# Iron in HCl(aq)

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Everyone's having such trouble with steel wool. I'm buying some reagent grade Iron Powder very soon.

@Jowrose: I also had the most luck with iron nails/ screws. However, most of the time they have a coating or another metal (I believe mine has zinc as they were galvanized) So it would be in your best interests to throw your nail in a "prewash" until you think all of the coating has been removed. When you put in Iron in HCl, eventually you'll actually see the very light FeCl2. As YT found out however, you cannot let this tube just sit around after he reaction has taken place as the Fe2+ is slowly oxidized to Fe3+ by oxygen. Also, some weird hydrolysis reactions happen. If you let the tube sit around, first you'll notice your (pretty much) colorless solution go yellow. Next, youll see a brown kinda-hydroxide? precipitate accumulate. So, let the reaction happen for a day, then isolate your FeCl2. I got interested in FeCl2 at least a month ago but I've yet to isolate it. Havn't tried either though..

YT, Your no beginner chemist, why don't you also get some reagent grade Iron.. If you do not already have it. I don't beleive you'll ever find very pure Iron in every day items as it has little value when it comes to durability compared to steel and other alloys

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I tried using the purest sample of iron I have; the iron from a thermite reaction. I always try and save the little "nuggets" of iron that form amongst the crumby aluminum oxide, and I figured those would be the best to use in an experiment (they're a little rusty, but that shouldn't matter). So I took two test tubes, put a few grams (I don't have an accurate scale yet, but it's in the mail ) of the iron in them, and filled them halfway with HCl. I come back a few hours later, the solutions are yellow (almost urine -yellow) and there is a large mass of the brown precipitate at the bottom, exactly as in the steel wool. What is this precicipitate?! I need to get some really pure iron. This is driving me nuts.

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http://81.207.88.128/science/chem/solutions/fe.html

Look at oxidation state 3. it talks about the hydrolosis reaction that you see.

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YT, Your no beginner chemist, why don't you also get some reagent grade Iron.. If you do not already have it. I don't beleive you'll ever find very pure Iron in every day items as it has little value when it comes to durability compared to steel and other alloys

no, I dont have any in stock, for the same reason you stated further on, I simply had no need for any. as for its use or not in everyday life, that would certainly worth looking into, Ill bet theres Some practical application for it in something, its a question of finding out where and in what

naturaly the 1st thing that springs to mind is the soft iron cores used in some old transformers, and I have boxes of them, some dating back to the 1940s, Ill have a sort out when I get some time

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also heat packs have pure iron powder mixed with vermiculite. Trouble is, it oxidized pre quick (since it IS a hand warmer). Maybe its not such a great source...

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Have a look at http://www.emovendo.net for pure elements. Iron is particularly cheap in the form of small scales at very high purity. These do not as quickly rust away as powder and with patience they can be dissolved in hydrochloric acid. You also can buy iron powder, but that should be stored in a VERY well closed container, otherwise it will soon degrade to rust. The powder, used for magnetic experiments is less suitable. It contains a lot of contaminants in order to make the powder less susceptible to corrosion. Very pure iron is not an easy element on storage, unless it is in the form of small scales. These scales look as follows:

http://woelen.scheikunde.net/science/chem/compounds/iron.html

I indeed have purchased quite some metals at high purity for the special purpose of chemistry experiments (aluminium, zinc, tin, lead, iron, nickel, chromium, silver). These metals are not that expensive. The only metal, which can be found in everyday items at high purity is copper as electricity wire. All other metal items simply have too much impurities. Besides the cheap metals mentioned above, I have purchased gallium, indium, antimony, vanadium, rhenium, ruthenium, molybdenum, niobium and cobalt for chemistry experiments. Sometimes it also is worthwhile to buy salts of metals. Ceramics/pottery shops have salts or oxides of many common metals at remarkably good purity (e.g. copper, cobalt, iron, nickel, chromium, vanadium) for just a few bucks.

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Yeah, I've used Fe203 in ceramics as a stain; I was sure there was some other bonding agent to help it permanently stick to the clay in the kiln, and now I realize that it was just a mixture of iron oxide and water. I still don't understand how it permanently fuses with the clay though.

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well Ive been making Iron last night, I used some of the iron sulphate dissolved in distilled water and added a chunk of magnesium, the iron sulphate in excess.

this morning the Mg was gone and the iron particles were all nicely stuck to the magnet at the base of the flask

a few good washes in clean water and the addition of the old HCl and reaction!

Ive left it through the day, most of the irons gone, it was filtered and then used in my "Listening to Precipitation" experiment, and Ide like to add, IT WORKED!!!!

nuff said

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There! I'm glad that you got SOMETHING to disolve ;-). I also did this, although I displaced FeCl3, not FeSO4. I guess that was kind of self defeating though as the reason you reduced the iron was to make pure FeCl2 solution.. By the way, is there a reason you wanted FeCl2 other than to have it? I'm going to try to isolate it as best I can sometime soon. My only chance is to make large crystals. By the way Woelen, My CuCl2 solution is still evaporating. I added way too much water. Maybe I'll saturate it with more CuCl2 tonight..

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no reason other than to just have some in stock really, I needed some for my latest experiment also, so I figure Ill make a little extra while Im at it.

btw, youll not beleive this, but that wirewool in the 20% H2SO4 soln is still over 50% there and magnetic, there also no significant color change to the soln other than a dirty look to it (probably carbon particles).

the only problem I can see with trying to determine the nickel content (if present) is that Ni is also magnetic, And gives a Greenish sulphate only a little deeper than iron sulphate does, I think Ill just keep this wire wool out the lab and with the engines instead

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Ok, I think I got a relatively pure sample of iron, and I've been using it to make FeCl3. I took some of my old FeCl3 (that's probably not what it is, it's so impure, but I'm calling it that anyways) and added little bunches of aluminum to it, keeping a magnet next to the test tube to attract all the iron particles to the side (thanks for that idea, YT, it worked). When all the FeCl3 was reacted, I dumped out the liquid and washed the iron particles several times. I suppose some of this could have been Nickel, but for optimism's sake I'm going to say it was pure iron.

I took a small sample of this iron and mixed it with some HCl in a test tube. This time I got a very clear solution, almost exactly the same color as Woelen shows on his website (with just a tint of green that is very hard to distinguish). I let this be, and after a few hours the top cm of the solution had oxidized to FeCl3 (a yellow-orange color).

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Excellent!

Ive gone a slightly different route for my Iron sample (the one I wish to use for Keeping in stock).

I have some iron sulphate (a little impure as its agricultural) this was then displaced with Magnesium chunks and a magnetron magnet at the base of the flask, after a day or so and occasional stirring, the liquid was poured out and replaced with water and washed 3 times, magnet in place all the time the liquid was being poured out.

the large chunk of unreacted Mg was removed, a litte water added to the flask again enough to cover the iron metal, and NaOH was added to react with any stray Mg metal or possible Alu in the metal.

after 2 hours kept warm, this was washed again several times in water and Dilute HCl added (roughly 5%) and left until there was no Fizz activity (about an hour).

this was then washed again with plain water 2 times and then kept under de-ionised water.

thats exactly where I am now, hopefully pure Iron under de-ionised water that will stay that way until I decide my next step

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...the only problem I can see with trying to determine the nickel content (if present) ....

A very sensitive reaction, which only is positive for nickel and not for iron, copper, cobalt, chromium and most other metals is the following:

Add an excess amount of a solution of NaOH to the liquid. This gives a precipitate. Then add a solution of a persulfate, such as Na2S2O8 or (NH4)2S2O8 or K2S2O8 (Feinätzkristalle) to the precipitate. If it contains even a small amount of nickel, then it turns black like carbon within a few seconds.

These persulfates can be purchased at electronics parts stores as printed circuit board etchant. They are very nice oxidizers in aqueous solution, however, not suitable for pyrotechnics.

I took a small sample of this iron and mixed it with some HCl in a test tube. This time I got a very clear solution, almost exactly the same color as Woelen shows on his website (with just a tint of green that is very hard to distinguish). I let this be, and after a few hours the top cm of the solution had oxidized to FeCl3 (a yellow-orange color).

This is a very nice result with the crude starting materials you have used. If the solution was very pale green, then there was no strong nickel contamination. Nickelic solutions have a much more intense green color.

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I can't recall offhand if the Ni needs to be in a metallic form or if it will work with ions, but bubbling some carbon monoxide gas would result in the production of Ni(CO)4 which will move out of solution and leave the iron behind. Problem is, Nickel carbonyl is VICIOUSLY lethal.

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Yes, woelen, the solution was a very pale green, nothing like the previous solutions made from steel wool. I think the main reason for this was that I used FeCl3 that I had obtained from steel nails (rather than steel wool) and had filtered out the solid carbon. These must be a whole heck of a lot more pure than the steel wool.

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• 1 year later...

Bumpity bump bump...

a Friend of mine asked if I had any Fe3O4 and Fe2O3, I said I have iron oxide in the Lab, Ill have a look and give him a sample (he does model trains and wanted some as paint).

I have Fe2O3, but no Fe3O4.

anyone know a Good and Reliable synth for this?

Ill be extracting it with a magnet so impurities arent THAT critical.

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You could try dissolving a mix of an iron (II) salt and an iorn (III) salt (e.g. iron (III) chloride and iron (II) sulphate). Then add sodium hydroxide to this. A dark precipitate of hydrous Fe3O4 is formed. Then boil the liquid for a while in order to make the precipitate more solid and easier to separate.

The isolated precipitate should be washed with water and then heated to dryness.

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• 6 months later...

What happens to a dilute solution fo FeCl4- if you add concentrated H2SO4? Am I correct in believing that the iron would not be seperated from the tetrachloride because it is such a strong bond (basically nothing would happen)? Or, would The iron react with the SO4 forming a ferric sulfate?

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• 2 years later...

I am not a professional chemist, but I believe the reason some people aren't getting the reaction creating FeCl<sub>2 is because these days, steel wool has a lot of other things in it other than iron and carbon. other compounds are added to make it more flexible, durable, and rust-proof.

I put a 1x1x3" wad of medium-fineness steel wool in a jar and covered it with 31.45% HCl. For a day or so nothing happened, but when i checked back on it nine days later, the steel wool had completely dissolved. There were a lot of incredibly fine black particles on the bottom, and when the jar was shaken, they stayed suspended for 2 days. I did put a metal screw-on lid on the jar, and after 9 days it was bent outward quite a bit due to the pressure of the hydrogen that was released from the HCl when the chlorine bonded with the iron.

aside from the black particles, which i believe to be carbon and other impurities, the solution was completely clear, no green, brown or yellow. Any reasons/thoughts?

The Golden Book of Chemistry describes this reaction a bit in its instructions on creating ferrous, and then ferric salt on page 68. The Golden Book can be downloaded legally and for free from this website:

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