The presence of even trace amounts of Chlorate in Perchlorates makes mixtures using the "Perchlorate" as sensitive or nearly as sensitive as mixtures containing ALL Chlorate. You must take ALL Chlorates out of Perchlorates if you are going to treat the product as a Perchlorate.

You should never mix any Chlorate with Ammonium Perchlorate as Ammonium Chlorate may form and you will have a very dangerous situation.
You should never mix any Ammonium compound (eg. Ammonium Nitrate) with any Chlorate as Ammonium Chlorate may form and you will get a very dangerous situation.
The presence of Chlorides in Chlorate has been shown to increase the sensitivity of mixtures of Chlorate and combustibles. If using Chlorate, remove all Chloride.
The finished product that you have made should be made neutral (this is best done before the final crystallization) using NaOH or KOH. Which you use is chemically unimportant but if you are concerned about the colour that your final product will make as it burns you must use the proper alkali.

"Bill Ofca's Technique in Fire Volume 10 Working Safely With Chlorate" lists Asphalt as one of the incompatible "chemicals that must NEVER be mixed with chlorates". It also lists "Sulfur or any compounds of sulfur, sulfides, and sulfates"

Skylighter carries this book (~$17). It is not a bad idea to own a copy if you choose to work with chlorates. It could save your life.
Ultraviolet light is also a problem when metal oxides are present which catalyze the reaction. This is one of the most common causes of spontaneous ignition especially with copper compounds. Asphalt typically contains many impurities.

Herbert G. Tanner
Instability of Sulfur-Potassium Chlorate Mixture: A chemical view
Journal of Chemical Education Vol. 36 No. 2, February 1959

Numerous amateur rocket enthusiasts suffered serious casualties from
premature sulfur-chlorate explosions, according to C. Burns (1). This is most
regrettable. These accidents should challenge elementary, chemistry teachers,
textbook writers, and civic-minded chemists to do their utmost to teach safety to
tyro chemists. Chemistry without safety is too frequently fatal.

The following exposition of the chemistry of sulfur-mixture is presented-with the
hope that the information will be conveyed to secondary school chemistry
students as a lesson in safety. Safety is not a series of "'don'ts," but is a matter
of foresight, with avoidance of hazards. In the discussion of the chemistry of
sulfur-chlorate mixture, an example of how safety can be designed into an
experiment is included for the benefit of the chemists of tomorrow.

Sulfur, by itself, is not a particularly reactive element at ambient temperatures.
Even in the finely divided state it is difficult to ignite. Similarly, chlorate is a
relatively stable compound at ordinary temperature. It decomposes only when
heated (2) almost to its melting point of 368'C. It is much less reactive at room
temperature than is pure oxygen, for example. These facts would not incline one
to predict that a simple mixture of sulfur and chlorate might ignite spontaneously
after being stored in a warehouse many months, or detonate when a fuse is
being stuffed into a rocket charge.

World accident records for the past, century show that sulfur-chlorate stability
is erratic and unpredictable. This fact indicates that the mixture might contain a
concealed chemical trigger. A close look, particularly at sulfur, discloses this
trigger and the manner in which it becomes cocked.

There are two basic grades of sulfur—crude and refined. Sulfur produced at the
mine is called crude sulfur (formerly “roll sulfur"). About half of the world of sulfur
is extracted from the ground in Texas by the Frasch process. Traditionally, this
raw product is called crude sulfur, even though it has the amazing purity of
99.8 to 99.9% sulfur. The chief impurities asphalt which is present in the dissolved
state. This mere trace of asphalt makes raw sulfur difficult to burn on a large scale.
Air oxidation at ordinary, temperature is detectable only after period of months.
Moisture in the air assists this slow “rusting” process.

Refined sulfur is produced by distillation. The purist is obtained conditions that
cause the sulfur to condense directly too the solid state, i.e., the sulfur is
sublimed. The product , called flowers of sulfur, when freshly prepared has an
average purity of 100.0% sulfur. Even when both grades of sulfur are pulverized
to the same degree of fineness, flowers of sulfur oxidize more rapidly because its
surface is cleaner.

Polythionic Acid, the "Trigger"

Air oxidation probably produces sulfurous acid on the surface of sulfur, but this
acid has not been detected (3) because, as H. Debus showed, sulfurous acid
reacts quickly with sulfur to form polythionic acids beginning with trithionic acid,
H2S306. The latter adds sulfur atoms successively to form tetra- and penta-
thionic acids. These acids, particularly pentathionic acid (5), are responsible for
the fungicidal value of dusting sulfurs. Aqueous polythionic acids accumulate
until the concentration reaches a limit (6) where evaporation or a temperature
rise will cause partial decomposition according to the equation:

H2S3O6 = H2SO4 + S02 + (n - 2) S

The reaction is irreversible, but loss of polythionic acids becomes repaired by air
oxidation of sulfur when the temperature subsides. Reaction (1) is significant in
that a sudden increase in temperature can, cause a significant amount of sulfur
dioxide to be formed, not by direct oxidation of sulfur, but by decomposition of a
concentrated solution of polythionic acids. This event, however, must await the
formation of a concentrated solution of these acids.

Reaction (1) is responsible also for the accumulation of sulfuric acid on sulfur.
The sulfuric acid, being hygroscopic, favors additional production of polythionic
acids. By capillary action, sulfuric acid coats the surface of chlorate and
produces chloric acid. Although dilute chloric acid at ambient temperature is a
very weak oxidizing agent, its presence has given rise to the theory that chloric
acid is the primary cause of sulfur chlorate instability.
Sulfur dioxide will react directly with moist chlorate to produce chlorine dioxide.

S02 + 2KC103 = 2CIO2 + K2SO4 (2)

Chlorine dioxide immediately attacks, sulfur, the chief reaction being,

2CIO2 + 4S = 2SO2. + S2Cl2 (3)

Expressing reactions (2.) and (3) as one reaction.

S02 + 2KClO3 + 4S = 2SO2 + S2Cl2 + K2S04 (4)

Reaction (4) is a chain reaction because more sulfur dioxide is produced than is
consumed. Initiation of reaction (4) requires an extraneous source of sulfur
dioxide. This trigger is cocked by the growth of polythionic acids on the sulfur,
and is “pulled" when the heat from friction, impact, sunshine, the like is sufficient
to release a threshold amount of sulfur dioxide from the "reservoir" of polythionic
acids. Heat from reaction (4) quickly ignites a portion of the mixture, and, if the
gases produced cannot escape readily, the mixture will explode.

Popular formulas for preparing sulfur-chlorate mixtures generally specify more
sulfur than can be oxidized by the chlorate to sulfur trioxide. The reasons for this
are not pertinent to this subject, but it is important to note that the excess sulfur
increases the instability of the mixture by favoring formation of a greater amount
of polythionic acids.

This “trigger" theory was capable of explaining some previously baffling industrial
accident conditions, but it needed experimental confirmation. It predicted that a
chlorate mixture prepared with an oxidation resistant crude sulfur would be more
stable than one prepared with flowers of sulfur. Comparative tests were made
using ignition temperature as an index of stability.

Experiment A: 2 g of flowers of sulfur, 4 g of reagent grade chlorate, and 4 g of
cleaned and dried sand were intermingled on unglazed paper with a metal
spatula; 1 ml of water was added dropwise, and the mixing was continued. The
product was pressed lightly into a thin open-ended cardboard tube made from a
paper match-cover after the striker strip had been cut off. The tube was laid on
an asbestos board in an empty ventilated hood and was heated at approximately
1oC temperature rise per minute by a heat lamp suspended above the tube. The
temperature was measured with a thermocouple inserted to the center' of the
mixture. Ignition temperatures on repeat experiments ranged from 82 to 91oC,
and averaged 85oC.

Safety precautions that were designed into the above experiment include mixing
the powders with a metal spatula and the use of unglazed paper to reduce
accumulation of static charges. The powder was dampened primarily to give it
sufficient cohesion to remain in the cardboard tube, but the dampness no doubt
helped to suppress static charges. The striker strip was removed from the
match-cover because it contains a sulfide of phosphorus which might ignite the
mixture prematurely by contact with stray chlorate. Porosity, contributed by sand,
and the short open-ended tube all owed easy escape of gases, thereby reducing
the probability of an explosion. Both ends of the tube were left open to reduce
fire hazard by canceling rocket effects when the mixture ignited. Fire hazard was
further reduced by the asbestos board and by the emptiness of the hood.

Experiment B: Mixtures similar to those in Experiment A were prepared with
crude sulfur that had been pulverized until the particles had a mass median
diameter of 44 microns, thereby approximating the particle size of flowers of
sulfur. These mixtures ignited ,sharply at 124oC. This temperature is above the
melting point sulfur and indicates that ignition may have been initiated by sulfur
dioxide produced by air oxidation of hot sulfur vapor.

Experiments A and B definitely demonstrated the predicated greater stability of
the crude sulfur mix. However, the chloric acid theory could predict that the
greater sulfuric acid content of the flowers of sulfur, instead of the polythionic
acids, would cause the lower ignition temperature. As a matter of fact, the
flowers of sulfur contained 0.02% "acidity" calculated as sulfuric acid. The
acidity of the crude sulfur was detectable but was too small to titrate.

As a test of the "chloric acid" theory, Experiment A was repeated using flowers of
sulfur dampened with sufficient 0.2 N H2SO4 to give it an acidity value of 0.50%.
When this highly acidified sulfur was mixed with chlorate, an abnormal amount of
chloric acid must have been formed. In spite of this excess, the ignition
temperatures averaged only 87oC in close agreement with Experiment A. These
results dismiss the "chloric acid" theory as a primary explanation of instability.

Final confirmation that sulfur dioxide triggers the ignition was obtained by
repeating Experiment A, but. instead of heating the mixture, a long capillary
glass tube containing sulfur dioxide flowing at the rate of 0.3 cc per second was
thrust into the mixture. Ignition occurred almost as quickly as though the flame of
match had been applied.

Sulfur--chlorate typifies many fuel-chlorate mixtures. Metallic sulfides and
polysufides behave qualitatively like sulfur. Phosphorus cannot yield sulfur
dioxide, but at room temperature it evolves reducing vapors of phosphorus oxide
and possibly some elemental phosphorus which react so energetically with
chlorate that an explosion usually occurs before mixing can be accomplished.

Carbon disulfide, rosin, turpentine, thiocyanates, aldehydes, sugars, tannin, and
numerous other materials that form volatile reducing agents on heating,
contribute instability to chlorate mixtures. Powdered metals such as oily
aluminum, zinc, and magnesium are sometimes used with or without other fuels
in chlorate mixture. These metals are corroded by chlorate. Their chlorates are
hygroscopic and decompose at relatively low temperatures. The high heats of
oxidation of these metals could cause local temperatures sufficiently high to
ignite the mixture. No wonder such concoctions are unpredictable in stability.

Much of the unsavory reputation assigned to chlorate should be reassigned to
those who have selected the fuels. Apparently availability and cheapness,
instead of chemistry, have dictated the choice of fuels. Any fuel that is subject to
air oxidation at ordinary temperature, or forms an unstable chlorate, or produces
an easily oxidizable vapor below 100oC, or reacts with chlorate below 150oC
should be excluded. Starch (7) is a fuel that survives these requirements.

Literature Cited

(1) BURNS, C., J. CHEM. EDUC., 33, 308 (1956).
(2) BROWN, F. E., AND WHITE, W. C. 0., Proc. Iowa .Acad Sci., 31,291(1924).
(3) MELLOR, J. W., "Comprehensive Treatise on Inorganic and Physical
Chemistry," Longmans, Green &, Co., Inc., London, 1947, Vol. 10, pg. S8.
(4) DEBS. H., Ann.. 244, 76 (1888),
(5) YOUNG, H.C. AND WILLIAMS ROBERT, Science, 67, 19 (1928)
(6) MELLOR, J.W. "Modern Inorganic Chemistry," Longmans, Green & Co., Inc.,
London, 1916, p. 458; EPHRAIM, FRITZ, "Inorganic, Chemistry," Gurney and
Jackson, London, 1926, p, 463.
(7) TANNER, H.G. U.S Patnet 1,966,652 (1934)

PGI Bulletin No.
41, March, 1984
COMPOSITIONS  By - The WiZ (donald j haarmann)

(The following article was in response to Ed Raszkowski's Question in PGI
Bulletin No. 38.)

Dr. McLain and Donald Lewis; Effects of Phase Change in Solid-Solid
Reactions-1966, [my copy courtesy of E.R.], determined that when a
mixture of potassium chlorate was doped with copper chlorate, and then
ground with purified sulfur [a brave move], the mixture detonated
spontaneously after being let stand undisturbed on a shelf for only thirty
minutes! The experiment was repeated using the same components,
simply placed in a plastic vial and mixed by "tumbling the vial
approximately ten minutes about its major axes." After approximately
twenty-four hours of standing behind a barricade, a similar detonation
occurred. [A less complete discussion of this work is reported in Dr.
Milan’s book; pg. 78.]

They concluded as a result of these experiments that; "Aside from the
spectacular, the finding that a potassium chlorate could be made which
when ;nixed with sulfur is spontaneously detonable at room temperature
there is a byproduct(s) of this experiment."

The doctrine of using brass, bronze or copper screens, tools and jigs for
pyrotechnic mixing and loading for non-sparking properties certainly
needs to be reexamined in light of these findings. Emphasis added]

The warning is repeated by Ellern.

J.C. Shumacher; Perchlorates: Their Properties, Manufacture and Use;
page 206+ff reports that the "***decomposition [of ammonium perchlorate]
was strongly catalyzed by powdered copper metal or cupric oxide, and to
a lesser extent by sodium chlorate, lithium or potassium dichromate."
"emphasis added] even more importantly he states on page 215:

For example a number of fires in the drying and packaging equipment of
the AMMONIUM perchlorate plant of Western Electrochemical Company
(now American Potash dc Chemical Corporation) were found to have
been caused when the perchlorate crystals came into contact with copper
tubing in the vibrating pan-dryer heat exchanger. When stainless steel
was substituted for copper the fires ceased. Ammonium perchlorate was
found to be most sensitive to ignition at a moisture content of 0.02 to 0.5
per cent, [how dry are your stars?], particularly when it is contaminated
with copper and possibly in the presence of other metallic contaminants.
This experience confirms other reports of the accelerating effect of copper
on the thermal decomposition of ammonium perchlorate, and pyridine
perchlorate. [Emphasis added.]

The development of cheap electricity during the later part of the 19th
century, was followed by the first commercial electrolysis of chloride
solutions for the production of sodium chlorate by the French in 1866. As
a result of this process large amounts of chlorates became available at
reasonable cost. It soon followed that because chlorate mixtures had "at
all times fascinated inventors on account of the large amount of oxygen
stored up in *** chlorate(s), which can be given off so readily," these
cheap and readably available chlorates (sodium, potassium, ammonium)
were soon used for the production of explosives. In 1890, electrolysis of
sodium chlorate, lead to the production of potassium and ammonium
Perchlorates for use in the production of explosives. Although few of these
chlorate/perchlorate based explosives ever found favor in this country,
they were widely used i!1 Europe. Their use having in all cases been
supplanted by the introduction of ammonium nitrate based blasting

Because of an increasing number of accidents resulting from the use of
chlorate explosives, the US Bureau of Mines performed "Frictional Impact"
testing of chlorate explosives between September, 1911 and February 1,
1919. [Why they took eight years to complete these tests is not recorded!]
The results were reported in: Bureau of Mines; Tech. Paper 234, 1919.

Investigated was: "A certain potassium chlorate explosive that is used in
the United States and is designated Chlorate Explosive B." [Potassium
chlorate, sugar, gum arable]. The mixture had come under suspicion
because a number of premature explosions had occurred during its use in
bituminous coal mines. Quoted from the report are the following

1. "While inserting copper needle in a charge *** the charge exploded."
2. "The charge exploded while the copper needle was being inserted in
the charge bore hole."
3. "While pushing a 12-inch cartridge into bore hole, with copper needle
inserted about 6 inches in it, *** charge exploded."

They felt that: In view of the frequency of these accidents and the
seeming ease with which they took place, not being explainable by the
sensitiveness of the explosive itself to friction, it was considered advisable
to determine, if possible, the cause of the excessive sensitiveness to

It was thought, as practically all bituminous coals contain pyrite, perhaps
the pyrite was the cause of the accidents. [Now time out for a little
mineralogy -- Pyrite or "Fools Gold" is iron-sulfide (FeS2), it can with time


Mixtures containing Sulphur and a Chlorate, Black oxide of Copper or Sulphide of Copper (and probable all Sulphides and Sulphites) with chlorate should be avoided. Purple fires which generally contain one of the above compounds of Copper have an especially bad name for going off spontaneously. - C.T. Brock, 1872. ----------------------- HAZARDS OF BLUE STAR METALS Lloyd Scott Oglesby First published in the American Fireworks News #59 & 60. August/September 1986 Republished in Best of AFN II (from which it was scanned) Jack Drewes Editor/Publisher of AFN has given permission for me to posted here. I have had several letters inquiring about metal fuel blue stars. The chemical potential for reactions between magnesium (or magnalium) and any copper-plus-two salt or compound is large. All of these reactions are quite exothermic, enough so to make a fella worry. All of these reactions are thermodynamically spontaneous and most are chemically spontaneous at 20°C., especially in the presence of good old water. For instance if you mix powdered copper sulfate pentahydrate and fine magnesium powder, pour the mix into a tube and squirt water into the tube, it will undoubtedly heat up, almost always steam, frequently catch fire with surprising violence and occasionally go bang with a convincing report. If the chemicals are not perfectly dry to begin with or the day is very humid, the reaction may start before you apply the water. TESTING REACTIVITY The reactions involved in the order of their appearance: Mg° + Cu++ -> Mg++ + Cu° Mg° + H2O (g&l) -> MgO + H2 Mg° + SO4 -> MgO + MgS To observe the first reaction safely, all you need is a good heat sink, so go to the sink and fill a glass with water. Dissolve a teaspoon or so of copper sulfate or any other soluble copper salt in the water, measure the temperature of the water. Stir the water vigorously and add slowly a teaspoon of magnesium powder. You will notice the almost instant appearance of a red, brown or black precipitate of copper powder. Take the temperature of the water again. By increasing the amount of chemicals used or decreasing the amount of heat sink, water, you can easily reach the boiling point of water. You can easily arrange for a surprisingly violent steam explosion which makes a nice outdoor demonstration, tossing the contents of the glass into the air a good many feet. If you don't mind poisoning the grass with copper, you might try it. Remember that the reactions are a bit difficult to regulate and the shower is boiling water with chemicals. To arrange for the sudden "bumping" type reaction, it is best to wipe the glass with oil, which eliminates the points on the glass where active sites help bubbles of steam form, and thus dissipate the heat harmlessly without the dramatic "bump'. Put water and magnesium in the glass, pour in suddenly a saturated solution of the copper salt and prepare for the sudden bump" explosion that will cause rather nasty rain. Don't do this in the kitchen, unless you want to paint your wife's ceiling, scrub walls, etc. The form of copper powder that those reactions produce is excellent for pyrotechnic use, if distilled water is used. Any excess magnesium present can be dissolved with acetic acid. Traces of acetic acid will cause the copper to react with air forming "vertigris" so very thorough rinsing is required.


DON'T BE FOOLED We have so far considered only chemical problems involving the ubiquitous water, magnesium and the copper salts. You might falsely suppose that these could be eliminated by simply using cellulose nitrate in acetone or MEK as a binder, but cellulose nitrate is not sufficiently insulating electrically, not a good enough water barrier, and both acetone and MEK solvents commonly contain water and attract water from the atmosphere. Acetone, in fact, must contain water to preserve it, and prevent it from polymerizing, or dimerizing. Also cellulose nitrate slowly decomposes to yield a very hygroscopic coating on the magnesium. AMMONIUM PERCHLORATE PROBLEMS As my old friend and teacher William McGavok, Ph.D (God rest him well) would say, "we aren't out of the woods yet". In fact, we haven't even met all of the dangerous beasts therein. Ammonium perchlorate gives such beautiful blues in low temperature flames that it would be nice to try it in high temperature flames. It is an excellent chlorine and HC1 donor. Since the flame of magnesium can consume chlorine and hydrogen chloride, we will need extra to achieve a good blue. Ammonium perchlorate is hygroscopic enough that the powder mixed with magnesium powder commonly evolves ammonia and makes a puddle of magnesium perchlorate. If you mix a tiny dab of both in the palm of your hand and close your hand on it for a few minutes, the reaction will start from moisture on the skin. In a few more minutes you can smell the ammonia. Quickly wash it off if you notice rapid heating!


-- Posted to rec.pyrotechnics by: donald j haarmann - independently dubious ----------------------------

Using new materials

The following was posted by Tom Perigrin to Rec.Pyro.

There is nothing inherently wrong with looking into the use of exotic
materials for fireworks.   Goodness knows I've done it.   But,
when doing so, there are some things you should consider:

1)  Safety
2)  Justification

1)  Safety -

    When working with new materials you should realize you are stepping
off into the unknown.   And unlike cake baking, that "first step" can be
a REAL DOOZEY in fireworks.    The very purpose of a firework mixture is
to react violently, at a rate which is just short of uncontrolled
mayhem.  It is very easy to make a small change in a formulation that
can push it past that point, and convert a star into a little bomb, or a
salute into an accident.

    An example is the lack of appreciation some people have for the
difference between chlorate (ClO3-) and perchlorate (ClO4-).    When I
was a kid I thought that they were similar, and that perchlorate just
had 4/3's as much oxygen to give off.  Thank goodness I did most of my
big experimentation with perchlorate, thinking that I'd get more oxygen
that way.   In fact, I was accidentally being safer through the use of
perchlorate.   Thats because chlorate is much more reactive than
perchlorate.   I used to make mixtures based on a simple theory;  to
make a firework you need an oxidizer and a fuel.   So I'd mix oxidizers
with sulfur or charcoal or magnesium powder, or whatever.   Sometimes
I'd make mixtures of fuels, such as Mg and sulfur.   Little did I know
at that time that chlorates and sulfur are much more dangerous than
perchlorates and sulfur.

    Even more dangerous was the fact that I owned both red phosphorous
and potassium chlorate.  Had I ever mixed those in any substantial
amount I probably would have lost limb or life.

    The bottom line is this - experimentation should be done from a
position of knowledge, not ignorance.  Of course, there is the old
saying "if we knew what we were doing it wouldn't be research".
However, one should know as much as possible about the topic before
jumping in willy-nilly.  Read up as much as you can, and learn about the
material.  You should also learn about the limitations of reasoning by
analogy...   For example - chlorate is not just perchlorate minus a
bit...   Azides (NaN3) and nitrides (Na3N) are not similar!  While
sulfur and charcoal are both fuels, they are VERY different in their
reactions, etc...

    Take for example, the work being done on basic copper chlorate.  Now
this guy strikes me as being sensible.  He made the stuff, and tested
it's sensitivity to hammering and flame before mixing it up with fuels.
The guy is proceeding in a exemplary fashion.

    Remember, start with small amounts and test for sensitivity,
stability, etc.  Don't ignite mixtures with a match - use fuse or
electric ignition.

2)  Justification -

    WHY are you doing this?   Okay, there can be many answers - ranging
from "I just want to" to "I think this will give a much better color".
But sometimes the justification should be examined in terms of
cost/benefit analysis.

    The cost should include your risk, but lets set that aside right
now.   The cost will include terms such as cost of materials, cost of
manufacturing, etc.   Benefits will include improved effect, size,
weight,  residue, etc...  Let's examine that for Boron

    Boron is interesting because the combination of it's very low atomic
weight (10.8) along with the high heat of formation of it's oxide (B2O3,
-280 kcal/mole) gives it a very high energgy density.  Compare the value
against aluminum (Al2O3, -399 kcal/mole, at wt Al = 27):

4 B + 3 O2  ->  2  B2O3  + 560 kcals
4 Al + 3 O2  ->  2 Al2O3 + 798 kcals.

At first it may seem like Al gives off more heat, and that is true.  But
you must consider how much fuel you had to burn

4 B = 43.2 grams
4 Al = 108 grams

So that means

B:  560 kcals/43.2 grams = 12.96 kcla/gram
Al:  798 kclas/108 grams = 7.38 kcals/gram

Boron is a much better fuel on a gram per gram basis.  This may make it
seem like a great choice, but in fact, this is only somewhat useful.
Consider a stoichiometric mixture of fuel and oxidizer...  B/KClO4 and

3 KClO4 + 8  B  ->  3  KCl +  4  B2O3
3 moles  * 138.5 grams/mole  +  8 moles  * 10.8 grams/mol   =  501 grams

4 moles * -280 kcal/mole = -1120 kcals
Energy density = -1120 kcals/501 grams = 2.23 kcals/gram
Ratio of oxidizer/fuel -  4.82

3  KClO4 + 8  Al  ->  3  KCl  + 4  Al2O3
3 moles  * 138.5 grams/mole  +  8 moles  * 27 grams/mol   =   631 grams
4 moles * -399 kcal/mole = -1596 kcals
Energy density = -1596 kcals/631 grams = 2.53 kcals/gram
Ratio of oxidizer/fuel -  1.92

Boy - this sure seems wrong when you first see it...  you would expect
the mixture containing the boron to be the winner for energy density
based on the numbers for the pure elements burning in air.   But think
of it this way:    The reactions above are mixed up according to
stoichimetries based on MOLES, not on grams.  While the energy density
of B is great on a gram per gram basis, you need far fewer grams of B in
the mixture.   Thus, the use of kcal/grams is sort of bogus - you should
use calculate the energy density of the reaction, not of the individual

The justification for using boron cannot be based on cost ($70/pound for
tech grade amorphous boron), on a green color (weak green), or on high
energy density.   So why use boron?  Among other reasons, it is highly
stable for long term storage, is less sensitive to accidental ignition,
and forms a slag which is more useful for igniting pyrotechnic objects.

So, lets return to a question I asked earlier this week:  Why use basic
copper chlorate?   What are the projected advantages?  Consider a normal
perchlorate based blue color formulation:

Red Gum
Copper oxide

The potassium perchlorate acts as an oxidizer, the red gum is a fuel,
the 1 or 2% of charcoal is an ignition and burning propogation
enhancer,  the parlon is a chlorine donor to create copper chloride in
the flame, the dextrin is a binder and the copper oxide is a copper
source and a weak oxidizer.  Since we are on the topic of perchlorate
oxidizers, lets consider copper perchlorate.   Compare that to the
hypothetical formulation:

Red Gum

In this formulation the copper perchlorate would be both an oxidizer and
the copper source.   Consider the potassium ions in the first
formulation - their only job is to hang around with the perchlorate ions
- they are place holders, and they are useeless weight as far as any
useful pryotechny is concerned.  The second formulation would be so much
better in that aspect.  Also, since the expected mode of decomposition
of Cu(ClO4)2 would be to give CuCl2 and oxygen,  we don't need the
parlon...  less wasted weight, perhaps better efficiency at creating
CuCl in the flame.  In the first one the CuO has to find the chlorine
being created and react to give CuCl...

Cu(ClO4)2 absorbs water to form the hexahydrate  Cu(ClO4)2 * 6 H2O.
That means that at least 29% of the weight of the oxidizer would be
useless water.  At this point, the percentage of useful oxygen in the
perchlorate would be 35%, compared to 46% in potassium perchlorate.
Still, this is a useful amount of oxygen.  The amount of copper in the
perchlorate itself is 17%, so a mixture containing 80% oxidizer would
contain about 13% copper.   This is a good improvement compared to the
roughly 8% copper found in many blue formulations.

So, whats the problem?   Copper perchlorate is deliquescent - it absorbs
so much water from the air that it dissolves itslef in a puddle of
water.   This is NOT conducive to making a good firework.   Thus, it's a
great idea, but technically it's not feasable for anything short of
military pyrotechny which can afford to use vapor barriers, etc.

What about basic copper chlorate - Cu(ClO3)2*3Cu(OH)2  (BCC) ???   Well

Mol Wt - 523
% oxygen - 18% from chlorate, 9% from CuO formed from Cu(OH)2   total
% copper -  48%   (WOW)
% required chlorine 25% (50% if one considers CuCl as color emitter)

The small amount of available oxygen is going to be a problem.   A
stoichiometric mixture of BCC and charcoal (assume it is carbon, AW =
12) would be either

Cu(ClO3)2*3Cu(OH)2  +  3  C  ->  CuCl2  +  3 CO2 + 3 CuO + 3 H2O


2  Cu(ClO3)2*3Cu(OH)2 + 9 C ->  2  CuCl2 + 3 Cu + 9 CO2 + 6 H2O

The former reaction would require  523 grams of BCC and 36 grams of
charcoal,   while the second would require 1046 grams of BCC and 108
grams of charcoal (roughly 6 and 10% fuel, respectively).  That low of a
fuel content generally means that the heat output of the reaction will
not be adequate to heat and decompose the oxidizer, thus producing a
very poor flame, if any.

Thus, BCC probably cannot be used as an oxidizer by itself.   So, lets
think of admixing it with another oxidizer, such as potassium
perchlorate.   Since the common formulations contain about 8 to 10%
copper, lets use that as a goal for BCC.   I'm not going to bother to do
a complete optimization, so as a guess lets look at:

KClO4   65
BCC   20
Charcoal   15

65 g KClO4/138 grams/mole = .47 moles
20 g BCC/523 grams/mole = .04 moles
15 g charcoal /12 grams/mole = 1.25 moles

assuming KClO4 -> KCl +  2 O2,  we get 1.88 moles of oxygen atoms (not
O2) from the KClO4
assuming BCC -> CuCl2*3Cu(OH)2   we get .228 moles of oxygen atoms from
Total of oxygen atoms 2.1 moles
So, C + 2 O   -> CO2,  we need 2.5 moles of O atoms,  this is just a bit
oxygen poor.  Thats good.

However, only 25% of the copper atoms in the BCC have chlorine attached
to them.  We have to add parlon or PVC...  that would require a
recalculation to determine how much.  But if we add parlon, what do we
replace - charcoal, BCC or KClO4?   Parlon is not a fuel, so replacing
charcoal will make it oxygen rich and reduce the temp of the flame.
So, maybe add 3% parlon in the following fashion:

KClO4 - 64
BCC - 19
Charcoal - 14
Parlon - 3

Okay,  this still won't be good, since red gum is better at making
colors than charcoal.  So, we replace charcoal with red gum.  Add

KClO4 - 64
BCC - 19
Red gum - 12
Charcoal - 2
Parlon - 3
Dextrin -  5% additional

So, what is the advantage?  Well, we have about 25% of the copper atoms
directly associated with chlorine - maybe.  But even so, 3/4's of the
copper will still have to undergo a reaction wherein chlorine in the
flame will have to find the copper, and form CuCl in situ...   But why
did I say "maybe"?  Because there is good evidence that chlorates
decompose via the following sort of reaction:

2 KClO3 ->  K2O + ClO2 + ClO3
2 ClO2 -> Cl2O2 + O2
Cl2O2 -> Cl2 + O2


The main point is tht the metal chlorine bond is broken FIRST, thus
giving no advantage to a copper chlorate over a potassium salt... at
least as far as the formtion of CuCl is concerned.   Obviously this
needs to be studied in greater detail.

So, if the material doesn't give a better color, then what is the
justification for using it?  Probably none.  It won't be cheaper, it
isn't more efficient, etc...    However, research must be done to see if
it will work better.  But it has to be done in a rational sane
fashion.   Not some willy-nilly method of slapping shit together to see
what happens.  And that is what separates the quick from the dead.

Tom Perigrin.