mrjeffy321
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Black Iron Oxide Differentiation
There are two types of black Iron Oxides, Iron (II) Oxide, FeO, and Iron (II, III) Oxide, Fe3O4.
I am trying to tell whether this black Iron Oxide powder I have is FeO or Fe3O4.
I know my black Iron Oxide is very magnetic. I know that Fe3O4 is magnetic, but I am not sure if FeO is or not, so I am not sure if this property can
be used to differentiate between the two.
I am trying to find some easy test I can perform on the powder. Since FeO is made up of just 1 Fe+2 ion and Fe3O4 is really 1 Fe+2 and 2 Fe+3 ions, I
figured that I might be able to perform some type of chemical reaction if I dissolved the oxide into a solution. Perhaps react the Iron Oxide with
Hydrochloric acid to produce soluble Iron Chloride (FeCl2 and/or FeCl3) and then precipitate out either an Fe+2 or Fe+2 salt. But the problem is,
most Iron (II) salts have the same or similar solubility as Iron (III) salts.
Any suggestions?
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12AX7
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Color? Not real great as iron solutions oxidize slightly in air anyway...
It's a pretty good bet you don't have FeO, since it disproportionates between 300 and 560C. Though if you bought it, it could be. Brauer doesn't
mention magnetism on FeO but does for Fe3O4, so it seems a good bet.
Tim
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mrjeffy321
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The dry Oxide powders are both black and I think Fe+2 and Fe+3 solutions are just slightly different shades of brown, Fe+3 is a dark brown and dilute
Fe+2 is a yellowish color that will probably turn brown if it is concentrated.
I did purchase this black Iron Oxide (as opposed to making it myself). It was sold as simply, "black Iron Oxide". I strongly suspect it is Fe3O4,
but at the same time, I dont have any evidence to lead me to believe it definatly isnt FeO.
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kyro8008
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If you have the equipment and experience then why not do a redox tiration?
Dissolve for example 0.5g of the oxide in excess dilute sulphuric acid.
If its FeO, then: FeO + H2SO4 --> FeSO4 + H2O
If its Fe3O4, then: Fe3O4 + 4H2SO4 --> FeSO4 + Fe2(SO4)3 + 4H2O
Thus, for 0.5g if FeO then: the solution after dissolving FeO in the acid will contain 0.006963788 moles of Fe2+.
And for 0.5g if Fe3O4 then: the solution after dissolving Fe3O4 in the acid will contain 0.00216076 moles of Fe2+.
Make sure there is a good excess of the sulphuric acid in the solution and titrate it with standardised potassium permanganate. This will oxidise the
Fe2+ to Fe3+.
From: MnO4- + 8H+ + 5Fe2+ --> Mn2+ + 5Fe3+ + 4H2O
Then if FeO, 0.001392757 moles KMnO4 will be needed to fully oxidise the Fe2+.
Or if Fe3O4, 0.000432152 moles KMnO4 will be needed to fully oxidise the Fe2+.
From the results, even with errors, this should work for this and give a good indication of what you have. Obviously you will need to work out
relative concentrations and amounts for reagents, this is just an example.
Btw, do not use hydrochloric acid to dissolve the oxides if you do the titration and permanganate will oxide it to chlorine.
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neutrino
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How would you tell the endpoint of titration? Is it just when the intense permanganate color disappears?
Iron (II) ions are bluish green. Iron (III) ions are almost colorless. The brown color we associate with iron (III) is from the color of the common
salt iron (III) chloride, which contains the brown FeCl<sub>4</sub><sup>-</sup> complex.
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mrjeffy321
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kyro8008, I do not have the equiptment or materials needed to do that titration, but your method seems to be a very good option. After all, if I
would need almost 3 times the amount of Potassium Permanganate to react with the Iron Oxide, then that is a pretty clear cut way to tell.
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Iron (II) ions are bluish green. Iron (III) ions are almost colorless.
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I was basing my solution colors on Iron Chloride solutions I have seen. Obviously, Iron (III) Chloride is dark brown, but an Iron (II) Chloride
soltion isnt anywhere close to being bluish green, it is a [dark] yellowish color.
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12AX7
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The FeCl2 solutions I've prepared are green (I think  ). They turn yellow and
brown in air.
IIRC, Woelen says FeCl2 is essentially colorless, when air is excluded.
I find the crystals (hexahydrate I think) are a light bluish green, turning drab brown on the surface (due to oxidation to ferric oxide and chloride).
For the titration, you'd find the endpoint when the purple color doesn't disappear, no? As long as there's Fe(II), it'll "bleach" the solution and
remove the intense purple. When Fe(II) is gone, MnO4- will remain, coloring the solution strongly. (Ain't it great when your titrant is its own
indicator?)
Tim
[Edited on 3-31-2006 by 12AX7]
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kyro8008
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Thats right; the permanganate is in the burette and as it hits the solution when Fe2+ is still around it turns colorless (which looks weird!).
When the endpoint has been reached, the permanganate will remain purple when entering the solution, and is deemed finished when the solution just
turns pink.
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mrjeffy321
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Alright, the procedure looks good,
all I need to do now is get a burret, some KMnO4, and some Sulfuric Acid.
I went to the pet store (Pet Smart) today and read the backs of all the bottles hoping they carried some fish products with KMnO4 in it, sadly, I
couldnt find any.
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