James Dudley
2004-01-19 20:06:00 UTC
I was browsing the archives and came across this very informative post
by Tom Peregrin. It should be of great use to any amateur pyro who is
in search of potassium nitrate, but is restricted by shipping, hazmat
prices or unavailability.
Ejoy,
James D.
------------
One of the most commonly asked questions I see on the internet is ³how
can
I get potassium nitrate²? The obvious answer is ³buy it from a
pyrotechnic chemical supply company². However, a lot of people don¹t
seem
to want to go that route... perhaps they are afraid of paper trails,
perhaps they are underage, perhaps perhaps perhaps... Whatever the
reason, they want to get it in a different way. Some people are lucky
-
they can find it at an agricultural supply house. K-power is a brand
of
potassium nitrate that often sells for between $15 to $20 per 50 pound
bag. One has to mill it, but it¹s cheap! However, not all stores
that
sell fertilizer stock potassium nitrate, and many won¹t order it.
Thus,
these people are often compelled to buy it from pharmacies at
exorbitant
prices, or try to extract it from stump remover or even from fecal
matter. It was obvious that a better source was needed.
The answer can be found in two other common agricultural chemicals -
potash and ammonium nitrate. I called three garden places and all of
them had these two chemicals. The prices are cheap - at the time of
this
writing I can get potash for $5.10/50 pound bag, and ammonium nitrate
for
$8.99/50 pound bag. One should be able to make around 55 to 60 pounds
of
potassium nitrate for around $14, plus labor, plus heating costs.
The process relies on a bit of elegant chemistry. Potash is
potassium
carbonate. When potassium carbonate is dissolved in water, it largely
disassociates to potassium ions and carbonate ions. The same is true
of
ammonium nitrate. Thus, a solution of potash and ammonium nitrate
will
consist of four different types of ions - potassium, ammonium, nitrate
and
carbonate. One can get the same mixture by dissolving potassium
nitrate
and ammonium carbonate. Thus, the solution can be viewed as some
sort of
mixture of all four chemicals - potassium carbonate (K2CO3) ,
potassium
nitrate (KNO3) , ammonium carbonate ((NH4)2CO3) and ammonium nitrate
(NH4NO3). It can be said that all four compounds are in equilibrium,
thusly:
K2CO3 + 2 * NH4NO3 <-- --> 2 * KNO3 + (NH4)2CO3
Long ago a chemist name Le Chatelier observed that when an equilibrium
is
disturbed, the system adjusts to re-establish the equilibrium. Thus,
if
one removes one of the components in the equation, a reaction occurs
and
more is created. For example, if one had an equilibrium solution of
the
four compounds shown above, and removed the ammonium carbonate, then
more
ammonium carbonate would be generated by the reaction of potassium
carbonate and ammonium nitrate. This would generate more ammonium
carbonate, but it would also generate more potassium nitrate. If one
could remove all of the ammonium carbonate, then the final solution
would
contain nothing but ptassium nitrate!
Fortunately, it is actually easy to do just that! Ammonium carbonate
decomposes to ammonia, carbon dioxide, and water at temperatures above
60
degrees Centigrade (140 F). Thus, dissolving potash and ammonium
nitrate
in water, and then bringing it up to the boiling point will cause
ammonia
and carbon dioxide to boil off, leaving behind water and potassium
nitrate.
The amount of potash and ammonium nitrate should be ³stoichiometric²,
i.e., there should be two molecules of ammonium nitrate for every
molecule
of potash. That can be accomplished by using the molecular weights
to
the calculate the weights of chemicals that will contain the
appropriate
ratios of molecules. However, there is one tricky part - both
ammonium
nitrate and potash absorb variable amounts of water. Thus, the
starting
materials must be assayed for their water content. This can be done
by
weighing out 100 grams of each compound and placing it in an oven at
about
150 degrees C (300 F). The samples are weighed every hour until the
weight stops changing. This allows one to calculate the percentage of
water, and to adjust the amounts of the starting materials as needed.
This can be done as such - let us say that it was found that the
potash
contained 16% water (a typical number). To determine the adjustment
factor, one takes 100/(100-water percentage) = 100/84 = 1.19 Thus,
to
obtain 250 grams of anhydrous potash, one would take 250 * 1.19 = 298
grams of ³wet² potash (one won¹t see or feel the water).
The appropriate ratios for the weights of anhydrous ammonium nitrate
and
potash are 160 grams and 138 grams respectively (these must be
adjusted
for water content). The room temperature solubilities of ammonium
nitrate and potash are about 130 grams/100 mililiters and 115
grams/100
mililiters respectively.
This idea was tested in two experiments. In the first laboratory
scale
experiment 169 grams of ammonium nitrate (which had been found to
contain
5% water) were dissolved in about 150 ml of water, and 164 grams of
potash
were dissolved in 150 ml of water. (Remember, the solubilities given
above are for anhydrous compounds, and so the water is not factored
into
the equation to determine the minimal amount of water needed.) When
ammonium nitrate dissolves in water it takes up heat, so the solution
gets
very cold (this is the basis for some instant ice-packs). The
solutions
were both warmed to about 50 degrees C, and then mixed in a large
beaker. The odor of ammonia was immediately evident. The beaker was
placed in a fume hood. The solution was then heated and stirred by
hand.
At about 60 degrees C a few small bubbles appeared, and at 80 degrees
C
the bubbles were forming rapidly. The solution appeared to be
boiling,
even though the temperature was far below the boiling point of water.
Heating was continued, and over the next half hour, the temperature
stayed
around 80 C. After that time, the rate of gas bubble formation
slowed
down and the temperature of the solution rose to 100 C, at which point
a
more normal looking boiling commenced. The volume of the solution
decreased steadily over 30 minutes, during which time the temperature
of
the solution raised to 113 C. At this point, the volume was around
100
ml, and a large rime of crystals was beginning to form around the
surface
of the boiling solution. The beaker was removed from the heat, and
allowed to stand for a minute, after which 100 ml of denatured alcohol
was
added to cause rapid crystalization of the potassium nitrate. The
solution was allowed to cool to room temperature, and the crystals
were
collected via vacuum filtration, and dryed in an oven at 150 C. The
final yeild was 194 grams of very finely divided pure-white
microcrystaline potassium nitrate (96% of the theoretical amount of
202
grams). The crystals were tested for ammonium ions by the
KOH/litmus/evolved-gas method, and for carbonate by attempted
precipitation of barium carbonate via barium chloride. Both tests
showed
less than 1% contamination.
The success of this laboratory experiment led to a ³production scale²
test. In this experiment 7.4 pounds (3.36 Kg) of ammonium nitrate
were
dissolved in 1 gallon of water in a 20 quart stainless steel pot, and
the
pot was placed on a grill over an outdoor woodfire and warmed to 50 C.
An
electric heater could probably be used as well. While that solution
was
warming, 7.2 pounds (3.28 Kg) of potash were placed in a 1 gallon
plastic
milk jug, and 2/3 gallon of hot tap water was added. The jug was
shaken,
and the undissolved potash was allowed to settle. The potash was
poured
into the hot ammonium nitrate solution, and the choking odor of
ammonia
gas caused the experimenter to retreat upwind at a rapid pace. The
rest
of the potash was dissolved in a similar fashion, and was added to the
stainless steel pot at arms length while standing upwind. The fire
was
stoked and the pot was allowed to heat up and begin the evolution of
ammonia and carbon dioxide. Once the rapid evolution of ammonia
commenced, the experimenter amused himself by seeing how far downwind
he
could detect ammonia with his nose. The odor was unbearable at
distance
at 50¹ downwind, and could easily be detected up to 500¹ away.
Obviously
this procedure on this scale is NOT suitable for normal urban
settings.
After about 3 hours the odor of ammonia diminished to a barely
noticable
level, and the temperature of the solution was found to be about 104
C.
The solution was allowed to boil until it appeared to be a little more
than 2 quarts. At this point the temperature as 110 C. The solution
was
allowed to cool and sit overnight. The next morning the mess in the
bottom of the pot was broken up using a heavy oak 2by2, and poured
into a
pillow case. The crystals were wrung dry, and allowed to dry in a
shallow
pan in the sun. After 2 days, they were weighed. The final yeild was
8.1
pounds (about 90% of total possible yield) of a mixture of large and
small
potassium nitrate crystals. The material would obviously have to be
milled before use. The material was tested for ammonium and
carbonate
ions, and was found to contain a slight contamination with carbonate
ions
(perhaps 1%).
The process is not terribly difficult, and could even be interupted
part
way through to make CIA black powder or non-milled Chrysanthemum type
mixes. For example, once the solution has been boiled to eliminate
the
ammonium and carbonate ions and concentrated to a minimal volume, it
should contain only potassium nitrate. Rather than isolate the
potassium
nitrate, one could add charcoal and sulfur, followed by a period of
boiling and then addition of alcohol (see articles in ³Best of AFN
II²).
Safety Notes: The process releases choking and poisonous gases. Do
not
breath the gas. This process HAS to be done outside or in a
laboratory
fumehood. Normal kitchen vent hoods would be inadequate to remove the
ammonia gas. Note that this also makes this a little hard to do on
the
sly in large quantities. The potash solution is caustic, and hot
concentrated potash solutions might cause chemical burns. As always,
rinse all chemical spills with copious quantities of running water,
and if
burning or irritation persists, see a physician. Ammonia readily
attacks
copper, so one cannot use brass or copper kettles. Potash solutions
attack aluminum. Finally, the Merck index remarks that while ammonia
is
generally regarded as non-flammable, that mixtures of air and ammonia
will
explode if ignited under favorable conditions. However, since the
ammonia
explosions only occur when the concetration is between 13% to 79% in
normal air, it is very unlikely that such huge concentrations will be
formed, especially as the ammonia will always be accompanied by equal
amounts of non-flammable carbon dioxide. Be aware.
(C) 1998 Tom Perigrin
Unauthorized commerical or webpage use prohibited
by Tom Peregrin. It should be of great use to any amateur pyro who is
in search of potassium nitrate, but is restricted by shipping, hazmat
prices or unavailability.
Ejoy,
James D.
------------
One of the most commonly asked questions I see on the internet is ³how
can
I get potassium nitrate²? The obvious answer is ³buy it from a
pyrotechnic chemical supply company². However, a lot of people don¹t
seem
to want to go that route... perhaps they are afraid of paper trails,
perhaps they are underage, perhaps perhaps perhaps... Whatever the
reason, they want to get it in a different way. Some people are lucky
-
they can find it at an agricultural supply house. K-power is a brand
of
potassium nitrate that often sells for between $15 to $20 per 50 pound
bag. One has to mill it, but it¹s cheap! However, not all stores
that
sell fertilizer stock potassium nitrate, and many won¹t order it.
Thus,
these people are often compelled to buy it from pharmacies at
exorbitant
prices, or try to extract it from stump remover or even from fecal
matter. It was obvious that a better source was needed.
The answer can be found in two other common agricultural chemicals -
potash and ammonium nitrate. I called three garden places and all of
them had these two chemicals. The prices are cheap - at the time of
this
writing I can get potash for $5.10/50 pound bag, and ammonium nitrate
for
$8.99/50 pound bag. One should be able to make around 55 to 60 pounds
of
potassium nitrate for around $14, plus labor, plus heating costs.
The process relies on a bit of elegant chemistry. Potash is
potassium
carbonate. When potassium carbonate is dissolved in water, it largely
disassociates to potassium ions and carbonate ions. The same is true
of
ammonium nitrate. Thus, a solution of potash and ammonium nitrate
will
consist of four different types of ions - potassium, ammonium, nitrate
and
carbonate. One can get the same mixture by dissolving potassium
nitrate
and ammonium carbonate. Thus, the solution can be viewed as some
sort of
mixture of all four chemicals - potassium carbonate (K2CO3) ,
potassium
nitrate (KNO3) , ammonium carbonate ((NH4)2CO3) and ammonium nitrate
(NH4NO3). It can be said that all four compounds are in equilibrium,
thusly:
K2CO3 + 2 * NH4NO3 <-- --> 2 * KNO3 + (NH4)2CO3
Long ago a chemist name Le Chatelier observed that when an equilibrium
is
disturbed, the system adjusts to re-establish the equilibrium. Thus,
if
one removes one of the components in the equation, a reaction occurs
and
more is created. For example, if one had an equilibrium solution of
the
four compounds shown above, and removed the ammonium carbonate, then
more
ammonium carbonate would be generated by the reaction of potassium
carbonate and ammonium nitrate. This would generate more ammonium
carbonate, but it would also generate more potassium nitrate. If one
could remove all of the ammonium carbonate, then the final solution
would
contain nothing but ptassium nitrate!
Fortunately, it is actually easy to do just that! Ammonium carbonate
decomposes to ammonia, carbon dioxide, and water at temperatures above
60
degrees Centigrade (140 F). Thus, dissolving potash and ammonium
nitrate
in water, and then bringing it up to the boiling point will cause
ammonia
and carbon dioxide to boil off, leaving behind water and potassium
nitrate.
The amount of potash and ammonium nitrate should be ³stoichiometric²,
i.e., there should be two molecules of ammonium nitrate for every
molecule
of potash. That can be accomplished by using the molecular weights
to
the calculate the weights of chemicals that will contain the
appropriate
ratios of molecules. However, there is one tricky part - both
ammonium
nitrate and potash absorb variable amounts of water. Thus, the
starting
materials must be assayed for their water content. This can be done
by
weighing out 100 grams of each compound and placing it in an oven at
about
150 degrees C (300 F). The samples are weighed every hour until the
weight stops changing. This allows one to calculate the percentage of
water, and to adjust the amounts of the starting materials as needed.
This can be done as such - let us say that it was found that the
potash
contained 16% water (a typical number). To determine the adjustment
factor, one takes 100/(100-water percentage) = 100/84 = 1.19 Thus,
to
obtain 250 grams of anhydrous potash, one would take 250 * 1.19 = 298
grams of ³wet² potash (one won¹t see or feel the water).
The appropriate ratios for the weights of anhydrous ammonium nitrate
and
potash are 160 grams and 138 grams respectively (these must be
adjusted
for water content). The room temperature solubilities of ammonium
nitrate and potash are about 130 grams/100 mililiters and 115
grams/100
mililiters respectively.
This idea was tested in two experiments. In the first laboratory
scale
experiment 169 grams of ammonium nitrate (which had been found to
contain
5% water) were dissolved in about 150 ml of water, and 164 grams of
potash
were dissolved in 150 ml of water. (Remember, the solubilities given
above are for anhydrous compounds, and so the water is not factored
into
the equation to determine the minimal amount of water needed.) When
ammonium nitrate dissolves in water it takes up heat, so the solution
gets
very cold (this is the basis for some instant ice-packs). The
solutions
were both warmed to about 50 degrees C, and then mixed in a large
beaker. The odor of ammonia was immediately evident. The beaker was
placed in a fume hood. The solution was then heated and stirred by
hand.
At about 60 degrees C a few small bubbles appeared, and at 80 degrees
C
the bubbles were forming rapidly. The solution appeared to be
boiling,
even though the temperature was far below the boiling point of water.
Heating was continued, and over the next half hour, the temperature
stayed
around 80 C. After that time, the rate of gas bubble formation
slowed
down and the temperature of the solution rose to 100 C, at which point
a
more normal looking boiling commenced. The volume of the solution
decreased steadily over 30 minutes, during which time the temperature
of
the solution raised to 113 C. At this point, the volume was around
100
ml, and a large rime of crystals was beginning to form around the
surface
of the boiling solution. The beaker was removed from the heat, and
allowed to stand for a minute, after which 100 ml of denatured alcohol
was
added to cause rapid crystalization of the potassium nitrate. The
solution was allowed to cool to room temperature, and the crystals
were
collected via vacuum filtration, and dryed in an oven at 150 C. The
final yeild was 194 grams of very finely divided pure-white
microcrystaline potassium nitrate (96% of the theoretical amount of
202
grams). The crystals were tested for ammonium ions by the
KOH/litmus/evolved-gas method, and for carbonate by attempted
precipitation of barium carbonate via barium chloride. Both tests
showed
less than 1% contamination.
The success of this laboratory experiment led to a ³production scale²
test. In this experiment 7.4 pounds (3.36 Kg) of ammonium nitrate
were
dissolved in 1 gallon of water in a 20 quart stainless steel pot, and
the
pot was placed on a grill over an outdoor woodfire and warmed to 50 C.
An
electric heater could probably be used as well. While that solution
was
warming, 7.2 pounds (3.28 Kg) of potash were placed in a 1 gallon
plastic
milk jug, and 2/3 gallon of hot tap water was added. The jug was
shaken,
and the undissolved potash was allowed to settle. The potash was
poured
into the hot ammonium nitrate solution, and the choking odor of
ammonia
gas caused the experimenter to retreat upwind at a rapid pace. The
rest
of the potash was dissolved in a similar fashion, and was added to the
stainless steel pot at arms length while standing upwind. The fire
was
stoked and the pot was allowed to heat up and begin the evolution of
ammonia and carbon dioxide. Once the rapid evolution of ammonia
commenced, the experimenter amused himself by seeing how far downwind
he
could detect ammonia with his nose. The odor was unbearable at
distance
at 50¹ downwind, and could easily be detected up to 500¹ away.
Obviously
this procedure on this scale is NOT suitable for normal urban
settings.
After about 3 hours the odor of ammonia diminished to a barely
noticable
level, and the temperature of the solution was found to be about 104
C.
The solution was allowed to boil until it appeared to be a little more
than 2 quarts. At this point the temperature as 110 C. The solution
was
allowed to cool and sit overnight. The next morning the mess in the
bottom of the pot was broken up using a heavy oak 2by2, and poured
into a
pillow case. The crystals were wrung dry, and allowed to dry in a
shallow
pan in the sun. After 2 days, they were weighed. The final yeild was
8.1
pounds (about 90% of total possible yield) of a mixture of large and
small
potassium nitrate crystals. The material would obviously have to be
milled before use. The material was tested for ammonium and
carbonate
ions, and was found to contain a slight contamination with carbonate
ions
(perhaps 1%).
The process is not terribly difficult, and could even be interupted
part
way through to make CIA black powder or non-milled Chrysanthemum type
mixes. For example, once the solution has been boiled to eliminate
the
ammonium and carbonate ions and concentrated to a minimal volume, it
should contain only potassium nitrate. Rather than isolate the
potassium
nitrate, one could add charcoal and sulfur, followed by a period of
boiling and then addition of alcohol (see articles in ³Best of AFN
II²).
Safety Notes: The process releases choking and poisonous gases. Do
not
breath the gas. This process HAS to be done outside or in a
laboratory
fumehood. Normal kitchen vent hoods would be inadequate to remove the
ammonia gas. Note that this also makes this a little hard to do on
the
sly in large quantities. The potash solution is caustic, and hot
concentrated potash solutions might cause chemical burns. As always,
rinse all chemical spills with copious quantities of running water,
and if
burning or irritation persists, see a physician. Ammonia readily
attacks
copper, so one cannot use brass or copper kettles. Potash solutions
attack aluminum. Finally, the Merck index remarks that while ammonia
is
generally regarded as non-flammable, that mixtures of air and ammonia
will
explode if ignited under favorable conditions. However, since the
ammonia
explosions only occur when the concetration is between 13% to 79% in
normal air, it is very unlikely that such huge concentrations will be
formed, especially as the ammonia will always be accompanied by equal
amounts of non-flammable carbon dioxide. Be aware.
(C) 1998 Tom Perigrin
Unauthorized commerical or webpage use prohibited