Sciencemadness Discussion Board

Anorganic peroxides

Mildronate - 12-9-2010 at 08:09

I am interesed in making anorganic peroxides like magnesium peroxide, calcium peroxide, cadmium peroxide, zinc peroxide maybe any else? Anbody had maded this? How? How good they are for energic materials?

kclo4 - 12-9-2010 at 09:00

Inorganic peroxides, I think would be more proper.

I have never made them, but alone they will not be explosive - if they are mixed with a fuel/combustible mixture they will likely react very energetically, or even with water to form hydrogen peroxide.

I believe barium peroxide is used, or at least used to be used to produce hydrogen peroxide on an industrial scale.

Sodium and potassium or the oxides if oxidized/burnt correctly can produce sodium peroxide and potassium peroxide or superoxide.

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


pjig - 12-9-2010 at 09:27

Barium peroxide..., can a salt of barium sulfate be used with 30% hydrogen peroxide to create barium peroxide? Or would it have to be a oxide ?

densest - 12-9-2010 at 11:07

Well, barium sulfate -is- a salt of barium and sulfuric acid... Heating BaSO4 to 1600C decomposes it to BaO + SOx. Between 450C and 600C BaO absorbs oxygen producing BaO2. BaSO4 is for most purposes insoluble in anything useful.

All from the Merck Index - a book I bought and suggest that anyone who can afford it buy it as an incentive to the publisher.

mnick12 - 12-9-2010 at 14:31

I have both barium and strontium peroxide, none of the two are very interesting for energetic purposes. When mixed with magnesium powder they produce a very bright strobe type effect, certainly not explosive though. When mixed with ammonium perchlorate sugar and starch they make wonderful flares with very nice colors. The best out of these was a mix of barium peroxide,barium chlorate, a small amount of finly powder Mg, and some sugar. It was by fare the most beautiful emerald green that I have ever seen, to bad I didnt write down the ratios it was a few years ago. So as far as energetic oxidizers they they blow, but for coloring pyrotechnics they are wonderful.

woelen - 12-9-2010 at 23:07

You can make very powerful inorganic peroxo complexes, e.g. potassium tetraperoxochromate(V), K3CrO8. Another very energetic peroxide is the dark blue K2Cr2O12. Both these compounds are stable and can be prepared relatively easily. They can be made to explode on their own, but with suitable reductors they are really amazing.

Another interesting one is Cr(NH3)2(O2)2, diammine chromium(IV) diperoxide. The latter is not stable though and should not be kept around in more than sub-gram quantities!

You also can make a very energetic vanadium compound, KVO4.2H2O. This also shows a very energetic reaction with suitable reductors.

Have a look in the "exotic oxidizers" thread for more info on these.

pjig - 20-9-2010 at 20:35

So are you saying that if I start with baSo4 I can end up with BaO2 through heating it at the described temps? Is this right?
Quote: Originally posted by densest  
Well, barium sulfate -is- a salt of barium and sulfuric acid... Heating BaSO4 to 1600C decomposes it to BaO + SOx. Between 450C and 600C BaO absorbs oxygen producing BaO2. BaSO4 is for most purposes insoluble in anything useful.

All from the Merck Index - a book I bought and suggest that anyone who can afford it buy it as an incentive to the publisher.

CZip - 21-9-2010 at 11:08
Quote:

So are you saying that if I start with baSo4 I can end up with BaO2 through heating it at the described temps? Is this right? Quote: Originally posted by densest Well, barium sulfate -is- a salt of barium and sulfuric acid... Heating BaSO4 to 1600C decomposes it to BaO + SOx. Between 450C and 600C BaO absorbs oxygen producing BaO2. BaSO4 is for most purposes insoluble in anything useful. All from the Merck Index - a book I bought and suggest that anyone who can afford it buy it as an incentive to the publisher.


Yea, thats possible, but this proces is for industry. Belive me, decomposition of BaSO4 is hard - not impossible but hard. If you dont have a good torch or somthing very hot (mabe stove..?), you dont have to try it.
If thing that better way is using Ba(NO3)2 or Ba(OH)2

The WiZard is In - 21-9-2010 at 13:59

Quote: Originally posted by CZip  

Yea, thats possible, but this proces is for industry. Belive me, decomposition of BaSO4 is hard - not impossible but hard. If you dont have a good torch or somthing very hot (mabe stove..?), you dont have to try it.

If thing that better way is using Ba(NO3)2 or Ba(OH)2


My copy of —

Pradyot Patnaik
Handbook of Inorganic Chemicals
McGraw-Hill 2003
[Which has a black line on its bottom - and you know what that means.] Sez —

...prepared by heating barium oxide with air or oxygen at 500oC.

Barium oxide from barium carbonate heated w/ coke, carbon
black or tar.

...by thermal decomposition of barium nitrate.

NB — Water soluble barium compounds are toxic me thinks
they are cumulative poisons.

Contrabasso - 21-9-2010 at 14:34

Pyrotechnics by Shidlovskiy does have slight mention of metal peroxides but not much and not in significant compounds.

pjig - 21-9-2010 at 16:50

Huh.., thats strange,,, I always wondered what some claim as a exotic flash broke down to...> 1:1 barium sulfate and fine aluminium.
It is claimed to be a sort of flash/thermite ... Im wondering if it is due to the decomposition of the baso4 turning to peroxide under extreme heat, then acting as a oxidizer for the aluminium.. I have tried this mix before out of curiosity, but never got it to do more than smolder in a molten form.

mnick12 - 21-9-2010 at 17:13

Well BaO2 is pretty cheap it can be bought from pyro suppliers. As for BaSO4 flash powder, BaSO4 alone is a terrible oxidizer, but in combination with perchlorate it makes some very impressive comps. I had one with mg powder and KClO4 that would self-confine in 1gr quantities.

The WiZard is In - 22-9-2010 at 07:34

Quote: Originally posted by pjig  
Huh.., thats strange,,, I always wondered what some claim as a exotic flash broke down to...> 1:1 barium sulfate and fine aluminium.

It is claimed to be a sort of flash/thermite ... Im wondering if it is due to the decomposition of the baso4 turning to peroxide under extreme heat, then acting as a oxidizer for the aluminium.. I have tried this mix before out of curiosity, but never got it to do more than smolder in a molten form.



Page 8 PGI Bulletin No. 46, March, 1985
THE FEW, THE PROUD, THE SULFATES
Donald J Haarmann [AKA — The WiZard is In.]

Having lost the original file! This was scanned in. And you
know what that means!!

Most sulfates are not water soluble, are geologically stable
and can be easily and cheaply obtained by mining, rather
than having to be produced through complicated and expensive
chemical processing. Therefore sulfates pass the first test
for possible inclusion in any pyro formula; they are
inexpensive. Indeed native sulfates such as barite (BaSO4)
and celestite (SrSO4) are the starting materials for other
barium and strontium compounds used in fireworks.

Sulfates certainly appear attractive because their oxygen
content compares favorably with that of metal chlorates,
perchlorates and nitrates, as Table 1 illustrates. Also a
comparison of the heat evolved from reaction of aluminum and
various oxidizing agents again shows that sulfates compare
favorably with more familiar pyrotechnic oxidizers. (See
Table 2.)

Table 1
Percent oxygen contained (percent by weight) for various
pyrotechnic oxidizers and sulfates, for the anhydrous
compound.

Nitrate Chlorate Perchlorate Sulfate
Ammonium 0.60 0.47 0.54 0.48
Barium 0.37 0.32 0.38 0.27
Calcium 0.58 0.46 0.41 0.47
Copper 0.51 0.42 0.49 0.40
Gadolinium 0.42 0.35 0.42 0.32
Lithium 0.69 0.53 0.60 0.58
Magnesium 0.65 0.50 0.57 0.53
Potassium 0.47 0.39 0.46 0.37
Sodium 0.56 0.45 0.52 0.45
Strontium 0.45 0.38 0.45 0.47

Table 2
Heat produced (cal/g) from a mixture of an oxidizer or
sulfate with aluminum. Values from AMCP 706-185(1967) and/or
Vasilev (1973) (*).

Sodium perchlorate 2,600
Lead nitrate 1,500
Sodium chlorate 2,500
Barium nitrate 1,400
Potassium perchlorate 2,400
Cu sulfate 1,400/1,560*
Potassium chlorate 2,200
Ca sulfate 1,300/1,470*
Sodium nitrate 1,800
Na sulfate 1,200/1,360*
Potassium nitrate 1,800
K sulfate 1,200/1,180*
Lithium sulfate 1,620*
Barium sulfate 900/910*
Magnesium sulfate 1,610*
Lead sulfate 800
Ammonium nitrate 1,600

However, low cost is not the only criteria for selecting
oxidizers for use in fireworks compositions. A quick check
of Table 1 reveals several oxidizers with high oxygen
content, for instance, calcium, sodium, and ammonium
nitrates, sodium chlorate, and magnesium perchlorate.
However, of these only sodium nitrate has found use, albeit
limited primarily to military pyrotechnics. All of these
compounds are hygroscopic and therefore unusable in the real
world. In fact, magnesium perchlorate is used as a drying
agent under the trade name of "Anhydrone".


There can be no doubt that the largest problem concerning
the use of sulfates as oxidizing agents is their waters of
hydration, for example:

Na2SO4-10H2O and CuSO4-5H2O. Although the ten extra oxygen
atoms in sodium sulfate raise its total oxygen content from
45% to 70%, this extra oxygen contained in the waters of
hydration is not available for productive work. In truth it
only gets in the way, since a large amount of heat is
required to first remove the water of hydration from a
composition's outer surface before the ignition temperature
can be reached. Then once the reaction becomes self
sustaining, even more heat, produced by a burning star for
instance, will be removed from the reaction in the form of
vaporized water. (It should be noted that the latent heat of
vaporization for water is 540 calories per gram of water at
100° C. This value represents heat that must be supplied by
the pyrotechnic reaction to change water at 100° C into
steam at 100° C.) There is also the possibility, in
magnesium containing compounds, of the water vapor reacting
with the magnesium forming hydrogen and magnesium oxide,
effectively removing a large amount of fuel, with little
gain in heat. In the case of sodium sulfate decahydrate,
where 56% of each molecule is water, 31,920 calories of heat
would have to be supplied simply to remove all the water of
hydration in the form of steam from each 100 grams of
sulfate. For example, in a composition using potassium
perchlorate as the oxidizer and aluminum as the fuel, 13.3
grams of aluminum and potassium perchlorate would be needed
just to remove the water from each 100 grams of sodium
sulfate decahydrate, before any useful work (heat and/or
light) would be produced!

As a further complication, the temperature at which waters
of hydration are liberated varies from sulfate to sulfate,
e.g., sodium sulfate decahydrate loses all its water at 100°
C while manganese sulfate monohydrate does not lose all its
water until the temperature reaches 400-450° C! And to
really complicate things, manganese(II)sulfate can exist as
either mono, tri-, tetra, penta, hexa, or heptahydrate!
Although the tetrahydrate is the most common form.

However, US Patent 2,885,277 claims to make use of the
waters of hydration in magnesium sulfate heptahydrate,
MgSO4-7H2O (Epsom salts), to produce hydrogen gas when the
sulfate is reacted with magnesium. It is further claimed
that this combination will function as either a torch or a
salute. It would be well to note that Ellern (1968, p. 272)
expresses doubt concerning the safety and utility of such
mixtures.

The use of sulfates as oxidizers suffers from yet another
problem. As Dr. Conkling (in press) has pointed out "In
pyrotechnics, the solid liquid transition appears to be of
considerable importance in initiating a self propagating
reaction. The oxidizing agent is frequently the key
component in such mixtures, and a ranking of common
oxidizers by increasing melting point bears a striking
resemblance to the reactivity sequence for these materials."
Unfortunately the melting point of most sulfates is much
higher than either chlorates, perchlorates or nitrates. Only
four sulfates (manganese, copper, zirconium and iron) have
melting points below that of barium nitrate, and these four
are well hydrated (tetra or penta). Melting points are
summarized in Table 3.


Table 3
Melting point for various anhydrous oxidizers and sulfates.
Values are from the CRC Handbook. d decomposes, sd slight
decomposition.

Copper perchlorate 82
Ag perchlorate 486
Iron perchlorate >100d
Thorium nitrate 500
Strontium chlorate 120d
Th perchlorate 501
Lithium chlorate 128
Ba perchlorate 505
Scandium nitrate 150
Sr nitrate 570
Manganese(III) sulfate 160d
Ba nitrate 592d
Americium nitrate 170
Zn sulfate 600
Copper sulfate 200sd 650d
Th(I) sulfate 632
Silver chlorate 230
Silver sulfate 652
Lead chlorate 230
Mn(II) sulfate 700
Lithium perchlorate 236
Lithium sulfate 845
Sodium chlorate 248
Nickel sulfate 848
Magnesium perchlorate 251d
Sodium sulfate 884
Lithium nitrate 264
Ytterbium(III) sulfate 900
Calcium perchlorate 270
Yttrium sulfate 1000
Sodium nitrate 307
Cesium sulfate 1010d
Rubidium nitrate 310
Rubidium sulfate 1060d
Potassium nitrate 334d
Potassium sulfate 1069
Calcium chlorate 340
Samarium sulfate (basic) 1100
Potassium chlorate 356
Magnesium sulfate 1124d
Potassium perchlorate 400d
Lanthanum sulfate 1150
Zirconium sulfate 410d
sulfate 1170d
Cesium nitrate 414
Calcium sulfate 1450
Barium chlorate 414
Barium sulfate 1480
Iron sulfate 480d
Sr sulfate 1605d
Sodium perchlorate 482

It is evident that getting compositions based on sulfates as
oxidizers to ignite while not impossible ... is not going to
be easy. There can be no doubt that it is going to take an
extremely hot ignition source!

Copper sulfate with its low melting point looks like a prime
candidate but again, the water of hydration is a problem.
Exposed to moist air, CuSO4 becomes CuSO4-H2O, and when
wetted, CuSO4-5H2O. Also, because copper sulfate is water
soluble, it is seldom found in native form (chalcanthite).
Therefore it is manufactured from copper metal and sulfuric
acid, and as a result fails the first test, it is not cheap.
It is also not safe with chlorates.

Although certainly attracting because of their low cost
oxygen content, sulfates have for the most part, not been
employed as oxidizing agents. However, them have found their
niche in strobe formulas.

Vander Horck (1974) reported on several formulas using
calcium and copper sulfates demonstrated to him by Bob
Winokur who later (Winokur, 1974) made additional comments
about them. Further Dr. Shimizu (1981) presents several
strobe ("twinkler") formulas using sulfates, i.e.,
strontium, barium, sodium and calcium. Advantage is taken of
the great difficulty of igniting and then sustaining
ignition in sulfate based compositions. Therefore flashes of
light are produced each time the sulfate reaches its melting
point or decomposition temperature, burning commences and
shortly thereafter extinguishes only to repeat, producing
the strobe light effect.

Sulfates have long been used in color flame compositions
more for their metal than oxygen content. However, for the
most part, the color produced by sulfate based compositions
not containing metal fuels such as aluminum or magnesium,
will be found to be less than satisfactory, since only metal
fuels are capable of producing the high temperatures
necessary to melt or decompose most sulfates. The use of
various sulfates is detailed below:

Copper sulfate: In older literature, e.g. Kentish (1878)
compositions for blue flames can be found using copper
sulfate and potassium chlorate, where the copper ion is used
to produce the blue color. THIS COMBINATION IS DANGEROUS.
Safer and more effective blue formulations are available.

Barium sulfate: Troy Fish (1981) recommends the use of
barium sulfate in parlon bound green stars. He notes that as
a result of barium sulfate's extreme insolubility (0.000413
grams per 100 ml of boiling water!), it is one of the few
nontoxic barium compounds. I have been able to locate only
seven formulas using barium sulfate, and all seven use
either magnesium, aluminum or magnalium.

Calcium sulfate: Despite the many obstacles noted above,
calcium sulfate hemihydrate (plaster of Paris) [CaSO4-
1/2H2O] has been used as an oxidizer in fireworks and
pyrotechnics: In combination with sodium and barium nitrate
in white light compositions (Ellern, 1968, formulas 36, 37
and 38), as an incendiary when combined with aluminum (US
Patent 2,424,937, Vol. 3 of the "Black Book", 1982), or
aluminum and magnesium sulfate (US Patent 4,381,207), and
when compounded with aluminum, Teflon, and sulfur (US Patent
4,349,396) as a metal cutting torch.

Calcium sulfate combined with either aluminum or magnesium
has been suggested as a "flash report" mixture! (Sanford,
1974)

This sulfate is found in pink tableau fire or star
compositions using potassium perchlorate as the oxidizing
agent. Weingart (1947) has the only modern for
mula I have been able to locate that uses calcium sulfate
without either aluminum, magnesium or magnalium.

Potassium sulfate: The Technico Chemical Receipt Book 1896
long ago recommended the use of potassium sulfate in blue
compositions. There is only one modern formula using
potassium sulfate, Dr. Shimizu's white "twinkler" using
magnalium as the metal fuel.

Strontium sulfate: This sulfate had long ago been used in
the production of red or purple flames. However, there are
no formulas using strontium sulfate in Lancaster, Ellern or
Weingart. There are however, three "twinkler" formulas in
Shimizu using strontium sulfate. All three contain
magnalium.

Sodium sulfate: I have been able to locate only four
formulas using sodium sulfate, all by Dr. Shimizu, who uses
sodium sulfate in combination with magnalium for yellow
strobe stars.

Manganese sulfate: Perhaps the most interesting use of
sulfate is the addition of manganese sulfate (MnSO4 H2O) to
aluminum sodium nitrate flare compositions. Farnell et
al.(1972) discovered that this compound alters "the
decomposition of sodium nitrate to form oxides of nitrogen
rather than its normal decomposition products of nitrogen
and oxygen." This change results in a 55% decrease in
burning rate, a 155% increase in luminous output, and a 466%
increase in luminous efficiency!

Although not a mainstays of the fireworks trade, sulfates
have found employment along with the proverbial kitchen
sink, used frying pans, oil of spike and philosopher's
wool!!!

Literature cited

AMCP 706
185, 1967, Engineering Design Handbook, Military
Pyrotechnics Series
Part 1; Theory and Application. NTIS AD 817071.

Black Book, 1982, Improvised Munitions Black Book, Vol. 3.
Desert Publications.

Conkling, J., (in press), The Chemistry of Pyrotechnics and
Explosives: Basic Principles and Theory. Marcel Dekker, New
York.

CRC Handbook of Chemistry and Physics, 1981, 62nd edition.

Ellern, H., 1968, Military and Civilian Pyrotechnics.
Chemical Publishing Inc., NY.

Fish, T., 1981, Green and other colored flame metal fuel
compositions using parlor. Pyrotechnica Vll, pp. 25
37.

Farnell, Westerdahl and Taylor, 1972, The Influence of
Transition Metal Compounds on the Aluminum
Sodium Nitrate Reaction. Third International Pyrotechnics
Seminar.

Kentish, T., 1887, The Pyrotechnists Treasury, The Complete
Art of Fire
Making. Chatto and Windus, London.

Sanford, R., 1974, Plaster of Paris flash powders, American
Pyrotechnist Fireworks News, p. 527.

The Technico Chemical Receipt Book 1896.

Merck Index, 1983, The Merck Index: An Encyclopedia of
Chemicals, Drugs, and Biologicals. Merck and Co., 10th
edition.

Shimizu, T., 1981, Fireworks: The Art, Science, and
Technique. Maruzen Publishing Co.

US Patent 2,424,937, July 1947, Incendiary Composition.

US Patent 2,885,277, May 1959, Hydrogen Gas Generating
Propellant Compositions.

US Patent 4,349,396, September 1982, Metal
Cutting Pyrotechnic Composition.

US Patent 4,381,207, April 1983, Pyrotechnic Composition.

Valsilev, A.A., et al., 1973, Combustion of mixtures of
metal sulfates with magnesium or aluminum. Translated from
Russian. NTIS AD 785988, 5 pp.

Vander Horck, M.P., 1974, Unconventional star compositions
demonstrated. American Pyrotechnist Fireworks News, 7(4),
issue no. 76, p. 506.

Weingart, G. W., 1947, Pyrotechnics. Chemical Publishing
Co., NY, pages 61 and 134.

Winokur, R., 1974, More on unconventional stars. American
Pyrotechnist Fireworks News, 7(5), issue no. 77, p. 516.


djh
---
This is what I thought: for the most banal event to become an adventure
you must (and this is enough) begin to recount it. This is what fools people: a
man is always a teller of tales, he lives surrounded by his stories and the
stories of others, he sees everything that happens to him through them; and
he tries to live his own life as if he were telling a story.

But you have to choose: live or tell. For example, when I was in Hamburg, with
that Erna girl I didn't trust and who was afraid of me, I led a funny sort of life.
But I was in the middle of it, I didn't think about it. And then one evening, in a
little cafe in San Pauli, she left me to go to the ladies' room. I stayed alone,
there was a phonograph playing "Blue Skies." I began to tell myself what had
happened since I landed. . . . Then I felt violently that I was having an
adventure. But Erna came back and sat down beside me, she wound her arms
around my neck and I hated her without knowing why. I understand now: one
had to begin living again and the adventure was fading out.

Nothing happens while you live. The scenery changes, people come in and go
out, that's all. There are no beginnings. Days are tacked on to days without
rhyme or reason, an interminable, monotonous addition. . . .

That's living. But everything changes when you tell about life; it's a change no
one notices: the proof is that people talk about true stories. As if there could
possibly be true stories; things happen one way and we tell about them in the
opposite sense. You seem to start at the beginning: "It was a fine autumn
evening in 1922. I was a notary's clerk in Marommes." And in reality you have
started at the end. It was there, invisible and present, it is the one which gives
to words the pomp and value of a beginning. "I was out walking, I had left the
town without realizing it, I was thinking about my money troubles." This
sentence, taken simply for what it is, means that the man was absorbed,
morose, a hundred leagues from an adventure, exactly in the mood to let
things happen without noticing them. But the end is there, trans-forming
everything. For us, the man is already the hero of the story. His moroseness,
his money troubles are much more precious than ours, they are all gilded by
the light of future passions. And the story goes on in the reverse: instants
have stopped piling them-selves up in a lighthearted way one on top of the
other, they are snapped up by the end of the story which draws them and
each one of them in turn, draws out the preceding instant: "It was night, the
street was deserted." The phrase is cast out negligently, it seems superfluous;
but we do not let ourselves be caught and we put it aside: this is a piece of
information whose value we shall subsequently appreciate. And we feel that
the hero has lived all the details of this night like annunciations, promises, or
even that he lived only those that were promises, blind and deaf to all that did
not herald adventure. We forget that the future was not yet there; the man was
walking in a night without forethought, a night which offered him a choice of
dull rich prizes, and he did not make his choice.

I wanted the moments of my life to follow and order themselves like those of a
life remembered. You might as well try and catch time by the tail.

Jean-Paul Sartre Nausea

pjig - 22-9-2010 at 11:57

This still doesnt show much on baso4 with aluminum. I know its a hard to find mix, and im sure its inferior, but still a bit interesting. :P