Sciencemadness Discussion Board - How do I recrystallize ferric chloride?

How do I recrystallize ferric chloride?

Carbon8 - 24-5-2018 at 12:42

Does anyone have any ideas about how to recrystallize FeCl3 from an acidic etching solution. I have had a fan blowing on an evaporation dish of the stuff for a week and the fluid level has barely budged.

I haven't found much information online about low-energy or low-cost methods of purifying FeCl3. Is that because there are none?

fusso - 24-5-2018 at 12:55

You simply can't, FeCl3 is extremely deliquescent and will absorb moisture from atmosphere.

[Edited on 24/05/18 by fusso]

VSEPR_VOID - 24-5-2018 at 13:28

I used this method to dry holmium oxide,

Place your solution into a crystallization dish, over a beaker of water. Boil the water and allow the steam to contact the bottom of the dish. This will keep your dish at about 60C and not exceeding 100C. This ensures that your solution will evaporate quickly without decomposing or being etched into the dish by heat. I highly recommend this method.

Carbon8 - 24-5-2018 at 16:18

Thanks to all for your comments.

The etching solution lists its ingredients as ferric chloride and HCl.

I don't have a compelling reason for wanting crystals, I just want to see if it's possible. I will give the warming treatment a try, although I've read that there is a risk that the ferric chloride will decompose to ferrous chloride and HCl.

VSEPR_VOID - 24-5-2018 at 17:02

If it does decompose find out at what temperature. If its above 100C then you will be fine using steam

fusso - 30-5-2018 at 00:22

Update: FeCl3 can be dried by fan according to Nilered
Seems like where you live is very humid or simply the weather is very humid at that time.

AJKOER - 1-6-2018 at 05:51

Comments, first avoid sunlight (or diffused light) otherwise with aqueous Fe(lll) at low pH (best at pH 3) as you could have some photo-fenton chemistry afoot (see Eq 4 at http://cdn.intechopen.com/pdfs/29377/InTech-Fundamental_mech... ) forming ferrous.

Next, limit oxygen exposure with light, as any created Fe(ll) (especially if complexed with chloride, see https://www.sciencedirect.com/science/article/pii/S001670371... ) with O2 and H+ will likely undergo an electrochemical reaction leading to a basic ferric chloride.

Unwanted ferrous can also be created in the presence of other transition metal salts (for example, a copper ion impurity). A redox couple equilibrium may form resulting in an equilibrium concentration of ferrous, which will be re-established as ferrous is consumed per any of the above.

In short, salts of increasingly anodic metals, and further those that are considered transition metals or related, may produce unexpected products (see, for example, http://www.sciencemadness.org/talk/viewthread.php?tid=83114#... ).

[Edited on 1-6-2018 by AJKOER]

aga - 1-6-2018 at 11:01

Unfortunately ferric chloride is one of those substances that simply refuses to crystallise.

Heating it to dryness decomposes it, so that's not an option.

If you leave it out to dry, it'll either sit there not losing any water, or suck more from the air and remain liquid.

A reaction using anhydrous ferric chloride as a catalyst was shown by NurdRage.
He stuck dried iron wool in a flask of dry toluene and bubbled dry chlorine into it - the anhydrous ferric chloride was formed in situ, catalysing the formation of chlorotoluene(s).

If you could keep the air out, perhaps that'd be a way to produce ferric chloride crystals, depending on IF it would ever crystallise out of toluene/chlorotoluene(s) at all.

That's all random speculation as i never tried it myself.

Edit:

Presumably if that was tried, the toluene would get chlorinated at every position, producing god-knows-what as a by-product. Probably something Bad.

[Edited on 1-6-2018 by aga]

walruslover69 - 1-6-2018 at 11:11

You could probably dry it with azeotropic distillation with toluene in a dean stark trap. Then dissolve it in anhydrous acetone or ethanol and crystallize it from that solution.

aga - 1-6-2018 at 11:32

I wonder ...

If it's soluble in toluene, perhaps you can get crystals by just distilling out most of the toluene, sealing the bottle and chilling it.

I've used ferric chloride for 37 years for making PCBs.
The first stuff i ever bought was military surplus.
It came double-bagged in thick plastic, and was dry (as i saw it then).

It certainly would be nice to find a way to get ferric chloride drier, and smaller.

Amateur gold refiners produce a lot of this stuff in their final 'stock pot'.

It'd be useful to find a way to process it into a small bag of orange-brown crystals (or mush) rather than a 5 gallon bucket of trouser-staining goop.

I do believe you've sparked something off here Carbon8

walruslover69 - 1-6-2018 at 11:55

I wasn't aware it was soluble in toluene but that makes since knowing what else it is soluble in. I assume something like hexane or other nonpolar could be added to lower the solubility and crystallize it out of toluene.

It just occurred to me that using toluene might be slightly problematic. Would the ferric chloride chlorinate the toluene to some degree while its refuluxing for 1-2 hours?

aga - 1-6-2018 at 12:47

Toluene distills quickly, so most of it should be gone in a simple distillation in about 30 mins.
(i'm imagining small quantities, not 60-gallon flasks)

Iron probably hangs onto the chloride ion far better than some weakling sub-metal organic could ever hope for.

Cool/stopper the flask as quick as possible and maybe FeCl₃ crystals ?

Not sure anyone knows .... now there is a Challenge !

Superb suggestions walruslover69.

Edit:

A simple experiment to see if it is soluble to any extent in toluene would be to stick a few mls of toluene in a test tube, add the same mls of water, then put a few drops of FeCl₃ in.

The water/toluene should separate into distinct layers.

If there's any brown colour in the upper layer after a bit of mixing, bob's your uncle.

Further Edit:

If anyone has a test tube, some water, toluene and ferric chloride, please try this and post a photo ASAP.

If not, i'll move the separation-from-paint-thinner to make more toluene to the head of the queue.

Oh, and then distill it. Important that part is.

[Edited on 1-6-2018 by aga]

walruslover69 - 1-6-2018 at 15:11

I didn't have any ferric chloride laying around so I just quickly dissolved some iron oxide in hydrochloric acid and treated it with h2o2. I took 5ish ml and added it to the cylinder along with 5 ml of toluene. I swirled it around and it appears that none of iron went into the toluene layer.

Wikipedia says that it is very soluble in diethyl ether so I tried the exact same thing and got the exact same result of no iron being present in the organic layer.

THE PLOT THICKENS!

phlogiston - 1-6-2018 at 16:03

Put it in a airtight box or bag with an even more hygroscopic substance.
Google 'dessicator' for inspiration.

violet sin - 1-6-2018 at 20:21

Method for preparing anhydrous iron chlorides
US1938461A https://patents.google.com/patent/US1938461

[This also describes some older processes that maybe useful for other routes]

:::::: since dehydration of the tetrahydrate results in hydrolysis and only basic salts are left. Likewise, attempts to dehydratethe ferric chloride hexahydrate, FeClsnGI-IzO, result in hydrolysis, There are methods for preparing the anhydrous ferrous salt from ferrous chloride tetrahydrate for instance the double ammonium salt may be formed by dissolving the tetrahydrate .in' an ammonium chloride solution, evaporating the solution of the double salt to dryness and heating the'residue in the absence of air. The waterand ammonium chloride escape and-leave the anhydrous ferrous chloride behind. ::::::

..."by evaporatinga ferrous chloride solution at an elevated temperature so as to precipitatethe dihydrate crystals and effecting a separation of the crystals from the mother liquor at a temperature above about 90 C. I have found that it is possible, in the substantial absence of air or oxygen, to dehydrate the ferrous chloride 'dihydrate compound directly without "hydrolysis, and without departing from the solid phase, in order to secure substantially amhydrous ferrous chloride...

... I have found that anhydrous ferric chloride can most advantageously be produced by simply chlorinating anhydrous ferrous chloride...

...The anhydrous ferrous chloride thus prepared may then be passed in intimate contact-with dry chlorine gas to .chlorinate it substantially to anhydrous ferric chloride. In effecting this chlorination, I have foundit expedient to pulverize the anhydrous ferrous salt so that substantially'all of it will pass through a standard Tyler 40-mesh screen. It is also preferable that. during chlorinationheat be supplied to the salt so as to raise its temperature to between about 140 and 190-C.', preferably about 175 0., since this materially reduces the time required for chlorination, and increases the percentage conversion to the ferric salt. This method for the preparation of. anhydrous ferric chloride possesses advantages not enjoyed bythe processes heretofore developed."

[Edited on 2-6-2018 by violet sin]

Melgar - 1-6-2018 at 20:54

Yeah, you can't get anhydrous FeCl3 once it's already hydrated. It's similar to AlCl3 in that respect. You might be able to get hydrate crystals though.

The trouble is that HCl will leave along with water as you try and dry it, and so it will leave a mixture of iron oxychlorides or whatever they're called, rather than solely FeCl3.

aga - 1-6-2018 at 23:05

Quote: Originally posted by walruslover69

... it appears that none of iron went into the toluene layer.

Thanks for the experiment and the great photo !

Shame it doesn't go into the toluene

Quote: Originally posted by Melgar

... will leave a mixture of iron oxychlorides or whatever they're called, rather than solely FeCl3.

Hmm. Good point.

Quote: Originally posted by phlogiston

Put it in a airtight box or bag with an even more hygroscopic substance.
Google 'dessicator' for inspiration.

Doesn't work for this compound, as stated above.

Google 'ferric chloride' for more inspiration

This patent: https://patents.google.com/patent/CN103991911A/en
says "ferric trichloride solution is heated, boiled, cooled and crystallized to obtain high purity ferric trichloride crystal."

Could be rubbish. Could be true.

Another googled snippet mentioned 'ferric chloride hexahydrate", which is presumably crystalline.

I suppose the conditions must be acidic and rammed with chloride ions to keep the compound intact, so the question is how to get to that state without having a lot of water lying around.

violet sin - 2-6-2018 at 11:07

Aga, "the ferric trichloride is dissolved in **water** and then filtered to prepare a ferric trichloride solution; the ferric trichloride solution is heated, boiled, cooled and crystallized to obtain high purity ferric trichloride crystal."

"The method of preparing a high-purity ferric chloride claim, wherein said ferric chloride to ferric chloride hexahydrate crystals."

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

So this is not what you want. The patent I linked does make anhydrous ferric (and ferrous) chloride. But the problem is, there is a need to chlorinate the ferrous salt, but it's under 200℃. Still it's not ideal. There was mentioned the decomposition of ammonium chloride/ ferric chloride double salt in vacuum to get ferric chloride anhydrous. Check it out.

I surely concede, nothing is perfect for every home chemist. but I am going to try mine at some point in prep for a project.

aga - 2-6-2018 at 11:55

Basically the OP wants crystals from already-dissolved ferric chloride.

Me too - got lots of it and very little usable space to keep it, hence my interest in reducing that solution to reusable crystals (even mush).

NurdRage has a few good videos about recycling PCB etching ferric chloride solution, but focus more on 100% recycling rather than retaining ferric chloride as such.

Bezaleel - 11-6-2018 at 09:16

An idea:
* Make the FeCl3 solution as concentrated as possible by evaporation.
* Add the double volume of 96% ethanol to it. (The solubility of FeCl3 in ethanol is only some 10% lower than in water.)
* Absorb the water by the addition of water free CuSO4 of MgSO4 or the like. (Tricky: does the water absorber have enough strength to push the FeCl3 into the ethanol?)
* Filter the solution to remove the hydrated water absorber.
* Distill off the ethanol. (This is tricky: will the dissolved FeCl3 now also decompose in one way or another?)

woelen - 12-6-2018 at 05:27

The reason that no FeCl3 from aqueous solution is extracted into an organic layer is due to the ionic nature of the hydrated compound in solution.

If you take anhydrous FeCl3, which is a black solid, looking very much like iodine, then you can dissolve this in quite a few organic solvents. The same is true for anhydrous CuCl2 and anhydrous AlCl3. These anhydrous compounds are covalent compounds.

In water you get metallic ions, partially coordinated to chloride ions and water molecules, giving rise to ionic species in solution and these do not extract into an organic (apolar) layer.

BTW, I have never seen crystalline hydrated ferric chloride. The stuff which you can buy at electronics shops as copper etchant is a solid, consisting of little amorphous spheres, the size of peanuts, which dissolve exceptionally well in water and give a solution which easily hydrolyses to basic chloride. The anhydrous compound is crystalline, it consists of many small glittering dark grey/black crystals, which also dissolve very well in water.

I expect that the best you can get with hydrated ferric chloride is a cake of amorphous solid when you boil down the solution of ferric chloride from strong hydrochloric acid and assure that the solid material has a lot of HCl-gas around it when dried. Maybe you can get a solid if you boil down to an acidic syrup from conc. HCl and the nearly solid syrup is stored in a tightly closed container in which also some NaOH and/or conc. H2SO4 is stored in a separate dish.

[Edited on 12-6-18 by woelen]

walruslover69 - 12-6-2018 at 07:28

Going off that. Shouldn't you be able to precipitate and possibility crystallize it out of solution by adding ethanol or IPA?

Bezaleel - 12-6-2018 at 16:11

Quote: Originally posted by walruslover69

Going off that. Shouldn't you be able to precipitate and possibility crystallize it out of solution by adding ethanol or IPA?

Well, that's the basis of my idea two posts above this post, but as woelen explains, you won't dehydrolyse the iron(III) by the addition of ethanol or the like. Put otherwise, ethanol has a weaker affinity for H2O than Fe3+ which exists as an aquo or chloroaquo complex and keeps water softly bound around that complex.

Probably, when you add water to an FeCl3 solution in ethanol, a layer will separate out, containing basic FeCl3 solution in water.

I wonder what happens when you add CaCl2.0H2O to an acid saturated FeCl3 solution.

happyfooddance - 12-6-2018 at 20:37

Quote: Originally posted by Bezaleel

I wonder what happens when you add CaCl2.0H2O to an acid saturated FeCl3 solution.

I am thinking a fair bit of HCl fuming would result.

I think the increased chloride saturation and dehydrating properties of calcium chloride would force HCl out of solution, but the FeCl3 will remain hydrated.

walruslover69 - 13-6-2018 at 05:41

Quote: Originally posted by Bezaleel

I was originally under the impression that hydrated FeCl3 would be insoluble in organic solvents like ethanol, IPA, acetone etc... but from some research it appears the hydrate is soluble in most semi polar solvents just not to the extent that the anhydrous form is.

Does anyone have any solubility data for FeCl3 6H2O? I have been trying to find some but I am unable to find any actual numbers except for the solubility in water.

I think my original idea of azeotropic distillation might be the most promising.

My rough idea of the procedure is as follows.

add a substantial amount of anhydrous IPA to a concentrated solution of FeCl3. Ideally all of the water should be able to distill over as the IPA-water azeotrope leaving only a solution of FeCl3 6H2O in IPA.

Evaporating the IPA might produce solid FeCl3 6H2O crystals.

If not, the addition of a miscible organic solvent like toluene or something could be added to precipitate the FeCl3 6H2O

Second thought- Doing the azeotropic distillation with toluene instead of IPA might work out better.

DraconicAcid - 20-7-2018 at 14:59

Dammit. I was all set to do a Friedels-Craft alkylation when I found out my ferric chloride is the hexahydrate. I'm pretty sure that won't work.

I can probably convert it to NBu4[FeCL4], but I doubt that will help.....

clearly_not_atara - 20-7-2018 at 15:07

If you can make a dry tetrachloroferrate, you can probably pyrolyse it to obtain FeCl3, no?

DraconicAcid - 22-7-2018 at 12:07

Possibly.

It seems you can also extract H[FeCl4] into ether or some other organic solvents. I might try that.

DraconicAcid - 25-7-2018 at 15:53

I checked the storeroom at the other campus, and lo and behold! two bottles of anhydrous ferric chloride! According to the label, though...one of them was 20% full of a dark liquid that sloshed around when I picked it up, and very little solid at all.

JJay - 25-7-2018 at 15:56

I have a large unopened jar of ferric chloride. I'd be interested in seeing details on your alkylation.

clearly_not_atara - 25-7-2018 at 18:27

Quote: Originally posted by DraconicAcid

Possibly.

https://pubs.acs.org/doi/abs/10.1021/ac60202a040?journalCode...

I was very surprised to find that tetrachloroferrate salts "may be dried at 110 C in air without decomposition", and can be prepared from the simplest starting materials: tetramethylammonium chloride and FeCl3 hexahydrate.

[Edited on 26-7-2018 by clearly_not_atara]

clearly_not_atara - 26-7-2018 at 19:32

Can't edit the above post. I think this is the paper you wanted.

https://www.researchgate.net/publication/231427245_The_Extra...

Attachment: laurene1955.pdf (580kB)
This file has been downloaded 469 times

MidLifeChemist - 28-9-2020 at 08:41

I received a 1lb (454g) "brick" of Iron (III) Chloride from US Pigments.

It would not break, it was rock hard. Covered with layers of newspaper and saran wrap. I donned my nitrile gloves and started to unwrap the saran wrap.

This brick was rock-hard, and wouldn't even fit into my 1000ml beaker, remaining lodged at the top of the beaker. However, after a few squirts of distilled water, the edges dissolved and the brick slid down to the bottom of the beaker. I added a couple hundred more milliliters of distilled water, and the solution got very, very hot and was an intense red-brown color. Also, I noticed that this "brick" stains everything around it a brownish-yellow color, even things it doesn't touch (half joking here).

So that was last night. I'm not sure what I'm going to do with it. The beaker is sitting there with 300ml of water in it, and some size chunk of undissolved FeCl3. The way I see it, I have 2 choices:

#1 Add more water until it all dissolves, store it as a concentrated FeCL3 solution is a couple of 16oz HDPE bottles that I have laying around...
#2 - Heat the solution until it all dissolves, and try to recrystalize it in a large glass tray and a fan like NileRed did. I realize will be difficult as FeCL3 is very hygroscopic and may take several days. and I don't have a dedicated rented lab like NileRed does, I have a house with a garage and kids.

I'm open to suggestions from my fellow ScienceMadness members!

teodor - 13-10-2020 at 13:36

Sorry, this is response to 2 years old post, but I found it by google few days ago and something made me interested.

Quote: Originally posted by woelen

The reason that no FeCl3 from aqueous solution is extracted into an organic layer is due to the ionic nature of the hydrated compound in solution.

I did a few experiment with diethyl ether extraction. As a starting material I used freshly precipitated Fe(OH)3 solution in 36% HCl (the full amount of Fe(OH)3 which could be dissolved). So, my initial solution of FeCl3 as I suppose didn't contain anything else except FeCl3 and water.

I red the extraction of FeCl3 from water solution is possible when we have it 6M in HCl.

Just to compare: the ether layer of initial FeCl3 (without HCl) and the same solution + HCl to make it 6M:

(the resolution is not enough to see but I can say the ether layer at the left is totally white).

Then I made some pictures that shows colour of ether layer depending on HCl aqueous solution molarity (at the right it is always 6M):

Of course it doesn't mean it is easy to extract solid from that ether - as we know ether always dissolves a bit of water and I am not sure whether HCl can change (let say, increase) this amount but I think it could.

Also I tested a bit the FeCl3 affinity to water (in organic solvents) and it is really bigger than few compounds I tried to use to dehydration.

But it shows that, possible, with displacing OH ligands with Cl the property of FeCl3 are changing from ionic to more covalent, if woelen's theory is true.

Edit:

The colour of FeCl3 0M in HCl and FeCl3 6M in HCl is completely different, but after first shaking it becomes almost equal, so I think it is either because the ether layer gets that HCl from water or it lowers the concentration of FeCl3.

Also you can see here at the right 3 distinct layer: ether, FeCl3 in contact with ether almost discolored (the same color as ether) and the initial FeCl3 concentration in the bottom layer.

[Edited on 13-10-2020 by teodor]

[Edited on 13-10-2020 by teodor]

[Edited on 13-10-2020 by teodor]

[Edited on 13-10-2020 by teodor]

Texium - 13-10-2020 at 17:04

Nice experiment, teodor. I suppose that what you're extracting into the ether layer is essentially "tetrachloroferric acid," H⁺ [FeCl₄]^- which you could expect to be a little more soluble in organic solvent than what is essentially a mixture of hydrated ferric ions and chloride ions.

Honestly, I know it's not going to catch on, but I think the anhydrous forms of ferric chloride, aluminum chloride, zinc chloride, etc, shouldn't even be thought of as different versions of the same compound as their hydrous forms. When you consider the difference in structure and chemical properties, it's practically as inaccurate as calling acetic anhydride "anhydrous acetic acid." I propose the use of more covalent sounding names for these covalent metal halides. How does trichloroiron sound?

teodor - 14-10-2020 at 02:42

Thank you Texium.

Indeed, in acc. with wikipedia:

FeCl3 * 6H2O has the structural formula trans-[Fe(H2O)4Cl2]Cl * 2H2O
FeCl3 * 2.5H2O has the structural formula cis-[Fe(H2O)4Cl2][FeCl4] * H2O.
FeCl3 * 2H2O has the structural formula trans-[Fe(H2O)4Cl2][FeCl4].
FeCl3 * 3.5H2O has the structural formula cis-[FeCl2(H2O)4][FeCl4] *3H2O.

So, I suppose the ether layer extracts H[FeCl4], the tetrahedral form and the water layer keeps the octahedral [Fe(H2O)4Cl2]2+ which then disproportionates more upon addition of HCl, that way all the iron could be extracted to the ether layer (according to literature).

After standing overnight in a freezer the ether layer was separated into 2 layers. I think the lower one now collected most of the water by formation of the octahedral form again. So, I suppose the upper layer is H[FeCl4].

It's amazing that this acid could be isolated as free acid solution (or is it Fe[FeCl4]3 ?).

If it is free acid, is it possible to make its ammonium salt? And what should be its decomposition product in anhydrous environment?

So, any information about [FeCl4]- acid is welcome.

Edit: Did somebody dissolve anhydrous FeCl3 in diethyl ether? Which color did you observe?

[Edited on 14-10-2020 by teodor]

[Edited on 14-10-2020 by teodor]

teodor - 17-10-2020 at 15:06

I did few experiments trying different organic solvents. I had idea that extraction ability is depending on the ability to dissolve HCl.

So, I found that n-butyl acetate has the same extraction properties as diethyl ether. Indeed, it has also good ability to dissolve HCl, so probably my hypothesis is right.

Unlike diethyl ether n-butyl acetate takes virtually no water from FeCl3/HCl/H2O solution. So it is possible to dry the organic layer completely which gives a yellow film on a glass. This film very fast changes the color from yellow to brown (several seconds), probably it is the action of oxygen. It is not possible to dissolve brown ("oxydised") part of the film in water, but it seams yellow film could be dissolved immediately upon contact with water.

I think it's worth trying to dry the solution in oxygen-free environment.

[Edited on 17-10-2020 by teodor]

ChemichaelRXN - 18-10-2020 at 02:54

What if you keep the beaker with the wet ferric chloride in a type of vacuum desiccator with heat somehow applied to the beaker with a heat strip wrapping? I think that might drive the water out and keep it stored in the desiccant. It might take at least a few days possibly, but it may work.

teodor - 18-10-2020 at 03:02

Yes, I thought about that. Probably the vacuum pump will go up in my shopping list.

But today I think the real reason of "oxidation" is loosing HCl in the drying process. I am thinking about possibility to remove HCl from n-butyl acetate solution without drying.

teodor - 7-12-2020 at 03:47

I did the experiment of FeBr3 extraction with ether and it was quite successful. Unlike FeCl3 or, more accurate to say, FeCl3 * n HCl * mH2O it is quite easy to get some nice-looking crystals by evaporating ether a bit.
It is interesting that product of Fe(OH)3 reaction with HBr has dark-brown color, the same color as bromine itself. (Reaction of Fe(OH)3 with HCl produces light-yellow color).

Fery - 7-12-2020 at 12:51

Teodor, nice experiments! I would be afraid of water present in diethylether as you already mentioned - I see you froze out as much H2O as possible in freezer overnight in FeCl3 experiment. And very probably with FeBr3 too. Were your products completely water free? Your method is cool way of preparing anhydrous compounds and avoids dangerous vigorous direct reaction of halogen with metal. I see you also used butylacetate which is more water free than diethylether. Couldn't be also methylisobutylketone used? I know it is used for precious metal extractions (but as cyanides), perhaps it could work with FeCl3/H[FeCl4]/FeBr3 too?

teodor - 7-12-2020 at 13:38

Thanks Fery.

In the case of FeCl3 the water goes together with H[FeCl4] molecules far beyond usual capacity of the organic layer, at least in case of ether. This process was studied by Ernest H. Swift in various works, for example here: https://doi.org/10.1021/ja01303a058

I didn't find any study of other solvents, so my result with n-butyl acetate is quite interesting. As a next step I plan to measure water amount which goes into organic layer as well as Fe, Cl. It is possible that water extraction is always proportional to Fe3+ extraction, so we either get water or no Fe.

Edit: I doubt that freezing the water out (as on my picture) freeze all the water. But evaporation of ether on air is not quite accurate test for that. But I didn't make more accurate measurement yet.

Also H[FeCl4] was never isolated as far as I know. But it is the same as make the products water free. So, I doubt it is easy. FeCl3 has bigger affinity to water than many desiccants and I didn't find much information about H[FeCl4] without water, usually it forms H[FeCl4] * nH2O (according to Swift).
I think I will give a try to AlCl3, SnCl4. Also have no idea about possible reaction with P2O5.

Edit: probably the easier is to test FeCl3 (anhydrous) solution in an absolute ether with different water-containing compounds.

I didn't study FeBr3 yet except one short experiment which immediately gave me some crystals from ether, probably FeBr3 * nH2O. But I already mentioned the color difference, that probably means also the difference of the Fe complex ion itself.

About methylisobutylketone - good idea thanks. I will compare with n-butyl acetate if will put my hand on it.

[Edited on 7-12-2020 by teodor]

[Edited on 7-12-2020 by teodor]

Fery - 8-12-2020 at 06:27

Teodor, I have 5 L of that solvent. If you have any idea let me know.
Maybe trying to mix anhydrous FeCl3 in excess with solvent in stoppered vessel for few days, then decant the liquid, e.g. 2 ml into testing tube using graduated pipette, the same with diethylether, n-butylacetate...
Then add 5 ml water solution of SCN- into every testing tube (ether/ester/ketone) and shake all tubes to extract Fe3+ from organic phase into water phase and compare visually the intensity of red Fe SCN complex in water phase so we know which solvent is more powerful (presuming all Fe3+ left the organic phase and entered the water phase)? But maybe H[FeCl4] partition coefficient differs from anhydrous FeCl3...

clearly_not_atara - 8-12-2020 at 07:21

Quote: Originally posted by Texium (zts16)

How does trichloroiron sound?

You mean trichloroferrane? :p

I think it might make more sense to say "ferric hydrate chloride" than "ferric chloride hydrate" if we rank the ligands by their proximity to the Fe3+ center. But this cute idea gets really complicated when we try to name the fluorides.

Quote:

H[FeCl4] was never isolated as far as I know.

Probably because it isn't a stable compound. The anhydrous etherate was prepared by combining the anhydrous precursors HCl, FeCl3 and iPr2O.

The interesting question is not "can we prepare anhydrous HFeCl4" (no) but "can we dry H11O5*FeCl4@iPr2O"? Natural choices to try in my opinion are CaCl2 and Na2S2O7, the former being far more efficient by weight but the latter possibly better at removing the last bit of water -- dehydrating this complex solution likely proceeds through several stages.

[Edited on 8-12-2020 by clearly_not_atara]

Bedlasky - 8-12-2020 at 07:51

Fery: Fe-NCS complexes can be quite easily extracted from water in to organic solvents. So your idea doesn't work. I used this technique to demonstrate aerial oxidation of Fe2+.

Fery - 8-12-2020 at 08:21

Hi Bedlasky, good hint. Fe3+ complex with SCN- should be better soluble in organic than in water phase. MIBK is used to extract cyanide complexes of precious metals (Au, platinum group) from water phase even they are present in water phase in very low concentrations.
So then better to evaporate the organic solvent and add water solution of SCN- to the remainder to know approximately the amount of Fe3+ extracted into organics by just looking how much dark color obtained to know which solvent is the best...
Also organohalogens (CH2Cl2, CHCl3) could be tried to extract FeCl3 / H[FeCl4] and then evaporating and finally determining the amount of Fe extracted using a naked eye (an approximate colorimetry).

Swinfi2 - 9-12-2020 at 09:05

I cant remember the source but I remember reading a patent a while back about crystallising a hexamine metal chloride then thermally decomposing the amine groups off the metal as a way around the irreversible hydration.

Does anyone know what metals that might work for?

itsallgoodjames - 11-12-2020 at 07:16

It would probably work with nickel and cobalt, though I don't have a source with that. I doubt it would work with ferric chloride though, as I don't believe that hexammineiron(iii) chloride exists.

teodor - 31-5-2021 at 00:31

I didn't make experiments with FeCl3 for a while but last few days I had enough time to investigate the behaviour of anhydrous FeCl3 in diethyl ether. The short report is: it is not easy to separate FeCl3 and the ether, it forms a kind of "stable" compound. I have no doubt we can break the bonds but the question is will we get FeCl3 + diethyl ether or FeCl2 + ethyl chloride.

Here is the story.

Thinking about this thread for a while I had an idea to check things from the other end: when we mix black anhydrous FeCl3 with dry ether and add water/HCl to this solution. I dried ether with sodium until the metal showed a shiny surface. So, I mixed ~25g of FeCl3 and 50ml of sodium-dried ether. The process is quite exothermic and I added FeCl3 in small portions because it causes explosion-like boiling of the ether. So, this fact suggests we create some kind of a bond. After realising that (and checking some publication I found by googling the topic) I did some experiments with separation.
I tried to boil-off the ether under aspiratory vacuum but it is possible to boil-off only part of that. Even keeping the flask in a 65C water bath under such kind of vacuum several hours doesn't convert the brown opaque liquid into a solid.
But putting this liquid (after boiling-of as much free ether as possible) into freezer starts polymerisation. It forms a brown jelly like that we can get from gelatin. This jelly is quite stable in air at room temperature and 5x5x5mm cube of it can resist a moisture almost a whole day - it thaws very slowly forming brown liquid which also doesn't react fast with water in the air. Only at the evening I found signs of hydration in form of yellow crystals under the layer of the brown liquid. The next day (keeping the open petri dish on a table) I found that the number of crystals was just increased.

So, I think the diethyl-ether is not a proper solvent to recrystallise anhydrous FeCl3 but can help to get a good crystals of a hydrated form.

The last photo shows what happens when we mix an etherial solution FeCl3 and a pinch of SnCl2*2H2O.

[Edited on 31-5-2021 by teodor]

maldi-tof - 8-6-2021 at 11:20

I have obtained many times FeCl3 6-hydrate using a rotary evaporator.
The synthesis is quite easy, Iron + HCl + H2O2. Then concentration with the best vacuum you can, and about 70 °C. You will see then than a thin layer of product is formed. Then stop the concentration, let it stand, and centrifugue.

If you want FeCl3 anhydrous...i can't help you, it has been impossible for me to obtain.

I know that you want to recrystallize, and I am not explaining how to do exactly that...and not even the same product maybe

teodor - 8-6-2021 at 15:20

I think it is the best answer for the original question, maldi-tof. My posts & experiments under this topic are not directly related to the original question, it is more about testing of changes in behaviour due to Fe atom coordination changes which is related to some discussion here but not to the topic. Doing experiments with FeCl3 in non-aqua solvents I got several results which raised my interest to the wider topic and I still doing & planning new experiments. But probably for me it is better to write about it in some different place.
The topic of getting anhydrous FeCl3 from solution is not only about detaching H2O, which is known to be very hard, but even detaching a solvent, which I think is something not so well known.

Chemateur80 - 21-2-2024 at 13:45

This looks like a good thread to bump when one’s interested in making some solid iron chloride.
Easier said than done, it seems, after reading about it.

I’ve made a solution of iron(III)chloride by dissolving 25g Fe-powder in an excess of HCl and then adding H₂O₂.
I evaporated the solution down a bit with some heat. That was fun. Some nasty HCl-fumes and small brown stains on my work bench from splashing.
I now have a solution containing ca 72g FeCl₃/250ml. The pH is around 1 and I can’t see any particles in it from hydrolysis.

I’ve watched the Nurdrage video on FeCl₃.
https://www.youtube.com/watch?v=43Xsh9J7S-g&ab_channel=N...
In the end he just quickly shows some really nice looking solid hexahydrate without mentioning how he managed to evaporate the solution. Too bad, considering how difficult it is.
Any ideas?

Nilered also has a video where he makes the hexahydrate and evaporates it down to a brown “cake”.
However, when he dissolves it in water, the solution is quite cloudy, so the product must have hydrolysed to some extent.
https://www.youtube.com/watch?v=BtnCynfmBnc&t=0s&ab_...

Anyway, it would be really nice to have some solid hexahydrate, or at least concentrate the solution as much as possible without too much hydrolysis.
Has anyone tried setting up a distillation apparatus and distilling off water while occasionally dripping in a litttle >25% HCl, or even slowly bubbling in Cl₂-gas, too keep the pH low and add chloride ions?

Would this work at all, or would it just be a waste of time?

Bedlasky - 21-2-2024 at 16:28

Do you want specifically FeCl3 or are you making it just because you want some soluble Fe(III) salt on hand? If the later, I recommend making NH4Fe(SO4)2.12H2O instead. It crystallize well and make beautiful violet crystals. Another interesting possibility would be (NH4)2[FeCl5]. These double salts are far less hygroscopic. From this reason I recently made Na2[BiCl5] and K2[SnCl6] instead of BiCl3/SnCl4.5H2O.

https://en.crystalls.info/Ammonium_pentachloroferrate(III)

I don't have experience with FeCl3 crystallization, but I would try evaporate it as much as you can and then place beaker with the solution in a jar with NaOH or KOH on the bottom (KOH is better in drying, but more expensive). This should absorb rest of water/HCl. Don't forget to change hydroxide from time to time.

bnull - 22-2-2024 at 06:47

You may try this one also. Evaporate the water until the decomposition begins. Then, while still hot, bubble HCl so the precipitate redissolves. Cool the solution to 0°C. The solubility of the hexahydrate in water falls from about 91.94 g/100 g at 20°C to 74.5 g/100 g at 0°C (source).

clearly_not_atara - 22-2-2024 at 10:27

I think that Bedlasky's idea is a pretty good one.

Ammonium chloroferrate was synthesized by mixing together an alcohol solution of ammonium chloride and ferric chloride hexahydrate in equivalent amounts by the reaction:

FeCl3•6H2O + 2 NH4Cl >> (NH4)2FeCl5•H2O + 5 H2O

with the product effectively crystallized by evaporation. The result can be dehydrated and then partially decomposed into NH4FeCl4, but further heating gives not FeCl3 but FeCl2 due to redox with the ammonia. The monohydrate described features octahedral iron, which is usually a sign of stability.

The reaction is a little surprising because ammonium chloride normally has quite limited solubility in alcohol, but since the alcohol is probably wet it should dissolve.

Attachment: Dyachenko2018.pdf (611kB)
This file has been downloaded 38 times

[Edited on 22-2-2024 by clearly_not_atara]

Chemateur80 - 22-2-2024 at 13:00

Thanks for the suggestions. There seems to be a few threads here about ferric ammonium sulfate, but I can’t find one about ferric ammonium chloride.

Iron salts seem tricky to make, at least if they’re not double salts. The ferrous ones get oxidized and the ferric ones hydrolyses and are hygroscopic.

I actually have around 100g of ferric nitrate nonahydrate and also a little ferric chloride hexahydrate. Didn’t make them myself though.

If double salts are the best way to get solid Fe(III)-salts I might try making them, but it would be nice to make some solid FeCl₃.

Chemateur80 - 22-2-2024 at 13:15

This is a bit off topic, but since I'm going to use my ferric chloride to try to make ferric ammonium oxalate/citrate I thought I'd ask.

As stated in a previous post, I have a solution of FeCl₃ at the moment, and I'm a bit unsure about the stoichiometry.
Should I use calculations with the molar mass of just FeCl₃, or the hexahydrate, or some sort of aquo-complex?

I believe I should use just FeCl₃, because it seems incorrect to include water of crystallization when there are no crystals, and calculating from some aquo-complex is a little ovekill.
Is this correct?

clearly_not_atara - 22-2-2024 at 19:37

The easiest thing to do is to titrate the iron content of the solution by taking an aliquot of solution (say 10 mL) neutralize with excess base, collect the iron hydroxide precipitate and heat it to 300 C (ish, some sources say lower) to decompose to mostly pure iron oxide. Then you can get the molar concentration of iron in the solution and use that to balance your equation.

This sounds like a lot of steps, but they're mostly pretty simple, because iron hydroxide precipitates readily and dehydrates at a relatively low temperature.

Texium - 23-2-2024 at 07:05

I have quite a lot of iron(III) chloride I’d be happy to sell. It was once anhydrous, though based on its appearance, is definitely not anymore.

Rainwater - 23-2-2024 at 14:29

I have been inadvertently successful in chloranating many different metals using dry ammonia chloride gas at high temperatures. Just mix everything up and heat over 450. Its really quite annoying, the only reactor material that can withstand the heat and corrosion ive found is quarts glass. But using fine mesh steel wook gives jet black solid when the gas stream is cooled. Could be of interest

Chemateur80 - 26-2-2024 at 08:44

Quote: Originally posted by clearly_not_atara

The easiest thing to do is to titrate the iron content of the solution by taking an aliquot of solution (say 10 mL) neutralize with excess base, collect the iron hydroxide precipitate and heat it to 300 C (ish, some sources say lower) to decompose to mostly pure iron oxide. Then you can get the molar concentration of iron in the solution and use that to balance your equation.

This sounds like a lot of steps, but they're mostly pretty simple, because iron hydroxide precipitates readily and dehydrates at a relatively low temperature.

Godd idea. Thank you. I'll do that.
I guess my question is how to calculate stoichiometry in general with salts in solutions, especially those that have water of crystallization.
Like in this case, for FeCl₃, is it OK to just go with FeCl₃? Or would it be more correct to do the calculations from the hexahydrate or some aquo-complex?

Quote: Originally posted by Texium

I have quite a lot of iron(III) chloride I’d be happy to sell. It was once anhydrous, though based on its appearance, is definitely not anymore.

Thanks for the offer Texium, but I think it would be too expensive shipping it from the US to Europe, with possible customs fees as well.

Quote: Originally posted by Rainwater

I have been inadvertently successful in chloranating many different metals using dry ammonia chloride gas at high temperatures. Just mix everything up and heat over 450. Its really quite annoying, the only reactor material that can withstand the heat and corrosion ive found is quarts glass. But using fine mesh steel wook gives jet black solid when the gas stream is cooled. Could be of interest

Interesting, but it's a bit too complicated with the equipment that I have.

Texium - 26-2-2024 at 09:24

Quote: Originally posted by Chemateur80

I guess my question is how to calculate stoichiometry in general with salts in solutions, especially those that have water of crystallization.
Like in this case, for FeCl₃, is it OK to just go with FeCl₃? Or would it be more correct to do the calculations from the hexahydrate or some aquo-complex?

It actually doesn’t matter. It’s a bit difficult to wrap your head around, but it’s simpler than you think. If you know how many moles of iron are in a given volume of solution, it doesn’t matter what the iron is attached to. If you made up 1000 mL of solution from one mole of anhydrous ferric chloride, and 1000 mL of solution from one mole of ferric chloride hexahydrate, both solutions will be identical. Each will contain one mole of iron. It’s all the same in solution.

The problem is when you’re uncertain of how much iron you have dissolved. This can happen easily with deliquescent salts, since the nominal hexahydrate might have absorbed extra water. Or in the case of my stuff, since it started off anhydrous, it might be a mixture of hexahydrate and lower hydrates now. In these cases, you have to analyze your solution using a method like atara suggested that will give you an iron compound with a defined stoichiometry. From that, you can calculate the true concentration of your solution.

Quote: Originally posted by Chemateur80

Thanks for the offer Texium, but I think it would be too expensive shipping it from the US to Europe, with possible customs fees as well.

Fair enough, I hadn’t noticed that you were in Europe.

[Edited on 2-26-2024 by Texium]

Chemateur80 - 27-2-2024 at 14:40

Quote: Originally posted by Texium

Alright, I understand. I'll just base my calculations on the amount of iron in solution. Thanks!