Sciencemadness Discussion Board - Tungsten metal plus piranha solution?

Tungsten metal plus piranha solution?

stamasd - 21-8-2019 at 14:29

Does anyone have experience with reacting metallic tungsten with piranha solution? My ultimate goal is to make WO3 as easily as possible from metallic tungsten. I haven't found much about the subject via a web search. The only mention is here https://www.tungsten.com/chemical-resistance-of-tungsten/ : a mix of ammonia and H2O2 (aka "basic piranha") reacts with tungsten "under certain conditions"

Panache - 21-8-2019 at 14:59

Good question (cf/ so many questions here....8).
No idea but it seems like something i should try today, give me a few hours, i have all the reagants at hand.

stamasd - 21-8-2019 at 15:09

Thanks for trying it. I have the chemicals too but it's night time here so I won't be able to try until tomorrow.

j_sum1 - 21-8-2019 at 16:21

Tungsten will dissolve in H2O2 but it takes time. Think days not hours. I think this is the standard method for making WO3.

I have not done this myself -- a fellow member sent me some WO3 when he was working on this. I have done it with Mo and found there were many different reactions when I used different acids with the peroxide. I don't think I tried with ammonia however.

Panache - 21-8-2019 at 17:47

Having never made basic piranha before i sought a reference, before simply using 28% ammonia soln and conc sulphuric in a 3:1 volume ratio.
The mit guide confirmed the ratios, but also that it has to be at 60C or higher foreffective use.
Thus ill try two tungsten dissolutions, one at rt one at 60C.
As a control ill use an incandescent globe to light the room.

stamasd - 21-8-2019 at 18:00

Good luck! Depending on how it goes I'll prepare my own experiment tomorrow.
Also I assume you meant ammonia and H2O2 above. Don't mix concentrated ammonia with sulfuric acid.

Tsjerk - 22-8-2019 at 00:14

Tungsten should indeed dissolve in ammonia / peroxide. You don't need much ammonia, it is is used to keep the pH up and to keep tungstic acid in solution. You can precipitate tungstic acid by acidification.

WO3 can be made from the acid by heating on a flame.

Panache - 22-8-2019 at 06:43

Its going well, ive intially used some primary gap electrodes as the tungsten source because they are thick, thus should react controllably? Checkedvyge density 19.3, close enough

This is the beginning, on the left is the solutio i had heated to 60, one right rt.

stamasd - 22-8-2019 at 08:43

So here's my ongoing experiment. Just started it but so far it's going very well.

2.35g piece of tungsten electrode (TIG)
25ml 29% NH4OH plus 10ml 27% H2O2
Dropped tungsten in, placed outside in the sun on a piece of rock that has been baking in the sun for a few hours. Without any extra heating, just the heat from the rock and the reaction itself brought the temperature to 55C.

Vigorous reaction, and after 5 minutes the solution has already turned pale yellow from dissolved WO3.

[Edited on 22-8-2019 by stamasd]

Tsjerk - 22-8-2019 at 09:55

Nice work! If it is so vigorous, maybe diluting it down a bit and heating it way less would help against the decomposition of the peroxide and the vaporizing of the ammonia. My feeling says little ammonia will stay dissolved at 55 degrees. Just take your time with a stoppered flask.

WO3 doesn't dissolve

it is ammium tungstate. You want to keep that in solution, otherwise it forms paratungstate which is much less soluble. Getting that back in solution is no fun.

When have tungstic acid, you have way more freedom in making WO3 compared to when making it from ammonium paratungstate.

All the more reason to keep it cool and diluted. Ammonium metatungstate is soluble up to about 0,1 g/ml.

stamasd - 22-8-2019 at 10:21

So the latest update: i checked at 1.5h, the bubbling had stopped. The flask was still warm though not as much as before. The place where it rested was now in the shade. I didn't have the thermometer with me so I don't have the precise temperature. I estimate around 40C. The color was the same pale yellow, and there was a minute amount of needle-like crystals floating at the surface.

At this point I moved the flask to another warm place and waited for a few minutes to see if the reaction restarts. It did not. It still smelled very strongly of ammonia, though this in itself can be deceiving as to how much of the NH3 was still present.

Next step, I added another 10ml H2O2. For a few seconds nothing happened, then all of a sudden the whole liquid started bubbling violently and foaming about half-way up the flask. But since the flask is intentionally much larger than the volumes involved, it was nowhere near to the spilling point.

The bubbling died down in about 10seconds. Completely, no bubbling at all after that. The color also disappeared completely and now it's clear and colorless like water. No trace of crystals left either. I think all of the extra H2O2 that I added decomposed at once, oxidizing everything it could in the process.

This is very instructive. I have moved the flask now to a cool place, put a stopper and will allow it to cool to room temperature. Then I will add an extra portion of H2O2, and possibly extra ammonia as well.

The piece of tungsten looks corroded. Dull gray from the shiny gray it was before.

[Edited on 22-8-2019 by stamasd]

stamasd - 22-8-2019 at 12:16

Next step: after a further 2h in which nothing happened, I added another 10ml H2O2 to the room temperature flask (meaning around 32C).

Immediately the liquid turned pale yellow again, and gas bubbles started evolving from the tungsten. Not violent like the previous time, and localized on the metal only. That tells me that there's probably still enough ammonia to drive the reaction.

I don't really know how to explain the color changes though.

Continuing to observe.

stamasd - 23-8-2019 at 04:27

So, lesson learned: the "basic" piranha solution can be every bit as violent as the acid one.

This morning I added again to the reaction (at room temperature) 20ml H2O2 and 5ml NH4OH.
It started bubbling around the tungsten, but over the next 2 minutes the bubbling expanded gradually to the whole volume of liquid and foam started rising. The flask became hot, and eventually overflowed (because the total volume was higher now than yesterday). When I noticed that the temperature was rising I wuickly moved the flask into a water bath I had ready nearby so the overflow occured in the water bath. I estimate that I lost 5-10ml of the liquid.

At this point I decided to stop the experiment. There was again almost no gas production at all after the overflow, probably because all of the H2O2 decomposed in bulk.

Total in:
2.35g tungsten
30ml 29%NH4OH
40ml 27%H2O2

Total out:
2.308g tungsten (a whopping 42mg decrease)
50ml reaction mix
The missing volume is due to the overflow and probably some evaporation.

Workup in progress. I neutralized the solution with 20% HCl to slightly acidic pH=5. No precipitate formed. Now I'm evaporating it slowly on low heat.

Notes for the follow-up experiment: probably most of the H2O2 was wasted during the two episodes where it decomposed in bulk. Instead of adding it in big batches, small and repeated additions would probably be best. A very slow running addition funnel? The problem is, my only addition funnel is way too big, 250ml. Now I wish I had a 25-50ml one.

[Edited on 23-8-2019 by stamasd]

Tsjerk - 23-8-2019 at 04:49

Maybe you can try with less ammonia. Sure the basic piranha is nice to wash flasks and fritted funnels, but to dissolve tungsten it only has to be basic, I would guess around pH 11.

Peroxide is acidic and the higher the pH the more of it is in the ^-O-O^- form. That might be why all your peroxide is decomposing.

I would try to see what happens when you leave it for a day or two at RT with enough peroxide to lets say dissolve half a gram W, with just enough ammonia to keep the pH at around 11.

I would also start with a volume that is big enough to dissolve the amount of tungsten calculated to dissolve.

The reaction does consume a bit of OH^-, but half a gram of tungsten can never use much.

[Edited on 23-8-2019 by Tsjerk]

stamasd - 23-8-2019 at 07:19

Yes, less ammonia should work. The reaction appears to proceed vigorously enough at RT when enough peroxide is around. The decomposition of which is what limits its efficiency.

The evaporation of the solution on low heat has produced a white powder, most of which I think is ammonium chloride, perhaps with a small amount of ammonium tungstate. I will try to calcinate it, but that's for another day. I have house chores to attend to for the rest of the day.
(edit) the weight of the residue is 1.31g. There is also a fine white powder coating the inside of the flask all the way to the top, most likely evaporated and recrystallized ammonium chloride. I will not try to recover that as it's irrelevant to the main experiment.
Based on the tungsten weight difference, at most 60mg ammonium tungstate should be in that residue. And at most 52mg WO3 should result from calcination.

[Edited on 23-8-2019 by stamasd]

Tsjerk - 23-8-2019 at 09:44

If you really wanted to you could put the flask on a flame to get rid of the ammonium chloride.

[Edited on 23-8-2019 by Tsjerk]

wg48temp9 - 23-8-2019 at 10:35

Plain H2O2 is all that's needed. See the note below. Adding alkali will accelerate the decomposition of the peroxide as its less stable in high ph.

"Tungsten wire of common sizes can
be dissolved in 30% hydrogen peroxide
in convenient times (3 hours or less)
at 60” C., the optimum temperature.
Powder, especially of finer sizes, is
best dissolved in more dilute (10%)
reagent at room temperature to decrease
the violence of the reaction.
Most samples of commercially pure
tungsten dissolved completely; however, certain less pure grades gave a
visible residue or turbidity which was
found by qualitative spectrographic
analysis to consist of insoluble oxide or
hydroxide impurities, principally of
Si, with smaller proportions of AI, Ti,
Mg, and Fe. The turbidity was cleared
by the addition of sodium hydroxide.
Tungsten may be converted into
various of its compounds via dissolution in hydrogen peroxide. Evaporation of the solution and drying of the
residue at not over 100” C. produces
yellow crystalline pertungstic acid which
is freely soluble in hot water. Higher
drying temperatures (ca. 180” C. and
higher) convert the pertungstic acid
to tungstic acid, or anhydrous tungsten
trioxide, which may be reacted and
dissolved in accordance with their
known properties. Alternatively, the
alkali tungstates may be produced by
addition of alkali hydroxide to the
pertungstic acid solution, followed by
boiling to expel both the free and combined hydrogen peroxide.
Hydrogen peroxide as a solvent for
tungsten has the advantages over the
commonly employed nitric-hydrofluoric
acid mixture of not requiring platinum
ware, and of not introducing a difficultly removable anion. "

From: Murau, P. C. (1961). Dissolution of Tungsten by Hydrogen Peroxide. Analytical Chemistry, 33(8), 1125–1126. doi:10.1021/ac60176a021

stamasd - 23-8-2019 at 13:07

See that's the kind of information that I could have used before starting.

^^
What I found by myself led me to believe that H2O2 alone, or even with ammonia, wouldn't be active enough.

diddi - 23-8-2019 at 23:27

i see you are using TIG electrodes as your source. i use just 50% H2O2 to dissolve the electrodes, but i collect the Thorium. the W is just a biproduct for me (but i do keep it). They take weeks to dissolve btw.

stamasd - 24-8-2019 at 03:00

I don't have access to 50% H2O2, the max I can get is 27%. Yes, TIG electrodes; the one I'm using is 2% lanthanated because I have a lot of scrap pieces of that (I have thoriated too but no scrap pieces of that and I'd rather not break a good electrode).

I have restarted another experiment last evening with the remaining tungsten piece, 20ml H2O2 and a few drops of ammonia at RT. Haven't checked on it today yet, but yesterday there was just a little gas generation around the tungsten.

AJKOER - 24-8-2019 at 05:39

Per this study https://apps.dtic.mil/dtic/tr/fulltext/u2/a358781.pdf Tungsten can be made to corrode, at least to a limited extent, in the presence of 1% aqueous Na2SO4 (or, apparently better, is NaCl, see https://books.google.com/books?id=Uuz1Qgi3qLwC&pg=PA18&a... ), relative to select alloys (for example, Cu(70%) /Zn(30%) and one alloy containing a small amount of Carbon), per Table 3.

As such, I would try (pending a cheap available W source, most likely the filament in an incandescent bulb) a galvanic cell approach employing Tungsten metal (low surface area) versus Carbon (as graphite with more relative surface area) with household H2O2 in an electrolyte of sea salt in a microwave assisted prep.
------------------------------------------------------------------------------------

Interestingly, came across this health warning sheet on Tungsten at https://www.atsdr.cdc.gov/toxfaqs/tfacts186.pdf .

[Edited on 24-8-2019 by AJKOER]

Tsjerk - 24-8-2019 at 06:58

Quote: Originally posted by AJKOER

Per this study https://apps.dtic.mil/dtic/tr/fulltext/u2/a358781.pdf Tungsten can be made to corrode, at least to a limited extent, in the presence of 1% aqueous Na2SO4 (or, apparently better, is NaCl, see https://books.google.com/books?id=Uuz1Qgi3qLwC&pg=PA18&a... ), relative to select alloys (for example, Cu(70%) /Zn(30%) and one alloy containing a small amount of Carbon), per Table 3.

As such, I would try (pending a cheap available W source, most likely the filament in an incandescent bulb) a galvanic cell approach employing Tungsten metal (low surface area) versus Carbon (as graphite with more relative surface area) with household H2O2 in an electrolyte of sea salt in a microwave assisted prep.
------------------------------------------------------------------------------------

Interestingly, came across this health warning sheet on Tungsten at https://www.atsdr.cdc.gov/toxfaqs/tfacts186.pdf .

[Edited on 24-8-2019 by AJKOER]

It is perfectly fine for everyone to completely ignore this post as irrelevant and imaginary.

AJKOER - 24-8-2019 at 09:36

Quote: Originally posted by Tsjerk

......
It is perfectly fine for everyone to completely ignore this post as irrelevant and imaginary.

Perhaps this is a correct assessment, but I have already secured two W filaments, and may be able to soon test my claim.

To be clear, I have to exceed only "to a limited extent".
--------------------------------------------------------------

Another point, working with NH3/O2 or H2O2 is, in my assessment, not a basic piranha reaction, actually electrochemical (and a bit more) in nature. See related chemistry with copper, ammonia and O2 (or H2O2) at https://www.sciencemadness.org/whisper/viewthread.php?tid=15... and elsewhere on this forum.

Implication of this point is adding NH4Cl (or the sulfate) likely improves the reaction rate as in the case of copper.

I believe this ammonia/H2O2 plus NH4Cl path to be the best path.
---------------------------------------------------------------------------

OK, did a quick run with an electric bulb filament based on a bleach battery (but I acidified NaOCl with NaHCO3 forming HOCl) using a graphite rod as an electrode.

Unexpected result (much bubbling at the carbon rods in a solution of HOCl, but only after 90 seconds of microwave heating) suggesting that electric bulbs are tested with a small amount of air presence in the bulb. This implies to me that the W wire actually has an oxide coating, which actually makes the filament more noble than carbon.

The picture here appears to confirm my thoughts https://en.m.wikipedia.org/wiki/Incandescent_light_bulb and is not the picture of the metal displayed here https://en.m.wikipedia.org/wiki/Tungsten .

Apparently, I will have to find another source of Tungsten metal.

[EDIT] Video of reaction attached below. I will place all other comments and experimental results in a new thread.

[Edited on 25-8-2019 by AJKOER]

AJKOER - 24-8-2019 at 14:02

Short clip showing, apparently in the presence of tungsten oxide filament, gas bubbles evolving from the carbon rods in HOCl after removal from the microwave following 3 treatments of 30 seconds (with the prior 30 second intervals indicating no reaction).

Attachment: W_Carbon_Run.mp4 (2.3MB)
This file has been downloaded 492 times

[Edited on 25-8-2019 by AJKOER]

Panache - 24-8-2019 at 18:58

Apologies for the late sub back into the game...life, other priorities ensued, some observations.
I think the 60C with the ammonia only reduces the extent of the reaction when dealing with tungsten samples of low surfAce are. Of the two I ran (one at rf one at 60) the one start bubbles continuously, although one would never say vigourously. However it ran (ie 'bubbled off the electrode)far longer than the hot one and ulitimately dissolved more tungsten by weight.
I'm about work them up by Tserk, thanks for the reference and even with this limited experience it would that it rings true.

AJKOER - 25-8-2019 at 05:06

Panache:

You can stop making comments supporting my postulated electrochemistry (a new thread/pictures/results soon available exploring a related, but I suspect inferior to the ammonia cell, which I proposed above is likely one of the better path).

Anodic corrosion apparently works more efficiently when the relative surface area of the metal anode, here Tungsten, in my proposed ammonia/H2O2 or O2 electrochemical cell scheme is small. To repeat, when the relative electrode surface area ratio is SMALL (hence low surface area metal Tungsten), say, to the solution's surface area contact with air, the anodic corrosion accelerates.

The cell may create an electric current and associated solvated electrons, e-(aq), so with preferentially added NH4Cl or the sulfate via:

NH4+ = H+ + NH3

And, the removal of H+:

e-(aq) + H+ = .H

the latter two reactions occurring in solution could supply added NH3 (and I would further guess that starting with the highest ammonia concentration is likely not optimal and having the right NH4+ concentration is a factor, see https://www.sciencedirect.com/science/article/abs/pii/S03043... which discusses optimal leaching factors).

[Edited on 25-8-2019 by AJKOER]

stamasd - 30-10-2019 at 07:23

Quick update on this after a couple of months.
I had originally set up an experiment with a 2g fragment of tungsten, plus H2O2 (27%) and NH4OH (30%). Then I forgot about it.
I found that beaker again last week. All liquid had evaporated (must have been 20ml or so) and the tungsten rod, which initially had showed just a little mass loss, had completely disappeared. What was left was a yellow mass at the bottom, which I thought at first to be crystalline. At closer examination they aren't crystals but irregular pieces that appear like dried-out and broken jello. They are translucent, and I assume they are a mix of ammonium paratungstate (which is white) and tungsten trioxide and/or tungstic acid (which are bright yellow).

I haven't worked that up yet. I'm surprised how well it worked to dissolve all of the tungsten.

But in the meantime I set another experiment up, larger scale. In 2 flasks I added 20g of tungsten pieces each, followed by 25ml of 27% H2O2 each, and in one of them also 5ml of 30% NH4OH. Initially the flask with ammonia heated up a lot and started boiling, almost boiling over despite being a fairly large (500ml) flask. The color of the liquid changed to pale yellow. The other flask with only H2O2 only showed very little gas bubble formation at the surface of the tungsten, which ceased by the next day.

Next morning the yellow color in the ammonia flask had faded, both flasks had a very slightly yellow solution.

Left with ammonia, right without.

I added 6ml H2O2 to each, and in the flask with ammonia the yellow color suddenly intensified. No color change in the one without ammonia. Also no gas bubble formation in either.

I kept the same pattern daily; each day adding 6ml H2O2 in each flask, and every other day 1ml of NH4OH to the left flask only (I keep that one covered with parafilm to minimize ammonia loss). The same phenomenon kept occurring, loss of color by the next day and color intensifies in the left flask as soon as more H2O2 is added. The color in the right flask doesn't change upon addition, just slowly becomes more yellow from day to day.

Flasks on the 3rd day before and after additions:

I noticed that in the left flask (with ammonia) a very fine, dust-like, off-white precipitate has started to form. The precipitate does not change color when I add more H2O2.

No pics on day 4. Here are pics from today, which is day 5. The color changes have become a lot more intense, and there is a lot more of the precipitate.
Before:

And after H2O2:

So far each flask has received 55ml of 27% H2O2, and the left flask only got a total of 8ml 30% NH4OH.

I will keep adding reagents in the same pattern for another 5 days then will just observe.

[Edited on 30-10-2019 by stamasd]

stamasd - 31-10-2019 at 04:34

Pics from this morning.

Before:

Added ammonia (1ml) in the left flask only; no color change.

Added H2O2 (6ml) to both; yellow color reappears in the left flask.

Herr Haber - 31-10-2019 at 04:47

Hello stamasd !

I have tungsten in almost every possible form except plate / foil.
I have all the reagents you used in your experiments and more.
Most important: I have a little bit of space

So if there's an experiment you are not doing for lack of space, time or reagents let me know.

stamasd - 31-10-2019 at 05:51

I have space and reagents, time however is another matter... I happen to have some time right now though. Thanks for the offer!
The next thing I plan to try is to see if I can speed up the tungsten dissolution using electrochemistry. But that won't be until this experiment is over.

wg48temp9 - 31-10-2019 at 07:39

Molybdenum can also be dissolved with H2O2. If the Mo is in excess when the H2O2 has all reacted to form Mo acid the Mo metal then reduces the acid to a Mo blue.

See the image below from a thread titled "blue tungstic or Molybdic acid ?"

files.php.jpg - 12kB

stamasd - 1-11-2019 at 16:53

Hmm. No pics today, but...
I spent a little time looking at the 2 flasks and I noticed something interesting, which I had not observed until today.
The second flask, the one with only H2O2 and no ammonia, which does not exhibit any color changes when I add more reagent to it, may be in fact more interesting than the other one.
What I noticed is that in that second flask there is now definitely less metallic tungsten remaining compared to the other one. They both started with the same amount (20g) of fragments from 2mm tungsten rods.
Well, the rods in flask 2 are now significantly thinner than the ones in flask 1. I'd say they're no more than 1mm in diameter.

If I'm right, they should completely dissolve in the next 3-4 days.
We'll see.

[Edited on 2-11-2019 by stamasd]

stamasd - 2-11-2019 at 11:29

Okay update with some very interesting information and some calculations.

I pulled out all of the remaining metallic tungsten in both flasks and dried it separately. I was right, the pieces in the H2O2 alone are much thinner and lighter. Here are some photos for comparison.

Left, H2O2+NH4OH; right, H2O2 alone.

Same but closer look.

I then weighed the metal separately. The initial weight was 20.0g for both samples.

The metal from H2O2+NH4OH:

The metal from H2O2 alone:

So in 8 days, the H2O2+NH4OH dissolved 1.3g, whereas the H2O2 alone dissolved 13.8g. Or 0.1625g/day and 1.725g/day respectively.

The NH4OH is actually detrimental to the rate of reaction, slowing it down more than 10x.

fusso - 2-11-2019 at 12:28

I wonder if you only add a little NH3 after considerable amount of W is oxidsed, to dissolve most of the oxides, will the dissolved tungstates affect the rate?
Also how does aq NH3 slow down the rxn? Is it due to the extra water or NH3/tungstates decomposing the H2O2?

stamasd - 2-11-2019 at 12:37

I don't have an explanation, these are just my observations. If I were to guess, I'd venture to say that the rate of decomposition of the H2O2 is probably higher in the presence of ammonia. Also apparently no ammonia is needed to dissolve the oxidation product as the solution in the flask with only H2O2 is perfectly clear, if only slightly yellowish but with no visible precipitates.

[Edited on 2-11-2019 by stamasd]

wg48temp9 - 2-11-2019 at 15:50

H2O2 is oxidising in acid but in alkali it can be reducing. So I guess if you put too much ammonia in the H2O2 will not oxidize the W or the oxidation rate will be reduced.

Tsjerk - 3-11-2019 at 04:01

How would hydrogen peroxide be reducing? I can only imagine it being oxidized by something like fluorine or chlorine trifluoride.

stamasd - 3-11-2019 at 06:16

2 KMnO4 + 3 H2O2 --> 2 MnO2 + 2 KOH + 2 H2O + 3 O2

Mn(VII) is reduced to Mn (IV) by the oxygen in H2O2, which itself is oxidized to O2.

[Edited on 3-11-2019 by stamasd]

Mesa - 8-11-2019 at 05:27

Fastest/easiest way to dissolve W... Make a thermite-like rxn with NaOH or KOH and dissolve the resulting ash in H2O2.
Alternatively, if you can get it as a powder, ignite a pile of it in open air with a lighter flame, and spread the pile around while it slowly turns red hot, then cools to a fluffy-er yellow WO3 powder.

Reaction is self sustaining with any fine grit W or carbide powder.
Note: if carbide is used, you might have to separate any trace cobalt or molybdenum compounds. Not quite sure though. I only saw it happen in workplace fuckups at my old job. Never tried igniting WC intentionally myself

stamasd - 8-11-2019 at 15:34

I don't have a cheap source of powdered W or WC. My W is in the form of broken 2mm rods. I'd also prefer if the resulting WO3 is not contaminated with other chemicals if possible.

Also, as someone else says, I like to explore the science of it. I haven't seen the color change yellow/colorless/yellow mentioned anywhere else, nor the kinetics of W dissolution in acidic vs alkaline conditions.

Mesa - 8-11-2019 at 15:53

You could probably just dissolve the ash in conc. NaOH to get sodium tungstate and whatever Na products formed, then treat with HCl to precipitate WO3, all Na products remain soluable.

Without trying it myself I'm not so sure it would dissolve. I'm pretty confident that the HCl step would effectively isolate WO3 product though. No other products could form from the thermite form from HCl except NaCl and H2WO4(decomposes to WO3 in H2O) right?

[Edited on 8-11-2019 by Mesa]

EDIT: There's a pretty easy method to purify any soluable tungsten compound.

Make the ammonium tungstate from HCl + tungstate, separate crystals and redissolve in ammonia solution, evaporate, then fire in open crucible til ammonia stops evolving.
This should work to purify the WO3 you've produced also. In case there is any impurities from electrodes.

I'll post patents when I get on home computer.

[Edited on 9-11-2019 by Mesa]

stamasd - 9-11-2019 at 03:00

But then I wouldn't be investigating which compound I get from dissolving tungsten in plain H2O2, as I plan on doing soon. You'd think it would be either WO3 and/or H2WO4, but these are bright yellow and completely insoluble. Whereas what I got is a perfectly clear solution that is almost colorless, only a very pale yellow. And that despite having almost 20g dissolved tungsten per 100ml. Whatever is the tungsten species I have (and I'm open to suggestions as to what it is) it's highly soluble.

Mesa - 14-11-2019 at 04:40

The solution will precipitate out WO3 over time. I'm not sure what the dissolved species is, except that its unstable in solution and eventually degrades to insoluble WO3 if left alone.

stamasd - 14-11-2019 at 19:51

Actually I'm not so sure about that. I haven't touched the flasks for 4 days and their appearance is unchanged as far as the liquid is concerned, only the little bit of tungsten left still undissolved is getting progressively smaller. The liquid is still perfectly clear and very slightly yellow. In a few places it has splashed on the walls and the drops evaporated leaving behind a white deposit with no yellow tinge at all.

[Edited on 15-11-2019 by stamasd]

Mesa - 14-11-2019 at 20:59

Try adding a few drops of HCl.

stamasd - 15-11-2019 at 05:31

I'll do that in a small sample when all the tungsten is dissolved. The rate of dissolution has slowed down considerably but there is less than half a gram left so it shouldn't take more than a few days.

stamasd - 20-11-2019 at 16:35

Only 2 small pieces of tungsten remain, about 0.5mm thick and 3-4mm long each.

Mesa - 21-11-2019 at 23:29

Likely some kind of peroxo tungstic acid such as described in this link

Evaporating the water at no more than 100*c will isolate the pertungstic acid. Heating above 100*c will decompose it to WO3.

Could explain the precipitate on the side of the flasks?

EDIT: This link seems to be more relevant
https://www.researchgate.net/post/How_can_I_get_a_clear_solu...

[Edited on 22-11-2019 by Mesa]

stamasd - 22-11-2019 at 07:56

I can't use the first link "you have reached a page that's unavailable for viewing"
As for the second link, it describes the second part of the experiment (the left flask in some of the pictures above) with H2O2+ammonia, which showed a dissolution rate 10x slower than plain H2O2.

But yes, it's probable that what I'm getting is an adduct between WO3 and H2O2 aka peroxopolytungstic acid.

[Edited on 22-11-2019 by stamasd]

stamasd - 26-11-2019 at 13:52

The tungsten dissolution is complete. I have set the solution to evaporate.

Fery - 16-2-2020 at 02:55

I've just finished some experiments and my opinion is that ammonia speeds up H2O2 decomposition.
You can buy unstabilised H2O2 (I believe food grade and maybe pharma grade) but most of it used in chemistry is stabilised (H3PO4, Na2[Sn(OH)6] etc.)
There is no need to have alkaline environment to dissolve W as pertungstic acid is formed in excess of H2O2 which is better soluble than poorly soluble tungstic acid. Alkaline environment just speeds up H2O2 decomposition.