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| | From:  MikeKL5 (Original Message) | Sent: 10/16/2004 6:18 PM |
Hey Steve,
I have a question for you regarding a compound we synthesized in Advanced Synthesis class. The compound that we made is (+) and (-)-Tris (o-Phenanthroline) Iron(II) Perchlorate Trihydrate. The formula for the compound is [Fe(o-phen)3](ClO4)2*4H2O, and the compound is a dark, blood-red color. The synthesis came from William Jolly's "The Synthesis and Characterization of Inorganic Compounds" Prospect Heights, Illinois: Waveland Press, Inc; 1991. page 469-470.
This is the most obscure compound that I've ever come across in any kind of lab class. lol. I have to write up a synthesis report in a week, and I can't find much of anything about this compound. I can't even find a standard IR spectra online for this stuff. Have you by some miracle encountered this compound before? If not, do you know of any other good sites / online databases that I can check. I've looked through the SDBS spectral database, and I couldn't find anything about it. Our professor gave us a standard 1H-NMR for the compound, but he doesn't give out IR specs.
Any help at all would be deeply appreciated.
Thanks in advance,
MikeKL5 |
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| | | From: ·Steve· | Sent: 10/16/2004 6:50 PM |
Hi Mike, good to hear from you! I'll do some digging and see if I have anything more on your compound. Somewhere around here I've even got an old pre-1991 edition of Jolly too - it's been a standard text for I don't know how long.
One of our General Chem II labs is "Colorimetric Determination of Iron", in which the students take a sample containing an unknown amount of iron(III), reduce it to iron(II) with hydroxylamine, and add excess o-phenanthroline to form the dark-red product. This absorbs very strongly at 510 nm, which they measure with a Spec 20, and with a Beer's Law plot they then determine the amount of iron in their sample. This is a good experiment except that the o-phenanthroline is costly.
L8R!
Steve
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| | | From: MikeKL5 | Sent: 10/20/2004 2:34 AM |
Hey Steve,
Sorry that it took me so long to get back to you. I appreciate you
being able to tell me anything about this stuff at all. I came across
another source which mentioned a similar complex being used for
spectroscopic analysis of trace amounts of iron ion. That's about the only
interesting thing that I could find out about it, besides the fact that it
doesn't melt, but rather explosively decomposes at ~255 C. I had fun
figuring that one out while doing the melting point. lol.
Anyway, thanks again for your input. It definately helped.
Take care bud,
MikeKL5
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Reply
| | | From: ·Steve· | Sent: 10/20/2004 3:15 AM |
Hi Mike, sorry I haven't had time to search my "archives" for an IR of this compound - you may have to go to the original journal literature, such as Inorganic Chemistry from around the 1960's or earlier, via a Chem Abstracts search, to find a published spectrum. I guess your Jolly text does not have a literature citation. Yes, perchlorates can keep you on your toes! I've seen a perchlorate explosion or two myself; resulting from simply scraping compound from the flask - dangerous stuff. Did your text say why the perchlorate salt was preferred over, say, nitrate? I'll do some looking, although if I turn anything up it will probably be too late to do you any good for your report. It would have been appropriate to have at least mentioned the IR features in Jolly, to aid in characterization. Perhaps the visible spectrum was of greater interest, given the vivid color of the compound. Steve |
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