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woelen
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What you have prepared could well be some anhydrous salt of titanium(III), which is nearly insoluble.
I myself have had similar experiences with vanadium, chromium and nickel.
I heated V2O5, added to sulphuric acid. When this is done, then at first the V2O5 dissolves, giving a deep red solution. On further stronger heating
at a certain point, gas is produced (this must be oxygen) and the red color of the V2O5 disappears. After a fairly long time of heating, a grey solid
remains, which is insoluble in sulphuric acid and also insoluble in water. According to some research I did on this solid, the insoluble material is
anhydrous VOSO4. Hydrated VOSO4 is deep blue, much like CuSO4.5H2O, but somewhat darker.
A similar thing exists for chromium. If you heat the dark purple chromium sulfate or dark purple chrome alum in concentrated sulphuric acid, then you
get a green solid, which does not dissolve in water, nor in acids.
Finally, I had a similar experience with nickel sulfate. The anhydrous salt does not dissolve in water, but if it is left in contact with water, then
after a few days, it is rehydrated and then it does dissolve.
Maybe you have a similar effect with your titanium salt. It may be anhydrous.
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bismuthate
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Thanks deltaH.
blogfast what happens when the blue titanium containing percipitate is mixed with a strong oxidising or reducing agent and heated? That may give us
some idea of the chemical porperties.
EDIT: woelen does this same thing happen to occur with niobium?
[Edited on 19-10-2013 by bismuthate]
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deltaH
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He's done the strong oxidant part, heated with peroxide and acid which yielded a bright red peroxide product. See he's test tube pics.
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bismuthate
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I did see that however that mixture also contained HCl which may have helped in the reaction. Maybe an oxidiser like a chlorate would react in a dry
mixture.
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blogfast25
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Quote: Originally posted by deltaH  |
I suggest you repeat the test with 6g Ba(NO3) to be on the safe side.
[Edited on 19-10-2013 by deltaH] |
It has occurred to me. I'll explain tomorrow.
I think it is too. Proving this may be harder.
| Quote: Originally posted by bismuthate | | I did see that however that mixture also contained HCl which may have helped in the reaction. Maybe an oxidiser like a chlorate would react in a dry
mixture. |
Bismuthate: it was shown beyond reasonable doubt that it is a Ti(III) compound. These respond very positively to oxidisers. Heating with oxidisers
will create Ti(IV) based compounds.
[Edited on 19-10-2013 by blogfast25]
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blogfast25
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Even though I had taken out a little insurance for the sulphate/chloride test by holding back a bit of the Ba(NO3)2 solution and then testing the
filtrate with it (it tested negative), the test for chlorides was repeated as follows.
1.0 g of product was treated with 10 ml of 33 % NH3 and the Ti(OH)3 filtered off. The clear filtrate was acidified with glacial
acetic acid, diluted to about 100 ml and heated up to near BP. To this was then added a 50 ml solution of 5 g Ba(NO3)2 in water, also
hot. The hot slurry was then filtered and allowed to cool.
To it was added a strong solution of about 5 g AgNO3 in about 20 ml of water and a strong precipitation occurred:
The quantity of precipitate is in line with yesterday’s observation and is difficult to explain other than by bound chloride in the product.
The remaining material (about 7 g) is now being dried carefully at low heat till constant weight, for use in more quantitative tests.
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deltaH
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Okay! Yes it looks nearly certain now to contain substantial amounts of chloride. Now to determine if this is an oxytitanium(III) bisulfate/chloride
as I suspect. I don't think it's a simple salt like Ti(HSO4)3 or TiCl3 which would be soluble, so it's reasonable to suspect that it's a polymeric and
probably amorphous -O-Ti(III)-O- material with mixed HSO4- and Cl- as counter ions. This would explain its insolubility and relative stability.
[Edited on 20-10-2013 by deltaH]
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bismuthate
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since it does contain chloride I wonder if it will react with H2SO4. Also you still haven't tested it with a dry reducing agent/ product mixture. You
may want to test its solubility in different solvents.
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blogfast25
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| Quote: Originally posted by bismuthate | | since it does contain chloride I wonder if it will react with H2SO4. Also you still haven't tested it with a dry reducing agent/ product mixture. You
may want to test its solubility in different solvents. |
Testing the product with oxidisers (that is what you meant, I suppose?) is useless: we know what happens.
Re. solvents, I’m more interested in volatile, non-reactive anti-solvents, for drying purposes. Acetone and glacial acetic acid seem to fit the bill
so far.
Thanks deltaH for your comments.
***************
Another batch of the product was attempted, this time leaving out the K2SO4. No precipitate was obtained. So 4.4 g of K2SO4 was then added anyway, and
some precipitate formed but much less than in the initial run.
Sensing that perhaps the potassium wasn’t important but the sulphate in the K2SO4 was, another run was made, this time with no K2SO4 but 10 ml
(instead of 5 ml) of 96 w% H2SO4 added when the titanium dissolution was all but over. No precipitate formed whatsoever, indicating again the K2SO4
plays some part in its formation, but after a bit of cooling something interesting happened: over the course of about 5 minutes a deep purple
crystalline product slowly separated out, eventually filling most of the volume:
The photo doesn’t do the colour justice, my camera is colour blind (possibly due to TL lighting in my lab): it’s more like a lighter shade of
chromium alum and very different from the sandy light blue precipitate.
Also interesting is the fact that the supernatant liquid is completely colourless, indicating it contains no Ti(III) (but Ti(IV) cannot be excluded).
I have not tested the crystalline matter yet but I’m hoping it is water soluble. With there being more than enough sulphate present, this could even
be a simple Ti2(SO4)3 hydrate. That would raise hopes of a Ti(III) alum again. But isolating and drying this material may prove challenging. Paper
filtration media won't do here...
[Edited on 21-10-2013 by blogfast25]
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Poppy
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Have you tried simply evaporating the mixture? I had a table showing double salts compatibility but its not on this computer, also, it may be the case
that the proportion of the reagents is exceeding/ requesting some molar ratio, thus failure is prominent.
Have you nuts to wait a month or so and make some evaporites??
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blogfast25
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| Quote: Originally posted by Poppy | Have you tried simply evaporating the mixture? I had a table showing double salts compatibility but its not on this computer, also, it may be the case
that the proportion of the reagents is exceeding/ requesting some molar ratio, thus failure is prominent.
Have you nuts to wait a month or so and make some evaporites?? |
Evaporating a solution containing Ti(III) would have to be done under inert gas blanket or in vacuum. Not that simple... It's a 'last resort', IMHO.
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bfesser
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Evaporating under a vacuum is trivial. Jam in a one-hole stopper with a tube connected to vacuum, and swirl gently in a warm water bath.
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ScienceSquirrel
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It seems that caesium titanium alum is known;
http://scitation.aip.org/content/aip/journal/jcp/107/20/10.1...
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blogfast25
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Aha. 'caesium titanium alum' yields quite a few references.
Thanks for that.
| Quote: Originally posted by bfesser | | Evaporating under a vacuum is trivial. Jam in a one-hole stopper with a tube connected to vacuum, and swirl gently in a warm water bath.
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Trivial only if you have vacuum on tap. My mini vac pump probably pulls only 70 % vacuum but that might be enough. And there has to be some kind of
cold trap between the pump and the evaporating liquid, certainly if you're evaporating acidic solutions, like a second vac flask in an ice bath...
[Edited on 23-10-2013 by blogfast25]
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blogfast25
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I repeated the experiment with 50 % H2SO4 and late addition of K2SO4 from above but with (NH4)2SO4 instead. So, 2.4 g of Ti powder, 15 g of water, 15
g of 96 % H2SO4 and addition of 3.3 g of (NH4)2SO4 when gas evolution started to wane.
After cooling to RT it looked like this:
The light blue precipitate is again intrinsically blue (not just coloured by the deep blue supernatant liquid) and filters to a filter cake of this
colour:
Interestingly, assuming this precipitate is the same as the one reported above, this was obtained in the absence of chloride anions.
Another test will now concentrate on trying lower H2SO4 concentrations.
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blogfast25
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Finally some progress on how to dissolve titanium in sulphuric acid to quite concentrated solutions of Ti2(SO4)3 (III), without a sandy precipitate
appearing. The problem of precipitation was solved by reducing the amount of sulphate in the solution.
I found that using approx. 40 w% H2SO4, with an acid reserve of 50 % of the stoichiometrically required amount of H2SO4, at reflux leads to almost
complete dissolution of the Ti powder, without any premature precipitation taking place.
For example, in the latest run 2.4 g Ti powder, 16 g of water and 12 g of 96 % H2SO4 were combined and refluxed until hydrogen evolution had all but
ceased (this took about 1 hour). This resulted in a deep purple but transparent solution which after cooling was carefully decanted off and weighed as
28.6 g of solution, just under 25 ml.
A small amount on unreacted metal was left behind and it was isolated, rinsed with DIW and dried. It weighed only 0.05 g.
This means that based on molar masses this solution is about 33 w% in Ti2(SO4)3 (III). Sigma Aldrich market a solution of about 45 w% Ti2(SO4)3 (III)
(nearly £100/100 ml, 99.9 % trace metals, exclusive of tax or shipping) .
To the cool solution about 3.3 g of solid (NH4)2SO4 was added (the stoichiometric amount for a Ti(III)/ammonium alum. This dissolved into the solution
easily, with very gentle heating. It’s now cooling on an ice bath where it will be kept overnight.
My hopes of finding NH4Ti(SO4)2.12H2O crystals aren’t very high though. Looks like I’ll have to break into my 10 g stockpile of CsOH…
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deltaH
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My thinking as well blogfast, the Ti ion is so small you need a bigger M(I) to stabilise that sucker.
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bismuthate
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deltaH do you think it would it be possible to use Rb instead of Cs?
blogfast25 be carefull with that cesium hydroxide it's dangerous and expensive.
Edit from a seach it seems that Rb Ti sulfates do exist.
http://pubchem.ncbi.nlm.nih.gov/summary/summary.cgi?cid=4413...
[Edited on 27-10-2013 by bismuthate]
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woelen
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Another less expensive option may be to use salts of the N(CH3)4(+) ion, tetramethylammonium. This is a large ion. The chloride or bromide of this
salt sometimes is offered on eBay.
I am not sure whether this eases the formation of an alum, but as this is a large ion, it may help somewhat.
[Edited on 27-10-13 by woelen]
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deltaH
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Oh my... just looked up ionic radii of these metals on wiki, turn out Ti(III) is 81 pm versus Al(III)'s 67.5, so taking K/Al alum as the base case,
looks like you need something slightly smaller, not bigger!!!!
I'd say go for sodium at 116pm compared to potassium's 152pm
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blogfast25
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| Quote: Originally posted by deltaH | Oh my... just looked up ionic radii of these metals on wiki, turn out Ti(III) is 81 pm versus Al(III)'s 67.5, so taking K/Al alum as the base case,
looks like you need something slightly smaller, not bigger!!!!
I'd say go for sodium at 116pm compared to potassium's 152pm
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Hmm. There's a KCr(III) alum, surely the ionic radii of Cr3+ and Ti3+ won't be a million miles apart?
Also, the CsTi(III) alum appears to be a reality, going by multiple references to this material and Cs+ is decidedly larger. Of Al the NH4,
Na and K alums are known (not sure about the Rb and Cs ones), so ionic radii don't seem to be a great issue there.
Personally I'm convinced this is a problem of not exceeding the solubility limit at near 0 C temperatures.
Thus evaporating solvent from the solution is a possibility but I think it needs to be done at fairly low temperatures because high hydrates don't
like high temperatures, as evidenced also by the low melting point of several alums: for these the crystalline hydrate cannot exist above that
temperature.
[Edited on 27-10-2013 by blogfast25]
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deltaH
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Indeed, well if CsTi(III) is known, then you can't exactly argue with that, now can you 
| Quote: | | Thus evaporating solvent from the solution is a possibility but I think it needs to be done at fairly low temperatures because high hydrates don't
like high temperatures, as evidenced also by the low melting point of several alums: for these the crystalline hydrate cannot exist above that
temperature. |
Good point.
Out of madscience interest, can one prepare alums from things like choline and a metal? The reason that I ask is because I happen to have quiet a bit
of choline left over from my choline soap experiments?
[Edited on 27-10-2013 by deltaH]
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woelen
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The Cs/Al-alum exists, I have appr. 150 grams of this (purchased from eBay).
If you have a soluble Cs-salt, e.g. CsCl or CsNO3, then add a few drops of a concentrated solution of that to a solution of K/Al-alum and then you
soon see many glittering crystals separate from the liquid. If you add a few drops of a soluble cesium salt to a solution of K/Cr-alum, then you soon
get many glittering purple crystals of Cs/Cr-alum. So, I can imagine that with Cs-salts you can get your Cs/Ti alum precipitating.
Cs-alums tend to be much less soluble than their K-counterparts.
An interesting experiment could be to add a little amount of a Cs-salt to your Ti(III) solutions. I would not want to use the rare CsOH for this
experiment. If you have CsCl or CsNO3, then use that in combination with a concentrated solution of Ti2(SO4)3, without any added K2SO4 or (NH4)2SO4.
[Edited on 27-10-13 by woelen]
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blogfast25
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| Quote: Originally posted by woelen | [
Cs-alums tend to be much less soluble than their K-counterparts.
I would not want to use the rare CsOH for this experiment, if you have CsCl or CsNO3, then use that in combination with a concentrated solution of
Ti2(SO4)3, without any added K2SO4 or (NH4)2SO4.
[Edited on 27-10-13 by woelen] |
Yes, I seem to remember the Cs alums are much less soluble.
I don't have any Cs salts, other than CsOH. Will take a peak in eBay.
An NH4Ti(III) alum would have been nice though: reasonably cheap, potentially more stable than most solid Ti(III) compounds and potentially useful. So
I might still go for low temp. vacuum evaporation, just to see...
@deltaH: alums from choline? Interesting thought but not the foggiest clue.
eBay uk:
Cesium nitrate: £80.99 for 25 g
Cesium chloride £7.99 for 5 g
And this one's strange:
http://www.ebay.co.uk/itm/Liquid-Cesium-Chloride-Plus-Rubidi...
£99 / L, claims to contain "96,000 mg CsCl / L", plus a gram of RbCl thrown in for good measure. Medicinal, allegedly!
[Edited on 27-10-2013 by blogfast25]
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DraconicAcid
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| Quote: Originally posted by deltaH |
Out of madscience interest, can one prepare alums from things like choline and a metal? The reason that I ask is because I happen to have quiet a bit
of choline left over from my choline soap experiments?
[Edited on 27-10-2013 by deltaH] |
I don't know. I do know that you can prepare a room-temperature molten salt with choline chloride and urea, which is a nifty solvent for metal oxides
(haven't tried it yet, but I've got the stuff).
Cite: http://pubs.rsc.org/en/content/articlelanding/2003/cc/b21071...
Please remember: "Filtrate" is not a verb.
Write up your lab reports the way your instructor wants them, not the way your ex-instructor wants them.
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