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Author: Subject: Potassium Trioxalatochromate III
Boffis
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[*] posted on 15-5-2019 at 12:58
Potassium Trioxalatochromate III


The preparation of the ferric iron analog of this compound has been widely discussed on this forum but not in detail has the chromium III complex been discussed before. I required some of this compound for testing organic bases that I am currently working on and decided to prepare some according to the instructions in the Handbook of Perparative Inorganic Chemistry by Brauer (SM library p1372). However, while I followed to instruction precisely all did not go well.

I had already done a few test tube scale reactions so that I had at least an idea of what the product should look like. The very dark coloured crystals are very soluble in water and at high dilution give a beautiful purplish violet colour.

The problem was after adding the prerequisite amount of oxalic acid and potassium oxalate the solution was still brown, very dark brown admittedly but not the violet colour I was expecting. A quick test with alcoholic diphenylcarbazide revealled the presence of much Cr VI was still present so I added more oxalic acid straight from the jar. I had to add about another 10g in order to generate a violet product and then about 3ml of 50% KOH solution to neutralize the solution and raise the potassium content.

I then sat down and began to examine the stoichiometry of the reaction and soon discovered that a great deal more oxalic acid was required that prescribed in Brauer. In the latter text it specifically states potassium oxalate monohydrate and oxalic acid dihydrate but it is now clear that these are in fact the weights of anhydrous material required and the reaction is roughly:
K2Cr2O7 + 2K2C2O4H2O + 7C2O4H2(H2O)2 -> 2K3Cr(C2O4)3(H2O)3 + 6CO2 + 17H2O

So my revised receipe is as follows:
Dissolve 15.0 grams of neutral potassium oxalate monohydrate and 36.5g of oxalic acid dihydrate in about 200ml of hot water (a stirrer hotplate helps). In a separate beaker dissolve 12g of potassium dichromate in 50ml of water with gentle heating. When the chromate salt has dissolved use a pipette to add this solution to the oxalate+oxalic acid solution by injecting the chromate solution as close to the bottom of the oxalic acid solution as possible; this helps ensure that all of the chromium VI has been reduced before it reaches the surface and risks being expelled as a fine mist from the foaming cause by carbon dioxide emmission. The reaction is fairly exothermic and the carbon dioxide generated causes vigorous frothing.

When all of the chromate solution has been added let the solution stand for a few minute until gas evolution ceases and then evaporate down to a small volume (about 45-50ml) in a shallow bowl on a steam bath or hotplate. Cover and allow to cool slowly and crystallize. A mass of very dark, lustrous crystals had formed after about 4 hours.

[Edited on 16-5-2019 by Boffis]
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[*] posted on 15-5-2019 at 23:17


It's not as risky as hexavalent chromium, but keep in mind that oxalic acid and oxalates also are (somewhat) toxic.

I myself made this compound several years ago. I weighed out a mix of finely crushed oxalic acid and finely crushed potassium dichromate in stoichiometric ratios. This can be done very accurately, because both oxalic acid dihydrate and potassium dichromate are not hygroscopic and the commercial products are very close (well within 99%) of the theoretical formula. I added 1% additional oxalic acid, just to be sure that there is no excess hexavalent chromium.

I mixed the chemicals well and then put them in a small beaker, which itself was immersed in cold water, on which a watch glass can be put to avoid loss of material. Next, I added a small amount of water to set off the reaction. Beware, the reaction is quite violent. Do this outside! A lot of carbon dioxide is produced and there may be small droplets of solution as well, which make it into the air. That's why I did this experiment outside. After the reaction you have a very dark purple or green liquid, which solidifies on standing. The precise perceived color depends on the type of light under which the compound is viewed.
Allow to dry in contact with dry air. The material is nearly pure as is and is absolutely free of hexavalent chromium (I tested that with acidified H2O2). I did not recrystallize it.

http://woelen.homescience.net/science/chem/compounds/pot_tri...



[Edited on 16-5-19 by woelen]




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[*] posted on 16-5-2019 at 03:35


The main risk with this reaction is the fine mist of droplets produced by the vigorous frothing. As Woelen says its best to do this outside if you don't have a fume cupboard. In Brauer he states that the oxalate solution is to be added to the dichromate solution but this causes the mist released to be full of CrVI throughout the addition. I found that the reaction works just as well if the dichromate is added to the oxalate solution and this way the mist only contains CrIII oxalates and oxalic acid, though still not healthy if inhaled.

There is not enough potassium in the dichromate alone (ratio Cr/K of 1/1) to satisfy the requirement for the final salt which has a Cr/K ration of 1/3, that's why I added the potassium hydroxide solution in my initial preparation. This isn't necessary if the correct amount of potassium oxalte is used in the first place.

K trioxalatochromateIII crystals.jpg - 43kB
Potassium trioxalatochromate III, the very dark crystals are rather hard to photograph.
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[*] posted on 16-5-2019 at 03:59


I tried the procedure written by boffis, but I used round bottom flask instead, to limit some of the mist getting away. I wonder if it's possible to crystallize it via room temp. evaporation, or is it gonna decompose over tíme?



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[*] posted on 16-5-2019 at 05:22


No its quite stable except perhaps over prolonged exposure to bright light. The Manganese III, iron III and cobalt III analogs are all light sensitive which is why I decided to use the chromium analog. I simply evaporated mine down on the hotplate, slowly, over about 1.5 hours in a shallow bowl, poured it into a small beaker and rinsed the bowl with 2ml of water. The photograph above is of these crystals with the liquid simply pour off. The crystals are glassy, almost black and practically opaque but are completely soluble again in a little water.

I am sure you can evaporate it at room temperature and grow really nice crystals. The solid doesn't seem to be hydroscopic or deliquescent so it should evapoarte to dryness in the end.

I have just repeated this experiment and found that to reduce all of the CrVI required slightly more than the theoretical amount of oxalic acid, 37.3g in total of oxalic acid.

[Edited on 16-5-2019 by Boffis]
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[*] posted on 6-12-2020 at 17:56


I had a go at making this. There used to be a site online called Thomas-something that listed many preparations and I think this is where I copied the procedure from. But that site seemed to have stopped working a while ago. What I copied said:

"To a solution of 23 g. of potassium oxalate monohydrate
and 55 g. of oxalic acid dihydrate in 800 ml. of water is
added 19 g. of powdered potassium dichromate in small
portions with vigorous stirring. When the reaction is
ended, the solution is evaporated nearly to dryness and
allowed to crystallize. Potassium .trioxalatochromiate
forms deep-green crystals with a brilliant blue iridescence"

I did the experiment on half scale:
- 11.5g potassium oxalate monohydrate and 27.5g oxalic acid (I assumed the dihydrate) added to 400ml water in a beaker, heat and stir to dissolve.
- Added 9.5g potassium dichromate in small portions with vigorous stirring.
- There was a very slow reaction - a gentle release of bubbles. I am not sure why in my case the reaction seemed much slower than described earlier in this thread, can it be because my solution was more dilute?
- The solution went from white to brown to eventually dark over 30 minutes or so, photos below.
1.jpg - 411kB 2.jpg - 431kB 3.jpg - 513kB

The gentle release of bubbles continued for hours. Eventually after 6 hours there was no more bubbles and that was also the end of the day.
- The next day the solution was steamed down. When it go to around 30-40ml crystals formed. The solution was then cooled and in the process became a wet solid.
- The black product was scooped out and dried under a steel dish in the sun. After 4 hours it was much drier; I ran out of patience by then and put it on the steam bath for 2 hours. The steam bath drying did not seem to change the color or appearance and there was no smell; i.e. no obvious decomposition.
- The crystals looked lovely, black and glittering. They formed some hard layers in the drying dish and had to be scratched out.
4.jpg - 752kB

Final recovery was 28.5g which is some 90% of the theoretical. Color is black with maybe a hit of green. It is easily soluble in water and a dilute solution appears violet when light is shined through it.
5.jpg - 304kB

From what I can see online most photos showed a similar black-ish color. Did anyone get this as descrived above "deep-green crystals with a brilliant blue iridescence"?

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[*] posted on 6-12-2020 at 23:33


What do you think succinic acid complex is possible? Is it worth doing?

[Edited on 7-12-2020 by vano.kavt]
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[*] posted on 7-12-2020 at 00:45


Hi vano.kavt I had to google succinic acid, never heard of it before. Probably worth a try!

I want to try to make copper gluconate; only reason being that I was given a lot of calcium gluconate....
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[*] posted on 7-12-2020 at 01:25


Quote: Originally posted by Lion850  
Hi vano.kavt I had to google succinic acid, never heard of it before. Probably worth a try!

I want to try to make copper gluconate; only reason being that I was given a lot of calcium gluconate....


I bought succinic acid on Amazon and it is very high quality. Definitely trying to look interesting. It is smaller than in the photo. Copper gluconate is very nice compound.
https://www.amazon.com/gp/product/B009JH2CME/ref=ppx_yo_dt_b...

[Edited on 7-12-2020 by vano.kavt]

received_766202377304170.jpeg - 183kB

[Edited on 7-12-2020 by vano.kavt]
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[*] posted on 7-12-2020 at 05:39


I did it. Potassium Trisuccinatochromate III. Its very nice and unusual yellowish-green compound. Its solluble in water.

received_299229761728508.jpeg - 209kB
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[*] posted on 7-12-2020 at 05:42


Did you make it by reducing Cr(VI)? If so, what is the reducing agent?

[Edited on 7-12-2020 by teodor]
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[*] posted on 7-12-2020 at 05:58


Quote: Originally posted by teodor  
Did you make it by reducing Cr(VI)? If so, what is the reducing agent?

[Edited on 7-12-2020 by teodor]


I did the same as the others. I used succinic acid instead of oxalic acid.
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[*] posted on 7-12-2020 at 06:07


Are the oxidation products of succinic acid by dichromate the same as oxalic acid, CO2 and water?
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[*] posted on 7-12-2020 at 06:29


Quote: Originally posted by teodor  
Are the oxidation products of succinic acid by dichromate the same as oxalic acid, CO2 and water?


I think it produced carbon dioxide. I heated solution and It was brown slightly black. I filtered it and evaporated any water. Maybe it is nonanhydrate.

[Edited on 7-12-2020 by vano.kavt]
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[*] posted on 7-12-2020 at 06:34


The malonic acid complex is also interesting. I have malonic acid in another house and I can not make it this time. I am very interested in how different colors they will have from each other.
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[*] posted on 7-12-2020 at 06:38


Vano.kavt: Very nice result. I really like your unusual ideas.

But I think that using succinic acid as reducing agent is waste of precious reagent. I think that mixing of solutions of metabisulfite(or ethanol)+succinic acid and acidified potassium dichromate is better. Metabisulfate or ethanol are cheap reducing reagents. So you need succinic acid just as complexing reagent. Be careful, reaction between dichromate and ethanol is quite exothermic.




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[*] posted on 7-12-2020 at 07:03


Quote: Originally posted by Bedlasky  
Vano.kavt: Very nice result. I really like your unusual ideas.

But I think that using succinic acid as reducing agent is waste of precious reagent. I think that mixing of solutions of metabisulfite(or ethanol)+succinic acid and acidified potassium dichromate is better. Metabisulfate or ethanol are cheap reducing reagents. So you need succinic acid just as complexing reagent. Be careful, reaction between dichromate and ethanol is quite exothermic.


Thank you so much Bedlasky. I was given Amazon gift cards so this and many other reagents are free for me. I can also buy succinic acid very cheaply in my city. That's why I'm not worried about this loss. Yes i know its quite exothermic. As for the reaction between chromium trioxide and ethanol is awful.
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[*] posted on 7-12-2020 at 07:52


Very interesting! I have sodium hydrogen succinate and had no real application for it. Now I know what to do with this compound! The color of this complex indeed is quite special for chromium(III). I also made the trisoxalato complex and its color strongly depends on the light under which it is observed. Under fluorescent light it looks grey, like your pictures, but under tungsten light it looks green.



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[*] posted on 7-12-2020 at 07:58


Quote: Originally posted by woelen  
Now I know what to do with this compound!


It would be nice if you could do the same and compare the photos. I am going to use tartaric acid next.
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[*] posted on 7-12-2020 at 08:09


Are there limitations which carboxylic acids could be used as a complexing agent? 2 examples before were about dicarboxylic acids and now you are talking about tartaric acid. Could straight-chain carboxylic acids be used for the same purpose then?
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[*] posted on 7-12-2020 at 08:19


Quote: Originally posted by teodor  
Are there limitations which carboxylic acids could be used as a complexing agent? 2 examples before were about dicarboxylic acids and now you are talking about tartaric acid. Could straight-chain carboxylic acids be used for the same purpose then?


I do not know I havent any information, but I will know when I try.
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[*] posted on 7-12-2020 at 08:23


I'd assume denticity comes into play here. Straight-chained, monocarboxylic acids will be poorer ligands than dicarboxylic acids, because you have less atoms on the molecule that can form a coordinate bond with the metal centre. Judging by the presence of additional hydroxyl groups, tartaric acid may be a better ligand than succinic acid.

I know this is the case with alcohols. You won't get a complex to form between Cu2+ and ethanol nor methanol. But ethylene glycol and glycerol on the other hand will form a coordination complex with Cu2+, thanks to additional hydroxyl groups on the molecule, which increase the denticity of these ligands.

[Edited on 7-12-2020 by EthidiumBromide]
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[*] posted on 7-12-2020 at 09:53


It would be interesting to combine this property with some organic synthesis where the desired product is isolated from the reaction by complexing with Cr(III) or other transitional metal.
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[*] posted on 7-12-2020 at 10:27


I did some experiments with Cu(II) complexes with polyhydroxy compounds. Polyhydroxy compounds also form complexes with boric acid, which increase its ionization - this is used in analytical chemistry in volumetric determination of boric acid, D-mannitol is often used as complexing agent for boric acid, but fructose or glycerol are also very good.



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[*] posted on 7-12-2020 at 10:36


I would expect malonic, maleic ans succinic acids to make similar complexes to the oxalates (I've had middling success with Cu(II) and these ligands, along with the Fe(III) malonate complex). With lactate, you should also get a chelate (it complexes copper nicely, but I was unable to isolate the complex), but it will probably give you stereoisomers (mer and fac), which may complicate the isolation.



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