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Author: Subject: Exotic thermites & analogs
DerAlte
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[*] posted on 20-8-2008 at 21:54


@blogfast25

A bit OT but...

I ought to have guessed you were from the UK. Things have indeed deteriorated badly. When I was a lad of 16-20, (1953-1957) I could get anything I wanted bar such things as arsenic and antimony oxides, mercuric chloride (but not mercurous, nor mercury metal), cyanide, and other compounds on the Poison Register. I used to get them from the local chemist (=pharmacist in UK) - even oleum. My father got the poisons by merely signing 'for experimental purposes'. Al powder. Mg powder, Na, K metal, chlorate, white P, red P, you name it. I had them all. Also any chemical glassware under the sun.

Yes, I was into pyro technics, too. My best ( worst?) effort was one Guy Fawkes night, when I set off a permangante, sulphur and Mg flash powder about 100 g worth in a cardboard tube reinforced by copper wire. It rattled the windows. Ibid, I tried same on the common in a forked tree and blew one fork off. In misty weather, the flash was spectacular!

I managed to gas myself with chlorine (2 days of lung ache), Bromine (worse!) and the carcinogenic chromyl chloride, to say nothing of phosphine, H2S et al. Happy days, denied to today's youth of all ages by the Nanny State, PC bullshit, and idiotic Greens (but not true environmentalists, of which there are few).

Enough, this thread is no the place to rant. Keep up the good thermite work, you and Chloric...

Best regards,

Der Alte
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[*] posted on 22-8-2008 at 09:43


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 MnCl<sub>2</sub> with powdered magnesium, I ran another 20 g test today, using MnO2 as a heat-booster.

The molar formulation was: MnCl<sub>2</sub> = 1; MnO<sub>2</sub> = 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 MnCl<sub>2</sub> also tests positive for NH<sub>4</sub><sup>+</sup> (add alkali and the smell of NH<sub>3</sub> is unmistakable). There is thus certainly residual salmiac present, so I'll have to assay the raw material for MnCl<sub>2</sub> 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 KMnO<sub>4</sub>, which I haven't ruled out as a heat-booster yet... :)
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[*] posted on 22-8-2008 at 10:06


How hot can you take MnCl2 before it melts or evaporates to nothingness? Most transition metal chlorides are rather volatile (for that matter, I wouldn't be surprised if a noticable amount co-sublimates with the NH4Cl, something to think about I guess). MnCl2 is an odd one out, perhaps resembling MgCl2 or CaCl2 more than others.

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[*] posted on 22-8-2008 at 10:42


Hi Tim,

MnCl2: MP = 654 C, BP = 1,225 C

This reaction [reduction of the chloride with magnesium] was used by E.Glatzel in 1889 to prepare the metal. I don't have the original article and the precise conditions unfortunately. But I reckon similar successful reductions have been carried out with more volatile oxidants. I'm aiming for yields of > 90 % of lump Mn metal. I'd prefer to work with the fluoride but can't for obvious reasons.

Also, since as I'm still below the MP of Mn (1246 C) at this point, significant blow-off of the chloride doesn't sound plausible...

Generally taking stock of the MPs and BPs of all reagents and all reaction products is of course good practice but perhaps we can push that too far too. Take the example of niobium thermite, the only method in use for industrial production of that metal (albeit followed by refining, of course): the BP of Al is 2,792 K, the MP of Nb 2,477K. To allow the metal to separate from the slag the reaction mix therefore needs to heat well beyond 2,477 K, in all likelihood beyond the BP of Al, yet Al boil-off isn't reported. It's to be assumed that the reduction reaction starts taking place at much, much lower temperatures...

As regards co-sublimation, the recorded weight losses from 4 or 5 fumings do not indicate any loss of MnCl2: if anything the weight losses seem too low compared to theo expectations. I might need to drive up temperature (or simply dry in a stream of dry HCl).

[Edited on 22-8-2008 by blogfast25]

[Edited on 23-8-2008 by blogfast25]
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[*] posted on 24-8-2008 at 08:10


It's also occurred to me that if I assume KMnO4 + 4 Mg ---> K + Mn + 4 MgO to proceed, then MnCl2 + 2 K ---> Mn + 2 KCl has also to be taken into account, as it's highly exothermic too...
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[*] posted on 24-8-2008 at 10:44


Ah yes, so in total you might have something like...ahhh that would be...
2 KMnO4 + (n+8) Mg + (n+1) MnCl2 = (n+3) Mn + 2 KCl + 8 MgO + n MgCl2 (for n >= 0)
Tally up the HOF for those, in terms of the proportion n, and you should be sitting pretty accurately...

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[*] posted on 25-8-2008 at 07:13


Yes. Nicely put.

But simpler:

MnCl<sub>2</sub> + n KMnO<sub>4</sub> + (1 + 4n - n/2) Mg ---> (1 + n) Mn + n KCl + 4n MgO + (1 - n/2) MgCl<sub>2</sub> (n >= 0)

Similarly for the thermite of Cr<sub>2</sub>O<sub>3</sub>, boosted with K<sub>2</sub>Cr<sub>2</sub>O<sub>7</sub> 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:

Cr<sub>2</sub>O<sub>3</sub> + n K<sub>2</sub>Cr<sub>2</sub>O<sub>7</sub> + 2 (1 + 2n) Al ---> 2 (1 + n) Cr + (1 + 2n) Al<sub>2</sub>O<sub>3</sub> + n K<sub>2</sub>O (n >= 0)

Aaaah... fun with stoichiometry! :)
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[*] posted on 9-9-2008 at 11:31


Work commitments and bad weather have so far prevented me from carrying out more tests with MnCl<sub>2</sub>, 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 MgCl<sub>2</sub> 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 MgCl<sub>2</sub> well past its boiling point.

Take the case of a generic chloride, MCl<sub>n</sub> with a standard (298 K) heat of formation (HoF) of ΔH<sub>f, MCln</sub>. The standard reaction enthalpy ΔH<sub>R</sub> for the reaction MCl<sub>n</sub> + n/2 Mg ---> M + n/2 MgCl<sub>2</sub> is:

ΔH<sub>R</sub> = - ΔH<sub>f, MCln</sub> + n/2 ΔH<sub>f, MgCl2</sub>

with ΔH<sub>f, MgCl2</sub> = -642 kJ

The BP of MgCl<sub>2</sub> is 1,685 K and the enthalpy to heat 1 mol of M and n/2 mol of MgCl<sub>2</sub> to that temperature is ΔH<sub>M</sub> + n/2 . ΔH<sub>MgCl2</sub>, 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<sub>R</sub> + ΔH<sub>M</sub> + n/2 ΔH<sub>MgCl2</sub> = 0

or: - ΔH<sub>f, MCln</sub> + n/2 ΔH<sub>f, MgCl2</sub> + ΔH<sub>M</sub> + n/2 ΔH<sub>MgCl2</sub> = 0

Reworked, the standard HoF of MCl<sub>n</sub> should be smaller than:

ΔH<sub>f, MCln</sub> < ΔH<sub>M</sub> + n/2 (ΔH<sub>f, MgCl2</sub> + ΔH<sub>MgCl2</sub>;)

or ΔH<sub>f, MCln</sub> < 50 - n/2 . 526 (kJ/mol)

as otherwise in open reactor (and adiabatic) conditions MgCl<sub>2</sub> will be heated past its boiling point.

With a HoF of - 481 kJ, it becomes immediately clear that MnCl<sub>2</sub> 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

ZnCl<sub>2</sub> would probably be the best choice, provided th reaction can be cooled a little with a heat sink.
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[*] posted on 17-9-2008 at 11:32


Finally another couple of tests designed to increase end-temperature of the MnCl2/Mg reduction.

One test was a 20 g batch of MnCl2/Mg, w/o any booster but pre-heated to 225 C for about 45 min in a gas fired oven. That burned remarkably fast but the slag was porous and no lump metal was found.

Secondly, a 20 g test with a high level of MnO2 booster. MnCl2 = 1 mol, MnO2 = 0.4 mol, Mg = 1.8 mol. This 20.0 g batch and (steel) crucible were accurately weighed before and after reaction.

The reaction was very fast and hot (comparably to a 'good' thermite) and VERY smoky. Weight loss during reaction was about 18 w%. The reaction products are 1.4 mol Mn, 1 mol MgCl2 and 0.8 mol MgO, or about 46.5 w% of MgCl2. It appears roughly half of the latter was blown off, presumably because MgCl2 has a relatively low BP (1,412 C) and thus a high vapour pressure at these temperatures. Evaporating MgCl2 would also cool the reaction because of the used latent heat of evaporation. (for KCl it's 79.1 kJ/mol but I haven't got a value for MgCl2).

The obtained slag/metal was extremely porous and could be crumbled by hand. Needless to say no lump metal was found.

I'm now convinced the volatility of the slag is the problem here, but I need first to eliminate with certainty volatility from residual ammonium chloride, the drying-aid.

If I'm right then obtaining lump metal from such a reaction would require a closed, pressure-proof reactor, to avoid the slag from volatilising. I'm thinking of a defunct RC plane engine, 10 - 20 cc. Or a lawnmower motor, something like that. With a sparkplug it's possible to engineer RC ignitions.

Ironically, the MnCl2/Mg reduction was meant to cure the problem of high Mn volatility in MnO2 thermites, while the remedy seems to present the "opposite" problem: evaporating slag!
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[*] posted on 20-9-2008 at 08:13


And as it happens, I've just found the abstract of Glatzel's original paper on the reduction of manganous chloride with magnesium:

Preparation of Manganese from Manganese Chloride and
Magnesium.

By E. GLATZEL (Ber., 22, 2857-2859).

-Manganese can be prepared by heating a mixture of finely divided, anhydrous
manganese chloride (100 grams) and dry, powdered potassium
chloride (200 grams) in a covered Hessian crucible until it just
melts, and then adding magnesium (15 grams) in portions of
3-4 grams, at intervals of 2-3 minutes; if the fused mass is too
hot a very violent reaction occurs, and the contents of the crucible
are thrown out. The crucible is covered again, heated more strongly,
and then allowed to cool slowly in the furnace. The yield of manganese
is 20-25 grams, the metal containing traces only of silica,
and being quite free from magnesium.
The specific gravity of manganese, as the average of four determinations,
was found to be 7.3921 at 22". F. S. K.

So Glatzel basically heated molten MnCl<sub>2</sub> with powdered Mg in the presence of (rather a lot) of molten KCl flux/heat sink.

Without a doubt, the initially formed manganese is powdered metal and the second heating phase is designed to take the mixture of slag, possibly unreacted MnCl<sub>2</sub> and KCl to above the MP of manganese (1,246 C) and allow the metal to separate out.

Also, the high molar ratio of MnCl<sub>2</sub> to Mg (0.8 / 0.57) would help explain the near-absence of Mg in the obtained Mn. Theoretical yield: 31 g of Mn.

[Edited on 20-9-2008 by blogfast25]
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[*] posted on 20-9-2008 at 10:35


Have yet to try the calcium oxide flux for my chrome thermite but recently read that calcium aluminates can melt as low as 1300 Celcius! After I originally posted this idea, it dawned on me that adding CaO(an alkali) to the dichromate(acidic) boosted thermite might cause side reactions before ignition. Of coarse siad calcium chromates should still yield appropiate calcium aluminate flux.

I can figure out the specific heat of calcium oxide and how much dichromate booster I will need to counter the heat sinking effect. Question is, do I really need stochiometric amounts of CaO to form the 1:1 aluminate or can I use substantially less to just "fluidize" the slag?




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[*] posted on 21-9-2008 at 08:24


Chloric1:

Use substantially lower quantities of CaO: for a1:1 ratio of Al2O3/CaO the amount needed of the latter would be impractically high. If the Ca aluminate (or parts of it) has indeed such a low MP, it will act as a slag-fluidiser, even at modest doses.

Personally I've never tried lime as an additive in thermites but many literature references to that use exist. It's recommended for instance for MnO2/Al...
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[*] posted on 21-9-2008 at 10:32


Well, it gets better. Supposedly, from a boron by thermite patent of 1962, calcium aluminate is easily soluble in dilute HCl. This was important considering boron does not separate in another phase. This would be really nice for getting clean vanadium,tungsten, molybdenum or silicon for that matter. If the aluminate can be dissolved with little fuss, it might be a source of CaCO3 and Al(OH)3 for latter.



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thumbup.gif posted on 22-9-2008 at 08:48


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 KClO<sub>3</sub> + 2 Al + CaO ---> Ca(AlO<sub>2</sub>;)<sub>2</sub> + KCl
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[*] posted on 22-9-2008 at 09:20


Quote:
Originally posted by chloric1
calcium aluminate is easily soluble in dilute HCl


Easily soluble, or easily reacts with?

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[*] posted on 23-9-2008 at 04:32


Tim:

There isn't, at first glance, much to be learned from the Tinkerwebs on the properties of calcium mono aluminate, other than that its hydrolysis behaviour appears to be complex. It would sound plausible though that in certain circumstances this aluminate would react with HCl, to produce CaCl2 and AlCl3, analogous to aluminates of Group I. Wait and see, I guess...

For the hypothetical reaction of silica with Al in the presence of lime (overall reaction:

SiO<sub>2</sub> + 4/3 Al + 2/3 CaO ---> Si + 2/3 CaAl<sub>2</sub>O<sub>4</sub>

boosted with:

KClO<sub>3</sub> + 2 Al + CaO ---> CaAl<sub>2</sub>O<sub>4</sub> + KCl)

The overall stoichiometric formulation composition would be:

SiO2 .......................... 1 mol
KClO3 ....................... x mol
Al ............................... (4/3 + 2 x) mol
CaO ........................... (2/3 + x) mol

Applying a few reasonable assumptions a thermocalc shows that in strictly adiabatic conditions at x = 0.2 this mixture would reach approximately 2,500 K (well above the MP of alumina).

I will test this 'shortly' [cough!]...
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[*] posted on 24-9-2008 at 07:26


Interestingly, the lime route may even open up a new possibility of solving the MnO<sub>2</sub> thermite problem of Mn volatility, by cooling the reaction with lime. Provided the calcium mono aluminate can form somewhat <i>below</i> the MP of actual alumina, this would be a way of keeping the reaction temperature lower than the BP of Mn, yet have a nicely fluid slag all the same. Industrially, Mn is produced by thermite reduction of Mn<sub>3</sub>O<sub>4</sub> in the presence of lime. Unfortunately I don't have a pret-a-porter formulation at hand...

It would also be interesting to test the lime in a TiO2 thermite boosted with CaSO<sub>4</sub>, largely by replacing the CaF<sub>2</sub> slag fluidiser with CaO.
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[*] posted on 24-9-2008 at 09:00


I was thinking that the calcium aluminate would lose the calcium to the HCL and leave part of the alumina as a hydrous jelly. Aqueous aluminum chemistry requires alot of free acid. So double theoretical amounts of HCl should do the trick.

The lime trick might work for manganese. What about using calcium chloride as a flux in small amounts?




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[*] posted on 24-9-2008 at 11:46


Quote:
Originally posted by chloric1

The lime trick might work for manganese. What about using calcium chloride as a flux in small amounts?


That's generally speaking far too volatile (BP below 2,000 C) and would be blown off. Could be useful in chloride reductions though...
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[*] posted on 27-9-2008 at 06:18


Well, the two first attempts at using CaO (quicklime)/CaAl2O4 as a slag fluidiser weren't very successful at all.

Both were calcium sulphate boosted formulations and in the first I set the mol of CaO to half the mol of Al, so that the CaO/Al2O3 ratio was exactly 1. The thermite burned well but considerably slower than its CaF2 analog and the slag was very viscous and didn't even make it to the bottom of the crucible. There was no lump metal whatsoever.

Reasoning the high CaO level was perhaps proving too much of a heat sink, in the next test I cut it right down to CaO/Al2O3 = 3/8, less than half the original dose. This one burned faster and hotter, with more (but not all) slag collecting at the bottom of the crucible but no lump metal was found at all...

In both cases the slag was properly fused and completely non-porous. In the first test the slag composition was approximately (in mol) CaS/Al2O3/CaO = 0.5 / 4/3 / 4/3. A lump of it, dunked in 32 w% HCl did react slightly at first and H<sub>2</sub>S could be smelled. On the plus side, 24 h later and the (approx.) 5 g lump of slag had completely disintegrated and mostly actually dissolved. The solution also tests strongly positive for the target metal.

It's not clear whether this dissolution is partly due to hydrolysis of the CaS or due to hydrolysis of the alleged Ca aluminate (or both), as I don't remember how this type of slag (but minus the CaO) normally behaves in HCl.

The next test will be in a chlorate (no CaS is then formed) boosted SiO2 thermite, possibly in conjunction with a small amount of CaF2. If lump metal forms and the slag does indeed dissolve in HCl then the Si-metal(loid) recovery problems I've been having may be solved by use of CaO. But right now I'm waiting for a new consignment of Al powder...

%%%%%%

Here's an interesting thermite based patent (EP19920106428) abstract, that, for once, actually gives some detail on the formulations used.

<i>"Ferroniobium was produced, using niobium ore (niobium oxide) and iron oxide as starting metallic oxides, aluminum as a reducing agent, sodium chlorate as an exothermic agent and a mixture of fluorite and quick lime as a slag forming material."</i>

In addition they also used small quantities of size-reduced ferroniobium (FeNb), as a heat sink (not entirely sure why).

One of the formulations gives a good idea of the ratios and amounts of CaF2/CaO used (quantities in kg):

Nb ore .......................... 1,000
Fe ore ........................... 162
Al ................................... 350
NaClO3 ........................ 50
CaO .............................. 60
CaF2 ............................. 80
FeNb ............................. 50

So the quantities of slag former and slag fluidiser used are really quite small (but Nb2O5 is a real scorcher, due to its high oxygen content).

[Edited on 27-9-2008 by blogfast25]

[Edited on 27-9-2008 by blogfast25]
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[*] posted on 2-10-2008 at 09:04


A few more tests with CaO at a lower level and in combination with CaF2 didn't really yield anything useful, at least not in sulfate boosted reactions.

I decided to give it one last shot in a silicon thermite, potassium chlorate boosted. The formulation was:

.......................................mol
SiO2 ............................... 1
KClO3 ........................... 0.36
Al ................................... 2.05
CaO ............................... 0.72
CaF2 .............................. 0.41

The CaO/Al2O3 target ratio was 0.70 (I was trying to get 1 but was a little short of lime).

A 20 g batch burned quite hot but no slag collected at the crucible bottom, rather a greyish, porous mass was formed. And then something strange happened.

Dunking the lightly crushed slag in 30 w% HCl, I noticed strong formation of gas, presumably hydrogen. This would be very unusual as silicon doesn't react with HCl and if the reaction had proceeded stoichiometrically there should be no free Al worth speaking of... I left the beaker and contents to stand for a few minutes. After I came back, a hot column of foam had started to flow over the beaker's rim and crackling noises could be heard. In fact, miniature 'explosions', including very small but clearly visible flames could be observed. I suspect that these were due to silane (SiH<sub>4</sub> and possibly higher silanes) spontaneously bursting into flame upon contact with air. There was also considerable generation of heat.

Possibly the lime reacts partly with the silica, forming a Ca silicate of sorts, which may be harder to reduce than pure silica. Some unreacted Al would then be able to alloy itself with any silicon formed, forming an aluminium silicide of sorts, which with HCl would evolve silane(s) and hydrogen...
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12AX7
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[*] posted on 2-10-2008 at 19:38


Aluminum and silicon form a simple eutectic system, no intermetallics ("silicide").

Possibly a calcium silicide?

Tim




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[*] posted on 2-10-2008 at 21:31


@Blogfast & chloric, the foremost exotic thermiters:-

I was recently surprised to read - in a "usually reliable" text - that both Ba and Sr can be produced by a thermite process (but no references given.)

Sr has MP=777C, BP=1382C (a bit low).
Ba has MP=727C, BP=1897C (better),
both according to CRC.

Haven't looked at the energy angle, but both metals are rather energetic.

Ever thought of these two?

Regards, Der Alte
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[*] posted on 3-10-2008 at 06:48


Brauer gives prep with aluminum-containing briquettes under heat and vacuum. Not really exothermic so not really fair to call it thermite.

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[*] posted on 3-10-2008 at 07:00


@Tim and DerAlte:

Ignoring for a minute the entropies, the reaction enthalpies (ΔH) for MeO + 2/3 Al ---> Me + 1/3 Al<sub>2</sub>O<sub>3</sub> (with Me = Ca, Sr or Ba, resp. HoFs - 635 kJ/mol, - 592 kJ/mol and - 548 kJ/mol) are all positive, so no spontaneous change of state is to be expected (but things might be different at much higher temperature if the entropic terms (- TΔS) kicks in - without gaseous reaction products the entropic term is usually really small though).

Certainly I wouldn't expect CaO, SrO or BaO to react with Al in backyard thermite conditions (what might be achieved by strongly heating such mixtures and by siphoning off the volatile earth alkali metal reaction product, assuming some forms in equilibrium, is a different matter). DerAlte, it may be possible that the thermite references to SrO and BaO may have been in reference to that: strongly heated mixtures of Al and MeO, with removal of any formed Me.

Magnesium, at high enough temperature and subject to removal of both the formed CO and Mg can be produced carbothermically from MgO and cokes, for instance... For an aluminothermic high temperature process to produce Sr or Ba, low BP could actually be an advantage...

Tim, I can't really see how else to explain the formation of silane(s), as formation of Ca in significant quantities seems highly unlikely in these conditions. Wouldn't an alloy of Al and Si, perhaps in particular ratios, generate some silane? Just a guess...

[Edited on 3-10-2008 by blogfast25]
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