picramic acid from picric

Many thanks for the ideas about producing sodium sulphide Markgollum! I will surely give the carbon reduction of gypsum a try one day. It is good to know that producing (presumably) high purity picramic acid is possible using Na2S. The polysulphide reduction (NaOH and S) never gave good results, either with homemade nitrite, (Lead+NaNO3 --> CO2) or reagent grade nitrite. The step to produce Na2S myself however was too much for me, given the substantial danger of working with H2S and the uncertanty whether the picramic acid was the problem or the diazotization itself.



[Edited on 24-1-2011 by nitro-genes]

I have been reading about DDNP for about 6 years now, off and on, and it has always seemed sort of illusive. Reading posts here and from the old Rogue science forum it was clear that the big problem was obtaining a suitable reducing agent. The most common and effective selective reducing agents seem to be sodium sulfide, sodium hydrosulfide and ammonium sulfide. After reading one of the excerpts from an early twentieth century chemistry book posted on this forum, I decided to try generating hydrogen sulfide gas and bubble it through a sodium hydroxide solution in order to obtain sodium sulfide or hydrosulfide. The method involves heating a mixture of low volatility hydrocarbon with sulfur. In this case Vaseline was used, but I have seen examples where paraffin was specified. A pdf of the excerpt is also attached at the end of this post.

I have taken a lot of pictures, which should do most of the explaining. A Vaseline and sulfur mixture was put in a mason jar, which had a perfect hole cut in its top to accept a rubber stopper. A piece of mild steel brake line was used to lead and cool the hot gas from the generator to the clear vinyl tubing and bubbler at the other end. An aquarium air stone or bubbler was used as the diffuser.

Safety was of utmost importance, as very small amounts of this gas can kill you. It was a very cold (-20C) and breezy day. The generator was outside and I was inside, behind a glass door that is reasonable airtight, controlling the power to the hotplate remotely with a variac (variable auto transformer).

I increased the power to the hotplate gradually over the course of 30 minutes or so until a fairly steady, but slow, stream of bubbles was coming from the diffuser. I didn't increase the power to the hotplate at this point because I believed slower production was safer. After an hour and a half I was forced to stop as the sulfide/hydroxide solution started to freeze and the diffuser also froze which filled the backyard with H2S stink for 10-15 minutes. It was such a cold day, however, that no one was outside anywhere near where I was operating which was fortunate. Judging by the volume increase at least 3/4 of the sodium hydroxide was converted to sulfide. According to Wiki the yellow colour of the solution is from polysulfide contamination.

The last picture shows the reactants which were to be used for the following partial reduction producing sodium picramate. The sodium (hydro)sulfide solution is on the left and picric acid made (days ago) from aspirin is on the right.




Attachment: Hydrogen Sulfide Gas Preparation_prothiere1903_abstract.pdf (100kB)
This file has been downloaded 1156 times


[Edited on 27-1-2014 by Hennig Brand]

Next, a reduction of the ortho nitro group of picric acid, the one in the 2 position, was performed. Five grams of picric acid (from ASA) was dissolved in a methanol-water solution and neutralized with sodium hydroxide. The solution was made from about 75% methanol and 25% water by volume.

I found a page with the reaction equations, going from TNP to DDNP, which was part of a Chinese primary explosives presentation. I have included a jpg of that reaction equations page, as well as a link to the whole presentation.

http://www.cecd.umd.edu/documents/presentations/Hong-Kong/Yi...

The article I was reading when I performed the reduction is attached, and was also found on this forum. I believe Rosco posted it. I didn't follow it accurately unfortunately, because of not knowing the process well and being distracted. The reaction time was much longer than recommended by the article and it took me a long time to get the temperature up to where it should have been. I also was unsure of the concentration of my previously prepared sulfide solution and I am sure I didn't use enough. A ferrous sulphate solution was used to test for reaction completion, but in retrospect it was obvious that the endpoint hadn't been reached. I obtained about a 60% yield, and the article reported yields in some cases in the high 90's. I still have at least 3/4 of my sulfide solution, so I will try again.

The colour change going through the reduction reaction is really very beautiful. The product in the last picture is (or should be) sodium picramate.

Conversion of the sodium picramate to picramic acid was easily accomplished by dissolving it in water and acidulating with sulfuric acid. Pictures of the picramic acid produced, as well as the DDNP produced from it, will be shown in the DDNP thread.



Attachment: Picramic Acid using Sodium (Hydro)Sulfide J[1]. Chem. Soc., 1945, 663 - 665,.pdf (344kB)
This file has been downloaded 1865 times


[Edited on 27-1-2014 by Hennig Brand]

Quote: Originally posted by Hennig Brand

I am pretty sure all, or most, of the NaOH was converted to at least sodium sulfide and probably mostly to sodium hydrosulfide actually. I didn't actually test it, just judging by the quantities needed to reduce a give quantity of picric acid and the color of its solution, etc. Yeah, that was my thought too, that a little extra NaOH in the reaction mixture wouldn't be a problem most likely. [Edited on 30-1-2014 by Hennig Brand]

Sodium Picramate - Experimental

Some more sodium picramate was made and this time greater care was used to record some of the reaction conditions and quantities used. This process is not optimized and neither are the quantities used.

Five grams of picric acid (from ASA), of greater than 98% purity as indicated by melting point, was dissolved in 100 mL of methanol and heated to 50-55C. Sodium hydroxide (0.9g) was dissolved in 20 mL of water and with stirring was slowly poured into the alcoholic picric acid solution neutralizing the picric acid and forming sodium picrate. While maintaining the temperature of the reaction mixture between 50-55C, sodium (hydro)sulfide solution was pipetted in, with magnetic stirring, a mL at a time over the course of 20-30 minutes. The sodium (hydro) sulfide solution was made by bubbling hydrogen sulfide gas through an aqueous solution of sodium hydroxide. Four grams of sodium hydroxide was used for the solution. The amount of Vaseline and sulfur placed in the H2S generator was at least 50% more than would be needed if all H2S was absorbed and converted to dissolved sulfide salt. The volume of sodium (hydro)sulfide solution obtained was ~45mL. About 15 mL was used to reduce the 5g of sodium picrate to sodium picramate. Ferrous sulfide solution was used to test for reduction reaction completion. After the last sulfide addition, heating was maintained for an additional 10-15 minutes and then the heating was stopped and the reaction mixture allowed to cool. Once the reaction mixture had cooled to about 40C, about 75 mL of ice cold water was added and the beaker was set into an ice-water bath and cooled very quickly to 10C or below. The sodium picramate was then filtered out and washed with small quantities of ice cold water.

From 5g of picric acid, 3.8g of sodium picramate was obtained or about a 79% yield. My last experiment produced only a 60% yield. The biggest difference between the two experiments was that the first experiment used less methanol and more water as the reaction solvent/medium. A little greater care was also used when testing for reaction completion, though I am not sure I really got it right yet. There is a lag time, governed by the reaction time, between when the reducing agent is added and when it is consumed.




[Edited on 4-2-2014 by Hennig Brand]

reduction of styphnic acid with tin chloride

I am not yet getting anywhere near 100% yields with the method I am using, but I believe my picramate is very pure given the yields of DDNP obtained from it. A melting point could easily be taken, to determine purity, but my melting point apparatus is not close enough to get to for the next little while.

The sulfur contamination issue made me remember something I forgot to point out earlier. During the hydrogen sulfide gas generation process there was some sulfur sublimation and deposition going on. You can see it in the attached picture. There was no evidence that the sulfur made it past the cold steel tube, however, as even after 1.5 hours the vinyl tubing between the steel tubing and the diffuser/bubbler was perfectly clear showing no signs of deposited sulfur. Ambient temperature was about minus 20⁰C, of course, which would have really helped keep the sulfur vapor from traveling very far in the steel tube.




[Edited on 6-2-2014 by Hennig Brand]

The following came out of a college or university level lab manual. The lab manual has a series of hydrogen sulfide experiments. A pdf of the whole manual is attached below as well.

"Experiment C. Reaction between H2S (g) and NaOH (aq). Hydrogen sulfide reacts readily with 6 M NaOH. The reaction is:

H2S(g) + NaOH(aq)-> NaHS(aq) + H2O(l)

1. Pour 25 mL 6 M NaOH into a 150 mL beaker. Use the H2S(g) that remains from Experiments A and B or prepare a fresh syringeful of H2S as described above. It is unnecessary to wash the gas for this experiment.
2. Remove the latex syringe cap from the H2S-filled syringe, and suction a few millilitres of NaOH(aq) into the syringe. Hydrogen sulfide reacts instantaneously with the NaOH(aq). The plunger may move rapidly inward and/or the NaOH solution will be drawn rapidly into the syringe. The reaction is so rapid, it could be surprising. The beaker is used because the beaker walls will contain any splashed NaOH(aq). This solution can be slowly discarded down the drain with large amounts of water. If possible, use a sink in the hood and wear gloves."

It seems I was making some wrong assumptions earlier in this thread. Sodium hydrosulfide is what I think is produced exclusively by bubbling hydrogen sulfide into a sodium hydroxide solution, even though many webpages say that it is sodium sulfide. It is also nice to know that hydrogen sulfide is converted almost instantly on contact with the sodium hydroxide solution. I am much less nervous to make sodium hydrosulfide solutions now. If made in small quantities, done outside and using a little common sense the risks are manageable. However, one should probably arrange to be somewhere else when the gas generator is producing. I used twice as much sodium hydroxide and water this time, because I wanted more sodium hydrosulfide and twice the column height meant much greater contact time between the hydrogen sulfide and the sodium hydroxide solution (though it may not matter much).

Also attached is the first page of the section on Sodium Hydrosulfide from "The Encyclopedia of Chemical Technology". I just took a picture of it with my phone and attached it as a jpg. I might get a better scan of the whole section later. From reading the sodium hydrosulfide and sodium sulfide sections, bubbling hydrogen sulfide gas into a sodium hydroxide solution is the process used to make sodium hydrosulfide not sodium sulfide.

Attached is a picture of about 90mL of sodium hydrosulfide solution freshly prepared. It is a little cloudy this time which is different than last time. A little sulphur vapor made it over into the solution maybe.


Attachment: H2S Hydrogen Sulfide Experiments.pdf (150kB)
This file has been downloaded 1340 times





The sodium hydroxide solution, which H2S was bubbled into, was about 80 mL and contained 8g of sodium hydroxide (about a 2.5 molar solution of sodium hydroxide).


[Edited on 18-2-2014 by Hennig Brand]

Ideas For Reducing The Poisoning Risk Associated With Hydrogen Sulfide Generation

Apart from dealing with toxic H₂S, the process is very straightforward giving high yields and, yes, a very pure product. From what I have read it seemed like the best option to obtain a high purity sulfide reducing agent. Hydrogen sulfide is extremely poisonous, though, and I won't be doing this again until I have found or developed a little better process. I still have a bit of a sore throat, and feel a little unhealthy, and that is from fairly mild exposure close to 2 days ago. Very dangerous stuff indeed!

In the absence of a good fume hood a strong electrical fan could be placed just over the generator as it was disassembled. What would be even better would be to have the type of blower fan that connects to large flexible air hose so that the H₂S gas could be lead safely away from the work area during disassembly. With a decent fan, and maybe also a H₂S rated gas mask, disassembling the gas generator would be much safer.

Ok, I think I have a simple yet effective solution to the poisoning by gas generator disassembly problem. Hydrogen sulfide is fairly soluble in water (ca. 0.12 mol/l or 4.1g/l at 20⁰C & ca.0.15 mol/l or 5.1 g/l at 10⁰C). Values taken from "The solubility of gases in distilled water and seawater-V. Hydrogen Sulfide", by A.A. Douabul and J.P.Riley

The volume of hydrogen sulfide in the glass mason jar used was about 450mL (or less) and was not under pressure. Using the simple 22.4 L/mol rule for ideal gases at STP, it was found that the mason jar could contain as much as 0.02 moles or about 0.68g of H₂S. Given that 170 mL of water at 20⁰C can absorb that amount of hydrogen sulfide, a 2 gallon bucket of cool water would absorb it easily. The idea was to let the reactor cool undisturbed and then immerse it in a bucket of water. Two bricks could be placed on the bottom of the bucket, with space between them. The generator vessel (mason jar) could be flipped upside down under water before the rubber stopper or top was removed and then placed over the two bricks so that the water could contact the gas. Another brick or dense object would need to be placed over the top of the jar to hold it down. A little sodium hydroxide could be added to the bucket of water beforehand if converting the hydrogen sulfide to sodium (hydro)sulfide seemed advantageous. A little sodium hydroxide in the water would greatly increase the speed of absorbtion.

The process could maybe be improved, but I think the general idea is sound and would greatly increase safety. I guess a person could hold the jar under a sprinkler system when it was opened. Or even better a sprinkler system that was spraying sodium hydroxide solution (ouch!) Shouldn't make jokes, but I couldn't resist.

The smartest thing, now that I am starting to fully understand the seriousness of the poison issue, would be to use a much smaller vessel so that a much smaller quantity of poison gas would be left in the vessel after production was finished. The reason a larger one was used was because it put the screw top with stopper well away from the heat source below. It also allowed most of the sublimed sulfur to deposit in the jar so that it didn't end up in the delivery tube. The 16 oz size is also the most common for a mason jar (I think).


Ok, I think this is the best idea yet. The top of the mason jar, or rubber stopper, could have installed a small hose barb the size that is needed to attach an aquarium pump air line. This would either have to have a small valve or some other way of keeping it closed during the generation process and be safely opened after generation. When gas production is complete, simply attach an aquarium air pump and pump air into the reactor which will dilute and carry (or force) the H₂S out of the generator and through the bubbler where it is safely absorbed. It might be a good idea to use a fresh sodium hydroxide solution since the one used during generation should be very saturated or close to it and therefore have less absorbtion ability. Also pumping air (oxygen) into the prepared sodium hydrosulfide solution will convert some of the sodium hydrosulfide to polysulfides, which is also undesirable.


[Edited on 20-2-2014 by Hennig Brand]

H2S Hydrogen Sulfide Toxicity Information

Picric Acid Solubility in Methanol-Water Mixtures

Attached is a jpg of a snip-it from a journal article regarding solubility of picric acid in various solvents. The article title is, "The solubility of Picric Acid in Mixed Solvents Part I. Water-Alcohol and Water-Acetone Mixtures".

The table should serve as a guide when trying to decide how much methanol to use, if the method involving methanol is used.



[Edited on 23-3-2014 by Hennig Brand]

Ferrous sulfate is cheap and safe.

9 moles FeSO4 solution and 18 moles NH4HCO3 or NH4OH or perhaps just use NaOH solution added separately and simultaneously in a parallel proportional molar addition with stirring per mole of warm 30C sodium picrate solution should be about right for the reduction, with subsequent raising to maybe 80C and decanting and filtering of the hot solution, and acidifying to slight acidity and near complete decolorizing of the solution to precipitate the only slightly soluble free picramic acid. The picramates are color dyes so the endpoint is visual on the acidification. I better check my stoichiometry because I could be off and just did this one in my head. Don't bother checking the literature because it probably isn't there.

The strategy that would be best I think is to calculate a drip rate for the ferrous sulfate into the slightly basic sodium picrate solution and then make the additions of ammonium bicarbonate as solid sprinkled into the stirred mixture at a rate that stays slightly ahead of the ferrous sulfate to keep the reaction mixture slightly basic. Similarly if using a liquid base the additions of base should lead by a small amount the ferrous sulfate addition in order to keep the reaction system alkaline. Whether magnetic stirring may be good here or not is unknown since the byproduct Fe3O4 is going to be magnetic. The magnetic interaction could be helpful or interfering, but my best guess it could be helpful and may present no issue.

[Edited on 25-3-2014 by Rosco Bodine]

Quote: Originally posted by Rosco Bodine

Ferrous sulfate is cheap and safe. 9 moles FeSO4 solution and 18 moles NH4HCO3 or NH4OH or perhaps just use NaOH solution added separately and simultaneously in a parallel proportional molar addition with stirring per mole of warm 30C sodium picrate solution should be about right for the reduction, with subsequent raising to maybe 80C and decanting and filtering of the hot solution, and acidifying to slight acidity and near complete decolorizing of the solution to precipitate the only slightly soluble free picramic acid. The picramates are color dyes so the endpoint is visual on the acidification. I better check my stoichiometry because I could be off and just did this one in my head. Don't bother checking the literature because it probably isn't there. The strategy that would be best I think is to calculate a drip rate for the ferrous sulfate into the slightly basic sodium picrate solution and then make the additions of ammonium bicarbonate as solid sprinkled into the stirred mixture at a rate that stays slightly ahead of the ferrous sulfate to keep the reaction mixture slightly basic. Similarly if using a liquid base the additions of base should lead by a small amount the ferrous sulfate addition in order to keep the reaction system alkaline. Whether magnetic stirring may be good here or not is unknown since the byproduct Fe3O4 is going to be magnetic. The magnetic interaction could be helpful or interfering, but my best guess it could be helpful and may present no issue. [Edited on 25-3-2014 by Rosco Bodine]

I should amend my earlier "first guess" stoichiometric ratios on this to 6 moles of ferrous sulfate being needed for reduction since the reported byproduct of the reduction is the red iron oxide Fe2O3 instead of the mixed ferrous-ferric Fe3O4 byproduct I was guessing would be the byproduct. It seems possible that some of both could form as the byproduct and if this is true then the actual amount of ferrous sulfate required could be a bit more than the 6 moles but less than the 9 which would be the worst case requirement.

Reading the period literature from 150 years ago for clarification on the terminology is helpful.

The reference by Aime Girard posted earlier here
http://www.sciencemadness.org/talk/viewthread.php?tid=433&am...



Terms that are archaic "protosulphate of iron" is ferrous sulfate and the "protoxide" is ferrous hydroxide and the abundant precipitate of "peroxide of iron" is Fe2O3 or red iron oxide which would likely form first as a perhaps gelatinous transiently stable hydrated ferric oxide then decomposing in the hot reaction mixture to the neat oxide which should be a particulate precipitate. Girard's mention of the "ammoniacal liquid" would indicate that the base being used was ammonium hydroxide or ammonium bicarbonate. But it would seem that other bases could be used. Sodium hydroxide should serve as well. Milk of lime has been described in the literature and has the barium hydroxide also been used for precipitation of the sulfate value from the ferrous sulfate. There could be advantage to removal of the sulfate value from conversion of the ferrous sulfate to ferrous chloride, using calcium chloride and filtering out the byproduct calcium sulfate. The needed amount of hydrated lime could be neutralized with HCl and then the solution of CaCl2 and FeSO4 could be mixed and the FeCl2 solution from filtration could be used for reduction, gradually added to a hot solution of sodium picrate, with an added 2 equivalents of sodium hydroxide, with the byproducts of reduction being sodium chloride remaining in solution, and a mixed precipitate of ferric oxide and sodium picramate. In the alternative using NH4OH as the base could have advantage due to the much greater solubility of the NH4Cl byproduct than NaCl. Additional dilution water could be added and the completed reaction mixture heated to dissolve the sodium picramate, decanted and filtered. Addition of HCl should precipitate the nearly insoluble free picramic acid from the residual NaCl or NH4Cl solution.







The reaction appears to be

C6H2(NO2)3OH + NaOH ----> C6H2(NO2)3ONa + H2O

( 6 X ) [ FeSO4 + 2 NaOH ----> Na2SO4 + Fe(OH)2 ]

6 Fe(OH)2 + C6H2(NO2)3ONa ----> C6H2(NO2)2NH2ONa + 5 H2O + 3 Fe2O3

To provide an excess of theory of the reducing agent to compensate for atmospheric oxidation and other possible process depletion losses it would probably be fine to use a 7 X quantity of the reactants for the second equation and use a couple of per cent excess of theory beyond for the NaOH to assure the reaction mixture stays distinctly basic. The sodium picramate end product is itself a pH indicator which will provide a shift from an intense dye dark red color to a nearly colorless supernatant liquid when the completed and diluted and filtered hot reaction mixture is acidified to precipitate the free and nearly insoluble dull orange brown color picramic acid.

It would also seem possible that a lesser amount of ferrous sulfate could be used if dextrose or fructose along with the needed amount of NaOH to enable its use as a reducing agent was used in substitution for the omitted quantity of the ferrous sulfate and NaOH, since the reducing sugar itself will accomplish the reduction of sodium picrate to sodium picramate directly, or will act indirectly and in parallel to reduce the ferric hydroxide byproduct back to ferrous hydroxide which should react again with the unreacted sodium picrate as a reducing agent regenerated by the reducing sugar. Likewise should occur for regeneration of the ferrous hydroxide regardless of its source being sulfate or chloride or acetate, it should be regenerated by basified glucose being freshly added to the reaction mixture.

The earlier estimation for a possible requirement of 9 moles of ferrous sulfate for the reduction of 1 mole of sodium picrate was uncertainty about the mechanism that would predominate for the reduction. There is a Schikorr reaction
which is 3 Fe(OH)2 → Fe3O4 + H2 + 2 H2O

http://en.wikipedia.org/wiki/Schikorr_reaction

To supply the 6 hydrogens needed for the reduction by this Schikorr mechanism would require 9 moles total of the ferrous sulfate. And the byproduct would be the mixed ferrous-ferric oxide Fe3O4.

[Edited on 28-3-2014 by Rosco Bodine]

Quote: Originally posted by Rosco Bodine

If I had to guess, the reduction of sodium picrate using sodium ascorbate it likely going to be the best method. [Edited on 28-3-2014 by Rosco Bodine]

It is a convoluted imprecise description article that could have been written better, and it becomes obvious there is error for example 6 in table 1 where the ratio should be 1.15 to 1 and it looks like a transposing error from example 4 above in the same table. The arithmetic I did showed a ratio of 2.273 to 1 for what is shown on the table as 2.3 to 1 rounded. Anyway I think the patent I posted earlier about the reactions involved when gassing NaOH with H2S is absolutely correct that the Na2S forms first and this would square with other process chemistry studies which have shown the trend of higher sulfides forming first preferentially and then the sulfur value lowering as forced by conditions favoring change to what would ordinarily be a less stable sulfur compound under ordinary conditions. The sulfur compounds have varying stability particularly for solutions which are subject to instability after bottling and what begins as a pure solution of one sulfur compound gradually changes to an equilibrium mixture of various sulfur compounds and can even precipitate sulfur in the process of reaching an equilibrium. The composition of the mixtures that are formed is algebra that is affected by time and temperature. I was trying to get a handle on this to guesstimate what may be the best strategy for sequencing a dry method reaction of solid NaOH and solid sulfur with minimal water and calculating the dilution and temperature and time for reaction, sequenced with subsequent additions of added base as hydroxide or bicarbonate or carbonate as would produce an optimized mixture of reducing sulfur compounds for the Zinin reduction, where the byproduct would be soluble thiosulfate and no precipitation or minimal precipitation of free sulfur as byproduct would occur, and no sulfur dyes as byproduct with picramate would be an issue. I still think there is an optimization possible for such a sulfide route, but the algebra for achieving it would likely come down to trial and error. The math only is helpful so far, and it complicates things when the math reported by others is wrong for typos or editing errors, leaving a mystery to figure out. That's what we have here.

The appeal of the sulfide reduction is the concentrated reaction mixture and the rapidity and high yields possible at 98% or possibly even the quantitative yield is attainable.
And I think high yields are possible even using dry method produced polysulfide Zinin reagents which are not reported so far as I know except by me And maybe my math is not so good sometimes. I shall continue to scratch my head and think on this some more. But there are possible also interactions with manganous and ferrous salt used concurrently with sulfide, and variants that suggest quantitative yields are possible with the ferrous salt. The reference I posted regarding the formation of basic ferrous salts I think provides information which may be useful ultimately in a strategy to utilize the ferrous salt as a regenerable reducing agent, possibly regenerated even by sugar. Avoiding the sulfide entirely even though it is efficient is desirable since it seems to be a "nasty" synthesis leaving some impurity even in those procedures where a high yield is produced. A cleaner and more pure product that could still be easily synthesized is what motivates examination of the alternative methods.

Sulfur chemistry is a journey into the tall grass While on safari there, probably good to bring a gun and matches.

[Edited on 31-3-2014 by Rosco Bodine]

Quote: Originally posted by Hennig Brand

The first attachment "Communications” is a very good article, thanks for posting. The conclusions they came to ring true for me, though they went into more depth than I did with regards to theory. I think when bubbling hydrogen sulfide into a dilute sodium hydroxide solution, sodium hydrosulfide is formed without going through sodium sulfide. Either that or the sodium sulfide is a very short lived intermediate. Sodium hydrosulfide is formed by the half neutralization of hydrogen sulfide. I also think that a large excess of hydrogen sulphide gas must be bubbled through the solution, with the simple setup used, or else all the sodium hydroxide will not be converted. I think that the excess sodium hydroxide causes undesirable reactions and loss of yield. The next time I prepare sodium hydrosulfide I am going to be much more careful and use a large excess of hydrogen sulfide.

Yes I agree the Clayton article Communications does pretty well specify conditions that would be useful for a scheme where the NaHS Zinin reagent is prepared by use of a H2S gas generator to treat an NaOH solution until weight increase indicates conversion to NaHS. This is a parallel type of gas / liquid phase reaction scheme as is used for producing sodium hypochlorite or sodium bicarbonate, monitoring the completion point by weight and possibly by the increasing blowthrough that is unreacted and reaches the scrubber. Bubblers are good for visual flow monitors for these type reactions. And having the reaction flask sitting on a scale for weight increase monitoring is greatly helpful.

Reviewing this thread I see that 7 years ago I was already getting a pretty good handle on the approaches that could be workable for a laboratory scale, and had an understanding of the milder conditions that would be provided by using H2S generation for producing either Na2S or NaHS. I described this method for Na2S on page 4
http://www.sciencemadness.org/talk/viewthread.php?tid=433&am...

And the last couple of references clear up any misinterpretation about the composition of the "hybrid mixture" of Na2S and NaHCO3 used to increase the yields to nearly quantitative. There is a buffering scheme at work there I think where the same equilibrium mixture of NaHS and Na2S could be made using the H2S gas generation scheme, but interrupting the conversion of the initially formed Na2S before it is completely converted to NaHS by measured weight gain. If you follow what I am describing, it would be even easier to make the hybrid reagent if the conversion of the initially formed Na2S was stopped when
approximately two thirds of the Na2S had been converted to NaHS. The amount of NaHCO3 to be added to the methanol and sodium picrate aqueous mixture would be reduced by the same two thirds as that amount would not have been needed for conversion of Na2S to NaHS in methanol, having been accomplished by variation of the composition of the hybrid mixture done deliberately by modulating the H2S flow to the Na2S solution formed initially.

Such a modification of "incomplete conversion" of the Na2S to NaHS by interrupting the H2S gas flow, combined with adding methanol to the aqueous sodium picrate reaction mixture, could increase the yields of 90% reported by Clayton using NaHS in the Communications article, to the 98% yields reported by Hodgson and Ward using the equilibrium "hybrid" mixture of NaHS and Na2S made in situ in methanol by adding NaHCO3 to Na2S. Such approach should shorten the time required for working the H2S gas generator and would be convenient to interrupt at a stage while it was still freely absorbing in the still reactive Na2S not yet nearly completed conversion mixture. The only modifications otherwise would be the reduction in needed NaHCO3 and the addition of methanol, to emulate the higher yielding reaction conditions described by Hodgson and Ward.

So I think that surveying the available literature with better comprehension in light of the most recently obtained references, allows for proposing such a modified scheme to improve the yields to nearly quantitative.

Such a scheme of using the H2S gas generator is the most technically precise method of performing this type reaction. And this is likely the approach that will produce the highest yields. However I still believe it is possible to further optimize the "dry method" of mixing solid NaOH and S and reacting with added heat and a sequence of manipulations that can provide a useful Zinin reducing agent of generally predictable composition as the result, if it is desired to avoid a scheme requiring use if a H2S gas generator. There would be an extensive experimental workup required for study and charting results of such a scheme that would be good to know as an alternative method, and what level of efficiency to which that less technically superior method is limited. Reduced yields may be acceptable for an improvised scheme where the additional equipment needed for the best method is not available or is inconvenient.

As an additional reference that further describes some of the reactions and crystallization aspects related to isolation of Na2S I am attaching US3666410 for its information value.

Science Madness has done it again here I think with a more comprehensive review of the literature regarding an obscure subject, researched over a period of many years to arrive at a good general picture of the chemistry involved. Google sure knows who we are when it comes to certain information and discussion of historically obscure and technically interesting topics for which compendiums of good information are simply otherwise unavailable Call this thread a keeper.

Attachment: US3666410 Na2S Crystallization Process.pdf (578kB)
This file has been downloaded 930 times

[Edited on 3-4-2014 by Rosco Bodine]

The relationship of Na2S and NaHS is the same as the relationship of Na2CO3 and NaHCO3 or Na2SO4 and NaHSO4 ......NaHS is the "bi" or acid salt as opposed to the neutral salt. Na2S is the neutral salt. NaHS is the acid salt, which could also be called Sodium acid sulfide, or perhaps sodium bisulfide, or sodium hydrosulfide. Sodium Sulfide and Sodium Hydrosulfide are not chemicals as frequently and commonly encountered as more familiar salts. Therefore it isn't immediately recognized as would be the case for salts like the carbonates or bicarbonates particularly.

Time for a song or two illustrating "aspect ratios"

<iframe sandbox width="640" height="360" src="//www.youtube.com/embed/78O6--THTF0?rel=0" frameborder="0" allowfullscreen></iframe>

<iframe sandbox width="640" height="480" src="//www.youtube.com/embed/GfJWqjoekow?rel=0" frameborder="0" allowfullscreen></iframe>

[Edited on 3-4-2014 by Rosco Bodine]

Sodium Hydrosulfide and Sodium Sulfide Equilibrium

It still seems logical that having the reducing agent being made up mostly of sodium hydrosulfide would be best. If you look at the reaction equation where sodium sulfide is used there is a mole of sodium hydroxide formed for every mole of sodium sulfide consumed. This seems to be the reason that sodium sulfide, when used by itself, results in such low yields. If sodium hydrosulfide is used as the reducing agent, the amount of sodium hydroxide in the reaction mixture will be kept very low since the sodium hydrosulfide (acid salt) will neutralize, or mop up, the sodium hydroxide produced by the reaction and form sodium sulfide in the process.

Still a little confused as to why sodium sulfide mixed with sodium hydrosulfide would actually produce higher yields than using just sodium hydrosulfide by itself.

Reaction equation taken from Chinese explosives presentation slides attached.





[Edited on 4-4-2014 by Hennig Brand]

Quote: Originally posted by Hennig Brand

It still seems logical that having the reducing agent being made up mostly of sodium hydrosulfide would be best. If you look at the reaction equation where sodium sulfide is used there is a mole of sodium hydroxide formed for every mole of sodium sulfide consumed. This seems to be the reason that sodium sulfide, when used by it, results in such low yields. If sodium hydrosulfide is used as the reducing agent, the amount of sodium hydroxide in the reaction mixture will be kept very low since the sodium hydrosulfide (acid salt) will neutralize, or mop up, the sodium hydroxide produced by the reaction and form sodium sulfide in the process. Still a little confused as to why sodium sulfide mixed with sodium hydrosulfide would actually produce higher yields than using just sodium hydrosulfide by itself. Reaction equation taken from Chinese explosives presentation slides attached.

The hydrosulfide ion works fine as the principal reducing agent but if used alone appears to be yield limited to 90% unless added methanol is the trick that would account for increasing the yield. However it is just my guess really that pH is what is having the effect and is allowing for a secondary and tertiary reaction system involvement when the reaction system is more basic. If you look at the 3 possible Zinin mechanisms, then you can see that any free sulfur byproduct from the reduction by the hydrosulfide ion would react immediately with any normal sulfide present to form sodium disulfide Na2S + S ---> Na2S2 and then
the Na2S2 would subsequently contribute to the reduction as well, but that mechanism would then have its byproduct oxidized sulfur tied up as thiosulfate, similarly as would occur for the byproduct oxidized sulfur of the normal sulfide.
By using a mixture containing the normal sulfide along with the acid sulfide, you are enabling all 3 possible Zinin mechanisms to be occurring, in a mixture where there is sufficient Na2S present that any free sulfur byproduct from the hydrosulfide ion mechanism to be recycled through formation of soluble disulfide instead of accumulating as an insoluble precipitate or impurity in the desired product. By having some Na2S present it assures that the sulfide oxidation byproduct from the reduction will occur as soluble thiosulfate or as unreacted soluble Na2S2 at the completion of the reaction, instead of having a cloudy sulfur milk of precipitated free sulfur byproduct as an undesired contaminant for the sodium picramate. The higher reaction kinetic value for the Na2S2 is likely what increases the yield beyond the 90% which is obtained using hydrosulfide alone as the reducing agent.

The article posted earlier, Reduction of p-nitrotoluene by aqueous ammonium sulfide provides a good insight into the 3 mechanisms which are involved and partly interdependent and sequential, having differing kinetics and occurring at different rates.



Therefore, for a mass reaction as would proceed from a mixing all at once of the reactants "in a lump" addition poured together to react, there is going to be a most favorable proportional mixture of NaHS and Na2S that should give the highest yield. Some mathematical prediction of the likely composition for such a mixture can be done, but the fine tuning of the mixture to identify the precise optimum ratio of acid sulfide to normal sulfide could be done by experiment.

[Edited on 5-4-2014 by Rosco Bodine]

I see what you mean about the elemental sulfur by-product when pure sodium hydrosulfide is used as the reducing agent. Also, I think I will save the methanol and use water instead for my next experiment. Between us, I think we may have most of the pieces needed for the efficient production and usage of picramic acid and DDNP. Getting this DDNP thing ironed out has been a long time coming for the hobby community. (Stops to pat self on back)

[Edited on 5-4-2014 by Hennig Brand]

So are you getting your glassware all dusted off for a series of experiments?

No, but I wish my other priority chores were out of the way so I could do the experiments. One aspect about experiments I like is that things that have been seen cannot be unseen. The Zinin reaction is complex and has never been truly identified as a single mechanism but can be 3 or more concurrent reactions, and I believe there are likely 4th and 5th reactions that add to the list that would be Na2S3 and N2S4 and I don't believe the 3 finite mechanisms identified are absolutely governing, but are generalizations of what are predominating reaction mechanisms. For example when a "dry process" reagent is made and modified in whatever subsequent ways that affect the sulfur rankings of the mix of reducing compounds that results, it would seem that a mixed system of various polysulfides and sodium sulfide results which can still be a useful Zinin reagent as a mixed system that defies analysis except according to inexact modeling as applies to the Zinin reduction itself. On a mathematical basis trying to identify with any precision what is the exact composition of a polysulfides mixture and how that mixture accomplishes a reduction is an abstract algebra that compares with trying to nail jello to a tree. Zinin reactions being explored comes down to experiments that are based upon some murky mathematics and end up being a whole lot of "try this and see what results"

For experimental results we have 3 possibilities:

[A] Yeah it worked just like I figured

[B] Well it worked on paper

[C] WTF!

And of course there may be a combined experiment result of indefinite proportions of the above three, or a not yet identified alphabet soup of other possibilities which follows.

Rubicks cube and the Zinin reaction have commonality as hair pullers, yet there are solutions where everything seems to line up and the jello does get nailed to the tree. A quantitative yield is that point, but the exact reactions are likely to remain enigmatic. So there is an appreciation of the difficulty of Hodgson and Ward giving precise explanation of what was actually occuring in their own reported experiment, understanding that they didn't really know

It is almost a tongue in cheek academic joke when a high yield Zinin reduction is reported, to with great interest query the reporter as to exactly what mechanism they attribute their results, knowing they are being assigned a perplexity.

Zinin likely understood the abstract mechanism aspect himself and probably derived great satisfaction from publishing a puzzler that would still be puzzling a hundred years and more later.

This is a convoluted description, but I will try to explain generally how it is that even though following the method of Clayton, it is unlikely that the actual reducing reagent predominating throughout the reduction will be the NaHS that is the starting material.

The potential variant reactions possible in a Zinin reaction system can become complex. For an example, take the third reaction mechanism of hydrosulfide and see that the byproduct of 3 S and 3 NaOH could subsequently react to form an indefinite mixture of Na2S3 + NaOH and IMO it likely does react further since the initial "nascent sulfur" byproduct would be highly reactive and would further react with whatever it could. And the Na2S3 + NaOH as only one example could react with unreacted NaHS present, such as the following reaction/s:

Products of mechanism 3 are 3 S and 3 NaOH from the 3 NaHS used to reduce 1 mole of sodium picrate

The NaOH will react instantly with unreacted NaHS to form Na2S which will further react with the "nascent sulfur" to form Na2S2 and the Na2S2 could also react with more "nascent sulfur" to form NaS3 and possibly Na2S4

(3) NaOH + (3) NaHS <-------> (3) Na2S + (3) H2O ( + 3 S ) -----> (3) Na2S2

According to this scheme then one half of the NaHS reacting according to mechanism 3 would convert in situ to Na2S2 the remaining half of the NaHS for that mechanism as a result of the byproduct NaOH causing conversion of the unreacted NaHS to Na2S and reacting with free sulfur to form Na2S2. So it is clear that even though you may start a reduction using only NaHS alone as the Zinin reagent, because of the complex subsequent chemistry involved, when only the first half of that NaHS has reacted, the remaining half is already converted to other higher sulfides that are not NaHS but mathematically according to theory would be Na2S2.

The Na2S2 formed is not a loss of reducing power for the second half-reaction of what was the original NaHS but is a multiplier of 3X the reducing power of the first half of the NaHS consumed. So the overall reduction becomes one where half of the NaHS is consumed accomplishing only one quarter of the overall reduction, while the remaining three quarters of the reduction are accomplished by conversion of the unreacted half of the NaHS reacting with decomposition byproducts of the first reacting half of the NaHS.

If the Na2S2 does not react in the reduction immediately,
then it can itself react with the nascent sulfur byproduct
and it can further react with other materials present as follows:

Na2S2 + S -------> Na2S3 However the Na2S3 can further react also and be converted back to Na2S2

Na2S3 + NaOH + NaHS ----> 2 Na2S2 + H2O

The potential variant reactions like this soon makes it clear that a Zinin reaction is complex by its nature where the reaction system is in flux with an indefinite mixture of reactions and subsequent reactions possible and probable.

I can guess that the reaction mechanism 3 is expected then
to provide a subsequent path to reaction mechanism 2 even for a reduction where the Zinin reagent being used for the reduction would be NaHS alone. The reaction would likely not just neatly complete according to mechanism 3 but would "recycle" according to mechanism 2. So where does the free sulfur byproduct come from? Obviously the "recycle" does not occur or does not occur completely, depending on the concentration and temperature and pH. But I would bet good money that thiosulfate would be detectable in the completed reaction mixture, even for a reduction done completely according to Clayton, bearing out what I describe.

How you can factor these variants into the estimated theoretical model is not yet nearly an exact science.

Way to go Zinin, nothing like keeping things simple, huh.

A Zinin reaction defies analysis in a way that is so artful that at the completion of the analysis and reaction modeling, the only firm conclusion is likely to be that yes there is definitely sulfur involved in the reaction, just like I thought I wasn't kidding when I said earlier that sulfur chemistry of this sort is a safari into the tall grass.

[Edited on 5-4-2014 by Rosco Bodine]

A deeper look into the Zinin reaction

There is some more analysis I have done in reviewing this topic and looking at my old notes. Mower time for the tall grass. To analyze this reaction further and avoid a lot of fractional odd number molar quantities that are harder to follow for visualizing these molecule parts is easier with whole lego blocks, so there is applied an algebraic multiplier across the reactions so that the stoichiometry comes out as whole number values easier for me to then inventory the results for these parallel mechanisms for the range determination of NaHS required by theory.

Based upon the earlier described process the Zinin reaction mechanism 3 can be added to the Zinin reaction mechanism 2 to better describe what is occuring, taking into account the recycle of byproducts from reaction 3 continued as reaction 2.
Here is the summary diagram again for the three main Zinin reaction paths. My multiplied values in equations below this summary diagram will correspond to these ratios as the same reactions stoichiometry held true but expanded by a multiplier to avoid fractional value results.



The summary equation resulting for reaction 3 with reaction 2

1 ArNO2 + 3 NaHS + H2O ----> ArNH2 + 3 S + 3 NaOH


3 NaHS + 3 S + 3 NaOH ----> 3 Na2S2 + 3 H2O

3 ArNO2 + 3 Na2S2 + 3 H2O ----> 3 ArNH2 + 3 Na2S2O3


4 ArNO2 + 6 NaHS + H2O -----> 4 ArNH2 + 3 Na2S2O3

If the reduction using NaHS as described by the Clayton Communications article proceeds according to the scheme above then for each 1 mole of sodium picrate there will be needed 1.5 moles of NaHS

However, this 1.5 moles NaHS requirement takes into account only the parallel reactions of Zinin reaction mechanism 2 and reaction mechanism 3. But it will also occur to unknown extent that a parallel reaction of mechanism 1 can occur. Formation of Na2S from the NaOH byproduct from reaction mechanism 3 will be instantaneous. Na2S thereby formed can either further react with the free sulfur byproduct or in the alternative the Na2S can react with the sodium picrate as reaction mechanism 1. While the earlier parallel reactions involving mechanism 3 and mechanism 2 would not greatly change the pH of the reaction mixture, this reaction mechanism 1 can (and does) account for increasing alkalinity of the reaction mixture and may be buffered by the presence of NaHCO3, moderating that rise in pH which is already identified by Hodgson and Ward to greatly reduce the yield.

4 ArNO2 + 6 Na2S + 7 H2O ---> 4 ArNH2 + 3 Na2S2O3 +
6 NaOH

This reaction mechanism produces the need for neutralization as is described in the fundamental processes of dye chemistry article.

If reaction mechanism 1 is considered to be an equal possibility for how the Na2S derived from reaction mechanism 3 may proceed, then perhaps half of the Na2S will then act directly according to reaction mechanism 1, while the other half of the Na2S will react with free sulfur and form Na2S2 which will proceed according to reaction mechanism 2.. The actual extent to which reaction mechanism 1 occurs is what will govern the elevation of pH and need buffering or neutralization.

A different scaling multiplier should be used to arrive at whole number quantities.

Reaction mechanism 1 as in Clayton Communications article

4 ArNO2 + 12 NaHS + 4 H2O ----> 4 ArNH2 + 12 S + 12 NaOH

Subsequent reaction of decomposition products with remaining half of unreacted NaHS

[ 12 NaHS + 12 S + 12 NaOH ----> 12 Na2S2 + 12 H2O ]
( if 100% of Na2S reacts with the free sulfur )

In the alternative, if 50% of the Na2S further reacts with half the free sulfur
6 NaHS + 12 S + 6 NaOH ------> 6 Na2S2 + 6 H2O + 6 S

And the other 50% of the Na2S instantly forms and acts as a reducer
6 NaHS + 6 NaOH ------> 6 Na2S + 6 H2O


4 ArNO2 + 6 Na2S + 7 H2O ---> 4 ArNH2 + 3 Na2S2O3 + 6 NaOH


6 ArNO2 + 6 Na2S2 + 6 H2O ------> 6 ArNH2 + 6 Na2S2O3


14 ArNO2 + 24 NaHS + 4 H2O -----> 14 ArNH2 + 9 Na2S2O3 + 6 NaOH + 6 S

For each 1 mole of sodium picrate 1.7143 moles of NaHS would be required by theory for the reduction based upon the proposition that 50% of the Na2S forming in situ from self-reaction of the NaHS decomposition byproducts then proceeded to act directly as a reducing agent, while the other 50% of the Na2S first reacts with half of the free sulfur to form Na2S2 and then proceeds to act as reducing agent. Ideally and most desirably it would be wished that 100% of the instantly formed Na2S would await reacting with the free sulfur and form Na2S2 which would then reduce the sodium picrate without having any affect on the pH since the only byproduct of that mechanism is sodium thiosulfate. But we know that some free sulfur is going to be present which indicates that part of the Na2S is reacting differently and is otherwise occupied and does not form Na2S2. As a result the reaction mixture moving in the direction of reduction is becoming more alkaline which tends to quench the reduction as that alkalinity increases. So the ideal reaction where all of the Na2S would be converted to Na2S2 does not occur and this increases the amount of NaHS required for the reduction of 1 mole of sodium picrate from the 1.5 mole of that ideal reaction to 1.7143 moles based upon a 50% conversion for the Na2S to Na2S2.

Next there may be examined the worst case scenario where 0% of the instantly formed Na2S reacts with free sulfur and instead proceeds to reduce the sodium picrate. Calculations based upon the worst case scenario should then provide the upper limit for the amount of NaHS required by theory per 1 mole of sodium picrate, so that the range of NaHS possible to be a required minimum is then known. For the worst case
scenario there are 2 moles NaHS required by theory for each 1 mole sodium picrate. So for range from ideal reaction 1.5 moles to worst case scenario 2.0 moles or 1/3 more NaHS on a molar basis would be required by theory, is not much difference.

8 ArNO2 + 12 Na2S + 14 H2O -----> 8 ArNH2 + 6 Na2S2O3 + 12 NaOH

Materials possibly useful as a buffer is an ammonium chloride / ammonia solution, or perhaps an ammonium chloride / ammonium acetate mixture.

My impression from observing the reaction done using the more crudely improvised polysulfide containing dry method reagent is that the reduction runs closer to ideal than to the worst case scenario based upon the amount of free sulfur not being great but seeming to be a nuisance contaminant, not a large amount being produced as a reaction byproduct, more like a significant but minor amount trace contaminant.

Reviewing the Hodgson and Ward article again it appears that for reagent (A) that it is indeed 4.6 moles of reducing agent used for reducing agent (A) and that it is an equimolar mixture of Na2S and NaHS and NaHCO3. Reducing agent (B) is Clayton’s reagent. A caveat of sorts is described in the Hodgson and Ward article where they describe testing the reaction mixture to determine when free sulfide was accumulating and showed the end of the reduction. Obviously a positive test would have caused or should have caused interruption of the addition of that reducing agent (A) when only half of the quantity had been added, since half that amount described being made up as reagent (A) is what value is charted as 2.3 and this is where the Hodgson and Ward article descrption becomes enigmatic. But regardless of Hodgson and Wards number discrepancy I am pretty sure my numbers and analysis are correct. If the reducing agent (A) is an equimolar mixture of Na2S and NaHS as described, then that mixture could be achieved by interrupting the H2S addition when the weight of the mixture showed that half of the initially formed Na2S has been converted to NaHS, and then the same molar amount of NaHCO3 would be added as 1/3 the molar amount of the NaOH used to produce the sulfides mixture.

However, it may be that an alternate buffer scheme using ammonium salts could perform better than NaHCO3.

It seems that Hodgson and Ward were in the general area of examining the effect of ammonium salts but their approach was not correct, as if they did not fully understand the Zinin reaction, which is understandable since no one does even after more than 150 years of chemists experimenting with it. The kinetic studies look at the byproducts and try to perform a post mortem on the reductions to guess what mechanisms predominated and led to the products found.

While doing this review there was attention renewed on a patent US2705187 which is classic process chemistry for useful reducing sulfur compounds being produced efficiently by a non-H2S direct reaction of NaOH and S. The temperatures are within operating limits of fluoropolymers which would seem workable as an alternative to expensive alloys. An FEP or PVDF or PFA lab bottle or cartridge filter housing in an oil bath may be workable as a reactor. Im not sure that even the etching would be too excessive for some heavier ceramics or pyrex over the short term since the exposure time is brief. Expendable glassware might be okay in a limited use for this. The etching could prepare it for later adherent coating with a hot melted film of FEP or PFA.

One of the aspects of the reduction schemes using iron and iron salts that would be a beneficial feature is the more pH friendly byproduct Ferric Oxide which doesn't leave a bumper crop of free radical hydroxide as a byproduct driving the reaction mixture alkaline. There are patent schemes where ferrous salts are used along with alkali sulfides in a kind of hybrid reaction where it is likely pH control is more easily done. So it is possible to combine reducing agents and schemes in the same pot and there may be advantage to those hybrid reducing schemes that are part Zinin and part other. The cleanest reduction out of the portfolio of various methods possible is probably the reduction using zinc powder, and that method would probably work similarly to the reduction of nitroguanidine to aminoquanidine. A few possibilities not mentioned are magnesium in methanol as magnesium methoxide, Clemmensen Zinc, or aluminum amalgam. Alkaline sugar and ascorbic acid have been mentioned already. So there is plenty of room left for experiments for this reduction.

[Edited on 8-4-2014 by Rosco Bodine]

Quote: Originally posted by Rosco Bodine

Looking at this now I am beginning to favor my original impression that the Hodgson and Ward description of reagent (A) is actually describing what I first thought it was, which is simply Na2S + NaHCO3 + H2O which is what they are regarding as "equivalent to" when they say "containing" sodium hydrosulfide 5.6g., 0.1 mole, NaHS. Hodgson and Ward reference Hodgson and Birtwell as precedent for reagent (A) and it is clear that Hodgson and Birtwell are describing just Na2S + NaHCO3 + H2O.

Yes, you are correct in my opinion. I think that article was deliberately vague and ambiguous, especially given how well the author expressed himself when he felt like it. Given everything I understand now, methanol, sodium bicarbonate and sodium sulfide is a great combination when used in the correct proportions. I spent half the day experimenting; made more sodium sulfide and reduced 5g of picric acid. I am going to wait until tomorrow until I am sure the product is really dry to report the yield, but it looks fantastic and the filtrate was a nice semitransparent sodium picramate color, not the deep red/purple color that is usually seen when there is a lot of bi-product dye(s).

Regarding hydrogen sulfide generation, I figured out why my generator quit early the last time. The steel delivery tube (brake line) was completely plugged. I used a propane torch and heated the tubing, burning away a lot of the blockage material (seemed to be mostly sulfur). It plugged again though after about 20 minutes of H₂S generation. I cut the power after the rubber stopper blew out (tighter fitting stopper this time) and waited half an hour or so for the H₂S to dissipate before retrieving the sulfide solution. I guess I need a bigger gas delivery tube, and a better way of cleaning it also.

Eight grams of sodium hydroxide was used with water to make about a 75 mL solution. About 150% more sulfur and Vaseline than theoretically needed to convert all the sodium hydroxide to sodium hydrosulfide was used in the generator , though of course it was not all used. The graduated cylinder was weighed before absorption and after. Even though the generator only ran for 20 minutes or so, the weight of the solution had increased by about 3g. This means that nearly all the sodium hydroxide had been converted to sodium sulfide. For simplicity the solution was treated as being composed only of sodium sulfide. About 2 moles of reducer was used for every mole of picric acid.


Monoreduction of Picric Acid - Experimental

Reactants:
5g TNP
0.9g NaOH
3.7g NaHCO₃
35mL aqueous sodium sulfide solution (described above)

Reaction Solvent:
50mL MeOH
water from sulfide solution

The 0.9g of NaOH was dissolved in 50mL of methanol in a beaker under magnetic stirring. Once dissolved, the 5g of picric acid was sprinkled in under strong stirring. The 3.7g of NaHCO₃ was dissolved in the 35mL aqueous sodium sulfide solution before being pipetted into the reaction flask over a 3 or 4 minute period. There was a thick, yellow/orange precipitate, but the stirrer was not decoupled and kept going. About 2 to 4 minutes after the last addition an induction point was reached and the temperature rose from 40^oC to 60^oC over the course of 2 minutes and the reaction mixture turned dark red. The reaction was allowed to cool back down to 40^oC or so over the course of 15 minutes under strong magnetic stirring. After the temperature had dropped below 40^oC an ice bath was added and the temperature brought down to 15^oC before 75mL of ice water was added to aid precipitation of the sodium picramate. After 10-15 minutes, and with the reaction mixture at about 10^oC, the sodium picramate was separated from the mixture by gravity filtration. Yield looks very good.

I was unsure of how well this would work because the amount of water was too large really (about double what it should have been), also there was some unconverted sodium hydroxide present. Looking at sodium bicarbonate and carbonate solubilities in methanol water mixtures, this extra water would result in at least ten times the sodium carbonate being in solution if common ion effects on solubility are ignored. Some sodium carbonate in solution is still a lot better than an equivalent (or more) amount of sodium hydroxide in solution though. If the gererator had not plugged, the 8g of sodium hydroxide would have been converted to sodioum hydrosulfide which later could have been converted to sodium sulfide by adding another
4g of sodium hydroxide to the solution.

I will post the yield tomorrow, once it's dry.





[Edited on 11-4-2014 by Hennig Brand]

Thanks for the good information on hydrogen sulfide. I guess I wasn't using nearly as much of an excess as I thought. That webpage you linked to above has some really good information about hydrogen sulfide.

Monreduction of Picric Acid - Yield

I finally got the sodium picramate from yesterdays experiment nice and dry. From 5.03g of picric acid 4.62g of sodium picramate was obtained, or about a 95% yield. A small amount was not recovered from the filter paper and a small amount was lost during the post filtration wash with ice cold water, but all in all a very good yield indeed. No I did not fudge the numbers, they are real.
A picture of the dry yield is attached below.





[Edited on 12-4-2014 by Hennig Brand]

Yes weight for weight DDNP may be half as good as Pb(N3)2 but the same weight of DDNP requires a substantially larger ID cap to comfortably cram enough of the much lower density DDNP in the detonator volume available to get the job done. DDNP is therefore disadvantaged by both its equivocality for substantial DDT critical mass, and disadvantaged by also being a lower density material requiring additional volume to compensate for both differences. Simultaneously the required adjustment of geometry will address the larger critical diameter for DDNP which is also in play. Usually these properties are interdependent. It is the same story generally for any detonator design, the very small dimension detonators require very hot and very fast high density unequivocal initiators like silver or lead azide to get the job done in a small dimension package, but using lower performance materials can easily double the geometry needed. Even really low performance initiators can work if the detonator dimensions are scaled up enough for accomodating the larger critical diameter and multigram initiator charges required for the high order velocity to be attained. For example a lead styphnate initiator or lead picrate initiator type of detonator can be made but they are inconveniently large so they are not really as practical as a more reasonably sized detonator using higher efficiency initiators. DDNP has better performance than those examples of what would be a really poor initiator. So DDNP is fairly identified as a kind of intermediate performance initiator among the various choices of possible initiators. But the literature seems to describe DDNP in more glowing terms and favorable review than would deservedly be a more conservative and practical review.
DDNP is workable, but the experimenter must rely upon their own test results and not rely upon what is described in much of the dubious "trusted literature" which is simply not so trustworthy or experimentally verifiable in its claims for DDNP. Read the literature with healthy skepticism. Then your own experimental results are less surprising when you see what is actually true.

I did an experiment and my experiment doesn't say what the book says. How about that. The book is Wrrrrr....wrrrr...WRONG! But your experiment is right.

[Edited on 12-4-2014 by Rosco Bodine]

Well I will try to not get into one of those my primary is bigger than yours arguments, but anyway.

My last experiment (before I ran out of reinforcing caps) involved the high order detonation of picric acid initiated with 0.4g of DDNP. The picric acid was in a 7.6mm casing and the DDNP was in a 1/4" reinforcing cap placed within the same 7.6mm casing. It looked as though I may have been able to drop the weight even lower than 0.4g, and still achieve high order detonation, but I haven't done the tests yet. Everything is based on some standard, but as far as I am concerned DDNP is a high performance primary. For instance in comparison to TATP, DDNP is a very high powered primary. It doesn't have the ability to detonate almost instantaneously unconfined like lead azide does, which makes it less convenient to use however. I think the literature may be correct, and that DDNP may be superior to lead azide for initiating explosives like picric acid and TNT. Then again it depends on how you define superior. In terms of ease of use and high tolerance for incorrect loading, etc, lead azide is far superior. I would go so far as to say that DDNP requires 10X the education to use effectively as what is needed for lead azide. Lead azide is extremely forgiving of a lack of understanding, with quite large departures from the ideal still producing acceptable results.

DDNP is much less dense than lead azide, which is true, but when you are using 2g picric acid base charges a little extra cap length isn't that big of a deal , especially for the hobbyist. DDNP has very high sensitivity to flame, which is an advantage over lead azide. DDNP is also relatively insensitive to most other forms of stimuli making it a relatively safe primary. Also as you pointed out, it is fairly "green"; it does not contain heavy metals.

I think lead azide is better when used in small caps with more sensitive secondaries like PETN and RDX. I think DDNP can really hold its own when used to initiate picric acid or TNT when used in a reinforced cap configuration.

I am going to have to try silver azide coated DDNP. Apparently a couple to a few percent of silver azide can be formed on the surface of other less unequivocal primaries producing a primary that is much more unequivocal. Maybe that would solve a lot of the inconvenience with DDNP.

I think lead azide is a much more convenient primary for most applications. I also think that very few people have successfully made DDNP, and even fewer people have a clue about how to effectively use DDNP.

hydrated protosulphide of iron