Sciencemadness Discussion Board

OTC source of thiocyanates or ferrocyanides?

Fulmen - 16-4-2018 at 15:04

What are the most common uses for thiocyanates? Any products out there that could contain a fair amount? I suppose I could synthesize it from ferrocyanide, but again I don't know where to look for it. Ebay isn't really an option (overseas shipping and customs etc).
I need some as a rectifier for caustic black baths, thiocyanates are the only rectifier I have found so far.

Bert - 16-4-2018 at 15:45

Are you blackening ferrous metals, or something else?

--------------------

Does anyone in your country still use old fashioned blueprints? Or cyanotype photographic processes?

If anyone there sells photographic chemicals, you might look there.


[Edited on 4-17-2018 by Bert]

Fulmen - 17-4-2018 at 01:34

Yeah, it's a plain caustic black bath for steel.

Haven't found many shops selling photo-chems, and none of them had any thiocyanates. Cyanotype seems to be the best option, but it still involves import. It's still better than photo-chems, I have found a couple of cases where they tested positive for GHB. I don't need the police kicking down my door over some bogus drug charge...

Bert - 17-4-2018 at 06:10

High or low temperature? I assume high (around 141 C), but just to be sure.

Swinfi2 - 17-4-2018 at 06:25

Isn't K2Fe(CN)6 use as an anti-caking agent in table salt?

That would be a inconvenient source but maybe you can buy it as the additive?

Fulmen - 17-4-2018 at 06:33

High temperature.

Boffis - 18-4-2018 at 01:42

Both ferrocyanides and thiocyanate can be prepared easily from Prussian blue pigment. This pigment is fairly cheap if you can find a supplier, it depends I suppose on which country you are in. Try a larger supplier of art materials or specialist pigment suppliers.

The process method has been described on SM on several threads already.

Fulmen - 18-4-2018 at 03:04

Prussian blue is an interesting option. I do have some in the form of marking compound for metal work, but it's fairly expensive. I'll start digging through art&crafts-suppliers to see if anyone sells the pure stuff at a reasonable price.

PirateDocBrown - 18-4-2018 at 07:50

You can get Prussian Blue by just evaporating off the water from laundry bluing. Around here you can get an 8 oz bottle of bluing, yielding about 100 grams of pigment, for about $5.

It will have small traces of gluteraldehyde and oxalic acid additives, (at least the brand I use) but otherwise essentially pure Prussian Blue.

Fulmen - 18-4-2018 at 08:07

I have never seen any laundry bluing products around here, and a web search didn't turn up any products either. But I have found a shop that might sell pure prussian blue pigment at a reasonable price, so I think I will try that route first.

Boffis: I couldn't find any posts on prussian blue to thio, but several using alkali salts. I assume those should work as well?

PirateDocBrown - 18-4-2018 at 09:39

You can easily find Mrs Stewart's bluing on eBay.

I also would like to hear of good synthetic routes to both thiocyanides and ferrocyanides.

[Edited on 4/18/18 by PirateDocBrown]

Fulmen - 18-4-2018 at 10:23

True, it's just so painful having to pay 3xprice for shipping. I'll give the local sources a chance first.

Boffis - 18-4-2018 at 10:55

Quote: Originally posted by PirateDocBrown  
You can get Prussian Blue by just evaporating off the water from laundry bluing. Around here you can get an 8 oz bottle of bluing, yielding about 100 grams of pigment, for about $5.

It will have small traces of gluteraldehyde and oxalic acid additives, (at least the brand I use) but otherwise essentially pure Prussian Blue.



If you already have the pigment in suspension and want to make ferrocyanide just add a strong warm solution of KOH or NaOH until the blue colour has turned to an even dark brown, stir periodically and allow to settle for a day or so to make it easier to filter. Potassium ferrocyanide is less soluble in the cold so crystallises more easily.

@Fulmen, Thiocyanates are prepared from either dried potassium or sodium ferrocyanide and sulphur. They are simply mixed with the appropriate metal carbonate or hydroxide and sulphur and fused, the mixture turn black from the formation of iron sulphide and the appropriate thiocyanate leached out with water and crystallised by evaporation. I am sure this is mentioned in a thread somewhere on SM.

There's even a chemplayer video https://www.youtube.com/watch?v=CKNBKCnnwLY

If not it is certainly covered in several practical inorganic chemistry books

Here's a link to an ebay supplier in the UK, I have bought this brand so I can vouch for the quality:
https://www.ebay.co.uk/itm/Cornelissen-Dry-Pigment-Prussian-...

[Edited on 18-4-2018 by Boffis]

Fulmen - 18-4-2018 at 11:30

Thanks, Boffin. I was searching for a direct route, but converting it to the sodium salt isn't much work.

I have found a shop selling Paris Blue PB27. It's listed as ferriammonium ferrocyanide, but I doubt if that matters.

Boffis - 18-4-2018 at 22:36

@Fulmen, There is also this thread on SM

http://www.sciencemadness.org/talk/viewthread.php?tid=21329#...

but I have to say the OP's "recipe" is a bit off the mark as Blogfast pointed out. It need a source of nitrogen. Now a days cyanuric acid left over from making chlorine from TCCA tablets is probably a better source of nitrogen, or perhaps calcium cyanamide if you want to try making the ferrocyanide yourself without Prussian blue.

There are numerous variations on the basic Prussian blue including so called "soluble" Prussian blue which is actually a colloidal preparation but all work; a little ammonia may be evolved though.

It might be possible to convert Prussian blue directly to thiocyanate but I can't find a description of such a process but maybe heating the pigment with potassium polysulphide (liver of sulphur which is sold as a patinized for bronze and other copper alloys) or just fused sodium sulphide might work.

Fulmen - 18-4-2018 at 23:12

Well, since polysulfide is made by heating sodium carbonate and sulfur the original reaction should work as well. But I don't mind one extra step to avoid any possible problems. Synthesizing it from scratch sounds like a fun challenge, I might just give it a try. I do believe I have some cast iron turnings, and I can just dig plain steel turnings out of the lathe. Any idea what temperature is needed?

Boffis - 19-4-2018 at 04:07

If you are going to try from scratch here are a couple of references that you might find useful. They are both old and out of copy write, both cover the preparation of ferrocyanide from scratch and can be downloaded from the internet. The files are too large to attach to this post.

The Cyanide Industry, Robine & Lenglen, 1906

and

The Chemistry of Cyanogen Compounds an the manufacture and estimation, H E Williams, 1915

Fulmen - 19-4-2018 at 05:03

Thanks, I love those old books. Downloading as I type.

Fulmen - 19-4-2018 at 07:35

Robine & Lenglen only lists one interesting direct route: "Igniting nitrogenous substance with potassium sulfate, or with sulfur and potassium carbonate". No further details are given, so I question the feasibility of this route. The others require either cyanide, carbon disulfide or other toxic or hard to get precursors. Williams did not list anything useful.

Fulmen - 19-4-2018 at 14:49

Seems like the same holds true for ferrocyanides, either Na/k+Fe+C+N at high temperature or cyanides in some shape or form. Using atmospheric N sounds like a recipe for poor yields, and N-bearing organics would probably make for a very smelly afternoon. Neither sounds appealing.

Fulmen - 20-4-2018 at 00:18

What are my options for safe, non-stinking sources of nitrogen (no proper ventilation for the furnace)? Cyanuric acid has been mentioned, but I can't find it anywhere local. I can get TCCA, but converting it in any real amount would mean handling a lot of chlorine.

Melgar - 20-4-2018 at 01:02

Cyanuric acid forms when you heat urea to about 200˚C, giving off huge amounts of ammonia gas in the process. Surely you can obtain urea somehow? In the presence of alkali metals, you get cyanate salts.

Fulmen - 20-4-2018 at 01:43

*facepalm*
I only found coldpacks (expensive and unknown) and 600kg big bags. But AdBlue exhaust additive should be pure urea, and it's not that expensive either.
I assume sodium cyanate would be the best choice, seems like CA decomposes at 400°C while the cyanate is listed with a mp of 550°C. The cyanate could also be used for nitrocarburizing, something I could have a real use for.

Fulmen - 20-4-2018 at 10:23

OK, 4,7l of AdBlue aquired (should be more than 1.5kg of urea).

There seems to be some disagreement over which alkali salt is best: http://www.sciencemadness.org/talk/viewthread.php?tid=77330
However this writeup seems solid: https://chemistry.mdma.ch/hiveboard/methods/000442105.html

Melgar - 20-4-2018 at 11:06

Either K2CO3 or KOH should work. I think K2CO3 is preferred because it doesn't release or generate any water during the reaction. And the K salts are supposed to be a lot easier to purify and isolate, plus I'm pretty sure melting point is lower.

Fulmen - 20-4-2018 at 11:29

Don't really have much K2CO3 left, and no real source either. Found some KOH-based paint stripper, but I don't know if it's pure or not.

As for the risks of cyanides I've found several vague references to cyanates breaking down to cyanides at elevated temperatures, but nothing definite. I guess it should be OK as long as I avoid extreme temperatures or reducing conditions.

Fulmen - 20-4-2018 at 23:26

I spent some time reading up on the hydrolysis in water to see if it's possible to recrystallize. It seems like it's most stable above pH9, but I also realized that purification isn't really needed. I now have two uses for cyanates, as a precursor to thiocyanates and as a C-N-source for nitrocarburization. Neither will require high purity, in fact both uses will require addition of more carbonate. So as long as the yields are good I should be able to use the product as-is.

The next real hurdle will be a reaction vessel. Due to the scale I will be aiming for I need to deal with the ammonia gas. Using carbonates would produce a mixture of ammonia and and ammonium carbonate I guess, not really that useful. So the hydroxide-route seems like a good choice, unless I'm missing something.

Fulmen - 22-4-2018 at 02:14

I might be on to a cleaner reaction. From the attached PDF:
A method of preparing pure NaCNO is described, based on isomerisation of urea to NH4CNO under anhydrous conditions. Sodium metal is dissolved in dry butanol and urea is then added in eqnimolecular quantities; the mixture is refluxed. Insoluble NaCNO separates in almost theoretical yield.

The conversion rate is temperature dependent, so a high-boiling alcohol should be used. Glycerol or perhaps a glycol should be viable solvents. One problem is the formation of carbonates due to hydrolysis with the water produced by the reaction.

The article mentions the use of sodium butoxide as a water scavenger, obviously I need something a bit more OTC. Or perhaps the high bp. of glycerol is enough to drive out the water fast enough to reduce the problem?

Attachment: 1150-THE-PREPARATION-OF-PURE-SODIUM-CYANATE-FROM-UREAae22.pdf (398kB)
This file has been downloaded 1068 times
Edit: Perhaps something like CaO could work?

[Edited on 22-4-18 by Fulmen]

Melgar - 22-4-2018 at 04:50

Okay, first of all, you're making wrong assumptions all over the place. The best way to purify alkali cyanate salts is to use the fact that everything that isn't an alkali cyanate salt will BURN when you get it to the salt's melting point. Do you know what ammonium carbonate does when heated to 400˚C? It's a trick question, ammonium carbonate can't exist at those temperatures. Unless maybe you're somewhere near the core of Jupiter.

The advantage of K2CO3 is that first, the cyanate salt melts at a lower temperature, so your reaction vessel is more likely to hold up. Second, sodium salts have an annoying tendency to damage most things when they're molten, and potassium salts are somewhat better in that respect. K2CO3 is very cheap, and much cheaper than KOH. And certainly a fuckton cheaper than sodium metal.

The disadvantage of making the sodium salt is that at the higher melting point, the cyanide salt is more likely to form in significant amounts. It's not hard to separate them using standard chemistry methods, but those are problems that can be avoided by just using K2CO3 in the first place.

S.C. Wack - 22-4-2018 at 06:22

You may want to read what the Inorganic Syntheses entry for the cyanates has to say. IIRC it's in the second volume that can be found at a libgen or archive.org.

Perhaps I should mention that when I tried urea to cyanate, something sublimed, blocking the ammonia outlet, and the container failed inside the furnace.

[Edited on 22-4-2018 by S.C. Wack]

Melgar - 22-4-2018 at 08:08

Quote: Originally posted by S.C. Wack  
You may want to read what the Inorganic Syntheses entry for the cyanates has to say. IIRC it's in the second volume that can be found at a libgen or archive.org.

Perhaps I should mention that when I tried urea to cyanate, something sublimed, blocking the ammonia outlet, and the container failed inside the furnace.

[Edited on 22-4-2018 by S.C. Wack]

That'd be isocyanic acid, which would then trimerize to cyanuric acid or polymerize to some other random crap:

https://en.wikipedia.org/wiki/Isocyanic_acid

Fulmen - 22-4-2018 at 09:27

Melgar: K2CO3 might be cheap in theory, but that doesn't help me when the availability is fuck-all. It wouldn't be a major issue if I wanted to make 10g, but I might need as much as a kilo of the stuff. Also, the sodium salt seems to be the preferred salt for nitrocarburization (not sure why yet).

S.C Wack: Thank you for that reference. It basically confirms and expands on what my other sources tell me.

But I must admit the solvent-based route seems to be more promising if I am to make any significant quantities. First of all the temperatures will be much lower, allowing the use of standard glassware:

"The rate of reaction appears to be determined by the boiling point of the solvent. Thus there is a steady increase in the series methanol (b.p. 65°; 3o% conversion in 4o h), ethanol (b.p. 78°; 4o% conversion 4° h), propanol (b.p. 97°; 52% conversion in 15 h), butanol (b.p. 117°; 72% conversion in 2 h), isoamyl alcohol (b.p. 131°; 8o% conversion in 1 h)."

The question is if one can get decent yields without using sodium metal. One idea is to use boiling glycerin or glycol at reduced pressures.

Melgar - 22-4-2018 at 10:45

Quote: Originally posted by Fulmen  
Melgar: K2CO3 might be cheap in theory, but that doesn't help me when the availability is fuck-all. It wouldn't be a major issue if I wanted to make 10g, but I might need as much as a kilo of the stuff.

Oh, well why didn't you say so? Are you familiar with cream of tartar? Commonly used for cooking? It's potassium bitartrate. A pound of it costs about $7 where I am. So what you do is, you take that shit and burn it, making sure there's plenty of oxygen available. If you have KNO3, (probably not, but I figured I'd mention it) then you can mix that in too, since it'll release oxidizing radicals while latching onto the CO2 that's produced.

Keep the heat on until there are no traces of brown or black left, and it's all white. That's K2CO3.

Although I mentioned potassium bitartrate, any other potassium salt of an organic acid will work. Potassium sorbate, for example. Potassium benzoate, even. It's just a matter of how much crap you have to burn off. Even potassium chloride will form potassium carbonate if you mix it with charcoal and burn it. Even some dishsoaps can be burnt to leave a K2CO3 residue. And all you have to do to isolate it is dissolve it in water and filter it. Even wood ashes can be soaked and filtered to get K2CO3, since that's the only carbonate in wood ash that's water soluble.

I guess I just never would have thought that lack of K2CO3 could be a problem for anyone.

BTW, it helps us here if you put your country or state in the "location" field, so we can figure out what laws you're having to deal with.

Fulmen - 22-4-2018 at 11:31

There is virtually no OTC source of potassium at all around here except low-sodium salt. And considering the amounts I might need it's not really practical to make it from alternative sources like ashes/soap etc.

Melgar - 22-4-2018 at 11:50

Just curious: where exactly do you live? I'm trying to think of a place where you have access to sodium metal and butanol, but not eBay.

Not sure if you have access to black powder or road flares, but they typically produce potassium carbonate. Potassium nitrate and potassium perchlorate are used in flares, and are usually mixed with some carbon-based fuel, as well as a bit of strontium nitrate to make them red. Edit: The white smoke is potassium carbonate though, and I guess you'd need a way to capture that.

Even failing that, they sell 20-pound bags of potassium chloride to regenerate water softeners for people on low-sodium diets. Pulverize that up with charcoal in a trash can with holes in the bottom, light it, and leach the ashes.

I'm not trying to be mean, I'm just trying to keep you from thinking that sodium metal is somehow going to be more cost-effective.

There's actually an expired patent that details a cost-effective way to produce potassium carbonate with very little in the way of resources. Interestingly enough, it's US Patent #1:

https://motherboard.vice.com/en_us/article/pggjyb/the-first-...

[Edited on 4/22/18 by Melgar]

Fulmen - 22-4-2018 at 12:58

LOL! I didn't say I had sodium metal, but around here the range of useful OTC chems is pretty limited. Doesn't really matter where I live, I prefer to keep a low profile as this is an open forum. I also avoid ordering chems if possible, and importing is completely out of the question. It does limit my options but does keep the storm troopers off my doorstep.
I actually did try to extract K from ash once, I was helping a friend with a soap project. It was messy and the yield was low. Producing larger quantities that way doesn't really get my juices flowing.

I don't think sodium butoxide is required, it's primary function is as a water scavenger. From the article: "In one experiment with alcoholic NaOH the final preparation contained 75% NaCNO, and approximately 25% Na2C03." I could probably live with this yield if I have to, but I also think it must be possible to overcome this problem somehow.

Melgar - 22-4-2018 at 13:16

You certainly CAN use sodium carbonate, if you had no other choice, but I've used both, and 200˚C difference between the melting points makes a pretty huge difference, as far as how much of a pain in the ass it is. Potassium cyanate melts at a lower temperature, and acts as a solvent for the reaction to continue, whereas sodium cyanate is much harder to melt, and meanwhile all the isocyanic acid is vaporizing and making a mess.

Homebrewing stores and bakery supply stores will certainly carry either the carbonate or bicarbonate salt of potassium. Potassium bicarbonate is used instead of baking soda for people who are on a low-sodium diet, for example. It's just so ubiquitous that it's hard for me to imagine not being able to source it. Like not being able to source wax or copper or something. Just seems weird to me. But I guess I'll let it go now.

Fulmen - 22-4-2018 at 14:45

It is what it is. That's part of the reason I'm looking at the solvent route, if one can limit the presence of water it should be able to produce a fairly clean product with decent yields. And at only 150°C. I don't see why you're so hell-bent on the thermal route.

Melgar - 22-4-2018 at 17:30

Because the thermal route is so easy! You just mix stoichiometric amounts of urea and potassium carbonate in a cheap glass beaker with thick walls, put some foil over the top loosely so the ammonia and CO2 can escape, then heat it from underneath with a hand torch or something. (I used a hose clamp to hold the beaker to the side of a metal workbench.) All the impurities either vaporize or pyrolyze, and you're left with potassium cyanate that barely requires any purification. With potassium as your alkali, you can do it in glass, since the damage to the glass is minimal, and it only needs like 400˚C. Sodium attacks glass much more already, and when you add to that the fact that you need to get it to like 600˚C, it's really not safe at all to use glass, and you need some special crucible.

Fulmen - 22-4-2018 at 23:23

OK, you do have a fair point. I don't really like pushing glass to 400°C, but that's not a deal breaker. It just sounds a bit impractical once you start scaling it up.
I will do another search for potassium salts, they are always useful to have around and I could at least make enough cyanate for some initial tests.

The solvent method certainly has it's own set of challenges, the primary being it's sensitivity to water: "With butanol containing 0.04% moisture, 1-2% carbonate was generally present in the sample. With butanol containing 0.3% moisture, the carbonate content rose to 10%.". Then again these numbers are probably from running the reaction over night, with a higher bp a few hours should be enough. Nor should it be necessary to run the reaction to completion, especially if one performs multiple batches.

I also realize that the sodium butoxide isn't a scavenger but rather an oxygen-free source. If a method for removing water from the reaction is found it might be possible to use sodium hydroxide. Reducing the pressure could work, an aspirator should be enough to prevent water from condensing back out. That also takes care of the ammonia, eliminating the need for a fume hood or doing the reaction outside.

Another benefit of this method is that it produces a powdered product, grinding down large quantities of fused salt isn't exactly done in 5 minutes.

Fulmen - 25-4-2018 at 14:07

I have found a source for K2CO3, but it's not local. I might be able to stop by in a few weeks time, but being impatient I decided to give NaOH a try.
I started by melting 20g in a hotplate in a stainless cup. The assumption was that if I could melt the NaOH I could dispense with the propane. I then added 35g of urea (5g in excess) in portions, not really the best idea I've had. Even outside with good ventilation the ammonia was a problem, and the mixture foamed severely. Before half the urea was added I was left with a tan, powdery solid, so I ground it up as best I could with the remaining urea and left it for 10 minutes.
The product was then broken up and washed twice with EtOH. The wash was yellow and tested strongly alkaline.

Net product: 32,65g (100,1% yield).

Now obviously this shouldn't produce perfect yields, so I might try another wash tomorrow. Or not, it doesn't really matter for now. I'll probably use this for making thiocyanate.

Fulmen - 27-4-2018 at 09:41

Next was a test to see if thiocyanate would form from cyanate and elemental sulfur. 10g of the crude sodium cyanate and 6g of sulfur was mixed together and heated on a hotplate. Initially some NH3 was produced, as the temperature increased SO2 began to escape. After 30 minutes the reactants had barely fused together, so I increased the temperature with a blowtorch. Once the mix melted it seemed to be self-sustaining, producing a lot of bubbling and some sparks that resembled burning charcoal. After the reaction had subsided the container was cooled and the product dissolved in warm water.
This produced a large portion of black insoluble matter and a yellow liquid with a very faint "sulfurous" smell. From what I could find thiocyanate should produce a blood-red complex with Fe(III) and a blue with Co(II), sadly I only got a red/brown precipitate with Fe and a whitish ppt with Co. Could it be polysulfides interfering?

I also tried the cyanate reaction with NaOH and urea again, this time with the ingredients finely ground and blended. The foaming was even worse this time, possibly due to more aggressive heating, so the reaction had to be stopped early. I assume it's because sodium cyanate is solid at the reaction temperature. Hopefully a larger reaction vessel and adding the reactants in portions will help.
This time the product was grey, understandable as the stainless steel was also blackened from the heat.

Fulmen - 28-4-2018 at 12:50

I also got around to test the sodium cyanate for nitrocarburization, and it shows real promise. Even a few minutes of heating with a blowtorch did produce a significant increase in surface hardness, next step will be to make enough for a test bath. Not sure how I should test the hardness, I do have a Rockwell tester but it's not well suited for this use.
Perhaps some sort of abrasive test would be easier to jerry-rig?
I should also try something like a salt spray test, the coating is supposed to improve corrosion resistance.

VSEPR_VOID - 28-4-2018 at 16:43

If you find a source of thriocyanates or ferrocyanides make sure to add it to the List of Chemicals and Materials Made by Sciencemadness.org Users

https://docs.google.com/document/d/1AoI2VA5L4bmFw2HwXS2OVYTV...

Fulmen - 29-4-2018 at 05:13

Yet another messy reaction with NaOH. This time I started with 1/3 of the urea mixed with the NaOH, it still foamed over making a righteous mess. I believe the production of solid NaOCN that's the problem, even with a blowtorch I could only get the material touching the sides of the vessel to melt. So while it's good enough to test nitrocarburization (shows real promise) it's not suited for large scale production. Guess I will have to get some K2CO3 after all.

The solution from the thio-reaction has started ppt out a yellow substance (I think, hard to tell as the solution is yellow as well), this will require more investigation. In theory it could be sulfur from polysulfides oxidizing in air, but there is barely any H2S evolved so I don't really think so.

Fulmen - 3-5-2018 at 10:26

Persistence pays off. I solved the foaming problem simply by adding the mixed reagents in small portions, I just use several stainless beakers and work slow. I can easily make 200g in an afternoon, so I should have more than a kilo in no time. I expect I will need some potassium salt as well, but I can at least start with the sodium salt.

I still need to figure out a suitable composition for the nitrocarburization bath. I can find examples for cyanide baths, but little on cyanate-based ones. The odd thing is that the baths are used in the 4-600°C range, and most ingredients have a melting point above this. Here are two examples of cyanide baths:
1: • NaCN, 30% • KCl, 39% • Na2CO3 or K2CO3, 25% • Moisture, 2%
2: • NaCN, 60% • KCl, 24% • K2CO3, 15% • Moisture, 1%

I'm suspecting some eutectic effects are in play, this could complicate the search a bit.

Melgar - 3-5-2018 at 21:34

If you're interested in OTC alkali cyanides, a popular preparation of mixed sodium/potassium cyanide (circa 1800s, anyway) used to be adding stoichiometric amounts of sodium carbonate to potassium ferrocyanide, then melting the mixture down. Supposedly iron oxide precipitates out, CO2 is released, and the molten salt is a mixture of sodium and potassium cyanide, which has a lower melting point than either one by itself.

Obviously, this is not one to fool around with, and should initially only be done at the smallest scale possible. But alkali cyanides are really not bad to work with at all, as long as you're very careful about having good ventilation and never exposing them to acids. (Unless you take all necessary precautions and know exactly what you're doing, anyway)

Cyanides tend to be used a lot in metallurgy because of how well they work for certain purposes. If you can't find any information on the use of cyanates, there's a very good chance it's because cyanides work a lot a better.

The fact that every preparation uses both sodium and potassium salts indicates it's definitely a eutectic type of molten salt mixture. If you feel comfortable attempting it, you could even do a one-pot preparation from potassium ferrocyanide, once you've ensured that this reaction does actually produce alkali cyanides.

Fulmen - 3-5-2018 at 23:39

I'm staying clear of cyanides. While they are more common (or at least have been around longer) it's actually the cyanate (formed in situ by oxidation) that provides the nitriding action. Cyanide-free baths exist and seems to be replacing cyanide, so I can't find any reason why it should be inferior.

The info I have found only mentions cyanates and carbonates, for all I know it could be as simple as that. The two cyanide baths I listed only contain 30-60% CN, but as the MP of all the constituents are above 550°C I suspect it's formulated to produce an eutectic. But it's also possible that the cyanate concentration must be kept low for optimal results.

And as I'm typing this I found data on the Sursulf-process (sulfur added for lubrication purposes):
CN-: <0,8%
CNO-: 36±2%
CO32-: 19±2%
K+: 24,5±2%
Na+: 20±2%
Li+: 1,25±0,2%
S2-: 2-10PPM

The Li is a catalyst and reduces MP, the Na/K-balance seems to be to reduce bath viscosity.

Then there is patent 4019928:
The bath preferably contains 25-57 weight % of cyanate calculated as cyanate ion, 0 to 30% alkali metal chloride and the balance carbonate and alkali metal ions. Preferably cyanide is omitted, but it can be present in an amount up to 5%. Without removing the waste salt, the bath is regenerated by adding melon, melam, or melem.

Now I'm getting somewhere.

Fulmen - 4-5-2018 at 23:47

I had a faint hope of finding a carbonate-free composition. Since the cyanate breaks down to carbonate during use a carbonate-free bath could simply be regenerated with an excess amide after use. But I assume there is a reason for it, so I think it's better to stick with something true&tested for starters. The Sursulf-bath ends up at 30% NaOCN, 30% KOCN, 20% Na2CO3, 20%K2CO3, I think this will be my starting point. I still need to get hold of some K2CO3, in the mean time I will research bath analysis.

There are 3 components of interest, CN-, OCN- and CO32-. The cyanide should oxidize to cyanate, but I might have to test for the buildup during testing. Aeration of the bath is generally required, but I have a faint hope that a small bath can work without it.
With a straight cyanate/carbonate-bath I should be able to balance the ratio simply by analyzing one, and the carbonate seems like the obvious choice.
"Practical Nitriding and Ferritic Nitrocarburizing 2003" by David Pye lists a method where the carbonate is precipitated with BaCl2, filtrated, dissolved with HCl and titrated with NaOH.
I would like to avoid the barium, perhaps calcium would work? And with a precipitate a gravimetric method should work as well, saving me a titration setup. A direct titration would be nice of course, but I assume there is a reason why it's not mentioned.
Analyzing cyanate seems too complicated for my taste, Pye's method is based on disassociation to ammonia and a Kjeldahl-type distillation.

As for regeneration urea will likely be a poor choice due to the baths operating temperature. Patent 4019928 mentions "melon, melam, or melem", but I can't find too much info on these. Cyanuric acid might work, I'll have to look into that.

eesakiwi - 5-5-2018 at 06:21

Quote: Originally posted by VSEPR_VOID  
If you find a source of thriocyanates or ferrocyanides make sure to add it to the List of Chemicals and Materials Made by Sciencemadness.org Users

https://docs.google.com/document/d/1AoI2VA5L4bmFw2HwXS2OVYTV...


I can get Potassium Ferrocyanide as a case hardening powder from a hardware store.
US$35 for 500gms (lb)

Fulmen - 5-5-2018 at 11:13

I did a quick test of my crude sodium cyanate by dissolving 10g in water and precipitating it with calcium nitrate. It seems like calcium cyanate has low solubility, as the initial ppt was far more than expected. After washing with a large amount of water I ended up with 2,5g (indicating 26% sodium carbonate) but I'm not sure if I got all the cyanate out.
I do have a fair amount of barium perchlorate, but I rather not use that if I have a choice. There are many other carbonates I could try, but I'm having a hard time finding solubility date for cyanate salts.

I do have another idea to test, and that is the hydrolysis of cyanate to carbonate and ammonia. This should give a weight loss of 18,5%, more than enough for a gravimetric analysis. I'm simply going to weigh a beaker, add a weighed sample and a bit of water, then boil it to dryness. In theory this should produce pure, dry sodium carbonate, right?

Boffis - 5-5-2018 at 11:22

All the data you need is in (for the nth time I quote this reference) "Cyanogen compounds, Their chemistry, detection and estimation" by H E Williams 1915 and 1948. It can be down loaded from the web easily.

Fulmen - 5-5-2018 at 13:24

I've looked through the 1915-edition (can't seem to find the 1948 edition), but I couldn't find anything really useful. Calcium cyanate is mentioned, but not any solubility data. And while it does list some methods for analysis they aren't what I'm looking for. I'm not really set up for analytical work, and I don't think I need very high accuracy for this. So I'm looking for a simple, rough test hat will give me a ballpark composition.
The direct hydrolysis idea might be a bit optimistic, but it's worth a try. Another possibility is to hydrolyze it in HCl, this should produce a mixture of NaCl and NH4Cl.

Update: The math for the last one looks promising.
The carbonate should convert to NaCl with a 10% increase in weight, while the cyanate should produce NaCl and NH4Cl with a 16% decrease in weight.
You have to understand that this isn't a bath that will be used very often, so setting up a lab for the analysis isn't really an option.

[Edited on 5-5-18 by Fulmen]