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Polverone
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[*] posted on 21-5-2002 at 09:03
Preparation of cyanides


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)?
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[*] posted on 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).
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[*] posted on 21-5-2002 at 09:04


Alternative method of production, by madscientist PREPARATION OF HYDROGEN CYANIDE FROM POTASSIUM PERMANGANATE, METHANOL, SULFURIC ACID, AQUEOUS AMMONIA, AND POTASSIUM HYDROXIDE

Notes:
-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

PREPARATION OF POTASSIUM FORMATE (HCOOK) AND MANGANESE FORMATE (Mn(HCOO)2):
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.


PREPARATION OF FORMIC ACID (HCOOH):
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


PREPARATION OF AMMONIUM FORMATE ( [HCOO-][NH4+] )
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


PREPARATION OF POTASSIUM CYANIDE (KCN):
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


PREPARATION OF HYDROGEN CYANIDE (HCN) FROM POTASSIUM CYANIDE:
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+]




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[*] posted on 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...
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[*] posted on 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.
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[*] posted on 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.




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[*] posted on 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
...
But!
(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).

PH Z
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[*] posted on 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

...

PH Z
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[*] posted on 2-6-2002 at 01:07


Quote:
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).

Quote:
...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
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[*] posted on 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

THE OXIDIZING POWER OF CYANATES AND THE FREE ENERGY OF FORMATION OF CYANIDES.,
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.
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[*] posted on 18-7-2002 at 21:47
Further down the cyanide trail...


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.
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[*] posted on 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.
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[*] posted on 31-7-2002 at 12:14
Finally, success seems in sight...


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.
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PHILOU Zrealone
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[*] posted on 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!).

PH Z
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[*] posted on 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.
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[*] posted on 21-8-2002 at 14:15


Maybe a typo?
Baryum carbide --> Baryum carbonate
????
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[*] posted on 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...



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[*] posted on 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)

PH Z
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[*] posted on 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.



\"To ignite, or not to ignite, that is the question.\"
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[*] posted on 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
and
HSCN vs HOCN
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!

PH Z
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[*] posted on 7-10-2002 at 10:24
Sorry to dredge up an old topic...


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...
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smile.gif posted on 7-10-2002 at 12:31
Bringing up old topics is a good thing!


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.



I weep at the sight of flaming acetic anhydride.
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[*] posted on 27-10-2002 at 12:07
Success


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.
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[*] posted on 27-10-2002 at 13:05


horray!

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?

/rickard
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[*] posted on 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 knives.com) 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 http://shorinternational.com/cyanodestruct.htm, 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.
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