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Chris The Great
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[*] posted on 24-10-2006 at 20:52
Vanadyl Diacetate


Hello

I've been interested in vanadyl diacetate, VO(OAc)2, as it forms a peroxide addition compound which is great for all sorts of selective, high yield oxidations.

However, the synthesis leaves a lot to be desired. I found only one version of it, and it's not fun. Taken from The one dimensional chain structures of vanadyl glycolate and vanadyl acetate, Curtis Weeks et al., J. Mater. Chem., 2003, 13, 1420–1423.
Quote:
In a typical preparation
0.224 g of a 1:1 molar ratio of V2O5 to LiOH was stirred for
24 hours in a flask containing 50 ml of glacial acetic acid. The
resulting orange solution was poured into a 125 ml Teflon lined
Parr reactor and heated in a conventional oven for 2 days at
200*C. After cooling for 12 hours the tan crystals were filtered,
washed with distilled water, and dried at 50*C. The lithium
hydroxide is not essential in the formation of VO(CH3COO)2
but its presence resulted in the formation of larger crystals.
When water was added to the acetate reaction medium, the
known compound V3O7-H2O was formed; this synthesis is,
however, simpler than those previously described.
Microwave heating at 200*C, in a CEM Model MDS 2100 oven,
produced VO(CH3COO)2 in one hour but only in the presence
of LiOH; the pressure of the vessels attained 80–120 psi. The
microwave products were always contaminated with unreacted
V2O5.

Sounds not fun.

I also have this reference, but checked the library and couldn't find it:
F.A. Cotton, Inorg. Syn., 13 (1972) 181

Anyway, what I'm thinking is reducing V2O5 to vanadyl, VO2+, using sodium sulfite. Woelen has this reaction on his (excellent) page. Could one then add acetic acid or perhaps sodium acetate, and perhaps filter out the product, or precipitate it with alcohol? Not a large amount is needed to funciton as a catalyst but scale the above procedure up to a few grams and look at how much glacial you'll be needing! :o
There's got to be an easier way, and I suspect their method was only because they wanted big crystals for analysis of them, not a useful amount for a synthesis.

Thanks a ton in advance!
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The_Davster
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[*] posted on 24-10-2006 at 21:20


Quote:
Originally posted by Chris The Great
I suspect their method was only because they wanted big crystals for analysis of them, not a useful amount for a synthesis.



Based on the journal, I can say you are spot on in thinking this, for some sorts of x ray diffraction you need large single crystals, there is some sort of advantage over powder XRD, it has to do with structure refinement, and atom position in the crystal lattice. I have had to read many an article full of latice constants, not my cup of tea, but others I know have devoted their career to it. I have a feeling that you should be able to find *somewhere* a synth for it in powder form.

I am 90% sure my library has that journal, I browse it often, I'll look up the specific article tomorrow or the day after.

As for ideas, if you can make vanadyl sulfate, you may be able to react this with silver acetate in suitable solvent leaving vanadyl acetate.




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[*] posted on 24-10-2006 at 21:21


Vanadyl sulfate is insoluble in ethanol also.



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[*] posted on 24-10-2006 at 21:24


WTF, I don't get it, why not just add LiOAc?!

I certainly don't see why you can't handle it like any other ion and form the salt, but I'm no left-side-transition-metal expert...

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[*] posted on 24-10-2006 at 21:26


They were going for very pure crystals, as well.

I've been researching this too, trying to find better routes. The route we had used in the past uses vanadyl sulfate trihydrate and strontium acetate in alcohol solution; the precipitated strontium sulfate is filtered off and the alcohol evaporated under vacuum. The VO(AcO)2 is a blue solid, when we made did this it didn't like to crystallise, we suspected one or both of the reactants were contaminating it.

I've tried making VOCl2 by heating V2O5, HCl, and ethanol, giving the green hydrated form. This is preperation that I read once and can't find again. Take an alcoholic solution of VOCl2 (hydrated) and add strontium or barium acetate in alcohol, filter off the ppt of Sr/Ba Cl2. Using lead acetate in a glycol or glycerin might work better, but Pb(OAc)2 isn't very soluble in simple alcohols. I keep meaning to try this, but most of my lab is packed away some distance from here...

[Edited on 25-10-2006 by not_important]
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[*] posted on 24-10-2006 at 21:49


Why an alcoholic solution?
Vanadyl sulfate is soluble in water, so just react it with barium or lead acetate. Or does it cause too much hydrolysis?




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[*] posted on 24-10-2006 at 22:08


Quote:
Originally posted by guy
Why an alcoholic solution?
Vanadyl sulfate is soluble in water, so just react it with barium or lead acetate. Or does it cause too much hydrolysis?


The original reference did so, and it does seem to give too much hydrolysis. Adding some acetic acid seemed to help, but we had no instrumentation to check the purity of the product, had to go with colour, transparancy, and crystal formation.

It seems to be difficult to get the VO(AcO)2 as a solid from solutions containing much water, and there are some reactions where you don't want a lot of water around.
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[*] posted on 24-10-2006 at 23:49


It is easy to make solutions of vanadyl salts, but I've had a really hard time isolating the salts. Vanadyl salts seem to be excessively hygroscopic and they are a pain to isolate from aqueous solution. I once tried and I ended up with some black vanadium (IV) oxide/sulfate crap, totally insoluble in water.

I have been looking for buying some vanadyl salts, and finally I purchased a few grams of vanadyl sulfate hydrate, but this stuff is insanely expensive ($1 per gram!!!), while vanadium (V) compounds, such as V2O5, and NaVO3 are cheap. So, there must be a reason for this high price of vanadium (IV) compounds, and that probably has to do with the difficulty of isolating them.

Here follows a picture of reagent grade vanadyl sulfate.



As you see, this stuff is very wet, and even this wet stuff costs $1 per gram. It is reagent grade, so it will be pure with respect to other metals and other anions, but it certainly is not pure with respect to water.

I also tried making other vanadium salts, but salts of weak acids are particularly difficult to make from aqueous solution. On heating, the weak acid is driven off, and some black gunk remains behind.




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[*] posted on 25-10-2006 at 00:38


Is your interest based on that article posted by Phel, "Unusual reactivity of peroxo vanadium complex for catalytic oxidation of aralkenes to benzaldehydes selectively (>99%) in conjunction with aqueous H2O2 as an oxidant is described here for the first time."?

The one that suggested styrene can be oxidized to benzaldehyde?

If so, I can say that forming this complex is quite simple. The way I did it was to first heat V2O5 with ordinary hydrated oxalic acid. There's a brief and vigorous reaction after the acid melts, yielding a beautiful blue glassy mass (I always used acid in some excess). This material readily dissolves in hot acetic acid and appears to give the red complex described in the paper upon addition of H2O2. It seemed to oxidize styrene as advertised, but my crude methods gave poor temperature control when I tried to go beyond test tube scale, and a lot of acetic acid is used relative to styrene, so I didn't repeat my experiments with it.




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[*] posted on 25-10-2006 at 00:59


Reduction with oxalic acid usually gives a complex with excess oxalic acid, I'm not sure what effect the complex might have. In this case the complex may be destroyed by the H2O2, oxidising to CO2 with the vanadyl grabbing onto the acetic.

A lot of vanadium complexes and compounds shift to red when hydrating/hydrolizing; the esters with alcohols are famous for this. So the formation of a red colour isn't a 100% proof of the particular complex being formed.

Also, many vandium salts and complexes will give some aldehyde along with oxide when reacted with alkenes and H2O2 or other peroxides; this particular one is interesting in its selectivity for the aldehyde. I haven't had a chance to try it, but if it matches similar V oxidations much less acetic acid will be needed for scaled up runs. A PTC may help in this aspect as well, and possibly using an alcohol-acid solvent to get the styrene to mix with the oxidant and catalyst. I don't know, perhaps even acetone-acetic acid could be used; I suspect some acetic acid needs to be there to keep the vanadyl behaving properly.

A good set of experiments for someone with a GC, as that would make it easy to determine the relative amounts of aldehyde, bensoic acid, and styrene oxide formed, plus any oxidation of alcohol if used as a co-solvent. It would also be nice to know the fate of the =CH2, does it end up as formaldehyde, formic acid, or CO2 + H2O?
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[*] posted on 25-10-2006 at 01:25


Quote:
Reduction with oxalic acid usually gives a complex with excess oxalic acid, I'm not sure what effect the complex might have. In this case the complex may be destroyed by the H2O2, oxidising to CO2 with the vanadyl grabbing onto the acetic.

Vanadyl also is oxidized by hydrogen peroxide. This reaction is very fast and complete. I think that the oxidation of vanadium (IV) to vanadium (V) is faster than the oxidation of the oxalate to carbon dioxide.




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[*] posted on 25-10-2006 at 17:10


I got the article, unfortunatly my new printer makes worse than horrible scans(or at least I can't get the settings right). But the gist of the article is dissolving V2O5 in refluxing acetic anhydride(slight excess) at 140C, untill the ppt is a constant grey colour. Washed with carbon tetrachloride to remove excess AA.

I'll try to find a better scanner to use.




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[*] posted on 26-10-2006 at 00:07


Acetic anhydride? Damn, I didn't order a bottle of that :mad: lol

Thanks for getting the general idea though. So what is the "slight excess" that they used? I'm not sure what the actual reactions will be, as some of the AA is going to be oxidized and some will react to form the desired compound etc.

Yes polverone, that was the article that caught my interest in Vanadium compounds as oxidizers, this one in particular. I'm also interested in what other things it can oxidize, the authors in that article had no luck with phenyl on the alkene chain but I'm sure it would get some better luck with other groups. I'm thinking piperine or piperic acid has a good chance of a decent yeild using this to get the aldehyde. I'm sure a lot of people would be quite happy then after trying the KMnO4 oxidation to poor results.
Did you happen to post this here? I did a search but turned up nothing on that topic.
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[*] posted on 26-10-2006 at 22:36


If the peroxide " addition compound " is what you want ,
and not specifically to isolate vanadyl acetate .....

Why not just gradualy add 27% or stronger H2O2 to a stirred suspension of vanadium pentoxide in warm glacial acetic acid and watch the pervanadic acid / peracetic acid
complex form ?

It would seem that ultimately that is what you will have anyway as the active reagent once the peroxide addition is made to whatever vanadium acetate compounds .
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[*] posted on 27-10-2006 at 07:56


Actual vanadyl acetate is wanted because the article describes it as giving a high yield, very selective oxidation to the aldehyde. Other vanadium compounds and H2O2 give a wider mix of products, including the epoxide, corresponding acids and esters, and polymers; all depending on what form the vanadium was in and other conditions.

It's not clear that the desired complex, which includes acetate, would be formed from V2O5 in acetic acid. The short cut you suggest could be tried, but as I lack a GC, IR spec, and NMR, it would take a lot of work to confirm the product mix. I'd thought about adding a vanadium ester, from refluxing V2O5 in dry alcohol, to acetic acid, but it's not clear to me that VO(OR)3 would transform to the proper complex to end up giving VO(O2)(OAc)2(H2O)2.
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[*] posted on 28-10-2006 at 11:51


I just have a feeling that in the peracetic acid solvent ,
which is a pretty active reagent by itself ....that all roads lead to Rome as it were , favoring the formation of the
same pervanadic acid / peracetic acid mixture from whatever vanadium contributing precursors are useful ,
and that either the reducing or oxidizing capability of
the peroxide would come into play as needed to get
the oxidation state of the vanadium " corrected " to
what is favorable for the formation of the complex .

I could be wrong about this ...but it would be my first guess that whatever oxide or hydroxide or carbonate or acetate of vanadium at whatever oxidation state would
end up forming the same reagent in reacting with
peracetic acid . Vanadium Pentoxide is used simply because it is a common commercial form for the active
principle ....which is the Vanadium ....but different
compounds should lead to the same result .

And given the specification for the acetate of vanadium ,
which is very likely to form anyway in this reaction system , it seems even more likely in this case that other vanadium compounds easily convertible to the acetate
in situ .....would all lead to the same reagent .
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[*] posted on 28-10-2006 at 23:58


I did a small amount of checking, and there have been alkane oxidations done with other vanadium salts in acetic acid, and those lead to mixtures of epoxide, aldehyde, carboxylic acid, and other products. Perhaps in those compounds the complex was stable enough that the acetate was not formed. They did form a red coloured complex with peroxide, which may well be a characteristic of the V->O2 interaction. I don't remember we ever got much aldehyde from V2O5 oxidiations, we were after styrene oxide for an alternative route to 2-phenyl ethanol and benzaldehyde would have shown up as benzyl alcohol

Guess I'll have to make some acetic acid from vinegar and try the direct route. I have my doubts it will work, vandium(V) just doesn't seem to give di-carboxylates directly.
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[*] posted on 30-10-2006 at 00:59


Vanadium(V) compounds, except for peroxides, would be substantially weaker oxidants than Cr(VI) compounds, which are commonly used in laboratory oxidations of primary and secondary alcohols to aldehydes and ketones. By contrast, permangate oxidizes further, to carboxylic acids, and can under rigorous conditions also break C-C bonds to form them (hence it can react with C=C double bonds and tertiary alcohols).

Vanadium found in natural ores is exclusively pentavalent, as vanadates. V2O5 is used as a catalyst for the oxidation of SO2 to SO3 in the manufacture of sulfuric acid, being alternately reduced to VO2 and re-oxidized to V2O5 by O2.

[Edited on 30-10-2006 by JohnWW]
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[*] posted on 24-11-2006 at 22:17


As requested: (scanner still on the fritz)


Quote:

Vanadium(V) oxide (18.2g) is added to acetic anhydride(50mL) in a round bottomed flask fitted with a reflux condenser and a silical gel guard tube. The ammount of acetic anhydride added should be in excess of that required theoretically for the reaction. This prevents the solid fluffy product from charring during refluxing of the contents. The contents are heated under reflux using an oil bath or electric mantle maintained at 140 +/- 5 C. The start of the reaction is indicated by a change in the color from the reddish-brown of vanadium(V) oxide to light gray. During the reaction, the contents may need to be shaken occasionally to prevent caking of vanadium (V) oxide. The reaction is complete in an hour as indicated by the separation of a product which does not change its gray color even on prolonged refluxing. The contents are cooled and filtered through a sintered-glass funnel. The solid product is then transfered to a 200mL round bottom flask containing carbon tetrachloride and the mixture refluxed for 10-15 minutes to remove traces of acetic anhydride from the compound. The product is again filtered and washed with carbon tetrachloride. After it has been allowed to stand in the air for 2h it is dried in vaccuum at room temperature for about an hour. The excess acetic anhydride may be recovered by distillation of the filtrate. Yield 35g(95%)


Properties:
The compound is a gray, nonhygroscopic powder, odorless powder. It does not melt but decomposes on heating. Its pyrolysis curve reveals that it is stable upt to 214C. It looes weight between 214 and 388C and attains a constant weight at 388C, leaving a residue of vanadium(V) oxide. It is insoluble in common organic solvents, eg., carbon tetrachloride, benzene, chloroform, and cyclohexane. It does not form any addition compounds with tertiary organic bases like pyridine, picolines, etc. Its IR absorbtion spectra has the following charactistic bands: 2854(s), 1495(s), 1450(s), 1065(w), 1040(m), 900(s), and 665(s).



[Edited on 25-11-2006 by The_Davster]




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Chris The Great
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[*] posted on 22-1-2007 at 03:55


First, the starting materials are measured out, 18.31g vanadium pentoxide (pottery supplier) and 50ml acetic anhydride (chemical supplier). The scale turned off automatically so the reading includes the wax paper as well, since I know somebody is going to read the numbers on it lol.


The acetic anhydride is poured into the 125ml RB flask, and then the vanadium pentoxide is added. The flask is fitted with a condenser and placed in an oil bath.


I overheated it at first and overloaded my condenser, the reaction boiled so much it foamed over and steam (corrosive acetic anhydride steam) slowly drifted out the top of the condenser for a few minutes while the reaction got under control. After that there were no problems at all. The colour changed from orange to a brown/gray.


I kept refluxing, and added another 10ml or so acetic anhydride after 1.5 hours since the colour was not gray! I refluxed for a total of 3.5 hours, before I clued in that the colour was not changing and the authors simply are not great at describing the final product, which is more of a brown than a gray. So I let it cool, and filtered with a double coffee filter. Methylene chloride was added to the flask and swirled to get the settled sludge out of the bottom, and added to the filter. This whole process involved a lot of crying at hot acetic anhydride is NOT fun for the eyes! Now I know why there is a little icon of a fume hood on the bottle :( There was a solid chunk at the bottom that fused together which couldn't be washed out. The product dried fairly nicely but still smells of acetic acid, it will need to be refluxed in DCM as the article says. But that is for tomorrow.


The chunk of stuff was dissolved into HCl to get it out. Here is a picture of the blueish solution it forms (it is slightly green because the large chunk probably contains some unreacted vanadium pentoxide- it's more blue in real life).


When hydrogen peroxide is added, it forms a very dark red-brown colour as it should. Here it is in my waste bucket, the flash made it much lighter than it actually is. It does show that the product is what was desired though.


[Edited on 22-1-2007 by Chris The Great]
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[*] posted on 22-1-2007 at 23:41


Nice procedure. It looks like you made it.

There is one remark, I want to add. The last check with the peroxide does not tell anything. The red/brown color is not specific for vanadyl acetate, it occurs with ANY vanadyl and also any vanadium (V) species in both weak and strong acid solution. This is some vavandium peroxo complex with vanadium in the +5 oxidation state. With vanadyl, first the vanadium is oxidized from +4 to +5 oxidation state and then the peroxo complex is formed.

Just do the following: Take a pinch of V2O5. Add some NaOH and some water. Heat a little and then all dissolves, you get a colorless solution. Next, acidify it. You get a yellow/orange solution of a vanadium (+V) species. VO2(+). Then add some hydrogen peroxide. This will give you the red/brown complex.

[Edited on 23-1-07 by woelen]




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[*] posted on 23-1-2007 at 00:00


The dried powder was placed in a 250ml flask and refluxed with approximately 100ml of methylene chloride for about 20 minutes. It was filtered, the flask being rinsed with more DCM to get out the last of the sludge, and then the powder was air dried. It left a fairly greyish powder, smelling very faintly of acetic acid. The yield was 26.25g (71%) which is pretty good considering some of the stuff solidified and caked up during the synthesis. I imagine that if one had stirring that the reaction would yield the 95% claimed in the paper.

Thanks for that reminder woelen, I didn't know that all the peroxide complexes had that colour. The acid solution to clean out the caked stuff was a very deep blue today, so it's clearly in the +4 oxidation state.

Now to make some alkenes to see if this stuff oxidizes as advertised :D
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