From: glhurst@onr.com (Gerald L. Hurst)
Newsgroups: rec.pyrotechnics
Subject: Re: Drug Store use of Saltpetre??
Date: 7 Jan 1996 02:07:53 GMT

In article <30ee639a.104007353@news.gnn.com>, drx@gnn.com says:

>        What Isle did you find this stuff in, I looked all over the
>store and couldn't find anything that said SALT PETRE.

I asked my local pharmacist if they get many teenybomber inquiries
re salt peter. She told me that youngsters come in all the time
asking for the material, but the store will not sell it to them
for understandable reasons.  Times have changed, of course. When I
was a youngster they would give you a knowing look but take your
money whether you were buying oil of anise for stink bait, clove
oil for toothpicks or KNO3 for whatever. This was in a time 
before everything became someone else's fault.

In those days I used to walk from the center of a major city
into the country carrying a rifle and ammo pouch in plain sight.
This was a time before we knew much about random shootings.

Later I remember a long period when you could buy dynamite and
blasting caps over the counter without a hassle and shoot the
material outside the city limits as long as you didn't make too
much noise on a regular basis. In those days I set up no fewer
than three explosives labs on the outskirts of major cities
and shot dozens of open tests daily without ever getting a 
visit from the sherrif or filing any papers.  This was a time 
before wanton terrorism. 

Today, even a chemist has to jump through hoops to purchase
glassware, not to mention chemicals because Uncle takes a great
interest in what you are planning to do or make.

We have lost a lot of freedoms and we will pay a price for the
guardianship of the government in lost creativity. Many times
new products and processes arise from risky activities by
individual entrepeneurs. Our freedom to take chances probably
accounted in large part for the disproportionate amount of new
technology which originated in 19th and early 20th century USA.

We make our choices about freedom versus risk and take our 
chances on how well we can live with the outcome.  I suppose
it depends on what we define as "a better society."

Jerry (Ico)

From: glhurst@onr.com (Gerald L. Hurst)
Newsgroups: rec.pyrotechnics
Subject: Re: NH4N03 to KNO3
Date: 15 Jul 1996 04:30:08 GMT

In article <31E7D273.34C@mail.wwdc.com>, taustin <taustin@mail.wwdc.com> says:

>I have been looking for fertilizer KNO3 and in the process I found
>ammonium nitrate. I was thinking that if I couldn't find a source for
>fertilizer potassium nitrate in 50 pound bags I could react the NH4NO3
>with KCl in the same manner as is used to make KNO3 from NaNO3.

You'd do better to react the AN with potassium hydroxide or  
potassium carbonate in water solution and  drive off the resultant
ammonia and water by evaporation with gentle warming.

NH4NO3 + KOH  --> KNO3 + H2O + NH3

2NH4NO3 + K2CO3 --> 2KNO3 + H2O + 2NH3 + CO2

I normally eschew giving out "recipes," but converting AN to saltpeter
is probably a step in the right direction for budding pyrotechnicians.

Jerry (Ico)

From: glhurst@onr.com (Gerald L. Hurst)
Newsgroups: rec.pyrotechnics
Subject: Re: saltpeter
Date: 12 Nov 1996 07:12:25 GMT

In article <3287C027.6A88@uiuc.edu>, Michael Brown <mdbrown2@uiuc.edu> says:
>
>Gerald L. Hurst wrote:
> 
>>> [Detailed KNO3 synthesis from KCl and NaNo3 snipped]
>>
>> I was wondering if it is really necessary to dilute your mixture to one
>> liter and then evaporate it back to 250 ml.  It seems to me that the prep
>> would also work if you simply used 250 ml and let the NaCl precipitate
>> out at will during the mixing process then heated and filtered the hot
>> solution prior to cooling and crystallization.  I guess I do not see the
>> necessity of getting all the material into solution merely to filter off
>> impurities.
>
>
>It's not to filter the impurities, it's so that you have enough
>water to dissolve the least soluble salt, NaCl. What happens is
>that as you dissolve the NaNO3 and KCl, NaCl is formed. When the
>volume of solvent (say 250 mL) becomes saturated, the NaNO3 and
>KCl cease to dissolve.

I forgot to ask.  Why do you want the NaCl to dissolve?  You are 
simply going to filter it off later.  let it precipitate and take
common ions with it. When the solution becomes "saturated" it will
be saturated with dissolved KNO3 in equilibrium with solid KNO3 
and dissolved NaCl in equilibrium with solid NaCl (at room  
temperature) until you heat it and dissolve the KNO3, which has a
sharper temperature coefficient of solubility.

Jerry (Ico)

From: glhurst@onr.com (Gerald L. Hurst)
Newsgroups: rec.pyrotechnics
Subject: Re: saltpeter
Date: 13 Nov 1996 05:03:38 GMT

In article <32894378.3388@mediasoft.net>, Rich Weaver
<weaver_r@mediasoft.net> says:

>If you find a better way, I'm all ears since the stuff isn't done after
>the above process but needs to be recrystalized to remove the remaining
>NaNO3, KCl and yes, even NaCl. This time freezing is in order.

Rich, I opined to Mike that your collective problem is probably surface
dirt and large crystal size.  You talk about how sloooowwwly the big 
KCl crystals dissolve.  Have you ever used the Old European large crystal
sugar to sweeten a drink? -- it takes forever for those crystals to 
disappear.

I hope you guys will try grinding up some of your starting materials
before I have to grumble my way into the garage to see if I've got
some NaNO3 and KCL ready for the coffee mill:)

Jerry (Ico)

From: glhurst@onr.com (Gerald L. Hurst)
Newsgroups: rec.pyrotechnics
Subject: Re: saltpeter
Date: 13 Nov 1996 18:18:51 GMT

In article <3289D23F.1B07@uiuc.edu>, Michael Brown <mdbrown2@uiuc.edu> says:

>Gerald L. Hurst wrote:
>
>> >If you find a better way, I'm all ears since the stuff isn't done after
>> >the above process but needs to be recrystalized to remove the remaining
>> >NaNO3, KCl and yes, even NaCl. This time freezing is in order.
>
>Yeah, what he said.
>
> 
>> I hope you guys will try grinding up some of your starting materials
>> before I have to grumble my way into the garage to see if I've got
>> some NaNO3 and KCL ready for the coffee mill:)
>
>I'm not trying to be facetious Jerry, but between Rich and I, we have
>quite a bit of experience at this little procedure and we've tried
>everything. Hell, I've even tried staring at it, but alas, the watched
>pot does boil but the KCl won't dissolve.
>
>I've done it with recrystallized KCl. Nice clean stuff, but no go.
>Jerry, it looks like you'll have to get the Mrs. out of the kitchen
>and give it a try.

OK guys, I'll fetch some salt substitute from the grocery.  I saw a
jar of NaNO3 in the garage.

BTW, did you try dissolving the KCl in 250 ml of hot water and adding
the solid NaNO3 to the hot solution?

Jerry (Ico)

From: glhurst@onr.com (Gerald L. Hurst)
Subject: Re: saltpeter
Date: Nov 12 1996
Newsgroups: rec.pyrotechnics

In article <3287C027.6A88@uiuc.edu>, Michael Brown <mdbrown2@uiuc.edu> says:

>Gerald L. Hurst wrote:
> 
>> I was wondering if it is really necessary to dilute your mixture to one
>> liter and then evaporate it back to 250 ml.  It seems to me that the prep
>> would also work if you simply used 250 ml and let the NaCl precipitate
>> out at will during the mixing process then heated and filtered the hot
>> solution prior to cooling and crystallization.  I guess I do not see the
>> necessity of getting all the material into solution merely to filter off
>> impurities.
>
>
>It's not to filter the impurities, it's so that you have enough
>water to dissolve the least soluble salt, NaCl. What happens is
>that as you dissolve the NaNO3 and KCl, NaCl is formed. When the
>volume of solvent (say 250 mL) becomes saturated, the NaNO3 and
>KCl cease to dissolve.
>
>That's why I calculated the amount of NaCl that would be formed
>and the amount of water (almost 1,000 mL) needed to dissolve it.
>Rich Weaver discovered this by observation...all the NaNO3 and
>KCl just wouldn't dissolve until he dumped in four times the
>amount of water that he calculated for the final amount of KNO3.
>
>After Rich told me how he finally got it to work, I sat down and
>did the calculations...and then wondered why I hadn't thought of
>that earlier. Lack of experience, I guess.

I don't think you're observing the process correctly and are perhaps
mistaking precipitating KNO3 and NaCl and for undissolved starting
materials.  I will admit that I have not run this reaction, but my
mental model of the process (paradigm) tells me the excess water is
not necessary.

Consider the reaction:

NaNO3(s) + KCl(s) <--> NaCl(s) + KNO3(s) where water is a catalyst.

The deltaG for this reaction is about -13 Kcal/mole so it must occur 
according to thermodynamics.  The only thing that might prevent this
(in your lifetime) would be slow kinetics.  However, we are here 
dealing with inorganic salts, all of high solubility, and the kinetics
MUST be fast by all common experience.

Another way of looking at the process without the specific numbers is 
to realize that the addition and later removal of 750 ml of water 
leaves the product mixture in the same thermodynamic state as though 
the water had never been there.  Again this assumes no kinetic 
hindrance, but for inorganic aqueous metathetical reactions involving
high solubility that is a given.

I suspect that when you mixed NaNO3 and KCl it gave the illusion of
relative inertness when, in fact, the reaction was chugging merrily
along substituting white crystals for other white crystals.

A prediction such as I am making is what Beavis likes to refer to as
"fancy ass chemistry," especially given that it is based solely on
thought experiments. Why not put it to the test to see if there's 
anything to this theoretical stuff?  Use 270 ml or whatever amount of 
water you normally wind up with after boiling it down.  Boil the 
slurry mixture for a while, filter it hot and then cool the filtrate
to RT and subsequently refrigerate and filter off the cold (KNO3) 
product.

Jerry (Ico)

Think of the glee Beavis will experience if you wind up with garbage.
Maybe there will still be an opportunity for me to become a welder and
make an honest living:)

From: glhurst@onr.com (Gerald L. Hurst)
Subject: Re: saltpeter
Date: Nov 13 1996
Newsgroups: rec.pyrotechnics

In article <32891E0D.557A@uiuc.edu>, Michael Brown <mdbrown2@uiuc.edu> says:

>> I don't think you're observing the process correctly and are perhaps
>> mistaking precipitating KNO3 and NaCl and for undissolved starting
>> materials.
>
>Not at all. I start with de-icer for KCl, real nasty looking stuff
>about 3/16" cubes (that's 0.0047625 meters for the Europeans) and
>the NaCl comes out of solution as fine as Morton salt..."When it
>rains, it pours." When the NaCl starts to come out of soln, it comes
>hard and fast, so if it was NaCl there would be a noticable build
>up.

I trust you are grinding or finely crushing the KCl before you run the
reaction.  Large crystals will dissolve exceedingly slowly at room
temperature, especially with product NaCl trying to coat them.

>There's no doubt in my mind that it's the solubility of NaCl that
>stops the displacement reaction. Okay, maybe a little doubt.

Why should it? The ACTIVITY of your liter of dissolved NaCl is quite
nearly the same as it would be in 250 ml since the soly does not
increase much with temperature. Until you dilute way past the soly of 
NaCl, it is in equilibrium with the solid phase and must have an 
activity of unity.  Remember, there is no reaction going on in the liquid
phase - you just have a lot of ions in equilibrium with their various
solids.  
 
>> The deltaG for this reaction is about -13 Kcal/mole so it must occur
>> according to thermodynamics.
>
>I'll have to admit that you caught me with my pants down on the thermo
>calculations, but I'll bet that the deltaG that drives the reaction is
>stopped by the energy required to crystallize the NaCl. Where's that
>slide rule....

The crystallization of NaCl is a slightly EXOTHERMIC process. 

>> I suspect that when you mixed NaNO3 and KCl it gave the illusion of
>> relative inertness when, in fact, the reaction was chugging merrily
>> along substituting white crystals for other white crystals.
>
>Nope.

Well, if your observation says no, then I'll bet you've got a surface
dirt problem that is physically inhibiting the dissolution, especially
if you've been working with large crystals or agglomerates with 
correspondingly small total surface area. Have you actually tried 
simmering or agitating the limited-water large-lump brew for a period 
of time?

Jerry (Ico)

From: glhurst@onr.com (Gerald L. Hurst)
Newsgroups: rec.pyrotechnics
Subject: Re: saltpeter
Date: 14 Nov 1996 00:16:06 GMT

In article <56968v$pvb@geraldo.cc.utexas.edu>, glhurst@onr.com (Gerald L.
Hurst) says:

>NaNO3(s) + KCl(s) <--> NaCl(s) + KNO3(s) where water is a catalyst.
>
>The deltaG for this reaction is about -13 Kcal/mole so it must occur 
>according to thermodynamics.  

Silly me.  That figure should be about -660 CAL per mole.  It makes
no difference in the outcome of the reaction but one should not 
anticipate scalding fingers on the beaker:)

Jerry (Ico)

From: glhurst@onr.com (Gerald L. Hurst)
Newsgroups: rec.pyrotechnics
Subject: Saltpeter from NaNO3
Date: 23 Nov 1996 08:32:52 GMT

Mike Brown and I have been having a discussion
about the preparation of saltpeter from sodium
nitrate and potassium chloride:

NaNO3 + KCl <--> NaCl +  KNO3

The solubility of all of the species in the
reaction increase rapidly with temperature except
for the NaCl, which is only slightly more soluble
at high temperatures.  It is this relatively low
and unvarying solubility of  SN that forms a basis 
for the preparative method.  One would righteously 
expect to be able to simply chuck the starting 
materials into just enough water to dissolve the 
product KNO3 at about 100 degC, filter off the NaCl  
then chill the liquid and filter off the KNO3.

Mike and a friend, however, repeatedly observed
the reaction and came to the conclusion that it
did not actually proceed from left to right unless
they added enough water to dissolve all the
product NaCl  (about a 4 or 5-fold excess) and
then boil it back off to precipitate the hot NaCl
prior to the chilling and second product filtering
step.

This procedure bothered me a lot because I could
not see any theoretical reason why the water
addition and boiling off would be necessary or
desirable.  I suggested that perhaps the visualize
of the reaction were simply creating the illusion
of non-reaction or that they merely needed to
grind up their rough starting materials to give a
bigger surface and thereby speed up the
dissolution step.

Mike suggested I try the reaction myself.  I have
not yet had the opportunity to do so, but I would
like to offer a quote from one of my old chemistry
books -- one copyrighted over the years 1905-1923,
that is back in the period when the common
industrial preparation of KNO3 in the US was still
from Chile saltpeter (NaNO3).  The Germans
switched by necessity to the synthesis via the
intermediate Haber process (ammonia) in WW1, but
the allies mostly still get their supplies from
Chile because of maritime dominance and also as a
result of being somewhat retarded in chemistry
until the Great War taught us just how dependent
we were on Germany for many of our other basic
chemical materials.

I do not mean to suggest that the following quote
proves anything.  I merely present it for the fact
of its existence, not for the truth of the matter.
Anyway, here is what James Kendall said in his
textbook, "Smith's College Chemistry":

     The supply of the natural nitrate being
     insufficient, the salt is made by the double
     decomposition from Chile saltpeter NaNO3
     [reaction as above shown with NaCl
     precipitate].  Sodium chloride is not much
     more soluble in hot water than in cold.  The
     three other salts, however, become very
     soluble as the temperature rises.  Hence,
     when  sodium nitrate and potassium chloride
     are heated with VERY LITTLE WATER, they
     dissolve, sodium chloride is precipitated,
     and potassium nitrate remains in solution.
     The mass is filtered quickly through canvas
     to separate the precipitate, and potassium
     nitrate crystallizes from the filtrate as it
     cools.

 Jerry (Ico)

From: glhurst@onr.com (Gerald L. Hurst)
Newsgroups: rec.pyrotechnics
Subject: Re: Reaction help
Date: 18 Jan 1997 06:45:59 GMT

In article <5bp96t$ig9@newsman.murdoch.edu.au>, Kerry Masters
<masters@cleo.murdoch.edu.au> says:

>could someone please complete the following equations-
>
>        KNO3 + C -->
>
>        
>        KCLO3 + C -->
>
>These are two equations that really bother me, as I have no idea what
>happens to the oxidizer.

4KNO3 + 5C --> 2K2CO3 + 3CO2 + 2N2

2KClO3 + 3C -->2KCL + 3CO2

Should be a pretty good models of the idealized reactions.

Jerry (Ico)

From: glhurst@onr.com (Gerald L. Hurst)
Newsgroups: rec.pyrotechnics
Subject: Re: sodium nitrate
Date: 4 Sep 1995 21:41:52 GMT
Organization: Consulting Chemist
Lines: 26

In article <42fm43$ajf@pcnet2.pcnet.net>, cam@pcnet.com (Cambridge 
Computer Corp.) says:

>i recently bought some sodium nitrate. I was looking for KNO3 but couldn't 
>find any. Is it a good oxidizer?  Does SN work as well as  KNO3 in smoke 
>bombs, and rocket fuel?  How about in BP instaed of KNO3? If it does work 
>in black powder could someone give me the ratios?
>Thanks 

Sodium nitrate is much more soluble in water than is the 
potassium salt and therefore more hygroscopic. This property
does not prevent its use in commercial explosives, but 
pyrotechnic formulations are much more prone to be affected
becaue they, unlike high explosive, rely on flame sensitivity.

Chances are good that mixes made from NaNO3, under high humidity,
will absorb enough water from the air to fulfill the prophecy 
inherent in the old admonition to "Keep your powder dry." The 
oxidizing properties of the sodium salt are similar to those of 
its potassium cousin, but the formula weight is about 84% that of 
saltpeter; therefore, the amount used in any given formulation may
be reduced by that factor from the standpoint of oxygen balance
stoichiometry.

Jerry

From: glhurst@onr.com (Gerald L. Hurst)
Newsgroups: rec.pyrotechnics,alt.engr.explosives
Subject: Re: Explosive SAWDUST!! How to make....
Date: 29 Oct 1996 20:01:20 GMT

In article <jbatistaE01xA2.Byt@netcom.com>, jbatista@netcom.com (Jose
Batista) says:

>Hrm... that is somewhat interesting... I might point out though, that the 
>sawdust is not "nitrated". There has been no chemical reaction. What you 
>describe is merely a mixture. I'd be interested to hear from some of the 
>pros as to the history (if any) of this stuff in contrast to that of BP...

From a standpoint of energetics, calcium nitrate is a tad better than
saltpeter because of the somewhat lower formula weight per unit of 
oxidizer.  The same could be said of sodium nitrate.  The hooker is that
the calcium salt is even more hygroscopic than Chile saltpeter and is 
thus very poorly suited for pyrotechnics work.

Calcium nitrate is manufactured by Norsk Hydro in sweden and is thus 
readily available in northern Europe.  It is imported into the US where
it is an important ingredient in certain water-based commercial 
explosives usually in a eutectic mixture with AN.  For these purposes the 
high water affinity is an asset.

Mixtures of any nitrate with cellulosic material do have low explosive
properties but they are relatively anemic because cellulose is a poor
fuel which carries a weight burden of bound oxygen and absorbed water.

It was common practice to dope cheaper dynamites with mixtures of 
cellulosic waste and a metal nitrate but these materials were second 
rate performers and their manufacture was based on low costs where high
performance was not required.

Jerry (Ico)

From: glhurst@onr.com (Gerald L. Hurst)
Newsgroups: rec.pyrotechnics,alt.engr.explosives
Subject: Re: Explosive SAWDUST!! How to make....
Date: 30 Oct 1996 02:27:58 GMT

In article <555nqg$cdj@geraldo.cc.utexas.edu>, glhurst@onr.com (Gerald
L. Hurst) says:

>Calcium nitrate is manufactured by Norsk Hydro in sweden and is thus 
>readily available in northern Europe.  

I should have said Norway rather than Sweden.  Norsk is a very large
corporation with extensive international holdings including some in
the United States.  They are currently defendants in the civil NY Trade
Center bombing case.

Jerry (Ico)

From: glhurst@onr.com (Gerald L. Hurst)
Newsgroups: rec.pyrotechnics
Subject: Re: Calcium Nitrate Question
Date: 18 Oct 1996 07:08:59 GMT

In article <546ota$hv8@carrera.intergate.bc.ca>, Rory Barr
<barrfarm@intergate.bc.ca> says:

>Hello, I am a beginner in the pyrotechnic field and I have a question. I 
>have recently been experimenting with homemade Black Powder and have 
>achieved a final product that I consider, by the standards of this crowd, 
>to be very good. Anyway I was just wondering if Calcium Nitrate has any 
>use in pyrotechnic mixtures because my source for Potassium Nitrate, the 
>local drugstore, is a bit too expensive for long term supplies. I did 
>however find a nice little hydroponics shop that sells cheap Calcium 
>Nitrate. A helpful response would be greatly apreciated.

Calcium nitrate is too hygroscopic to be of use in fireworks, but it
would be child's play to convert it into potassium nitrate.  Remember
that calcium carbonate and calcium sulfate are both quite insoluble,
which means that the potassium salt of either of these anions can be used to 
precipitate out the respective calcium salt from aqueous solution, leaving
a solution of potassium nitrate.  Wood ashes are largely potassium 
carbonate so there you go.

Jerry (Ico)