Sciencemadness Discussion Board

How to convert ferrovanadium (82%) in vanadium pentoxide

thorazine - 14-5-2010 at 12:55

Hi! I bought some ferrovanadium in ebay and what i need is vanadium pentoxide. Anyone know any route or literature where i can find a way to extract and purify vanadium? Thanks!

12AX7 - 14-5-2010 at 13:48

Mmmm...

Dissolve in sulfuric, filter impurities, base extract and oxidize, filter Fe(OH)3 and impurities, precipitate V2O5? Or crystallize sodium vanadate.

Tim

not_important - 15-5-2010 at 03:40

Actually, use hydrochloric acid to get it into solution, it's better at dissolving the iron and forms soluble vanadium oxychlorides. You can use tech grade or hardware store HCl as the main contaminate is iron, hardly a concern in this case; dilute so as to be 10% to 20% HCl.

Once you've done that filter the solution, then treat it so as to oxidise all the iron to Fe(III) and all the vanadium to V(V). You could slowly bubble Cl2 through the warm solution, or bring it to near boiling and drip in fairly concentrated H2O2 (25% to 35%); the solution will change from a greenish to a yellow to yellow-orange as this is done. Calculate the needed amount of Cl2 or H2O2 based on the assumption that all the iron is originally Fe(II) and all vanadium as V(IV), use a distinct excess. Finish off by boiling the solution for a few minutes.

Then use Na2CO3 to neutralise the solution, the remaining HCl and the FeCl3 and the VOCl3, and add enough excess of Na2CO3 so as to dissolve the vanadium as sodium metavanadate NaVO3. You can calculate the excess amount based on 18 g anhydrous Na2CO3 to 25 grams V2O5 or 14 grams of vanadium in the ferrovanadium, plus however much would neutralise off the HCl you started with. There should be about 125 ml of water for every 18 g Na2CO3 (14 g V) going to making NaVO3 so add more water if the filtered acid solution is more concentrated than this (in terms of grams V to ml H2O). Keep the solution gently boiling will doing this, both the keep the NaVO3 in solution and to drive off CO2 - as the Na2CO3 is added the CO2 will create foaming so add it slowly and stir after each addition. After all the Na2CO3 has been added, boil for several minutes more.

Filter the hot solution to remove the ppt of iron hydrated oxides. If the filtrate is not clear, run it through the same filter again as needed to remove any solids.

At this point you can simply drop the pH to around 2 to get V2O5. However the V2O5 prepared this way is often somewhat colloidal and difficult to retain on a filter, so the following alternative may be better.

Allow the solution to cool to 60-70 C. Now add a strong solution of an ammonium salt, 75 g of NH4Cl in 125 ml of hot water for every 14 g of starting vanadium (multiply as needed to get the matching weight of ferrovanadium) Ammonium sulfate or nitrate can also be used, the amount of water is adjusted so as to insure solution of the ammonium salt and the corresponding sodium salt - don't forget the amount of water in both the ammonium and vanadium solutions. Mix well, then cover and set aside to cool for at least 6 hours. Collect the pale yellow crystals of NH4VO3 on a filter, then wash with first a small amount of the ammonium salt solution, them with several small amounts of ice cold distilled/deionised water - until no more chloride can be detected in the wash water if you've AgNO3 to test with. Allow the NH4VO3 crystals to air dry. Yield should be about 70% of the vanadium in the ferrovanadium.

You can heat the ammonium metavandate to 60-670 C, with stirring, in air to get V2O5. The NH4VO3 is a useful reagent on its own, as are the sodium and potassium metavanadates.





blogfast25 - 15-5-2010 at 09:36

Or you could just buy some reasonable quality V2O5 from a pottery shop, it's used as a yellow pigment in glazes.

I've never tried dissolving ferrovanadium in strong HCl but it sounds like a slow process, even at 20 % HCl. Let us know how you get on.

What do you want the V2O5 for?

thorazine - 15-5-2010 at 16:17

Thank you all!

not_important, i will try that this week! Thank you very much, i'm a beginner and your post is very detailed, its a big help!! :) I just have H202 3%, and you said that you need at least 22%. Here I have sodium hypochlorite (13%), will work as an oxidant? I found this paper* where they use it in oxidative precipitation.
* http://rapidshare.com/files/387794634/Recovery_of_vanadium_f...

I don´t have a fumehood, but i have this air extractor (300m2/h). I think if i put it over the bench there will be no problem, right?

An i don´t have NH4Cl but i have Ammonia (25%) and HCl, i think i can prepare it with this.. :)

Thanks, again!

blogfast25,
in portugal is hard to find, and online i didn't find V2O5. I have no real purpose. I like chemistry as i like metals, found it a good subject to start playing. I've tried to prepare FeCl3, but failed - i've just get a FDA approved Pigment Yellow 42 http://www.sciencemadness.org/talk/viewthread.php?tid=13794 ! Now i bought 250 g ferrovanadium in ebay. Maybe I'll convert 50g of Fev in ammonium metavandate and V2O5. Then, i will thhink in what I can explore more, I am searching for info on V what i can do with it, to learn and have fun :). I have noticed that I can do with ammonium metavandate Mendelin Reagent. Do you have more sugestions of things i can do with it?

Thanks!

thorazine - 15-5-2010 at 18:02

Humm..iwhat i said is wrong, in paper they said they used NaClO3 and not NaOCl!But i was searching and i have maybe 40g of Sodium chlorate ..can use that?



[Edited on 16-5-2010 by thorazine]

not_important - 15-5-2010 at 19:17

The problem with using more dilute oxidants is that you drop the concentration of vanadium in the solution, this is particulary true with your H2O2. Ammonium metavanadate is somewhat soluble, even in solutions saturated with ammonium salts. So too dilute means both leaving more vanadium behind in solution, and having to use more ammonium salt to precipitate the portion of the vanadium that can be forced out of solution.

The paper is targeting V2O5 as the precipitated substance, and it is much less soluble than the vanadate. But they're talking about commercial production where they can do a better job of filtering than you likely can.

You could go through the steps up past precipitating and filtering off the iron, using one of the oxidisers you have access too. Right off the top of my head I don't know if hypochlorite will do the oxidation - you could try it on a small scale, if if pushes the vanadium to the +5 state you'll see V2O5 dropping out of solution if the pH i in the 1-6 range. It might well work, as it will make Cl2 by reacting with the excess acid. Anyhow, once the iron is gone you could evaporate down the solution until it's just damp (steambath or similar heat source) and then add water to make a solution with the proper concentration of vanadium. The sodium metavanadate is stable, but doesn't crystallise readily and is somewhat hygroscopic so it's not too convenient of a compound for storing. Making the NH4VO3 gives a stage of purification, and is handy for storage.

If your HCl is technical grade, not water-white in appearance, distill it into aqueous ammonia to make the NH4Cl, you don't want to be adding iron back into the vanadium.

woelen has done a few things with vanadium, such as this http://81.207.88.128/science/chem/exps/vanadium/index.html


A note - while NaOH can be used to neutralise the acidic solution from dissolving the ferrovanadium, Na2CO3 has an advantage related to purification. The likely contaminates in the ferrovanadium are C, Si, Al, and Ca. Carbon isn't going to dissolve, likely remaining as a mix of carbon and metal carbides, but Si and Al can be brought into solution by strong base and Ca(OH)2 is slightly soluble. Using Na2CO3 pretty much avoids those problems.

As for the air extractor, I'll let someone else comment on that but it sounds rather small to me; do a forum search on "fumehood". If you avoid open containers of HCl in concentrations above 20%, and don't use Cl2 or hypochlorite with acid, you should be OK. The fumes from more concentrated HCl are hard on most metals, tools rust like crazy near it, and the fumes could damage the air extractor.

Remember that vanadium is somewhat toxic, although not persistent.

unionised - 16-5-2010 at 03:41

I'm fairly sure that, under alkaline conditions, air will oxidise Fe to Fe(III) which will precipitate and V to V(V) which will dissolve.
Air is generally considered to be cheap and of low hazard.

not_important - 16-5-2010 at 06:14

The iron in particular will, but air oxidation can be slow and the reference procedures that I've seen never rely on it. Part of the problem with air oxidation is that both iron and vanadium will have precipitated to some degree before conditions are alkaline enough to get decent oxidation rates, and the resulting solid mixtures tend to be slow to fully oxidise and leach out. The other likely contaminates don't help either.

And I selected the more common oxidants. Several procedures for NH4VO3 or V2O5 from crude V2O5 or ferrovanadium use KMnO4 as the oxidiser, others use HNO3. Both are harder to get now than back when the procedures were written, while the pool and spa industry has made it fairly easy to be able to generate chlorine.

Because I've never read or do such a process, I can't say "do thus and thus and thus, and most of the vanadium will be in solution while all the iron will have been removed." If I were to try it, I'd slowly drip the acid solution of V+Fe into a hot solution of Na2CO3 that was being strongly aerated, and stirred as well to prevent buildup of a precipitate layer on the bottom. If they have both a mechanical stirrer and an air pump that doesn't contaminate the air with oil, then it certainly could be tried.

thorazine - 16-5-2010 at 14:01

not_important,
thanks a LOT! :)

I was testing if a solution NaClO3 worked as an oxidant, but I got stuck on the first step. Added approximately 80 mL of HCl (% 20) at 8.11 g FeV, about an hour on a magnetic stirrer and not dissolving. Perhaps the process is very slow, as blogfast25 said. I feel the presence of a funny smell, so something must be happening .. :)






[Edited on 17-5-2010 by thorazine]

woelen - 16-5-2010 at 23:25

Yes, dissolving of ferrovanadium is very very slow. Ferrovanadium is known for its corrosion resistance and its strength (for these reasons it is used in better quality tools).

You could try adding some solution of NaOCl to the HCl (just a few drops, no gas should escape) and see if this helps. It also helps if you hit the ferrovanadium into small particles with a heavy hammer. Add the NaOCl very slowly, dropwise, while stirring. Let us know if this helps. Another option is strong heating (boiling) of the 15% solution of HCl. Beware of the nasty fumes from this boiling solution, do this outside. But in any case, patience will be needed.

The method of non_important for working up the solution indeed should work. The only point is that if you use H2O2 that a brown peroxo complex is formed at low pH, which is not that unstable and it may require quite some time of boiling before all of it is destroyed (this brown complex is much more stable than the similar deep blue chromium complex). You must destroy this complex before adding a base, so keep on boiling until the dark redbrown color of the complex is gone and the solution has turned more yellow than brown.

[Edited on 17-5-10 by woelen]

not_important - 16-5-2010 at 23:36

Thanks, woelen. And I'd forgotten how stable the perox complex can be.

Your ferrovanadium may have a higher percentage of vanadium than the stuff I worked with, and/or less aluminum, and/or more silicon. Pure vanadium isn't appreciably attacked by HCl, but the alloy I had did dissolve in warm HCl.

The alternative dissolution methods use hot HNO3 + H2SO4, or hot HNO3 + HF; the second isn't something you want to mess with.

Try heating and adding small amounts of H2O2 or HNO3, or do as woelen suggested using NaOCl; they are all based on getting small amounts of oxidiser into the solution (but not too much at once). In all cases there will be some chlorine given off, so have good ventilation, and if you do heat it to boiling then there's the HCl fumes too. If you heat that much then it is best to do it in a Erlenmeyer flask to keep in spray and act as a condenser.



[Edited on 17-5-2010 by not_important]

unionised - 17-5-2010 at 13:13

Incidentally, I know you can extract Fe(III) from solutions containing a lot of HCl using ether as an extraction solvent. Does vanadium extract?

I realise this is only useful if you can get ether but I's still like to know if the vanadium extracts too (take your pick of vanadium's oxidation state, provided it's compatible with Fe(III) compounds)

JohnWW - 17-5-2010 at 15:38

Quote: Originally posted by unionised  
Incidentally, I know you can extract Fe(III) from solutions containing a lot of HCl using ether as an extraction solvent. Does vanadium extract?
I realise this is only useful if you can get ether but I's still like to know if the vanadium extracts too (take your pick of vanadium's oxidation state, provided it's compatible with Fe(III) compounds)
In such solutions, Fe(III) would be present as [FeCl4]-. Although ether is polar, enabling ionic species to have some solubility in it, it is aprotic, meaning that the Fe(III) will not be hydrolyzed by it.

But vanadium is both more electropositive than Fe and more stable in higher oxidation states. If it is present in solution in the (V) state (which is most common in ordinary conditions), usually as the oxy-anion [VO4]---, strongly acid solutions are liable to hydrolyse it to polymeric anions and then precipitate it out as V2O5. If it can be gotten into solution from ferrovanadium in a strong HCl solution, it is most likely to be in the (III) oxidation state (Fe would initially be in the (II) state until further oxidized), which may be similarly as [VCl4]-; this would be less stable with respect to oxidation than [FeCl4]- but less liable to hydrolysis.

The main industrial use of vanadium, other than in high-strength ferrous alloys such as used in tools, is as solid V2O5 catalyst in the manufacture of H2SO4 via SO2 from SO2 gas. Do you want to make H2SO4 by this method? In this industrial reaction, the V2O5 is reduced to VO2 by the SO2 gas (from combustion of sulfur, obtained from mining (Frasch process mostly) or by recovery in refining from S compounds such as thiophene in crude oil), which is oxidized to SO3, and then the VO2 is oxidized back to V2O5 by heating in O2.

woelen - 18-5-2010 at 00:28

JohnWW, it is not true that vanadium in strongly acidic solution would be in the +3 oxidation state. That oxidation state is very strongly reducing. Especially if the help of some oxidizing agent is needed to get the vanadium in solution, it almost exclusively will be in the +5 oxidation state. In 20% HCl this will be in solution as a cationic species, the light yellow so-called pervanadyl ion, VO2(+), which must not be confused with the bright blue vanadyl ion VO(2+). The name 'pervanadyl' is bad, because this has nothing to do with peroxide or peroxo complexes, the prefix 'per' just tells that it is in a higher oxidation state than vanadyl (which has vanadium in oxidation state +4).

Vanadium in oxidation state +5 has a very complicated aqueous chemistry. At very high pH (around 14), it mainly exists as colorless VO4(3-), the orthovanadate ion. When pH is lowered, this ion is hydrolyzed and species like VO2(OH)2(-) are formed, which when crystallized from solution forms the metavanadate ion VO3(-), but the metavanadate ion probably is not a single simple ion, but more like a polymeric species [VO3)n](n-), very similar to the metaphosphates (calgon). When pH is reduced further, then further hydrolysis and condensation occurs, anionic species are formed with many vanadium atoms in it, with bridging O and OH groups. At pH in the range 2 ... 3 (IIRC) the hydrolysis and condensation becomes so strong that macroscopic neutral particles are formed, which can be formulated like nV2O5.mH2O, with n and m being very large. This is a deep orange/red precipitate. When pH is lowered further, then further protonization and subsequent splitting off of water occurs which leads to formation of deep orange polyatomic cationic species but finally at pH below 0 (strong acid at fairly high concentration) the splitting up is complete and simple aqueous VO2(+) ions go in solution.

Some info with pictures: http://woelen.homescience.net/science/chem/solutions/v.html

[Edited on 18-5-10 by woelen]

thorazine - 18-5-2010 at 07:36

Yesterday when I added a few drops of sodium hypochlorite (13%) the solution turned bright green. I left about an hour and a half on the magnetic stirrer, and not dissolved but the solution turned a green darker. Today when I see the solution was blue (as in the pic).



Uploaded with ImageShack.us


Right now I'm outside boiling the solution. When it reached the boiling point of the solution turned dark green again.

I have a question: as the solution is evaporating, as it boils, can i add water to keep the solution?

not_important - 18-5-2010 at 08:29

Yes, you can add water to maintain the volume. However the vanadium compounds formed have high solubilities, so some concentration will not cause problems.

Normally the green colour means that the solution is mostly V(III), while blue is V(IV). Ferrovanadium will function as a reducing agent, when the NaOCl had been consumed the vanadium in solution was being reduced, air can oxidise it. However aqueous VOCl2 in strong HCl is a bright green hue even though it is a V(IV) compound. Heat can also change the coordination sphere, changing the colour.



thorazine - 19-5-2010 at 04:29

Yesterday, after the solution being boiling about an hour i've get a red / brown solution and for what i read in woelen's site maybe its V4O92-. Filter and today added sodium hypochlorite until the entire solution becomes a dark green. Today i'll boil the solution to concentrate (the actual volume is maybe 350mL with 8,11g of FeV) and add Na2CO3 to neutralize the HCl and precipite Fe.

[Edited on 19-5-2010 by thorazine]

[Edited on 19-5-2010 by thorazine]

woelen - 19-5-2010 at 06:47

The brown ion V4O9(2-) only exists at high pH, in alkaline solution when vanadium is in the +4 oxidation state. So, if you have a red/brown color in acidic solution, then you certainly don't have V4O9(2-). At low pH, vanadium in oxidation state +4 forms the bright blue (very much like copper ion as in copper sulfate) vanadyl ion VO(2+).

thorazine - 19-5-2010 at 07:22

Correct. I read more carefully and really could not be V4O9 (2 -). What is the justification for that color, then? It was exactly that red / brow like in picture. However the solution is now dark green. I am boiling it to remove excess solvent because I have a small problem: the larger borosilicate glass I have is an 250 mL Erlenmeyer, i need to work with small volumes. When I was adding sodium hypochlorite it formed an intense yellow color, which after dissolving turned the green. The idea I have is to decrease the volume boiling and then add NaOCl (13%) to oxidize any Fe (II) to Fe (III) and all V (IV) to V (V) - until the solution turn yellow, then.

thorazine - 19-5-2010 at 08:52

I was experimenting and i add a solution of 10g of NaOCl3 in 15ml of water while boiling and it turned red/brown again, but a more bright red. And the foam on the topo turned yellow..

[Edited on 19-5-2010 by thorazine]

not_important - 19-5-2010 at 09:12

red-brown is often polymers of hydrated V(5) oxide, or mixed V(5)-V(4) Yellow is V(5) in solution at various pH values. Green can be solutions of V(4)+V(5), of V(4) in high Cl(-) concentrations.

VinwaterPourbaixdiagram.png - 9kB

The upper dashed line marks where the stuff will oxidise water to O2, the lower dashed line where water is reduced to H2. Those reactions may be slow for various reasons, and concentrations as well as the formation of complexes can shift things, but this should give you a feel for what species may be present. You should be able to find a similar Pourbaix diagram for iron - such as is here: http://en.wikipedia.org/wiki/Pourbaix_diagram

thorazine - 19-5-2010 at 10:49

When added NaOCl3 got a brown precipitate. Maybe V2O5 quite contaminated. I'll start again and do exactly as no_important said. Tomorrow I will try to go to a drugstore to get H2O2 (30%).

Thank you a lot! I'm a newbie, thanks for your patient. I'm learning a lot .. :)

woelen - 19-5-2010 at 22:45

NaOCl3? What do you mean with that? You mean NaOCl? or Na2CO3?

V2O5, when precipitated from aqueous solution, has a beautiful bright orange color, as you might have noticed from the pictures in my webpage. However, if iron in oxidation state +3 is present as well, I can imagine that a more dirty brown color is obtained and that the precipitate is a mixed hydrated ferric/vanadium(V) oxide, which also could be formulated as some form of hydrated ferric vanadate.

I think that the only way to get rid of the iron is to add a large excess of NaOH, such that all vanadium goes into solution as orthovanadate, while the iron precipitates as ferric hydroxide. The clear liquid then is decanted from the ferric hydroxide and then carefully acid is added, until the orange precipitate of hydrous V2O5 is formed. In this way you might be able to recover most of the vanadium. Don't add too much acid, as that leads to redissolving of the hydrated V2O5 and formation of VO2(+).

not_important - 19-5-2010 at 23:48

I think it's better to use Na2CO3 than NaOH. Hydroxide could pull some of the contaminates into solution, depending what is the the alloy besides V and Fe. The carbonate is commonly used in the various procedures I've read for getting pure vanadium compounds from the old style crude V2O5, and for making NH4VO3. Aqueous ammonia is alkaline enough to do the job, but as that compound is not very soluble using ammonia directly doesn't do well for purifying.

I think that NaClO3 was meant, it's been mentioned before.

thorazine - 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]

12AX7 - 20-5-2010 at 19:34

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

Tim

woelen - 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.

reaction of V2O5 and Na2CO3

tetrahedron - 14-10-2012 at 07:49

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]

blogfast25 - 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?

tetrahedron - 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]

blogfast25 - 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]

tetrahedron - 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]

blogfast25 - 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...

AJKOER - 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]

tetrahedron - 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).

blogfast25 - 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)!

tetrahedron - 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.

blogfast25 - 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...

elementcollector1 - 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...

blogfast25 - 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]

elementcollector1 - 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...)