Chemistoftheelements

Yes, I've figured since that its was some tungstic oxide hydrate. Wasn't at all sure at the time, though.

It surprises me just how lasting these names are, and even chemical manufacturers occasionally perpetuate this confusion, as I’ve witnessed with a bottle of “tungstic acid”.

This admitedly isn't the best example, but it concludes on a probable structure, i.e. there still remains an uncertainty. However this may well be explained by the difficulty of producing material which is sufficiently crystalline and uniform in composition, as your second link hints, thanks. I take your point about the molecular formulas, but the H3AsO4. 1/2H2O in the wikipedia link I took to imply a compound with As single bonded to 3 OH groups, and double bonded to an O, analogous to an orthophosphoric acid unit, apart from the ½ H20 bit. I read it this way as wikipedia phrases it as the "hemihydrate form"(of arsenic acid). The structure As2O5.4H2O which I suggested is hydrous As2O5, i.e with no OH groups. This highlights the confusion, and part of my frustration, that hydrous acids and hydrous oxides are banded together with the same name- cf "perrhenic acid": http://en.wikipedia....Perrhenic_acid.

Precisely. But that would be a challenge....you thinking of giving it a crack?

For an example of what I mean: http://www.nrcresearchpress.com/doi/abs/10.1139/v72-106 if XRD etc, cannot answer the question of the structure of this compound, what can? Also, although I have limited knowledge of this sort of thing, the structure H3AsO4.½ H2O (http://en.wikipedia.org/wiki/Arsenic_acid) seems a little strange to me; is a water molecule shared between two H3AsO4 units? Do the XRD experiments on this compound (if any have been done) rule out As2O5.4H2O ?

There seems to be much uncertainty of the structure of the oxoacids of certain elements, particularly metalloids, but also of metals such as molybdenum, tungsten (i.e. "molybdic" and "tungstic" acids), and such terms as "perrhenic acid" and "rhenium heptoxide" are frustratingly used interchangeably. Orthosilicic acid certainly exists in solution, but when crystallized from aqueous solution under standard conditions, no-one really knows what happens- perhaps the acid polymerizes, perhaps a hydrous silica is produced. An example of this uncertainty can be illustrated by the compound "Arsenic acid" when crystallized; wikipedia* talks of the structure H3AsO4 ½ H2O, but how do we know that the structure isn't actually the stoichiometrically identical As2O5.4H2O? What are the limitations of Raman spectra, X-ray crystallography and other techniques for determining the structure of those substances which result from crystallization from aqueous solution of such acids? What my question is is this: How can we be sure of the structure of these substances who's structures have so far proved so elusive? * http://en.wikipedia.org/wiki/Arsenic_acid

All I know is that when I've neutralised Fe(II)+ containing solutions with aqueous ammonia, I've gotten a precipitate. Whether that precipitate is Fe(OH)2 or whether it's a hydroxide- ammonia complex, I don't know- the observation that some ammonia- containing complexes are soluble whilst others aren't makes it difficult to tell, and a frustration for me. I'd like to be able to predict theoretically how some complexes (and not just ammonia- containing ones) dissolve in water under standard conditions. Anyway, here's very likely what the ammonia- nickel chloride complex which I mentioned is: http://www.periodictable.com/Elements/028/index.html It's about the fifth picture from the bottom.

Thank you, that was interesting and informative. Yes, it's insoluble enough for the NiCl2 to be amost completely recoverable, as I recall. You can drive the ammonia off in a fume cupboard, which was, infact, part of the practical. I wasn't sure of the formula for a long time, but it's a very striking experiment, and the kind of demonstration which could be used more widely in schools and colleges to encourage people to participate in chemistry.

Hi, I remember performing a practical once whereby ammonia solution was added to a solution of Nickel(II) chloride, with a purple precipitate appearing. This has always stuck with me as an example of an insoluble complex, and noting that many other ammonia complexes are soluble, it has now got me wondering why some complexes of any kind dissolve whilst others don't. Could anyone who has more theory please recommend to me something which would help me understand why some complexes dissolve and others don't in water under standard conditions? I realise this may be a bit involved, so I was thinking of a textbook or an area of inorganic chemistry which covers this? Many thanks.

Thanks. Briefly, withought trying to break out into another thread, could you recommend something which would help me understand why some complexes dissolve and others don't ( e.g. Cu2+ ammonia complexes versus Ni2+ ammonia complexes), in water under standard conditions? Many thanks.

"And chemistoftheelements, would you be in a position to repeat your experiments with stirling silver and pure silver and with warmed 0.88M ammonia in a sealed container?" Hi Greg; after a bit of encouragement from 3 forum members, I've decided to stick around a bit. In answer to you question, not at the moment, but I may have access to a fume cupboard in the not- too distant future- hot ammonia is nasty stuff, so this is essential. What interests me most is that the results so far are strikingly different between fine silver and a high- silver alloy. Maybe it's Ag2S which forms on fine silver, and a mixture of copper compounds on sterling silver, with some silver compounds. As copper is higher up the reactivity series than silver, it may be that it preferentially corrodes and "protects" the silver, in the manner that the zinc in zinc plate on iron will do if this coat is scratched through to the iron. On the solubility of Ag2S, it's always been my understanding that the solubility of a substance per say is not a barrier to forming soluble complexes, but if Ag2S- ammonia complex in water is insoluble, then the tarnish layer on fine silver, if it does consist of Ag2S, will not be removed. For example, TiO2 is extremely insoluble in water, yet is dissolved by hydrofluoric acid, because the complex formed is soluble. Hypervalent_iodine, what's your take on this? Thanks.

In the previous thread, I posted some of my conclusions from an experiment I conducted on the removal of the tarnish formed on 99.99% fine silver, and on the alloy sterling silver. My intention at the time, was to add a little to the discussion based on my experience, not to further an argument. For the sake of science, I will add the results of a second experiment, using the same piece of 99.99%, i.e. fine, silver (as I don’t have a great deal of this available to me), and some more tarnished sterling silver, after treatment with 9% ammonia solution: 1) In the second experiment on the fine silver, there was no further removal of tarnish, and the original, unpublished observation of H2S evolution (which came as a surprise, as there’s no obvious mechanism for this to occur) was not replicated in the second experiment. 2) The sterling silver again, in the second experiment, markedly de-tarnished, to the extent that this could be used as an effective method of at least helping to completely remove the tarnish from sterling silver, at least of it’s more usual composition, in which copper is the alloying element. No H2S evolution took place. The only firm conclusions I reach from the above experiments are as follows: 1) The tarnish layer formed on fine silver is of a different chemical nature of that formed on sterling silver. The species in question were not identified by the author of these experiments. 2) More study is warranted into the differences in the observations between the first and second experiments on fine silver, and into the nature of the tarnish films formed on fine, and sterling, silver. Onto the topic of humility and politelness now, and may I add that after my original post on the topic, which was intended for the purpose of encouraging scientific discussion, I gained the distinct impression that John Cuthber had formed some sort of grudge, as evidenced by his unwarranted behaviour in the Neodymium chloride thread in the Inorganic chemistry section. Although I don’t necessarily condone the methods which Greg used in this thread, it is clear that he was provoked by the same behaviour in the previous, now closed, thread on silver tarnishing, as I was in the Neodymium chloride thread. I can’t say I blame him, as I also have a low tolerance for arrogance and lack of humility. It has been demonstrated by further discussions in this thread (after the conversation between Greg and John Cuthber, that Greg is ammenble to civilised discussion with others who use a more civil approach; it is largely dependent on the manner in which any correcting and commenting is done, and this applies to most people. Please refer to my point about politeness on page 2 of the Neodymium chloride thread in Inorganic chemistry). I joined this forum for the sake of discussing chemistry, and I really am not interested in arguing with people who cannot grasp the basic rules of social interaction. Regretfully for me at least, I’m out of here.

OK, I goofed whilst numbering the points- there are infact 9 of these! Of course, if Pr(IV) reduces instantly, i.e. is extremely unstable relative to Pr(III), then you will get only one precipitate, containing both metals. If this is the case, which it may well be, then you are no worse off than you were to begin with.

Hi elementcolloctor1, yes, I know the feeling, it can be a little hard to keep track of these things. I think we're getting a little confused her, so I think it might be helpfull if I re- post a conversation we had about 3 days ago. Note that as far as I can tell, whether the 8- point method I suggest works is a matter, I think, of whether you have any Pr in there in the first place, and just how long Pr(IV) will hang around in acid solution; you MIGHT, might be lucky and get a small amount of Pr(IV) precipitate, but as John Cuthber has rightfully pointed out, it is pretty unstable. Here was our ealier conversation: elementcollector1, on 31 December 2011 - 12:45 AM, said: Hah, speaking of sodium bicarbonate, I just lost a large beaker to the stuff. Poor thing has a massive chunk missing. Anyway, I know oxalic acid is used as wood bleach, so I'll look into that. I'd like to go all the way to Nd and maybe Pr metals if I could, but we'll see. Process of getting to Nd from neomagnets, in short: 1) Dissolve in acid. 2) Add oxalic acid or salt to precipitate insoluble oxalates. 3) Ignite in air to produce oxides. (for separation of Pr) 4) Somehow convert the Pr6O11 that formed into PrO2. 5) (for separation of Pr) Place in 5% acetic acid? I read below on Sciencemadness that PrO2 is insoluble in 5% acetic acid, and apparently Nd2O3 is. Elementcollector1, 1) Happy new year! 2) To answer this and your latest question: There’s some good logic to this, but I’m not sure that this would work, for several reasons: 1) At the oxide stage, the Pr would be in solid solution in the Nd2O3, and probably being only a minor constituent , would be prevented from oxidising to higher than Pr(III). In other words, the Nd(III) would stabilise the Pr(III), in the same way that Th(IV) stabilises U( IV) in high ( and low) U thorianites. 2) Oxidation states of Pr higher than 3 would be unstable in acid solution- think CeO2, so adding acetic acid would probably cause what little Pr in a >III oxidation state to revert pretty quickly to PrIII and go into solution. 3) Lanthanoid (and related oxides) don’t easily dissolve in acid once they’ve been ignited. Working on a theory that whether a metal oxide dissolves or not depends on the ionic radius, the coordination number of the metal, and possibly it’s affinity for oxigen, I produced some Yb2O3, Sc2O3 and Sm2O3 and found that all 3 were very reluctant to dissolve; after ignition at 1000 degrees C (and then cooling), only the Sm2O3 dissolved, but very slowly. If, however you heat them only minimally ( just enough to ignite the oxalate in your case), then you will probably succeed. 4) By using H2O 2, you may end up with a product containing O-O bonds, which would be bad, because such a compound may prove thermally dangerously unstable. You would have to work on very small quantities to find this out, or alternatively oxidise with something else. What I would suggest is the following, which circumvents problem 1, but not the others, so it’s not guaranteed to work: 1) Dissolve in acid 2) Add H2C2O4 to precipitate Nd and Pr and separate them from Fe, etc. Add NH3 to raise pH to increase precipitation if necessary, or dilute with water. 3) Gently heat precipitate to create Nd/Pr oxides. 4) Dissolve oxides in H2SO4, NOT HCl, as this is reducing. 5) Oxidise this resulting solution to produce Pr2 O(SO4)3(?), a higher basic sulfate of praseodymium which will be insoluble. 6) Filter. You have now recovered Pr. 7) Ignite your P2 O(SO4)3 7)To remaining solution, add NH3 to precipitate Nd(OH)3. 8) Ignite the Nd(OH)3. From here, you’ll have to figure out a way to reduce the Pr and Nd oxides. Good luck, and let us know how you get on. In the mean time, I’ll try to figure out how to create subscripts!

Hi elementcollector1 "Please explain. When did the oxide crystals form again?" This is at the stage after you ignite your oxalate, which you were wondering how you should convert, in terms of it's praseodymium content, to PrO2, and you suggested H2O2..<br style="mso-special-character:line-break"> <br style="mso-special-character:line-break"> You wrote: “ Hah, speaking of sodium bicarbonate, I just lost a large beaker to the stuff. Poor thing has a massive chunk missing. Anyway, I know oxalic acid is used as wood bleach, so I'll look into that. I'd like to go all the way to Nd and maybe Pr metals if I could, but we'll see. Process of getting to Nd from neomagnets, in short: 1) Dissolve in acid. 2) Add oxalic acid or salt to precipitate insoluble oxalates. 3) Ignite in air to produce oxides. (for separation of Pr) 4) Somehow convert the Pr6O11 that formed into PrO2. 5) (for separation of Pr) Place in 5% acetic acid? I read below on Sciencemadness that PrO2 is insoluble in 5% acetic acid, and apparently Nd2O3 is.”<br style="mso-special-character:line-break"> <br style="mso-special-character:line-break"> Then later on: “My main problem with this is how to convert the Pr6O11 to PrO2. Would hydrogen peroxide work? Excess heating? From the same forum as earlier, members have mentioned that on ignition, praseodymium oxalate turns brown, but I'm not sure if this is PrO2 (in which case, yay!) or something else. Nd (III) cannot be oxidized any further by igniting in air, so that's safe.” This was before I came up with my prodedure suggestion. Good luck! "<br style="mso-special-character:line-break"> <br style="mso-special-character:line-break"> " This bit, of course, is complete nonsense. I think this sort of thing is happening because I'm copying from microsoft word before I post.