Toma's acid salt: ammonium guanidinium bisulfate sulfuric acid salt

[deltaH]

WARNING: This preparation involves boiling strong concentrated acids that would corrode flesh instantly if it lands on you and also causes terrible thermal and chemical burns. It should only be attempted while wearing a plastic apron, face shield, thick rubber gloves and labcoat while working in a fume cupboard and having a water shower source nearby and by skilled practitioners. The product when wet with attack many hydrolysable plastics such as polyesters and polyamides, as well as many metals. The use of polypropylene and glass is strongly advised. If you carry out this reaction, you do it at your own risk!

Introduction:

From my investigations into trying to decarbonate/hydrolyse urea into ammonium bisulfate in sulfuric acid media at high temperature, I appear to have stumbled upon the formation of a peculiar highly acidic salt and will report my continuing investigation into it here and also cordially invite members to comment and advise. For the time being I am naming it Toma's acid salt or alternatively ammonium guanidinium bisulfate sulfuric acid salt.

 

Hypothesis:

The salt's proposed formula is NH₄•C(NH₂)₃•2HSO₄•2H₂SO₄

The proposed chemical equation is: 2(NH₂)₂CO(s) + 4H₂SO₄(l) + heat => NH₄•C(NH₂)₃•2HSO₄•H₂SO₄(l) + CO₂(g)

 

Experimental:

I prepared it as follows (original work):

 

The reaction was conducted with a mole ratio of sulfuric acid to urea of 2. To 1000 ml dilute sulfuric acid (37% H₂SO₄, estimated by density) was added 113g urea prills in a large 5l heat and acid resistant glass pot. NB: it is essential that one has at least five times the volume a container as the mixture will form a rising effervescent froth later on and you need the freeboard space to prevent spillover! This was then boiled on a medium flame until the volume was greatly reduced and boiling subsides. Then a slight effervescence is noticed by the formation of smaller bubbles that don't form vapour condensation (steam). At this point the mixture is carefully watched as the effervescence will gradually accelerate and then run away creating a rising froth. When this occurs, the heating is immediately turned off to manage the effervescence which subsides soon afterwards. Thereafter, gentle heating is continued until a clear solution with no effervescence is observed. Heating is then turned off and the clear melt is allowed to solidify upon cooling.

 

WARNING: The product is extremely acidic, corrosive and probably hygroscopic/deliquescent, so store and handle with great care (handle as if it were conc. H₂SO₄)!

 

Observations/results:

 

 

Figures from top to bottom: (i) Boiling subsiding and effervescence begins; (ii) effervescence strengthens; (iii) effervescence is about to run away; (iv) effervescence subsided; (v) reaction done and no further changes apparent on heating and (vi)crystal mass obtained upon cooling

 

Some observations and tests so far (very rough and unverified):

1. The product is extremely water soluble, the whole mass nearly completely dissolved when 200ml of water was added at room temperature.
2. The product is slightly discoloured amber, but otherwise appears as glassy moist transparent crystals.
3. The product required more than twice the NaOH followed by Na₂CO₃ to neutralise to the point where ammonia is smelled persistently, compared to it all being a bisulfate salt (crude estimation for now... the amount required caught me by surprise so I had to back calculate a little).
4. Further (excess) base addition results in a strong strong ammoniacal smell.
5. The effervescence stage of the reaction subsides sharply when it ends.
6. Repeating the reaction with 1/4 acid withheld (i.e. an acid to urea mole ratio of 1.5) yielded a much faster, almost instantaneous effervescence ran away when it occurred. Upon cooling and adding the remaining 1/4 portion of acid, it was observed that most of the transparent crystals quickly dissolved while warming, but a mass of milky glassy/bubbly material was left behind which dissolved more slowly. Reaching the same hot stage in the melt now produced only a partial effervescence compared to the 3/4 acid step and doesn't run away.
7. Nothing further appears to occur on prolonged heating after effervescence subsides. The reaction appears to have reached a conclusion and the product melt appears relatively inert even though very hot.

Discussion:

Since I required more than twice the amount of base to neutralise my salt compared to if ammonium bisulfate was formed only (based on a crude estimation, but nevertheless) I conclude that my salt appears to have an empirical acid composition of X⁺ HSO₄^-.H₂SO₄ (since it only partially 'wetted' by acid, I must conclude that the acid is structural). Since decarbonation was observed, I am faced with the challenge of proposing what X could be in light of the all the observations made and the reaction stoichiometry used. My only hypothesis that fits all the observations that were made is that X is a 1:1 mix of ammonium and guanidinium ions, making my salt's formula: NH₄.C(NH₂)₃.2HSO₄.2H₂SO₄ (to be confirmed!) as the all the urea appeared to react completely to form something containing ammonium and evolving CO₂ gas, but couldn't only have been ammonium bisulfate based on the estimation of free acidity, so guanidium is the only thing I can propose that fits the observations. For now, until confirmed, I can only speculate that some hydrolysis/decarbonation of urea occurred to form ammonium ions, but that these also attacked other urea to form guanidine and water giving a mixed ammonium guanidinium bisulfate product and excess acid (that happened to crystallize in the double salt as a solvent of crystallisation) instead of the expected ammonium bisulfate alone. The water consumed by urea hydrolysis to ammonia and generating CO₂ is replenished by the elimination of water from the attack of ammonia on urea, so the net effect is that the reaction is water neutral.

 

Remaining questions:

1. Does my product consist of sulfamate?
2. Does my product consist of a urea salt?
3. To perform a more accurate neutralisation to determine acid content.
4. Can I specifically demonstrate the presence of guanidine in my product?
5. General characterisation (solubilities, m.p. determination, etc.)
6. Mechanism of its formation?

Additional questions:

• Is this salt novel and is the name appropriate?
• I encourage members to reproduce my investigation if they have the safe means to do so and help characterise this interesting salt! I am highly limited in terms of my equipment for now (but not on enthusiasm!), so collaboration would be greatly appreciated and please post your findings here!

Proposed uses:

• A solid, relatively inert and economical source of sulfuric acid.
• Highly acidic electrolyte for electrolysis.
• Homogeneous acid catalyst.
• Highly acidic ionic liquid medium (molten)

General comment:

These are early days into my investigation of this salt, so apologies in advance for the lack of information. As I repeat the investigation and try other tests, as well as improving my accuracy (which is very crude at the moment), I will update. I will constantly try to falsify my hypothesis and remain critical of it for a while still as I admit it is complex.

 

Thank you in advance for your comments and participation!

Edited November 5, 2013 by deltaH

[hypervalent_iodine]

!

Moderator Note

In addition to the warning by the OP of the dangers of working with the chemicals mentioned here, we urge everybody to look up all the MSDS (Material Safety Data Sheets) for the proper personal protection when experimenting with these chemicals. Mistakes with these chemicals can result in permanent injury or even death.

 

If you don't know what you're doing, STOP doing it.

[deltaH]

Thanks hypervalent_iodine and so true! Unfortunately I don't think MSDS' exists for my product as I believe it is novel, however, urea's and sulfuric acid's MSDS' certainly do exists and as for the product, familiarising oneself with those (particularly conc. sulfuric acid) as well as for ammonium bisulfate and guanidinium bisulfate or sulfate should come close enough!

 

I have also found some literature that may be pertinent to this topic, namely, the known preparation of guanidine sulfates from urea and sulfamic acid. If my formula hypothesis is correct, then these studies may suggest that it proceeds via the intermediate formation of sulfamates/sulfocarbamates, cyanamides intermediates and ammonium bisulfate off course.

 

From a thread posted by user Axt on another forum on the preparation of nitroguanidine from sulfamic acid and urea, I've pulled the following references about the formation of guanidine sulfates from sulfamates and urea:

 

1. J. S. Mackay “Preparation of Guanidine Sulfates” US patent #2464247 (1949)

2. J. L. Boivin and A.L. Lovecy “Mechanism for the Formation of Guanidine from Urea and Ammonium Sulphamate” Canadian Journal of Chemistry; 33[7]; pg. 1222-1225; (1955)

3. J. L. Boivin and M. Tremblay “Synthesis of Guanidine from Urea, Ammonium Benzenesulphonate, and Ammonium Sulphamate” Canadian Journal of Chemistry; 36[2]; pg. 378-382; (1958)

 

I believe the most urgent priority in my investigations is to try to investigate the sulfamate content of my product (in case that is a major component) and also try to expressly test for the presence of guanidine and urea.

 

For a guanidine test, I am thinking of employing guanidine salts' property of being a powerful protein denaturant (chaotropic agent). For this test I would neutralise a small amount of my mystery salt using sodium bicarbonate solution, then add a few drops of egg white solution and hopefully see it turning white, but perhaps this type of denaturation causes albumin to remain transparent? The problem is that urea is also a chaotropic agent, though not as strong as the guanidinium ions I believe. I am hoping that any biochemists here can advise me on this dilemma and also whether egg white would work and how it would appear.

 

Secondly, as for a sulfamic acid test, I am considering adding a small amount of KNO₃ to a tiny portion of my salt dissolved in a little H₂SO₄. If significant amounts of sulfamic acid or sulfamates are present, then I would expect to see gas bubbles produced. I think chemists may understand why I do not want to discuss the details of this test (unless given permission to do so).

Edited November 6, 2013 by deltaH

[John Cuthber]

I'm surprised you didn't cite this thread from the same site

 

http://www.sciencemadness.org/talk/viewthread.php?tid=26993#pid306040

 

which includes comments like "This is not chemistry, not even speculations. Just plain brainfart."

 

 

"Secondly, as for a sulfamic acid test, I am considering adding a small amount of KNO₃ to a tiny portion of my salt dissolved in a little H₂SO₄. If significant amounts of sulfamic acid or sulfamates are present, then I would expect to see gas bubbles produced."
Nope, try KNO2

[deltaH]

I'm surprised you didn't cite this thread from the same site

 

http://www.sciencemadness.org/talk/viewthread.php?tid=26993#pid306040

 

which includes comments like "This is not chemistry, not even speculations. Just plain brainfart."

 

 

"Secondly, as for a sulfamic acid test, I am considering adding a small amount of KNO₃ to a tiny portion of my salt dissolved in a little H₂SO₄. If significant amounts of sulfamic acid or sulfamates are present, then I would expect to see gas bubbles produced."

Nope, try KNO2

Kudo's for your creative attempt to enflame by reference, it didn't work there, it doesn't work here either. How can it... it's just an insult, you don't even make a logical argument to counter? What's worse is that link pertains to my investigation to find an OTC route to prepare ammonium carbamate from urea, the discovery of the salt I discuss here is in fact one of the reasons that route I investigated is low yielding!

 

As for your suggestion of KNO2... this is particularly dangerous, trying to get me to blow myself up with azo salts... nice

 

The action of nitric acid (generated in situ by addition of KNO3) upon sulfamates is a known route for producing nitrous oxide gas and it is far less likely to explode (unlike nitrites and azo salts). This is not because KNO2 wouldn't produce a gas with sulfamates (it does, N2), but it's more the issue of what diazotization can do on other things that may be present here as well as the inherent bad danger of any reaction that produces N2 (these tend to run away far too easily). Have you considered what would occur if my hypothesis is in fact correct and I added KNO2?

Edited November 7, 2013 by deltaH

[Phi for All]

Kudo's for your creative attempt to enflame by reference, it didn't work there, it doesn't work here either. How can it... it's just an insult, you don't even make a logical argument to counter?

 

!

Moderator Note

Attacking the idea is what we do in science. Attempts to refute your ideas are NOT insults, as long as they're done in a civil manner. Attacking the person is against SFN rules of civility. Please understand the distinction and if you feel you're being attacked personally, don't hesitate to report the post.

[deltaH]

Hi Phi for All, absolutely! Attack on my idea is welcome by employing scientific arguments, however, John Cuthber makes no such attempt, merely references an insult and an unrelated post (as I've explained), and yes, I've reported it almost the moment I saw it

 

Now for those who have scientific arguments to make... I will be EXTREMELY grateful to them.

 

*****************************************

 

I would like to make a correction to my original post, there is a typo in my equation:

2(NH₂)₂CO(s) + 4H₂SO₄(l) + heat => NH₄•C(NH₂)₃•2HSO₄•H₂SO₄(l) + CO₂(g)

Should read:

2(NH₂)₂CO(s) + 4H₂SO₄(l) + heat => NH₄•C(NH₂)₃•2HSO₄•2H₂SO₄(l) + CO₂(g)

 

 

i.e. The "2" in front of the right-hand-side H₂SO₄ was accidentally omitted, however, this is correctly reported in the rest of the text.

Edited November 7, 2013 by deltaH

[John Cuthber]

Delta

It's clear to me that you have little or no idea what you have made.

If you think that adding a stronger oxidant (like HNO3) rather than HNO2) will be less likely to provoke a violent reaction, you have not thought it through.

 

If you use KNO2 you can run the reaction in dilute acid, yet you think that's somehow less safe.

 

The point is pretty much moot. Ammonium ions will give rise to gas evolution with nitrate or nitrite.

 

Guanidinium nitrate doesn't play nicely with other children either.

The real problem with your assertion is that the hydrolysis of guanidine by boiling it with acid is well known and has been for nearly a hundred years.
http://pubs.rsc.org/en/content/articlelanding/1926/JR/jr9262901213#!divAbstract

 

So, what you are saying is that you are making guanidine in circumstances where it is known not to persist.

 

 

 

Anyone who looked through the thread I cited will understand why I did so.

[deltaH]

"It's clear to me that you have little or no idea what you have made."

 

I have as far as I can tell obtained an unknown product that is not described by the literature to the best of my knowledge. So what do we do in such a case? We perform experiments and propose a theory that fits the observations. I have started this process and the current best theory I have for it is NH₄•C(NH₂)₃•2HSO₄•H₂SO₄ based on the experimental observations. By the scientific method, this is correct until proven otherwise, but it's not necessarily the Answer/Truth. If you dispute why my observations do not fit this proposed formula, then please be specific. I will also be trying to invalidate my formula by performing specific tests. These will be ongoing for a while still.

 

"If you think that adding a stronger oxidant (like HNO3) rather than HNO2) will be less likely to provoke a violent reaction, you have not thought it through."

 

My main issue with KNO₂ is that diazonium salts that may form as intermediates are notoriously unstable, nitro products (nitroguanidine, nitrourea) are far less so.

 

"The point is pretty much moot. Ammonium ions will give rise to gas evolution with nitrate or nitrite."

 

Good point, do you have an alternative suggestion?

 

"...the hydrolysis of guanidine by boiling it with acid is well known..."

 

From the first page of that article you cite: "CN₃H₅ + H₂O <=> CON₂H₄ + NH₃", I do not have access to the remainder of the article.

 

The decarbonation reaction only kicks off once the salt melt stops boiling. It thus appears that conditions are pretty anhydrous at this point. As I've discussed in my opening post, this reaction is water neutral. While water is transiently produced by guanidine formation from urea and ammonia, it is consumed by urea hydrolysis that generates more ammonia. In the end of the day its the thermodynamics that counts and the loss of CO₂ here is the main driving force that drives this whole thing over. The only way that the mass balance for this system can be satisfied is if an equimolar amounts of ammonia and guanidine formed from the urea component. I challenge you to suggest an alternative that fits the mass and elemental balance and experimental observations reported.

 

"So, what you are saying is that you are making guanidine in circumstances where it is known not to persist."

 

As before, observational evidence suggests anhydrous conditions at the end of the reaction (also at its onset). So I don't see the problem. The remainder is simply chemical equilibrium between all the species and that is driven by the loss of CO₂ to completion by Le Chatelier's principle... by coincidence, also known for about a hundred years (well a little over).

 

"Anyone who looked through the thread I cited will understand why I did so."

 

You are still strawmanning, please state your point plainly.

Edited November 8, 2013 by deltaH

[John Cuthber]

Come back when you have an IR spectrum or something indicating a guanidinium salt.

[deltaH]

Come back when you have an IR spectrum or something indicating a guanidinium salt.

Sadly IR is out of the question for my means at the moment, if somebody would like to assist in this regard, it would be greatly appreciated! Failing this, I can only try a 'wet' chemical test for guanidine salts, as a matter of priority.

 

*********************************************

Okay, I had some time today and so decided to prepare another batch because I need material for testing (I had neutralised the material from my first batch a while ago).

 

Unfortunately, my retailer was out of the good quality battery acid I bought last time (I estimated it's H₂SO₄ conc. to be about 37%). I had to buy a rather dubious no-name brand battery acid from a 'back alley' automotive parts vendor. I determined it's SG to be about 1.15 which is shockingly low. That would make it about 21% H₂SO₄, however, we also start encroaching on the concentration where my crude SG to concentration determination starts to have large errors (I only have a gram scale and kitchen volumetric graduated measure).

 

Anyhow, I soldiered on, I used 500ml of my battery acid and calculated that 64g urea would give me an acid to urea mole ratio of 2. Proceeded with the experiment as before. The solution took an exorbitantly long time to boil down this time. The boiling was also sluggish compared to before. My gut feeling was telling me that there was even less sulfuric acid in this solution than what I determined and this was a concern, because I observed a violent reaction before when I had used an acid to urea ratio of 1.5, so having too much urea in the mix is dangerous for this reaction.

 

Anyhow, my gut feeling proved correct, luckily because of my observation, I had the wisdom to turn the heating to low and observe from afar. This time, when runaway effervescence occurred, it was very rapid, almost instantaneously, produced a lot of white fumes and instantly formed a solid mass of white and brown product that appears poorly crystalline (see photo).

 

 

If I assume that this battery acid was merely fudged by too much dilution, then this suggests that the concentration of the starting acid used may be important for the selectivity of this reaction! I therefore conclude that it's important that 37% H₂SO₄ be used or closest to!

 

Some additional observations on the material I made here:

 

• It is bone dry, appears brown/white, hard and brittle (excess acid does not appear to be present) amorphous looking mass at the bottom and a slightly more crystalline white mass on the top and sides.
• While it is water soluble, but much more sluggish in dissolving compared to Toma's acid salt which was extremely fast dissolving upon addition of even a small amount of water.
• It doesn't appear particularly hygroscopic.
• Adding some Na₂CO₃ to a solution prepared from the bottom more amorphous looking mass does produce effervescence, but not a great deal... it doesn't appear to contain much acid. Solutions of it didn't appear particularly acidic to the touch either.
• An ammoniacal smell is present on excess carbonate addition, but it is not very strong but this is very subjective.

Combining these observations, I would say that this material might be rich in sulfamates and/or sulfates suggesting that my acid to urea molar ratio was way too low resulting in far too violent a runaway reaction and very high temperatures favouring sulfamate production. In any event, whatever happened, there was definitely insufficient acid present!

Edited November 8, 2013 by deltaH

[John Cuthber]

Sodium dodecyl sulphate will form a precipitate with guanidinium ions but not most of the other plausible ions present.

Also, learn to do titrations, that will let you produce much better quantitative data.

[deltaH]

"Sodium dodecyl sulphate will form a precipitate with guanidinium ions..."

Intriguing and thanks for the suggestion. As I struggled to verify this, do you have a reference for me please?

 

"Also, learn to do titrations, that will let you produce much better quantitative data."

I would need a burette and at least a two decimal scale for that. This would cost me more money than I can afford at the moment. I admit it would come in very handy, but it's something that I will need to put off for a while.

[John Cuthber]

Even with a scale that only weighs to the nearest gram you can do weight titrations - you just need to use dilute solutions so you end up weighing a lot of water.

http://www.uclmail.net/users/dn.cash/GravTitr1.pdf

[deltaH]

Even with a scale that only weighs to the nearest gram you can do weight titrations - you just need to use dilute solutions so you end up weighing a lot of water.

http://www.uclmail.net/users/dn.cash/GravTitr1.pdf

 

Indeed and thanks. Have you a reference for the guanidine precipitation you mentioned? I still cannot find info on that.

Edited November 8, 2013 by deltaH

[John Cuthber]

It's from this side comment "As a final note on chaotropic salts, guanidinium-HCI has the undesirable property of forming an insoluble precipitate (guanidinium dodecyl sulfate) in the ."

here

http://books.google.co.uk/books?id=4Yp2p-PqDHwC&pg=PA313&lpg=PA313&dq=%22precipitate+guanidinium%22&source=bl&ots=cTift7vo8S&sig=uP4r0rT76h2N0DBAEscLWHA5-L8&hl=en&sa=X&ei=TBh-UtGlI8bA7AbjkoCoBA&ved=0CC4Q6AEwAA#v=onepage&q=%22precipitate%20guanidinium%22&f=false

[deltaH]

It's from this side comment "As a final note on chaotropic salts, guanidinium-HCI has the undesirable property of forming an insoluble precipitate (guanidinium dodecyl sulfate) in the ."

here

http://books.google.co.uk/books?id=4Yp2p-PqDHwC&pg=PA313&lpg=PA313&dq=%22precipitate+guanidinium%22&source=bl&ots=cTift7vo8S&sig=uP4r0rT76h2N0DBAEscLWHA5-L8&hl=en&sa=X&ei=TBh-UtGlI8bA7AbjkoCoBA&ved=0CC4Q6AEwAA#v=onepage&q=%22precipitate%20guanidinium%22&f=false

Thank you very much for the link and advice. I was concerned because ordinarily a larger organic cation on an anionic surfactant tends to increase solubility of the surfactant. I myself have prepared choline palmitate soaps that are extremely soluble (compared to their sodium counterparts).

 

I took some liquid dishwashing detergents and mixed it with some of the partially neutralised solution that I had sitting outside evaporating from my first successful batch experiments. In theory, there should be guanidine sulfates in there.

 

This caused immediate precipitation upon squirting in some of the liquid detergent, but I believe this is simply a salting out effect. This lead me to believe that the effect described in that book is also a salting out effect. Normally when guanidinium hydrochloride solutions are used as denaturants, they are used at very high concentrations.

 

So I decided to test this out and prepared a salt solution and squirted in some liquid detergent and indeed observed the same precipitation (even worse). So this may well be a salting out effect and so not a good test for guanidine presence specifically.

 

************************************************************

Results of raw egg white denaturation test

 

I have now tried a egg white denaturation test making use of guanidinium salts' property of being powerful chaotropic agents.

 

I used my partially neutralised left over solution evaporating outside (left overs from the first two successful preparations) as is because I read for guanidine hydrochloride as a denaturant, they advise the use of 6M solutions, so wanted to give mine the best shot at it and use it as concentrated as possible. I mixed a little raw egg white with a small amount of water just to lower the viscosity and then poured this into my concentrated partially neutralised solution of the salt. The denaturation was immediate yielding a congealed white solution:

 

 

The test was therefore positive. Also, this test requires high guanidinium salt concentrations to affect rapid denaturation, so I would say I have a lot of it in there!

 

Control 1:

 

I was then concerned about the acidity of my solution (it was not fully neutralised and still caused Na₂CO₃ to fizz) possibly giving a false positive result, so I decided to do a control and use my ca. 20% battery acid solution and added raw egg white to it. Nothing happens, it simply dissolves and remains clear.

 

Control 2:

 

I was also concerned that this test might not differentiate between urea and guanidine (as urea is also a chaotropic agent, albeit much weaker). So I dissolved a couple of teaspoons of urea in about 50 ml water and added a few milliliters of battery acid to acidify and the added my raw egg white solution.

 

The solution remained colourless indicating, as I had suspected, that urea is not as powerful a chaotropic agent under acidic conditions compared to the action of guanidine.

 

I am now satisfied that there is additional and specific evidence for the presence of guanidine and in fair amounts. Again, please note, this is not definitive, but merely extra evidence that is in agreement with my hypothesis.

Edited November 9, 2013 by deltaH

[John Cuthber]

Congratulations, you have proved that the stuff either contains guanidinium or ammonium sulphate, both of which will precipitate proteins.

 

http://web.mnstate.edu/provost/AmmoniumSulfateProtocol.pdf

But I didn't think that was a matter of contention.

 

You have shown that egg white proteins are soluble in 20% sulphuric acid.

Again, that's probably not news.

 

What you need is something a bit more definitive.

Do you have a pH meter?

A plot of pH vs amount of added sodium hydroxide would give a pretty clear indication of the presence of Sulphuric acid, bisulphate, and ammonium ions.

Even if you don't know the exact concentration of the hydroxide, the ratios of the amounts would be informative.

[deltaH]

"Congratulations, you have proved that the stuff either contains guanidinium or ammonium sulphate, both of which will precipitate proteins.

http://web.mnstate.e...ateProtocol.pdf But I didn't think that was a matter of contention."

Thank you for highlighting a possible problem in my guanidine test and attaching your protocol article. There is, however, a big difference between precipitating proteins and denaturing proteins. As per your article, ammonium sulfate precipitates without denaturing. This should redissolve again upon addition of large volumes of fresh water. I will repeat my test and investigate this point specifically. If they are indeed denatured and not simply precipitated, then the solution will return to being colourless (protein will redissolve), especially as acidity helps the solubility of albumin by what I have observed in my acid dissolution test.

"Do you have a pH meter?"

Unfortunately not.

"A plot of pH vs amount of added sodium hydroxide would give a pretty clear indication of the presence of Sulphuric acid, bisulphate, and ammonium ions. Even if you don't know the exact concentration of the hydroxide, the ratios of the amounts would be informative."

I plan to do a mass titration as you have suggested when I can once again purchase reasonable quality battery acid and prepare a fresh batch of relatively pure salt (right now I only have a partially neutralised mess that is only good for qualitative guanidine tests). I'm waiting on retailers to restock... this can take weeks here. While not as informative as pH plots, it will at least tell me if my empirical formula X⁺ HSO₄^- H₂SO₄ is correct.

I have already evidence that part of X⁺ is ammonium, I need to know if there are appreciable amounts of guanadinium as well, which is why a guanidine specific test is important. After that, I will attempt to gather some evidence that this is indeed a true double salt and not mixture of two salts, e.g. something like a mixture of separately crystalline NH₄HSO₄.H₂SO₄ and C(NH₂)₃HSO₃.H₂SO₄.

*********************************************************************************************************************

Okay it appears you may be correct in part I took egg white and added it to a mass of undiluted raw egg white and stirred. It formed many congealed white masses as before. I let it sit for a while this time to give it maximum chance to denature. Then I added lots of fresh water. Most of it appeared to dissolve quickly.

This would indicate that a salting out had indeed occurred because of the re-dissolution.

However, there were many little gelatinous blob skins that didn't redissolve. I washed this repeatedly with fresh water and stirred, they sink to the bottom and resolutely don't dissolve.

So it appears that some denaturation had occurred as well as a lot of salting out.

Here is a photo of the gelatinous blobs that don't dissolve:

It would seem that the denaturation happens on the surface of the raw egg blobs and so forms these skins, while the inside part doesn't denature well and hence simply dissolves upon adding more water. Maybe this is why that book procedure you quoted earlier used surfactants, to help the guanidine salt penetrate?

Edited November 10, 2013 by deltaH

[John Cuthber]

It's possible that some denaturing of the protein is due to the acid rather tan guanidinium.

.

A lot of the recipes for meringue include vinegar or cream of tartar as a weak acid to coagulate the proteins.

Try neutralising some of the product and then see what it does to the egg white.

If you ad an excess of calcium carbonate to a solution of the material then filter off the calcium carbonate and sulphate you should get a fairly nearly neutral solution. If you boil it down that will drive off any ammonium carbonate and you should get a solution with any guanidinium salts and a little calcium sulphate.

(Barium carbonate would be better but calcium carbonate is much easier to get.)

"I have already evidence that part of X⁺ is ammonium, I need to know if there are appreciable amounts of guanadinium as well, which is why a guanidine specific test is important."

A titration would tell you not just that it is present, but how much.

 

You should also do a similar egg-white experiment wit ammonium sulphate (It's cheaply available as a fertiliser)

[deltaH]

"It's possible that some denaturing of the protein is due to the acid rather than guanidinium."

Yes exactly what I suspected, which is why I did that control test with some diluted battery acid. But the egg white simply dissolved more easily (instantly). It doesn't curdle like milk does upon acidification. I'll be honest, qualitatively testing for guanidine is a tough problem and I cannot call it neither here nor there with these protein tests anymore. There is some evidence for a positive test, but it's not dramatic enough to be a resounding yes.

 

"Try neutralising some of the product and then see what it does to the egg white."

I did that too, it doesn't work as well, slightly acidic works better. But I wouldn't jump to a conclusion here just yet because again, egg white in plain acidified water dissolves very well and doesn't seem to do anything. It could be that guanidinium bisulfate works better than the sulfate. Microbiologists, if I am not mistaken, use the chloride or thiocyanate... apparently those work best, so the anion does seem to have some effect on the effectiveness of the guanidine.

 

"If you add an excess of calcium carbonate to a solution of the material then filter off the calcium carbonate and sulphate you should get a fairly nearly neutral solution. If you boil it down that will drive off any ammonium carbonate and you should get a solution with any guanidinium salts and a little calcium sulphate. (Barium carbonate would be better but calcium carbonate is much easier to get.)"

I like this idea a lot. It just so happens that I have a lot of calcium carbonate powder. It's not reagent grade or anything, it's a finely milled white marble powder that has been purified by flotation. This should be fine. So yes, if I'm successful, I might be able to isolate some guanidinium carbonate by this... I'll give it a shot when next I have some spare time. Barium carbonate is out of the question, but yeah, it would have been better, of course.

 

"A titration would tell you not just that it is present, but how much."

I'll give a titration a shot next I can.

 

"You should also do a similar egg-white experiment wit ammonium sulphate (It's cheaply available as a fertiliser)"

It's not OTC here in small quantities. It's available from a bulk chemical supplier, but then I'd have to buy 40 kg bag or something like... too much, I'm already sitting with a 40kg bag of urea... what started this investigation in the first place lol

[John Cuthber]

"Try neutralising some of the product and then see what it does to the egg white."

I did that too, it doesn't work as well, slightly acidic works better."

What does that tell you?

The guanidinium ion will be present as that ion in acid, neutral or slightly alkaline solution so if it's doing the denaturing, the pH shouldn't matter.

 

 

Are there shops that cater for home gardeners where you are? That's where I get ammonium sulphate from.

[deltaH]

"Try neutralising some of the product and then see what it does to the egg white."

I did that too, it doesn't work as well, slightly acidic works better."

What does that tell you?

The guanidinium ion will be present as that ion in acid, neutral or slightly alkaline solution so if it's doing the denaturing, the pH shouldn't matter.

 

 

Are there shops that cater for home gardeners where you are? That's where I get ammonium sulphate from.

I don't think it's as simple as that John, especially as I've already seen that dilute H₂SO₄ made albumin dissolve very well and not cause it to denature. As for the ammonium sulfate, I've never seen it at our nurseries here, but I will check next time I am nearby.

 

Nevertheless, I am not happy with the denaturation test as whole anymore, it's just too neither here nor there for my liking. I think I need to concentrate on isolating the guanidine.

 

A good friend of mine from the SM forum has made a suggestion to add excess sodium hydroxide and use that to drive off the ammonia, provided the guanidine can survive and not be hydrolysed.

 

I've also been thinking about your suggestion about the excess calcium carbonate. While I think this would work well for removing excess sulfate, I think it will start to fail when the pH reaches neutral. But the idea of using the calcium to scavenge sulfate is good.

 

So, I'm going to try a hybrid of both these suggestions. I intend to use lime to scavenge sulfate and generate free base, but I will need to wait until I have fresh salt prepared, because my evaporating messes outside have way too much sodium in them to be good (I've neutralised them fully now with washing soda).

 

Okay, my plan is this, starting from freshly prepared acid salt:

 

1. Neutralise with limestone until effervescence ceases. Noting the mass required here to do so will give me an idea of the acid content of my salt. I should now have a solution of guanidinium sulfate and ammonium sulfate (more or less) if my hypothesis is correct.
2. Add burnt lime flakes and ice, then decant when all's done. The solution should now have had the majority of sulfate removed and be highly caustic. If necessary, I can pass it through a short column of slaked lime afterwards just to mop up and drive the point home, but I would need the majority of the sulfates removed before doing so or my column will become a solid impervious slug of plaster lol.
3. Finally, evaporate the solution for a few days to remove most of the ammonia and concentrate the guanidine hydroxide solution which will absorb atmospheric CO₂ and form guanidine carbonate solution.

 

Hopefully, guanidine hydroxide solutions are stable enough to survive the transitory steps here... I will do some digging...

 

Now I can do pretty much whatever I please with the guanidine carbonate if it's there

 

One idea is to neutralise with HCl, then evaporate and isolate guanidine hydrochloride as a solid... but anyhow, one step at a time, there's a lot to do in the meantime.

Edited November 10, 2013 by deltaH