Sciencemadness Discussion Board - Oxalate from formate?

Oxalate from formate?

Fulmen - 28-11-2018 at 07:49

I noticed something interesting on Wikipedia:
On heating, sodium formate decomposes to form sodium oxalate and hydrogen
Since I happen to have a few tons of the stuff at work I figured I could give it a try.

Some relevant literature:
https://www.researchgate.net/profile/Jerry_Kaczur/publicatio...

Apparently temperatures around 350-450°C should do the trick, either with straight formate or in combination with NaOH or Na2CO3. Doesn't sound to hard to pull off...

Fulmen - 29-11-2018 at 08:18

This looks very promising. Here are some excerpts from the paper:

Levi and Pita(3) ... stated that formate began to decompose at 300°C, and that the decomposition is violent at 400°C, with the reaction complete at 550°C. In later work, they also claimed that the use of small amounts of sodium hydroxide had a beneficial effect, producing technical grade sodium oxalate below the melting point of sodium formate.

Leslie(4) reported the experimental work in determining the factors in the conversion of sodium formate to sodium oxalate. ... at 2 in. of Hg (6.8 kPa) at 26°C, 14 in. of Hg (47.4 kPa) at 300°C, and at atmospheric pressure (101 kPa) at 360°C, the optimum NaOH concentration ranged from about 1 to 2 wt% to obtain nearly quantitative conversion to oxalate.

Górski (9): The formate yields to oxalate in the various atmospheres did not exceed 54% with no catalyst or with the addition of sodium carbonate for all three alkali metal formates. The addition of a base, such as NaOH to sodium formate, and KOH to potassium formate in a 1: 0.05 molar ratio achieved a 92 – 93% conversion to the corresponding alkali metal oxalate. He also reported that the yield of sodium oxalate was shown to decrease as the heating rate was lowered.

There are also some conflicting data:

Meisel’s(7) results showed a high variability with the shape and type of crucible holdersubstrates that were used, such as platinum, corundum, aluminum oxide, and others.Meisel(7) could not achieve the oxalate high conversion yields reported in the literature.

Fulmen - 30-11-2018 at 06:08

I'm recrystallizing some HCOONa now, it's a real PITA but at least I don't have to worry about yields.

The conversion seems simple enough, my plan is to add 2-3% NaOH, melt it on a hotplate to ensure a uniform mixture and then put it in the furnace at 400°C. My biggest worry is the H2 produced. I don't have any N2 or other inert gases, and CO2 seems to lower the yield. Any good ideas to prevent my furnace from blowing up?

Also, does anyone have suggestions for other uses for formates?
Calcium formate can apparently produce formaldehyde (https://library.sciencemadness.org/library/formaldehyde.html), and according to Goldschmidt tin formate decomposes almost quantitatively to formaldehyde.

[Edited on 30-11-18 by Fulmen]

Fulmen - 1-12-2018 at 06:28

I noticed a slight smell of formaldehyde when drying the formate (120°C), this could be an issue as I don't have good ventilation for my furnace. At least if I'm going to produce larger quantities. I'm having trouble finding much on this reaction, which I assume goes something like this:
2 HCOONa => H2CO + Na2CO3

I noticed that calcium formate decomposes at 300°C, while the conversion to oxalates requires temperatures around 400°C. Formaldehyde decomposes from around 150°C (http://www.inchem.org/documents/hsg/hsg/hsg057.htm). Wikipedia mentions the "Cannizzaro reaction in the presence of basic catalysts to produce formic acid and methanol", if I'm lucky the NaOH will suppress the formation of formaldehyde.

Fulmen - 2-12-2018 at 11:33

I did a preliminary test with 48,5g of NaOOCH and 1,5g NaOH, using a stainless beaker and a blow torch. After melting the melt started to produce flammable gas that quickly caught fire. Within a few minutes the melt foamed up and formed a grayish/tan solid mass, at this point the heating was discontinued. After cooling the mass was broken up and dissolved in appr. 350ml of 10% sulfuric acid with some slight effervescence. It was then filtered and left to cool, forming a white precipitate.

I will leave it in the fridge overnight to ppt out, then we'll see what we've got. I figured I would try to form iron oxalate, that should be fairly indicative.

Fulmen - 3-12-2018 at 09:44

I got a fair amount of precipitate (29,2g dried @80°C, 26,7g @120°C), the first to form was a fine white powder while larger, more transparent crystals formed over night. I worry that the initial ppt could be Na2SO4, guess I should have used HCl.

I tested the filtrate with FeSO4, but got no ppt. So instead I ppt it with calcium nitrate and treated it with conc. HCl to redissolve the oxalic acid. This gave me a slightly yellow solution (which I assumed were due to traces of iron). However, after heating it turned a strong orange color, don't really know what that means yet.

Fulmen - 4-12-2018 at 09:40

Yay! I redid the experiment, dissolving the product in 200ml of 15% HCl. Rather than trying to dissolve the whole batch in one go I leached it multiple times with the warm acid and precipitating it by cooling. The result was 27g of what I assume is the dihydrate. The iron test was positive (insoluble orange ppt that produced pyrophoric iron on heating).

Yield was lower than the literature suggested, but that was expected considering the crude setup. No big deal really as sodium formate is cheap as dirt.

Fulmen - 5-12-2018 at 13:30

Another batch, so far it seems like I got similar yields. I dropped the other batch on the floor, so no more results from that test. This time I hope to test the product by weight loss during dehydration and finally on decomposition.

I do get a significant (several grams by the looks of it) of insoluble matter, which has me perplexed.

The only products I can think of are:

Sodium oxalate.
Sodium formate. Unlikely as no smell was noted during acidification. Also it's by far the most soluble candidate.
Sodium carbonate. The most likely candidate (as indicated by effervescence), most things end up there if anything goes wrong.
Sodium oxide/hydroxide. Unlikely and should end up as sodium chloride together with any carbonates.
Sodium chloride (after acidification). This is guaranteed to be present, but should be soluble enough to remove.

Fulmen - 6-12-2018 at 22:39

Dang, seems like I still have the sodium salt. That fits better with the solubility, but how do I convert it to oxalic acid? I was sure acidification would do the trick.

Fulmen - 7-12-2018 at 12:15

I have tried to convert it again by dissolving it in a minimum of boiling water, then add conc. HCl until it dissolved. Sodium oxalate might be less soluble, but the Ka of HCl is 10³ greater than that of oxalic acid. Wouldn't that force oxalic acid out of solution?

Anybody? Helloooo? Am I talking to myself here?

AJKOER - 8-12-2018 at 20:12

Look at the reference paper, a bit dated in terms of science.

A possible low temperature aqueous approach would be to first create the carbon dioxide radical anion from the hydroxyl radical acting on the formate ion:

HCO2- (aq) + .OH --> H2O + .CO2- (see source below)

Then, speculation over the self reaction:

.CO2- + .CO2- =?= C2O4(2-)

Note, no air or oxygen presence as this is a known path to the superoxide radical anion:

O2 + .CO2- = .O2- + CO2

Source: See https://books.google.com/books?id=mckVFtJ7YecC&pg=PA14&a...
-------------------------------------------------

Paper: 'Reduction of carbon dioxide to oxalate by a binuclear copper complex' at https://www.ncbi.nlm.nih.gov/pubmed/25522935 .

See also: http://science.sciencemag.org/content/327/5963/313.full

[Edited on 9-12-2018 by AJKOER]

WGTR - 8-12-2018 at 20:59

Hmmm...well, you could try taking some of the sodium oxalate and precipitating the oxalate with some calcium chloride. Then, collect the precipitate and wash, and then with heavy stirring boil an excess amount with sulfuric acid solution. After some hours, filter out the excess calcium salts, and evaporate down what’s left. I’m not sure if that would even work, but that’s all I’ve got. Google Books will be your friend on something like this. Search for the free books and arrange them by date, then look for the ones from the late 1800’s or early 1900’s.

Edit: Just for fun, I took a cruise through Google Books. Perhaps boiling oxalic acid is a bad idea. Maybe this will help:

[Edited on 12-9-2018 by WGTR]

Fulmen - 8-12-2018 at 23:58

Thanks, that was useful. The calcium/sulfuric acid method is covered on page one, so that is definitively an option. But I was hoping to avoid the extra work, and a quick test with more HCl shows promise. Heating the product to an red/orange heat produced appr. 20% residue, suggesting a fair amount of oxalic acid present.

So my next step will be to try a larger run with HCl. The plan is to use a minimum of water (just enough to dissolve the NaCl) and enough HCl to dissolve everything while boiling.

unionised - 9-12-2018 at 01:34

Quote: Originally posted by AJKOER

A possible low temperature aqueous approach would be to first create the carbon dioxide radical anion from the hydroxyl radical acting on the formate ion:

I will just have a look on eBay and see if I can buy a pound of hydroxyl radicals.

Fulmen - 9-12-2018 at 02:06

LOL.

After some rough calculations and a little guesswork I started with 25g of sodium oxalate and 30ml of water, then added appr 75ml of 30% HCl while boiling (appr. 50% excess). A little bit of undissolved matter remained, so the liquid was decanted and left to cool. It seems to have worked, the precipitate was noticeable different from the previous attempts. The sodium oxalate formed a fine white powder while this time I got larger, somewhat transparent grains.
According to your reference oxalic acid is easily dehydrated, so I will air dry the product first. Subsequent dehydration and decomposition should indicate the purity.

[Edited on 9-12-18 by Fulmen]

unionised - 9-12-2018 at 02:13

Glad you appreciated the joke.
I was kidding- obviously- about buying OH radicals from eBay.
But, even if I had, they couldn't be used as a reagent to convert formate to oxalate.
They react with oxalate- so they would destroy it, rather than creating it.
https://link.springer.com/article/10.1007/s11172-008-0148-y

fusso - 9-12-2018 at 07:04

System alert: Congrats to Fulmen! You have unlocked the "talk to self" achievement!

[Edited on 181209 by fusso]

unionised - 9-12-2018 at 07:24

Quote: Originally posted by fusso

System alert: Congrats to Fulmen! You have unlocked the "talk to self" achievement!

[Edited on 181209 by fusso]

If he was talking to himself, how come you heard?

fusso - 9-12-2018 at 08:25

Quote: Originally posted by unionised

Quote: Originally posted by fusso

System alert: Congrats to Fulmen! You have unlocked the "talk to self" achievement!

[Edited on 181209 by fusso]

If he was talking to himself, how come you heard?

You can't get the point of this message?!

Fulmen - 10-12-2018 at 13:21

Fusso:

Drying at ambient takes it's sweet time. It feels dry, but it's still loosing weight ever so slowly. The dihydrate is supposedly pretty unstable, so I won't know for sure until I decompose it. But it looks like a yield of 30-40%. It's lower than hoped, I will have to dig further into the literature to see if I can improve it. Problem is it can't be done in my furnace, the hydrogen is just too flamable. It's not dangerous in an open vessel, but it can be a bit loud when it ignites. The sodium colors it intensely yellow, it's actually a pretty cool reaction. Soon after the mix melts it begins to bubble vigorously, at this point the reaction seems self-sustaining. The clear melt turns opaque and thickens before quickly rising up to fill the beaker as a white solid sponge.

I wonder if it could be used to produce sponge from other, perhaps refractory materials. If this sponge is added to a concentrated solution of CaCl it should convert into calcium oxalate which can then be decomposed into the carbonate.

Fulmen - 11-12-2018 at 12:52

The numbers are in, and they have me somewhat confused. Weight loss upon drying is consistent with the dehydration of oxalic acid dihydrate, but the decomposition produced 20% residue. If I assume the residue is Na2O then we''re talking 40% sodium oxalate. Could it be anything else?

unionised - 11-12-2018 at 13:16

It almost certainly isn't Na2O.
Redo the calculation for Na2CO3

Fulmen - 11-12-2018 at 13:43

Idunno, I heated it until it melted to a glass. There seemed to be some black insoluble matter in it, I guess strong heating can produce components that could char. I have had the same during the oxalate reaction. Some black residue coloring the batch black, insoluble in conc. HCl.

Fulmen - 15-12-2018 at 08:35

I ended up reprocessing everything via the calcium/sulfuric acid-route. It's a bit more wasteful, but it does produce a much purer product. The oxalic acid did darken a bit upon drying, guess I need another recrystallization to remove it.

unionised - 15-12-2018 at 08:58

Quote: Originally posted by Fulmen

Idunno, I heated it until it melted to a glass.

Sodium carbonate melts perfectly well...

Fulmen - 15-12-2018 at 10:56

I'm sure you're right. It was nevertheless too impure for my taste, so I didn't see any other choice than to blow some of my precious SA on it.

AJKOER - 16-12-2018 at 01:10

Quote: Originally posted by Fulmen

I noticed something interesting on Wikipedia:
On heating, sodium formate decomposes to form sodium oxalate and hydrogen
Since I happen to have a few tons of the stuff at work I figured I could give it a try.

Some relevant literature:
https://www.researchgate.net/profile/Jerry_Kaczur/publicatio...

Apparently temperatures around 350-450°C should do the trick, either with straight formate or in combination with NaOH or Na2CO3. Doesn't sound to hard to pull off...

Found a source (see 'Bonding Analysis of the [C2O4]2+ Intermediate Formed in the Reaction of CO22+ with Neutral CO2' by Ferran Feixas, et al., in J. Phys. Chem. A 2010, 114, 6681–6688 6681, link to full text: https://pdfs.semanticscholar.org/aee0/2cc27bee5b913b0a1584d6... ) that basically describes an unstable intermediate [C2O4]2+ that decomposes into CO2(2+) formed in a thermal reaction.

Bottom line, a questionable path to oxalate.

Fulmen - 16-12-2018 at 08:39

Whatever, I don't see it as germane to this topic. My initial focus was to investigate formates as a precursor to other useful chemicals. Oxalic acid was the first thing to come up, and last time I looked at oxalates I could only find the nitric acid/sucrose synthesis.

The OA synth is pretty much nailed, I don't think I will spend too much effort on improving the yield. So what else could we make? Formaldehyde sounds interesting, but it poses some additional challenges as the product is gaseous. Also I don't care much for the smell, so I'd like to find a synth that can use the formaldehyde directly. Urea-formaldehyde might be worth some study.

WGTR - 16-12-2018 at 19:24

Quote: Originally posted by Fulmen

So what else could we make?

Glyoxal. If you want a challenge. But maybe that was a rhetorical question.

In fact, oxalic acid is one oxidation product of glyoxal. If you found a way to reduce it back to glyoxal, that would be very cool, but undoubtedly very difficult and probably impractical.

Gyoxal is normally made by the careful nitric acid oxidation of ethanol/acetaldehyde. If made in the gas phase it can remain so. If, however, it is allowed to condense with a tiny bit of moisture it instantly polymerizes, and remains non-volatile up to its decomposition temperature.

Purification of glyoxal made by the nitric acid process is difficult, since the di-aldehyde is non-volatile. As the acid is expended the reaction stops. Upon gentle evaporation the dilute acid concentrates further, and the reaction starts again. I've seen the reaction run away with dense red fumes after it was seemingly completed.

The product contains numerous impurities: excess nitric acid, oxalic acid, glycolic acid, and glyoxylic acid are the main ones. Some of these side products are interesting in and of themselves. The reaction mix can be neutralized, but only with calcium carbonate. Other carbonates or hydroxides will perform a nasty little Cannizzaro reaction on the aldehyde.

I've read of ion-exchange membranes being used to remove the salts from glyoxal. The old literature from the early 1900's used barium and lead salts, if I remember correctly. That produces nasty waste products. I was thinking of using a form of capacitive deionization, but haven't gotten around to trying yet. I would have to make the electrodes for it.

What is glyoxal used for? Making N-methylimidazole via the Debus-Radziszewski imidazole synthesis. Some of the N-methylimidazolium chloride salts can be purified by melt-zone crystallization, using a hot wire looped around a slowly moving glass tube containing the salt. The purified salt is used to make N-methylimidazolium ionic liquids. These are used as room temperature molten salts for battery and supercapacitor chemistries. If you try to buy them, they are pretty expensive in small quantities ( I wonder why?).

Fulmen - 17-12-2018 at 01:24

Sounds interesting, but it's a little too ambitious for me.