From: nreitzel@lonestar.utsa.edu (Norman L. Reitzel   )
Newsgroups: rec.pyrotechnics,alt.engr.explosives
Subject: Re: Fact-checking: C4 stability?
Date: 23 Aug 1994 09:39:07 GMT
Message-ID: <33cg3r$go9@ringer.cs.utsa.edu>

In article <SIMONIS.72.777582988@stpc.wi.LeidenUniv.nl> SIMONIS@stpc.wi.LeidenUniv.nl (Kris) writes:
>>: Yeah, right.  I don't think the pasticizers and the raw ingredients of
>>: the explosive part of C4 (RDX?) are particularly likely to show up at
>>: the same "abandoned" chemical factory.

>>Actually I don't think that RDX is that hard to make.  If anyone is
>>interested I might be able to dig up some papers in JACS that give
>>the correct synthesis.  I think that hexamethylenetetraammine or
>>whatever you call it is easy to get as hexamine for starting barbecues or
>>something like that.  Fuming nitric acid would pose a problem but
>>just maybe an abandoned chemical factory might have that.
>
>Nitric acid is easy to make out of concentrated sulphuric acid and salpeter(
>KNO3), just mix em and start distilling, red fumes will appear and when
>cooled will turn into dark red drops of liquid, this is nearly pure nitric
>acid(about 96-99%, I think).

Hmm, I hope if you ever decide to make explosives you don't do it anywhere
near -me-.  The "dark red drops of liquid" you are describing are known
(strangely enough) as RED fuming nitric acid, which contains anywhere from
10% to 30% free N2O4.  It is definitely -OUT- for manufacturing
explosives, though it works quite well as missile fuel.  If you read
Urbanski on nitration (and you should), he points out that oxidative side
reactions (bad) are catalyzed by the presence of nitrogen oxides, and
gives several methods for removing said oxides from nitric acid.

The "fuming" nitric acid used in explosives manufacture is water white,
or at most a pale straw colour, and contains 98% to 100% nitric acid with
almost no free nitrogen oxides.  Since nitric acid of this concentration
is only marginally stable at room temperature, many plants make it on
site, or alternately use a weaker 90% acid sold commercially as "strong"
nitric acid.  The 90% acid is normally kept refrigerated at close to zero
degrees if it needs to be stored for any length of time.  Many commercial
explosives operations make their own 90% acid on site.

Explosives manufacture using nitration acid mixtures start off with 70%
nitric acid, sold as "concentrated" nitric acid, and make an anhydrous
mixture by mixing with oleum, a mixture of sulfuric acid and free sulfur
trioxide.  Oleum is commercially available with up to 65% free sulfur
trioxide from three or four manufacturers in the United States.

White fuming nitric acid can be made from potassium nitrate and sulfuric
acid, the only catch is that they must be distilled at low pressure.
This is not a hard-vacuum distillation; water aspirators work just fine.
The distillation can be carried out at 50C, and the resulting nitric acid
is close to 100% and perfectly clear.

  ------------------------------------------------------------------------
  Norman L. Reitzel, Jr.       |    "When you live beside the graveyard,
  nreitzel@lonestar.utsa.edu   |     you can't cry for every funeral."
  Blue Water Ventures, dba.    |                     Russian Proverb
  ------------------------------------------------------------------------

Newsgroups: rec.pyrotechnics
From: nreitzel@lonestar.utsa.edu (Norman L. Reitzel   )
Subject: Re: Concentrating HNO3 ???
Message-ID: <1994May11.185116.3942@ringer.cs.utsa.edu>
Date: Wed, 11 May 1994 18:51:16 GMT

In article <hpeens.1.768651781@rkw-lan.cs.up.ac.za> hpeens@rkw-lan.cs.up.ac.za (HENK PEENS) writes:
>Greetings Netters
>
>I am currently studying Chemical Engineering, and just got my grubby little
>hands 5 liters of a 70 % HNO3 solution. Is there an effective way to
>concentrate it in my home lab ???
>
You are close enough to the water-nitric acid azeotrope that there is no
simple way to concentrate the acid.  In order to get anhydrous nitric
acid, you add concentrated sulfuric acid and distill the ternary mixture.
Preferred method is to mix concentrated nitric and sulfuric acids 1:1,
then distill under mild vacuum (30-60 Torr) in an all glass apparatus,
through a column that has 2-3 plates.  This will give you 96-98% nitric.

Please note that anhydrous (or nearly) nitric acid is quite a different
substance from the 70% laboratory solution.  It starts fires, and it's
unstable at room temperature, among other things.

By the way, after you distill off the nitric acid, you can then distill
off the water fraction, ending up with your sulfuric acid reconcentrated.

  ------------------------------------------------------------------------
  Norman L. Reitzel, Jr.       |    "When you live beside the graveyard,
  nreitzel@lonestar.utsa.edu   |       you can't cry for every funeral."
  Blue Water Ventures, dba.    |                         Russian Proverb
  ------------------------------------------------------------------------

From: sfaber@cbnewse.cb.att.com (steven.r.faber)
Newsgroups: rec.pyrotechnics
Subject: Re: Nitric Acid/NH4NO3+H2SO4?
Keywords: Dangerous, cheap.
Message-ID: <1991Aug19.213334.9697@cbnewse.cb.att.com>
Date: 19 Aug 91 21:33:34 GMT

In article <1188@xroads.UUCP>, garbled@xroads.UUCP (Rightnour) writes:
>
>   The reason I want to use Ammonium Nitrate(which I will refer to as AN for
> the sake of my fingers) Is that AN is generally CHEAP.  I can get it at any
> store, in it's purified form, for only $1 a pound.


Ammonium Nitrate decomposes to N2O (laughing gas) when heated just
above the melting point 170-200 deg C.  The gas is unsuitable for
breathing though, since it contains traces of the other oxides of
nitrogen so don't get any ideas.  This may be a problem if you try
use this to make nitric acid, although nitric acid boils at 83 deg.
and should be produced at a lower temperature than this.

The usual apparatus described in old chemistry books for making
nitric acid by distillation of saltpeter with sulfuric acid shows
a retort sticking into a flask with cold water running over the
flask.
I made it this way before successfully, but the acid is a little
brown from some NO2 impurities.
It can be made with less NO2 impurities by vacuum distillation
at 0 deg. C, but this is harder to do in your home lab.


As far as hazards go, a fume hood is needed of course.
The hazards of nitrogen oxide gasses have been mentioned in
many previous postings.  A few wiffs can kill you etc. if over
200 ppm.  You should be able to smell it at 5 ppm, before it is
dangerous, though.  100ppm can be smelled and is not that terribly
irritating but very dangerous.
I don't know if I buy that one about "once lung tissue is
nitrated, it is no good".  If the concentration is high enough
to result in scar tissue damage, then it is true in a sense, but
I have seen no reference to cumulative damage otherwise.

I'm not sure what the apparatus mentioned in the Impovised
Munitions .. Handbook was in a previous posting, but in this case
cracking of the pyrex retort should not be a problem as long as
an asbestos flame spreader pad is used over the bunsen burner
flame.
Electric heating mantles have become standard in most colleges,
with many balking at any open flames as being a fire hazard, but
bunsen burners have been used for years in some organic labs
without incident.

Steve

From: ahahma@polaris.utu.fi (Arno Hahma)
Newsgroups: rec.pyrotechnics
Subject: Re: nitric acid question & soapbox
Message-ID: <1991Aug16.171332.21588@polaris.utu.fi>
Date: 16 Aug 91 17:13:32 GMT

In article <1173@xroads.UUCP> garbled@xroads.UUCP (Rightnour) writes:

>   I know that HNO3 can be made by the reaction of H2SO4+KNO3 in a heated
>retort..  my question is, can the reaction be done, with the replacement of
>KNO3 with NH4NO3 ? it seems logical, but on second thought, might be
>impossible, or at least VERY dangerous..

Yes, it is possible and works even better than those methods based on
saltpeter or sodium nitrate. Just don't heat more than up to 120..125 oC or
the AN will start decomposing to N2O and H2O leading to diluted
nitric acid. For the same reason, the molar ratio of AN to H2SO4 has
to be about one to one, not less than 1 to 0,8. Ammonium hydrogen
sulphate does not react with AN at sufficient rate at 120 oC. If you
can wait for a long time, it is possible to use 1:0,5 ratio. Moreover,
the mixture will not any more remain liquid but solidifies.

The procedure can be made continuous, if the AN is fed as a hot,
saturated solution. This also reduces the need of sulphuric acid,
although some additional acid has to be used to bind the extra water.
The reactions obviously run better, if some water is present. Do not
use water, if you are going to mix all the ingredients at once, i.e
use a batch process. The mixture will boil out of whatever vessel and
you are guaranteed to get it on your hands!

Dangerous? Yes, if you inhale the nitric acid fumes and spill the
mixture all around. The danger of an explosion is virtually nil, all
of the AN will be converted to nitric acid or N2O, before it reaches
the ignition temperature.

>\  /  C r o s s r o a d s  C o m m u n i c a t i o n s
> /\   (602) 941-2005 300-2400,9600 PEP Baud 24 hrs/day
>/  \  hplabs!hp-sdd!crash!xroads!garbled, garbled@xroads.CTS.COM

ArNO
    2

From: ornitz@kodak.kodak.com (Barry Ornitz)
Newsgroups: rec.pyrotechnics
Subject: Re: nitric acid question & soapbox
Message-ID: <1991Aug12.221414.13867@kodak.kodak.com>
Date: 12 Aug 91 22:14:14 GMT

In article <ckeH71w164w@bluemoon.uucp> levi@bluemoon.uucp
(Brandon Moore) writes:
>  To all my knowledge, I know not of preparing nitric acid with sulfuric
>acid and ammonium nitrate.
.....Then why did you post?

>  A better question is why would you want to
>waste ammonium nitrate, a decent, "safe", high explosive, on making
>nitrice acid when potassium nitrate and sodium nitrate are so readily
>availible to almost anyone in one form or another (pure, fertilizer, etc.)
>the reactio for sodium nitrate and sulfuric acid is
>H2SO4 + NaNO3 -> HNO3 + NaHSO4.  NaHSO4 is sodium bisulfate, aka nitre
>cake which can be added to more H2SO4 and heated to produce HCl acid gas.
                                                             ^^^
>I assume the reaction is similar with potasium nitrate, but not with
>ammonium nitrate, because the polyatomic ion NH4 is probably quite
>different than the K or Na atom in the conditions for the replacement
>reaction to occur, if it occurs in a similar why at all.  I don't profess
>to know for sure, but I'd stick with KNO3 and NaNO3 even if it did.

Now where did the halogen come from?  Does Brandon's post really sound like
he understands the chemistry?

PLEASE PEOPLE....Understand what you are writing about before you post!

Making nitric acid can be very dangerous without the right equipment, even if
you know what you are doing.  One good lung full of NO2 and you are likely
to not last very long on this earth.  And the pain even if you do not die will
make you wish you were dead.

Look in the military Improvised Munitions Handbook (or whatever the title is
today) if you want to see a method for making nitric acid that is virtually
guaranteed to end in a tragic accident.  If you follow this procedure, it is
almost impossible to not break the glassware due to thermal stresses.  If you
break the glassware, count on having the contents of the glassware spattered
on you.  It would be hard to find a more corrosive mixture to your skin.  If
you use a fire to provide the heat, count on even more damage when the glass-
ware breaks.

The key to this entire discussion is the issue of safety.  If you know what
you are doing, understand the hazards and how to minimize them, and take the
proper precautions you are far more likely to do things safely.  This means
you need information, accurate information - and this newsgroup is not always
to be trusted for this.  There are lots of good reference books available;
read them first.  If your library does not have them, spend your own money.  It
will be worth it.
				Barry L. Ornitz   ornitz@kodak.com

Newsgroups: alt.engr.explosives,sci.chem
From: "Norman Reitzel" <NormReitzel@swbell.keinspam.net>
Subject: Re: Anhydrous Nitric Acid Production Attempts
Date: Wed, 01 Nov 2000 19:47:50 -0600 (CST)

Sam

Not such a bad attempt, actually, considering the nature of the stuff you are
trying to produce.  You have already noticed that anhydrous nitric acid is
vile and nasty stuff, attacking rubber, cork, hands, ethanol, oil, you name
it, with great abandon.

Try the procedure in Inorganic Syntheses - works first time, every time.

In essence, dehydrate commercial nitric acid with equal volume of sulfuric
acid (note, equal volume is about 1.5x weight).  Distill under vacuum.

Now, some comments.  First, you don't have nearly enough vacuum.  You should
have a water blast going through your aspirator, and the outlet should lead
directly to the drain.  With a decent water pressure, you should be able to
get a vacuum of 40mm or lower.

Second, your condenser and your receiver need to be chilled with ice-water.
Cool water is not adequate, as anhydrous nitric acid has a rather flat vapor
pressure curve, and at 20C (cool water), it's busy eating away at your
aspirator, your tubing, your corks, etc.  I use an ice bath with circulating
water for both receiver and condenser.

Thirdly, you should be using an all-glass apparatus.  Anything organic in the
vapor path will be attacked and liberate nitric oxide, which will then
combine with oxygen in air and turn your acid yellow or orange.

Finally, you might be interested in some measurements of concentration.  When
distilling 62% (technical) nitric acid with 93% (technical) sulfuric acid,
the resulting acid contains around 4% water by weight.  For a lot of uses,
this is sufficiently strong for many purposes - it burns with fuels quite
well.  By distilling 70% (reagent) acid with 98% (reagent) sulfuric acid, one
obtains a product that has around 0.6% water, by weight.  To produce any
better acid, the procedure in Inorg Synt calls for redistilling the crude
anhydrous acid over a second, equal volume of sulfuric acid.  Repeating the
procedure using technical acids gives a nice anhydrous acid containing about
0.3% water by weight.  Using reagent acids produces something with less than
0.1% water.

Finally, I -must- repeat a warning.  Anhydrous nitric acid does not in the
least resemble the 70% concentrated acid.  It's an entirely different animal,
and a relatively vicious one, at that.  It attacks darn near anything, and
worse, it's a great solvent for lots and lots of things.   An anhydrous
nitric acid spill across your hand is a very bad thing.

-- Norm
    Blue Water Ventures (a HazMat Remediation Company)

On Tue, 31 Oct 2000 02:01:26 GMT, mail@powerlabs.org wrote:

>  This is my second attempt at distilling nitric acid to produce white
> fuming (anhydrous) concentrated HNO3. I am aiming for 98% +
> concentration for a test on my liquid fuelled rocket (previously N2O /
> Propane / Butane (60/40 welding mix)) and for some experimental (for
> research purposes only, sub 1 gram) nitrations.
>  I am fortunate enough to have built, over the years, a respectably
> well-equipped laboratory where such procedures can be carried out with
> proper equipment and in good control. I have also been experimenting
> with such things for a long time now and that, together with my
> education in chemistry, makes me confident in what I am doing.
>  So, without further ado, I will get straight to the point: I have
> spent 3 days searching part messages in these and other forums, and
> reading over all my books and internet resources on the topic, and I
> have not yet been able to come up with ANY detailed procedures
> describing how to carry out the distillation of nitric acid to result
> in the product I am looking for. Being left with no other option, I
> decided to do it using the best of my knowledge and the little I found
> on the Internet.
>  Two attempts were made so far. Both utilizing a Pyrex 250mL
> distillation flask with a =9610 =96 250C mercury thermometer attached to
> the flask by the top stopper and glass balls for smooth boiling, and a
> 40cm long Grahan (coil) pyrex condensor. The temperature of the mixture
> was kept constant by boiling it in an oil bath that was heated by a 0 =96
> 350C, 360W thermostat controlled hot plate.
>  My first attempt was to form nitric acid by mixing 29 grams of
> dehydrated, powdered sodium nitrate with 16mL of 98% analytical grade
> Sulphuric acid (read the ratios somewhere). Not all of the NaNO3 would
> dissolve, so I added an extra 10mL H2SO4. Than I proceeded to heat it
> to 89C (the boiling point of the anhydrous acid, as per MERK). Nothing
> happened. Than 105C (the boiling point of the negative azeotrope it
> forms with water at 68% concentration, as per Fisher MSDS). Still
> nothing, but by than the nitrogen tetroxide fumes which had been
> accumulating inside the flask and making it yellowish became denser.
> Increasing the temperature to 150C caused it to boil and decompose into
> dense N2O4 / NO2, so I reduced the temperature. Finally, after 3 hours,
> a millilitre or so had accumulated inside the beaker. It deflagrated
> with ethanol, indicating it was, in fact, anhydrous acid (the azeotrope
> form does not support combustion).
>  I decided that this acid cannot be distilled under atmospheric
> pressure=85 Hence, attempt number two: I mixed 20mL of HNO3 68%
> analytical grade with 20mL H2SO4 98% analytical grade, and proceeded to
> boil them in the same equipment as above, except that the output from
> the distillatory was into a distillation Erlenmeyer, and a low vacuum
> was produced inside it continuously by using a venturi-effect water
> aspirator (GREAT stuff, since it sucks all the noxious nitrogen oxides
> and acid vapours, dilutes them in large amounts of water, runs the
> water through the condenser to keep it cool, and than sends it straight
> down the drain:). Well, this time I got much better results: Over 10mL
> of the concentrated acid. Nothing like using the right equipment for
> the job:) No spills, no smoke, no problems=85
>  Still, what came out was very contaminated with nitric oxide vapours
> (looks yellow), and I had to heat the acids in the distillation flask
> it to 200C to get ANY boiling at all; and when it did boil, it did so
> very slowly, taking over 2 hours for any sizeable amount of acid to
> accumulate. Also, the cork that holds the thermometer was dissolving
> and starting to run down the sides of the flask (uh-oh!).
>
>  So=85 Now onto the questions=85 First of all I would like to know if
> anyone has ever been successful in concentrating HNO3, and if so, than
> how did you do it? Are my ratios right? I know excess H2SO4 has to be
> used, but I could not find any mentions as to how much excess=85 I
> suppose it is better to use the metal nitrate/sulphuric acid method as
> opposed to attempting to dehydrate the HNO3 directly=85 Also: Can the
> concentrated acid be stored in a glass stoppered brown flask or will
> its decomposition produce enough pressure to pup the stopper out? (I
> had a bad experience a while ago with the stopper in a flask of 30%
> H2O2 blowing out a while ago). Any other comments or information are
> more than welcome.
>  Well, any help would be greatly appreciated. Please contact me off
> forum on this (at mail@powerlabs.org).
>
>  Eagerly awaiting for your reply,
>
> Sam Barros
> ICQ#15156975
> Visit my Webpage: http://www.powerlabs.org/
> Science in Action!

Newsgroups: alt.engr.explosives,sci.chem
From: "Norman Reitzel" <NormReitzel@swbell.keinspam.net>
Subject: Re: Anhydrous Nitric Acid Production Attempts
Date: Thu, 02 Nov 2000 08:37:37 -0600 (CST)

On Thu, 02 Nov 2000 03:37:45 GMT, donald haarmann wrote:

> "Norman Reitzel" <NormReitzel@swbell.keinspam.net
>
> Try the procedure in Inorganic Syntheses - works first time, every time.
>
> In essence, dehydrate commercial nitric acid with equal volume of sulfuric
> acid (note, equal volume is about 1.5x weight).  Distill under vacuum.
>
>
> -------
> Sorry plain sulphuric acid is incorrect - Inorganic Syntheses (4:52-53)
> says to use a type of sulphuric acid whose commercial name is that of
> a Bohemian city at the South foot of the Harz Mountains, NNW of Erfurt.
> (First mentioned 927 A.D.; it became a city in the 12th century, part of
> Prussia 1803.) Or by the Latin for "oil".
>
> Also - Inorganic syntheses specifies "atmospheric pressure in a glass
> apparatus."

No problem, Donald, I just committed a cardinal sin.  I quoted a reference
without checking the reference.  The only real inconvenience is that now I'll
have to make a library trip to check the quoted ref in my lab notebooks.
I've only done this a few score of times.

Get back to you

-- Norm