Sciencemadness Discussion Board

Violuric acid salts (fantastic colors and variety)

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RustyShackleford - 16-12-2020 at 12:49

Screenshot 2020-12-16 214303.png - 70kB

Violuric acid can be prepared from barbituric acid (via NaNO2) or from alloxan (via hydroxylamine chloride). Barbituric acid can be prepared from condensation of diethyl malonate and urea, but i have not tried this, i simply purchased the barbituric (from S3 chemicals).

To make the violuric acid i followed the preparation on illumina from barbituric acid: https://illumina-chemie.de/viewtopic.php?t=5502

"
Procedure:
Sodium violurate:
6.40 g (50 mmol) of barbituric acid were dissolved in 100 ml of hot water in a beaker. A solution of 3.80 g (55 mmol) of sodium nitrite in 10 ml of water was added to the hot solution, and it immediately turned deep purple. A large amount of sodium violurate precipitated when it cooled down. To complete the reaction and precipitation, 10 g of sodium chloride (because of the addition of equal ions) as well as 1 g of sodium acetate and 3 ml of acetic acid were added to bring the pH to approx. 4-5 (the reaction is de facto complete in the pH range) and stirred for about 1-2 hours at room temperature. The pH was then brought back to the weakly basic range (lowest solubility) with 2.5 g of NaOH and allowed to crystallize in the refrigerator for a few hours.
The product was suction filtered and washed with a little cold water and air dried for a few days.
Yield: 9.85 g of sodium violurate dihydrate (84.5% of theory)


violuric acid, from sodium violurate:
7.8 g (33.5 mmol) of sodium violurate dihydrate were suspended in 20 ml of water and 10 ml of conc. Hydrochloric acid added. The deep violet suspension turned brown-pink in color. It was stirred well for about 1 hour at room temperature, suction filtered and washed thoroughly with dilute HCl, the color becoming a little lighter (cream-colored). The product was air dried for a few days.

Yield: 5.70 g of violuric acid monohydrate (88.2% of theory)
Note: if you touch the damp precipitate with a metal spatula, blue-violet spots will appear at the point of contact after a short time. Obviously, even the smallest traces of metal lead to a salt / complex formation."

I found that i got 86% yield on converting barbituric acid to violuric acid.


Most of the salts themselves were produced by simply heating the acid with the corresponding base in water and evaporating dry, but some (like barium and ethylene diamine) were prepared by adding salts of them to the solution of sodium violurate left behind in the preparation of the acid.

So far i have made: Na, K, Fe, Cu, Ba, ethylene diamine, hexamine, Sn, NH4, Mn, Zn and creatinine.

im planning to make: Sr, Ca, Pb, Al, aniline, benzylamine and quinoline

Let me know if there are particular salts youd want to see, so far it has made great colors with everything i have tried

Screenshot 2020-12-16 211809.png - 1.8MB

Screenshot 2020-12-16 211930.png - 1.4MB

Screenshot 2020-12-16 212001.png - 1.4MB

My favorites are these: Ethylene diamine, Mn and hexamine (bottom to top)
IMG_20201203_233709.jpg - 3.2MB
as well as the copper one.
IMG_20201203_233849.jpg - 2.1MB

The iron one isnt bad either, has a rich dark blue color in solution, but solid is so dark its almost black, not very fun

IMG_20201216_213811.jpg - 2.9MB

ShotBored - 16-12-2020 at 13:13

Those colors are awesome, dude. Only iron compound I've ever seen blue like that is Prussian blue. I'd be interested in seeing Pb, Sr, or guanidine salts! Or anything further on this subject really. These colors are astounding.

RustyShackleford - 16-12-2020 at 13:20

Quote: Originally posted by ShotBored  
Those colors are awesome, dude. Only iron compound I've ever seen blue like that is Prussian blue. I'd be interested in seeing Pb, Sr, or guanidine salts! Or anything further on this subject really. These colors are astounding.

Pb and Sr are in the works, guanidine i do not currently have, but maybe some day. maybe put a thing in your calendar to check back in a week or two, ill update the post as i produce new salts.

DraconicAcid - 16-12-2020 at 13:24

Gorgeous! I'm interested in seeing the nickel complex.

ShotBored - 16-12-2020 at 13:31

Quote: Originally posted by RustyShackleford  

Pb and Sr are in the works, guanidine i do not currently have, but maybe some day. maybe put a thing in your calendar to check back in a week or two, ill update the post as i produce new salts.


Derp for me, reading comprehension helps since you clearly state which ones you are working on above...oops.

Other compounds I would think would have interesting colors would include Cr and Mn...and As, though I usually stay away from As compounds myself. Hydrazinium might have vibrant colors too.

Finally, I did a lot of work with Nickel coordination chemistry in college. I always thought Nickel compounds had incredibly beautiful hues, mostly greens and deep reds are what I saw.

RustyShackleford - 16-12-2020 at 13:43

Quote: Originally posted by ShotBored  

Other compounds I would think would have interesting colors would include Cr and Mn...and As, though I usually stay away from As compounds myself. Hydrazinium might have vibrant colors too.


Mn i already did lol :D. Cr probably not hard, hydrazine might also be possible, but if i could make it from the sulfate that would be best, the freebase is kind of a pain to deal with and make pure.
Ni i would certainly do, just need to get some.
I basically want to make the violurate salt of every cation i have or can reasonably get

Bedlasky - 16-12-2020 at 14:33

Nice! Differences in colours of Na, K and NH4 salts are quite unusual. I would like to see Co2+ salt, cobalt salts in general have very nice colours.

Boffis - 16-12-2020 at 14:35

@Rustyshackleford, I did a load of work on these compounds years ago and posted a few microphotographs of the the crystals. Most M+ and M2+ metals form amazing coloured salts. They were the subject of much research back in the early 1900s because of the tendency of the colourless violuric acid to form strongly coloured slat with ions that normally form white salts. There are a few interesting exception Fe2+ forms a deep blue complex and cobalt and orange one. The shape of the crystals under the microscope is also highly characteristic. I used to use these reactions to identify alkali and alkaline earth metals in easily soluble minerals and Mn oxides.

The barium and Strontium salts are quite stunning under the microscope.
http://www.sciencemadness.org/talk/viewthread.php?tid=14644&...

I will have to dig out the other photos from a back-up disc. I even tried to encourage a thread dedicated to this type of microchemical test but to no avail.

I also starting preparing barbituric acid derivatives N-methyl, N,N'-dimethyl, thio- etc. and also related isonitroso-isoxazoles. Interesting to see that someone else is finally interested in these most unusual salts.

RustyShackleford - 16-12-2020 at 15:30

Quote: Originally posted by Boffis  
@Rustyshackleford, I did a load of work on these compounds years ago and posted a few microphotographs of the the crystals.


Thats very cool! with the amazing colors this acid produces im suprised its not more widely known about.
Do you have any tips for me in preparing these salts or the derivatives of the acid?

Bezaleel - 16-12-2020 at 15:50

Quote: Originally posted by Bedlasky  
Nice! Differences in colours of Na, K and NH4 salts are quite unusual. I would like to see Co2+ salt, cobalt salts in general have very nice colours.

Different colours for the alkali metals is special indeed!
@RustyShackleford What about Caesium?

Lion850 - 16-12-2020 at 17:30

Thanks for the post! I will be making some :)

Pumukli - 16-12-2020 at 21:13

Well, nice!

What about ammonium? Or silver? Cadmium? Mercury 2+?

woelen - 17-12-2020 at 00:56

This is really remarkable. I missed this thread up to now, but what most strikes me is the different colors for different alkali metals and ammonium. I do not know any other anion, which gives such striking different colors when e.g. Na(+) is exchanged with K(+).

It's also interesting to see Ba(2+) vs. Sr(2+) vs. Ca(2+). Usually, salts of these three anions also have more or less the same color.

I do have diethyl malonate. Up to now I did not have any real use for it, I obtained it as part of a package of old chemicals, years ago, and it still is in its unopened bottle. Now I have a use for that :D

Tsjerk - 17-12-2020 at 01:13

How about caesium as the cation? If you don't have any I could send you send you some CsOH for shipping costs. In what part of the world are you located?

[Edited on 17-12-2020 by Tsjerk]

RustyShackleford - 17-12-2020 at 01:55

Quote: Originally posted by Tsjerk  
How about caesium as the cation? If you don't have any I could send you send you some CsOH for shipping costs. In what part of the world are you located?
[Edited on 17-12-2020 by Tsjerk]

i sent you a U2U, hopefully we can work something out!

Heptylene - 17-12-2020 at 02:23

Incredible colors! I'm adding this compound to my todo list

Rubidium would be cool to see as well if you are going for cesium!

Also, if you have an interesting results with Al3+, maybe the rare earth would be interesting to investigate. They usually all react the same way and form pale-colored compounds, so a difference from this would be noteworthy.

Boffis - 17-12-2020 at 02:38

The colours you see can vary a lot even with a single cation depending on the hydrate and the ratio of acid to cation, pH, temperature etc. There are a whole series of papers on these salts by Hantzsch and others from 1909-1910. I hhave translated some of them but I can't find them just at the moment. Caesium forms a pure blue salt under warm neutral conditions. I spent a lot of time looking at organic buffers to try and find one that doesn't form a coloured salt with violuric acid. In the end the best was anion exchange resin beads washed with NaOH then acetic acid. Lithium is another interesting one, normally a bright red salt but then occasionally an insane fushia pink! Mn, Zn and some other divalent transition cations also react, but Cu2+, Fe2+ and Co react to produce soluble coloured complexes that are hard to crystallise. Interestingly UO2++ also reacts to form intensely yellow blades. I orginally examined this reagent as a microchemcial reagent for investigating secondary uranium minerals and Mn oxides both easily soluble in dilute acetic acid (the best medium for salt formation)

valeg96 - 17-12-2020 at 07:18

What about the lanthanides? I'm also curious about unusual metals such as U, Th and the likes.

RustyShackleford - 17-12-2020 at 07:41

Quote: Originally posted by valeg96  
What about the lanthanides? I'm also curious about unusual metals such as U, Th and the likes.

I myself do not have any pure lanthanides nor U or Th. I only have the mixture from dissolving lighter flints. If you have lanthanides or U/Th , i could send you some violuric acid, as long as you promise to post pics in this thread :P

itsallgoodjames - 17-12-2020 at 08:20

I wonder what violurate esters might look like. Also, metal complexes as the cation might be interesting. There's always so many possibilities with interesting ions.

valeg96 - 17-12-2020 at 08:51

Out of curiosity, can you post a full procedure for the preparation of the metal salts?

RustyShackleford - 17-12-2020 at 08:57

Quote: Originally posted by valeg96  
Out of curiosity, can you post a full procedure for the preparation of the metal salts?

i just measured stoichiometric metal base and violuric acid by mass, heated together in water solution and evaporated dry with heat. wasnt really anything advanced. I was doing 2-3 at a time.

valeg96 - 17-12-2020 at 11:53

Is that a proven procedure, backed up by literature? Where do the anions go? What do you mean with "metal base"?

DraconicAcid - 17-12-2020 at 11:56

Quote: Originally posted by valeg96  
Is that a proven procedure, backed up by literature? Where do the anions go? What do you mean with "metal base"?


A metal base would be a metal carbonate or a hydroxide. Do you need us to explain to you where the anions go when you react an acid with a metal hydroxide or carbonate?

valeg96 - 17-12-2020 at 12:05

Oh, well, it makes sense then. I'm just not used to call CuCO3 or Cu(OH)2 a "metal base", and I believe I've never heard anyone call transition metal hydroxides or carbonates "bases".

You're called DraconicAcid but you're caustic as hell today, bruh.

DraconicAcid - 17-12-2020 at 12:15

Quote: Originally posted by valeg96  
Oh, well, it makes sense then. I'm just not used to call CuCO3 or Cu(OH)2 a "metal base", and I believe I've never heard anyone call transition metal hydroxides or carbonates "bases".

You're called DraconicAcid but you're caustic as hell today, bruh.


Sorry- I've been spending the week trying very hard not to say what I really want to say to some of my students, and it's leaking out.

valeg96 - 17-12-2020 at 12:22

No worries. I was double asking because if the "base" route is the only one, I'm not going to attempt turning my salts into bases. If it was some kinda of "add salt, compound precipitates" thing I would've tried, but I don't want to fiddle around with them too much.

RustyShackleford - 17-12-2020 at 13:16

Quote: Originally posted by valeg96  
Oh, well, it makes sense then. I'm just not used to call CuCO3 or Cu(OH)2 a "metal base", and I believe I've never heard anyone call transition metal hydroxides or carbonates "bases".

Sorry for being unclear, i said metal base because i used a variety of different basic forms of the metal ions, like CO3, HCO3, OH

RustyShackleford - 17-12-2020 at 13:23

Today i produced some more salts.
Sadly had to borrow the calcium one from Diachrynic because my calcium source i attempted it with was impure.
Very nice to see such good variety in the 2 group

dasd.jpg - 2MB
I also produced some new amine ones, the color of the benzylamine salt is quite close to the Sr interestingly enough.
IMG_20201217_220634.jpg - 639kB

I also attempted to produce a urea salt, aswell as previously attempted a biuret salt. Neither of these worked, the violuric just crashes out alone, but in solution they are both a dark purple.

[Edited on 17-12-2020 by RustyShackleford]

IMG_20201217_215527.jpg - 3.1MB

[Edited on 17-12-2020 by RustyShackleford]

valeg96 - 17-12-2020 at 13:30

What about Co, Ni, W, Cd, Hg, Pb? I'd like to try them myself but I don't want barbituric acid in my lab lmao

Also, I'm wondering what these compounds look like structurally. They can't be all just simple ionic salts, some are probably coordination compounds. Also, are these compounds perfectly soluble, or do they crash out at some point?

[Edited on 17-12-2020 by valeg96]

RustyShackleford - 17-12-2020 at 13:42

Quote: Originally posted by valeg96  
What about Co, Ni, W, Cd, Hg, Pb? I'd like to try them myself but I don't want barbituric acid in my lab lmao

Also, I'm wondering what these compounds look like structurally. They can't be all just simple ionic salts, some are probably coordination compounds. Also, are these compounds perfectly soluble, or do they crash out at some point?

Co, Ni and Pb will definitely be made. i dont have any Cd though and not really comfortable making the Hg one. If you have either of those and would like some violuric to try i could send you a couple grams if you cover the shipping and promise to post pics in this thread.

Some of the salts are really soluble (like the iron) and some are quite non soluble, like barium. As for the coordination and structure, the big variety in colors likely stems from the resonance structures and that it has many ways of coordinating to metals. This is covered in the illumina article and im sure you could find a lot of old research if you know how to search for that stuff.

Tsjerk - 17-12-2020 at 13:59

Quote: Originally posted by valeg96  
I'd like to try them myself but I don't want barbituric acid in my lab lmao


Barbituric acid is very friendly toxicity wise, the toxicity is so benign the LD50 for rats is usually given to be > 5 gr/kg (read: couldn't be determined).

It is the derivatives you have to be careful with.

valeg96 - 17-12-2020 at 14:06

Toxicity does not concern me. I'm not sure if it's controlled, over here.

RustyShackleford - 17-12-2020 at 14:07

Quote: Originally posted by Tsjerk  

Barbituric acid is very friendly toxicity wise, the toxicity is so benign the LD50 for rats is usually given to be > 5 gr/kg (read: couldn't be determined).
It is the derivatives you have to be careful with.

Toxicity is very benign, and i dont think youd get in trouble over it (unless there are really stupid laws) because its not really feasible to produce barbiturate drugs with barbituric acid.

[Edited on 17-12-2020 by RustyShackleford]

valeg96 - 17-12-2020 at 14:26

Quote: Originally posted by Tsjerk  
]
Barbituric acid is very friendly toxicity wise, the toxicity is so benign the LD50 for rats is usually given to be > 5 gr/kg (read: couldn't be determined).

It is the derivatives you have to be careful with.


Read: "barbituric acid causes death by indigestion" :)

Boffis - 17-12-2020 at 14:42

Here are some photos of violurate salt crystals.

Ba violurate xacetate s2.jpg - 81kB Barium violurate note the two different forms

Sr violurate xacetate2.jpg - 77kB Strontium violurate, again note the two forms

Ca violurate series 2.jpg - 86kB Calcium violurate, brick red spherules and bow-ties but sometime orange rhombs form, the yellow blades are UO2++ violurate, the greenish colour cause by Fe2+ contamination of the mineral.

Na violurate xacetate 1.jpg - 78kB Sodium violurate forming at the edge of an evaporating drop

Its worth noting that it seems rather difficult to grow only one specific form so Rusty's compounds may not be pure hence the colours may vary. I am not sure what the difference is between the spherulitic or coralloid form and the euhedral crystals is. But I do remember from the Hantzsch paper that some form "acid" and normal salts while other sometimes form hydrates. I will try and find the Hantzsch paper that I translated years ago.
[Edited on 17-12-2020 by Boffis]

[Edited on 17-12-2020 by Boffis]

DraconicAcid - 17-12-2020 at 15:26

I was starting to think that this could be a really cool experiment to add to the curriculum for my organic class, but since making the barbituic acid to start with requires a 7 hour reflux, it probably won't happen this year. But I should be able to buy some over the summer.

woelen - 18-12-2020 at 01:14

Quote: Originally posted by woelen  
[...]I do have diethyl malonate. Up to now I did not have any real use for it, I obtained it as part of a package of old chemicals, years ago, and it still is in its unopened bottle. Now I have a use for that

After studying the synthesis of barbituric acid, I unfortunately come to another conclusion. I have found info in one of my books and also online sources. For making 100 grams of barbituric acid you need appr. 25 grams of sodium, 1 liter of absolute ethanol, 150 ml of diethylmalonate and refluxing for appr. one working day (7 hours, 8 hours, that kind of times). This is a few tens of euros of chemicals and a full day of babysitting a refluxing mix. I think that I also should buy it (appr. EUR 25 per 100 grams). Something for the beginning of next year.

DraconicAcid - 18-12-2020 at 02:03

If you have nothing else to do with diethyl malonate, you could try making acac-like complexes with it.

clearly_not_atara - 18-12-2020 at 19:05

Amazing work, had no idea this anion even existed, or that it was so colorful.

Quote: Originally posted by woelen  
I do not know any other anion, which gives such striking different colors when e.g. Na(+) is exchanged with K(+).

indeed, it seems Li+ would be a particularly interesting continuation


[Edited on 19-12-2020 by clearly_not_atara]

MidLifeChemist - 19-12-2020 at 09:54

>> Amazing work, had no idea this anion even existed, or that it was so colorful.
That is what I wanted to say.

This is amazing work. I had no idea this anion existed, or that it was so colorful. I'm sure I speak for others too.

IMHO this is a great example of what makes chemistry so interesting. There is so much for all of us to learn. Many of us are doing things that 99.9999938% of the world has never seen or done.

RustyShackleford - 19-12-2020 at 11:03

Quote: Originally posted by clearly_not_atara  
Amazing work, had no idea this anion even existed, or that it was so colorful.

Quote: Originally posted by woelen  
I do not know any other anion, which gives such striking different colors when e.g. Na(+) is exchanged with K(+).

indeed, it seems Li+ would be a particularly interesting continuation

Thank you! I found out about this acid from Diachrynic so some credit to him for inspiring me to do this.
Lithium will likely be made by end of next weekend, as i have mailed some violuric to a friend to make it with. He will probably attempt rubidium also and post results here.
Nickel is in the works by another user.

Justin Blaise - 20-12-2020 at 16:34

Really cool that these are nicely colored and easily isolated. Do you have any Ce, Nd, or Bi salts to try to make some complexes?

RustyShackleford - 20-12-2020 at 17:14

Quote: Originally posted by Justin Blaise  
Really cool that these are nicely colored and easily isolated. Do you have any Ce, Nd, or Bi salts to try to make some complexes?

Sadly none of those

DraconicAcid - 20-12-2020 at 17:50

Is the en salt with enH+ or enH2(2+)?

I imagine choline might also be interesting.

[Edited on 21-12-2020 by DraconicAcid]

Lion850 - 23-12-2020 at 00:47

Advice needed please. My source did not have barbituric acid but he did have diethyl barbituric acid. So I got some to try. I dissolved 5g diethyl barbituric acid in 100g hot water (with difficulty), and 3.5g sodium nitrite in 12g water. I added the sodium nitrite solution to the diethyl barbituric acid solution......and nothing. No purple color change; the solution stayed clear. I left it stirring.

I have zero knowledge of organic chemistry so help will be appreciated. The obvious question being can diethyl barbituric acid be a starting point to arrive at sodium violurate?

artemov - 23-12-2020 at 01:10

Quote: Originally posted by Lion850  
Advice needed please. My source did not have barbituric acid but he did have diethyl barbituric acid. So I got some to try. I dissolved 5g diethyl barbituric acid in 100g hot water (with difficulty), and 3.5g sodium nitrite in 12g water. I added the sodium nitrite solution to the diethyl barbituric acid solution......and nothing. No purple color change; the solution stayed clear. I left it stirring.

I have zero knowledge of organic chemistry so help will be appreciated. The obvious question being can diethyl barbituric acid be a starting point to arrive at sodium violurate?


I probably can't help you, but are you talking about 1,3-diethylbarbituric acid or 5,5-diethylbarbituric acid :o?

If it's the latter, I dun think the oxime will form at the 5th position cos the acidic alpha hydrogen is gone? My very limited knowledge ....

If it's the former, maybe 1,3-diethylvioluric acid will form? Just guessing.

Lion850 - 23-12-2020 at 01:25

Hi Artemov it only says Standard Diaethylbarbitursaure / Diethyl-barbituric acid and a formula (C2H5)2CCONHCONHCO. A Swiss product made by Fluka AG Buchs SG.


artemov - 23-12-2020 at 01:32

Quote: Originally posted by Lion850  
Hi Artemov it only says Standard Diaethylbarbitursaure / Diethyl-barbituric acid and a formula (C2H5)2CCONHCONHCO. A Swiss product made by Fluka AG Buchs SG.



Based on the formula, it seems that this is 5,5-diethylbarbituric acid aka barbital https://en.wikipedia.org/wiki/Barbital

This means the oxime cannot be formed at the 5th position to give violuric acid, unfortunately.

Do be careful with it cos it's a powerful sedative, and a controlled drug?

Lion850 - 23-12-2020 at 01:48

Artemov thanks for the feedback. It will be returned to the source, no use to me then.

Tsjerk - 23-12-2020 at 01:50

I don't think this would give colored complexes if 1,3-diethylvioluric acid was formed. The 1 and 3 positions are important in the different resonance structures, which are causing the different colors.

https://sci-hub.do/https://doi.org/10.1016/j.ccr.2014.01.002

The formula you posted does indicate it is 5,5-diethyl barbituric acid. Chemically there is not very much interesting to do with it.

[Edited on 23-12-2020 by Tsjerk]

artemov - 23-12-2020 at 02:12

Quote: Originally posted by Lion850  
Artemov thanks for the feedback. It will be returned to the source, no use to me then.



Erm ... you might want to wait for verification from more knowledgeable members, cos I might be wrong:P

RustyShackleford - 23-12-2020 at 10:08

Much requested cobalt violurate turned out to be quite bad, just brown.
6-methyl-quinoline turned out quite nice, light pink-red and acetone solution of it is a deep orange.

qqq.png - 4MB

Bedlasky - 23-12-2020 at 22:29

I think that cobalt salt is Co(III) violurate.

6-methyl-quinoline violurate look really nice.

RustyShackleford - 24-12-2020 at 02:22

Quote: Originally posted by Bedlasky  
I think that cobalt salt is Co(III) violurate.

6-methyl-quinoline violurate look really nice.

What makes you think its Co(III)? i used Co(II)CO3 to make it but it may well have oxidized i guess.

Bedlasky - 24-12-2020 at 03:07

Brown colour isn't typicall for Co(II) salts. If some cobalt salt is brown, it have usually some higer oxidation state.

Some cobalt complexes (especially wtih N donor ligands) have tendency to get oxidized by aaerial oxygen quite easily. For example [Co(NH3)6]2+ is slowly oxidized by oxygen to [Co(NH3)6]3+. Or Co(OH)2 is slowly converted by oxygen in to CoO(OH). [Co(CN)6)]4- is oxidized even by water to [Co(CN)6]3-.

[Edited on 24-12-2020 by Bedlasky]

RustyShackleford - 24-12-2020 at 03:57

Quote: Originally posted by Bedlasky  
Brown colour isn't typicall for Co(II) salts. If some cobalt salt is brown, it have usually some higer oxidation state.

Some cobalt complexes (especially wtih N donor ligands) have tendency to get oxidized by aaerial oxygen quite easily. For example [Co(NH3)6]2+ is slowly oxidized by oxygen to [Co(NH3)6]3+. Or Co(OH)2 is slowly converted by oxygen in to CoO(OH). [Co(CN)6)]4- is oxidized even by water to [Co(CN)6]3-.

[Edited on 24-12-2020 by Bedlasky]

Seems very possible, how would i go about preventing the oxidation? seems a bit difficult since i need to heat the cobalt with the acid in water and then evaporate it dry.

DraconicAcid - 24-12-2020 at 03:57

Quote: Originally posted by Bedlasky  
Brown colour isn't typicall for Co(II) salts. If some cobalt salt is brown, it have usually some higer oxidation state.

Some cobalt complexes (especially wtih N donor ligands) have tendency to get oxidized by aaerial oxygen quite easily. For example [Co(NH3)6]2+ is slowly oxidized by oxygen to [Co(NH3)6]3+. Or Co(OH)2 is slowly converted by oxygen in to CoO(OH). [Co(CN)6)]4- is oxidized even by water to [Co(CN)6]3-.

[Edited on 24-12-2020 by Bedlasky]


Yeah, but you need strongly-splitting ligands to do that, and there's only two violurates per cobalt centre. I have my doubts.

Isn't [Co(NH3)6]2+ a rather dingy brown-yellow?

Boffis - 24-12-2020 at 10:56

Some transition metals form true complexes, Co2+ doesn't seem to form a sparingly soluble salt like Zn and Mn but the solution slowly turns deep orange which, given that some oxidizing agents accelarate its formation, is likely to be a Co3+ compound. It is likely to be a triple ligand complex and since its formation seems to be insensetive to the other ions present and Co3+ tends to form 6-fold complex and I suspect violuric acid is a bidentate ligand. The complex is orange and very soluble.

Fe2+ complex is deep indigo blue and very soluble while the Cu2+ complex is olive green and less soluble. I have seen a paper comparing the complexes of Fe and Al with violuric acid, the latter forms a three fold octohedral colourless complex.

h0lx - 29-12-2020 at 08:49

Did some aminoguanidine violurate using the violuric acid RustyShackleford sent me.
UYeaRE2.jpg - 894kB
I reacted
136mg (1mmol) aminoguanidine bicarbonate
and
175mg (1mmol) Violuric acid
in some warm dH2O and it precipitated out.

TheMrbunGee - 30-12-2020 at 09:56

Hey, Do You have alloxan procedure, You mentioned? Colors look amazing, could try alloxan route, since I have sample with no use.

Boffis - 30-12-2020 at 13:39

Alloxan reacts with 1 molar equivalence of hydroxylamine to give violuric acid. In practice a hydroxylamine salt and sodium hydroxide are usually used.

Diachrynic - 5-1-2021 at 02:31

Quote: Originally posted by TheMrbunGee  
Hey, Do You have alloxan procedure, You mentioned? Colors look amazing, could try alloxan route, since I have sample with no use.

I would try the improved procedure by Guinchard from 1899:

violuric_from_alloxane.png - 34kB

Quote:
Violuric acid (or rather, pseudovioluric acid)
For the preparation of violuric acid I recommend the following, more convenient and higher yielding modification of the impractical method by Ceresole;
20 g alloxane (1 mol weight) and 10 g hydroxylamine hydrochloride (1.5 mol weight) [translators note: this is not a 1:1.5 molar ratio, it is more a 1:1.1 molar ratio] are dissolved in 150 g of water and heated for 1 hour on the water bath, after cooling one obtains 12.5 g of pure violuric acid directly in glittering, almost colorless crystals and another 2 g from carefully evaporating the mother liquor.



[Edited on 5-1-2021 by Diachrynic]

TheMrbunGee - 5-1-2021 at 14:29

Quote: Originally posted by Diachrynic  
Quote: Originally posted by TheMrbunGee  
Hey, Do You have alloxan procedure, You mentioned? Colors look amazing, could try alloxan route, since I have sample with no use.

I would try the improved procedure by Guinchard from 1899:



Quote:
Violuric acid (or rather, pseudovioluric acid)
For the preparation of violuric acid I recommend the following, more convenient and higher yielding modification of the impractical method by Ceresole;
20 g alloxane (1 mol weight) and 10 g hydroxylamine hydrochloride (1.5 mol weight) [translators note: this is not a 1:1.5 molar ratio, it is more a 1:1.1 molar ratio] are dissolved in 150 g of water and heated for 1 hour on the water bath, after cooling one obtains 12.5 g of pure violuric acid directly in glittering, almost colorless crystals and another 2 g from carefully evaporating the mother liquor.



[Edited on 5-1-2021 by Diachrynic]



This looks more like a procedure, thanks, I will try it when I find a moment.

RustyShackleford - 9-1-2021 at 08:24

Today i have produced some more salts of violuric acid thanks to SM user Tsjerk.
The ceasium salt, like the potassium, forms a blue hydrate, a slight bit darker than the potassium. On dehydration it turns to a nice teal, while the potassium turns purple.
The nickel salt is quite crummy, its a dirty light green, significantly worse than copper which in comparison was a wonderful moss green
IMG_20210109_171337.jpg - 1.3MB

valeg96 - 23-1-2021 at 10:54

This is the Rb salt. It slowly hydrates on air and turns blue like the K salt.

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Tsjerk - 23-1-2021 at 12:00

Nice! So now only lithium left for the alkali metals?

RustyShackleford - 23-1-2021 at 13:20

Very nice! thank you for making it! Could you post a picture of the hydrated form also?

valeg96 - 31-1-2021 at 03:31

Here it is.

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h0lx - 7-2-2021 at 15:23

Ethanolaminium Violurate

here it is compared to the aminoguanidinium one:



[Edited on 7-2-2021 by h0lx]

Housane - 8-2-2021 at 03:52

What amazing colours, I want to try and make sum but I can't find any sodium nitrite over here as it is a restricted chemical, could it be substuted for any easier to obtain chemicals here in the UK?

vano - 8-2-2021 at 11:28

Where did you buy this acid? Its salt's have very nice colours.

RustyShackleford - 8-2-2021 at 13:32

Quote: Originally posted by Housane  
What amazing colours, I want to try and make sum but I can't find any sodium nitrite over here as it is a restricted chemical, could it be substuted for any easier to obtain chemicals here in the UK?


Im afraid atleast from barbituric acid you will need nitrite, if you want another route you can go from diethylmalonate via sodium metal and hydroxylamine.

Quote: Originally posted by vano  
Where did you buy this acid? Its salt's have very nice colours.


I made it myself and have sent out a couple grams to several users on this forum. If you have an interesting new idea and are willing to pay a few bucks in shipping ill send you some too.

vano - 8-2-2021 at 14:24

I gladly tried the following metals. Scandium, cadmium, mercury 1 and 2, Gallium, indium, copper 1, samarium, silver, beryllium and others. Probably all is possible. Is it possible to send a parcel from your country to my country?

RustyShackleford - 8-2-2021 at 14:57

Quote: Originally posted by vano  
I gladly tried the following metals. Scandium, cadmium, mercury 1 and 2, Gallium, indium, copper 1, samarium, silver, beryllium and others. Probably all is possible. Is it possible to send a parcel from your country to my country?


Ill gladly send you some, i sent a PM

vano - 8-2-2021 at 15:12

Thank you

Housane - 9-2-2021 at 03:36

Ok thanks, where did you find your nitrite? As it is a explosives precursor here

RustyShackleford - 9-2-2021 at 03:53

Quote: Originally posted by Housane  
Ok thanks, where did you find your nitrite? As it is a explosives precursor here

i purchased a pound of it from ebay in 2019, it was from britain so idk if youre correct, or maybe the laws have changed since then for nitrite

artemov - 10-2-2021 at 02:07

Quote: Originally posted by RustyShackleford  
if you want another route you can go from diethylmalonate via sodium metal and hydroxylamine.


Can you elaborate more on this, or can you point me to a resource or something? I have malonic acid and hydroxylamine hcl ...

artemov - 10-2-2021 at 02:09

Quote: Originally posted by Housane  
What amazing colours, I want to try and make sum but I can't find any sodium nitrite over here as it is a restricted chemical, could it be substuted for any easier to obtain chemicals here in the UK?


Do you have access to poppers? It's an alkyl nitrite, but I'm not sure how pure are commercial poppers.

[Edited on 10-2-2021 by artemov]

h0lx - 28-2-2021 at 00:50

Pyridinium violurate. Strangely in solution it's much like permanganate but the crystals are orange
https://imgur.com/gallery/xrDb5el

Tsjerk - 28-2-2021 at 02:29

Nice!

Can you embed the photos in your post? Otherwise the are lost when the link breaks.

Amos - 28-2-2021 at 09:22

I don't believe I've seen anyone post an aniline salt yet, that would be neat. Piperidine too if anyone has it.

Thanks a bunch for bringing this to everyone's attention, Mr. Gribble

Boffis - 28-2-2021 at 09:51

Aniline may not be a strong enough base but piperidine should form a salt

RustyShackleford - 28-2-2021 at 11:46

Quote: Originally posted by Amos  
I don't believe I've seen anyone post an aniline salt yet, that would be neat. Piperidine too if anyone has it.

Thanks a bunch for bringing this to everyone's attention, Mr. Gribble

I did make the aniline salt but it was kinda ugly, possibly due to decomposition impurity in the aniline. Its on page 2 of this thread.
As for piperidine, i do have some 200mg of the HCl salt, but the freebase smells so horrible (like rotten cum) i dont want to do it. Perhaps when ive built a fumehood.

Bezaleel - 6-3-2021 at 17:06

I received some violuric acid from RustyShackleford yesterday in order to test what colours its rare earth salts have - thanks!. It's already after midnight now, but I could not resists the temptation to do one test.

Violuric acid is sparingly soluble in cold water, but much better in hot water. I noticed a somewhat sweetish smell coming off the warm solution (~60C).

0.291g of violuric acid was dissolved in ~15ml warm water, giving a clear purple solution (the dry salt is off white). This was added to a solution of 0.229g EuCl3.6H2O in a few ml water, yielding an orange solution!

IMG_3529_adj.jpg - 286kB

The solution is evaporating in a desiccator over NaOH. I'm charging up my UV lamp batteries to see whether the solution still fluoresces. EuCl3 solution (as well as the solid compound) fluoresces with a reddish colour.

RustyShackleford - 7-3-2021 at 09:58

Remade the aniline violurate with distilled clean amine, its a very dark red.
Anthranilic acid violurate is a light orange.
Im planning to make some substituted anilines and try with those (N- acetyl, N-benzyl, N-benzylglycine)

IMG_20210307_185022.jpg - 1.1MB

Side note: i have noticed that the color of the salts seem to "develop" over time. Like for example the Ni salt i made is now a more consistent and nicer shade of green than it was when i initially made it. No good idea of what the cause is.

IMG_20210307_191532.jpg - 363kB

[Edited on 7-3-2021 by RustyShackleford]

Bezaleel - 12-3-2021 at 14:38

Adding violuric acid to solutions of rare earth chlorides proved not a good idea. I had hoped to evaporate the HCl in solution over NaOH in a desiccator and obtain the rare earth violurates, but this did not happen. Instead, the Eu-violurate shown in my 7 March post lost its colour and a creme coloured solid was obtained.

So I continued to combine carbonates and hydroxides with violuric acid instead of chlorides. Sm2(CO3)3 gave a deep red solution which on drying gave red-brown granular crystals. I assumed the carbonate to be a 2.5 hydrate as posted here (bottom of page 8), but seeing I have some Sm-carbonate coloured residue, my carbonate may just have been dry.

Pr(OH)3 also gave a deep red solution. Pr(OH)3 was added in excess and intensely stirred for 15 minutes on a 70C hotplate. The solution was then filtered using a fine fritte. It's is now drying in a desiccator over NaOH.

Li2CO3 gave a strongly permanganate coloured solution, which on drying gives an intense almost fluorescent pink solid.

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Edit: fixed link

[Edited on 12-3-2021 by Bezaleel]

RustyShackleford - 12-3-2021 at 14:54

@Bezaleel
Wow that lithium salt looks amazing!
I have also noticed that its not really possible to make the violurate from strong acid salts of the metal, unless the metal violurate is quite insoluble. The technique i have been using for almost all of the salts i made is to make the carbonate or hydroxide of the metal, adding water + stoiciometric violuric acid, and evaporating that with the hotplate on 120C in a 50ml beaker. For ~0.2g violuric + corresponding metal carbonate/hydroxide ive been using 10-15ml of water. If you get a crummy looking product i have found that washing the salt w some acetone helps.

Bezaleel - 12-3-2021 at 15:30

So we ended up using the same method, apart from the use of a desiccator vs. a hotplate to evaporate the solution. I was not sure whether the violurate would stand temperatures over 100C, so I chose low pressure over temperature.

It seems like all the rare earths form red-brown compounds, so it may be more interesting to make a few non-rare earth ones instead. I can choose from indium, bismuth, antimony, vanadyl, maybe uranyl, rhenium, chromium, molybdenum, tungsten, and diethylamine. Of molybdenum, tungsten, and rhenium I don't expect much, but you never know, of course. (Other rare earth choices would be erbium, dysprosium, neodymium, holmium, lanthanum (~95%), and gadolinium.)

Please let me know your preferences.

Doe anyone know whether cyanuric acid could be used to make violuric acid from in a home lab?

RustyShackleford - 12-3-2021 at 16:04

Vacuum is better, i just dont have a pump good enough for that.
I would like to see all of those done, but if i had to pick a few i think uranyl>rhenium>indium.
A full rare earth series would be very cool, but ofc thats a lot of work.

Bezaleel - 16-3-2021 at 11:27

Updates

Samarium
I redissolved the Sm violurate and filtered off the undissolved carbonate, obtaining a warm clear solution. As it cooled down, a beige film formed on the flask. The solution was poured off and dried in a desiccator over NaOH, forming a deep red solid.

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Praseodymium
A new sample was made using Pr(OH)3. The hydroxide was stirred powerfully with demiwater, filtered again and then used. 0.300g of violuric acid was added to an excess of Pr(OH)3 and the undissolved hydroxide filtered off. When left to crystallise over NaOH, an orange-yellow film formed on the walls of the beaker and after everything had dried up completely, a mix of dark red and yellow-orange crystals had formed. As far as possible, the dark crystals were separated and the rest was redissolved - the yellow crystals did not dissolve on heating, but the red ones did. The solution is now crystallising again.

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A comparison between the Sm and Pr salts is in the last picture. It looks like the colour of the compounds is mainly caused by interactions inside the violuric acid ligand and not so much by the metal-ligand interactions. In particular the weak green of the Pr forbidden transitions seems to be retained, and causes a minor difference in colour of the salt as compared to Sm.

IMG_3643_adj2.JPG - 312kB

Indium
0.096g of indium were dissolved in 10% HCl. This took a long time, although the acid was heated to near boiling. The indium was precipitated with ammonia solution and the supernatant liquid pipetted off 3 times. 300mg of violuric acid were added. The solution turned purple immediately, which suggests that not all of the ammonium had been removed, although the pH of the solution with the In(OH)3 was 8.
On heating hardly any of the precipitate dissolved. The beaker was covered with cling film and put on a stirrer at 1200 rpm and heated at 70C plate temperature and stirred for 1 hour. Still much In(OH)3 remained undissolved, but the mix had obtained a colour resembling molybdenum blue.

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It seems that indium is chemically too basic to easily form a salt with violuric acid. My estimate is that this will count all the more for Re, Mo and W. I decided that I will not make the uranyl salt, because of the precautions necessary to safely work with it. I consider cerium, which is interesting because it is a coloured RE and also has a stable +IV oxidation state, and aluminium.

DraconicAcid - 16-3-2021 at 12:32

So violuric acid is made from barbituic acid....I wonder if something similar could be made from Meldrum's acid (which seems easier to make than barbituic).
https://en.wikipedia.org/wiki/Meldrum%27s_acid

Texium (zts16) - 16-3-2021 at 19:01

Quote: Originally posted by DraconicAcid  
So violuric acid is made from barbituic acid....I wonder if something similar could be made from Meldrum's acid (which seems easier to make than barbituic).
https://en.wikipedia.org/wiki/Meldrum%27s_acid
I think it should be possible! Meldrum’s acid is a very nice compound. I’ve worked with it a couple of times and it loves to undergo those sort of reactions. Without the nitrogens though, it may not have the same interesting coordination properties. It’s also not the most stable thing since it’s a geminal bis-lactone. Still, definitely worth a try.

RustyShackleford - 17-3-2021 at 06:44

Shame that the indium didnt form a solid, the solution looks real nice though. Also im quite suprised the Sm and Pr are so similar in color, from all the other salts i woulve expected a bigger differece. They both look quite nice though, thank you for taking the time to produce the salts and post pictures!

Lion850 - 17-3-2021 at 10:55

I follow this thread with much interest. Sadly up to now I did not manage to find any precursors to make violuric acid in Australia. And local prices I was quoted for violuric acid is astronomical.

Boffis - 17-3-2021 at 12:48

Depending on what other basic chemicals you have access to (ie nitric acid) you can prepare violuric acid from bird shit!

Guano -> uric acid -> alloxan -> alloxan-5-oxime (violuric acid)

The first step is the most difficult and complex because of the need to filter intensely alkaline solutions that result from the extraction of guano with caustic soda

vano - 27-3-2021 at 03:55

Hi. RustyShackleford gave me violuric acid and will make some violurates. Today i made scandium violurate, i used scandium carbonate. Nice yellow- brown colour.

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RustyShackleford - 27-3-2021 at 04:40

Quote: Originally posted by vano  
Hi. RustyShackleford gave me violuric acid and will make some violurates. Today i made scandium violurate, i used scandium carbonate. Nice yellow- brown colour.

Very nice! Thats the first yellow violurate salt i believe

vano - 27-3-2021 at 04:50

Thanks. This is indium violurate. Brown- dark red colour. I used indium carbonate. Carbonates are better.


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[Edited on 27-3-2021 by vano]

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