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Theoretic
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Does anybody know of electrochemical preparation af cyanides, e.g.,elctrolysing NaNO3 with a carbon anode or something like that?
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blip
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<a href="http://www.nidlink.com/~jfromm/chem201/chem204l.htm" target="_blank">NaNH<sub>2</sub> + C
<sup><u> <font face="symbol">D</font> </u></sup>> NaCN +
H<sub>2</sub></a>
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I am a fish
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| Quote: | Originally posted by Samosa
Well, the reaction between Chloroform and Ammonia is a well-known way to produce Hydrogen Cyanide...
NH3 + CHCl3 --> HCN + 3 HCl
If you don't mind working with straight HCN, then this seems like a feasible route. At least, it would be for me, since I have fairly easy
access to Ammonia Gas and Chloroform...
The only trouble would be separating the HCl. |
If you're prepared to work with HCN, it would be preferable to react potassium ferricyanide (available from photographic suppliers) with a strong
acid. Unlike ammonia and chloroform, the product wouldn't be mixed with HCl.
1f `/0u (4|\\| |234d 7|-|15, `/0u |234||`/ |\\|33d 70 937 0u7 /\\/\\0|23.
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DDTea
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That reaction to produce HCN also works by reacting Potassium Ferrocyanide and a strong acid, correct?
If not, would it be possible to convert the Ferrocyanide into the Ferricyanide?
Also, I do not like to order from suppliers...it seems like "cheating" to me; the struggle to obtain chemicals is part of the fun for me
 .
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blip
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| Quote: | | NH3 + CHCl3 --> HCN + 3 HCl |
Isn't there a way to get HCN to become gaseous? I thought it would come out of solution on it's own maybe at some good concentration, then
collect it in another vessel with more water to dissolve it in if needed.
Here's a reaction posted on this site somewhere if you need chloroform and have acetone and bleach:
CH<sub>3</sub>COCH<sub>3</sub> + 6 NaOCl <strike> </strike>> CHCl<sub>3</sub> +
NaCH<sub>3</sub>COO + 2 NaOH + 3 NaCl
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I am a fish
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| Quote: | Originally posted by Samosa
That reaction to produce HCN also works by reacting Potassium Ferrocyanide and a strong acid, correct?
If not, would it be possible to convert the Ferrocyanide into the Ferricyanide?
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It will work with either. However potassium ferricyanide is cheaper and contains more cyanide ions per unit mass.
1f `/0u (4|\\| |234d 7|-|15, `/0u |234||`/ |\\|33d 70 937 0u7 /\\/\\0|23.
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chloric1
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others?
Have yet to try the charcoal, cyanuric reduction but what about exposing solutions of thiocyanates to sunlight over months to precipitate sulfur?
Thiocyanates are acteone soluble while cyanides are not! Good way to separate for purification.
[Edited on 11/30/2003 by chloric1]
Fellow molecular manipulator
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unionised
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I'm a little suprised that nobody has tried the "grow your own" approach to cyanide based on things like this.
http://www.coopsjokes.com/amz/amzmisc4.htm
And there are much better sources too.
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chloric1
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Yes I am aware of this
This is agood suggestion but you need a lot of pips to extract reagent quanties of cyanide. I am going to look for a patent about electrolytic
oxidation of thiocyanates to cyanides and sulfides.
[Edited on 11/30/2003 by chloric1]
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chloric1
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Here it is!
Ok searching my archives I found it! Here it is in Word format.
Attachment: US Patent 4,519,880.doc (78kB)
This file has been downloaded 2394 times
Fellow molecular manipulator
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fritz
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I did some research in my books and notices and will now share the results with you:
Preparation of potassiumferrocyanate:
in my old Beilstein there is a notice that it had been produced by melting together animal related substances(???) with K2CO3 and Fe. The mass was
extracted with water and the K-ferrocyanated was recrystallized. I found also an experiment about the historical production of this substance and
there they melt equal parts of K2CO3, Fe and Urotropine (or urea or soj-flour) in this reaction ammonia is produced and the whole mass becomes more or
less black. reading this reminds me of a test for nitrogen-containing organic compounds (Lassaigne-test): there you melt Na (Hm, and here we got a new
problem! Could be perhaps replaced by Li from Li-batteries) with your organic compound and if it has nitrogen in it your melt contains NaCN (which
could be reacted with Fe(III)/Fe(II) ....)
Unfortunatelly my only attempt on making KCN (by melting K-ferrocyanide) failed. But if I would try again I would produce HCN by boiling 200g
K-ferrocyanide with a cold mixture of 160g sulfuric acid and 250g water. the HCN is dried by bubbling through 2n sulfuric acid and twice through CaCl2
( both warmed at 40°C). The dry HCN is introduced in a solution of one part KOH in three parts EtOH. The KCN will separate and is immediatelly
filtered off, washed with EtOH and dried over sulfuric acid. This should be more or less safe since the poisonous HCN is kept in a closed apparatus.
May be that the KOH solution should be used in great excess to avoid dilution of the EtOH by the water wich is build in the neutralisation-process. I
don´t know the solubility of KCN in diluted EtOH (it should be nearly insoluble in absolute EtOH)
When I tried to make KCN I tested the purity by precipiating AgCN with Ag-nitrate. But I think a better way would be titration with o,1n Ag-nitrate
solution: Some KCN is (exactly!) weighted and solved in water. Some KOH is added (pH should be slighly >7) and Ag-Nitrate solution is added until
the solution becomes to be cloudy (1ml o,1n Ag-nitrate sln. = 13,024mg KCN) It could also be titrated with K-chromate as indicator (like chloride)
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unionised
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The first drop of Ag+ solution will preciptate AgCN and go cloudy. I'm not sure if the chromate would work but I think it would (anybody got the
solubillity data to hand?)
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vulture
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That Ag test is tricky, because excess CN will dissolve Ag as Ag(CN)2 -
One shouldn't accept or resort to the mutilation of science to appease the mentally impaired.
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fritz
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No, there would not be a precipiate when you add Ag+ to CN- if there is a great excess of CN- (which you will definetivly have if you titrate a CN-
sln. with Ag-nitrate) because Ag+ reacts with excess of CN- to a complex anion ( [Ag(CN)2]- ) which is soluble. and this vulture is the whole trick at
this titration. You add Ag+ until every CN- ion is bonded in the complex. When there is more Ag+ in the solution the AgCN will precipiate because the
complex is destroyed. and this is the end-point of the titration. So one mole Ag+ is equivalent to two moles KCN ( or one ml 0,1n Ag-nitrate sln.
(=0,1mmol) is equivalent to o,2mmol KCN (= 13,024mg))
Solubility of AgCN at 20°C is 0,023mg per 100ml water. The solubility of Ag-chromate is 1,4mg per 100ml water at 0°C.
So when all CN- is bonded in the complex the red Ag-chromate will precipiate, which could be really easy realized.
(I did this titration with chloride and it did very well!!)
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unionised
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I hadn't realised the formation constant for the complex was high enough for you to be able to do the CN- titration that way.
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Jay Maity
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Another Method
Keeping nitrogen in very low pressure and creating electric spark we can get active nitrogen.
This nitrogen if mixed with Methane we can get Hydrocyanic acid and Hydrogen .
This gas mixture if passes through water HCN will solute.
This HCN can react with KOH and We can get KCN.
2CH4 + 2N = 2HCN + 3H2
HCN + KOH = KCN + H2O
[Edited on 4-1-2004 by Jay Maity]
jaymaity
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DDTea
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If you have a vacuum aparatus, this seems like it's the way to go forcontinuous generation of HCN. However, instead of Methane, I would use
Acetylene, because the reaction is much cleaner (no side products):
C2H2 + N2 --> 2 HCN
This gets me thinking, though-- could activated Nitrogen react with Carbide salts to form the corresponding Cyanide salt? e.g.:
K2C2 + N2 --> 2 KCN
Of course, the Carbide salt most easily available is Calcium Carbide, and I don't know how useful Calcium Cyanide would be.
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DDTea
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Loosely related, but Interesting
To keep predators away, Millipedes secrete Hydrogen Cyanide . I did not know
this, and just learned it today in my Chemistry book. But on further investigation, I came upon this excerpt from a Yahoo forum:
| Quote: | Dear Friends and Colleagues:
The purpose of the following account is to provide my personal experience
with cyanide gas, and what it did to me. This is just my anecdotal report,
and should not be used for planning for the results of the release of
highly-concentrated clouds of cyanide gas (especially potassium, sodium or
hydrogen cyanide).
I taught biology and physiology at Central High School, in Little Rock,
Arkansas in 1975-76. I often collected animal and plant samples to enhance
the classroom training (since I couldn't get funding for the samples). On
one late fall day I collected some 300 large millipedes (about the size of
an adult thumb) and packed them into a plastic container. I then stored the
container in a freezer to keep the millipedes from destroying each other in
a closed environment. Millipedes can recover after exposure to severe cold,
so this seemed a reasonable process. Four hours later I took the container
from the freezer and opened the lid wide, looking inside to see if the idea
had worked. I woke up the next morning, finding the millipedes lying all
over the floor...they were all dead. The refrigerator door was still open.
My head was very sore from where it had hit the floor. Why were the
millipedes all dead, and why was I on the floor?
What I had forgotten during the collection project (although I knew the
fact) was that some millipedes produce a small drop of cyanide (or other
chemicals) when they curl up. This defensive action makes them very
distasteful to predators. What I did not know was that piling hundreds into
a closed container, and then putting them in the high stress state being
frozen, would produce huge volumes of this liquid. The liquids usually
volatilize quickly in nature. In the plastic container they were trapped
and concentrated.
I never smelled almond. I never had a chance to react. My collapse was
simultaneous with opening the lid.
That gives you a sense of how little time a responder might have in a
concentrated cloud of cyanide gas.
An excerpt of a discussion about a millipede's chemical capacity is
provided below.
Sincerely,
Rick Tobin,CEM
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Among other things, Millipedes secrete Benzaldehyde, Acetic Acid, Chlorine and Iodine.
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Rosco Bodine
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| Quote: | Originally posted by madscientist
For preparing HCN from formamide, I recommend heating it in a flask, which sends the vapors down through a glass tube into another borosilicate glass
flask (which is being heated by intense flame); vapors from that flask then should be composed of HCN and water. |
That setup sounds reasonable as a practical method for for cyanide . The thermal cracking of formamide vapor could
be done by a "ketene lamp" tube , a long
coil of nichrome heating element within a
glass tube . Possibly a water heater heating element with a pipe threaded mounting flange could be mounted in a short length of pipe , having vapor
inlet
at one end and an outlet near the other
end , as an improvised sort of internally
heated "tube furnace" . This could be wrapped with fiberglass batting to help
prevent heat loss . The exiting vapors
could be bubbled through a hydroxide solution to produce the desired salt .
A coil of stainless steel tubing heated
by a burner would probably work also for the thermal cracking of formamide vapor .
With regards to sodium formate , a method I saw reported is that formaldehyde in alkaline solution with hydrogen peroxide , when heated produces the
alkali formate . This would
seem ideal for use with paraformaldehyde
which is depolymerized by alkali in warm
aqueous solutions .
With regard to cyanates as an intermediate for cyanide , it is the conditions of temperature and possibly
the atmosphere and catalytic impurities ,
as well as the particular metal cyanate being decomposed , which determines
whether there will be produced a cyanide
or a cyanamide . The presence of added
materials can also favor the formation of
one product or the other as the end product , as the cyanamide itself can be a
second intermediate for cyanide . There
is a method of producing cyanates under
relatively mild conditions , from a carbonate and urea , using DMSO as a solvent . The cyanate is formed in pure
condition and high yield , and the DMSO is
recyclable . See the patent for details .
US3167387
I have been looking at a possible route for
Zinc Cyanamide , since it reportedly forms
at lower temperature than is required for
Calcium Cyanamide . Zinc Sulphate is
cheaply available as a garden supply item , and its solution mixed with a solution of baking soda , precipitates Zinc Carbonate in fine crystalline
form . This
may be converted to the cyanate by the
method of the patent . At about 500 C
the Zinc Cyanate should decompose with
evolution of carbon dioxide , to leave a
residue of Zinc Cyanamide . This material
should be a good precursor for aminoguanidine , and dicyandiamide ,
and even for cyanide , with the zinc recoverable from the process .
To produce a cyanide from a cyanamide ,
there is added a carbonate and free carbon and the mixture is fired to a much higher temperature , above 600 C to produce a fusion of the mixture .
See
GB437614
There are different fusion mixtures and methods which are probably easier for cyanides , but one or another method may
be preferable for the economy or purity
of the product .
[Edited on 13-11-2004 by Rosco Bodine]
[Edited on 13-11-2004 by Rosco Bodine]
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Eclectic
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Text from U2U with Polverone's blessing:
Polverone
To: Eclectic
Sent: 15-11-2004 at 05:40 PM
Message:
quote:
I meant ferrocyanide. I was just thinking a one pot (dutch oven) reaction might be a simpler route to cyanides. Maybe as simple as cooking the hell
out of iron turnings, urea, and possibly sodium carbonate and ending up with prussian blue or a ferrocyanide.
If you cook the hell out of it, you will get cyanides, maybe some carbonate, and metallic iron. If you cook not quite so hard, you may indeed end up
with ferrocyanides. Actually, the historical literature is ambiguous about this, indicating that ferrocyanides may be formed from cyanides and iron
when the product of fusion is treated with water, but I never saw a modern, comprehensive treatment of this old industrial process. don't think
you will end up with prussian blue this way under either condition, though I did note the formation of some prussian blue in the neck of my test tube,
presumably from the interaction of iron oxalate that stuck to the glass and cyanide offgases.
Iron turnings was a traditional source of iron for ferrocyanide production (along with the sacrifice of the iron chambers themselves), but that was
chosen for economy, not convenience. I think an iron salt of an organic acid would provide iron in a more rapidly reactive form, and my test tube
experiments seem to support that view.
It's an interesting question to what extent my charcoal-fired steel cans were damaged by reactions leading to ferrocyanide and to what extent it
was simply attack by atmospheric oxygen. I would like to someday try this reaction with even better air exclusion, perhaps using an empty disposable
propane cylinder or other narrow-necked container.
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Eclectic
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Earlier Text from U2U
Polverone
To: Eclectic
Sent: 15-11-2004 at 05:15 PM
Message: Oh, you should feel free to ask this in the thread. We don't mind old threads being revived.
First, you will not directly form ferricyanides this way. You can form ferrocyanides with iron, but to get ferricyanides the ferrocyanides need to be
oxidized (chlorine is the traditional oxidizer used).
The other day, I did try making ferrocyanides on a small scale. I did it by mixing cyanuric acid, potassium carbonate, and iron oxalate and fusing
them all together in a pyrex test tube over a propane flame. Initially, a rather dark reddish-orange melt is formed. The color lightens after stronger
heating, and you obtain a yellow substance that looks much like the purchased sodium ferrocyanide that I have on hand. I presume that it was (impure)
potassium ferrocyanide. Stronger/longer heating destroys the color, and free iron metal appears as gray particles. This means that the ferrocyanide
has been decomposed, leaving cyanide and iron. Iron oxalate decomposition may deposit some free iron even if the oxalate is not in excess, but the
iron mostly seems to go into solution.
If you want pure ferrocyanides, you might be better off first making a cyanide (which requires only high temperatures, not careful temperature
control) and then reacting it with a mixture of iron II/III compounds to form a precipitate of insoluble Prussian Blue. The Prussian Blue can then be
reacted with KOH or NaOH to form the respective ferrocyanides, which can then be oxidized with chlorine to ferricyanides.
I would say that it is simple and easy to make a small quantity of cyanide from cyanuric acid or urea and carbonates in a test tube with propane flame
heating, except for one thing: the mixture froths and foams considerably in the early stages of strong heating, so that you can only start with a very
small amount of reactants without the tube overflowing. The final amount of cyanide produced is only half a gram or so when working at this scale.
Actually, my problems may be due to using such narrow test tubes. I know for sure that larger quantities can be produced in a steel can in a charcoal
fire, as documented in the cyanides thread, but the test tube reaction can be easily done indoors.
quote:
I'd ask in the forum, but the thread is pretty old:
Did you try reducing cyanates with iron turnings to make ferricyanides?
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Idunno
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Formulas
Can someone post the formulas for these reactions:
1) Potassium Ferricyanide and Sulfuric acid and heat,
and
2) HCN and NaOH
Has anyone tried this setup? (Outside I'd assume.) Any advice?
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Chris The Great
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An attempt was made to prepare sodium cyanide from urea, charcoal and sodium hydroxide, by the following reactions (this was attempted before in the
thread, but no reactions where posted):
OC(NH<sub>2</sub> <sub>2</sub> --> HOCN +
NH<sub>3</sub>
HOCN + NaOH --> NaOCN + H<sub>2</sub>O
NaOCN + C --> NaCN + CO
The pictures aren't really that important, and somewhat repetitive, so if you have dial up there isn't much reason to look at them, in my
opinion.
Charcoal was from an old fire. It was crushed as well as possible in a plastic bag with a hammer, although it was moist and still had quite a few
chunks in it. It was heated in a pot on low heat until it became free-flowing. It had a very fine dust, that got everywhere, and larger chunks which
I couldn't get rid of no matter how many times I ground it with the hammer. This was used in the following procedure.
40 grams of sodium hydroxide, 20 grams of charcoal and 65 grams of urea where measured out and added to a small stainless steel pot. They where
quickly mixed, but because the bits where all different sizes they didn't mix well at all. The pot was slowly heated with a butane burner. The
urea soon began to melt, and give off ammonia gas. The mixture rapidly turned into a liquid, which began to foam from the ammonia gas being produced.
The pot was often removed from the heat to prevent the reaction from proceeding to fast, as it continued even without a heat source. Stirring was
essential to keep it from foaming over. It was also left for a short time as I left to get my camera.
http://img280.echo.cx/my.php?image=liquid9uc.jpg
After another few minutes of stirring, the mixture rapidly solidified into large chunks of hard foam. Ammonia gas was still being evolved in large
amounts, and the constantly shifting breeze meant I was constantly getting gassed in the face with it. It was very unpleasant.
Stirring and heating was continued and the large chunks slowly broke down, all the while generating lots of ammonia.
http://img280.echo.cx/my.php?image=solid2pc.jpg
The colour gradually lightened to a light grey colour. Heating was stopped for a while while I crushed them up as well as I could with a hammer, and
refilled my butane cylinder as it ran out.
http://img280.echo.cx/my.php?image=solid26sz.jpg
The burner was then turned on full and stirring was continued. Ammonia slowly stopped being evolved. When the bottom of the pot was a dull red heat
(as far as I could tell, it was a sunny day outside), the mixture that was directly in contact with the metal started to melt, and occasionally a
large piece of charcoal would ignite and become an ember for a few seconds. The mixture slowly started to darken again, and then my burner ran out of
fuel. I burned the wooden table instantly with the bottom of my pot, but luckily I had a clay piece to rest it on while I refueled the cylinder with
a bit more butane.
http://img280.echo.cx/my.php?image=solid30jn.jpg
Heating on full was resumed, and there was no smell of ammonia at all. After a while it got back up to reddish heat, and the mixture in contact
started to melt. It was very slippery towards my copper stir stick. The mixture slowly changed from the dark grey to a black. Heating was continued
with stirring for a while, and then I left it on full heat and had something to eat.
http://img280.echo.cx/my.php?image=solid40oo.jpg
The burner ran out of fuel, and when I returned some time later the pot was only warm to the touch. There where several pieces of caked material on
the bottom of the pot, that was a light grey underneath the charcoal on the surface.
http://img280.echo.cx/my.php?image=result5gu.jpg
A small bit was added to some water, and it became very slippery, indicating that it was basic. A small amount of
H<sub>2</sub>SO<sub>4</sub> caused a small amount of bubbling. A piece of the caked lighter material was added, and this also
caused bubbling. I was unable to smell anything, probably because all the ammonia had decided my sense of smell would not work for the rest of the
day.
It does not appear to have formed much cyanide, although that's what I assume the bubbling to be as cyanic acid is a solid, hydrogen cyanide is a
gas. Nothing else in there should be able to bubble. However, it appears to be very little cyanide, because my burner isn't hot enough. Since
sodium cyanate has a mp of 550*C, that means I only got up to 550*C, as it was just melting. I will try heating the mixture more with a
propane torch, which will get it much hotter, hopefully enough to convert it all to cyanide. In that case, I should get about 49 grams sodium
cyanide, assuming 100% yield.
Here is a thread at SM on the topic, which is very informative.
https://sciencemadness.org/talk/viewthread.php?tid=23
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Polverone
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Try the Prussian blue test. Mix some of your crude material with a slightly acidified mixture containing Fe (II) and Fe (III) and look for a blue
color. Seeing blue means that you have formed some cyanide, but being sure that all cyanate is gone is another matter.
I think you will have a hard time heating a container like that pot to the needed temperatures with an ordinary propane torch. You need a furnace or a
wood/charcoal fire. If you want to make a reference sample to compare larger batches with, do it on a test tube scale, since your torch will easily
raise the test tube to the necessary temperatures and hold it there.
Do not be so sure that the bubbling you saw was HCN. Cyanic acid is easily hydrolyzed under acid conditions to ammonia and carbon dioxide. It's
also hydrolyzed in neutral water, though not as fast. Boil some of your crude mixture with water. Do you smell ammonia? If so, the conversion to
cyanide was incomplete.
Another simple test that can at least verify that most of your sample is cyanide is to make a strong solution of filtrate from your crude mixture,
then chill it and add some citric acid. If it bubbles immediately, you know you have a substantial contamination with carbonate. It should bubble off
some HCN as it is warmed, though.
So, for triple verification:
-Perform Prussian blue test (to confirm presence of cyanide)
-Perform hot water test (to confirm absence of cyanate)
-Perform strong, chilled solution acidification test (to confirm absence of carbonate)
[Edited on 5-31-2005 by Polverone]
PGP Key and corresponding e-mail address
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Polverone
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more CN- stuff
I know I've told some people about it, but I can't seem to find my more refined NaCN preparation writeup here.
The way I do it now, for making larger quantities, is to mix powdered urea or cyanuric acid with NaHCO3 and heat, covered, until it's melted.
Then I add the charcoal and raise the heat. Pre-mixing the charcoal seems to make it take longer, perhaps because it makes the mixture more
insulating/less fusible. Getting the initial large mass of powder to melt down is, IMHO, the most annoying part of it all and the limit to scaling up
further. A large volume of powder melts down to a much smaller volume of liquid, so you have to start with a large container for a decent batch size.
You need to hold the mixture at a dull red heat for some time to make the reaction complete. Heating it hotter will complete things faster. In a
larger batch, you should see/hear the vigorous popping/burning of sodium-tinged carbon monoxide bubbles when it's reacting at its most vigorous
in the open air. Of course it's my opinion that I should shield the reaction vessel (a can) with another can to keep too much oxygen from getting
to my hot mixture. I don't want to be oxidizing cyanides right back to cyanates.
The biggest improvement on the workup of the crude mixture is to make a strong solution of cyanide by swirling your broken-up charcoal mass in hot
water (smash up the glassy mass of charcoal with a hammer in some bags before soaking), and then instead of evaporating the filtrate, pouring it into
an excess of denatured alcohol with stirring. This will precipitate most of the cyanide. It will also leave you with a nice free-flowing mixture (when
dry) instead of the rocky masses you get at the bottom of a heated evaporation dish. Much of the water is carried away by the alcohol, and the
precipitate is faster to dry.
I have been told by others, though I have not confirmed it myself, that the reaction goes faster and easier if you start with potassium compounds
instead of sodium, perhaps due to the lower melting points involved.
You can easily produce what seem to be mixtures of cyanides and ferrocyanides by heating together urea, alkali carbonate, and ferric or ferrous
oxalate. The melts go through a variety of colors, starting near orange and going through red to pale yellow and finally colorless (when heated
strongly, also yielding metallic iron). The vapors from the reaction are evidently cyanide-bearing to some extent as they may turn traces of iron
oxalate clinging to the reaction vessel walls a Prussian blue.
The reaction of chloroform with ammonia in presence of alkali is simpler in that it doesn't need a lot of heat, but it's also harder to get
pure cyanide from it. The reaction with 5% aqueous household ammonia is sluggish to nonexistent, while the reaction with 28% aqueous ammonia was
almost violent. Ethanol may be needed to improve mutual solubility of aqueous ammonia and chloroform.
If you want extra-pure cyanides, Muspratt offers a helpful hint: dissolve KOH in ethanol, then introduce HCN to the solution. KCN precipitates in pure
form. The HCN could come from an acidified carbonate-free batch of KCN or NaCN formed by the chloroform-ammonia interaction or a high temperature
method. Of course this seems highly hazardous and unnecessary for most purposes.
You can experience the joys and wonders of the various cyanide-production methods on a test tube scale first to get a feel for them, before trying to
make dozens to hundreds of grams at a time.
PGP Key and corresponding e-mail address
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