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blogfast25
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Titanium (III) Potassium Alum: failed attempt
Alums need little introduction here, I think. Essentially double salts with the generic formula NM(SO4)2.12H2O with N = NH4, Na, K, Rb or Cs and M =
trivalent Al, Fe, Tl, Cr and some other exotic metals.
I’ve often wondered if M = Ti<sup>3+</sup> would be possible. Searching, admittedly not very deep, only yielded one Google book
reference (<a href="http://books.google.co.uk/books?id=PpTi_JAx7PgC&pg=PA684&lpg=PA684&dq=titanium+(III)+sulphate+alum"
target="_blank">link</a> but it’s quite a general book and perhaps not
very authorative on the subject.
So I decided to have a go myself.
About 2.0 g (0.025 mole) of TiO2 was dissolved in an excess of 96 % H2SO4 and the solution (of TiOSO4) diluted slightly to about 40 ml (so about 0.6 M
TiOSO4). A small amount of undissolved TiO2 was eliminated by pipetting off the supernatant liquid. To this solution was then added 2.2 g of K2SO4
which was dissolving by vigorous stirring. Slight cloudiness developed on heating (to help dissolve the K2SO4), pointing to hydrolysis of the TiOSO4
and was immediately stopped. Then about 1 g of ultra-fine, very pure zinc powder was added to effectuate the reduction of the
TiO<sup>2+</sup> ion:
TiO<sup>2+</sup>(aq) + ½ Zn(s) + 4 H<sub>3</sub>O<sup>+</sup>(aq) → Ti<sup>3+</sup>(aq) + ½
Zn<sup>2+</sup>(aq) + 6 H<sub>2</sub>O(l)
The solution turned a very deep purple, characteristic of Ti<sup>3+</sup> cations, almost immediately and hydrogen evolved.
At that point theoretically the solution contained an estimated over 30 g of KTi(SO4)2.12H2O per 100 g of solution and was carefully double
clingfilmed to at least provide a first barrier to air oxygen. But on cooling and overnight ice bath no crystals materialised.
It is possible that the solubility limit of the alum has not been exceeded. It may be difficult however to make solutions of TiOSO4 that are more
concentrated without running into hydrolysis problems.
Maybe I should make another attempt with ammonium sulphate? Or maybe this alum just cannot exist…
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[Edited on 13.10.13 by bfesser]
[Edited on 13-10-2013 by blogfast25]
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deltaH
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Very interesting work blogfast, well done. I would guess that your solution of TiO2SO4 is very acidic, so perhaps only part of the TiO2+ ions reduced
and the remainder was simply water reduction (zinc in acid). One possiblity might be to employ more than stoichiometric amounts of zinc as some
reduction did take place.
But I think there is a better way, what about other reducing agents, specifically one that also contributes Ti3+?
Yes... I am saying use titanium metal as the reducing agent, conditions certainly seem acidic enough and you would need to employ the additional
amount of sulfuric acid off course, half reactions from wiki:
TiO(2+) + 2 H+ + e− <=> Ti3+ + H2O Estd. = +0.19V
Ti(s) <=> Ti3+ + 3 e− Estd. = +1.37
So: 3TiO(2+) + 6H+ + Ti(s) <=> 4Ti3+ + 3H2O E = 1.56V
Maybe you would need to do this before diluting your TiOSO4/H2SO4 solution and using a cleaned (scoured) strip of titanium metal to remove surface
oxide layer and speed up the onset of dissolution. You could also weigh the strip before and after to make sure you reacted the necessary amount
before proceeding.
PLEASE be careful of runaway as I am sure the reduction will produce a lot of heat and might drastically accelerate in concentrated H2SO4!
Only once you have a deep purple solution of Ti3+ in conc. H2SO4 do you proceed with K2SO4 addition and dilution.
[Edited on 13-10-2013 by deltaH]
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blogfast25
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deltaH:
I'm fairly sure all Ti (IV) was reduced to Ti (III), I used to do this all the time with Al. Of course here I couldn't use Al.
I like the idea of reducing Ti(IV) with Ti(0). I think it will be sluggish though. Also, Ti(0) does dissolve noticeably in strong HCl or H2SO4. I used
to do that too, but it's slow even at reflux. I have some fairly dilute TiOCl2 solutions (about 0.1 M) from other exploits so I might test the idea
that way.
Adding the other sulphate (be it NH4 or K) after reduction is asking for introducing some oxygen into the system, Ti(III)'s mortal enemy. That's why I
added it before the reduction: hydrogen would flush out the oxygen.
Another possibility is to try and get as concentrated a solution of Ti2(SO4)3 by direct dissolution of Ti metal in dilute H2SO4. In hot 37 % HCl Ti
dissolves fairly well but it's not what I would call fast. Solution could be further concentrated under hydrogen flux.
[Edited on 13-10-2013 by blogfast25]
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deltaH
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Yeah, kinetics could be a problem but as long as it goes, it will all go with patience. However I don't think it would work with dilute solutions.
What I don't understand is that you dissolve TiO2 in conc. acid with no kinetic issues, surely then Ti would react just fine as it's the TiO2 layer
that passivates it?
[Edited on 13-10-2013 by deltaH]
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deltaH
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I really think water can be a big problem in trying to form Ti3+ as Ti really wants to go to TiO2, ie. Ti(IV), the only way to prevent this is to make
sure there is enough excess H2SO4 to sink the water. For example, the reaction:
(i) 3TiO(2+) + 6H+ + Ti(s) <=> 4Ti3+ + 3H2O
produces 3 moles of water, if this not sunk by a second reaction then a parallel reaction that can occur is:
(ii) Ti(s) + H2O + 2H+ => TiO(2+) + H2(g)
Note reaction (i) produces water and reaction (ii) requires it, so an effective strategy here is to sink the water by ionisation to H3O+
So you need a further 3 moles of H2SO4 to fully ionise the water produced by reaction (i), the balance equation now becomes:
3TiO(HSO4)2 + Ti(s) + 9H2SO4 => 4Ti(HSO4)3 + 3HSO4(-) + 3H3O+
This way you really force the issue and force the oxidising titanium to be oxidised by TiO(2+) and not H2O!
NOTE: I used HSO4(-) as the speciation for sulfate, not SO4(2-) as it's only the pKa1 of H2SO4 that is really extreme enough to guarantee things go to
completion.
Furthermore, if you are using 95% H2SO4, then you also need to account for the 5% water there as this would reduce the amount of H2SO4 species
available is probably far less than 95%, i.e. part of it is HSO4- and H3O+.
So if you start with TiO2(s), you need a wopping 11 H2SO4 + whatever excess to correct for the water already in 95% H2SO4, and then to once dissolved,
to this you add your titanium strip.
Finally, I think the issue of sluggish kinetics of titanium metal reaction will be overcome iff TiO(2+) is in high concentration making the solution
oxidising AND H2O is trace (ok if it's there as H3O+) else TiO2 can form which hampers the kinetics. Besides with such concentrated H2SO4, making the
solution really hot is possible should the need arise that should take care of kinetic issues, though this would be also be significantly more
hazardous!
Sorry for the long discussion, just found the topic interesting 
[Edited on 14-10-2013 by deltaH]
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unionised
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Just a thought.
If you crystallise ordinary (potassium +aluminium) alum from a solution containing Cr(III) you get mixed crystals (nice amethyst colours).
Could this approach get a crystal that at least contains some Ti(III)?
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blogfast25
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deltaH:
Thanks.
My next attempt will be based on H2SO4 + powdered Ti metal. I'm fairly convinced the problem here was too low concentration of the alum. See my post
on ferric alum, a few posts down: there the hot concentration of the alum was about 2.5 M!
Unionised:
What do you mean by mixed crystals? Al and Cr randomly distributed throughout the lattice? Or a mixture of Cr alum crystals and Al alum crystals?
This is certainly worth a try with the solution I have: I have plenty Al alum and it's easy to recrystallize (2 parts alum to 1 part water). I could
try and substitute the water (or part of it) with the solution that didn't crystallise any Ti alum. Definitely worth a shot!
[Edited on 14-10-2013 by blogfast25]
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deltaH
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All the best blogfast25, can't wait for your results!
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blogfast25
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Today I tried to dissolve some 'pyro grade' titanium powder (about 60 mesh, what the pyros like to call 'titanium sponge') in 50 w% H2SO4.
2.4 g of Ti powder (0.05 mole), 15 g of water and 15 g of 96 % H2SO4 (twice the stoichiometric amount) were combined in a 100 ml RBF. Reaction started
surprisingly quickly and was very vigorous, much more than I expected. The solution became purplish/blue very quickly.
Unfortunately it became clear that also some TiO2 was being formed. Presumably H2SO4 is too oxidising at 50 %.
After about 10 minutes near BP, I added some 5 g of K2SO4, simmered for some more and then allowed to cool. Quite quickly the mass solidified into a
two phase system, with a clear, colourless top layer (K2SO4? KHSO4?) and a purplish, opaque mushy mass:
Some of the bottom layer was dissolved in 37 % HCl and it was clear that there was undissolved white matter, presumably TiO2, in that layer.
[Edited on 17-10-2013 by blogfast25]
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deltaH
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Sounds very promising indeed and loved the colours! A small cooler box is a handy place to put a flask if you want it to cool down very slowly (over a
day)... would help the crystallisation!
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blogfast25
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Quote: Originally posted by deltaH  | | Sounds very promising indeed and loved the colours! A small cooler box is a handy place to put a flask if you want it to cool down very slowly (over a
day)... would help the crystallisation! |
Hmmm... personally I don't think it's very promising at all. I don't think the purple layer is alum but I don't know what it is. The added amount of
of K2SO4 was sub-stoichiometric, yet much of it seems to crystallise out on it's own. The excess H2SO4 and formed TiO2 are complications I didn't
need.
Tomorrow I'll make a comparison with 37 % HCl dissolution.
[Edited on 18-10-2013 by blogfast25]
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blogfast25
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2.4 g (0.05 mole) of the same Ti powder and 25 ml of 37 w% HCl (about twice the stoichiometric) was loaded into a 100 ml RBF and refluxed with gentle
heating on a hot plate. A primitive gazometer served as HCl scrubber and indicator of hydrogen evolution (right large test tube):
Reaction was very swift, taking approx. 30 min for the gas evolution to more or less die down. For most of the time 5 or 6 large bubbles of hydrogen
could seen floating to the top at the same time.
Somewhat before hydrogen evolution had ceased, 0.025 mole of K2SO4 was added and the solution simmered for a few more minutes. When gas evolution had
all but ceased heating was stopped and the solution would have been approx. 2 M in Ti<sup>3+</sup> but not enough sulphate was present for
KTi(SO4)2.12H2O stoichiometry. So after a bit more cooling about 5 ml of 96 w% H2SO4 was added very slowly to make up for the deficiency in sulphate
molarity.
Unfortunately this caused the solution which was hitherto clear (but very dark) to lighten in colour and to cloud over strongly. I suspect again
oxidation of the Ti<sup>3+</sup> by the sulphuric acid, to a hydrolysed Ti(IV) species. After cooling this was obtained:
I’ll know for sure tomorrow if the precipitate is TiO2 but it sure looks like it. Hopes of salvaging anything from this attempt are slim.
Another attempt tomorrow.
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deltaH
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| Quote: | | I suspect again oxidation of the Ti3+ by the sulphuric acid, to a hydrolysed Ti(IV) species. |
If the addition
of the acid did indeed oxidise the Ti3+ species, then it must have also produced H2 upon this addition, which it seems it didn't.
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blogfast25
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| Quote: Originally posted by deltaH | | If the addition of the acid did indeed oxidise the Ti3+ species, then it must have also produced H2 upon this addition, which it
seems it didn't. |
It's not necessarily easy to see, if it doesn't evolve all at once. If the whitish stuff shows to be insoluble later on, I can't see what else could
have happened.
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deltaH
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Look, I think you need to take stock of what you observed. In all intents and purposes, gas was probably not evolved, that would mean that something
precipitated upon adding the acid. Since it precipitated because of the acid addition, it's logical to conclude that it's a sulfate or bisulfate rich
salt.
Things I can think of are: an alum or Ti(HSO4)3 if this even exists? It's possible that the strongly dehydrating conditions of adding the extra H2SO4
conc. strips the Ti3+ of any aqua ligands and so forces the precipitation of the anhydrous salt Ti(HSO4)3? The fact that aqua ligands were stripped
might even explain why it's not white and temporarily insoluble.
If you dissolve your white product in dil. H2SO4 and get a blue solution again, then I would say this is some evidence to support such a theory.
[Edited on 18-10-2013 by deltaH]
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blogfast25
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It's possible, deltaH. It would be hope giving if the precipitate was indeed a Ti(III) compound. Incidentally, even anhydrous Ti(III) compounds, or at
least some, like TiF3 are coloured. Whatever is causing the problem, it might be fixable by adding the acid as 50 % or less.
All should be revealed tomorrow.
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deltaH
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| Quote: | | Incidentally, even anhydrous Ti(III) compounds, or at least some, like TiF3 are coloured. |
Interesting, I did
not know that, I must admit that Ti in any state other than IV is completely foreign to me, which is why I find this so fascinating
But your thread has made me realise that I prepared a Ti(III) complex accidentally once. I was condensing a quinone with
bis(trimethylsilyl)carbodiimide to form the quinone-dicyanoimide and this was done by first reacting the quinone with TiCl4, then adding the imide, so
that the titanium would pull the oxygen from the quinone to form some solid white Ti(OCl)x mess that precipitates. Anyhow, somehow when washing
glassware, I noticed these blue ink-like stains all over the sink from those white solids the next day.
Hang on... maybe what you have is that same titanium oxychloride mess.
[Edited on 18-10-2013 by deltaH]
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deltaH
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Actually a titanium oxochloride animal makes perfect sense, you will find that upon adding water/dilute acid, that it does indeed dissolve and will
give you a blue solution again as the chloride hydrolyses.
Look, strictly speaking, this white precipitate was supposed to be TiO2Cl2, but I definitely had lots and lots of blueing forming in the sink after
some time, so something was reducing it.
Maybe amorphous TiO2 formed by the hydrolysis of TiO2Cl2 was getting reduced by conc. HCl forming locally as the TiO2Cl2 was
hydrolysing from the moisture?
What I'm very surprised about is how Ti(III) survived exposure to air for so long!!!
This titanium chemistry is perplexing!
[Edited on 18-10-2013 by deltaH]
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blogfast25
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Well, well, looks like deltaH could be right.
This material does not appear to be TiO2 (IV).
It was filtered off on Buchner and washed profusely with cold water and didn’t lose any colour. It’s a sandy, non-gelatinous, light blue powder.
Here it is just after sucking dry, the filtrate of Ti(III) can also be seen:
Here’s another photo of it, after having recovered it from the filter:
Some tests were carried out on the material. In four test tubes about a quarter of a teaspoon of material was loaded. Then various additions were
made:
1) Strong ammonia solution: precipitates black Ti(OH)3. After a while gas started to evolve.
2) Strong NaOH solution: precipitates black Ti(OH)3
3) Some drops of 37 % HCl, some water and a few drops of 35 % H2O: oxidation of Ti(III) to Ti(IV) and formation of red Ti(IV) peroxo complex
4) Pure 37 % HCl: no change, no dissolution
All this strongly points to the material being a Ti(III) compound, insoluble or very poorly soluble.
A pinch of it was put on Al foil on a max setting hot plate. The material darkened somewhat but no substantial change was observed.
Then about 2 g of the material was treated with 10 ml of 33 % NH3 and the Ti(OH)3 filtered off. To a blank (virgin NH3 solution) and a few ml of the
clear and colourless filtrate, a ml of Ba(NO3)2 solution (2 g in about 10 ml of water) was added:
Left the blank, right the filtrate: the latter tests very strongly positive for sulphate ions. This suggests there’s plenty bound sulphate in the
material. Remember that the precipitate had been washed carefully with DIW.
I now assume that the precipitate obtained in the post above was in all likelihood identical to this material. That would suggest it will not contain
any potassium.
Some further tests will now be carried out to test for chloride and potassium. If a sulphate based insoluble Ti(III) compound exists, the preparation
of a Ti(III) alum may not be possible or very difficult to do.
[Edited on 19-10-2013 by blogfast25]
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deltaH
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Fascinating and extremely interesting chemistry.
| Quote: | | 4) Pure 37 % HCl: no change, no dissolution |
This is significant, because clearly you have some kind of
Ti(III) sulfate or bisulfate, BUT if it was simply Ti(HSO4)3 or Ti2(SO4)3, then one would expect it to dissolve in HCl, but it remained resolutely
insoluble suggesting that this is some polymeric and probably highly amorphous partial oxide, for example (TiO)(HSO4), the Ti(III) version of
(TiO)(HSO4)2
This might be the sulfate version of the white precipitates I made with my chloride derivatives back in the day...
I think it might help to consider the chemistry of aluminium oxychlorides as I think they are analogous.
Good luck with your investigations!
[Edited on 19-10-2013 by deltaH]
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blogfast25
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There's more. I'm now pretty convinced the material also contains bound chloride.
2.0 g of the material was again treated with 10 ml of 33 % NH3 and the Ti(OH)3 filtered off. To the filtrate was added 2 g of Ba(NO3)2 (99.4 %)
dissolved in about 20 ml of pure water. The sulphate crashed out as BaSO4, as before. This slurry was then filtered to clarity and acidified to pH 5
with glacial acetic acid. On adding a few g of AgNO3 (p.a.) dissolved in pure water, lots of AgCl precipitated. The amount is difficult to explain
away other than as chloride chemically bound in the material.
So are we perhaps talking about a Ti(III) double (bi)sulphate-chloride?
Tomorrow, time allowing, I'll see if I can test the bisulphate idea.
To address the stability issue of Ti(III), I say 'how long is a piece of string?' Of course Ti3+ is a powerful reducing agent and very easy to
oxidise. But I've kept TiCl3 solutions for years, away from air and found no significant deterioration. Obviously conditions of storage will affect
lifetime of a titanous solution.
[Edited on 19-10-2013 by blogfast25]
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deltaH
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Hmmm... I think you may have used too little barium nitrate to precipitate all the possible sulfate and so you might just have been forming silver
sulfate precipitate. Say for example you unknown precipitate has a formula Ti(HSO4)3, then MW = 339. That would have produced 17.7mmol sulfate
necessitating 17.7mmol Ba(NO3)2, thus 4.6g!
I suggest you repeat the test with 6g Ba(NO3) to be on the safe side.
[Edited on 19-10-2013 by deltaH]
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bismuthate
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Take the percipitate from the AgNO3 with a BaNO3 solution , then filter again and finally add NaCl if the percipitate was AgSO4 then a white
percipitate should form. Here is the series of reactions that I believe will happen.
AgSO4(aq)+BaNO3(aq)==>BaSO4(s)+AgNO3(aq)
AgNO3(aq)+NaCl(aq)==>NaSO4(aq)+AgCl(s)
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deltaH
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If your chloride test is still positive, then you have a choice between two things: it's either a mixture of two salts (I would think unlikely), or
it's an amorphous -O-Ti-O- (III) polymeric/framework material where the the residual positive charge is neutralised by mixed bisulfate and chloride
ions, as if it were an ion exchanged material.
If it is this, you should be able to, by repeated washing at the filter with concentrated acids, exchange it more fully with the one or the other. The
you should get nearly complete, only sulfate positive test or only chloride positive test, depending on the acid you used to do the repeated washings.
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deltaH
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Nice idea bismuthate, I like your way of thinking!
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