I have ordered some of this, costing an arm and a leg at over $260 per gram.
Does anybody have any practical advice about handling it. I will be oxidising double bonds to diols, using catalytic amounts of the OsO4 plus a
secondary oxidant.
But what does "catalytic amounts" mean in this context? Throw in a small crystal? Or add a ml of millimolar solution? Practical advice would be
appreciated.
And given the expense and the toxicity, is it possible or practical to recover the osmium during workup?
Edit WOOT I am now, appropriately, a hazard to self.
[Edited on 30-3-2009 by Paddywhacker]panziandi - 29-3-2009 at 17:50
Use the OsO4 oxidation in that preparation for ideas for yours.
Handle OsO4 under a fume hood, it is volatile and is corrosive to your eyes and respiraory tract. I usually see it supplied in vials or ampoules break
it with care, transfer what is required and seal the rest tightly.
You can recycle Os waste, depends how uch you will be wasting each time and disposal means as to whether you do or not. You could collect Os waste in
a waste bottle and then purify it when you have a sizable quantity to recycle.UnintentionalChaos - 29-3-2009 at 20:05
Edit WOOT I am now, appropriately, a hazard to self.
Absolutely. OsO4 is the last reagent I would turn to for making diols if I had to. I'd have just stuck to ring-opening epoxides. Peroxyacids are a
walk in the park by comparison.
[Edited on 3-30-09 by UnintentionalChaos]not_important - 29-3-2009 at 20:13
The " corrosive to your eyes" is in part the OsO4 being reduced by the surface of the eye to finely divided black OsO2 or even Os metal, dimming
vision or even causing blindness.
Recovery is possible, given the small amounts used this usually is done by reducing to OsO2 (ethanol will do) and adding a solid carrier to collect
that; eventually the collected carrier is processed to recover the Os content, which is converted back to OsO4.
OsO4 forms a complex with HMTA which is much easier to handle and works as well as pure OsO4 in most cases.
and other places. Paddywhacker - 29-3-2009 at 20:43
I have seen that a solution of the complex with hexamine is considered safer to handle, and that is probably the way I will go.
While not being an apologist for its toxicity, I will point out that it is used in catalytic amounts, so it should be easier to handle than
stochiometric amounts of peracids or some improvised ozone generator.
In asking for advice from anyone who has gone this route I am trying to minimize the probability of my having an unfortunate incident.
I seem to remember being employed, years ago, by a lab that used OsO4 as a catalyst in Kjeldahl digestions. Those were the days.
Edit:- Tracked down the Osmed reference. It is a safe-to-store solid that regenerates OsO4 when dissolved. I will convert all my purchase to this
form, considering the advice I have recieved from forum members.
Much appreciated, thanks.
[Edited on 30-3-2009 by Paddywhacker]Fleaker - 30-3-2009 at 12:37
Are you seriously using OsO4 for this? General rule is that if you can smell it, it's not good. It has a very high vapour pressure, since it sublimes
at room temp. Just make certain that your hood works extremely well and don't let it leave the hood until you're sure it's been reduced.
I had a bottle of it a long time ago. It was bubble wrapped up and stored in an air tight container in a safe. Over a few months bubble wrap was
metallized black with Os. That could be your lungs or your eyes.
If I were you, I would have just made it. It is very simple to make from osmium or a salt, and easily stored if you are the least bit proficient in
ampouling. garage chemist - 30-3-2009 at 15:01
I have a little Os powder, and even this produces enough OsO4 in air to slowly blacken then inside walls of a plastic container in which the opened
glass ampoule was stored.
The smell was minimal though, it seems like the OsO4 was being reduced by the plastic as fast as it was formed.
In OsO4-catalysed organic reactions, the Os values can be precipitated at the end by addition of a reductant (I think it was either sodium sulfite or
dithionite). This precipitates a black powder of unknown (to me) composition, which contains the Os.
Metallic Os and its compounds of lower oxidation states can be dissolved in a solution of NaOH and sodium persulfate with warming, giving a deep
orange solution (perosmate (VII)). Acidification then releases OsO4, which can be distilled off as its azeotrope with water, which boils below 100°C,
if desired.
OsO4 is quite soluble in water, so the distillate may come over homogenous if only small amounts of OsO4 are being produced.
For the production of anhydrous OsO4, oxidation of Os powder in a stream of dry oxygen in a tube furnace (at a certain temperature, too high and OsO2
is the only result) followed by condensation with dry ice in a trap is the only option.
Metallic Os is available on ebay in the form of massive vacuum-melted buttons at somewhat less than 100$/gram.
I have no idea on how those could be turned into Os compounds- dissolution in NaOH/persulfate is most likely far too slow, and combustion in oxygen
probably as well.
PHILOU Zrealone - 31-3-2009 at 02:09
KMnO4 diluted in basic media can also make diols out of alcenes...must be much armless than OsO4...sparkgap - 31-3-2009 at 03:57
Indeed, even cold dilute neutral permanganate solution is sufficient for making diols out of alkenes.
Anyway, I also want to say that I got the osmeth article from here. It was too big for attachment so I uploaded it here instead.
sparky (~_~)Fleaker - 31-3-2009 at 10:11
Osmium can be had much cheaper than 100/g. Osmeth is popular for storing OsO4 for stains, although I seem to remember being able to saturate hexane
(not hexamine) with it and use that for fixing/staining/killing/what have you.
