Exotic thermites & analogs

Quote:

Originally posted by Mardec xD I hate when something doesn't work... I made TH-3 thermite today, the same shit the military uses. I lit perfectly like it should. Burned very intense. Left a big pile of molten metal (Iron probably). But it didn't get through a small (empty) gas container for my bunsen burner.. The iron stayed on top... I don't get it, why didn't it go through it? Everything in the facinaty was charred. I used this comp: -12 g Fe2O3 -4 g Al -6 g Ba(NO3)2 -0,2 g sulpher. I lit it with 50/50 kno3/mg and the whole charge was packed in a 1,5 cm ID carbboard tube with the burning side directly aimed against the bottom of the gas container. And it didn't even made a small hole. The container was emty yes, I am not a noob/stupid. It was one like this: http://content.answers.com/main/content/wp/en-commons/thumb/... What are these things?! there was an aluminum can holding the tube in place, that completly vanished. [Edited on 19-9-2007 by Mardec]

Thermite issues

Nothing like a good physics class on sat. lol.

I offered to show my physics teacher what thermite can do this up coming sat. I was looking through this book that i bought about all sorts of chemicals and how to make them. Now i have Al And FeO2 but they are both 300 mesh, and the book says that i need 2 g of Iron oxide for every 1 g of aluminum powder that are 400 mesh? will that change the density enough to make it so that i will have a different reachtion???

and how much should i make?? enought to scare the CRAP! out of him?!?!?! lol

HgO..

So whos gonna try a HgO thermite?

Useful thermodynamic data on thermites and intermetallics

NOt to mention iodine pentoxide is a REAL oxidizer compared to the other oxides so I would shy away from this too. Unless it made up less than 20% of a composition and you increase reaction rate and get a purple cloud as a bonus!

Manganese thermite using Mn2O3/MnO blends

Making manganese metal by thermite reduction of manganese oxides is one of the more frustrating oxide reductions I've ever carried out.

To get a better understanding of what the issues are I'd firstly recommended reading my blog post on this subject.

The issue with manganese thermite, especially from the higher oxides MnO2 and Mn3O4, is that these reactions generate so much heat that in adiabatic conditions the temperature of the reaction products will exceed the boiling point (BP) of Mn (2,061 C, 2,334 K). It's this characteristic that causes MnO2 thermites to deflagrate often violently, almost 'explosively'. The end-temperature of a thermite reaction (in adiabatic conditions) can be estimated quite accurately from the molar reaction heat (ΔH, reaction enthalpy), and the molar heat capacities and molar heats of fusion of the reaction products, basically by applying [url=http://en.wikipedia.org/wiki/Hess's_Law]Hess's Law[/url]. If of interest to readers I can give an example of such a calculation (on request).

The obvious solution to the problem would be to create conditions in which the reaction enthalpy is either lower or partly dissipates away (non-adiabatic conditions) but there's another problem complicating this approach.

For a thermite reaction to yield liquid metal, separated from the liquid slag (both solidify on cooling of course), the melting point (MP) of both (whichever is highest) the produced metal and the by-product alumina (Al2O3) has to be reached at the end of the reaction. For alumina the MP is 2,327 K (2,054 C), for Mn, 1519 K. But as stated before, the BP of Mn is also only 2,061 C, perilously close to the MP of alumina.

This creates a real lose-lose situation: to achieve metal-slag separation of Mn/Al2O3 this mixture has to reach a temperature that's really close to the BP of manganese metal. Even higher temperatures will cause much of the Mn metal to simply boil off. At temperatures somewhat below the MP of alumina, the vapour pressure of the Mn will be less but metal/slag separation will not be able to occur, resulting in powdered, sintered metal frozen in the slag.

Accurate control of the end-temperature in near-adiabatic conditions is therefore essential to obtain any lump metal from such a reduction.

I therefore started out on a series of experiments designed to cool the MnO2 thermite by co-reducing MnO2 and MnO. The reaction enthalpies per mol of oxide for both reductions are respectively - 597 kJ per mol of MnO2 and - 173 kJ per mol of MnO. Initial results with a blend of 1 mol MnO and 0.4 mol MnO2 and a high level of CaF2 (calcium fluoride, fluorite, fluorspar) as a heat sink and slag fluidiser, showed that this kind of mix with a stoichiometric amount of Al powder is capable of making nice blobs of clean Mn metal, albeit at low yields.

In the mean time I've switched from MnO2 to Mn2O3 because thermochemical calculations show that an Mn2O3 runs a little cooler than the corresponding MnO2 reaction, mainly because the Mn2O3 reaction generates more moles of reaction products (per mol of Mn2O3, 2 mol of Mn and 1 mol Al2O3, against 1 mol Mn and 2/3 mol Al2O3 for MnO2).

I've run several small (20 g and 50 g batches) thermites using blends of Mn2O3 and MnO, always using a high level of CaF2. I set the level of CaF2 as a constant molar ratio of CaF2/Al = 0.225, the alumina slag therefore contains always the same molar fraction of CaF2.

On the whole the results indicate that obtaining yields (recovered metal/metal present in the oxide x 100 %) is hard to push much beyond 35 % or so. The last two batches, both with 50 g stoichiometric (and CaF2 = 0.225 molar ratio) mixes, gave the following yields:


...........................mol....................mol

Mn2O3................1.........................1
MnO....................1.........................0

Yield....................37 %...................19 %

The 1/0.5 blend gave the highest yield of Mn metal I've ever achieved (out of probably over 20 or so reactions) and the metal is clean skinned and solid. The largest regulus was 7.3 g. Both reactions ran well-contained, leaving a molten slag/metal puddle at the bottom of the crucible. The 1/0 batch yielded metal that was significantly more oxidised, yet on the whole of passable quality.

It's clear though although these results constitute a great improvement to the usual 'explosive' MnO2 thermite, there is only so much cooling the Mn2O3 thermite by blending it with the much cooler MnO can actually achieve in terms of yield improvement.

The only real solution to reducing Mn oxides (or halides) with higher yields is by using a reductant with a much lower melting oxide (or halide).

One such reaction that springs to mind is the reduction of anhydrous MnCl2 with Mg. The reaction enthalpy of MnCl2 + Mg ---> Mn + MgCl2 is unfortunately only a measly - 161 kJ/mol of MnCl2, about a 100 kJ short of success. Thermocalcs show that in adiabatic conditions the reaction products would be heated to about 1,200 K, well above the MP of MgCl2 (987 K) but about 300 K short of the MP of Mn (1,519 K). Pre-heating the mixture by about 300 K or simply heating it to spontaneous ignition could work to obtain liquid Mn and liquid MgCl2.

As regards using a reductant with an oxide of lower MP, that excludes both Mg and Ca, as both have oxides with insanely high MPs. That then really only leaves the alkali metals, in particular Li and Na.

Thermochemical calculation for the reaction Mn2O3 + 6 Li ---> 2 Mn + 3 Li2O (ΔH = - 838 kJ per mol Mn2O3) shows that in adiabatic conditions the estimated end-temperature would be 2,320 K which is still too high and too close to the BP of Mn.

But here we could blend with MnO again. Setting a target end-temperature of 2,000 K (well above the MP of Mn, yet well below its BP as well), the blend composition of a stoichiometric mix has been estimated to be about 1 mol Mn2O3 + 2.6 mol MnO.

One small problem: I haven't got any Li...

Sodium, with a ΔH = - 295 kJ per mol Mn2O3 for Mn2O3 + 6 Na ---> 2 Mn + 3 Na2O could also be a good candidate... Haven't got any Na either...

[Edited on 23-7-2008 by blogfast25]

[Edited on 23-7-2008 by blogfast25]

[Edited on 24-7-2008 by blogfast25]

You know, it's never occurred to me to ask my hardware store for WHITE-HOT BOILING stainless steel. I'll ring them right up and ask.

I'm not trying to develop a production process, just an interesting demonstration. Shopping at the hardware store can be interesting, of course, but not in the same way.

Quote:

Originally posted by -jeffB You know, it's never occurred to me to ask my hardware store for WHITE-HOT BOILING stainless steel. I'll ring them right up and ask.

LMAO

You put it like that, after leaving my neighborhood hardware store with my white hot stainless, I'll head off to the local pub and have myself a molotov cocktail.

Cr2O3 + Mg

Ouch! Would be nice to have a regulas of chromium

Actually, I am sticking to using aluminum here and I wanted to use potassium dichromate as a booster. I have both 300 mesh and 400 mesh Al, so I think 15% dichromate is all that should be needed. I worked out the dichromate reduction in thiis equation:

K2Cr2O7 + 2Al >2KAlO2 + Cr2O3
Cr2O3 +2Al > 2Cr +Al2O3
NET: K2Cr2O7 + 4Al >2KAlO2 +2Cr +2Al2O3

Yep, that seems right.

Quote:

Originally posted by chloric1 Ouch! Would be nice to have a regulas of chromium Actually, I am sticking to using aluminum here and I wanted to use potassium dichromate as a booster. I have both 300 mesh and 400 mesh Al, so I think 15% dichromate is all that should be needed. I worked out the dichromate reduction in thiis equation: K2Cr2O7 + 2Al >2KAlO2 + Cr2O3 Cr2O3 +2Al > 2Cr +Al2O3 NET: K2Cr2O7 + 4Al >2KAlO2 +2Cr +2Al2O3 Yep, that seems right.

Chloric1:

Firstly, the end-temperature to which a thermite reaction runs depends on various factors, the thermochemistry itself being a prime factor. Thermites based on TiO2, SiO2 and some other very, very stable oxides either don't burn at all or burn slowly and not very hot. This is certainly also true of a few Period 5 and Period 6 transition metals, where reduction of halides is usually preferred.

The Gibbs Free Energy released during the reaction (ΔG = ΔH - TΔS) indicates whether the reaction can take place or not and the reaction enthalpy (heat) ΔH, together with how much reaction products are formed, their heat capacities and heats of fusion (if applicable), as well as to what extent the reaction vessel is thermally insulated from the rest of the world, determines the end-temperature precisely and unequivocally.

Regards the Chromium thermite, remember that to obtain good quality metal reguli, usually fluorite is needed to obtain good metal/slag separation (but it isn't always indispensable). Most of my thermites, designed to produce good quality metal, contain fluorite (CaF₂. Here's my Cr formula (this is a stoichiometrically balanced mix):

in mol: Cr₂O₃ = 1 mol; Al = 2.82 mol; KClO₃ = 0.41 mol, CaF₂ = 0.6345 mol

in gram: Cr₂O₃ = 152 g; Al = 75.8 g; KClO₃ = 50.2 g, CaF₂ = 49.5 g

The fluorite level could be reduced or it could be eliminated altogether but bear in mind that fluorite is a slag fluidiser as well as a heat sink: without fluorite the reaction will run hotter (and that's not necessarily an advantage, if metal production and not pyrotechnics is your goal).

Let me know how you get on...

[Edit]

Also, the reaction between K2Cr2O7 and Al is likely to run to elemental K, not KAlO2, because K2O + 2/3 Al ---> 2 K + 1/3 Al2O3 is exothermic by -195 kJ/mol of K2O.

The overall booster reactions can then be balanced as:

K2Cr2O7 ---> K2O + 2 CrO3
K2O + 2/3 Al ---> 2 K + 1/3 Al2O3
2 { CrO3 +2 Al ---> Cr + Al2O3 }

K2Cr2O7 + 14/3 Al ---> 2 Cr + 2 K + 7/3 Al2O3

The third part is the one likely to produce enormous amounts of heat.

If you want to use 240 g (1.58 mol) of Cr2O3 and 36 g (0.122 mol) of K2Cr2O7, the amount of Al should then be increased slightly to about 107.4 g to respect stoichiometry.

The level of heat booster in your dichromate boosted formulation is quite low: 0.077 mol of K2Cr2O7 per mol of Cr2O3. For my chlorate based formulation it is 0.41 mol of chlorate per mol of Cr2O3.

But since as I have no thermochemistry data for K2Cr2O7 or Cr (+VI) oxide at hand, it's difficult to say whether the chromate boost will be enough to obtain molten slag and metal. The advantage of using dichromate is of course that it also produces Cr and not just alumina and KCl. I think it's the safest way for now to test the low level of dichromate first.

[Edited on 3-8-2008 by blogfast25]

Thanks blogfast! I appreciate that and I am taking notes. I got the dichromate idea from a couple of older digitized inorganic chemistry books. One uses Chromium trioxide and the more modern one uses potassium dichromate! The hexavalent chrome makes up significant portions simular to your chlorate formula but they are using 30 mesh Al!! So I toned it down to accommodate for finer powders.

OUCH! I did not factor in the fact Potassium metal is formed. Certain that it will be a purple vapor flame!
I really need to get a grasp of theorical thermochemistry. I am planning on studying in the University soon.

I do not have fluorspar at this moment but plan to get some. For this, and the fact I have spent the weekend entertaining guest I have yet to do my trial. I will let you know when I get around to it(possibly Monday or Tuesday evening). My next big order will include some fluxes so I can
compare the run with/without flux.

I would think slightly less fluorspar could be used because the potassium vapor should lower temperature somewhat but how much I don't know.

[Edited on 8/3/2008 by chloric1]

Chloric1:

Thermochemistry of thermites (and assorted reactions) isn't very complicated at all: it requires only basic algebra and invoking a couple of laws.

Here's an example of a thermochemical calculation for a boosted TiO2 thermite (the post is mine)

Another one on ferrotitanium, as well as a general calculation of the estimated end-temperature for a generic thermite (the post is mine too)

The cooling effect of the potassium will be small but could be accurately estimated using the heat capacities (Cp,l and Cp,s), heat of fusion and heat of evaporation of K. In your proposed formulation the cooling effect is likely to be negligible because the amount of K formed is small.

But in the KClO3 boosted formulation, where 0.41 mol of KCl are formed, heated, fused and evaporated (following Hess), the cooling effect has to be taken into account even though compared to the heat generated by KClO3 + 2 Al ---> KCl + Al2O3 is enormous. It all depends how accurate you want the calculation to, of course...

[Edited on 4-8-2008 by blogfast25]

[Edited on 4-8-2008 by blogfast25]

[Edited on 4-8-2008 by blogfast25]

Well the more I thought about superheated potassium vapor, the more I realized how little of an effect it will have. Not only that but I am considering the chemistry here. You will have molten Aluminum Oxide not to mention heated air surrounding reaction. So any losses of potassium would result in immediate combustion releasing white clouds of potassium oxide(don't breath),or immediate combination of said oxide with the molten alumina hence my original deduction of potassium aluminate as a product intermixed in a matrix of the alumina slag. I appreciate the post on thermochemistry of thermites. With your permission, I would like to copy them and save them as a Word doc.

Well, it definately ignited I will post a link to video when I get a chance to upload it. I use a flower pot half fill with sand. Truthfully, I was worried the pot would fly apart but it just made a "ping" when after the permanganate/glycerol hypergolic subsided and the thermite ignited. I really am getting tired of this ignition system because the smoke obscures my view and it smells like burnt sugar. Anyhow on to the mix. It gave off green smoke Obviously some chrome oxide was mechanically lifted by the vigor of the reaction. The slag/mess looks porous so I am not too optimistic. I will break her open around 9 PM GMT-5 to look for metal.

Since I prepared like 383 grams of thermite I still have 70-80% left so we can optimize together then compare to chlorate bosster and I will order fluorspar this week.


Update: Just uploaded my videohere

I did not find a distinctive globule or regulas. More like a broccoli sprout It is very hard but very brittle. I think it is chromium in a alumina matrix. I also seen some unreacted chrome oxide so I think it was not hot enough. I will post more findings later.



[Edited on 8/5/2008 by chloric1]

Well I have magnesium ribbon coming this wek I hope. I ordered 225 feet! I plan on selling most of it to finance research Next I do this I will measure out a 100 gram charge. I will figure out the dichromate addition based on your numbers of .15 tp .2 moles. Today I have to decide who to order fluorspar from. Not all pottery suppliers are created equal. I also need to order black iron oxide and possibly something else. This all helps to justify the shipping costs.

Why not get some vanadium pentoxide (pottery grade)? That's a real nice thermite, very hot and nice metal too. Cobalt's nice too, from CoO... The one I'm most interested in is niobium pentoxide: no boost needed, burns straight to about 2740 K! But I can't get any Nb2O5 here...

Instead of CaF2, cryolite (Na3AlF6) should also do the trick but it's more expensive of course.

I'll probably be moving away from oxides and move to halides (mostly chlorides). I've a MnCl2 + Mg ---> Mn + MgCl2 on the drawing board.

Thats just it, I got numerous suppliers for CaF2 and cryolite but I need the supplier with reasonably priced vanadium pentoxide. So I can thermite it and sell some too. I have one supplier that wants $45 per pound!! I have a couple others that range $13-15 per pound! There is definately no set standard on these type of item hence money can easily be made. The niobium pentoxide will probably have to be purchased directly from China or India in 25Kg quantities or more. I am definately unable to swing that as of now

Why do you want to move away from oxides? Want a new challenge? The problem with halides is that many are VERY soluble and hydroscopic. Many Fluorides are insoluble and would be easy to obtain anhydrous.

Personally, I wanted to make up some lead bromide and make a thermite with that and try to catch the aluminum bromide smoke in a bell jar or?

[Edited on 8/6/2008 by chloric1]

Yes, a new challenge. Most chlorides are indeed hygroscopic but drying them is usually possible. Anhydrous fluorides are hard to make: usually hydrofluoric acid or anhydrous HF and in some cases fluorine is needed. They're also much more expensive and more restricted for sale to private individuals (at least in Europe).

High grade Nb2O5 is available from ChemSavers but needless to say only at ridiculous prices.

The reduction PbBr₂ + 2/3 Al ---> Pb + 2/3 AlBr₃ is only just barely exothermic: ΔH = - 64 kJ/mol of PbBr2 (at 298 K) (for chlorides and bromides, Mg is much preferred to Al but for fluorides Al also works. This is mainly because of the fact that in the series of AlX3, AlF3 has by far the highest heat of formation).

I'm not sure (but it can be calculated) whether the temperature would be high enough to evaporate the AlBr3 and in any case it would be heavily contaminated with lead, due to mechanical entrainment of the nascent lead. AlBr3 must also be very prone to hydrolysis. Like for AlCl3, reaction of Al with dry HBr in a glass labware set up with condenser is to be much preferred here.

I just took receipt of my Mg powder, so the MnCl2 reduction project can now go ahead.

Na₃AlF₆ 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 AlBr₃?

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

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]

Quote:

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

amazingrust is indeed your site! 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.

First reduction of MnCl2 with Mg powder

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 ).

Dichromate boosted chromium thermite

Dichromate Boosted Thermite

Chloric-

Yes, all these formulations are fluxed. I set the CaF₂ to a level that is equal to (in mol) CaF₂ = 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

Cr₂O₃ ................................................. 1
K₂Cr₂O₇ ............................................. x
Al ........................................................ 2 + (14/3) x
CaF₂ .................................................. 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 CaF₂...

-----------

A test trying to use Mg + I₂ ---> MgI₂ 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 KMnO₄ 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 KMnO₄ + 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]

@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]

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 KClO₃.

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 CaF₂ + 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 AlF₃ + 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' MnO₂ (wild horses couldn't drag me away!) according MnO₂ + 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: MnCl₂ + Mg ---> Mn + MgCl₂, ΔH_{298 K} = - 161 kJ/mol (of MnCl₂

Booster reaction: MnO₂ + 2 Mg ---> Mn + 2 MgO, ΔH_{298 K} = - 684 kJ/mol (of MnO₂

I set the target end-temperature at 1,800 K, well above the MPs of both Mn and MnCl₂, but well below the BP of Mn (and below the MP of MgO).

Now we need to find out how many mol of MnO₂ is needed per mol of MnCl₂ 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 MgCl₂ 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_Mn = 58.75 kJ/mol, for MgCl₂ ΔH_MgCl2 = 133 kJ/mol, for MgO ΔH_MgO = 74.4 kJ/mol.

The total heat of reaction ΔH_R = - 161 - 684 x and in adiabatic conditions:

ΔH_R + (1 + x) ΔH_Mn + ΔH_MgCl2 + 2x ΔH_MgO = 0

Or:

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

and x = 0.06 mol of MnO₂.

Since as we're not in adiabatic conditions and some heat losses are inevitable, a good starting point formulation (in mol) would be MnCl₂ = 1; MnO₂ = 0.1; Mg = 1.2

[Edited on 16-8-2008 by blogfast25]

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

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

[Edited on 8/16/2008 by chloric1]

Chloric-

I'm pretty convinced the smoke you see is K₂O/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^m+ and O^2- form a crystal lattice M₂O_m 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^m+ is much higher than for M⁺.

Ca(NO₃₂ is super deliquescent, so count that one out.

MoS₂/Al is a very interesting proposition. The heat of reaction ΔH_R for MoS₂ + 4/3 Al ---> Mo + 2/3 Al₂S₃ 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_p or ΔH_fusion for The Smelly One, so a little, simple experimentation would be needed. Caution not to get H₂S 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]

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!

________________________________________

Update! Bonus video!

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]

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 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

Thanks, DerAlte, I've lived in several European countries and they're all quite Nanny state-ish when it comes to "better safe than sorry", but this one (UK) really does take the biscuit...

Anyroads, continuing with my "highly dangerous project" of trying to obtain lump metal from the reduction of anhydrous MnCl₂ with powdered magnesium, I ran another 20 g test today, using MnO2 as a heat-booster.

The molar formulation was: MnCl₂ = 1; MnO₂ = 0.2, Mg = 1.4

In terms of burn rate and heat this is starting to resemble what a successful thermite (this is more a hybrid, of course) looks like: very fast and regular, much more so than both previous attempts. Slag/metal had flowed more or less to the bottom of the crucible (an eggcup) but not completely and it was irregularly shaped and somewhat porous. Still, considerable fusion took place but no metal reguli were found.

Crushing the slag/metal mix, glistening metal could be seen in it and the crushed slag reacts with water as before: strong release of heat and quite some gas. Adding 32 w% HCl to the slag/water slurry results in a strong reaction and lots of hydrogen being released. The dissolved manganese formed will be recovered.

The reaction also produced lots of smoke. I'm wondering if residual NH4Cl is the (partly) cause. In previous reactions a lingering smell of ammonia could be observed (but not this time). The starting material anhydrous MnCl₂ also tests positive for NH₄⁺ (add alkali and the smell of NH₃ is unmistakable). There is thus certainly residual salmiac present, so I'll have to assay the raw material for MnCl₂ content (gravimetrically via carbonate, I guess). Residual salmiac would also cool the reaction a bit, mainly because of the heat of evaporation...

Weather permitting, another test tomorrow with a higher level of KMnO₄, which I haven't ruled out as a heat-booster yet...

Yes. Nicely put.

But simpler:

MnCl₂ + n KMnO₄ + (1 + 4n - n/2) Mg ---> (1 + n) Mn + n KCl + 4n MgO + (1 - n/2) MgCl₂ (n >= 0)

Similarly for the thermite of Cr₂O₃, boosted with K₂Cr₂O₇ because the K from K2Cr2O7 + 14/3 Al --->2 K + 2 Cr + 7/3 Al2O3 can also react as Cr2O3 + 6 K ---> 2 Cr + 3 K2O (the latter is highly exothermic). Here too an adjustment of the stoichiometry is needed to account for the reactivity of the K:

Cr₂O₃ + n K₂Cr₂O₇ + 2 (1 + 2n) Al ---> 2 (1 + n) Cr + (1 + 2n) Al₂O₃ + n K₂O (n >= 0)

Aaaah... fun with stoichiometry!

Work commitments and bad weather have so far prevented me from carrying out more tests with MnCl₂, despite having a new batch ready. But I found time for a little theorising, in particular in an attempt to find other anhydrous chlorides suitable for reduction by magnesium, <i>in open reactor conditions</i> (and not bomb reactor conditions).

One drawback of Mg is that MgCl₂ for open crucible reactions is its relatively low boiling point (BP) of 1,685 K (1,412 C). It would therefore be impossible to obtain lump metal without the slag boiling (or "exploding") off if the melting point of the metal pursued was higher than 1,685 K, (unless in bomb conditions).

But even for metals with lower MPs there remains the possibility of a very energetic reaction, exothermic enough to heat the MgCl₂ well past its boiling point.

Take the case of a generic chloride, MCl_n with a standard (298 K) heat of formation (HoF) of ΔH_{f, MCln}. The standard reaction enthalpy ΔH_R for the reaction MCl_n + n/2 Mg ---> M + n/2 MgCl₂ is:

ΔH_R = - ΔH_{f, MCln} + n/2 ΔH_{f, MgCl2}

with ΔH_{f, MgCl2} = -642 kJ

The BP of MgCl₂ is 1,685 K and the enthalpy to heat 1 mol of M and n/2 mol of MgCl₂ to that temperature is ΔH_M + n/2 . ΔH_MgCl2, which from the relevant NIST Shomate equations can be calculated to be resp. ≈ 50 kJ (approx. for many metals) and n/2 . 122 kJ.

In adiabatic conditions: <i>∑ΔH = 0</i>, so:

ΔH_R + ΔH_M + n/2 ΔH_MgCl2 = 0

or: - ΔH_{f, MCln} + n/2 ΔH_{f, MgCl2} + ΔH_M + n/2 ΔH_MgCl2 = 0

Reworked, the standard HoF of MCl_n should be smaller than:

ΔH_{f, MCln} < ΔH_M + n/2 (ΔH_{f, MgCl2} + ΔH_MgCl2

or ΔH_{f, MCln} < 50 - n/2 . 526 (kJ/mol)

as otherwise in open reactor (and adiabatic) conditions MgCl₂ will be heated past its boiling point.

With a HoF of - 481 kJ, it becomes immediately clear that MnCl₂ satisfies this condition.

Alas, few relatively easily accessible chlorides do meet it, in fact I've yet to identify a single one!

Here's a few that don't meet the criterion:

CuCl2: HoF = - 206 kJ/mol, n = 2
ZnCl2: HoF = - 415 kJ/mol, n = 2
PbCl2: HoF = - 359 kJ/mol, n = 2
FeCl2: HoF = - 342 kJ/mol, n = 2
FeCl3: HoF = - 399 kJ/mol, n = 3
AlCl3: HoF = - 706 kJ/mol, n = 3
CoCl2: HoF = - 313 kJ/mol, n = 2
SnCl2: HoF = - 333 kJ/mol, n = 2

ZnCl₂ would probably be the best choice, provided th reaction can be cooled a little with a heat sink.

Very interesting [Ca aluminate being soluble in HCl].

The only problems I can see is that to obtain Ca aluminate only you need quite a bit of this inert (and heat sink) substance (CaO + Al2O3 --> Ca(AlO2)2, assuming this is the simple structure of Ca aluminate) and that this will have to be factored into the thermo calcs (end temperature!) The heat to carry the CaO to the MP of the metal (or slag, whichever is highest) will be very considerable and cannot simple be ignored...

Also, high levels of such (basically half a mol of lime per mol of Al) an inert substance may well mess with the kinetics, may make ignition harder and such like.

But for Si this is something I need to try, absolutely, I may in fact calc it tonight.

For V, the reaction with just Al (no CaO) works extremely well and leads to excellent metal coalescence, no need to remove slag chemically (see one spectacular V2O5 thermite over at AmazingRust.com, results very, very close to my own).

For Mo it's also an interesting proposition.

For W, it could be a way to produce the powdered metal, as the MP is far too high to obtain lump metal.

One could make some Ca aluminate (to study its properties) by igniting a mix of Al, KClO3 and CaO: overall reaction KClO₃ + 2 Al + CaO ---> Ca(AlO₂₂ + KCl