Alchemist
Hazard to Self
Posts: 93
Registered: 22-6-2002
Location: Hostton Texas
Member Is Offline
Mood: No Mood
|
|
guanidine salts
Hello all,
has anyone had any real success in making any kind of guanidine salts? Yes, there has been some posts on the subject (example; urea explorations and
etc.), but has anyone really made any usable quanity in there home lab. Time to start experimenting. I hope to start this weekend. Anyone else wish to
help. Thanks...
P.S.; does anyone know of a good test for guanidine ?
P.S.2; I ment from OTC chemicals (like urea , etc.)
[Edited on 1-11-2004 by Alchemist]
[Edited on 1-11-2004 by Alchemist]
|
|
|
Mephisto
Chemicus Diabolicus
Posts: 294
Registered: 24-8-2002
Location: Germany
Member Is Offline
Mood: swinging
|
|
Guanidine picrate
As you asked about any kind of guanidine salts, I made guanidine picrate from guanidine nitrate and ammonium picrate. But the guanidine nitrate was
from another guanidine compound and not form cyanoguanidine, as you probably look for.
↑ Here a picture of the guanidine picrate.
Here you can find the German synthesis-text of guanidine picrate and a description of its properties (a part of LambdaSyn.com).
[Edited on 24-12-2005 by Mephisto]
|
|
|
JohnWW
International Hazard
Posts: 2849
Registered: 27-7-2004
Location: New Zealand
Member Is Offline
Mood: No Mood
|
|
I remember seeing a news item somewhere about guanidine, or a derivative of it, being used as an experimental anti-viral drug to treat and cure
influenza, administered as a nasal spray. Someone in Australia was doing research into it. Is that what Alchemist had in mind? Does anyone know
anything more about this?
|
|
|
chemoleo
Biochemicus Energeticus
Posts: 3005
Registered: 23-7-2003
Location: England Germany
Member Is Offline
Mood: crystalline
|
|
Alchemist, I am not sure what you are asking.
You can easily make the respective salts via metathesis, i.e. pretty similar to what mephisto did. This by the way I'd rather consider an unusual
synthesis, I suspect for a certain purpose...
Anyway, I handled GdnHCl and GdnSCN (thiocyanate) before, both are highly soluble salts.
The free base exists too, it can be prepared in an alcoholic solution to which NaOH is added. But I don't know how long these solutions last.
Anyway, what are you after? The synth. of guanidine salts, where you aren't really after the salt, but the guanidine itself?
Or on the various salts? Because making the salts is about as difficult as making various sodium salts.
PS wasn't the original process of making guanidine via guanidine thiocyanate? I.e. I think this used to be one of the main guanidine salts.
[Edited on 3-11-2004 by chemoleo]
Never Stop to Begin, and Never Begin to Stop...
Tolerance is good. But not with the intolerant! (Wilhelm Busch)
|
|
|
Quantum
Hazard to Others
Posts: 300
Registered: 2-12-2003
Location: Nowhereville
Member Is Offline
Mood: Interested
|
|
I was reading about the synth of guanidine nitrate using silica gel, NH4NO3, and urea. Just how risky is it? Could I carry out the reaction in a
regular oven to make say 50 grams or so? What are guanidine salts good for exactly?
BTW I also saw this on google:
Sponsored Links
Sexy Guanidine Singles
Free photos, personals and hot
profiles of local singles. Free
www.infobert.com
Perhaps they should work on the smart add system a bit more
|
|
|
Alchemist
Hazard to Self
Posts: 93
Registered: 22-6-2002
Location: Hostton Texas
Member Is Offline
Mood: No Mood
|
|
guanidine
AS a source of Guanidine, for the Nitrate,
Perchlorate, Picrate, and etc..
Maybe a route to Calcium Cyanamide would be easier.
Has anyone had any success making a far quanity of it?
Thanks....
|
|
|
Rosco Bodine
Banned
Posts: 6370
Registered: 29-9-2004
Member Is Offline
Mood: analytical
|
|
Urea , lime , and charcoal mixtures heated together to a high temperature yield calcium cyanamide . A higher intermediate which may be conveniently
substituted for urea is cyanuric acid , which is sold as a chlorine stabilizer for pools . I have not tried the reaction to
see if this works but it seems logical .
US3499726
US3173755
3 Ca(OH)2 + 2 (HOCN)3 + 6 C ----->
3 CaC(N)2 + 3 CO2 + 6 CO + 6 H2
The powdered lime , cyanuric acid , and charcoal are thoroughly mixed and placed in an iron or steel container with cover .
set over a high output burner and fired to
a low red heat . The progress of the reaction can be followed by the flareoff
of escaping gases . When exiting gases
cease to burn , the reaction is complete
and after cooling the residue should be
a high proportion of calcium cyanamide .
Pretreating the inside surfaces of a cast iron container and cover by painting the surfaces with water glass , sodium silicate , and pre-firing to
bake it on , may be necessary to protect the iron from corrosion if it is a low silica iron .
[Edited on 6-11-2004 by Rosco Bodine]
|
|
|
chemoleo
Biochemicus Energeticus
Posts: 3005
Registered: 23-7-2003
Location: England Germany
Member Is Offline
Mood: crystalline
|
|
Very interesting reaction, Rosco. It should be fairly easy to do, providing one can raise the mix up to 700 or so deg C. A normal coal fire should do
I guess.
Anyway, I found a patent that details the production of guanidine nitrate, from calcium cyanamide, and ammoniumnitrate, and water.
This one, unlike other related patents, has no danger of detonation, which is definitely a bonus. Temperatures are also quite low, <100 deg C.
It's GB507498, see http://v3.espacenet.com/origdoc?DB=EPODOC&IDX=GB507498&a... for details.
So quantum, this may be interesting to try in stead. Calcium cyanamide is also contained in certain fertilisers. Should be trivial just to heat the
fertiliser with ammonium nitrate and water.
Oh, and I found a patent relating to aminoguanidine, which can be obtained straight from calcium cyanamide and hydrazine - not that anyone will be
making this I guess. Nonetheless it's interesting to see how many uses calcium cyanamide has!
Never Stop to Begin, and Never Begin to Stop...
Tolerance is good. But not with the intolerant! (Wilhelm Busch)
|
|
|
Rosco Bodine
Banned
Posts: 6370
Registered: 29-9-2004
Member Is Offline
Mood: analytical
|
|
| Quote: | Originally posted by chemoleo
Very interesting reaction, Rosco. It should be fairly easy to do, providing one can raise the mix up to 700 or so deg C. A normal coal fire should do
I guess.
|
Should be easy enough to tell when you get it hot enough ,
by the flareoff of venting gases . It will be too hot for
aluminum , but stainless may be fine .
It seems to me that anything that would char to carbon at the high temperature
might substitute for charcoal . Perhaps
a carbon yielding equivalent of sawdust ,
starch , or even table sugar would work
as a source of carbon for the reaction .
I think 500 C would do it , but I'm not sure . Substituting Zinc Oxide or possibly
Zinc Carbonate and adjusting the other components if needed should give the
lowest temperature reaction for a cyanamide , Zinc Cyanamide of course being
the product . This should give a
filtered solution of sodium cyanamide after
digesting with sodium carbonate solution .
| Quote: |
Oh, and I found a patent relating to aminoguanidine, which can be obtained straight from calcium cyanamide and hydrazine - not that anyone will be
making this I guess. Nonetheless it's interesting to see how many uses calcium cyanamide has! |
Actually aminoguanidine is my interest for the cyanamide ,
via reaction with hydrazine sulfate .
[Edited on 6-11-2004 by Rosco Bodine]
|
|
|
Rosco Bodine
Banned
Posts: 6370
Registered: 29-9-2004
Member Is Offline
Mood: analytical
|
|
Update :
The reaction of a mixture of calcium hydroxide and cyanuric acid dihydrate and charcoal has been performed and the reaction product is heavily
contaminated with graphite formed from the carbon which was not oxidized . So
an alternative method is being contemplated as a cleaner
and probably more efficient synthesis , which would seem
more suitable as a small scale method for calcium cyanamide .
Calcium cyanurate , CaH(OCN)3 is decomposed by high temperature to calcium cyanamide . So it is logical that
simply forming and isolating the calcium cyanurate as a
first step would simplify a small scale synthesis , which
could then be completed by heating the calcium cyanurate
intermediate to high temperature and collecting the residue
of pure calcium cyanamide .
I have done an experiment to produce the calcium cyanurate
which appears to be successful .
16.5 grams ( .1 mole ) cyanuric acid dihydrate was added to
200 ml stirred hot water . 8 grams of NaOH in 50 ml water
was added and the mixture stirred for 15 minutes just barely at the b.p. . Then 11 grams ( .1 mole CaCl2 ) in 50 ml water
was streamed into the mixture with stirring and heating continued for 30 minutes . Theoretically  the white precipitate is calcium cyanurate in a supernatant solution of sodium chloride byproduct .
The reactions are :
(HOCN)3 + 2 NaOH + CaCl2 ---> CaH(OCN)3 + 2 NaCl + 2 HOH
After filtering and drying , the calcium cyanurate should
yield calcium cyanamide upon heating to decomposition at red orange heat in the absence of air by the following reactions :
CaH(OCN)3 ---> Ca(OCN)2 + HOCN
Ca(OCN)2 ---> CaC(N)2 + CO2
Here the progress of the first half reaction should be observable by the burning of escaping gases from the loosely sealed container in which the
calcium cyanurate is being heated , and it will be the cyanic acid gas which flares
off for this reaction at a lower temperature than the completing second reaction . For the second reaction ,
of course there will be an evolution of carbon dioxide so long as the reaction is proceeding . I think the advantage for
this proposed method , aside from elimination of graphite as a byproduct , is that it involves no gas / solid reactions to
form an intermediate . Such reactions are not generally efficient outside of an industrial process which uses rotary kilns or fluidized bed sorts of
apparatus for the reactions and
microfine powders and long reaction times . Those methods are difficult to implement as a lab scale synthesis , where a more straightforward thermal
decomposition of an intermediate already formed in advance is much simpler .
|
|
|
NJF
Harmless
Posts: 11
Registered: 23-12-2005
Location: England.
Member Is Offline
Mood: No Mood
|
|
The method involving urea and lime proceeds via calcium cyanurate, which (if you can't find cyanuric acid) may be conveniently prepared by heating
lime with urea at ~200-300*C:
Ca(OH)2 + 2 CO(NH2)2 --> Ca(OCN)2 + 2 H2O + 2 NH3
What do you want the aminoguanidine for? I recommend that you try aminotetrazole derivatives....
Formerly \"Nick F.\"
|
|
|
Rosco Bodine
Banned
Posts: 6370
Registered: 29-9-2004
Member Is Offline
Mood: analytical
|
|
The method using urea will not result in as pure a product because of the incomplete reactions leading to the intermediates , and also due to losses
from the moisture
which decomposes the intermediates as a competing reaction .
The aminoguanidine is intended as a starting material for
tetrazole experiments . But the calcium cyanamide is of
course much more general interest and the aminoguanidine
is just one of the many derivatives .
|
|
|
Rosco Bodine
Banned
Posts: 6370
Registered: 29-9-2004
Member Is Offline
Mood: analytical
|
|
A second experiment has been done on a 1 mole scale to
produce the hoped for calcium cyanurate . In a 4 liter beaker containing 3 liters of stirred hot water was added
165 grams cyanuric acid dihydrate . After 10 minutes most
of the cyanuric acid was dissolved and to the stirred suspension was added 80 grams solid NaOH . After another 15 minutes there was added a solution
of 110 grams CaCl2 in 500 ml hot water and the stirring continued for 30 minutes . After standing and cooling overnight the mixture was filtered
through cloth and the
cloth gathered and twisted down to squeeze excess
liquid as much as possible from the solids which compacted to a dense ball of microcrystalline and
slightly damp material which should air dry easily
when broken up .
It would appear that this method is adaptable to use
of soluble zinc and magnesium salts , the sulfates
would probably give the corresponding acid cyanurate via disodium cyanurate . And these cyanurates may then
lead to the respective metal cyanamides by thermal
decomposition in the absence of air .
Looking at the decomposition reactions which are predicted at refractory temperatures for the metal acid cyanurates , it appears to me that these
cyanurates could also have interest for their possible usefulness as a component of sacrificial ( ablative ) heat shielding materials .
|
|
|
Rosco Bodine
Banned
Posts: 6370
Registered: 29-9-2004
Member Is Offline
Mood: analytical
|
|
Quote: Originally posted by chemoleo  | Very interesting reaction, Rosco. It should be fairly easy to do, providing one can raise the mix up to 700 or so deg C. A normal coal fire should do
I guess.
Anyway, I found a patent that details the production of guanidine nitrate, from calcium cyanamide, and ammoniumnitrate, and water.
This one, unlike other related patents, has no danger of detonation, which is definitely a bonus. Temperatures are also quite low, <100 deg C.
It's GB507498, see http://v3.espacenet.com/origdoc?DB=EPODOC&IDX=GB507498&a... for details.
So quantum, this may be interesting to try in stead. Calcium cyanamide is also contained in certain fertilisers. Should be trivial just to heat the
fertiliser with ammonium nitrate and water.
Oh, and I found a patent relating to aminoguanidine, which can be obtained straight from calcium cyanamide and hydrazine - not that anyone will be
making this I guess. Nonetheless it's interesting to see how many uses calcium cyanamide has! |
GB507498 is also interesting because reportedly Disodium Cyanamide can be made from fusion of Sodium Cyanate and Sodium Hydroxide. What temperature
is required for the reaction is unknown. If this happens at a lower temperature
than the red heat required for production of Calcium Cyanamide, the process would be greatly simplified.
The Drechsel reaction referenced by Thorpe is
3 NaOCN + 3 NaOH ------> Na2CN2 + 2 Na2CO3 + NH3
Drechsel article attached. Please anyone who understands the German offer your insight here regarding the reaction conditions for the fusion of
cyanate and hydroxide. Page 13 of the pdf is page 89 of the article where the reaction is described although the reaction equation is incorrect in
the article and is shown corrected above to account for the evolution of NH3 reported by others.
Attachment: php019VGM.pdf (1.9MB)
This file has been downloaded 879 times
See related thread
http://www.sciencemadness.org/talk/viewthread.php?tid=14267&...
[Edited on 12-6-2011 by Rosco Bodine]
|
|
|
![[*]](images/xpblue/default_icon.gif){kind=icon}
