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Author: Subject: Preparation of FeO homestyle.

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[*] posted on 20-10-2003 at 15:49
Preparation of FeO homestyle.

I am not sure if anybody would need FeO anytime soon, but I accidently stumbled upon a fast way to make some high purity FeO when I was making Fe2O3 and Fe3O4 for thermite. As usual, I went out and bought a box of steel wool which I dumped into a bucket of water. At this point you let it sit and normally use techniques of siphoning water and seperating the crap on the bottom out until which will make a very nice Fe2O3 powder. Something got me to thinking, why not do a little test with NaOH to see if I can oxidize Iron faster, it works with Al. (Of course I wasn't considering the reasons it helps oxidize Al are irrelevant to Fe.) The result, it did not make Fe2O3 form faster and actually inhibited it production. (To make the reaction for rust go faster one, one should use a little acid or maybe some salt.) Something was collecting at the bottom however and it seemed very fine. Thinking that perhaps I had found a way to particilize iron, I seperated off the solution and the unreacted steel wool. After isolating the crud on the bottom I checked it over... It was black and very fine, like 400 Al mesh. To decide if it was iron, I put a flame to it. In a pile, it didn't real light on fire so much as it immediately became the nice powderized Fe2O3. If you dump it on a flame it looks like dumping flint dust over a bunsen burner. Anyhow, when you burn iron it usually will make oxides with higher oxidation states. If you burn steel wool for example, you will notice alot of bluish remains but not any visible Fe2O3. Because of this and the fact that the color was just too dark to be iron, I decided to research the subject more. I played around on paper for a bit and finally decided I needed a mechanism for iron oxidation. I also needed to learn what other oxides or products from my reaction might be possible. Two oxides of iron were black, FeO and Fe2O3. Iron Hydroxides are typically greenish and as the steel wool remained in tack this seemed unlikely. I learned that in the mechanism of iron rusting in water, iron must first attain a +2 state BEFORE it can go on to a +3 state. Alot of this has to do with hydroxide/hydronium ion concentration. In my reading I learned that the extra OH- in solution will impair oxidation to the +3 state. However, you would still have the Fe 2+ ion in solution. Ultimately, long story short is that when you evaporate the solution Fe 2+ pairs up with an oxygen and leaves you FeO. I should also mention that I did try to test if FeO was magnetic. It was slightly, but not as appreciable as Fe2O3. Remember, natural magnetism has a great deal to do with unpaired electron charges. Generally, the more unpaired electrons you have, the more magnetic it is. Fe 3+ is particularly magnetic as it has a half filled d-shell with all electrons unpaired. Fe 2+ however has only 4 unpaired electrons. Ultimately, I decided that the fact that it readily converts to Fe2O3 and is less magnetic than rust made it unlikely that this could be Fe3O4. However, to be sure one might want to run density tests, maybe some spectroscopy - Not stuff I can do with equipment from my home. If it is FeO, it seems to be of very fine particle sizes and depending on how you isolate it, mostly pure. Maybe someone out there can make use of this? Also apparently, water + iron + gluconic acid is a great way to make sustainable hydrogen for a fuel cell.
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Now celebrating 18 years of madness

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[*] posted on 20-10-2003 at 18:38
editorial note

In the future, please post your messages in the appropriate area (this was moved from Energetic Materials) and break the text into shorter, more readable chunks.
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Biochemicus Energeticus

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[*] posted on 21-10-2003 at 11:35

In order to be really sure that you got FeO, and not Fe3O4 or whatever, you could try and convert the oxide to a salt, like FeCl2. THis would have to be done with degassed water/acid however, and no other catalysing ions (distilled water). For one thing, FeCl2 is green (if I remember correctly, just like the vitriol FeSO4*6H20), while FeCl3 is dark yellow when dry, and reddish yellow in solution.
In addition, you could try and make the hexacyanoferrate. This is created simply by the addition of excess KCN. Yellow hexacyanoferrate corresponds to Fe2+, while deep dark red corresponds to the Fe3+ version of the hexacyanoferrate.

Yet another one is to add to your FeClx solution hydroxide (again in distilled degassed H2o). This will produce a green precipitate of Fe(OH)2, which will oxidise slowly via black (Fe3O4*xH2o) to brown (Fe(OH)3) ) at air contact. For instance, you could take your Fe(OH)2 precipitate, filter it and leave it on the filterpaper. ..good luck :)
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