalate?

Maybe take the same route as the persulfate derivative.... Mixing an excess of ammonium oxalate (4.45g/100ml @20C) solution with aq. ammonia, add copper sulfate solution, hopefully precipitating the complex instead of just copper(II) oxalate (~2x10-10g/100ml @20C). I can't find solubility data on ammonium bioxalate so I do not know for sure if this is going to work. I'll try it this week if nobody else gets around to it.

bismuthate - 16-6-2014 at 11:43

Quote: Originally posted by Praxichys  
Anyone try tetraaminecopper II oxalate?

No but I plan to soon. I'll post the results when done.

Praxichys - 16-6-2014 at 15:21

I just tried tetraaminecopper II oxalate with no success so far.

1. I mixed up a CuSO4 solution with some oxalic acid solution and immediately got a milky, pale blue precipitate which I presume is copper oxalate. Adding 10% ammonia solution slowly darkened the solution until it was deep blue, exactly like tetraaminecopper II sulfate. Mixing with 50% MeOH resulted in a precipitation of some deep blue crystals. I scaled it up and I am letting it precipitate right now to investigate. My gut tells me that this is just tetraaminecopper II sulfate. It looks exactly like it.

2. Oxalic acid was dissolved with aqueous ammonia until the solution remained basic, indicating an excess of ammonia. A solution of copper II sulfate was added, giving the same result as solution 1. A MeOH precipitation was not tried as the solution was fairly dilute at this time and appeared identical to solution 1.

3. Some basic copper II carbonate was added as a powder to an oxalic acid solution, producing some bubbles and a pale teal precipitate, appearing that not all of the carbonate had reacted. Extra oxalic acid was added to be sure all copper carbonate was consumed. The color did not change. It is speculated that the carbonate was encapsulated by the formation of oxalate, or perhaps teal is the actual color of copper II oxalate. Addition of aqueous ammonia to the solution produced a medium-blue cloudy mixture which will need scaling up and filtering for examination.

thesmug - 16-6-2014 at 21:27

I think you are right about getting the sulfate ions back. You might try eliminating the sulfate before attempting to complex with ammonia. I'm unsure of how you might do this.

bismuthate - 17-6-2014 at 02:26

Barium nitrate?
Or you could just filter it and then try to dissolve it in ammonia.

Praxichys - 17-6-2014 at 04:34

This morning yielded a tiny but recoverable amount of deep blue precipitate which held its color on drying. The sample was heated with decomposition releasing ammonia first and remaining a pale blue powder reminiscent of anhydrous copper II sulfate. However, further heating to ca. 400°C yielded some tiny particles of what appeared to be copper metal, indicating that the precipitate might be some amino-complexed double salt of copper II sulfate and oxalate.

I think the most promising method might be to make up a pile of clean and dry copper oxalate dihydrate, then gas a thin water slurry of it with ammonia with vigorous stirring until everything is dissolved, followed by crashing out into MeOH or EtOH. This will take considerable experimentation to determine the correct solubility ratios. I'm going to a concert tonight so I will not be able to try this until tomorrow.

DraconicAcid - 17-6-2014 at 08:25

Quote: Originally posted by thesmug  
I think you are right about getting the sulfate ions back. You might try eliminating the sulfate before attempting to complex with ammonia. I'm unsure of how you might do this.

Just add enough oxalate to precipitate the copper, filter out the copper(II) oxalate, then dissolve in aqueous ammonia.

papaya - 17-6-2014 at 09:54

Quote: Originally posted by DraconicAcid  
Quote: Originally posted by thesmug  
I think you are right about getting the sulfate ions back. You might try eliminating the sulfate before attempting to complex with ammonia. I'm unsure of how you might do this.

Just add enough oxalate to precipitate the copper, filter out the copper(II) oxalate, then dissolve in aqueous ammonia.


What is the structure of the supposed complex, any data available?

Justin Blaise - 17-6-2014 at 19:44

Found some info on preparing oxalate complexes of some trivalent metals in Inorganic Syntheses Volume 1, page 35.

Although more facile methods to make the chromium and iron complexes exist, the preparation of the cobalt complex seems interesting. If I can procure some lead dioxide I'll be sure to give it a shot.

"C . POTASSIUM TRIOXALATOCOBALTIATE
K3[Co(C2O4)3].3H2O
H2C2O4 + CoCO3 -> CoC2O4 + H2O + CO2
2CoC2O4 + 4K2C2O4 + PbO2 + 4HC2H3O2 ->
2K3[Co(C2O4)3] + 2KC2H3O2 + Pb(C2H3O2)2 + 2H2O

Twenty-three and eight-tenths grams (0.2 mol) of cobalt
carbonate is dissolved in a solution of 25.2 g. of oxalic
acid (H2C204*2H20) and 73.7 g. of potassium oxalate
(K2C2O4-H20) in 500 cc. of hot water. When the solution
has cooled to 40°C., while it is vigorously stirred, 23.9 g.
of lead dioxide (see synthesis 16) is added slowly, followed
by 25 ml. of glacial acetic acid added a drop at a time.
The stirring is continued for an hour, during which time
the color changes from red to deep green. After the
unused lead dioxide is filtered out, the trioxalatocobaltiate is
precipitated by the addition of 500 ml. of alcohol. The
material appears as emerald-green needles, which are
sensitive to both light and heat; The yield is 70 g. (71 per
cent). "




Attachment: inorganic-synthesis volume 1.pdf (7.4MB)
This file has been downloaded 1237 times

sasan - 17-6-2014 at 21:25


Very useful infomations,you can use H2O2 inatead of PbO2

Isn't it air sensitive?amazing that you obtained the complex in presence of air and light:o

I go for cobalt salt and making the complex,now I don't have enough ccobalt salt

Justin Blaise - 17-6-2014 at 21:38

I don't think it's air sensitive, as the complex is formed in the presence of an oxidizing agent. The cobalt is already in the +3 oxidation state, so I don't imagine oxygen from the air will oxidize it to the +4 state under normal conditions.
The ferrioxalate complex is also light sensitive, but I guess these complexes are light stable enough to be isolated at least.

I've made other Co (III) complexes with H2O2, so I'll try the synthesis with H2O2 instead of PbO2 as per your suggestion. Hopefully I'll have some results to post about tomorrow

sasan - 17-6-2014 at 22:09


Justin Cobalt is already in +2 state not +3.for it hydroxides ,oxidizes to +3 and for this complex reduced to +2.I will use your procedure to make this.thank you very much

Justin Blaise - 18-6-2014 at 21:00

I meant the Co+3 in the complex is unlikely to undergo oxidation by oxygen in the air under normal conditions.

I attempted the procedure for K3[Co(C2O4)3]-3H2O on 1/10th the scale in the paper. I replaced the PbO2 with 20 ml of 17% H2O2. The reaction proceeded as outlined in the paper until I added 50 ml denatured ethanol, which did not cause the product to precipitate. I thought this might be due to the added water from the peroxide, so I added an additional 20 ml of alcohol. This did not cause precipitation of the product either, so I am chilling it in a freezer overnight.

The synthesis appears to proceed with lead (IV) oxide replaced with hydrogen peroxide, but it does make the isolation a little more challenging due to the additional solvent.

sasan - 18-6-2014 at 21:46


For trisoxalato chromate,I have the same problem,it did't precipitate,I dont know why,even in addition of alcohol and freezing
I think it would b better to use PbO2,your water of hydrogen peroxide cause the disolution of complex,.maybe MnO2 would work,if you dont have PbO2

Justin Blaise - 18-6-2014 at 22:51

Luckily, that same document I posted has a procedure for PbO2, so this will be my next endeavour.

Update on the current reaction; a light colored (maybe white, but it's hard to tell because of how deeply green the solution is) solid has precipitated. My guess is potassium oxalate, due to its limited solubility in alcohols. No sign of green crystals yet.

woelen - 19-6-2014 at 00:06

This link has a description of how to make oxalatochromate(III) complexes. I used these methods. They are based on reduction of dichromate to chromium(III) in the presence of oxalic acid.

http://wwwchem.uwimona.edu.jm/lab_manuals/c10expt15.html

rumunpl1995 - 3-7-2014 at 06:18

Hey guys;
Recently I have done a project all around oxalate complexes. If you were interested I could post pictures of Mn(III),Cr(III)(trisoxalate + cis-dioxalate + trans-dioxalate),Cu(II)(dihydrate),Co(III),Fe(III) and their absorption graphs (in water and in solid state). My team including our chemistry teacher put a lot of effort in this project. We`ve also managed to prepare a poster but it is in Polish (If ya were interested I can translate !!!!) .

When it comes to Co(III) complex I ve managed only to make it with PbO2 (Pb3O4 also worked but despite stechiometric amounts yield was pretty bad :( ). Peroxide worked slightly I could see a green tint for a few seconds but after a while all Co(III) was reduced to (II). I ve done this complex like twenty times since first synthesis. If you want to get high yield you have to do it as quick as plausiable (the best way is to chill solution in fridge for like 15-30 mins and then filtrate).

And by the way I am a big FAN of YOUR SITE WOELEN. THANKS for such a magnificant work !!!!!!!!! :)

Waiting for replies. :D

Justin Blaise - 3-7-2014 at 06:52

I'd be very interested in seeing your work. I actually just finished making lead (II) acetate, for conversion to PbO2, in pursuit of the Co (III) complex.

rumunpl1995 - 3-7-2014 at 12:55

@sasan
Co(III) complex is not air sensitive it decomposes water (to oxygen) (like Mn(III) ) while reducing to +II oxidation state but it is light sensitive and should be stored in amber vials.

@Justin Blaise
I`ll keep my fingers crossed for you. ;) (dont be suprised if your solution will be very intense dark green in colour almost black :P )

Pictures: (31 MB PDF file may take about 5 min to download )
<a href="http://speedy.sh/uusNF/Ox-pictures.pdf">Ox_pictures</a>

Stay tuned , more goodies are coming :)



[Edited on 4-7-2014 by rumunpl1995]

sasan - 4-7-2014 at 00:30

Maybe you are right rumunp.thank you for pixs.I will post pictures of potassium bisoxalato chromate and trisoxalatochromate I have made.

sasan - 4-7-2014 at 02:25



IMG_20140704_141553.jpg - 717kB IMG_20140704_141937.jpg - 810kB IMG_20140704_141430.jpg - 652kB IMG_20140704_141205.jpg - 803kB

sasan - 4-7-2014 at 02:31


Top left is potassium bisoxalatochromate(lll) it has a dark and very shiny green color comparing to the dull color of trisoxalatochromate

top right is potassium bisoxalatocuprate(ll) one of the most beautiful compounds I've made:)

bottom left is crystaline potassium trisoxalatochrmtae(lll) it is has dark dull green color

bottom right is pulverized potassium tisoxalatochromate(lll) green under sunlight blue under fluorescent lights

rumunpl1995 - 4-7-2014 at 08:09

Absorbtion spectrum (300-800nm) for trisoxalate compunds in solution:
<a href="http://speedy.sh/jj9vd/Ox-chart-solution.pdf">Ox_chart_solution</a>

Spectrum in solid state will be post soon :) (are you interesed in seeing more or less identical graph or should I use raw reflectance data instead?!?!)

@sasan
Make Co(ox)3 this compund crystalizes in amazing star alike structures :D

[Edited on 4-7-2014 by rumunpl1995]

Justin Blaise - 4-7-2014 at 12:15

@rumunpl1995 , those pictures are beautiful. Nice graph, as well. Thanks for sharing! I would be interested in any additional information you'd like to share; even if it is a nearly identical graph.

Would you be willing to share your procedure for the Mn complex?

rumunpl1995 - 4-7-2014 at 14:00

@Justin Blaise
K3[Mn(ox)3]*3H2O synthesis (our modified method) :
5 grams of H2C2O4*2H2O was dissolved in 35 ml of water in a beaker placed on a magnetic stirrer equipped with hot plate and warmed up to 65 C (temperature can not raise above 75 C ) 1 gram of KMnO4 was very carefully added to the reaction mixture (in very small portions , check your temperature frequently , you must keep your temps below 75 C). White precipitate appeared, solution during stirring changed color. After complete addition of KMnO4 reaction mixture was let to stir until it was fully mixed and completely clear. 1,1 grams of K2CO3 was added in portions , the heat was turned off, beaker with solution was placed in an ice bath and cooled to 20C - 30C after that 25 grams of crushed ice was further added (to the ice bath I think so :P) and when reaction mixture was 2C 0,24 grams of KMnO4 was added. Solution turned red and was passed through Büchner funnel with a sintered glass disc. Eluate was collected and treated with 35 ml of ice cooled ethanol and placed in freezer. After 45 minutes when no precipitate was observed excess of ice cold ethanol was added until the precipitate crystalize out. Precipitate was obtained by filtration through Büchner funnel with a sintered glass disc flushed earlier with a couple HCl drops. Precipitate was flushed with cold ethanol then cold acetone and was left to dry in a DARK place.
:cool:


[Edited on 4-7-2014 by rumunpl1995]

rumunpl1995 - 6-7-2014 at 01:08

Absorption spectrum (200 - 900nm) for oxalate complexes Mn(III),Co(III),Fe(III),Cr(III),Cu(II),Ni(II) in solid state:
<a href="http://speedy.sh/Fpqg2/Ox-chart-solid-state.pdf">Ox_chart_solid_state</a>

Note: Ni(II) complex (light blue powder) did not disolved in water (Cu(II) also "decomposed" after 2-5 minutes depending on concentration, Ni(II) complex didnt disolved at all), maybe thats only [Ni(H2O)4]C2O4 or [Ni(NH3)4]SO4 (adding ammonia was part of procedure) ??
(I really do not know why it didnt disolved at all and why copper complex decomposed on contact with water after couple minutes)
Any proposals??

That`s the last piece of data. ;)



Justin Blaise - 10-7-2014 at 07:43

Thanks for that procedure rumunpl1995! There appears to be a little more going on in the solid state spectra.

I followed your procedure and was able to achieve about 60% yield of the Mn complex on the left. On the right is the Cu complex. Still working on the Co (III) complex.



Additionally, I was heating the Co (II) complex in a test tube to see if it produced Co metal (like Fe (II) oxalate) or the oxide and observed some
thermochromism from it. It was pink at room temperature and turned purple when lightly heated with a flame. On cooling, it reverted back to pink.
I was able to repeat the color change 3 times before the compound decomposed into a black powder.


teodor - 8-1-2022 at 12:58

I've got a good yield of K3[Co(C2O4)3] by the method from "Inorganic syntheses" (0.5 scale). While the main batch is being dried in a dark place (also I used a red photo light at the last stages) I can show the part which I've got adhering to the funnel. I will use it to check the light sensitivity (and will compare it with the crystals from the main batch in a few days).
By the way, which method could you propose to measure/compare the light sensitivity?

CrIII_1.jpg - 215kB CrIII_3.jpg - 241kB

Comment on the photos: the color reproduction of my (phone) camera is incorrect here. The color is emerald-green/uniform. I believe the effect is due to high reflection ability of the crystals.

[Edited on 8-1-2022 by teodor]

It should be mentioned that the water solubility of the complex is very high, for this reason, I performed washing with ethanol/methanol.

[Edited on 8-1-2022 by teodor]

DraconicAcid - 8-1-2022 at 14:55

I have had some luck making similar complexes with malonates, but I found with the iron(III) complex, I had contamination with potassium malonate.

teodor - 11-1-2022 at 04:44

DraconicAcid interesting, I will try to find a way to try it with malonic acid.

What I know at the present moment:
1. the formula is K3[Co(C2O4)3] * 3.5 H2O. I didn't find any publication whith the structure investigation (yet). Some publications which I found say that the water is "probably coordinated". It would be interesting to know the actual structure, especially how water molecules are distributed.
2. after 2 days of keeping the Petry dish under direct sunlight there are visible things of decomposition on the crystal surface and edges (clearly visible under a microscope).
3. I didn't find yet any other solvent apart from water for this complex. Question 1 in this regard is interesting.
4. Observing the crystals under a microscope reveals 3 types of them - green, blue and bluish-green. Bluish-green is the most abundant color but it often has a green form as inclusion (like 2 sides of the same crystal can have a different color). Also, green and blue form forms separate crystals, also blue crystals are different in shape. I think the blue form is a result of impurities (Nickel?) and the green form is a complex isomer.

Now the most interesting questions for me are 1 (structure) and 3 (alternative solvent).

Update: Methanolic HCl works as a solvent giving the same color of the solution as water (green) but after some time it changes to Co(II) color (blue). When the water is added drop by drop, the color remains blue but at some point, it changes from blue to almost clear but slightly grayish-green in one step.
The reduction of Co(III) to Co(II) doesn't go at the same speed in water HCl as methanolic HCl (hours vs 1-3 minutes)


[Edited on 11-1-2022 by teodor]

[Edited on 11-1-2022 by teodor]

teodor - 12-1-2022 at 10:16

OK, summarizing my experiments with K3[Co(C2O4)3] * 3.5 H2O.

There are a lot of articles studying this compound. The best one which answers most of my questions is:
"Copestake T. B. and Uri N. 1955. The photochemistry of complex ions: photochemical and thermal decomposition of the trioxalatocobaltate III complex. Proc. R. Soc. Lond. A228252–263
http://doi.org/10.1098/rspa.1955.0047"

So, reducing Co(III) to Co(II) in methanolic HCl is not because of the action of Co(III) upon methanol but due to thermal decomposition of the complex which is speeded up by H+ concentration and presence of an organic substance. But in all the cases Co(III) acts as an oxidizer on oxalic acid only.
The mechanism by which the thermal decomposition rate is increasing with the presence of the organic substance is not clear. In my experiment, even diethyl ether (which doesn't dissolve the complex even in the presence of HCl and always forms a top clear level) causes almost instant reduction to Co(II).
The properties of "organic substances" which increase decomposition rate are unknown.

And the structure of the complex is really very complex. The anion is 2 Co(C2O4)3 centers bound together by some interaction of oxalic acid ligands. The position of water is unknown/unclear. In my experiments, mixing with acetyl chloride doesn't change the crystals.

[Edited on 12-1-2022 by teodor]

AJKOER - 13-1-2022 at 05:40

For the formation of the copper amine oxalate, I suggest trying the action of H2C2O4 on the base copper amine hydroxide. I have prepared copper amine sulfate from the action of this base on aqueous MgSO4. In the current context, just use H2C2O4 for copper amine oxalate and water.

The copper amine hydroxide can be created in a mixed standard and electrochemical reaction with copper, pumped in oxygen (or some starting H2O2 followed by aeration as I am not sure if residual H2O2 will decompose the eventually created oxalate salt) along with dilute aqueous ammonia, and for an electrolyte, a touch of NaCl.

General reference on the associated chemistry, see https://www.academia.edu/292096/Kinetics_and_Mechanism_of_Co...

Note, with a fall in pH associated with the consumption of ammonia, a problematic side reaction with the formation of some NH4NO2. The latter can decompose at less alkali pH to liberate suddenly large volumes of N2 gas. This means use an open large vessel and be mindful of a possible spillage event.

AJKOER - 13-1-2022 at 05:42

For the formation of the copper amine oxalate, I suggest trying the action of H2C2O4 on the base copper amine hydroxide. I have prepared copper amine sulfate (and nitrate) from the action of this base on aqueous MgSO4 (and Mg(NO3)2). In the current context, just use H2C2O4 for copper amine oxalate and water.

The copper amine hydroxide can be created in a mixed standard and electrochemical reaction with a small solid piece of copper metal, pumped in oxygen (or some starting H2O2 followed by aeration as I am not sure if residual H2O2 will decompose the eventually created oxalate salt) along with dilute aqueous ammonia, and for an electrolyte, a touch of NaCl. Note: this is a spontaneous electrochemical cell reaction, one does not have to attach any wires.

General reference on the associated chemistry, see https://www.academia.edu/292096/Kinetics_and_Mechanism_of_Co... which is employed in the commercial dissolution of copper metal ore.

Note, with a fall in pH associated with the consumption of ammonia, a problematic side reaction with the formation of some NH4NO2. The latter can decompose at less alkali pH to liberate suddenly large volumes of N2 gas. This means use an open large vessel and be mindful of a possible spillage event.

[Edited on 13-1-2022 by AJKOER]