Sciencemadness Discussion Board

Thiourea oxidation at low pH

woelen - 3-12-2012 at 07:23

I did a nice experiment with thiourea in which this compound is oxidized.

The first experiment I did is the following:
- Dissolve some thiourea in water
- Add some dilute H2SO4
- Add some 10% H2O2
When this is done, the solution heats up somewhat. The solution remains colorless.

I repeated this experiment with dilute HClO4, dilute HCl and dilute HNO3. In all cases the solution heats up somewhat after adding the H2O2 and the solution remains colorless in all these experiments.

With HNO3, however, a special effect occurs. After a while, glittering crystals separate from the liquid. Everywhere in the liquid, beautiful glittering crystals are formed and these slowly settle at the bottom and a nice compact layer of solid white material is formed.

I repeated the experiment with other oxidizers and the result is similar in all cases. With nitric acid, however, a white crystalline solid is formed, not at once, but slowly over the course of a few minutes, while the other acids simply produce a clear colorless solution.

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I separated the white solid, rinsed with icec old water (it does not dissolve in cold water) and allowed to dry in contact with air. While drying, it turns somewhat yellowish.
I put aside some of the material in a dry and clean test tube and set this aside. To my surprise, several hours after the material was put in the test tube, it suddenly started smoking and burning (with invisible flames, but with a LOT of smoke).
The dry material also is very sensitive for heat.

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I have done some research about oxidation of thiourea. Unfortunately I did not find a nice article about this subject, but from scattered data on internet I found that oxidation of thiourea can lead to many different products, depending on pH and concentration of the oxidizer. At low pH, however, it seems that normally 2 molecules of thiourea combine to one molecule of formamidine disulfide, C(NH2)(=NH)-S-S-C(=NH)(NH2), and acid according to the following half reaction:

2 thiourea --> formamidinedisulfide + 2H(+) + 2e

Now, I am wondering what the white crystalline solid is. Do the H(+) ions combine with formamidine disulfide to form a salt? I am wondering what would be the structure of this salt. Could the H(+) connect to the NH2-groups, or to the =NH?

Another very striking thing is the instability of the solid. It decomposes in a very exothermic reaction, hours after its dry preparation and isolation. I only stored 100 mg or so. It is scary that this small amount can produce so much heat that it can decompose so violently without being ignited!
Making this compound also requires only dilute aqueous solutions (e.g. 10% H2O2, 10% HNO3). It certainly is not a nitration product like nitrate esters or nitro compounds made from nitric acid, sulphuric acid and some organic compound.
If this is an ionic nitrate of a charge +2 cation and two nitrate ions, then this would be extra surprising to me. Ionic nitrates normally are quite stable and need quite a lot of heat before they can be ignited.

Anyone knows more about thiourea and this rather special and interesting compound?

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As a final test I prepared a solution of thiourea and added dilute HNO3 to this. Even after a long time, no precipitate is formed at all. As soon as a solution of an oxidizer is added (I tried with solution of Na2Cr2O7, with Na2S2O8 and with H2O2) then a white crystalline solid is formed. This experiment shows that this is not simply thiourea nitrate, it really is an oxidation product of thiourea combined with nitrate ions.

[Edited on 3-12-12 by woelen]

DJF90 - 3-12-2012 at 08:15

I know a little of the chemistry of thiourea(s). Thiourea itself can be oxidised to the sulfonic acid, a decent guanylating agent for amines. IIRC the reagent for this is peracetic acid. Substituted thioureas can be converted to carbodimides under certain conditions. Thiourea also makes for an excellent primary reductimetric standard, and it reacts quantitatively with some common oxidisers under acidic conditions to yield the formamidine disulfide. I've attached a paper detailing this use as I believe it is the most useful for you.

[Edited on 3-12-2012 by DJF90]

Attachment: thiourea as reductimetric standard.pdf (276kB)
This file has been downloaded 1145 times

ScienceSquirrel - 3-12-2012 at 08:42

It is a known reaction, see here on page four, section ten;

http://www.hillbrothers.com/msds/pdf/n/thiourea.pdf

and there is a reference here;

http://www.chemicalbook.com/ChemicalProductProperty_EN_CB985...

[Edited on 4-12-2012 by ScienceSquirrel]

Very violent reaction of aqueous solution

woelen - 4-12-2012 at 12:48

Both references mention the combination of thiourea, nitric acid and hydrogen peroxide, but I think that this combination is too specific. This suggests that the white self-igniting compound is some peroxide species, but this is not the case. I get formation of this compound with any oxidizer. As I wrote before, if I dissolve thiourea in dilute nitric acid (e.g. 10% by weight) and then add another oxidizer than hydrogen peroxide, then the white solid also forms. I tested with a solution of CrO3 in water and with Na2S2O8. The white solid only forms with nitric acid, not with any other mineral acid.

So, the MSDS and the chemicalbook references do not really give more information. I personally am inclined to think that the white compound must be something like the nitrate salt of formamidine disulfide, with H(+) attached to one or both of the NH2 groups in the formamidine disulfide.

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I also had another exciting experience with this compound.

I dissolved a spatula full (100 to 200 mg) of thiourea in 3 ml of dilute nitric acid (appr. 10% HNO3 by weight). This is a clear colorless solution, no special reaction occurs.
To this, I added a solution of apr. 300 mg of sodium peroxodisulfate in 3 ml of water. When this solution is added, then hardly any reaction occurs (peroxodisulfate is a very sluggish oxidizer at room temperature). I carefully heated the liquid, such that I just could bear that temperature on my bare hands (I think it is somewhere around 50 C). When this is done, then quickly a lot of white crystalline precipitate is formed. I allowed the crystalline mass to form and then strongly swirled, such that I obtained a white milky liquid and then I heated this liquid until it becomes quite hot (not boiling hot, but quite close to that). The white solid dissolved again. Suddenly, the liquid started producing very fine bubbles of a colorless gas, the liquid turned red/brown and then suddenly a very violent reaction occured. Light brown fumes were produced, the liquid almost was ejected out of the test tube. I quickly dropped the test tube in a sink and stepped back, not wanting to hold such an out-of-control thing in my hand. This violent reaction is quite scary. So, even in aqueous solution this compound can lead to very violent decomposition reaction.

[Edited on 4-12-12 by woelen]

ScienceSquirrel - 4-12-2012 at 14:45

I think that you may be right about the structure of the product and the original authors were wrong.
The quotes do speculate on the structure and maybe it really is a formamidine disulphide nitrate or dinitrate.
Maybe they guessed at the structure of the compound formed and did not realise the generality of the reaction?
A copy of the paper woud be nice :-)

Nicodem - 5-12-2012 at 02:17

Woelen, perhaps the best you can do to check if the compound is formamidinedisulphide nitrate salt, is to first prepare formamidinedisulphide via a known literature method and then check if it forms an insoluble salt/complex with ammonium nitrate (nitric acid is more likely to cause redox reactions and is thus not the best source of nitrate anions).

Nitrate salts rarely ever have poor aqueous solubility, but formamidines, just like ureas, are strong H-bond donors and can form complexes which are not necessarily as soluble as nitrate salts would be (one such typical complex is urea nitrate).
Quote: Originally posted by ScienceSquirrel  
Maybe they guessed at the structure of the compound formed and did not realise the generality of the reaction?
A copy of the paper woud be nice :-)

"Bjorklund G. H. et al., Trans. R. Soc. Can.,1950, 44, p. 28"

Unfortunately, only the volumes 1-16 are available digitally (via archive.org), while the Royal Society of Canada homepage says the best option is an interlibrary loan. This usually costs less than a dozen euros, but also some additional motivation and a visit to a library.

woelen - 5-12-2012 at 02:52

Quote:
Woelen, perhaps the best you can do to check if the compound is formamidinedisulphide nitrate salt, is to first prepare formamidinedisulphide via a known literature method and then check if it forms an insoluble salt/complex with ammonium nitrate (nitric acid is more likely to cause redox reactions and is thus not the best source of nitrate anions).
For this reason I used dilute acid (at most 10% by weight). My experience with dilute nitric acid is that it reacts much like an acid without anionic redox properties.
But I'll try your suggestion. I'll make the formamidine disulfide with dilute sulphuric acid and an oxidizer (e.g. Na2S2O8 or H2O2) and then I'll add an aqueous solution of NH4NO3 or KNO3 and see whether I still get the white precipitate or not. According to what I found on internet, when pH < 1, then formamidine disulfide is formed on oxidation of thiourea.

[Edited on 5-12-12 by woelen]

Nicodem - 5-12-2012 at 04:45

Apparently, formamidine disulphide does not exist as a compound. Any attempts at its isolation result in its decomposition. Only its protonated form is stable and its salts can be isolated. This explains why it can be prepared from thiourea oxidation only in acidic media.
Quote:
Formamidine disulphide is a very weak base, and has not been obtained in the free state; addition of alkali, or even sodium acetate, to its salts produces immediate precipitation of sulphur, indicating, in conformity with “ Fromm’s rule” (Annalen, 1906, 348, 144), the presence of the grouping -C-S-S-C- in its constitution

cited from:
CCXXVII.—The interaction of iodine and thiocarbamide. The properties of formamidine disulphide and its salts
Emil Alphonse Werner
J. Chem. Soc., Trans., 1912, 101, 2166-2180
DOI: 10.1039/CT9120102166 (attached)

The same article goes on to describe the dinitrate salt as a solid melting and decomposing at 110 °C.
Quote:
one of its most characteristic salts, namely, the dinitrate,* is almost insoluble in water containing about 5 per cent. of nitric acid, and hence is precipitated in crystalline form when an oxidising agent is added, even to a dilute solution of thiocarbamide containing an excess of this acid.

* This compound was originally obtained by McGowan (Trans., 1886, 49, 195), who assumed it to be a “dithiocarbamide” dinitrate, (CSN2H4)2(NO3)2, analogous to the compound (CSN2H4)2CI2 prepared by Claus, and which is really formainidine disulphide hydrochloride.

It also reviews older reports of this dinitrate salt and compares synthetic methods. They found the best method to be the oxidation of thiourea dissolved in cold diluted nitric acid using these oxidants: iodine (99% yield), NaNO2 (93%), H2O2 (90%), KMnO4 (74%).
And finally it gives this warning:
Quote:
The dinitrate is very unstable when in a fine state of division, and decomposes spontaneously when dry after a very short time, in several instances violent decomposition took place within two hours from the time of its preparation



Attachment: CT9120102166.pdf (229kB)
This file has been downloaded 869 times

Nicodem - 5-12-2012 at 05:52

Woelen, here is also the follow up report from the same author, which is not on the topic of the interaction with nitric acid, but on the interactions with nitrous acid instead. You might find it interesting anyway.

CCXXVIII.—The action of nitrous acid on thiocarbamide and on formamidine disulphide. A new structural formula for thiocarbamide
Emil Alphonse Werner
J. Chem. Soc., Trans., 1912, 101, 2180-2191.
DOI: 10.1039/CT9120102180

Attachment: CT9120102180.pdf (193kB)
This file has been downloaded 804 times

woelen - 5-12-2012 at 06:06

Nicodem, many thanks for all this information. It clears up things for me and also opens new directions of experimenting. What you describe is exactly what I observed. The spontaneous decomposition with a lot of smoke is quite spectacular. I only had 200 mg or so of the compound, so nothing serious happened, but I must not think of what could have happened if I did the experiment on a scale of household jars instead of test tubes.

I myself certainly am not able of getting access to that kind of information. For me, getting chemicals is easier than getting information about them :D (unless I am willing to spend a lot of money on paid papers).

woelen - 20-12-2012 at 11:26

I found some time to continue experimenting with the thiourea oxidation in the presence of nitrate ions.

I again made some of the white solid and decanted most of the liquid above the white solid. I did not dry the material, just left it wet. Then I added a large excess of bleach (appr. 12% active chlorine). When this was done, a lot of gas was produced. It was a colorless gas with a strong acrid smell (no chlorine, it smells very different).
I kept a flame near the open end of the test tube (I expected the flame to extinguish, believing that this gas is nitrogen), but when I did that, I heard a loud WHOOP sound and a nearly invisible flame very quickly went into the test tube. The gas I made is flammable and easily ignited. After burning of this gas, there was a strong smell of chlorine.

What could this gas be? I am not an organic chemistry experiment and I am quite surprised to get a highly flammable gas from such a strongly oxidizing environment.

[Edited on 20-12-12 by woelen]

woelen - 22-12-2012 at 13:22

I did some more testing and the formation of the flammable gas is unique to the oxidation product of thiourea. Thiourea itself does not show such a reaction. I did the following experiments:

1) Add some solid thiourea to bleach. When this is done, no gas is evolved, the liquid becomes turbid and it looks like there is formation of elemental sulphur. The precipitate is pale yellow.
2) Add some solid urea to bleach. A colorless gas is produced and all of the urea quickly dissolves. The gas is not flammable. Probably this gas is N2.
3) Prepare some of the solid material from an oxidizer, dilute HNO3 and thiourea. Rinse with some water to get rid of acid and other impurities and then add bleach. The liquid becomes yellow/green and a colorless gas is produced. If a flame is kept near the gas, then it burns with a popping sound.

The gas is not a sulfide, there is no smell of rotten eggs. This is something which I did not expect neither, because in the presence of bleach, sulfides are oxidized extremely quickly.

Adas - 23-12-2012 at 05:43

I think that the gas may be some chloramine. What else could it be if you smelled chlorine after the explosion...

plante1999 - 23-12-2012 at 08:11

What was the flame color?

You could do some test to test the gas proprieties.

Is it solube in basic solution, does it bleach humid colored paper?

After some tests it would be easier to know the gas composition.