# Iron in HCl(aq)

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well heres an unexpected surprise' date=' my iron in HCl has now developed a distinct orange band just above the liquid level, strongly resembing Rust, the liquid itself is still a green color, but has some orange tinge in it now also.

I think its perhaps gone Too far and started breaking down somehow?[/quote']

On Woelen's site (Don't know it offhand, just check his profile (you probably already have it anyhow)) on the Iron solutions page, he talks about Fe3+ hydrolysing into Hydroxide and other weird things that can appear brown. By this point, you have mostly Fe2+, but it's still possible that that is what has happened.

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Im really disapointed in it now, I think Ill filter it off and displace the metal out with Magnesium in a magnetic bottle, I should end up with a reasonably pure Iron sample after, and enough to experiment further with, so its not All a waste

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Good Call on the Crystals, that'd make sense. I still have some liquid left in my tub of CuCl2. I know that it becomes pure because I set aside smaller amounts to crystalize and they showed promise. It may be another week before my large tub evaporates, but I will let you know when this happens.

With the CuCl2 crystals, do not let all liquid evaporate on your nice and clean crystals. Crap and impurities build up in the liquor above the crystals, so when you let is all evaporate, your final part will be quite impure and it would be sad if that impure stuff mixes up with your pure crystals. So, I would advice to decant the liquid from your nice crystals and to have your nice crystals drying (that may be quite hard actually, CuCl2.2H2O is quite hygroscopic).

The liquid can be put aside for other crystals, but these will be less pure and should not be mixed with your purer crystals.

Drying CuCl2.2H2O can be done by putting your crystals in a watch glass or something like that, putting drying agent (anhydrous CaCl2, available at hardware stores) besides your crystals in another watch glass and put both these watch glasses in a closed vessel (a well-closed tupperware box is suitable very well for this purpose). The CaCl2 will dehydrate air in the box and due to the dry air, the CuCl2.2H2O will get dry, without loosing its water of crystallization. Within a few days you should get reasonably dry crystals.

I threw a screw that I had in acid already prevoiusly to get rid of any coating. After Putting in a test tube with HCl, The Fe2O3 layer melted off (this was actually kind of neat) and the solution turned yellow. All at once it bubbled decently quickly and eventually the solution turned clear and then light green. I left it to react over night and even though the reaction vessel was loosely capped, the solution is now back to yellow. I also have a decent amount of whitish precipitate on the bottom, but no floaty crap! This seems to be a lot more pure than the scrubby stuff. So it looks like the white stuff is my FeCl2.

Have a look at this MSDS: http://ptcl.chem.ox.ac.uk/MSDS/IR/iron_II_chloride_hydrate.html

Ferrous chloride is light blue/green, not white. If you want to be sure that your white crystals are ferrous chloride, then you could decant the liquid above them and add the crystal mass to some acetone, shake a little and then decant the acetone. Repeat this another time. In this way, you rinse off the water quickly and you make the crystals dry. Acetone is very volatile and your crystal mass is dry in no-time when you put the solid mass, wet from acetone, on a central heater radiator. Rinsing with clean water is not wise to do, because a lot of your crystalline mass may dissolve again.

Before you put all your crystals in acetone, first try with a small amount, just to be sure that this does not spoil things. Btw, you should not use the acetone trick with your CuCl2 crystals. CuCl2 dissolves in acetone and forms a yellow/brown solution with a greenish hue!

As a side note, it amazes me at the difference in composition of two "iron"s It would make sense, because the scrubby iron needs to be very bendable without breaking. Iron itself would not fulfill this purpose. So the dark green liquid, white floaty precipitate are apparently added metals to increase tensile strength.

it is hard to say what all that added crap is. Indeed, usually a lot of other metals and sometimes carbon and even tellurium are added to the metal.

Also, is my precipitate of FeCl2 decently pure? You talk about Fe3+ hydrolysing on your website. Your website gives me the impression that FeCl2 is decently Stable under neutral conditions. If I washed the precipitate in a liquid insoluble by FeCl2 and then added distilled water to disolve the FeCl2, colul I then crystalize it leaving the other crap behind? Or would this ALSO oxidize.

If your precipitate is really crystalline (hard pieces with shiny edges), then it is decently pure, certainly enough for the kind of experiments you as a home chemist intend to do with it. If the precipitate is more like a paste and more flocculent, then I doubt whether it is FeCl2 and then I think it is sheer crap.

In general, flocculent precipitates are VERY impure, due to a process called co-precipitation.

Just as an example, suppose you have a solution of CuSO4 in water (light blue) and you add solution of NaOH, then you get a blue flocculent precipitate. On high-school one learns that this precipitate is Cu(OH)2. In reality, things are MUCH more complex. The precipitate also will contain SO4(2-) ions, Na(+) ions and who knows what more (oxo-species, water molecules). The physical processes in which such flocculent precipitates are formed, almost certainly cause other ions to be trapped (clathrated) into networks of -Cu-(μ-OH)-Cu- bridges.

When a crystalline precipitate is formed (usually these are formed slowly), then the processes are quite different. Slowly, ions (or molecules) are deposited on the crystal and the deposition is such that the structure is nicely regular. That is why recrystallization can be used to make purer compounds. Flocculent precipitates make things impure, crystalline precipitates make things purer.

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@YT. I will do the same with some of my sample. I have magnesium ribbon finally and this seems ot be a decent way of getting pure iron.

@Woelen: Good idea with the decanting the solution of my CuCl2. Unfortunatly, I still have a good 50 mL un evaporated, so I'll wait a few days and then decant.

I don't have acetone at the moment so I will do that later

I beleive my precipitate is crystaline, when shaken, al precipitate falls back down and makes a clear yelow liquid in less than 30 seconds. This leads me ot beleive that there is a large partical size (also looks that way) and therefore it must be crystaline.

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I attached a picture of my final result. It nicely shows the green solution.

I was really startled to see what happened with the liquid in a period of 2 days. A beautiful blue/green crystalline mass had settled at the bottom of the test tube. The liquid above it was green with a yellowish hue. Apparently, still some air was trapped in the test tube and some oxygen made the liquid yellowish.

I decanted the liquid and quickly rinsed the nice crystalline mass with some acetone. This preserved the crystal, but the color of the crystal quickly changed from the nice blue/green color to a purer green color, with a yellowish/brown hue . This is due to aerial oxidation, Fe(2+) being converted to Fe(3+).

I made some pictures of the crystalline mass. It looks really neat. The diameter of the piece is approximately 1 cm. You can still nicely see the round bottom part of the crystalline mass, due to the round bottom of the test tube.

In the first picture I also have drawn a small rectangle in the picture. That is close to the original color of the crystalline mass, before I took it out of the air-tight test tube.

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That is VERY neat. How did you get your crystals in such a mass? Mine are about 1/2 the size of standard NaCl crustals. They are very fine. (I'll have my CuCl2 for you to look at next week (My room is colder than I'd like for evaporation purposes, but nicer crystals can be made))

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I did not anything special at all. I dissolved the iron in the acid. In order to have it all dissolved in an acceptable time, I heated close to boiling. When (almost) all iron was dissolved, I stopped heating and put the test tube aside. Two days later, the large crystalline mass was at the bottom of the test tube.

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Well, my crystals are done drying. I coudln't get a clear picture like you and I was using a Minolta camera.. Anyway, you can kind of make out the crystal structure here. The color of the crystals is a light blue, it's a little darker in the picture

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Well done Xeluc ! You have done a VERY good job of making such pure CuCl2.2H2O. The color looks very good and it is an indication of high purity. This is very nice stuff for further experiments.

What is the length of one such a needle-like crystal?

Now there is one thing left. Make a better picture with your Minolta camera .

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Haha.. The largest crystals are ~1mm thick and 1-2 cm long.

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Wow, that really is great. The reagent grade CuCl2.2H2O I have has tiny crystals, just around 1 mm of length and 0.1 mm thickness or so.

Doesn't your camera have a macro-feature? The crystal of FeCl2.4H2O I had has a maximum diameter of approximately 1 cm and I could obtain a sharp picture of that, so with so many crystals of the mentioned size you should be able make a real stunning picture.

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Well, I now have my crystals in a holding container douh. Most of the crystal structure is gone but I DO plan on making more of these crystals. CuCl2 seems to be a VERY often used chemical for me. I'm glad to have found a way to make pure crystals of it. Yes, the camera has a macro feature but I never thought to use it. Obviously the picture would have turned out better. These crystals realyl WERE stunning. Sadly, most of them are broken now but after school tomorrow I'll definastly take a better picture of what I have for you. If your still interested in a week or so, I will do another batch. I encourage you to do this yourself though, The structure is pretty neat.

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Yes, Xeluc, I am interested in this. If you can make new pictures after a week or even later, then still this would be great. Such things are very nice and it is even better if you show them to the world .

I also intend to do something like this myself with CuCl2, although I already have several hundreds of grams of this compound. I did not know that such nice large crystals could be obtained.

Another very interesting suggestion for you. If you have a heat resistant glass beaker (e.g. Schott Duran or Pyrex glass) or a corrosion resistant crucible, then it may be very interesting for you to carefully heat the CuCl2.2H2O crystals, such that they loose their water. If the heating is done carefully, then they do not hydrolyse and loose HCl, but then you get pure anhydrous CuCl2. That latter stuff opens up a whole new world of copper chemistry in non-aqueous solvents. I just started exploring that. CuCl2 is mostly covalent and it dissolves in methanol, ethanol, acetone, DMSO. It dissolves with all kinds of colors (green, yellow/green, brown, yellow, blue). You even can get very weird precipitation reactions, such as CuCl2 + H2SO4 --> CuSO4(s) + 2HCl, where white CuSO4 precipitates. Also complex formation is very different in non-aqueous solvents.

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I did not anything special at all. I dissolved the iron in the acid. In order to have it all dissolved in an acceptable time, I heated close to boiling. When (almost) all iron was dissolved, I stopped heating and put the test tube aside. Two days later, the large crystalline mass was at the bottom of the test tube.

well I finaly filtered my iron in HCl and ended up with a green liquid and rust covered flask (a few drops of H2SO4 soon cleaned that up) I left this liquid on the heater overnight and got a mat of small crystal (about 2mms) in a dark brown/yellow Not clear soln and more rust stains, needless to say I got rid of the lot, I wasnt at all happy with the results.

Im going to try the above method, and test my luck that way

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@YT: Eww, looks liek it all oxidized. I tried to crystalize some FeCl2 but I added too much water and after heating, all of the FeCl2 stayes in solution. Of course now its yellow FeCl2-. I have noticed that the nasty brown stuff accumulates very slowly AND it is easy to get out of solution. I have found that shaking the lot a whoel lot and waiting for the FeCl2 to fall down, but the brown hydroxide, or whatever complex chemical Woelen called it stays in the water for a much much longer time. Just shake it up and decant, I forget is FeCl2 is soluble in acitone (I do not beleive so) but adding some of that in and shaking will get the rest of the brown stuff out. I Just disolved off the (probably) zinc coating from a screw and threw it in HCl. I could find little impurities (certainly nothing I'd care about). Eventually (as you know) Crystals form. I've done all of this without filtering as the screw is large enough to be removed form the test tube. Man, I got to stop with the parenthesis. I wouldn't leave the solution in a heater.. It's really hard to evaporate the FeCl2 solution in a heather without oxidizing it, although I'm sure it's possible. I was going to do something of the liek with CuCl but I havn't gotten a chance to yet as I do not have a suitable container. Just heat it in a test tube to disolve crystals and they should reform larger. I just thought to myself that I am being a little elementary here. So sorry if I have insulted your intelligence YT ;-)

@Woelen: Well, I have some more solution in a petri dish to oxidize. I'm doing it in a smaller vessle so I don't have to use as much. I will include a "size-key" or whatever so that you can see approximate crystal sizes. Right now the solution is brown (I poured a solution of CuCl2- into a dish) In a day or so it should turn green. I will let you know when I'm finished. The very large color change of anhydroud CuCl2 as oposed to hydrated was actually one of the reasons that CuCl2 was the first chemical I decided to synthesize. I have done this very crudely before. The first time I made it last sumer and I whined becase it was a light green color, not blue, was because I boiled the solution . I got anhydrous CuCl2 alright. Probably lots of Hypochlorite too... Anyhow, it never occured to me to do experiemtns with anhydrous CuCl2. I'll get on that.

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• 4 weeks later...

Next. I have no idea what you guys are talking about with your Iron not disolving at any appreciable rate. I figured steel wool wasnt tremendously pure' date=' so i went and bought some of those steel scratch pads made from the coiled steel. Pop one of those in HCl and it bubbles very vigorously. (31% HCl).

[/quote']

what steel wool did you use? Ive just aquired some Grade 1-2 steel wool, its very fine, and tried to replicate your findings.

theres no reaction at all happening, I have half a gram of this wire wool, uncompressed in the bottom of a test tube and have added 10ml of 30% HCl to it, 20 mins later still no reaction?

any idea whats up?

edit: the 3rd one down on this site http://www.hardware-ironmongers.com/results.cfm?rand=0.6301274569848134&ct_ref=1260 is the one Im using.

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I've been messing around with metals in HCl (mainly because HCl is the only acid I have right now...) and one of my endeavors has been to create a solid form of Iron Chloride.

So here's what I did:

First I took some steel wool and put it in a beaker of ~30% HCl. Immediately after placing the metal in the solution, there was bubbling and it began to turn a greenish hue. After leaving it outside (I don't like the fumes) overnight, it looked like this:

So I removed the liquid (leaving the black gunk (which I guess is just the carbon from the steel wool) and the grey precipitate) and placed it in a reagent bottle. After a day or so it had changed to a yellow-brown color.

After obtaining this solution, I tried to create a solid by evaporating off the water. I put some solution in a 100 mL beaker, put it on a hot plate, and kept the temperature just below boiling point. As the water evaporated, yellow crystals began to form around the water line of the beaker. Eventually, I was left with a brown paste at the bottom, and I removed the beaker from the hot plate. As I stirred this brown paste, it turned a bright yellow color and became much finer particles. This solid is, I'm assuming, iron chloride hexahydrate. I'm basing this assumption solely off the fact that my solid almost exactly like pictures of pure iron chloride hexahydrate I've seen. So I put some of this product in a test tube, sealed it with a stopper, and left it overnight. When I came back, the solid had turned a light pale green color.

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Ok, I just tried boiling down some of the grey precipitate from the first picture (the iron-HCl reaction). I got the exact same product as my light green solid; but this time there was no intermediary yellow solid. Instead of turning into a brown mush and slowly turning yellow, it turned to yellow and then the pale light green. What is this final product? Is it a different hydrate of iron chloride? Or a different compound altogether? For that matter, what is the precicipitate of the steel wool reaction?

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At YT: I jsut used some regular steel wool bought at a hardware store..... I found out that my steel wool was much more pure than the scratch pads. The scratch pad was an alloy to make the steel more bendable. Anyhow. I can't imagine why your steel wool isn't bubbling away. It reminds me of your troubles wiht that piece of aluminum foil. Maybe things jsut don't want to disolve for you . When I procure more HCl I would be happy to take a picture for you of the reaction taking place..

@ Jowrose: Your findings are.... weird. When you first disolve iron in HCl, you would expect a near-colorless solution, because Fe2+ ions are colorless.... Well a LITTLE green, but not as green as your first solution. Also, is there like an inch layer of white prec. in your first jar? If so, your using horribly impure iron. When I disolved my scratchpads in HCl I had similar findings.. My solution turned dark green and I had a white precipitate. Pure Iron in HCl will not make such a green color nor will it create that precipitate. Lets look at your second picture. As I would expect, the solution would turn yellow... or brown I guess... As Fe2+ oxidized with Oxygen to Fe3+. But then we look at your evaporate solid. If you slowly evaporated a solution of light green FeCl2, you would end up with Yellow/brown FeCl3 as the Fe2+ would all be oxidized. However, you have a light green solid and this REALLY puzzles me. Maybe Woelen could come over and help... I would have to say that your iron is very impure though.

Also, you said you left your solution out overnight. I would expect heavy oxidation (meaning no more green FeCl2) after that long. Weird..

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As I left the green solution overnight, i think there must have been steel wool still undergoing a reaction. Yeah, it must be horribly impure. But what's odd is that when i boiled down the precicipitate, I got a solid of almost identical color and texture. I am confused.

I have tried doing the experiment with steel nails, and the resulting solutions are much more yellow (so I'm assuming it's far more pure than the steel wool) but I still get the orange color from oxidation. Right now I have maybe 30 mL of this "nail-obtained" FeCl3 sitting in a dessicating box thing (with some calcium chloride dessicant). It's been there about a week, I don't know how much longer it will take.

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well youre not going to beleive this, but after leaving it overnight, the wool Under the acid level remains unchanged, everything Above it, is rusty to the point of crumbling appart, and theres no green liquid at all?

Ive taken out this sample and washed it, then added it to 20% sulphuric acid, its now gently bubbling away quite happily, so the metal stock this wire wool came from seems to be perfectly ok.

my HCl checks out just fine too.

I have absolutely NO idea at all whats going on here, it defies logic!

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I've read all your observations and I also have a hard time to explain all these things. I really urge you to try to find sources of pure iron. These scratchpads and steelwool things and so on are so increadibly impure (at least where I live). They contain other metals (sometimes even small amounts of tin) and they contain a lot of grease-like stuff and detergents. Not good at all to do experiments with them.

In acidic environments these detergents are converted to the corresponding organic insoluble acids (YT, remember what happens when acid is added to sodium benzoate, a similar things happens with the detergents in soaps). Other metallic impurities may give rise to all kinds of colors (dark green for chromium, green for nickel, white flocculent precipitate for tin).

@jowrose: Take some of your green powder and add some of this to water and shake well. Do you get a clear solution? Add some HCl. Does this make the solution clear? If acidification results in a clear solution, then you don't have detergent impurities nor tin impurities.

Also another word of warning: I received a bottle of 50 grams of reagent grade FeCl2 (anhydrous) from an old German lab. This bottle was still sealed, but it was almost 15 years old (it was prepared in the former GDR, just before "Die Wende", 1989). When I opened the bottle, the contents is totally brown/yellow and it does not dissolve clear anymore. It also does not reduce acidified dichromate anymore. So, all FeCl2 is oxidized in these 15 years, even in the sealed bottle . So, if even such a pure reagent grade compound is oxidized in a sealed bottle, then you can imagine how hard it will be to make pure FeCl2, which keeps for a reasonable time. The stuff I have can best be regarded as basic ferric chloride.

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I tried taking a few grams of the green solid and putting it in a test tube with water. It remained a light green/yellow color, so I added some HCl, and the color remained pretty much the same. I am going to search around for a pure source of iron, this reaction interests me more and more. I do have a purer (I hope) solution of iron chloride in the dessicating box. I used steel nails instead of steel wool; hopefully there was less excess stuff in the nails.

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Nickel metal is a common additive to iron to increase its workability and durability. The fact that you are seeing a greenish solution makes me think that there may be some nickel in there.

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well I just came into the lab this morning for my wake-up ritual (coffee and a cig) and the wire wool`s still there in the sulphuric acid, surely at 20% and being so fine a wire it should be all gone by now!?

so I suspect it is the wire wool also at fault here, either that or something VERY odd about things not wanting to corrode or react in this place is happening LOL

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