bereal511 - 14-12-2006 at 22:22
I've been having trouble on finding information on titanium sulfate and how to manufacture it. Most of the processes for making titanium sulfate have
been from the ore FeTiO3, ilemenite(sp), using sulfuric acid, but what's the minimum concentration of sulfuric acid that's required to form titanium
sulfate directly from titanium dioxide (if it's any different from the ore process)? And does titanium sulfate react with filter paper?
woelen - 14-12-2006 at 23:46
I have some TiO2 and up to now I have not managed to find any solvent, capable of dissolving this. I tried with hot concentrated H2SO4, hot
concentrated solutions of NaOH and a whole bunch of other acids and mixes (HNO3, HCl, HClO4). No success. Probably you need molten alkalies, but I did
not try that.
YT2095 - 15-12-2006 at 01:43
or a thermit reaction perhaps?
Eclectic - 15-12-2006 at 06:55
I've made Ti(III) sulfate from scrap Ti turnings and H2SO4 drain opener, diluted to 50% or so. I think I needed a little 30% hydrogen peroxide to get
through the surface oxide. Bright orange solution at first, then turns to inky blue as the reaction progresses to reducing conditions. Add 100% H2SO4
as the reaction slows until all the turnings are dissolved. It took about a week.
You aren't going to be filtering this stuff with paper filters. You'll need a glass frit buchner funnel, or glass fiber, plastic filter material.
You might be able to get a fairly stable Ti(III) alum by adding saturated ammonium sulfate solution.
[Edited on 15-12-2006 by Eclectic]
bereal511 - 16-12-2006 at 17:49
I wonder though, does hydrofluoric acid attack titanium dioxide? I'm not at all proposing for myself or anyone else to find that out, but I'm just
curious.
chromium - 18-12-2006 at 16:09
Titanium dioxide is soluble in molten KHSO4 (this is used for etching TiO2 surface) but i am not sure if this can be used to prepare titanium salts.
Alkali fusion can be used to make titanates from TiO2 but, again, i am not sure how these can be used to make titanium salts.
I have book which tells that TiO2 is soluble in aqueous HF but no data how fast this process is and what exactly happens.
[Edited on 19-12-2006 by chromium]
The_Davster - 18-12-2006 at 16:50
Arg lost this post when I opened a djvu file....
TiO2 dissolves in warm 20-40% HF (Brauer)
Handbook of inorganic chemicals states that it dissolves in HNO3 and aqua regia.
Fusion with acid sulfate followed by soaking the crucible in sulfuric acid gives a solution of TiOSO4.
Fusion with alkali gives titanates which form titanium compounds on extraction of the melt with acid.
Chemistry of the Rarer elements says it dissolves in hot excess H2SO4 being reduced to Ti2(SO4)3. Apparently it forms alums with alkali sulfates.
Could always reduce to crude metal with Ca or Na (Mg forms Mg-titanate) then use the metal for dissolution. Or for purer metal use CaH2
I have some TiO2 too. I should play.
not_important - 18-12-2006 at 19:08
The solubility may be affected by the 'history' of the TiO2, as with a number of other oxides. Heating to several hundred degrees C often makes an
oxide less reactive than the form produced by mild heating of a hydroxide, carbonate, or nitrate.
I believe that ilmenite is easier to attack with H2SO4 than TiO2 is. I also think that you can use TiO2 and concentrated H2SO4 to get TiOSO4, use a
slight excess of acid and heat slowly to 150 C.
Solutions of TiSO4 in water must be made in the cold and kept cool to avoid hydrolysis; addition of a small amount of free H2SO4 helps as well.
And, yes, you're going to want to use fritted glass or some other non-organic filter. It is possible that by doing the H2SO4 + TiO2 reaction, then
adding excess TiO2 and continue heating, you may able to get a solution low enough in acid that filter paper can be used.
Given tah TiO2 and H2SO4 are fairly low cost, some test tube scale experimentation might be worthwhile. A clear solution that, upon dilutation and
warming to near boil, gave a white ppt is likely to be TiOSO4.
Endo - 19-12-2006 at 05:50
TiO2 can also be dissolved (to Ti3+) in a mixture of ammonium sulphate and sulfuric acid. This prep is used for preparation of analytical samples of
suncreens and makeup containing TiO2. (can be done in a quartz crucible or even a beaker)
We also routinely digested TiO2 in a mixture of 50% HF and 37% HCl 1:3 by volume. This digestion would clarify after around 10 min at a simmer
leaving a clear solution. (platinum crucible) We would then warm this solution to dryness (didn't want HF in the instrument) and redissolve it in HCl
(~10% by volume). If in drying the solution the solution was heated beyond dryness it often took additional heating with HCl or a second addition of
HF (followed by drying again) before it would become soluble.
I know that the Ti 3+ was stable in a solution of ~2% Sulfuric with around 0.1% ammonium sulfate. Literature information exists indicating that Ti
3+ remains stable for a long time in 10% by volume HCl. Addition of hydrogen peroxide to the HCl solution forms a complex that varies in color from
light yellow to a dark red orange depending on the Ti 3+ concentration. (has been used for colorimetric assays)
In any case pushing the pH basic while having a clear solution of Ti 3+ and SO4 may be something to try.
unionised - 19-12-2006 at 06:12
Are you sure it's Ti3+ not Ti4+?
Eclectic - 19-12-2006 at 06:28
Easy to tell. Ti(III) solutions are dark inky blue (just a trace in alumina crystal is the blue of sapphire). AFAIK, Ti(IV) solutions are
colorless, and the peroxo complex is a really strong orange.
Oxalic acid is used to strip surface oxide in some Ti electroplating operations.
[Edited on 19-12-2006 by Eclectic]
woelen - 19-12-2006 at 07:04
Here follow some pictures and some properties of titanium ions. Indeed, Ti(3+) remains stable in dilute HCl for quite a long time:
http://woelen.scheikunde.net/science/chem/solutions/ti.html
The color of titanium (III) also can be quite different. In concentrated HCl it is much more blueish. With fluoride, the metal dissolves MUCH faster
in acid, and then a green complex is formed, which on contact with air (or any oxidizer) first becomes brown, and finally turns colorless:
http://woelen.scheikunde.net/science/chem/riddles/titanium+f...
[Edited on 19-12-06 by woelen]
Endo - 19-12-2006 at 08:22
Yes there was a tint of blue, as most of our solutions were ~50ppm it was not that noticable. I am sure it would be much darker as concentration
increases.
Dr. Beaker - 19-12-2006 at 13:12
you can allways imitate the industrial process and convert it to TiCl4 by heating with carbon under a flow of chlorine.
from there you need to distil the raw TiCl4 to remove Al and Fe chlorides and then the door is open to almost any Ti(IV) derivative imaginable.
bereal511 - 19-12-2006 at 16:22
Eh, that's not something I'm especially equipped for. It sounds like more trouble than its worth, especially with a chlorine flow generator under
high temperatures.
bereal511 - 6-2-2007 at 19:24
I have some following up on this subject of dissolving titanium dioxide.
A few days ago, I took about 20 grams of titanium dioxide and boiled it in 300 ml of 18 Molar sulfuric acid (in excess). I kept the reaction at 230
degrees C, which is when the mixture began to turn to a sulfur-like yellow as the sulfuric acid began to throw off heavy white fumes. I assumed this
was Ti2(SO3)3. So I left it to cool overnight. But when I came back the next day, the solution became a pale murky yellow and a white precipitate
was left on the bottom of the beaker. I haven't tested anything yet, so I'm not too sure if there's still any titanium salt of whatever sulfate
species in solution, but I'd like to know if it's possible that the titanium salt hydrolyzed somehow and reverted to a titanium oxide of some sort.
[Edited on 7-2-2007 by bereal511]
Fleaker - 7-2-2007 at 13:21
At 230 C it might have been the sulfuric acid giving off SO3 gas (which immediately combines with moisture in the air, giving you your white fumes).
bereal511 - 7-2-2007 at 19:16
Well, that's not exactly what I meant. I must have editted the sentence wrong twice, because I meant to say that the yellowing precipitate in the
beaker might have the titanium sulfate, not the white fumes. I figured the fumes were decomposition products of sulfuric acid before hand.
JohnWW - 10-2-2007 at 14:15
If there is any Ti(IV) sulfate, it could only be a basic sulfate, like TiOSO4 (which may be formable from direct reaction of TiO2 with SO3 gas) or
Ti(OH)2SO4. The Ti++++ cation would have too high a charge density for it to exist in aqueous solution, or to coexist with any complex oxy-anion. The
hydrated Ti+++ cation is an intense purple color, as I understand.