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Author: Subject: How to convert ferrovanadium (82%) in vanadium pentoxide
thorazine
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[*] posted on 20-5-2010 at 15:01


Hi!

Woelen, sorry, my mistake! What added was NaClO3 like not_important said. I tried it because I saw a paper where they use NaClO3 as oxidant in the extraction of vanadium.

As I said I started all over again. I found a drug store here that sells H2O2 (30%). But even it didn't went very well.

I added 12.29g of FeV to 120 mL of HCl (30%). Let boil for about an hour and a half and filtered.

When filtered it there was a lot of (metal?) residue, as you can see:


I obtained a reddish solution, iodine-like , which filtered. Here's the photo:


I figured it would take about 10 mL H2O2, but with this amount just turned dark - green. Continued to add up to a total of 35 mL, until the foam becomes bright orange in the top (its difficult to see the real colour of solution). Boiled a few minutes as not_importande said, and in the end the solution was too dark. While boiling slowly added 14 g of Na2CO3. In the end I got a car oil-like solution!!

Thats here:


I filtered and i've got a mud-like ppt.




What might have gone wrong? The Vanadium not been dissolved in its entirety? Oxidation was not complete? Other metals precipitated at the end - i know that V is on the solution, but whats that crap in the filter? Any ideias?

Thanks! :

PS - Next week the guy in the drug store will receive nitric acid (60%). Could it be an option to dissolve this, right?

[Edited on 20-5-2010 by thorazine]
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[*] posted on 20-5-2010 at 19:34


If it wasn't completely oxidized, you may have magnetite. Is it magnetic?

Tim




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woelen
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[*] posted on 21-5-2010 at 04:51


The dark red iodine-like color can perfectly be explained. This is due to a peroxo complex in strongly acidic medium. If you have this, then certainly you have the vanadium in solution.

The powder you still have after dissolving the ferrovanadium could well be (impure) carbon. Many metal alloys contain considerable amounts of carbon, which do not dissolve when the rest of the alloy does dissolve. This reminds me of dissolving aluminium from a hardware store in hydrochloric acid. I always have a dark grey precipitate left from such aluminium and you probably have the same here.

The dark iodine/like liquid now should be boiled for quite some time in order to destroy the complex. This indeed can make the liquid blue or green. The vanadium in oxidation state +5 is capable of oxidizing hydrochloric acid to chlorine and then itself becomes blue vanadyl. Did you smell any chlorine gas during the boiling.

Now, if you add Na2CO3 then iron will precipitate, together with other metals, but your vanadium also will precipitate if it is in oxidation state +4! This is not what you want. Vanadium in oxidation state +4 only dissolves in strongly alkaline solution, giving the brown hypovanadate ion, V4O9(2-). I can imagine that sodium carbonate is not sufficiently alkaline and that you get a precipitate of vanadyl hydroxide (VO(OH)2), which is dark grey.

You either have to assure that the vanadium is in oxidation state +5 before adding Na2CO3 (e.g. add NaClO3, but not peroxide as that rebuilds the peroxo complex), or you have to add NaOH in order to have the iron precipitated and the vanadium in solution again.

You could try the latter with your mud in the filter.


If I were you, I first would try things on a test tube scale. In that way you don't loose as much of the chemicals as you do now and you produce less environmentally problematic waste.


Summarizing:

vanadium in oxidation state +4 forms a blue solution in neutral and acidic media, which contain VO(2+).
vanadium in oxidation state +4 forms a dark grey precipitate of VO(OH)2 in neutral to moderately basic media.
vanadium in oxidation state +4 forms a brown soluble anionic species V4O9(2-)

vanadium in oxidation state +5 forms a pale yellow solution in strongly acidic media, which contain VO2(+).
vanadium in oxidation state +5 forms a deep orange precipiate of V2O5.nH2O at pH around 3.
vanadium in oxidation state +5 forms a pale soluble yellow anionic species in moderately alkaline solutions.
vanadium in oxidation state +5 forms a colorless solution in strongly alkaline media, which contain VO4(3-).

With peroxide, vanadium in oxidation state +5 forms a deep red/brown peroxo complex at low pH and a bright yellow peroxo complex at high pH and a dark blue peroxo complex at very high pH and very high concentration of H2O2.
Vanadium in oxidation state +4 does not form peroxo complexes. This is oxidized to state +5 by hydrogen peroxide and if there is excess hydrogen peroxide then the peroxo complex is formed.




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tetrahedron
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[*] posted on 14-10-2012 at 07:49
reaction of V2O5 and Na2CO3


V2O5 + Na2CO3 ---> 2NaVO3 + CO2

28.50g V2O5 (pottery grade; excess) and 17.48g Na2CO3 (also pottery grade; not sure about the hydration state) were weighed and added into a 200ml Erlenmeyer flask. about 150ml deionized H2O was rapidly added. the mix started bubbling immediately but did not heat up. at this point the color was still the dark yellow/light orange of V2O5. the mix was left for about 1h and shaken occasionally. then it was heated in a water bath to about 50°C, shaken, and left for another hour. finally, the mix was brought to a gentle boil and swirled until the reaction looked complete. during this step the color changed to a dark orange/reddish brown, while some chunks/flakes were observed floating around (these had a "charred" aspect: partly yellow, partly black). after cooling, the flask looked as in the picture. does this look normal?

vanadate.jpg - 142kB

[Edited on 14-10-2012 by tetrahedron]
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[*] posted on 14-10-2012 at 08:42


@tetrahedron:

You may consider yourself lucky that some V2O5 dissolved anyway because these pottery grade oxides are often calcined and not very reactive. And sodium carbonate isn’t a strong base to start with. I have some pottery grade V2O5 but have never tried to dissolve it in anything, so I’m just talking in general and not from personal experience with that oxide.

The bubbling also suggests that some carbonate was neutralised.

I would filter off the solution and try dissolving the filter cake (presumed unreacted V2O5) in hot (simmering), strong NaOH or KOH. Also, try using an excess alkali, not an excess V2O5. I assume you want a solution of orthovanadate?




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sad.gif posted on 14-10-2012 at 10:19


thanks blogfast for the info. no, i'm interested in the metavanadate, that's why i avoided the strong NaOH. the vanadium pentoxide from the pottery supply had a nice "yellow cake" appearance, just slightly dark, definitely not black like someone else's. i used an excess because i didn't want any carbonate in the product. it's also possible and desirable that the metal impurities present precipitate as carbonate (rust color). of course i plan on recycling the unreacted vanadium.

the filtrate has a mandarin peel color, with a chrome yellow meniscus:

filtrate.jpg - 150kB

edit. in a cylinder it's obvious that a darker "phase" is settling on the bottom of a lighter one..confusing

[Edited on 14-10-2012 by tetrahedron]

settling.jpg - 123kB

edit. the lower layer is rather dense..addition of another 76mL H2O went to the upper layer, which turned a bright yellow. the cylinder was thoroughly shaken to mix the layers, then 20.86g (NH4)2SO4 dissolved in ~80mL H2O, but contrary to expectation no sudden precipitation occurred. the solution slowly turned cloudy. left in a salt-ice-water bath an orange deposit formed (~6mL), while the whole solution turned bright yellow. if there's ammonium metavanadate in there, it must be heavily contaminated.

btw with a similar procedure the home scientist seems to obtain quite a different result:

http://www.youtube.com/watch?v=KUHO1DKKyG4

[Edited on 14-10-2012 by tetrahedron]
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[*] posted on 15-10-2012 at 12:05


His result is different, possibly because his pottery V2O5 seems to have dissolved better than yours, thus giving higher yield?

Try dissolving the orange precipitate is as little an amount of hot water as possible, then allow to cool and ice. You might get better quality ammonium metavanadate that way...

So the initial reaction appears to be V2O5 + Na2CO3 == > 2NaVO3 + CO2.

Nice video, BTW...

[Edited on 15-10-2012 by blogfast25]




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[*] posted on 15-10-2012 at 15:17


coating.jpg - 183kB

i was surprised to find that overnight a whitish precipitate covered the whole length of the cylinder. the liquid phase (light yellow/green) was decanted off. after shaking with ~100mL hot H2O the coating went into suspension, yielding an opaque yellow liquid looking very much like orange juice. after a while, on top of the old orange sediment there was a distinct pale yellow layer, fading more and more to white toward the top. i cannot explain why this precipitate (arguably NH4VO3) did not form immediately upon cooling, but took several hours instead. pics follow asap.

[Edited on 16-10-2012 by tetrahedron]
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[*] posted on 16-10-2012 at 07:30


Quote: Originally posted by tetrahedron  


i cannot explain why this precipitate (arguably NH4VO3) did not form immediately upon cooling, but took several hours instead. pics follow asap.

[Edited on 16-10-2012 by tetrahedron]


I wouldn't worry about it. 'Delayed action crystallisation' isn't uncommon.

Going by Wiki's solubility data (temperature dependence) on NH4VO3 something might be gained from icing the solution, to squeeze the last bit of product out...

I think I might try this synth sometime soon...




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[*] posted on 16-10-2012 at 08:16


Personally, I would use Vinegar (acetic acid) and liquid Bleach (NaOCl). The HOCl formed readily attacks and slowly dissolves Iron especially well, even in dilute solutions, in the presence of the acetate to produce an obvious soluble Ferric salt (red-brown). Note, without the acetate and in a closed bottle (no air), a Ferrous salt (green solution) is formed (verify by forming the HOCl from the addition of H2CO3 to NaOCl).

Alternatively, one can also use the common etching solution mix of HCl/H2O2. Then, proceed as outlined by Woelen.

[Edited on 17-10-2012 by AJKOER]
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[*] posted on 18-10-2012 at 14:08


Quote: Originally posted by blogfast25  
I wouldn't worry about it. 'Delayed action crystallisation' isn't uncommon.


thanks, that's relieving.

i finally managed to isolate two distinct precipitates: a heavy orange one, and a light white one. the white one really likes to stick to the vertical walls of the container.

i also noticed a puzzling phenomenon: the yellow supernatant turns clear upon heating (orthovanadate? i only used sodium carbonate so the solution shouldn't be very alkaline).
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[*] posted on 20-10-2012 at 07:00


Quote: Originally posted by AJKOER  
Personally, I would use Vinegar (acetic acid) and liquid Bleach (NaOCl). The HOCl formed readily attacks and slowly dissolves Iron especially well, even in dilute solutions, in the presence of the acetate to produce an obvious soluble Ferric salt (red-brown). Note, without the acetate and in a closed bottle (no air), a Ferrous salt (green solution) is formed (verify by forming the HOCl from the addition of H2CO3 to NaOCl).

Alternatively, one can also use the common etching solution mix of HCl/H2O2. Then, proceed as outlined by Woelen.

[Edited on 17-10-2012 by AJKOER]


AJ:

Has this been experimentally verified or is this one of your ‘HOCl’ based theoretical schemes? If you have verified it, what’s the ratio of vinegar/bleach used? Commercial bleach is alkaline and vinegar is only about 0.8 M acetic acid: get the ratio of vinegar/bleach wrong and you end up with a solution containing sodium acetate and sodium hypochlorite, neither use nor ornament…

You talk about ‘HOCl formed’. How, prey, tell? HOCl, although quite unstable, is a strong acid. Now you’re claiming that mixing a weak acid (acetic acid) with the dissolved (and dissociated) salt of a strong acid (sodium hypochlorite) somehow yields that strong acid. But it doesn’t. Equilibrium theory shows very clearly that if you mix, for instance weak acetic acid with sodium chloride solution, nothing really happens other than the weak dissociation of the acetic acid, the chloride ions being essentially spectator ions.

You’d also like me to believe that the hypochlorite will not oxidise Fe (II) to Fe (III), yet hypochlorite is one of the most powerful oxidising agents we know of, so much so that, as shown in another thread, it can oxidise Cr (III) to chromate (VI)!




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sad.gif posted on 7-11-2012 at 04:12


due to impure reagents and many losses i was only able to isolate 3.16g of crude NH4VO3

NH4VO3.jpg - 189kB

maybe someday i'll attempt this exhausting procedure again, this time with technical grade V2O5 and excess Na2CO3, as in Thompson's video.
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[*] posted on 7-11-2012 at 06:45


Quote: Originally posted by tetrahedron  
due to impure reagents and many losses i was only able to isolate 3.16g of crude NH4VO3


Nice looking product but disappointing yield...




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[*] posted on 3-6-2013 at 11:58


Would it be practical to isolate vanadium from high-speed tool steel (1-5% V)? The nearest pottery shop is a long way away...



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[*] posted on 3-6-2013 at 13:29


Quote: Originally posted by elementcollector1  
Would it be practical to isolate vanadium from high-speed tool steel (1-5% V)? The nearest pottery shop is a long way away...


The trip to that pottery shop is worth an awful lot of HCl! :D

Seriously, if V was a precious metal, I'd say 'go for it' any day. But even at the 5 % level? Get yourself some used vanadium steel spanners from eBay, if you're gonna go the recycling route...

[Edited on 3-6-2013 by blogfast25]




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[*] posted on 3-6-2013 at 14:30


Hmm. Well, okay... I'll save up so I can pick up some other stuff on the way there. (Nd2O3, etc...
http://www.seattlepotterysupply.com/Merchant2/merchant.mvc?S...)




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