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blogfast25
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[*] posted on 8-8-2008 at 05:07


Na<sub>3</sub>AlF<sub>6</sub> thermodata:

here they are. MP ≈ 1,000 C.

For my own experiments I would simply substitute CaF2 by Na3AlF6 mol per mol, then reduce the amount by 20 - 30 %: because of cryolite's lower MP compared to CaF2 its slag fluidising effect should be stronger, as liquid Na3AlF6 at thermite end-temperatures should be even less viscous than CaF2.

Out of curiosity, what's your interest in AlBr<sub>3</sub>? :)
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[*] posted on 8-8-2008 at 14:13


No REAL interest in aluminum bromide since I prefer inorganic chemistry over organic. I guess its I want to take a video and say hey I made this and I did so unconventionally. Besides aluminum bromide vapors will look cool on video:cool:

I received my magnesium ribbon today so I can use this to ignite my thermites now. I will try the chlorate/aluminum comp although I will be extra cautious. At what mesh does this mix be come flash?

Thanks for the data!

P.S. Please keep me informedon your halide explorations. If you can shoot a video I would like to see it.

[Edited on 8/8/2008 by chloric1]




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[*] posted on 9-8-2008 at 06:28


Quote:
Originally posted by chloric1
Besides aluminum bromide vapors will look cool on video:cool:

I received my magnesium ribbon today so I can use this to ignite my thermites now. I will try the chlorate/aluminum comp although I will be extra cautious. At what mesh does this mix be come flash?

Thanks for the data!

P.S. Please keep me informedon your halide explorations. If you can shoot a video I would like to see it.

[Edited on 8/8/2008 by chloric1]


The problem with thermite reactions (and analogues) that produce gaseous reaction products in significant quantities is that it's precisely that that can cause the reaction mix to deflagrate ("explode" as some would call it). MnO2 thermites are an example: with very fine ingredients such termites simply go "poooofff!" because the heat generated vaporises the formed manganese. Other thermites with quite low HoF oxides/relatively low boiling metals tend to do that too.

In the case of the PbBr2/Al reaction, reaction speed should be low and temperature accordingly low because of the small heat of reaction.

Chlorate/Al is in my experience very safe. I use fine potassium (reagent grade) with 400 mesh Al powder: never had a problem. Not sure whether there is a mesh size at which this mixture tends to flash. I did read somewhere that including sulfur in the mix can lead to a shock-sensitive mixture. Yet for some time (but not anymore) I did used to spike the chlorate/Al mix with small amounts of S, to increase heat output. I had no problems doing this.

If you're a little jittery about using chlorate I suggest to mix the chlorate with the oxide first, then add the Al powder. But I usually simply mix everything together w/o problems. Chlorate/Al booster systems are used to spike thermites (for metal production) in industry all the time.

With Mg powder it becomes a different story: chlorate/Mg is supposed to be a flasher. I've never tried it but I did once try an SiO2 thermire with magnalium (50 Al/50 Mg), boosted with chlorate. The thermite did flash (no explosion though), leaving an empty crucible....

As regards vids, I'll try. But my experience is that most successful reactions (success being defined here as 'obtaining good, clean metal') look very similar. They run fast, to white heat (to well above the MP of alumina and possibly up to 2,500 C), well contained and at the end most of the slag and metal are found as a flat, white hot, molten puddle at the bottom of the crucible. Below are some vids of the TiO2 experiments, unboosted and boosted with KClO3. Scroll down a bit to the "530 grams of KClO3-boosted..." and watch the video: this is what thermite heaven looks like: this one is shot by Jeffrey using my TiO<sub>2</sub> formulation (slightly adapted). See the white-hot, flat puddle of molten alumina and molten metal gathered neatly at the bottom. That's my baby!

All my successful thermites with transition metals look strikingly similar, even though reaction times may vary a little. Temperature is the great leveler here: if the reactions are designed to achieve more or less the same end-temperatures (as mine are) they will tend to run quite similarly.

The MnCl2/Mg reaction will be the exception to the rule because here I need to keep the end-temperature purposely low (considerably lower than the BP of Mn: that's why the chloride/Mg route was chosen in the first place).

[Edited on 9-8-2008 by blogfast25]
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[*] posted on 9-8-2008 at 08:31


amazingrust is indeed your site!:D I like to visit this alot. Just want to let you know that for some reason your vid links are down. I get Internet Explorer cannot load page. Let me know if you can get the vids back up.



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[*] posted on 9-8-2008 at 09:40


No, no. AmazingRust.com isn't my website, it belongs to a guy called Jeffrey with whom I'm friendly, that's all. When I checked a few vids less than an hour ago they were fine. Temp glitch, I feel...
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[*] posted on 9-8-2008 at 17:40


Oh OK my mistake I wonder if there is something wrong on my end. I mistook your enthusiasm for the titanium vid as a praise of your own creation. None the less you have been a tremendous help to me and I sincerely thank you. I am starting to really look at the thermochemistry of this to get a better feel for this. One thermite I have been thinking about and it might need coarser aluminum becuase of the activity, is barium sulfate. One product barium sulfide is easily soluble in hot water and can be reacted with excess KOH or NaOH to precipitate the barium hydroxide on cooling to near zero celsius.



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[*] posted on 10-8-2008 at 05:22


Yes, the video is his but the development of the TiO2 thermite formulation is mine. At the time I had emailed Jeffrey about it almost immediately and the same day he enthusiastically reproduced my initial results and made the video. The initial find was reported on this SienceMadness thread, where Jeffrey also commented as mrjeffy321 and reported confirming my results (looking through the thread I see that you were there too).

The barium sulfate thermite will almost certainly work, no question about it: I've run countless such reactions with dried plaster of Paris according to:

CaSO4 + 8/3 Al ---> CaS + 4/3 Al2O3

That reaction releases a hell of a lot of energy and I've even used it as a booster reaction for TiO2 thermites, successfully. In terms of ΔH of the overall reaction it's more or less the same as the chlorate reaction and thermochemically speaking both can be used more or less interchangeably.

But to get access to the reaction product BaS may prove difficult, mixed in with alumina as it is. The slag from the CaSO4 reactions positively reeks of H2S because of hydrolysis of the sulfide, so if you treated it with HCl you'd leach out some CaCl2, no doubt about that. The question is how much? It may well be worth looking into that, using CaSO4 as a model, before you start sacrificing the more expensive BaSO4... The yields may be surprisingly disappointing, not sure though... Size-reducing the slag for the purpose of extraction will also be difficult: fused alumina is incredibly hard.

Perhaps unusually high levels of CaF2 (or Na3AlF6) may cool the reaction enough, so that a porous, rather than compact slag heap results. CaF2 (and Na3AlF6) is also much softer than Al2O3 which would make a softer slag, easier to size reduce, essential to successfully extract as much BaCl2 as possible...

An approximate thermocalc will allow you to estimate the amount of heat-sink (CaF2 or Na3AlF6) needed to cool the reaction, say to just below the MP of alumina, thereby obtaining a porous 'slag muffin', much easier to break up, size reduce and extract the BaS from. The level of heat sink needed may however be so high that the mixture proves difficult (or impossible) to ignite because the heat balance equation makes no pronouncements about the kinetics of the reaction.

All in all a rather smelly method to access soluble Ba salts but well worth contemplating I feel... It will generate a mol of H2S per mol of soluble Ba, that's a helluvalot of rotten eggs!

And thanks for your appreciation. Frankly I love nothing more than to expand a little about my experiences with thermites and assorted reduction thermochemistry because I feel far too many backyard experimenters (but not all, of course) take an approach of "let's stick a bit of magnesium ribbon in it and keep our fingers crossed". But as with all reactions, there's a more than a bit science needed to get the results wanted. This is as much true if you want to create spectacular pyrotechnics or if producing small amounts of relatively pure metal is your purpose...

[Edited on 10-8-2008 by blogfast25]
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[*] posted on 10-8-2008 at 07:30
First reduction of MnCl2 with Mg powder


I ran the first 20 g batch (stoichiometric mix 1:1 moles) of MnCl<sub>2</sub> + Mg today.

Lit with a few g of KClO3/Al/Mg ribbon fuse, it burned right through, slowly but nevertheless quite hot and with a continuous yellow flame, a few cm high. There was no unreacted mix left, at least going by visual inspection.

The slag/metal mixture was very predictably of the 'porous slag muffin' type, so typical of heat-starved thermites. The slag looked like what I would imagine sintered anhydrous MgCl<sub>2</sub> to look like. Predictably, no Mn reguli to be found.

I recovered all the slag and ground it down to powder in my granite mortar and pestle. I transferred the powder into a Pyrex measuring jug and added dionised water. Considerable heat evolved, presumably the heat of hydration of MgCl<sub>2</sub>. There is also gas evolution which I presume to be hydrogen, from Mn + 2 H<sub>2</sub>O ---> Mn(OH)<sub>2</sub> + H<sub>2</sub> (Mn is known to react with hot water). Right know this thin suspension is grayish. If the gas evolution stops, I'll filter and test for Mg and Mn. The black-gray solid matter, assuming it is indeed powdered Mn, should react strongly with HCl.

All in all, this mixture, without any heat boosting measures, does what it says on the thermochemical tin.
Much work now lies ahead to try and heat boost it to obtain lump Mn metal.

The production of anhydrous MnCl2 will also have to be stepped up and refined: I know the product contains residual NH4Cl (from the drying process) and that it contains small amounts of oxide (going by the colour).
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[*] posted on 11-8-2008 at 08:16


I have about a hundred grams of UO3, and out of curiosity would like to try making some U metal, but only a gram or two at a time (to reduce wastage if it doesn't work).
Obviously though the charge will have to be above a minimum size to allow it to stay hot enough for long enough for the metal droplets to sink and coalesce.
I was thinking of putting UO3, Mg and some sort of flux at the bottom of a large graphite crucible, with a charge of KClO3, Al and Al2O3 (to reduce burn rate) on top.
What I hope would happen would be that the charge on top would burn down and produce lots of very hot liquid, and then ignite the thermite mix at the bottom. The Mg would then reduce the UO3, and due to all the alumina slag on top everything would stay hot for long enough to get a nice blob of U. I also hope that the slag on top would reduce the amount of U that escapes in the fumes.
What do you think of this plan? Also, any ideas for a suitable flux? My initial thought was KF, but then I worried about the volatility of UF6.
Thanks for any advice!
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[*] posted on 11-8-2008 at 09:02


Nick F:

I'm a little pressed for time right now so I'll have to keep this short.

IMHO and at first glance, for the reduction of UO<sub>3</sub>, Al would be a much better reductant: Mg is likely to be far too energetic, yet the insanely high MP of MgO makes it difficult to obtain both the nascent U metal and the MgO in liquid form. This is strictly necessary if lump metal is what you're looking to make. With Al, the experimental set-up would be much simpler that what you propose here...

Again and as always, only a thermochemical calculation can tell you what the expected end-temperature of such a reaction would be. If I find the time, I'll have a look later on or tomorrow.

The heat of formation (at 298 K) for UO3 is - 1224 kJ/mol (of UO3), that makes the reaction enthalpy for UO<sub>3</sub> + 2 Al ---> U + Al<sub>2</sub>O<sub>3</sub> (at 298 K) about ΔH = - 452 kJ/mol of UO<sub>3</sub>. That's a respectable value and it might be enough to heat the one mol of U and the one mol of alumina to above the MP of alumina. I have the heat capacities and heats of fusion of both U and alumina in my files so I can calculate the estimated end-temperature for such a reaction (stoichiometric mix) in adiabatic conditions.

[Edit]

A quick thermocalc shows that in adiabatic conditions, a stoichiometric mix of UO<sub>3</sub> and Al would burn to an estimated end-temperature of about 2,300 K (2,030 C), close to the melting point of alumina (2,054 C) (and well above the melting point of U i.e. 1,132 C). To obtain a liquid metal/slag mixture and to allow the U to separate out and settle at the bottom of the crucible, small amounts of KClO3/Al may be needed to boost overall heat generated.


For safety reasons, such a reaction should IMHO, always be carried out in a bomb-type reactor, such as a bomb calorimeter or similar, because thermites will always emit some smoke, which in the case of a uranium based reaction could contain quite lethal microparticles of U or its oxides.

Alternatively a much tamer reaction can be obtained with UCl<sub>4</sub> and Mg (or Li, Na or K). Some basic thermocalcs starting with UCl4 can be found here.

The safety concerns would remain the same but halide reductions can be run at lower temperatures...

[Edited on 11-8-2008 by blogfast25]

[Edited on 11-8-2008 by blogfast25]

[Edited on 11-8-2008 by blogfast25]
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[*] posted on 14-8-2008 at 07:02


Thanks!
I'll let you know if I ever get round to trying it, but it might have to wait until I visit my parents because they have a bigger test area (garden :)).
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[*] posted on 14-8-2008 at 07:29
Dichromate boosted chromium thermite


I ran two K2Cr2O7 boosted Cr2O3 thermites today, both stoichiometric mixes, one with 0.2 mol K2Cr2O7 per mol of Cr2O3, one with 0.4 mol K2Cr2O7 per mol of Cr2O3.

Somewhat to my surprise, neither yielded good quality Cr reguli.

The first one ran fast but clearly not hot enough because a porous slag/metal muffin was obtained. Inside glistening Cr could be seen but all metal was sub mm.

Increasing the dichromate level to 0.4 mol made the thermite run decidedly faster and the slag/metal was largely molten and collected (more or less) at the bottom of the crucible. Breaking open, clear areas of Cr metal could be seen but poorly separated from the alumina slag. Poor metal coalescence due to slightly too low end-temperature (premature freezing of the alumina) is clearly the cause here (I've seen many such cases before).

Success should be possible at 0.5 - 0.6 mol of dichromate but that's much higher than the estimated - 2,800 kJ of booster heat (per mol of dichromate) would suggest...

To be continured...

[Edited on 14-8-2008 by blogfast25]
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[*] posted on 14-8-2008 at 16:01
Dichromate Boosted Thermite


Well this is a better result but my metal yield was paultry. I was going to throw out the the byproduct but I have not and will not. I am going to put small amounts of it where ever chromium may be usefull. Hoping the heat from another charge will help separate more chromium out of its cermet. You can see my improved video here

@blogfast-did you add flux? I am not using flux because it has not been delivered yet. I will post pictures of my metal later. Only one nugget was significant size at just under 1cm.
Also, what kind of metal yeilds do you get from chlorate boosting?
[Edited on 8/14/2008 by chloric1]

[Edited on 8/14/2008 by chloric1]




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[*] posted on 15-8-2008 at 04:15


Chloric-

Yes, all these formulations are fluxed. I set the CaF<sub>2</sub> to a level that is equal to (in mol) CaF<sub>2</sub> = 0.225 Al. This ensures the slag (Al2O3 + CaF2) contains a constant molar fraction of CaF2, for purposes of comparison. Sometimes I use CaF2 = 0.1125 Al, it depends...

The generic formulation of the dichromate-boosted chromium formulation then becomes:

Species ............................................. mol

Cr<sub>2</sub>O<sub>3</sub> ................................................. 1
K<sub>2</sub>Cr<sub>2</sub>O<sub>7</sub> ............................................. x
Al ........................................................ 2 + (14/3) x
CaF<sub>2</sub> .................................................. 0.225 [2 + (14/3) x]

Yields (RECOVERED metal/stoichio metal times 100 %) from chlorate boosting are about 50 - 60 %. It has to be borne in mind that considerable amounts of the metal get locked into the slag that freezes when it hits the cold walls of the crucible. That helps explain why larger thermites tend to yield more recoverable metal: the fraction of metal/slag mix prematurely frozen on the crucible walls tends to be smaller. Also, closed reactors with some degree of thermal insulation (alumina or magnesia lining for instance) should give higher yields.

Your latest test seems to run much hotter: which variable did you change?

I would strongly advise against piling up one charge on the remains of the previous one: thermodynamics being what it is one can show very easily that unless the succeeding charge is much larger or much hotter than the preceding one, the remains of the preceding one will simply not melt (heat up, yes, melt? NO!!)

Now I've got an appointment with an MnCl2/Mg reaction... :)

Update:

Two more tests at x = 0.6 and x = 0.5 were both disappointing. At this level the slag/metal mass was really completely molten, yet metal coalescence was piss-poor. No real Cr reguli whatsoever. I'll try one with a much reduced level of CaF<sub>2</sub>...

-----------

A test trying to use Mg + I<sub>2</sub> ---> MgI<sub>2</sub> as a heat booster for MnCl2 + Mg ---> Mn + MgCl2 failed: the iodine simply fumes off in a spectacular purple cloud, but the main reaction still proceeded.

Ca + I2 ---> CaI2 is sometimes used as a booster in calciothermic reductions but obviously this must only work in closed reactor conditions, where the iodine has nowhere to go but react with free Ca.

Another test using permanganate as a heat booster at 0.04 mol (per mol of MnCl2) was inconclusive. I know, I know: Danger! Flash powder! But at this very small level this is w/o danger, IMHO.

The mixture was for some unknown reason difficult to light and when it finally did start it went on quite irregularly, sputteringly. At one point it seemed to go to white heat but by then the reaction mix had been used up. The slag metal/mix clearly shows a higher end-temperature was achieved: more fusion of the MgCl2 was noticable.

At such low levels of KMnO<sub>4</sub> it's actually quite difficult to mix it in homogeneously. Another attempt at a slightly higher level and with better mixing, soon. Assuming the overall booster reaction is KMnO<sub>4</sub> + 4 Mg ---> K + Mn + 4 MgO, the estimated reaction enthalpy should be in the order of about - 2,000 kJ/mol (of permanganate).

Other tests planned are pre-heating the mix (no booster) by about 300 K, prior to ignition and heating the mixture (no booster) to auto-ignition.

[Edited on 15-8-2008 by blogfast25]

[Edited on 15-8-2008 by blogfast25]
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[*] posted on 15-8-2008 at 15:03


@blogfast- Point taken on the thermodynamics of heating previous runs. I may someday try to make a pot of molten boric acid or borax and add the slag to see if it will help me separate more metal.

As far as run #2 goes:

As opposed to the 0.07mol of potassium dichromate I simply doubled the dichromate and added the additional aluminum to compensate. So I took 100grams of my previous mix and added 9 grams of potassium dichromate and 4 grams of aluminum for a total mass of 113 grams

Also, my ignition system changed as I only used magnesium. This was done in two staged a ribbon burning into partially imbedded shavings in the mix itself.

I have not measured the yield of my chromium yet but I know its not 50 or 60%.

I will do one last trial with remainder of the dichromate/chrome thermite adding the same amounts of dichromate and aluminum but submerging the reaction vessle in a much larger pot filled with sand to hold in heat a little longer. I wish I had some vermiculite around.

After this I will use my other 220+ grams to make chlorate boosted Chrome thermite. I do not know why chlorate would yield more chromium. Unless molten potassium chloride is a really good flux or insulator.

BTW my cryolite arrived yesterday after I posted to this thread:)

[Edited on 8/15/2008 by chloric1]




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[*] posted on 16-8-2008 at 05:14


Chloric-

So you were at 0.14 mol of dichromate.

I'm at a loss as to why this booster system, which clearly works in terms of boosting heat output and thus end-temperature, seems to impede metal coalescence. I've noticed the slag seems a little more glassy compared to what is being obtained with KClO<sub>3</sub>.

I also thought the elemental K might be causing problems, but how??? One suspect here could be interference of the hot, free K with the fluorite, according CaF<sub>2</sub> + 2 K ---> Ca + 2 KF. This has a positive heat of reaction (is endothermic) at 298 K of + 88 kJ/mol (of CaF2), so at first glance cannot proceed. But in many cases such endothermic reactions (at 298 K) do proceed at sufficiently elevated temperatures. Hell, the most used reduction reaction on Earth, the reduction of iron ore with CO in blast furnaces, proceeds only at the cauldron conditions of the furnace!

What's more, at thermite temperatures, KF, with a BP of a mere 1505 C, would be volatile, pushing the equilibrium to the right by removal of the reaction products.

It sounds far fetched but right now it's the only thing I can come up with (LOL). Right now it's merely a hypothesis, nothing more

If I'm remotely right about this then cryolite should suffer even more from 'potassium attack' because AlF<sub>3</sub> + 3 K ---> Al + 3 KF is exothermic at 298 K by about - 200 kJ (per mol of AlF3).

And on KCl in chlorate boosted mixes: KCl has a boiling point of about 1500 C and boils off. Chlorate boosted thermites do indeed tend to be quite smoky.

------------------

On the MnCl2 reduction side of things, I'm thinking of using good ole' MnO<sub>2</sub> (wild horses couldn't drag me away!) according MnO<sub>2</sub> + 2 Mg ---> Mn + 2 MgO, ΔH = - 684 kJ/mol (of MnO2, @ 298 K).

I'll be thermocalcing that this afternoon. If favourable that's me back to making some MnO2. Some things never change, I guess...

+++++++++++

For those interested, here's another example of a thermochemical calculation, applied to an MnCl2/MnO2/Mg stoichiometric mix (in adiabatic conditions).

Main reaction: MnCl<sub>2</sub> + Mg ---> Mn + MgCl<sub>2</sub>, ΔH<sub>298 K</sub> = - 161 kJ/mol (of MnCl<sub>2</sub>;)

Booster reaction: MnO<sub>2</sub> + 2 Mg ---> Mn + 2 MgO, ΔH<sub>298 K</sub> = - 684 kJ/mol (of MnO<sub>2</sub>;)

I set the target end-temperature at 1,800 K, well above the MPs of both Mn and MnCl<sub>2</sub>, but well below the BP of Mn (and below the MP of MgO).

Now we need to find out how many mol of MnO<sub>2</sub> is needed per mol of MnCl<sub>2</sub> for the burn to reach this target temperature.

Assume the amount of dioxide needed is x mol per mol of chloride. The reaction products will then be: (1 + x) mol of Mn, 1 mol of MgCl<sub>2</sub> and 2x mol of MgO.

NIST thermochemical data allows to calculate how much enthalpy is needed to heat these reaction products to 1,800 K, using the relevant Shomate equations.

For Mn we obtain ΔH<sub>Mn</sub> = 58.75 kJ/mol, for MgCl<sub>2</sub> ΔH<sub>MgCl2</sub> = 133 kJ/mol, for MgO ΔH<sub>MgO</sub> = 74.4 kJ/mol.

The total heat of reaction ΔH<sub>R</sub> = - 161 - 684 x and in adiabatic conditions:

ΔH<sub>R</sub> + (1 + x) ΔH<sub>Mn</sub> + ΔH<sub>MgCl2</sub> + 2x ΔH<sub>MgO</sub> = 0

Or:

-161 - 684 x + (1 + x) 58.75 + 133 + 2x 74.4 = 0

and x = 0.06 mol of MnO<sub>2</sub>.

Since as we're not in adiabatic conditions and some heat losses are inevitable, a good starting point formulation (in mol) would be MnCl<sub>2</sub> = 1; MnO<sub>2</sub> = 0.1; Mg = 1.2

[Edited on 16-8-2008 by blogfast25]
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[*] posted on 16-8-2008 at 17:36
Fluoride delema


Blogfast- The idea you proposed does not seem far fetched to me. In fact it seems like a logical reason. The reason I say this is that after the magnesium was consumed I was puzzle by the fact white smoke was forming. I did not care to breath it so I avoided it. Of coarse, as noted before, I had no fluoride at my disposal so I am wondering if the smoke was actually a cloud of potassium hydroxide formed by potassium oxide and water vapor. The physical properties of both potassium compounds are:

Potassium Hydroxde mp 380C bp 1324
Potassium oxide mp 350C with decomposition

I seriously doubt that the smoke is potassium oxide. Next and last run with dichromate booster I will try to condense some the smoke into an inverted beaker in hopes I can do a pH test to see if it is alkaline.


You make very convincing thermodyanic statements regarding possible outcomes of this system. I, now caught up in the whole thermodynamic calculations aspect of it now realize things on a new level. I have some conclusions of my own based on my limited research.

1) Alkaline earth oxides have higher heat of formation than their respective group one cousins.
2) oxidizers of group1 elements that may result in alkaline residues may not function well in thermite boosting compositions (sulfates, permanganates, chromates, persulfates, nitrates etc)

My proposal is if it is desirable to use an oxidizer with one of the above mentioned anions, maybe it should be of the group 2 ilk. Such as: Barium Chromate, calcium nitrate, calcium permanganate etc. Not sure about the sulfates though , will have to calculate that. to explain what I mean here, I have not yet looked at the heat of formation of calcium or barium sulfide. I know aluminum sulfide has lower heat of formation than the oxide but still is high enough that most sulfides can be reduced by aluminum exothermically.

One such composition I want to evaluate it molybdenite(MoS2) with aluminum:D

Of coarse for simplicity and cost one may use the more available sodium/potassium chlorate/perchlorate.

[Edited on 8/16/2008 by chloric1]




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[*] posted on 16-8-2008 at 18:29


BTW, you should be getting potassium vapor. Alkaline oxides (and hydroxides) react with magnesium and aluminum, yielding alkaline metal vapor, which obviously will promptly burn in air. The thermite may simply be too bright to see the purple flame as it burns.

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[*] posted on 17-8-2008 at 06:03


Chloric-

I'm pretty convinced the smoke you see is K<sub>2</sub>O/KOH because the K vapour would oxidise immediately after leaving the crucible. The reaction isn't noticeable because the amounts are small. That's my take on it.

Oxides of metals of higher valence tend to have higher HoF all round, as a general and quite reliable rule. I suspect the lattice energy (Coulomb attraction - energy released when M<sup>m+</sup> and O<sup>2-</sup> form a crystal lattice M<sub>2</sub>O<sub>m</sub>;) is the main cause here: it's much higher for higher valence ionic lattices than for lower ones, leading to higher HoFs, despite the fact that the ionisation energy to obtain M<sup>m+</sup> is much higher than for M<sup>+</sup>.

Ca(NO<sub>3</sub>;)<sub>2</sub> is super deliquescent, so count that one out.

MoS<sub>2</sub>/Al is a very interesting proposition. The heat of reaction ΔH<sub>R</sub> for MoS<sub>2</sub> + 4/3 Al ---> Mo + 2/3 Al<sub>2</sub>S<sub>3</sub> is (at 298 K) a whopping - 841 kJ/mol of MoS2 (MoS2 HoF = - 276 kJ/mol, Al2S3 HoF = - 651 kJ/mol)!!

This might even be enough to reach the MP of Mo (2,896 K). If not, boosting with extra Al and S should do the trick. I have no C<sub>p</sub> or ΔH<sub>fusion</sub> for The Smelly One, so a little, simple experimentation would be needed. Caution not to get H<sub>2</sub>S poisoning. Bleach is a good way to 'neutralise' the rotten eggs gas because hypochlorite oxidises the S (-II) back to elemental S (0), so treating the aluminium sulfide slag with copious amounts of thin bleach is the way to get to the metal, make some sulfur flour in the process and avoid H2S poisoning. Win-win-win, I say...

$$$$$$$$$$$$$$$$

Update:

Dichromate boosted Cr thermite:

Well, well, well: if at first you don't succeed, try, try, and try again...

I repeated the formulation with 0.4 mol of K2Cr2O7 but with the level of CaF2 cut to one tenth (CaF2 = 0.0225 Al), The formulation was thus: Cr2O3 = 1 mol; Al = 3.8666... mol; K2Cr2O7 = 0.4 mol; CaF2 = 0.087 mol. A 17.3 g batch of this mixture was mixed and lit.

It ran very fast and smoothly, resulting in a very flat slag puddle at the bottom of the crucible (that's always a very good sign). Breaking open, I found 1 regulus of Cr metal about 1 cm across (well formed), 1 of about 3 mm across and one of about 1 mm across. In the frozen slag even smaller droplets could be seen, too small to recover. The metal is clean-skinned, non-oxidised. It's chromium metal alright...

The mix contained about 6.60 g of Cr and the reguli weighed in at 4.40 g, a 67 % yield. Not bad, not bad at all for such a small reaction.

This would appear to confirm (but not prove) that interference of K with CaF2 might have been the problem here.

I will repeat a much larger batch with this formulation 'soon'.

[Edited on 17-8-2008 by blogfast25]
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[*] posted on 17-8-2008 at 16:34


Well something to consider. I might try calcium oxide as a flux since the calcium aluminates have melting points between 1500 and 1600 degrees C. The potassium should not interfere with this. not sure what the melting point of the fluidized slag is but I am sure its comparable since I am sure the alumina raises the melting points of cryolite and fluorite.

I still am considering the molybdenite reduction but molybdenite BEGINS to sublime at 450 C according to Merk. Hopefully the reaction is fast enough before too much ore flies away!:o

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Well, all of this browbeating and calculating has got me to kik back for some down to earth laughs and good times. Watch this video here This was some silly toy that came in a happy meal and my 3 year old HATED it calling it stupid bear.

[Edited on 8/17/2008 by chloric1]




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[*] posted on 18-8-2008 at 03:31


MoS2 begins to sublime at 450 C??? That sounds quite impossible, given its polymeric structure. Remember, synthetic MoS2 is used as a high temperature lubricant, an alternative to graphite no less... No, I can't see sublimation being a problem here...

CaO: definitely something to try...

What was the thermite used in the video? Looks like Classic Thermite (Fe2O3) to me...
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[*] posted on 18-8-2008 at 14:02


That is Fe3O4 thermite in that last video. I have yet to try the Fe2O3 thermite. Interesting that in all these videos I am using aluminum powder that is 9 to 10 years old:o I know aluminum powder has been noted to react with water so I figured this stuff wouldn't burn. Well, I was wrong. I tried thermite when I first bought this powder in 1998 or 1999 and failed so put this stuff aside and forgot about it for awhile. Now I obviously know better and I don't have ignition problems.

I forgot to comment on aluminum sulfide oxidation with bleach in my last post so here goes. I don't know how long I will be able to obtain sulfur so this method could be quite important. I hoping the hydrolysis of the sulfide yields a gell like aluminum hydroxide so it can easily be removed from the sulfur residue with 16% HCl. Next run of Chromium thermite will be soon and it will be 0.14 mole potassium dichromate like the last run but will install reaction vessel in larger sand filled one. But a modification will possibly involve a CaO flux. After this it will be a chlorate boosted run. Then I will be almost out of chromic oxide for awhile and any remainder will get mixed with magnetite to make iron/chromium alloys.

I got two pounds of vanadium pentoxide and cuprous oxide in the mail today:D




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[*] posted on 19-8-2008 at 03:40


Aluminium powder is pretty resistant to degradation because it passivates.

Yep, it should be perfectly possible to separate the sulfur from the Al(OH)3 by dissolving the latter in dilute HCl, especially if you get to it fairly quickly... But it'll be a smelly business... You can't get garden grade sulfur in the US???

For chromium, my tests show 0.14 mol dichromate to be too low. Good results with 0.4 mol.

Hmmm, V2O5, one of my favourites: burns very fast and hot because it's a higher oxide (lots of alumina formed) - no booster needed...

This MSDS on MoS2 mentions an MP of 2375 C but says nothing about sublimation...

[Edited on 19-8-2008 by blogfast25]
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[*] posted on 19-8-2008 at 08:37


I might try the higher level of dichromate to see. Last night I was shuffling through my CRC handbook and it gave the high melting point but ALSO said that sublimation begins at 450C. I am quite aware of molybdenites high temp lubrication properties. I roasted it a year or so ago to make molybdates and I did not notice any vapors exept SO2. I do not know why I am finding the data unless the amount that sublimes is miniscule.

AFAIK garden sulfur is still available but I am thinking of the future and this forum has been illustrative about how bad things are getting for home chemists outside the US especially in Australia and Europe. It seems the US is about 5 or years behind but we will catch up soon enough. Might next a few pounds of garden sulfur as growing season ends in September. Also would like to add that in desperation, very fine sulfur can be obtained from thiosulfate solutions with a little HCl added. Thiosulfate complexes many heavy metals forming complexes that are unstable to heat. They easily can decompose to there respective sulfides. This might of interest in thermite chemistry especially since one does not need to roast carbonates, hydroxides, or hydrated oxides to obtain a suitable oxidant for Al or Mg. Just filter and dry. Garden sufur needs to be dissolve by boiling toluene then crystallized. I can still order pure sulfur online but I should do the crystallization route just to get used to it. If they take away sulfur from online sources garden sulfur will soon follow.

[Edited on 8/19/2008 by chloric1]

[Edited on 8/19/2008 by chloric1]




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[*] posted on 20-8-2008 at 07:41


I'm not convinced that the US will follow suit on the restriction of chemicals for private usage: in Europe, but in the UK in particular, these restrictions have gradually grown and grown over the past decades, even before terrorism. In the US, AFAIK, there is more emphasis on personal freedom, at least in that respect. Here in the UK, knee-jerking is the standard reaction to even the minutest (or potential) problem. For those unfortunate enough to be interested in pyrotechnics (for instance) life has become almost impossible (and all this to try and prevent a few fools from blowing off their faces): in UK forums on pyro, legal matters have now become standard debating points. Sad but true...
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