Sciencemadness Discussion Board - red soluble metal chloride??

red soluble metal chloride??

Fantasma4500 - 19-4-2013 at 14:38

ok, before you all think back on metal x....

i just want to hear if any of you knows something about a metal that gives off red chlorides
the source i have is some magnets i found inside a LCD tv screen, behind these square pieces there was some really nice spoons of dense copper wire

i put a few pieces in 30% HCl
and now i got a really nice dark brown/reddish/orange solution
i can see theres iron in it aswell.. i have tried searching up all kinds of rare metal chlorides but no luck..
doesnt look like a boring metal chloride

simba - 19-4-2013 at 14:56

Palladium chloride is red and soluble in lots of different solvents.

[Edited on 19-4-2013 by simba]

Lambda-Eyde - 19-4-2013 at 15:05

Cobalt(II) can be red.

12AX7 - 19-4-2013 at 19:10

I'm thinking strontium or barium ferrite, in which case the color is iron(III).

It would be wonderful if they made magnets with palladium, as it would be very easy to salvage and recover

Tim

Finnnicus - 20-4-2013 at 01:18

Yeah, cobalt. Probably this http://en.m.wikipedia.org/wiki/Cobalt(II)_chloride from samarium colbalt magnets, samarium(III) chloride solutions are clearish.

[Edited on 20-4-2013 by Finnnicus]

Fantasma4500 - 20-4-2013 at 01:55

palladium chloride, thats very expensive...
well .. by any means i have some interesting metal in these magnets.. i think i have +200g of them, its like, theyre not really magnetic, they loose magnetic power over time..
think ill take a picture of it.. very clear and nearly flouroscent colour
cobalt could be possible, i have checked up on samarium chloride, IIRC the solution doesnt have any colour as its a white salt
judging by this picture it could potentially be iron-cobalt-samarium magnets i have!

http://www.rareearth-permanent-magnet.com/china-strong_sinte...

ill take some pictures of it..

edit: link to pictures
http://imgur.com/a/qFs8X

[Edited on 20-4-2013 by Antiswat]

Fantasma4500 - 20-4-2013 at 02:23

i found the density to be 4.4-4.7
both by measuring a piece then weigting
another one by measuring how many mL its volume is in a 100 mL measuring cylinder and then weighting it
its a hotpressed metal powder after all, which might explain that it doesnt come near the density of cobolt, samarium, iron or copper which is what its apparently composed of..
i can seperate the iron by boiling it down and decomposing the FeCl3 into Fe2O3 and Cl2 removing with NH4OH (fantastic stuff)

blogfast25 - 20-4-2013 at 04:50

Quote: Originally posted by Antiswat

i can seperate the iron by boiling it down and decomposing the FeCl3 into Fe2O3 and Cl2 removing with NH4OH (fantastic stuff)

Careful with notation: FeCl3 with ammonia solution yields Fe(OH)3.nH2O and NH4Cl (not Cl2).

Fantasma4500 - 20-4-2013 at 06:35

Quote: Originally posted by blogfast25

Quote: Originally posted by Antiswat

i can seperate the iron by boiling it down and decomposing the FeCl3 into Fe2O3 and Cl2 removing with NH4OH (fantastic stuff)

Careful with notation: FeCl3 with ammonia solution yields Fe(OH)3.nH2O and NH4Cl (not Cl2).

oh.. no no not that way.. i will put a beaker nearby to react the Cl2 that has gotten out of solution with the ammonia

blogfast25 - 20-4-2013 at 08:49

Quote: Originally posted by Antiswat

oh.. no no not that way.. i will put a beaker nearby to react the Cl2 that has gotten out of solution with the ammonia

Cl2 that has gotten out of what solution? How actually?

[Edited on 20-4-2013 by blogfast25]

Vargouille - 20-4-2013 at 08:51

Do you mean HCl? Heating FeCl3 will not give you Cl2 and Fe2O3, but rather Fe2O3 and HCl. In one case, a redox reaction occurs, while in the other a simple acid-base reaction occurs.

Fantasma4500 - 20-4-2013 at 10:12

Quote: Originally posted by Vargouille

Do you mean HCl? Heating FeCl3 will not give you Cl2 and Fe2O3, but rather Fe2O3 and HCl. In one case, a redox reaction occurs, while in the other a simple acid-base reaction occurs.

are you sure?
i smelt it as Cl2 when i did this last
and the gas was in the glass thing i was heating it in, it didnt escape..
i also saw that it was well atleast a heavy gas by opening a bottle of ammonia, squeezing it and then blowing in the direction of the glass container just above the bottle, directing ammonia fumes into the (chlorine?) gas.
I dont understand where the hydrogen could come from when decomposing FeCl3..

chemcam - 20-4-2013 at 10:34

FeCl3 reacts with moist air to give the HCl. IIRC FeCl3 is deliquescent.

blogfast25 - 20-4-2013 at 10:49

Quote: Originally posted by Antiswat

Ermm... YES! The smell of Cl2 and HCl can be mistaken for each other because they're both quite pungent, if it's smell you're going by.

It is HCl vapour (or gas) that reacts with airborne ammonia to give fumes of NH4Cl.

The 'hydrogen' comes from the hydrolysis of the FeCl3

FeCl3 + 3 H2O === > Fe(OH)3 + 3 HCl (oversimplified).

FeCl3 would only decompose to Cl2 at very high temperature and NEVER from FeCl3 solutions.

Fantasma4500 - 20-4-2013 at 10:51

hm.. well the rust still gave off some gas that reacted with ammonia, and at that point it was totally dry, guess it couldnt react with all the water in the air

12AX7 - 20-4-2013 at 16:45

SmCo alloys have density 8.2-8.4, so it is unmistakably not SmCo. It would also give off a lot of hydrogen, or require a lot of oxidizer: cobalt isn't an especially reactive metal, and may require some effort to dissolve (I don't know if this is the case; it is with nickel, but isn't with iron). It would not simply dissolve in acid.

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. Wash and filter the precipitate (wash with mild acid until it's white and clean), then perform a flame test. Strontium will glow red, barium green; a mixture will look yellow, but don't mistake this color for the pervasive color of sodium's yellow-orange color (don't use sodium sulfate, if you can avoid it, because that would be a source of contamination).

Tim

12AX7 - 20-4-2013 at 16:51

Also, FeCl3 does give off a small amount of chlorine. The redox potential of the pure stuff is actually quite powerful, such that starting with a pure compound of Fe(III) (such as the ferrite in question) can cause the solution to bubble a small amount of Cl2, leaving an equilibrium of Fe(II) and Fe(III).

Tim

Fantasma4500 - 21-4-2013 at 04:13

Quote: Originally posted by blogfast25

Quote: Originally posted by Antiswat

oh.. no no not that way.. i will put a beaker nearby to react the Cl2 that has gotten out of solution with the ammonia

Cl2 that has gotten out of what solution? How actually?

[Edited on 20-4-2013 by blogfast25]

my mistake, i meant when i evaporated all the water and then started to decompose the solid (:

Fantasma4500 - 21-4-2013 at 04:18

Quote: Originally posted by 12AX7

SmCo alloys have density 8.2-8.4, so it is unmistakably not SmCo. It would also give off a lot of hydrogen, or require a lot of oxidizer: cobalt isn't an especially reactive metal, and may require some effort to dissolve (I don't know if this is the case; it is with nickel, but isn't with iron). It would not simply dissolve in acid.

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. Wash and filter the precipitate (wash with mild acid until it's white and clean), then perform a flame test. Strontium will glow red, barium green; a mixture will look yellow, but don't mistake this color for the pervasive color of sodium's yellow-orange color (don't use sodium sulfate, if you can avoid it, because that would be a source of contamination).

Tim

strontium or barium??? this is getting interesting...

it doesnt really dissolve, its more like... its just laying in solution and then very slowly the solution turns more and more red until all HCl is used up
but i have H2SO4 so that should be interesting

i guess also i could detect that it has actually reacted by putting ammonia gas over the solution while its boiling to see if theres HCl formed tho it wont give me any clue on what metal ions theres in there..
interesting..
i think before i move on to making the sulfates ill scale it up abit so i can get massive precipitate, perhaps electrolysis of HCl with the pieces..

blogfast25 - 21-4-2013 at 05:02

Quote: Originally posted by Antiswat

my mistake, i meant when i evaporated all the water and then started to decompose the solid (:

Just to be clear (it really helps in science to be concise!), when you evaporated the water, hydrogen chloride started to come of the solution, NOT Cl2. To get Cl2 you'd have to thermally decompose anhydrous FeCl3.

[Edited on 21-4-2013 by blogfast25]

Fantasma4500 - 22-4-2013 at 11:33

Quote: Originally posted by blogfast25

when you evaporated the water, hydrogen chloride started to come of the solution, NOT Cl2. To get Cl2 you'd have to thermally decompose anhydrous FeCl3.

[Edited on 21-4-2013 by blogfast25]

actually i didnt think about it that way.. i do have some FeCl3, a few grammes.. i dried this slowly evaporating it.. i will try to heat this stuff up to be very very sure.. smell should determine if Cl2 or HCl forms (dont worry.. very carefully smelling, not snorting liquid chlorine..)

edit: update on FeCl3 decomp.
im unsure if i produced anhydrous HCl, it was heated in a candle bottom of aluminium, which it went through (??) not sure if this had something to do with the FeCl3 whatsoever
later heated on knifeblade i smelled the gasses, one time it went straight to my brain, very sharp smell..
it didnt smell like chlorine, so i guess it makes anhydrous HCl somehow, the crystals are all dry incase you were wondering about that..

also back when i got the HCl from heating it, it was all dry, and i was decomposing it into Fe2O3, i didnt know about Fe3O4 > Fe2O3 back then.. (:

[Edited on 22-4-2013 by Antiswat]

blogfast25 - 23-4-2013 at 05:13

Quote: Originally posted by Antiswat

[edit: update on FeCl3 decomp.
im unsure if i produced anhydrous HCl, it was heated in a candle bottom of aluminium, which it went through (??) not sure if this had something to do with the FeCl3 whatsoever
later heated on knifeblade i smelled the gasses, one time it went straight to my brain, very sharp smell..
it didnt smell like chlorine, so i guess it makes anhydrous HCl somehow, the crystals are all dry incase you were wondering about that..

[Edited on 22-4-2013 by Antiswat]

The FeCl3 you're talking about is in fact FeCl3.6H2O, ferric chloride hexahydrate (the stuff they use to etch copper with). On heating that further hydrolyses. Very simply put:

FeCl3.6H2O === > Fe(OH)3 + 3 HCl + 3 H2O

On further heating the hydroxide then dehydrates to Fe2O3:

2 Fe(OH)3 === > Fe2O3 + 3 H2O

Anhydrous ferric chloride (FeCl3) is made by reacting dry iron powder or filings with dry Cl2, while heating. This product would decompose to iron and chlorine (Cl2) but only on very strong heating (forget a candle!)

This is all very similar for AlCl3 and its hydrates.

Aluminium reacts with ferric chloride hydrate by redox displacement: Fe³⁺ + Al === > Fe + Al³⁺. This probably explains why your 'crucible' got eaten up. Aluminium also has a low MP: about 660 C, so even a candle isn't really safe to use with it! At a minimum use steel soup cans, or invest in cheap ceramic crucibles.

Fantasma4500 - 23-4-2013 at 08:22

hm.. well i used a propane blowtorch on low gas output, weak flame, no platinum in it just airintake through the sides of it..
but yes this might explain why it just went broke nearly instantly.. very interesting (:
could be used for ripping over my aluminium tubes in the future..
but i didnt actually know iron chloride was bound to water atoms..? makes perfect sense now why it forms HCl, very very concentrated HCl..

abit offtopic and yet not..
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..

MrHomeScientist - 23-4-2013 at 09:23

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.

blogfast25 - 23-4-2013 at 12:14

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]

Fantasma4500 - 23-4-2013 at 13:07

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..

Fantasma4500 - 23-4-2013 at 13:09

Quote: Originally posted by blogfast25

Quote: Originally posted by Antiswat

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 (:

blogfast25 - 24-4-2013 at 04:30

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.

Fantasma4500 - 25-4-2013 at 06:05

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..

12AX7 - 25-4-2013 at 17:06

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

blogfast25 - 26-4-2013 at 04:59

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]

Fantasma4500 - 26-4-2013 at 08:13

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..

Fantasma4500 - 26-4-2013 at 08:14

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

blogfast25 - 27-4-2013 at 05:00

Quote: Originally posted by Antiswat

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.

Vargouille - 27-4-2013 at 07:50

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.

blogfast25 - 27-4-2013 at 11:57

Well, I stand corrected Vargouille. I've got quite a bit of ferrous oxalate, so I might give that a try...

Fantasma4500 - 28-4-2013 at 13:18

Quote: Originally posted by blogfast25

Well, I stand corrected Vargouille. I've got quite a bit of ferrous oxalate, so I might give that a try...

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