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Fantasma4500
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Quote: Originally posted by MrHomeScientist  | You could do the gas tests for HCl and Cl2 to put an end to the speculation.
Put an open bottle of ammonia next to your setup when you generate the gas - a white mist of NH4Cl will form if it is HCl.
Chlorine is (supposedly) the only gas with a bleaching effect, so hang a colored piece of paper or flower petal in the gas and see what happens. HCl
might bleach a bit too, so the first test would be more conclusive. |
about chlorine, if i can get to form it in large quanities i can do a simple test with steel wool instead, burning steel wool reacts very well with
Cl2 making iron chloride, forms some red smoke.. looks pretty fancy aswell 
but as blogfast stated its bound to water, so it reacts to form HCl and not Cl2, it was very stinging, not the saturated choking feeling as chlorine
gives..
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Fantasma4500
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| Quote: Originally posted by blogfast25 | | Quote: Originally posted by Antiswat | HCl.. it can be bought seldomly as 37%.. what is the maximum concentration you can acquire from say 98% H2SO4 + NaCl?
it seemed as being very very strong concentration when coming from decomposition of FeCl3.. |
What comes off FeCl3 is NOTHING like what you get from NaCl + H2SO4. With the latter, assuming your H2SO4 is 95 % or better you basically get pure HCl
and if you lead it through water and get your numbers right, you can make 37 w% HCl.
Hint: without heating the reaction goes to:
H2SO4(l) + NaCl(s) === > HCl(g) + NaHSO4(s)
There are plenty of posts on H2SO4/NaCl-based HCl gas generators on this forum, so UTSF.
[Edited on 23-4-2013 by blogfast25] |
oh.. im sorry for not searching it up..
mean i have seen some .. oh.. graph i think its called over HCl potential concentration where it was capable of going as high as 70% some 50*C below
0*C
anyways i have searched around for maximum concentration of HCl a few times and never really got any results.. guess i missed the magical word for
what you call it, just as you get many times more useful results if you include the word synthesis in a search (:
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blogfast25
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At room temperature the maximum achievable HCl concentration is about 37 w%: that is known as 'concentrated hydrochloric acid'. At higher
temperatures solubility of gases in solvents (in general) decreases.
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Fantasma4500
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| Quote: Originally posted by 12AX7 |
A density of around 4 is much more suggestive of a ferrite (density 4.8-5.0).
You can determine if it's strontium or barium ferrite by adding a weak solution of a sulfate (sulfuric acid, potassium or ammonium sulfate, etc.) to
the magnet solution.
Tim |
i just tried this and i got no ppt.. :S
i used 37% H2SO4 (~+1m/L)
1.5 mL to around... 0.5 mL of concentrated XxClx solution
it just turned yellow (due to dilution)
could it be that its just iron..? i mean it certainly doesnt feel like iron, it snaps if you put too much pressure on it..
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12AX7
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Hmm, perhaps it's a simpler compound then. Don't know what they would use other than Sr or Ba ferrite though. Perhaps performance wasn't a priority
in its design.
Ferrites aren't very strong, they are technical ceramics and much weaker than, say, a ceramic formulated for strength and beauty, like porcelain. It
could also perhaps be a resin bonded type (think flexible fridge magnets, without the flex), though these have really poor performance.
Tim
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blogfast25
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To convincingly test for iron you need a bit of peroxide solution (to fully oxidise the iron to Fe(III)) and some ammonium or potassium thiocyanate
(aka 'rhodanide') solution (1 to 2 M).
Fe(3+) + SCN(-) === > FeSCN(2+), the latter is a deep, wine red complex.
With an excess of oxalate Fe(3+) forms a green, slightly fluorescent complex: FeOx3(3-) (trioxalatoferrate (III)) but it's harder to see at low
concentrations.
[Edited on 26-4-2013 by blogfast25]
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Fantasma4500
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| Quote: Originally posted by blogfast25 | To convincingly test for iron you need a bit of peroxide solution (to fully oxidise the iron to Fe(III)) and some ammonium or potassium thiocyanate
(aka 'rhodanide') solution (1 to 2 M).
Fe(3+) + SCN(-) === > FeSCN(2+), the latter is a deep, wine red complex.
With an excess of oxalate Fe(3+) forms a green, slightly fluorescent complex: FeOx3(3-) (trioxalatoferrate (III)) but it's harder to see at low
concentrations.
[Edited on 26-4-2013 by blogfast25] |
i dont have any thiocyanates really, but i have seen it described that you can make FeC2O4 (iron oxalate) and then decompose that into iron powder,
which i think i might do later on to be very sure there is iron in..
it does have a bright yellow tint when you swird the solution around and theres only a thin layer of the solution on the jar's sides so im very sure
there is at least some iron in it..
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Fantasma4500
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| Quote: Originally posted by 12AX7 | Hmm, perhaps it's a simpler compound then. Don't know what they would use other than Sr or Ba ferrite though. Perhaps performance wasn't a priority
in its design.
Ferrites aren't very strong, they are technical ceramics and much weaker than, say, a ceramic formulated for strength and beauty, like porcelain. It
could also perhaps be a resin bonded type (think flexible fridge magnets, without the flex), though these have really poor performance.
Tim |
well...
if this isnt a metal, and its conductive it could perhaps work for electrolysis purposes, tho im pretty sure it wont last as it contains metal.. when
you put it next to a magnet it gets attracted or well the other way.. so it does contain some metal im sure
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blogfast25
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| Quote: Originally posted by Antiswat | i dont have any thiocyanates really, but i have seen it described that you can make FeC2O4 (iron oxalate) and then decompose that into iron powder,
which i think i might do later on to be very sure there is iron in..
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I doubt that, but could be wrong on this, ferrous oxalate would decompose to iron. All transition metal oxalates that I know of, on heating decompose
to their oxide, at least in the presence of air.
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Vargouille
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It is the case that ferrous oxalate decomposes to pyrophoric iron. There are sites and videos that demonstrate this. It could be that the decomposition of oxalate to carbon dioxide, and the resulting presence of carbon dioxide,
prevents significant oxidation of the iron. I don't believe that ferrous oxide behaves so spectacularly when being shaken out as does pyrophoric iron.
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blogfast25
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Well, I stand corrected Vargouille. I've got quite a bit of ferrous oxalate, so I might give that a try...
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Fantasma4500
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please try and see how fast you can make this stuff react
ive put this on my neverending list of chemsitry stuff to do
its fairly simple to do and ive got what i need so i should get around it some time..
i expect to see interesting stuff with sulfur
a bit offtopic tho
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