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teodor
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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.
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Chemateur80
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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?
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Bedlasky
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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.
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bnull
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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).
Quod scripsi, scripsi.
B. N. Ull
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clearly_not_atara
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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]
[Edited on 04-20-1969 by clearly_not_atara]
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Chemateur80
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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.
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Chemateur80
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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?
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clearly_not_atara
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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.
[Edited on 04-20-1969 by clearly_not_atara]
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Texium
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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.
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Rainwater
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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
"You can't do that" - challenge accepted
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Chemateur80
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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.
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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.
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Texium
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| 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.
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Fair enough, I hadn’t noticed that you were in Europe.
[Edited on 2-26-2024 by Texium]
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Chemateur80
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| Quote: Originally posted by Texium | 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.
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Alright, I understand. I'll just base my calculations on the amount of iron in solution. Thanks!
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