Sciencemadness Discussion Board

Preparation of cyanides

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Polverone - 21-5-2002 at 09:03

I have the Poor Man's James Bond (all volumes) and one of them has instructions for preparation of cyanides. The preparation instructions basically go something like: heat charcoal and potassium/sodium carbonate in a steel vessel with iron filings/turnings at high heat, overnight (in a homemade furnace) to obtain sodium/potassium ferrocyanide (which one would separate from wastes by dissolving/filtering). Then heat your newly created ferrocyanide with more carbonate, again in a steel vessel in your homemade furnace, to obtain the cyanide salt. Pour off the fused cyanide salt from the vessel onto something like a slab of marble to cool it, then break it up and store it.
HOWEVER, there seem to be a few problems with this procedure. The Kirk Othmer Encyclopedia mentions that cyanides rapidly oxidize to cyanates when heated in the presence of air and iron(!) So why would the instructions say to use a steel vessel and not mention any measures to isolate from the air... Hmmm. Second, the original sources on 19th century chemical production that I have access to (coming soon to a website near you!) specifically use nitrogen-rich organic materials with the carbonates, not plain old charcoal, and do so in the absence of air. So, has anyone tried the procedure from the PMJB? I made a go at it one afternoon, but only heated the mix for an hour or so, with a large gas burner, and did not obtain anything resembling ferrocyanide. I will make an attempt using the original 19th century procedures once I have enough free time to cobble together a little charcoal furnace. I wonder: is Kurt "maimed myself with Armstrong's mixture because I didn't read the directions" Saxon mistaken in his procedure? I am inclined to think so, especially since he says 50 mg of KCN will kill a man (a very optimistic statement). But maybe had I followed his instructions to the letter it would have worked. Anyone out there with further comments (or better yet, experience)?

Polverone - 21-5-2002 at 09:03

YIKES! I found the following information on another site:

a patent on making metal cyanides from nitrates or nitrites and
carbon; US patent 579988.
KNO3 + 4C -> KCN + 3CO
KNO2 + 3C -> KCN + 2CO

I was unable to access the patent since I'm temporarily banned from the database for running too many queries (oops).

So I decided to try just forming a pyrotechnic mixture with the right ratios. 10 grams KNO3, 4.8 of charcoal, place in stainless steel vessel and ignite with gas heating from below...

As expected, the mass of what remained was much reduced, from loss of gas, solid particulates, things flung from the vessel by the reaction, etc. There was little material left in the bottom. I figured there had to be more to the method than this; after all, nobody talks about pyrotechnic formulas leaving cyanide lying around, and this is pretty much the same thing.

Anyway, not having an analytical method for detecting cyanides at hand and being too stupid to look one up (and also expecting failure), I added a bit of vinegar to the residue left in the bottom. It fizzed vigorously and I caught the distinct odor of almonds... At which point I backed the heck away from there. I now intend to find a method for assaying KCN that is not so suicidal, and also to try making some more and purifying it (I have no idea what purity I obtained with this first test.) This method seems to be a vastly superior route to cyanides for the home experimenter, compared to the laborious steps given in the PMJB and the 19th century texts from which they were derived. I hope to view that patent soon and see if it contains any additional refinements (compared to crude ignition).

madscientist - 21-5-2002 at 09:04


-all potassium chemicals can be substituted with their sodium parallel, if mass ratios have been properly adjusted
-it is highly recommended that nbk2000 dismiss all described and inferred safety precautions

126.4 grams of potassium permanganate (KMnO4) is added to 32 grams (approximately 40.2mL) of concentrated methanol (CH3OH):

10(CH3OH) + 8(KMnO4) --} 10(HCOOH) + 10(H2O) + 8(MnO) + 4(K2O)
10(HCOOH) + 10(H2O) + 8(MnO) + 4(K2O) --} 10(HCOOH) + 8(KOH) + 8(MnO) + 6(H2O)
10(HCOOH) + 8(KOH) + 8(MnO) + 6(H2O) --} 8(HCOOK) + 2(HCOOH) + 8(MnO) + 14(H2O)
8(HCOOK) + 2(HCOOH) + 8(MnO) + 14(H2O) --} 8(HCOOK) + Mn(HCOO)2 + 7(MnO) + 15(H2O)

Mixture is then filtered to remove the manganese oxide (MnO), and the filtered solution is then allowed to evaporate. What is left is a ratio of eight : one of potassium formate : manganese formate. The remaining crystals should weight approximately 81.77 grams if you acheived a 100% yield.

The mixture of potassium formate and manganese formate is added to concentrated sulfuric acid. That is, all 81.77 grams of the potassium formate and manganese formate crystals are added to 49 grams (26.5mL) of concentrated sulfuric acid. The remaining mixture is heated, and the vapors, which are composed of formic acid, are condensed. WARNING! FORMIC ACID IS TOXIC. PURE FORMIC ACID IS A COLORLESS FUMING LIQUID WITH A PUNGENT ODOUR; IT IRRITATES THE MUCOUS MEMBRANES AND BLISTERS THE SKIN.

8(HCOOK) + Mn(HCOO)2 + 5(H2SO4) --} 10(HCOOH) + 4(K2SO4) + MnSO4

Formic acid is added to an aqueous solution of ammonia ( [NH4+][OH-] ). The remaining solution is evaporated; the crystals left are ammonium formate crystals. Crystals should weight about 64 grams if you have been achieving 100% yields.

HCOOH + [NH4+][OH-] --} [HCOO-][NH4+] + H2O

The ammonium formate crystals are heated by flame in an environment containing as little oxygen gas as possible. The ammonium formate decomposes into formamide (HCONH2) which then decomposes into hydrogen cyanide.

[HCOO-][NH4+] --} HCONH2 + H2O
HCONH2 + H2O --} HCN + 2(H2O)

The gas given off is condensed in in a rubber, plastic, or, preferrably, glass tube that has one end immersed in a beaker containing a solution of potassium hydroxide (KOH). The tube should be positioned so that any liquids forming in it will run off into the beaker of potassium hydroxide. Some of the gas given off may not be condensed; that is why the tube is immersed in the beaker of potassium hydroxide. That will prevent a loss of much cyanide. The hydrogen cyanide will quickly react with the potassium hydroxide to form potassium cyanide. The hydrogen cyanide is reacted with the potassium hydroxide because the hydrogen cyanide will evaporate off quickly, which is both extremely dangerous and will cause the loss of a lot of cyanide. About 56.1 grams of potassium hydroxide should be used if 100% yields are expected. About 65.1 grams of potassium cyanide should result if 100% yields are achieved. The solution in the beaker, once all of the ammonium formate crystals have been converted into various gasses, should be evaporated off. The remaining crystals are potassium cyanide crystals.

HCN + KOH --} KCN + H2O

The potassium cyanide is then treated with an acid. This will form the potassium salt of the acid, and hydrogen cyanide. DO NOT ATTEMPT TO STORE HYDROGEN CYANIDE! IT WILL ALMOST CERTAINLY CAUSE THE DEATH OF AN UNINTENDED VICTIM SUCH AS YOURSELF! HYDROGEN CYANIDE SHOULD ALWAYS BE USED IMMEDIATELY AFTER IT IS MADE, OR CONVERTED IMMEDIATELY INTO POTASSIUM CYANIDE! It is recommended to use an acid that can be found concentrated. Concentrated sulfuric acid is believed to be the best acid to use.

KCN + [H+] --} HCN + [K+]

Polverone - 21-5-2002 at 09:05

HCN is easy... Hydrogen cyanide can easily be prepared by warming acidified potassium ferrocyanide. Potassium ferrocyanide can easily be purchased with little or no suspicion. It can be made, as well, but it requires a significant amount of time at elevated temperatures. So the whole involved process of producing formates and decomposing them is not necessary (although interesting.)

Oh, BTW, that patent I mentioned in the 2nd post in this thread? It involves using electrified carbon rods in molten KNO3/KNO2, so it's still not the easiest thing ever...

The "easiest thing ever" that I have found, from my good friend the Hive, is that when certain chlorine-containing solvents are gently heated and stirred for a long time with a mixture of aqueous ammonia and sodium (or potassium) hydroxide, they will form NaCl or KCl and NaCN or KCN (look it up for balanced equations and specific directions.) Sadly, this leaves you with a mixture of salts, and I am obsessed with purity. If I were to prepare large quantities perhaps I could separate the salts by recrystallization, but that sounds hazardous (because of larger quantities.)

Speaking of hazardous, I really don't think that any method which involves HCN gas is suitable for home preparation of cyanide salts. Not unless you have a really good fume hood, which no house I've ever seen does. Cyanides are a real PITA. You certainly don't need sulfuric acid to decompose them. Virtually any acid will work. HCN is a very weak acid, and its corresponding salts are of course strong bases. Atmospheric CO2 will liberate HCN from the damp salts. Strong acids are overkill.

Have you tried all or part of what you've written up? Apart from the minimal safety instructions about handling HCN, the formate preparation seems a little odd... Specifically, you're adding a considerable amount of potent oxidizer to a considerable amount of flammable liquid, with no mention of any cooling precautions or predictions as to how long the reaction takes to complete...

Polverone - 21-5-2002 at 09:06

Experimentation and private communication revealed a few things: you can't heat ammonium formate to make HCN, and if you're making formic acid you should dilute and/or cool the methanol/permanganate mixture unless you WANT it to boil.

madscientist - 21-5-2002 at 09:08

Yes, that is true. I realized that ammonium oxalate would probably decompose into ammonia and oxalic acid, rather than into oxamide and then cynogen. I double-checked this hypothesis by heating around 10g (COONH4)2 outside in a glass beaker with my propane burner. It decomposed as follows...
(COONH4)2*2H2O --> (COONH4)2 + 2H2O
(COONH4)2 + 2H2O --> (COOH2)2 + 2H2O + 2NH3

With extensive heating the oxalic acid melts, then decomposes into carbon dioxide, carbon monoxide, and water vapor. This information means that heating ammonium formate will not form formamide, HCONH2, which I know decomposes into HCN when heated. Formamide can be prepared via a different method. I have not attempted to prepare it; I have prepared some oxamide; the process for preparing formamide supposedly is similar. I first prepared ethyl oxalate by mixing the proportional amount of ethanol / oxalic acid, adding a small amount of concentrated sulfuric acid, and heating gently. It soon esterified, resulting in the oily liquid, ethyl oxalate. Ethyl oxalate slowly reacts with water.

(COOCH2CH3)2 + 2H2O --> (COOH)2 + 2CH3CH2OH

Reaction of ethyl oxalate with aqueous ammonia forms oxamide, (CONH2)2.

(COOCH2CH3)2 + 2NH3 --> (CONH2)2 + 2CH3CH2OH

Oxamide is not water soluble and can easily be filtered.

The process for preparing formamide should be similar to this outlined process for preparing oxamide.

Information on formamide from my chemical dictionary:

formamide (methanamide) HCONH2
Properties: Clear, colorless, hygroscopic oily liquid; sp. gr. 1.146; b. p. 200-212 C with partial decomposition beginning about 180 C; m. p. 2.5 C. Soluble in water and alcohol.
Derivation: By the interaction of ethyl formate and ammonia, with subsequent distillation.
Method of purification: Rectification
Uses: Exceptionally good solvent, softener, intermediate in organic synthesis.

Formamide's ability to dissolve in water, and the fact that it doesn't decompose readily into HCN before it boils presents some interesting challenges. It probably could be more easily purified for the home chemist by absorbing some of the water with MgSO4 (not the hydrate!), and / or letting most of the formamide / water / ethanol solution evaporate (assuming you prepare formamide in a manner similar to how I prepared oxamide). 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.

PHILOU Zrealone - 21-5-2002 at 09:09

True that the reaction you mention about carboacid ammonium salt dehydration will lead to amides!
CH3-CO2NH4 -heat-> CH3-CO-NH2
HCO2NH4 -heat-> H-CO-NH2
(NH4)2CO3 -heat-> 2NH3 + CO2 + H2O and no urea!
Further dehydration to cyano/nitriles compounds is hard and requires Acetic anhydride or dry 100% P2O5!

Esters solvolyse by dry NH3 is also a good way to get amides (but no cyano compounds):
CH3-CO2-CH3 + NH3(dry l or gas) --> CH3-CO-NH2 + CH3OH (amonolyse)

In aqueous acid or basic media cyano and amide compounds hydrates to ammonium salts!
HCN + H2O -H(+)/OH(-)-> HCO-NH2
HCO-NH2 -H(+)/OH(-)-> HCO2NH4
(this explains why H2SO4 and wet P2O5 can't be used to dehydrate amides to cyano or that NH3 dry (liquefied) gas has to be used in amonolyse of esters to get amides).

HCN is produced by high voltage sparks in a flow of cold dry NH3 gas in N2 between C eletrodes!
Being endothermic HCN needs to be cooled fast to get tiny % yield (usually exothermic way is favourised)
NH3 + C + energy --> HCN + H2
N2 + 2C --> NC-CN
NC-CN + NH3 --> NH2-CN + HCN
NH3 + HCN --> NH4CN
Results are HCN(l/g), NH4CN(s), C2N2(g) (cyanogen) and cyanamide(s).

Best way to get HCN is via HCl + excess K4Fe(CN)6 (ferrocyanide of K-hexacyanoferrite of K) or K3Fe(CN)6 (ferricyanide of K-hexacyanoferrate of K).Upon mild 40°C heating collect the vapours inside NaOH, KOH or NH4OH solution or inside an hermetic falsk (cold trap) at max 0°C (the lower the best since HCN boils arround 20°C).


PHILOU Zrealone - 21-5-2002 at 09:10

Yes, I forgot to say (it is sometimes hard to remember what I had written the first time when my post was denied on the former forum):
The typical solvants for the dehydration of ammonium salts of carboacids into amides is glycol and/or glycerol between 170 and 250°C!
Under reflux and cold trap to collect the amide (if volatile).
Try to make the reflux under N2 atmosphère or as minimum O2 as possible, otherwise the glycol/glycerol oxydise into aldehydes (accrolein, ...) that has very accrid and lacrymator fumes (their boiling point is also lower than the related bp of glycol and glycerol!).

Glycol: HOCH2-CH2OH
bp= 197°C under 760 mm (atm press); mp= -13°C

Glycerin: HOCH2-CHOH-CH2OH
bp= 182°C under 20mm (reduced press); mp= 20°C

Acrolein: CH2=CH-CH=O
bp= 53°C (760mm); mp= -87°C

This would allow you to distill and collect various amides as liquids (distillable or not) or as solids (amides often form solids due to strong H bondings):

Acrylamide: H2C=CH-CO-NH2
bp= 125°C (25 mm); mp= 85°C

Forma mide: H-CO-NH2
bp = 210°C (760 mm); mp= 2,5°C

Acetamide: CH3-CO-NH2
bp = 221°C (760mm); mp= 80°C

N-Acetylethanolamine: CH3-CO-NH-CH2-CH2-OH
bp = 152°C (5mm) (mp?)

Oxamide: NH2-CO-CO-NH2
mp >300°C

Malonamide: NH2-CO-CH2-CO-NH2
mp = 173°C

Benzamide: C6H5-CO-NH2
mp = 129°C



Rhadon - 2-6-2002 at 01:07

The Kirk Othmer Encyclopedia mentions that cyanides rapidly oxidize to cyanates when heated...

Do you know from what temperature on the formation of cyanate takes place in significant amounts?
Anyway, the cyanate can be decomposed again to a cyanide by heating it (KCNO decomposes at 700 - 800° C).

...especially since he [Saxon] says 50 mg of KCN will kill a man...

The lethal dose of potassium cyanide is, according to Römpp chemistry lexicon, 120 - 250 mg (human).

When working with HCN (I didn't do it yet, but if I would...) I'd use an old hoover as a fume-hood. The device itself would be placed outside of the house, the hose reaching through the small opening of a window to the inside of the room. If the hose end exactly where HCN is supposed to be given off you're likely to not even getting in contact with the smallest amount.
I tested this "fume hood" by boiling 25% NH3 solution inside of my room - I didn't smell anything as long as the hoover was turned on, thus I assume that it works perfectly.

I don't know if this is new to you anymore, since I already postet this on the Explosives And weapons Forum, but here is a method of making (yellow) potassium ferrocyanide:

Fe4[Fe(CN)6]3 + 12 KOH ==> 3 K4[Fe(CN)6] + 2 Fe2O3 + 6 H2O

Note: Fe4[Fe(CN)6]3 = prussian blue

Polverone - 2-6-2002 at 13:09

The improvised fume hood idea is a good one. Personally, though, I'd just do the work outside and have an extra scrubber bottle at the end of my apparatus (if I could manage it). I don't think I saw a mention of specific temperatures for cyanides oxidizing to cyanates, but I only have access to the concise encyclopedia, not the full one. According to

Gilbert N. Lewis, Thomas B. Brighton;
J. Am. Chem. Soc.; 1918; 40(3); 482-489.,

"it is known that fused cyanide is readily oxidized by the air to cyanate." The authors also state that carbon dioxide will oxidize cyanides to cyanates, but that carbon monoxide will reduce cyanates to cyanides "in part"; the equilibrium between cyanides/cyanates is a significant part of what they write about.

Further down the cyanide trail...

Polverone - 18-7-2002 at 21:47

Although I should just give in and distill HCN into aqueous NaOH, I continue to tilt at windmills and attempt to find a method of preparing relatively pure cyanides with *no* HCN involved.

Recent paths I have taken:

Philou Zrealone (I certainly wish he was still posting here) communicated privately to me that sodium sulfide can be used to precipitate the iron from sodium ferrocyanide or ferricyanide, leaving NaCN in solution. I tried using sodium sulfide with potassium ferrocyanide with no success. I wonder if he's actually used this method, if I'm missing a necessary condition, if I really need to start with the sodium salt instead of the potassium salt, or what. I have tried a number of different variations on the same basic theme and have yet to obtain anything resembling pure KCN or NaCN. On the most recent attempt I prepared a concentrated solution of potassium ferrocyanide and kept adding sodium sulfide, muttering at it to "precipitate, already!" with no such luck.

Another tack I took was to make another attempt at preparing a sodium cyanide/ferrocyanide mix by dissolving blood meal in molten NaOH. I had done this once before and obtained a mess that turned dark blue and evolved at least some HCN with the addition of sulfuric acid. This time I tried to add more blood meal to convert the totality of the NaOH to something useful.

I was limited, though, by heavy foaming whenever I added more blood meal. I might have done better with a large vessel to better contain everything. The multi-stage reaction was interesting, but I didn't get much interesting stuff as an end product. After dissolution and filtration of the fused mass I had some foul-smelling brown liquid (should have raised the temperature to more fully decompose all that blood). It foamed with the addition of citric acid, but I smelled nothing new, so I don't think there was any free NaCN. I didn't bother adding sulfuric acid before throwing it out because I'm not really looking for a route to ferrocyanides.

I decided it was time to do some more research in the ACS archives. I found one old article that mentioned in passing, rather depressingly, that the KCN produced by 19th century methods (heating bone/blood/leather with potassium carbonate to get ferrocyanide, then fusing ferrocyanide with carbonate to get cyanide) was very impure, rarely exceeding 38% KCN. So even if I built a furnace it wouldn't solve my quest.

Then I found another old article, a very interesting 2-part article from 1879 all about cyanogen and cyanides. All of the modes of cyanide formation mentioned required high temperatures. Most of them required *very* high temperatures. But there was a mention in passing (can't recall if it was from the 2-parter or another article) that some authorities believed that barium carbonate readily formed cyanide in a reducing atmosphere containing nitrogen, even at a cherry-red heat. Even better, barium carbonate has an extremely low solubility while barium cyanide's is quite high, so I would be getting pure Ba(CN)2!

I improvised a setup to maintain a high-temperature reducing atmosphere around a small metal dish filled with powdered charcoal and barium carbonate. I looped steel wire around the dish and tied the ends of the wires to a ring stand so that the loop of wire was in the middle of the ring. I also tied a second wire between the ring stand support and to the existing loop of wire to support the wire loop when it began to sag at high heat. I cut three narrow slots in a soup can to pass the wires and inverted it over the metal dish containing the powders. I also punched a few small holes in the top of the soup can for gases to slowly pass through. I placed a large gas burner beneath this arrangement and heated the assembly for about 30 minutes.

Visual inspection showed that the dish always was anywhere from dull red to bright orange from the heat. I adjusted the gas/air balance to the burner until I could see pale flames coming out of the vent holes in the top of the soup can, indicating that the atmosphere inside the can was reducing.

After the 30 minutes or so of heating I shut off the gas, removed the soup can, grabbed the dish containing the reactants, and wrapped it as quickly as possible in aluminum foil (to extinguish the charcoal that had ignited on air exposure and to let it cool). After it had cooled somewhat I added the dish contents to water, stirred, and filtered. The filtrate was perfectly colorless. It gave no reaction with citric acid. The attempted synthesis was a failure. I don't know if my atmosphere had too much carbon dioxide or was otherwise defective, or if I didn't reach high enough temperatures despite the dish's appearance, or if I didn't wait long enough, or if the original authorities were wrong about how easily barium carbonate formed the cyanide.

I then decided I'd try another method, one very unlikely to give me pure product yet interesting anyhow. There is a patent whose number escapes me, the basic premise of which is KNO3 + 4 C = 3 CO + KCN. It was actually done with carbon electrodes in an arc-furnace arrangement, but I had had some success before conducting this in a purely pyrotechnic manner.

I tried to do a slightly larger batch tonight than I had on the previous occasion. I prepared 40 g of a 5:1 molar ratio of charcoal and KNO3 (I wanted to ensure that there would be excess carbon). I used finely ground charcoal but coarse KNO3 powder since the faster reactions between fine powders drives more material away as smoke. After ignition I had a mass of fused liquid mixed with excess charcoal at the bottom of a can. I again added water and filtered to remove the charcoal. The liquid fizzed vigorously with citric acid but had no scent of HCN (and, yes, I've smelled HCN before) so I must conclude that I had potassium carbonate. This was especially annoying since my previous tiny batch *had* given the telltale scent of HCN on addition of mild acid. I don't know what went wrong this time.

Now I'm again going back to the drawing board. Potassium cyanate is the salt of cyanic acid. It can be reduced to potassium cyanide at relatively high temperatures. Cyanuric acid, the trimer of cyanic acid, is readily available as a chlorine-level stabilizing compound for swimming pools. However, I haven't been able to find much information on inorganic cyanurates. Is it plausible that potassium cyanurate, perhaps mixed with carbon, might also be reduced to potassium cyanide? I don't know. Neither can I do any further investigation at the moment since I have to leave for Texas Saturday and have to spend Friday preparing for the trip. Neither do I particularly want to buy 5 pounds of a chemical (cyanuric acid) only to find that I have no use for it. Any comments or insights would be appreciated.

Rhadon - 20-7-2002 at 13:16

Polverone, unfortunately I can't help you. But I can come up with a few questions :)

I make NaCN by decomposing ferrocyanides. Thus, it is mixed with another decomposition product: iron carbide. In order to get pure NaCN, I'll react the mixture of sodium cyanide / iron carbide with sulfuric acid, leading the evolved HCN through an aqueous solution of NaOH. Nothing new that far.
Now I asked myself if the iron carbide could react with sulfuric acid, especially if it is hot. I've never worked with the chemistry of carbides up to now, and finding information that goes beyond superficial things seems to be quite difficult.
And: Will I have to heat the mixture of H2SO4 / NaCN / Fe3C in order to separate the HCN? I presume not, because halogen halides have a very low solubility in H2SO4, so this should also be true for pseudohalogen halides.

Finally, success seems in sight...

Polverone - 31-7-2002 at 12:14

Since I have returned from Texas I have been able to experiment further, with encouraging results. First, I found a British patent, (710143), that relates a method of preparing cyanates from cyanuric acid. In the patent they are concerned about avoiding cyanide, but I am obviously NOT.

Quick summary of the patent: powdered alkali carbonates are mixed with powdered cyanuric acid and heated to about 520 C. If this is done in a carbon dioxide atmosphere, there is no detectable amount of cyanide formed. The reaction is carried out in a closed steel vessel at atmospheric pressure. Only about 2/3 of the stoichiometric amount of alkali carbonate should be used. If more carbonate is used, some remains unconverted to cyanate. If less is used, some ammonium carbonate and other products form.

I was using an open vessel so I used a slightly larger excess of cyanuric acid than the patent recommends, especially since I was planning on heating the mixture strongly enough to drive off any ammonium carbonate.

Here's how my latest experiment went:

I strongly heated sodium bicarbonate (baking soda) to produce sodium carbonate. I have sodium carbonate on hand but it is in the form of coarse granules containing some moisture, and I wanted a fine, anhydrous powder. I measured out 20 grams of the freshly prepared sodium carbonate and 27 grams of cyanuric acid granules. The granules were obtained as a swimming pool supply - "chlorine stabilizer, 100% cyanuric acid." I reduced the cyanuric acid to powder in a mortar and thoroughly mixed it with the carbonate. I also powdered 5 grams of charcoal and set it aside.

I poured the powder mix into a stainless steel dish, put the dish in a ring stand, and took the stand outside. I heated the dish with a large laboratory burner using propane as a fuel. Considerable "smoke" was given off as the mixture was heated. I don't know if this was volatilized cyanuric acid, ammonium carbonate, or a mixture of substances. It took about 10-15 minutes for the powder to completely melt down to a fluid. This occurred at a temperature so low that the reaction vessel was not glowing at all, so I am sure that the sodium carbonate (or at least a large proportion) was converted to cyanate. I then added the 5 grams of charcoal (somewhat in excess of what is theoretically needed to reduce cyanate to cyanide) and increased the heat by placing the burner closer to the vessel.

The charcoal powder does not readily mix with the molten salt, but it gradually absorbs and is wetted by the fluid to form a sort of paste. Gas evolution was fairly rapid at first, with lots of large bubbles forming and popping. As time went on the bubbles became fewer but the gases leaving the mix must have changed because the gas jets would ignite and burn with a sodium-yellow flame. I am unsure about this 2-phase gas evolution. What is the first gas that doesn't burn, and what is the second gas that does? I expected the reaction NaCNO + C = NaCN + CO, which could be the source of my flammable gas, but I'm not sure about the first part of the reaction.

The whole time this was going on, the liquid was slowly creeping up the sides of the vessel, forming interesting patterns. It was bubbling a bit on the metal. Near the top of the dish it was forming patterns that resembled toad skin. It was also turning white and infusible at the top - converted, I fear, back to sodium carbonate from my burner's carbon dioxide.

I continued the heating for 20-30 minutes after I added the charcoal. I wanted to heat it until all gas evolution ceased, but I wasn't sure how long that would take and didn't want to run out of propane. Plus, I feared that I would eventually be working counterproductively as CO2 converted my cyanide back to carbonate. Perhaps in future runs I should cover the dish with something to minimize CO2 intrusion.

I then withdrew heat and scraped the paste in the bottom into a lump while it was still hot (experience showed that it was very hard to remove if left as a uniform layer until cold). The lump, once it had cooled somewhat, was added to water. Stirring and heat, over the course of 1-2 hours, broke up the glassy lump and allowed me to filter the liquid to remove the charcoal.

The liquid was evaporated in a shallow dish over the course of a night. There is a faint cyanide odor to the granular masses I have, but I have no idea as to purity. This morning, again consulting the concise Kirk-Othmer, I learned that sodium cyanide can considerably hydrolyze to formate and ammonia above 50 C. Whoops! In the future I will use cooler water. That could definitely explain the strong ammonia scent over the dish in the later phases of evaporation. I thought it was just leftover cyanate hydrolizing and releasing that NH4.

These results seem fairly encouraging. I seem to have made sodium cyanide (of unknown purity, unfortunately) without a furnace, any special chemicals, or handling HCN gas. I don't know if I have enough propane left to do another run before refilling. I would really like to discover if the flammable-gas-evolution ever ceases, and if that also marks the complete conversion of cyanate to cyanide. I would like to try running the reaction at a higher temperature to see if the conversion is appreciably faster. I would also like - good lord would I like - to be able to perform a more sophisticated analysis on my end product to see what is really in it, and in what proportions.

PHILOU Zrealone - 15-8-2002 at 16:56

BaCO3 + 4C --2000°C --> BaC2 +CO2 + CO
BaC2 + N2 -heat-> Ba(CN)2

By reductive atmosphère they don't mean reductive flame, but N2/NH3 without O2!

HCN -H2O-> H-CO-NH2 -H2O-> H-CO2-NH4

Another interesting reaction is:
Hg + C2N2 --> Hg(CN)2 (explosive when dry!)

Sodium dichloroisocyanurate is a derivative of cyanuric chloride (C3N3Cl3):
C3N3Cl3 + 3 H2O --> C3N3(-OH)3 + 3 HCl
(-C(-OH)=N-)3 (cyanuric acid) <--> (-C(=O)-NH-)3 (isocyanuric acid)
(-C(=O)-NH-)3 + 3Cl2 --> (-C(=O)-NCl-)3 + 3HCl
(-C(=O)-NCl-)3 + 2NaOH --> (-C(=O)-NNa-(C(=O)-NCl-)2 + NaOCl + H2O
(this last reaction explains why it can be used to clean pools since it frees NaOCl (hypochlorite like in Javel water!).


Polverone - 17-8-2002 at 15:37

[QUOTE]BaCO3 + 4C --2000°C --> BaC2 +CO2 + CO
BaC2 + N2 -heat-> Ba(CN)2

By reductive atmosphère they don't mean reductive flame, but N2/NH3 without O2![/QUOTE]

I was hoping I could get away with a reducing flame since it would at least exclude oxygen, and contain considerable N2 from the air. I figured I could live with a low yield since barium cyanide/carbonate would be very easy to separate. I found an older reference that indicated that the conversion of barium carbonate to cyanide would happen well below 2000 C under the right conditions, but I was able to obtain nothing.

PHILOU Zrealone - 21-8-2002 at 14:15

Maybe a typo?
Baryum carbide --> Baryum carbonate

vulture - 8-9-2002 at 01:47

I wouldn't want to be handling Ba(CN)2 if you offered me money for it! Extremely poisonous cyanide plus very poisonous Ba 2+ ions. That is, if it's soluble...

PHILOU Zrealone - 8-9-2002 at 11:04

Yes Ba(CN)2 is soluble and toxic!
If you want to stay logical with yourself:
Pb, Hg, Fe, Cu, Ni, Co, N3(-), CN(-), Ba, Li, Sr, NO3(-), NO2(-), aceton ,toluen, ethanol, methanol, .... nearly all chemicals are toxic even NaCl it only is a mather of use, safety (linked to the amount of knowledge and good sense you have) and quantity.

BaC2 is the tween brother of CaC2; it is made the same way with approximatively the same amount of energy.Now they differ a bit on properties and even if they free C2H2 upon water contact, they display different affinity for N2:
CaC2 + N2 --> CaN-CN (calcium cyanamide used as fertiliser) + CxNy
BaC2 + N2 --> Ba(CN)2 (baryum cyanide)


raistlin - 8-9-2002 at 11:30

I have read several places that there is such a thing as SCN, but that it isnt toxic. I want to know if anyone out there has ever heard about it, and if it is toxic or not.

PHILOU Zrealone - 9-9-2002 at 14:02

In chemistry it is not rare to have S in place of O!
H2S vs H2O
Hydrogen sulfide vs hydrogen oxyde

CH3-SH vs CH3-OH
methyl thiol vs methylol
CH3-S-CH3 vs CH3-O-CH3
dimethyl sulfide vs dimethyl ether
CH3-S-S-CH3 vs CH3-O-O-CH3
Dimethyl dissulfide vs methyle ether peroxyde
thiocyanate vs cyanate!

HS-C#N vs HO-C#N thus
S-CN(-) display similar properties with Cl(-), CN(-), OCN(-), N3(-) and belongs to the family of the speudo halogen!
It is supposed to be an energy rich fuel
and many combination of thiocyanates (as with cyanides) and oxydisers are high explosive mixes (much more explosive than average pyrotechnic binary mixes)!
The free acid is unstable and polymerises!


Sorry to dredge up an old topic...

Polverone - 7-10-2002 at 10:24

I've been reading some old chemistry books (thanks, a_bab) and I discovered that yes, you can prepare hydrogen cyanide and cyanogen (mixed with water in both cases) by the thermal decomposition of ammonium formate and ammonium oxalate. Madscientist's initial guess was therefore correct. However, the experimental difficulties he encountered serve to illustrate another valuable principle: lab work isn't as simple and easy as paper work! It's especially difficult to verify your results when you work with crude, low-cost materials and can perform only crude qualitative analysis. Of course, this makes success all the sweeter for the amateur experimentalist...

Bringing up old topics is a good thing!

madscientist - 7-10-2002 at 12:31

The difficulties I had encountered that Polverone speaks of have to do with heating ammonium oxalate to try to yield oxamide. It resulted in some oxamide, ammonia, and little bit of carbon - the oxamide gradually vanished, leaving just a small quantity of carbon behind. There probably wouldn't have been nearly so much ammonia liberated if the ammonium oxalate hadn't been a hydrate. The carbon was most likely the result of the thermal decomposition of cyanogen, (CN)2, which is endothermic.


Polverone - 27-10-2002 at 12:07

I get the feeling that other people don't care even half as much as I do about cyanides ;-) but I've got to share my latest experiment anyway.

I was treading the path of cyanate reduction again - see my earlier posts. This time I started with about 50 grams each of NaHCO3 (subsequently converted to Na2CO3 by heat) and cyanuric acid. The NaHCO3 was obtained as baking soda; I used it (as before) to obtain finely powdered Na2CO3 without a lot of manual labor.

This time I did a few things differently: I used charcoal instead of gas for my heat source - which allowed me to maintain a high temperature for a long time - and I mostly excluded air from the reaction vessel. I'm not sure if that mattered very much, though.

In any case, I built a fire in a little charcoal barbeque, using a mixture of briquettes and large lumps of mesquite charcoal (since I had both on hand). Once the fire had taken to the charcoal a little bit, I put the NaHCO3 in an empty soup can and heated it over the fire until it appeared that it had turned to carbonate. I then added the powdered cyanuric acid and 10 grams of powdered grapevine charcoal (this is well in excess of the amount theoretically needed for the reduction; I'd use less next time).

I placed the can back on the fire and loosely sealed the top by setting a tapering stainless steel dish weighted with sand in the mouth of the can; gases could escape, but little carbon dioxide from the fire would be making its way to the interior of the can.

I then invigorated my fire and heated the can up with the aid of an electric hairdryer directed at the burning charcoal. When all the charcoal was burning well and heaped up around the can, this was no longer necessary. The can maintained a healthy red-orange glow even without the extra air.

When I removed the air-blocking dish at the top of the can, I was greeted with a small spontaneous yellow fireball as the hot flammable gases inside finally met sufficient oxygen for combustion. I left the dish off for a bit just to watch. There were places in the pasty mass of chemicals where a continuous stream of flammable gas issued forth. I assume that the gas was carbon monoxide, but it burned with a vivid yellow, probably due to picking up sodium compound vapor from the hot melt.

I replaced the air-blocking dish and waited. At periodic intervals I removed it to check on the progress of the mixture. It wasn't long before the continuous streams of burning gas disappeared, but I continued to see bursts of flame when I removed the dish for about an hour. After that time, there was no sign of further reduction when I exposed the hot interior gases to the air. Nevertheless, I continued to let it heat for another 2 hours since my charcoal fire was proving so long-lived.

After that time I removed the can from the fire and let it cool. I then cut the can up with tin snips - the metal was extremely brittle and was heavily oxidized - until I retrieved the bottom portion with the hardened mass of charcoal and salts. I smashed the mass free with a small hammer, crushed it into small chunks, and covered it with room temperature water in a small jar. I swirled the jar periodically and 2 hours later filtered the liquid to remove the charcoal. The bulk of the filtrate is now drying at room temperature in a large glass pan covered with a grocery bag (to prevent dust/other junk from falling in).

I would love to dry this material more rapidly but hot water will accelerate the hydrolysis of the cyanide. If I used forced air to dry it, I would be exposing the liquid to increased carbon dioxide from the atmosphere, again decreasing yields (and presenting a possible hazard as well). I suppose what I really need is a vacuum distillation setup so I can remove the bulk of the water rapidly at low temperatures. Alternatively, a solvent that dissolves NaCN without decomposition and evaporates more rapidly than water would be used, if I knew of one. Or a solvent that is miscible with water but unreactive with and a poor solvent of NaCN could be used to force crystals out of solution so they could be collected by filtration.

Anyway, how do I know I have NaCN of reasonable purity? The solution has the slipperiness of a base, more intense than Na2CO3. It has the characteristic odor of HCN (don't rely on this alone - many can't smell HCN). Small amounts that I have evaporated have yielded cubical crystals resembling table salt - a crystal structure that NaCN has and Na2CO3 and NaCNO don't. I have obtained a lovely prussian blue on adding drops of the filtrate to a slightly acidified mixture of Fe(II)/Fe(III) salts. Anyway, that's my tale. I hope you enjoyed it.

rikkitikkitavi - 27-10-2002 at 13:05


time to fire up that old coal oven.

Polverone, you said the reaction mixture was subjected to a heat enough to make it glow red-dull? That would place it around 750 C ?

what would be a good way to dispose off any reaction solutions?


Polverone - 27-10-2002 at 14:02

No, it was hotter than dull red. It was an orange-red, and (guessing based on the color vs. temperature chart at I would say it was around 850 C. You don't even really need any sort of oven/furnace for this. You could build a charcoal fire in a perforated coffee can and probably achieve the temperatures I did.

CN- is easily oxidized to the harmless CNO-. According to, sodium hypochlorite is a good choice for this task. I imagine that other (alkaline!) oxidizing agents would work as well, but the hypochlorite is inexpensive and readily available.

I don't have much time for experimentation nowadays due to schoolwork - I shouldn't even have spent time performing this weekend's experiments - but I now want to try making various cyano-complexes of transition metals. I've only seen iron compounds thus far - what other colorful materials await me? Mmm, I love mad science.

Nick F - 27-10-2002 at 14:55

I'm glad you got it to work! It means I can copy you and avoid all the frustrating failures!! :)
I'd love a good way to make cyanides, and now it seems there is one, although I'll be making my cyanates with urea.
There are a few things I'd like to try with cyanides, mainly the production of 1H-5H tetrazole with NaN3 and NaCN (rather than 5-R tetrazoles made with nitriles), also cyanides as fuels (especially amine cyanides in gas generating compositions), and silver/copper (etc.) cyanide oxosalts as energetic materials.

For determining cyanide content, maybe you could find a metal cyanide that is insoluble with water, probably lead or silver cyanide would be. Dissolve a known weight of your product in water, add excess Pb(NO3)2 or AgNO3 or whatever, collect the ppte, dry and weigh. You work out the number of moles of ppte, and you therefore know the number of moles of cyanide in your initial sample, and from there it's easy to work out the purity.

Nick F - 27-10-2002 at 14:56

Damn, I forgot hydroxides, carbonates etc will probably also ppte :(

Polverone - 27-10-2002 at 15:18

I haven't yet come up with a good method of analysis. I should pay a visit to the library, but I've been busy. One idea that comes to mind is precipitating any carbonate with calcium or barium nitrate, then precipitating cyanide with silver nitrate. I'm quite sure there are no significant amounts of hydroxides in my solution. Unless I work with large quantities, though, I won't have very much accuracy in my work. My electronic balance is sensitive only to 0.1 g. That's fine for most of the stuff I do, but terrible for analysis.

Anyway, I'd like to hear about it when/if you get around to making cyanides starting from urea. I was going to try that myself as a followup to the cyanuric acid method, but I have been trying to be good and do some homework this afternoon.

Another tip

Polverone - 27-10-2002 at 17:24

I stained some of my glassware with prussian blue while doing cyanide tests. Conc. HCl didn't seem to have any effect on it. Dilute HF took it off like a charm, though, and didn't harm the glass either.

And speaking of cyanide tests, here's an extremely simple way of obtaining the mixture of Fe(II) and Fe(III) salts you need to precipitate cyanide as prussian blue:

Dissolve copper (II) sulfate in water. Add an excess of steel wool or iron powder, so that you are left with a mixture of iron (II) sulfate, metallic copper, and metallic iron. Filter to remove metal bits. Keep the iron (II) sulfate solution stored well protected from the air, since it's easily oxidized to iron (III) sulfate. In fact, that's how it's used in the cyanide test.

Mix a bit of the green iron (II) sulfate solution with a sample of your suspected cyanide. You may need to acidify it, too, with a little HCl or similar (obviously you need to limit the amount of suspected cyanide you test because of this step!) Swirl the liquid around to agitate it and expose it to the air. If there's cyanide in there, the liquid will take on a beautiful shade of blue (and/or precipitate blue particles) as the iron (II) sulfate is oxidized by the oxygen in the air.

I'm sure you can also prepare iron (II) sulfate from iron and sulfuric acid, but CuSO4 is safer and easier to work with and clean up. Also, the CuSO4-iron reaction is quickly completed, whereas you can wait days for the last traces of acid to finish reacting when you just mix an acid with an excess of metal.

Nick F - 29-10-2002 at 04:44

Polverone, I posted my results in your thread on the E&W. They're quite promising.

notagod - 29-10-2002 at 14:25

I think I read somewhere that the discoverer of HCN, Scheele, later used ammonium chloride when preparing potassium ferrocyanide, insted of dried blood and animal hides. Maybe it's one way to go?

madscientist - 29-10-2002 at 15:37

Nick, could you post your results here as well? :)

Polverone - 29-10-2002 at 16:21

I think I read somewhere that the discoverer of HCN, Scheele, later used ammonium chloride when preparing potassium ferrocyanide, insted of dried blood and animal hides. Maybe it's one way to go?

I don't see how ammonium chloride could be directly substituted. If you heat ammonium chloride with a metallic base, you're going to end up with a metallic chloride. Was it some other ammonium salt that you were thinking of?

And, yes (to Nick F.), it would be nice to see that message that you posted to E&W, as it seems that is taking an extended leave of absence...

Nick F - 30-10-2002 at 05:20

Oh yes, sorry I forgot access has been restricted... here it is:

16g of urea, 16g of NaHCO3 and 5g of charcoal were ball milled for half an hour, to form an intimately mixed dark grey powder.
The mixture was heated on an oil bath at around 150*C with stirring until the reaction had stopped (lots of ammonia is given off, so do it outside!), as indicated by the effervescence susbsiding. The pasty mixture was stirred as it cooled to break it up.
The black, gravel-like substance remaining was put into an iron crucible with a narrow neck (used to be a CO2 canister), and placed in a very hot (bright yellow) fire. I used a coal fire, fed with air at about 400*C from a paint stripping gun. After a short amount of time a flammable gas was produced (CO), and heating was stopped 10 minutes after this gas was no longer produced. Charcoal was poured into the crucible, which was then left to cool, with the charcoal hopefully excluding most air.

The residue was extracted with cold water and filtered, leaving a clear, colourless, strongly alkaline solution. A drop put into citric acid solution produces HCN smell :)
I've made some CuSO4, so I'll do the prussian blue test too.

(I haven't got round to doing the prussian blue test yet...)
So it would appear that it has worked, but I have no way of doing a quantitative analysis :(

Polverone - 30-10-2002 at 09:23

It certainly sounds like you've encountered success too. Try to dry yours a little faster than I've done with mine. I've left mine at room temperature for days, and I fear that when it's finally dry I'll mostly have sodium carbonate just from atmospheric CO2. I'm going to have make another batch, I think. On the other hand, the Kirk-Othmer Encyclopedia says that above 50 degrees C, NaCN solutions undergo some irreversible conversion to formate and ammonia. So I have to be careful to use gentle heating if I want accelerated drying in the future.

P.S.: It's not just that E&W has gotten more restrictive in who's allowed in. I still have my account and everything. It's just that for three days now I've been getting the message "Rogue Science is closed for October 28" and a prompt for a username/password. My usual name/pw combo doesn't work, either.

Nick F - 31-10-2002 at 01:34

That's what I meant by restricted, only staff have access at the moment.

My NaCN solution has also been sitting around for a while now, I'll do another batch and vac. dry it straight away.


LiveDestroyer_23_X_ANAH - 31-10-2002 at 01:58

Be careful and open you eyes ! :cool:
Are you sure the mirrored E&W forum
was (are) not a place to isolate people by lousy
law-hunters and this sort of lackeys ?


Ramiel - 31-1-2003 at 21:32

How would one synthesise or purchase cyanuric acid?

trinitrotoluene - 14-2-2003 at 23:59

I say purchase it. Today I went to the HomeDepo which is a pretty big hardware store and checked out the pool chemicals section, I had found cyanuric acid, I forgot what I said its used for. Never brought it though because I was short on money and time. From what I remember I think its in 1 pound (454 gram) bottles. The cost is around $15 pretty cheap.Is it possable to prepear cyanides with cyanuric acid?

Polverone - 15-2-2003 at 10:56

Yes, it is used to protect swimming pool chlorination agents from photodegradation, IIRC. And yes, it can be used to prepare cyanides - see earlier in this thread. Sorry Ramiel, I would have answered much sooner but I hadn't noticed your question. I am glad someone did.

I read above

Boob Raider - 30-3-2003 at 12:10

that someone tried to ppt. the Fe2+ from the Na Ferrocyanide by the addition of Na2S. I for some reason (which I can't quite conceptually explain) think NH4OH will probably ppt the Fe2+ out of soln. Leaving a mix of NaCN, NH4CN. Now X-# of moles of NaOH are added (X = # of moles of NH4+). This theoretically should leave quite pure NaCN.

[Edited on 30-3-2003 by Boob Raider]

I also found

Boob Raider - 31-3-2003 at 19:17

out that Na ferrocyanide decomposes at ~ 485*C to NaCN, Fe, C and N2. But a lot of CN- is lost as C and N2, although it is a way around HCN and NaCN obtained can very easily be purified.

Making cyanides

Theoretic - 18-6-2003 at 04:07

You can make cyanides just by mixing an alkali or an alkali-eart carbonate, carbon and heating that strongly on air (no reducing atmospere needed)

Theoretic - 1-7-2003 at 07:16

The original quote:
"Cyanides.—The salts of this acid, known as cyanides, may be prepared by the action of cyanogen or of gaseous hydrocyanic acid on a metal; by heating the carbonates or hydrooxides of the alkali metals in a current of hydrocyanic acid; by heating alkaline carbonates with carbon in the presence of free nitrogen: BaCOI + 4C + N2= Ba(NC)i+ 3C0;"
This may be useful as well:
"Ammonium cyanide, NH4NC, a white solid found to some slight extent in illuminating gas, is easily soluble in water and alcohol, and is very poisonous. Its vapour is inflammable. It is obtained by passing ammonia gas over hot coal; by subliming a mixture of ammonium chloride and potassium cyanide; by passing a mixture of ammonia gas and chloroform vapour through a red hot tube; and by heating a mixture of ammonia and carbon monoxide:

"heating strongly" is putting it mildly

Polverone - 1-7-2003 at 11:14

Before I used the cyanate method I tried heating powdered charcoal with barium carbonate in air. I got the container to glow a nice cherry red and held it there for 20 minutes, but after the mix had cooled I couldn't find even a trace of cyanide.

DDTea - 1-7-2003 at 13:42

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.


Theoretic - 3-7-2003 at 04:41

How about the ammonia-coal route?
I think NH4CN is convertible into alkali-metal cyanides by passing it through an alkali-metal hydroxide.:o

Iv4 - 7-7-2003 at 07:58

I'll try tht.The coal/ammonia thing.Recomend any tests?

Theoretic - 8-7-2003 at 06:41

Yes, (quote)
"One of the oldest, simplest, and best
qualitative tests is the Prussian blue test.
This test combines the unknown with a mixture of iron (II) and iron (III) salts in acidic media. The formation of a blue solution or precipitate indicates the presence of cyanide."
I add myself: Fe(II) and especially Fe(III) salts undergo strong hydrolysis, so
they're already self-acidified.:cool:

Iv4 - 8-7-2003 at 07:52

I guess I forgot to mention that I have next to nothing.How about trying to make copper cyanid?

Another thing.When you say ammonia you mean ammoina /water or the gas?
I sucesfully made some sodium cyanide and ran an old fashion gold test(suposed to get paid for it).200 samples(all negative,btw).

The sodium cyanide was prepared by heating with sodium carbonate and urea.After crushing the suspected ore I added water/NaCN to exces(first few times anyway).After about 1 minutes of stiring I filtered out the rock/crap/whatever.

Now I know activated carbon would have been better for this case but all I had was ion exchange resin.No precipitate was observed in any of them.Just for the hell of it I actually used this with a few spec of gold(no not mine)and precipites were observed.

Just tought you might want this and I'm sorry I dont hav better detrails.Just got it from some prospcters guide.

[Edited on 9-7-2003 by Iv4]

Theoretic - 9-7-2003 at 03:57

Does anybody know of electrochemical preparation af cyanides, e.g.,elctrolysing NaNO3 with a carbon anode or something like that?:o

blip - 26-7-2003 at 05:44

<a href="" target="_blank">NaNH<sub>2</sub> + C <sup><u>&nbsp;<font face="symbol">D</font>&nbsp;</u></sup>> NaCN + H<sub>2</sub></a>

I am a fish - 26-7-2003 at 05:50

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.

DDTea - 26-7-2003 at 09:09

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 seems like "cheating" to me; the struggle to obtain chemicals is part of the fun for me :D.

blip - 26-7-2003 at 16:07

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>&nbsp;&nbsp;</strike>> CHCl<sub>3</sub> + NaCH<sub>3</sub>COO + 2 NaOH + 3 NaCl

I am a fish - 27-7-2003 at 08:01

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?

It will work with either. However potassium ferricyanide is cheaper and contains more cyanide ions per unit mass.


chloric1 - 29-11-2003 at 17:28

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]

unionised - 30-11-2003 at 05:19

I'm a little suprised that nobody has tried the "grow your own" approach to cyanide based on things like this.
And there are much better sources too.

Yes I am aware of this

chloric1 - 30-11-2003 at 08:18

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]

Here it is!

chloric1 - 30-11-2003 at 08:29

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 2239 times

fritz - 18-12-2003 at 10:37

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)

unionised - 18-12-2003 at 12:39

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?)

vulture - 18-12-2003 at 13:26

That Ag test is tricky, because excess CN will dissolve Ag as Ag(CN)2 -

fritz - 18-12-2003 at 14:24

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!!)

unionised - 21-12-2003 at 14:55

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.

Another Method

Jay Maity - 25-12-2003 at 20:15

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

[Edited on 4-1-2004 by Jay Maity]

DDTea - 26-12-2003 at 14:28

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.

Loosely related, but Interesting

DDTea - 27-12-2003 at 14:57

To keep predators away, Millipedes secrete Hydrogen Cyanide :D . 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:

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.

Rick Tobin,CEM

Among other things, Millipedes secrete Benzaldehyde, Acetic Acid, Chlorine and Iodine.

Rosco Bodine - 12-11-2004 at 17:04

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 .


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


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]

Text from U2U with Polverone's blessing:

Eclectic - 15-11-2004 at 14:48

To: Eclectic
Sent: 15-11-2004 at 05:40 PM
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.

Earlier Text from U2U

Eclectic - 15-11-2004 at 14:58

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.

I'd ask in the forum, but the thread is pretty old:
Did you try reducing cyanates with iron turnings to make ferricyanides?


Idunno - 25-11-2004 at 06:55

Can someone post the formulas for these reactions:

1) Potassium Ferricyanide and Sulfuric acid and heat,


2) HCN and NaOH

Has anyone tried this setup? (Outside I'd assume.) Any advice?

Chris The Great - 30-5-2005 at 20:33

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.

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.

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.

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.

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.

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.

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.

Polverone - 30-5-2005 at 21:01

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]

more CN- stuff

Polverone - 30-5-2005 at 21:34

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.

chloric1 - 31-5-2005 at 16:09

Well Polverone, that is just plain delicious!:D Buy the way, after you solvate your raw product do you filter it before adding to alcohol? I know it may seem to be a smartass question but I remeber the cyanurate process involved filtering through a little diatamaceous earth. Did I even spell that right?:o

Polverone - 31-5-2005 at 17:21

Yes, the liquid should of course be filtered. Coffee filters will do okay but a layer of diatomaceous earth will filter out even the finest charcoal particles. Denatured alcohol definitely works for the precipitation, and since I'm in the US my denatured alcohol is more than half MeOH. I think ethanol should work too. I've also tried isopropanol and acetone to force the precipitation, but they are immiscible with the cyanide solution and no precipitation takes place.

Also, as I mentioned in another thread, I discovered more recently from <A HREF=""><i>Industrial Nitrogen Compounds and Explosives</i></A> that the high-temperature reduction of cyanates was at one time used as a patented industrial source of cyanides:

<b>Siepermann's Process</b> (see English Patents 13,697 of 1889, 9,350 and 9,351 of 1900).

One part of sodium carbonate and two parts of charcoal (that is, sufficient charcoal to keep the mass from fusing during the process) is heated to dark redness in the upper part of a vertical iron tube while a current of ammonia gas is sent through the mixture. Potassium cyanate, KCNO, is formed thus :

K2CO3 + NH3 = KCNO + KOH + H2O.

The mixture is then allowed to fall to the bottom of the tube, where it is heated to bright redness. The cyanate is decomposed with formation of cyanide:

KCNO + C = KCN + CO.

The final product is thrown into air-tight vessels, cooled, lixiviated with water, the solution being evaporated in vacuo until the KCN crystallises out. <b>KCN is soluble with difficulty in the presence of much KOH or K2CO3, and crystallises out before these salts in the form of anhydrous crystals.</b> As first made the KCN was a damp deliquescent mass, which had to be fused with the product of the ferrocyanide process. The working of the process is difficult. It has been worked at Stassfurt since 1892.

<b>Bielby Process</b> (see English Patent, 4,820 of 1891). The principle is much the same as the Siepermann Process, but differs in important details. Much less carbon is used, so that at the end only slight excess remains. The charcoal is added gradually during the operation, so that the material is always present as a molten liquid through which the ammonia gas is forced under a slight pressure, when the following action takes place:

K2CO3 + 4C + 2NH3 = 2KCN + 3CO + 3H2.

The final molten product is filtered from the small excess of unchanged charcoal, and thus a white saleable product is directly obtained without the difficulties of lixiviation. However, since the melting point of the pure potassium carbonate is inconveniently high (about 890 C), ready-made cyanide is added to it in order to reduce the temperature of fusion.

The Beilby process has been worked since 1892 by the Cassel Gold Extracting Co. at Glasgow, has achieved remarkable success. In 1899 Beilby's process was estimated to supply fully 50 per cent of the world's output of high-strength cyanide.

Despite the mix-up with sodium and potassium carbonates near the beginning of the text, the description offers an interesting method of forming pure cyanide without the danger of working with HCN: evaporating KCN solutions containing excess K2CO3 or KOH <i>in vacuo</i> will cause the KCN to precipitate out as pure crystals before the contaminants. Since home experimenters are concerned more with convenience than with cost and throughput, it is of course more convenient for us to form the cyanate from urea or cyanuric acid than gaseous NH3, but the final reduction should work the same either way.

Edit: these tidbits of information nicely illustrate some of the WWW's limitations. The above text says that the Beilby cyanide process once provided half of the world's cyanide demand, yet there are only two unique references to it on the whole web, one from the Sciencemadness Library and another from the Digital Library of India.

[Edited on 6-1-2005 by Polverone]

garage chemist - 14-6-2005 at 08:18

I just got two nice small crucibles (30ml capacity, easily heated to red glow with a bunsen burner), one made of iron, the other one made of nickel.
I purchased them primarily for small- scale cyanide production (and also for molten salt electrolytic processes).
Which one would be better for cyanide production? I'd use the iron one since everyone here uses iron.

Your method for analyzing the cyanide product (test for cyanate) is interesting, Polverone. An easy test for cyanate was something I've been looking for.
The precipitation of crystalline cyanide with alcohol is also an enormous improvement.

Would it be more convenient to produce NaCN or KCN? I'm thinking about the solubility in alcohol (Ethanol), which should be as low as possible.
Does anyone know which one has lower solubility in ethanol?

EDIT: I found out: KCN dissolves in ethanol only 4,5g per liter! Perfect for this application.
NaCN has "low solubility" in ethanol- it's no further specified what this means.
Maybe someone has got reliable data on NaCN solubility in ethanol?

[Edited on 14-6-2005 by garage chemist]

Polverone - 14-6-2005 at 23:25

The iron crucible should work fine. Remember to keep it covered from air while you are heating it, as much as possible. I always used old food cans rather than any proper crucibles, and by the end of the reaction the cans were always very embrittled and thinned near the bottom. It may be simply because this is where the heat is applied more strongly; I don't know if the vessel contents contribute to the deterioration of the metal.

I have produced mostly NaCN. The first time I did it, I didn't heat it long/strong enough and got considerable cyanate along with my cyanide, as evidenced by the ammonia smell of my warm aqueous extract. Keep in mind that aqueous cyanide solutions too can hydrolyze to form ammonia (and formate), but this is considerably slower than with the cyanate. So I would be suspicious of your cyanide extract if it immediately smells like ammonia, but not so suspicious if it smells like ammonia after being kept warm for an extended period of time. This degradation, as well as interaction with atmospheric CO2, makes me prefer the precipitation with alcohol over plain evaporation to dryness.

For many purposes, some residual cyanate and carbonate in your cyanide will not matter very much, but it does complicate measurements when you are unsure of purity.

Empty gas cylinders look like they would be almost perfect makeshift crucibles -- CO2 cartridges for the small scale, or disposable propane cylinders for the larger scale. Their narrow necks and larger bodies should easily exclude air.

In the future, I would like to try compressing the powder into the vessel before heating it. It's possible that I could get better heat transfer and a larger batch size that way. Another variation that I would like to try is using a hydroxide, instead of carbonate, in the initial mixture, since this too should lead to easier melting.

Marvin - 15-6-2005 at 01:20

I'm far from an expert in this reaction but my understanding is when formed in this way molten cyanides tend to dissolve iron metal from the container to produce ferrocyanides. Maybe a graphite or clay slurry bound with a little sodium silicate coating the inside of the container would help?

Cyanates tend to hydrolyse very easily, under the circumstances water produced by use of hydroxide might cause problems.

garage chemist - 15-6-2005 at 02:04

Yesterday I tried it out with 6,9g K2CO3, 6,2g urea and 1,6g Charcoal.

I first dissolved both the urea and K2CO3 in water, boiled this down to dryness and heated it until no more ammonia was evolved (now I see that this was no good idea, and I likely have considerable formiate contamination in my product because of hydrolysis).

Reduction was carried out in the iron crucible. I first melted the cyanate and then stirred in the charcoal, this was very convenient.
CO evolution was evident at red glow, since when I lifted the lid to stir it a bit, there was a quick bluish flame and the mix was evolving gas streams.

After cooling and extraction with water, I saw that my charcoal was way too coarse and a lot was left behind.
However, the extract strongly smelled like almond and I didn't smell ammonia.

I'll do another batch and melt the urea + carbonate dry instead of first dissolving in water. I'll also use finer charcoal powder.

Saugi - 22-8-2005 at 02:40

I read this CN-Thread with interest. Some Years ago, I nearly killed myself by heating a very small amount of ferrocyanide with accumulator-acid (H2SO4, 37%) in a testtube on the balcony. So I made my first experience, how damned toxic this shit is...

But it's still fascinating. Things that are forbidden by law, are always fascinating.;)

The best method to get real pure CN ist to neutralize a solution of NaOH or KOH with the right amount of HCN, (as already described) there is no other way.

But there are different methods to produce HCN. The easiest way ist to heat some ferrocyanide in dilluted H2SO4, but I wonder if prussian blue will work for this too.

In a painter-shop you can order kg's of pigmented prussian blue (pure! also called as milori blue) = Fe4(FeCN6)3 which decomposes in contact with strong acids.
I'm not sure if HCN will be produced by mixing and heating it with dilluted H2SO4, cause i found no information on this on the web, but I believe so.

You also can produce Calciumcyanide (ca. 40% pure) by heating Calciumcyanamide with Calciumcarbide and NaCl. This will work without arc furnace too - lowers only the yield. This "black cyanide process" ist described on this page:

It's easily possible to produce amounts of blackcyanide, that can be turned into pure cyanide by mixing with acid and bubbling the gas into an -OH solution. Unfortunately it's quite dangerous, but can be done safe if you know how.

To produce temperatures up to 1800°C you can use a burner like this:

if the mixture is given in a quartz tube, it can be reacted by heating than without having contact with the surrounding air.

Calciumcyanide is cheap and available as an black prilled fertilizer (called Kalkstickstoff in GER)

A way to form Calciumcyanamide
If somebody has N2 in flasks, just let it flow over CaC2 in a tube and heat up to 1000°C. It will react: CaC2 + N2 --> CaCN2

[Edited on 22-8-2005 by Saugi]

garage chemist - 22-8-2005 at 04:04

The Calciumcyanamide process is interesting, since this substance is available as "Kalkstickstoff" easily where I live.

It is made by reaction of calcium carbide with N2: CaC2 + N2 ---> CaCN2 + C (this doesn't produce cyanide, but cyanamide! Seems like a typo in your above post).
The C isn't removed from the commercial product, the black mass is just prilled and sold as fertilizer.

This fertilizer therefore already presents an intimate mixture of CaCN2 and C. We only need to add NaCl and heat it.

Saugi - 22-8-2005 at 15:25

You're right, sorry for my mistake! I corrected it.

And I forgot to write that "Kalkstickstoff" is industrial grade Calciumcyanamide. It's an important info, cause not everybody will know this.

If one of you crazy guys has tested the cyanide synthesis with CaCN2 + NaCl - please let me know the results!

I'm sure this will work, but it depends on the correct amounts. If I had the time, I would make the test for my own. But be aware, the product and even it's fine dust will be evil toxic!

kmno4 - 7-9-2005 at 09:08

I think that the only way to obtain pure (>98%) cyjanides is introducing mixture HCN+H2O to akoholic solution KOH or NaOH. I have made many grams of KCN using this methode (more than 1kg :o)
Mixture HCN/H2O was obtained by heating K4[Fe(CN)6]+H2O+H2SO4 or H3PO4. Both acids work good and concentration(%) has no special meaning: 30%-60% is ok.
Reaction is:
2K4[Fe(CN)6]+3H2SO4 -> 6HCN + FeK2[Fe(CN)6]+3K2SO4
With a simple distilling aparaturus, during haeting, at about 50*C (temperature of vapors) the mixture HCN/H2O starts to evolute, and after condesation it is introduced into KOH/C2H5OH solution.
KCN starts to precipitate,the solution is getting warm, but there is no need to to cool it. This mixture MUST be continual stirrig by magnetic stirrer - in another case a "cake" is collecting at the surface and there is no full reaction between HCN and KOH - it is a little dangerous.
The distillation is interputted at about 95*C - just too many water is then evoluted.
KCN is filtered, washed 2-3 times and dried in stainless steal pot, at the gas burner.
Yield is about 80%, counting at given theoretical equation, very pure KCN. Do not dry it in the air, it is hygroscopic. I used about 100g K4[Fe(CN)6]*3H2O per one run.
In distillation flask there is a lot of green "shit" - it is FeK2[Fe(CN)6] and it is useless, not dissolvig in acids and easly oxidicing itself.
By reduction KOCN or NaOCN is very hard to obtain pure cyanides. In the best case you will have something what is a poison, being mixture of KOCN, KCN and K2CO3...
Also thermal decomposition of K4[Fe(CN)6] is not good, because KCN is very easily oxidate by atmospheric O2, yielding KOCN.
DO NOT TRY THIS If you are not sure of yourself or your lab equipment.

[Edited on 7-9-2005 by kmno4]

meyer - 29-1-2006 at 10:04

Hi I am a new User looking for preparation of NaCn. I find a very easy method at
http://www.hypatia- There is written the formula:

Na2Co3+4C+N2+ Heat ( Fe Catalyst)------> 2NaCn+3Co

The Bucher Process for the Preparation of Sodium Cyanide:
To prepare sodium cyanide, simply heat 5.41 grams of sodium carbonate [Na
2.45 grams of powdered charcoal or lampblack [C] with 4 grams or so of iron filings
[Fe] in a stream of nitrogen [N] (air, being 78% nitrogen, will suffice).

Can tell me someone if it will work??????

Organikum - 31-7-2006 at 02:24

I asked myself if polyethylene could not be used for reducing the cyanate to the cyanide. PE is known to reduce many metal oxides to the metal and if it works for reducing cyanates it might be a cleaner alternative to charcoal.


Polverone - 31-7-2006 at 18:40

That would be interesting to try as a variation. The greatest problems I've had are with frothing/bubble formation, which limits the amount that can be produced in a given volume and also impedes heat transfer. I fear that using PE would make the gas production even more of a problem, and I wonder how well the molten PE or its vapor would mix with molten cyanate/cyanide. The charcoal-reduced cyanide does seem to have traces of sulfur in it (faint smell of H2S), which PE would avoid.

Organikum - 1-8-2006 at 00:03

Frothing, bubbling, heat transfer? Coarse steelwool?

Nicodem - 1-8-2006 at 01:12

Wouldn't polyethylene start "cracking" at the temperature required? Most of it might escape in form of vapors of higher olefins.

If using pre-prepared tripotassium cyanurate the powder is fine enough (and chemicaly homogenous = as much as it can be) to form KOCN at lower temperatures (250-350°C).* This yields a highly homogenous and very fine product. Perhaps such product would be more ready to react with something like polyethylene at somewhat lower temperatures.
Perhaps the reduction would work even better with thoroughly homogenized crude KOCN with sugar:D, which then carbonizes to some fine carbon that should do the job. However, at the begining there would be a lot of annoying froating due to evolved H2O and CO2 from such a "caramel candy". Starch does not melt, so perhaps it would be more apropriate even though it would be less homogenous.

* This is how I once supposedly prepared some KOCN: I prepared "tripotassium cyanurate" (due to pKa3, I doubt one can prepare it stoichiometricaly pure from KOH) from cyanuric acid and 3 eq of KOH in IPA solution, vacuum filtered the white voluminous paste and while still wet (to avoid CO2 absorption) heated it on a hotplate well covered with alu-foil until up to some 250-350°C and left there for about 2h. There was some cracking of the powder at the beginning and later no more notable change. I haven't analyzed the product but given that no cyanuric acid precipitated after acidification with HCl of an aqueous solution of the product, I assume only KOCN could have formed. This was a modification based on GB710143 where Na2CO3 with cyanuric acid is used to produce NaOCN. If someone is terribly curious if this was truly KOCN, I might bother to check with IR (if someone is so kind to provide a reference spectra:P).

Analytical methods

Ozone - 8-11-2006 at 20:47


Polverone, on page one you indicated an interest in some analytical tests for cyanide (besides the nose...). I have attached EPA 9014 which includes the usual protocol for both titrimetric (more likely to find use here) and spectrophotmetric (better) methods.

This assay methodology is quite general, and so long as the matrix is not really nasty, there should be no distillation required (which is commonplace in more elaborate methods).
These methods are designed for trace levels though, so large dilution factors will be required, viz. 30,000 or more, for a relatively pure product.

Hope this helps,


Ozone - 8-11-2006 at 20:49


My apologies, I forgot to attach the file; here it is.


[Edited on 9-11-2006 by Ozone]

Attachment: EPA 9014.pdf (73kB)
This file has been downloaded 1320 times

kmno4 - 11-11-2006 at 18:25


KOCN can be obtained just by simply heating urea and K2CO3. Heating mixture easy melts (CO2+H2O goes out). Molten substance is KOCN. NaOCN (from Na2CO3) melts in much higher temp. and is harder to obtain as pure substance.
KOCN does not react even with Mg or Al at about 700 C.
[I wrote it all somewhere in this forum, but it was a long time ago :P]

[Edited on 12-11-2006 by kmno4]

"Murder Most Foul"

Magpie - 11-5-2007 at 16:26

I watched this Agatha Christie film today on TCM, the film classics channel. One person was murdered by what I think Jane Marple said was "Prussic acid," which she deduced by the "smell of almonds." Apparently some sodium cyanide mixed with candle wax was placed in a pot on a stove. A timer was set for the burner to go on and off when the victim was in the kitchen. This generated the HCN thereby killing said victim. The victim, btw, was unintentional, Miss Marple being the intended victim.

Pyridinium - 11-5-2007 at 21:14

Not that I want to work with cyanides (still not as scary to me as HF, but scary enough), but I have some use for the [much safer] cyanate. I'm working on some experiments with chemical microscopy of iron compounds. So I was looking at the old urea / K2CO3 method of making KOCN, mentioned above.

I came up with a balanced equation (OK I'm tired so it might not be all balanced):

2H2NCONH2 + K2CO3 -----> 2KOCN + CO2 + 2NH3 +H2O

Now, it's late and I'm getting sleepy, so I'm not sure if this is the correct equation of what happens, but it seemed to fit. Urea and potass. carbonate going into just KOCN, CO2 and water was unable to be balanced.

I have some old refs. for qualitative tests for both CN- and OCN-, some of which are pretty interesting. It's a pity CN- is so toxic, because it is used in a lot of qualitative tests I'd like to try. Not to mention it's still hard to beat for gold & silver recovery. But cyanate itself is interesting for a few reasons, I might edit this post later for a couple uses of it.

I may still have a jar of old crabgrass killer that's kicking around somewhere, I recall it was mostly KOCN. From the looks of it, I think the jar was from the 1960's or even the 50's. I should dig that out and test its purity.

EDIT: strange, but last night I checked and there was one copy of this post. Today I came back and there was a double, a quote of this post from myself. I deleted the double. Still not sure what happened.

[Edited on 12-5-2007 by Pyridinium]

Axt - 12-5-2007 at 05:59

I'm not sure if this holds any value:

"Carbonyl cyanide has been obtained[2] from <a href="">diisonitrosoacetone</a>. It is a colourless liquid with a boiling point of 65-5° C. at 740 mm. Density, 1.124 at 2O°C. By hydrolysis it decomposes forming carbon dioxide and hydrocyanic acid." M. satori, <a href="">The war Gases</a>, pg. 58.

2] MALACHOVSKY and coll., Ber., 1937, 70, 1012.

[Edited on 13-5-2007 by Axt]

kmno4 - 15-5-2007 at 01:42

Thermodynamic says that reaction:

KSCN + Pb -> KCN + PbS

is possible, dG is about -20kJ.
I am not able to check it, because my KSCN turned out to be NH4SCN ....

[Edited on 15-5-2007 by kmno4]

12AX7 - 15-5-2007 at 14:08

How do you seperate the lead from the sulfide? In solution is not going to work, it will passivate. A fusion akin to the nitrite/lead process might work.


garage chemist - 15-5-2007 at 16:27

Hmm, I will soon need a sizeable quantity of a crude cyanide in order to prepare the liquid anhydrous HCN that I need for the TCT synthesis.
I originally planned on purchasing 1 kg of potassium ferrocyanide and melting this with K2CO3 (liberates 5 of the 6 bound cyanide ions as KCN, one is turned into KOCN). But this would be quite expensive, and I also would have to buy more K2CO3 which isnt cheap either.
K4(Fe(CN)6) also comes as the tetrahydrate, adding dead weight.
Heating ferrocyanide with acid is out of the question, that is very wasteful of material due to hydrolysis of HCN and because only half of the cyanide is liberated as HCN in the first place.

I now decided upon melting soda with urea to make NaOCN and reducing this with charcoal powder in a soup can. The heating will be done with burning charcoal, I'm simply going to place it in a pile of lit charcoal in the grill and blow air at it.
I know that this makes some NaCN since I did this in a steel crucible in small scale.
The problem is that an awful lot of side reactions are going on, cyanates produce cyanides while melting even without any reducing agent by disproportionation into things like cyanamide, carbonate and also with nitrogen loss which is highly undesirable.
I will probably not be able to exceed a few ten percent of NaCN content in the resulting material. Not that this would be a big problem, I can simply make more, but I cant really dose the HCN for the synthesis.
I will probably react a weighed amount of raw cyanide with acid, condense the evolved HCN and weigh it to determine the HCN yield of my raw product.
Cyanate will liberate HOCN which immediately hydrolyses into carbonate, carbamate and ammonium, giving no rise to impurities in the liquid HCN.

Can you suggest better methods for determining the NaCN content of a solid crude material, in the presence of cyanate, cyanamide, carbonate and other gunk?

12AX7 - 15-5-2007 at 16:37

Why not colorimetrically with Fe(II) / Fe(III) solution? Redox titration perhaps?

Raw gravimetric or whatever methods are probably the only way to go for the impoverished chemist, unfortunately.

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