bereal511
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Vanadium oxides
These probably seem like simple questions, but is it possible to simply obtain V2O4/V2O3 from V2O5 by roasting it? And does V2O4/V2O3 absorb oxygen
and reforms into V2O5 over time?
As an adolescent I aspired to lasting fame, I craved factual certainty, and I thirsted for a meaningful vision of human life -- so I became a
scientist. This is like becoming an archbishop so you can meet girls.
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woelen
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Not very well. V2O5 volatilizes, it is a covalent compound. V2O5 is used in ceramics and coloring glass, and as such it remains stable even at the
temperatures of molten glass and the temperatures, used in baking ceramics.
According to my literature, V2O5 can be decomposed by heating, but unfortunately no temperature is mentioned. But because it is stable in the hot
ceramics oven, I expect its decomposition temperature to be very high, well above 1000 C, and also well above its volatilization temperature (this
volatilization also is a problem for ceramics-people, because of the toxicity of the vapor).
[Edited on 16-11-06 by woelen]
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bereal511
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Is there a relatively simple way to process V2O5 into it's lower oxides chemically?
As an adolescent I aspired to lasting fame, I craved factual certainty, and I thirsted for a meaningful vision of human life -- so I became a
scientist. This is like becoming an archbishop so you can meet girls.
-- Matt Cartmill
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woelen
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Yes, that is quite simple.
1) Dissolve the V2O5 in a solution of NaOH. You'll obtain a colorless or slightly yellow solution. Filter out any solid stuff. Now you have a solution
of NaVO3 and/or Na3VO4.
2) Add some 30% H2SO4. The liquid must be strongly acidic and yellow. When acid is added slowly, then you first see a deepening of color, and then it
becomes lighter again: colorless --> yellow ---> orange ---> red precipitate ---> deep yellow clear again ---> yellow. Now you have
yellow VO2(+) ions in a strongly acidic solution.
3) Add some solid sodium (bi)sulfite or potassium (bi)sulfite. Heat the liquid. The liquid will darken, become green and at a certain point it will
become beautifully bright blue. If this point is not reached, then you added insufficient (bi)sulfite. When the liquid is bright blue (without any
green shade), then add a small pinch of (bi)sulfite just for sure, and you should have a clearly observable pungesnt smell of SO2. If you have this,
boil the solution for a while to get rid of all SO2.
At this point you will have a beautiful bright blue solution, containing the vanadyl ion, VO(2+). The yellow VO2(+) is reduced to blue VO(2+),
vanadium has gone from oxidation state +5 to oxidation state +4.
The solution is still acidic, and also contains sulfate ions and sodium ions. I use this solution as is, without isolating the vanadyl. Isolating that
from the solution is very difficult. What can be done is neutralizing the acid, such that the pH is only slightly acidic and then evaporate the
liquid. Then you obtain a mix of Na2SO4 and VOSO4, but don't make this totally dry.
If you want further reduction of the vanadium, then take your blue solution of vanadyl-ions, add some extra acid and a spatula full of granulated
zinc. The solution then becomes green at first, due to formation of V(3+), and later it becomes lavender, due to formation of V(2+).
The following pages may be interesting for you:
http://woelen.scheikunde.net/science/chem/solutions/v.html
http://woelen.scheikunde.net/science/chem/exps/vanadium/inde...
Making the oxides from these solutions is not that difficult for the vanadium (IV) species, but it is very hard for the other two. These are oxidized
VERY easily, by oxygen from the air.
For the +4 oxidation state, you can precipitate VO(OH)2, by careful addition of NaOH to the blue solution (a picture of that precipitate is shown in
the first page I mentioned above). Do not add too much, that results in formation of brown V4O9(2-) ions.
Also, it is of great importance that you use distilled water, otherwise you obtain a very impure precipitate, apparently calcium and/or magnesium
precipitate with the brown V4O9(2-) very easily.
VO(OH)2 will decompose on heating, giving black VO2.
[Edited on 17-11-06 by woelen]
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bereal511
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Ah, thank you very much for the procedure.
As an adolescent I aspired to lasting fame, I craved factual certainty, and I thirsted for a meaningful vision of human life -- so I became a
scientist. This is like becoming an archbishop so you can meet girls.
-- Matt Cartmill
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SAM4CH
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If V2O5 reduces "volatilizes" at high temperature what is the oxide produce (is it V2O3, how can I return it to pentoxide?
What about carrying V2O5 on inorganic silica or on thermal ceramic to be use as catalyst in SO2-->SO3??
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Nerro
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I believe the ammonium salt of metavanadate (sp?) NH<sub>4</sub>VO<sub>3</sub> can be deposited on crushed rockwool or
something similar and heated untill all NH<sub>3</sub> is gone to deposit V<sub>2</sub>O<sub>5</sub>
[Edited on 11-5-2007 by Nerro]
#261501 +(11351)- [X]
the \"bishop\" came to our church today
he was a fucken impostor
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courtesy of bash
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SAM4CH
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how to perpare ammonium salt from V2O3 or from other vanadium oxides?
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woelen
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Preparing the ammonium salt actually is quite easy. You need potteries V2O5 of reasonable quality. Add this to aqueous ammonia and dissolve all of it
in excess ammonia. Heating is required to dissolve all of it. Filter away any insoluble matter while the liquid is still hot and keep the clear
liquid. Beware of the stink of ammonia, do this outside or at least in a very well ventilated room.
Next let the liquid evaporate in a petri dish for one or two days on a warm radiator (50 C or so). Soon, the ammonium metavanadate separates from the
liquid. The nice thing of this prep is that both the water and excess ammonia simply evaporate and no other impurities exist if good quality pure V2O5
is used as a starting point. Solid NH4VO3 is not hygroscopic and only sparingly soluble in water. The really pure solid is white. I have some of the
stuff, but that is very pale light yellow, probably due to some higher polynuclear vanadium (V) compounds, but that is of no concern.
If you want to make catalytic V2O5, then you might even skip the evaporation phase. Simply dissolve as much as possible of V2O5 in aqueous ammonia and
then soak some glass-wool with this solution, let dry and then heat to drive off the ammonia, leaving behind glass-wool, covered with V2O5. I however
must admit that I have not done this last thing personally. I expect it to work, but what its real value as catalyst for SO2 to SO3 conversion will
be, that should be determined by means of experimenting.
[Edited on 11-5-07 by woelen]
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not_important
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| Quote: | Originally posted by SAM4CH
If V2O5 reduces "volatilizes" at high temperature what is the oxide produce (is it V2O3, how can I return it to pentoxide?
What about carrying V2O5 on inorganic silica or on thermal ceramic to be use as catalyst in SO2-->SO3?? |
What about it? Standard industrial process, the catalytic effect is promoted by K or Cs salt, usually the sulfate, in small amount. The alkali
sulfates form acid sulfates and then a melt with vandanium ions, the liquid film is where most of the action occurs.
These catalysts are not run hot enough to cause loss of oxygen and reduction of the V(+5), plus they are in oxidising conditions.
There are a lot of patents covering the subject.
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not_important
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| Quote: | Originally posted by SAM4CH
how to perpare ammonium salt from V2O3 or from other vanadium oxides? |
V2O3 will dissolve in HCl to give V(+3), VCl3 aq. Air will oxidise it to the +4 state, oxidising agents such as H2O2 or HNO3 will take it all the
way to +5, with H2O2 you'll need to boil the reaction mix to decompose the peroxi compounds.
You're not too likely to find the lower oxides of vanadium. The pentoxide is the standard form in the trade, it and ammonium vanadate are stable and
easy to handle.
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not_important
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Needed to check my refs, including Sidgwick's Chemical Elements & their Compounds. They say that V2O5 is stable up to at least 1700 C,
near its boiling point, heating the solid in air or oxygen at 600 C is used to insure full conversion to V2O5.
V2O5 and strong hydrochloric acid will generate Cl2 and V(4)OCl2.aq, heating mixtures of V2O5 with alcohol and hydrochloric acid will do the same
without generating chlorine. The hydrate VOCl2 may be crystallised by concentrating the acid solution as green crystals. It hydrolysis in weak acid or
neutral solutions, adding Na2CO3 to the VOCl2 solution in HCl precipitates greyish 2 VO2.7H2O (or V(OH)4 1,5H2O).
Heating V2O5 with carbon at 1000 C gives V2O3, this is better done by using carbon monoxide or hydrogen as the reducing agent. The oxide is slowly
oxidised by air and can be ignited by heating.
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SAM4CH
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I tried to react V2O3 with HCl but the reaction then I added some nitric acid but no change in color occure!! I noticed the reaction is not active
...??!
I am not sure that it is V2O3 "it resulted from V2O5 at high temperature which was carried on organic silica!!
I need advice!!
Sam
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not_important
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V2O5 is pretty stable up to above 1700 C in air, so unless you had heated it in a reducing atmosphere you still have V2O5. The colour of the oxide
should be a clew as to which oxide you have.
What is "organic silica"? Some sort of New Age unicorns and butterflies quartz sand?
Boiling V2O5 with strong hydrochloride acid will generate Cl2 and reduce the vanadium to V(IV), you end up with hydrated green VOCl2. As you've
indicated no mole amounts, I've no idea what have happened with that reaction or any later one.
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phj
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I know for a fact that V2O5 is used as a catalyst in the production of sulfuric acid.
Hence, treating it with SO2 will yield a lower oxide, I guess.
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woelen
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Whether V2O5 is reduced by SO2 depends on the conditions used. Under certain conditions it catalyses the reaction between SO2 and O2 to give SO3.
Under these conditions, the V2O5 is not used up (it is catalyst).
In aqueous solution, however, V2O5 can be fairly easily reduced to vanadyl ion. Add some solid V2O5 to a solution of Na2SO3 in 10% H2SO4 and heat for
a while (assuring that the SO2 is not quickly driven off). The liquid becomes beautifully blue, due to formation of vanadyl ion.
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