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Author: Subject: Converting TiO2 to TiCl4 ?
AndersHoveland
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[*] posted on 16-5-2012 at 14:46
Converting TiO2 to TiCl4 ?


Is there any way to directly convert titanium dioxide to titanium tetrachloride, without reducing the TiO2 to titanium metal?

I am trying to develop a lower temperature method to make titanium metal from the ore.

I am thinking it might be possible to react CCl4 with TiO2, with some heating. TiCl4 is a liquid at room temperature, so the reaction might be able to proceed under relatively mild conditions. CBr4, being much more reactive, might react with TiO2 at room temperature, although CBr4 has a higher melting point at 91 °C.

TiO2(s) + CCl4(l) --> TiCl4(l) + CO2(g)

Carbon tetrachloride, for example, reacts with aluminum tribromide at 100 °C. (N. N. Greenwood, A. Earnshaw. Chemie prvku (Chemistry of the Elements). Informatorium, Prague, 1993.)

Heats of formation (kJ/mol)
TiO2 -944.7
TiCl4 -804.2, TiBr4 -619.7
CCl4 -139.5, CBr4 -150.0*
CO2 -393.5

There seems to be several different conflicting values for the heat of formation of CBr4. This one seems to be the most plausible.

Lower melting point titanium alloys:
Quote:

Ti15Cu15Ni and the newly developed Ti21Ni14Cu.
Zr-Ti-rich side of the Zr-Ti-Ni(Cu) alloy system were investigated for brazing of titanium alloys. Low-melting ternary and quaternary eutectic alloys with melting temperatures below 800°C were discovered. Using eutectic as well as off-eutectic braze alloys, CP-Ti and Ti-6Al-4V alloys were successfully brazed at 830°C and 850°C.


[Edited on 16-5-2012 by AndersHoveland]
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[*] posted on 16-5-2012 at 15:12


I have been waiting for a way to make TiCl4 at home, and if your idea works I would be a happy person. I would test it myself but i don't have any or any easy wat to make any right now.
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[*] posted on 16-5-2012 at 15:44


The Wikipedia page on TiCl4 references a process of producing the chloride directly from the oxide. Rather than continuing to speculate, it would more useful for everybody to report on what's already being done. The Wikipedia page isn't referenced. It would be a public service to provide one. I'd start with Kirk-Othmer.
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AndersHoveland
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[*] posted on 16-5-2012 at 15:58


Apparently my idea is not new. :(
http://en.wikipedia.org/wiki/Hunter_process
http://en.wikipedia.org/wiki/Chloride_process

Here is another idea: heat calcium chloride together in an electric furnace with the ore, distill off the vaporized TiCl4 gas. It would be an endothermic reaction. I would think the temperature would have to be very very high, however.

3 TiO2 + 3 CaCl2 --> 2 CaTiO3 + TiCl4(g)

or perhaps

3 TiO2 + SO2Cl2 --> 2 TiOSO4 + TiCl4

(sulfuryl chloride) (titanium IV oxysulfate "hydrate" is made by reacting conc. H2SO4 with TiO2. There does not seem to be any information about the anhydrous compound)

[Edited on 17-5-2012 by AndersHoveland]
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[*] posted on 16-5-2012 at 17:19


Simply melt pyrosulphate with the ore in presence of NaCl, it will make HCl + TiCl4

Sources: my own work

In fact I already made most inorganic titanium compounds and many organic ones, I think that I have some(a lot) knowledge in this field.

[Edited on 17-5-2012 by plante1999]

[Edited on 17-5-2012 by plante1999]




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[*] posted on 16-5-2012 at 18:03


I believe Sauron posted some stuff about reacting benzotrichloride with TiO2 to yield both TiCl4 and benzoyl chloride.



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[*] posted on 16-5-2012 at 19:25


May I suggest (speculation) exposing TiO2 to Silicon tetrachloride. On cooling, SiO2 precipitates out.

The reaction:

TiO2 + SiCl4 ---> TiCl4 + SiO2 (s)

--------------------------------------------

Yet another speculation:

TiO2 + 4 BrCl ---> TiCl4 + 2 Br2O

and, as we are interested in a lower temperature synthesis, I would try the photolysis of BrCl, which is known to form atomic Cl (see, for example, http://imk-aida.fzk.de/CMD/final_report/summaries/Summ-HEP2.... ). The reaction may, however, proceed slowly (assuming it proceeds at all).

Also, I would expect:

2 Br2O --> 2 Br2 + O2

so make allowances for a possible rapid gas expansion with an energy release. EDIT: I found a reference which states that Dibromine oxide is indeed unstable decomposing when warmed above -40 C and vigorously at its melting point of -17.5 C into Bromine and Oxygen (Reference: "Inorganic Chemistry" by Egon Wiberg, A. F. Holleman, Nils Wiberg, page 464, link: http://books.google.com/books?id=Mtth5g59dEIC&pg=PA464&a... ). Note, this decomposition could further move the reaction to the right.

As you may be aware, BrCl is especially irritating to the eyes even in very small dose exposures and, of course, toxic.



[Edited on 17-5-2012 by AJKOER]
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[*] posted on 16-5-2012 at 21:18


Plante1999, please tell me more on how you made TiCl4 from TiO2, pyrosulfate and NaCl. Quantities, apparatus,conditions?
I'd be highly interested in this method.

According to Vanino's book, a mixture of 1000g TiO2 and 350-400g charcoal can be melted down in an electric arc furnace (35 volts, 200-250 amps), forming lumps of titanium carbide. This carbide reacts more readily and at lower temperature with chlorine than the usual mixture of TiO2 and carbon.
If natural rutile ore is used instead of pure TiO2, laborious purification from SiCl4 and VOCl3 is necessary.

I once made a few tens of grams of TiCl4 in my tube furnace by reaction of Ti metal with chlorine. The Ti was used in the form of hard porous lumps, called "titanium sponge" and bought from a chemical supplier. I think you could use cutoffs of Grade 1-4 titanium sheet just as well.
I first flushed the quartz reaction tube containing the Ti with dry chlorine gas at room temperature and then heated it up while continuing to slowly feed chlorine through it.
At a certain temperature (sorry, don't remember the number) the Ti metal suddenly started reacting, producing intense heat and bright yellow incandescence. I increased the chlorine stream and turned off the furnace since the reaction produced plenty of heat itself, the metal was basically burning in chlorine.
The crude TiCl4 collected as a turbid liquid in the cooled receiver and was distilled to give the pure clear and colorless substance.
What I remember most vividly about this preparation is the
insane amount of dense white smoke that was released upon disassembly of the still warm TiCl4-wetted synthesis apparatus. The smoke was even thicker and more opaque than that from handling pure SO3 in air.
Warm TiCl4 is the most potent smokescreen chemical that I've ever seen. When it's at room temperature the smoke is not that strong, i.e. weaker that SO3, but heat it up and it's very impressive.






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[*] posted on 17-5-2012 at 10:19


You’re a braver man than I am, Garage Chemist, to pull that off in homelab conditions.

The starting idea of trying to make titanium or titanium compounds like TiCl4 at lower temperatures appears to me to be a contradictio in terminis: look at these HoF’s, ferchr*ssake! Anything that would start from TiO2 would need fairly insane temperatures to get going. You could of course hope for a catalyst: good luck with that!

But I'd like to hear more about Plante's pyrosulphate method...


[Edited on 17-5-2012 by blogfast25]




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[*] posted on 21-5-2012 at 19:46


Quote: Originally posted by garage chemist  

What I remember most vividly about this preparation is the
insane amount of dense white smoke that was released upon disassembly of the still warm TiCl4-wetted synthesis apparatus.



That's why, in the glovebox world, the QAD test for ppm levels of H2O is to open a bottle slightly inside the box. Very handy when you have reason to doubt the instruments.




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[*] posted on 26-9-2013 at 18:45


I am also eagerly awaiting details of Planet's preparation.
edit-sorry about your name. I am typing on my phone and it is improperly auto correcting.

[Edited on 27-9-2013 by Chemosynthesis]
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AndersHoveland
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[*] posted on 26-9-2013 at 23:04


Quote: Originally posted by plante1999  
Simply melt pyrosulphate with the ore in presence of NaCl, it will make HCl + TiCl4

That is a good idea.
I think the reaction would still be endothermic though, and would take a large amount of heat to get the ore to dissolve in the flux. Might be better to use potassium salts, which have a lower melting point.

4 Na2S2O7 + 4 NaCl + TiO2 --> 2 Na2SO4 + TiCl4 (vaporized)


Another idea might be to react TiF4 with CaCl2 in a flux of molten ZnCl2

Probably pyrosulfate with NaF and TiO2 would very easily make TiF4. very dangerous fumes
titanium has a strong affinity towards fluorine, even stronger than towards oxygen

Titanium dissolves in hydrofluoric acid to form TiF3. once the element titanium is separated from the oxygen, electrolysis becomes easy, because halogenated titanium can form low melting point fluxes.

Also, an obvious thing to mention, titanium tetrachloride spontaneously reacts with water, whereas the trichloride is soluble (forming violet solutions). Titanium will also dissolve with hydrofluoric acid in the presence of an oxidizer. TiF4 will form, but I believe it is impossible to separate it from the water.

[Edited on 27-9-2013 by AndersHoveland]
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[*] posted on 27-9-2013 at 00:19


Funny, 2CaCl2 + TiO2 --> TiCl4 + 2CaO will work, but, you'd be refluxing the CaCl2 -- molten chlorides have fairly low boiling points (relative to melting) and high vapor pressures.

Fluorides: isn't there a hexafluorotitanate?

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[*] posted on 27-9-2013 at 01:59


Plante can the pyrosulphate be safely heated in borosilicate glassware, or do I need to find an alternative reaction vessel,such as a steel container?



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[*] posted on 27-9-2013 at 08:46


There is a patent for the synthesis of anhydrous TiF3, via pyrolysis of an intermediate [(NH4)2Ti(OH)F4]:

http://www.freepatentsonline.com/WO2009090513.html

Another link (that right now I can’t find) started from aqueous TiCl3 to produce a similar fluoride based Ti (III) intermediate.

TiF3 could be reduced easily with Mg powder but extra external heat would almost certainly be needed to obtain both metal and slag in the liquid state and so allow gravitational separation of metal and slag. About 2000 C would be required to reach this goal, through a combination of reaction enthalpy and added heat.

[Edited on 27-9-2013 by blogfast25]




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[*] posted on 27-9-2013 at 14:32


Here is a partial extract from Atomistry.com on TiCl4 (http://titanium.atomistry.com/titanium_tetrachloride.html ) to quote:

"Titanium Tetrachloride, TiCl4. - Titanium and chlorine combine when heated together to 350° C., forming Titanium Tetrachloride, TiCl4. In place of the pure metal that containing carbon, or the carbide, may be employed. This chloride is also conveniently prepared, like non-metallic chlorides, by passing chlorine over a heated mixture of titanic oxide and carbon, as well as by leading the vapour of carbon tetrachloride or chloroform over the heated dioxide. Ferrotitanium may also be used as a source of the tetrachloride. Most of the iron is first removed by hydrochloric acid, and the residue is heated in a porcelain tube through which chlorine is passed. Ferric chloride condenses in the cooler parts of the tube, and titanic chloride is obtained by further cooling and then fractionated. In another process rutile is reduced by aluminium according to the Goldschmidt reaction, and the product heated in a current of chlorine; the crude titanic chloride thus obtained needs to be separated by fractional distillation from silicon tetrachloride, derived from the silica of the rutile."

AndersHoveland, you mentioned a path via CCl4, apparently one can also use CHCl3.



[Edited on 27-9-2013 by AJKOER]
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[*] posted on 27-9-2013 at 21:47


If I had spare titanium lying around and my weekend weren't full, I would strongly consider asking for advice on attempting an analogous procedure to garage chemist's excellent sulfur trioxide production method with pyrosulfate and sodium chloride in the presence of ore.
https://sciencemadness.org/talk/viewthread.php?tid=5495

That or perhaps packing part of a distillation apparatus with the ore and passing distilled chloroform or carbon tetrachloride over it with gentle application of a propane torch.

[Edited on 28-9-2013 by Chemosynthesis]
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[*] posted on 28-9-2013 at 10:16


Here’s an abstract from a paper describing the synthesis of NH4/K fluorotitanates (III), from aqueous TiCl3 solutions:

http://www.sciencedirect.com/science/article/pii/S0022113900...

Abstract

Three different fluorotitanate (III) complexes of ammonium and potassium have been synthesized directly by the reaction of TiCl3 in 10% HCl with ammonium- and potassium-fluoride. The ratio of the reagents was: TiCl3: NH4F(KF)  1:5, 1:6 and 1:8, respectively. All compounds were obtained as precipitates of characteristic colour: with a Ti/F ratio  1:5, a brownish precipitate of A/TiF3(OH) (H2O)2/ (ANH4, K), was obtained. Red-violet colour of precipitate appeared when a compound A2/TiF5/ was formed, and violet colour was produced when A3/TiF6/ precipitated. When washed with ethanol and dried in vacou over concentrated sulfuric acid, no formation of peroxytitanates (IV) occurred.
Fluoride ion exchange in acidic solutions of fluorotitanates(III) complexes is described together with some preliminary work on structural characterization.




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[*] posted on 28-9-2013 at 15:48


I came across a reference ("Inorganic Chemistry" by James House, page 365, link: http://books.google.com/books?id=dDrnzWoGQC8C&pg=PA365&a... ) that gives a different product than you cited as follows:

TiO2(s) + CCl4(l) --?--> TiCl4(l) + CO2(g)

Namely, to quote:

"TiO2 + 2 CCl4 ----> TiCl4 + 2 COCl2"

and one could look at the formation of Phosgene as problematic. It could be that both reactions are correct, it is simply a matter of a difference in concentration, so as such I would avoid an excess of CCl4 relative to TiO2 here to be safe. Also, having the out gases heated to over 200 C might be helpful as COCl2 decomposes to CO and Cl2 (but could reform on cooling in the presence of carbon).

Interestingly, as CCl4 decomposes at 700 K (427 C) to Carbon and Chlorine, this preparation is the same as passing chlorine over a heated mixture of titanic oxide and carbon (as I cited above per Atomistry.com). Now, upon cooling of gases in the presence of oxygen employing the CCl4 preparation path, some Phosgene could form between 50 to 150 C via the exothermic reaction (see http://en.wikipedia.org/wiki/COCl2 ):

CO + Cl2 --Activated Carbon-> COCl2

So, I guess the creation of Phosgene is conceivable during the preparation of TiCl4 via CCl4 depending on reaction conditions.

Similarly, possible reaction (speculation) with CHCl3 could accordingly be:

TiO2 + 2 CHCl3 --?--> TiCl4 + H2O + C + CO + Cl2

assuming again a thermal decomposition of Chloroform first and then the C, Cl2 and TiO2 reaction to form TiCl4. However, as CHCl3 decomposes per one source (see http://link.springer.com/article/10.1007/s11814-012-0086-0#p... ) at around 700 C (forming CCl4, C2Cl4 and other gases depending on the inert gas used in the thermal decomposition) and per another source around 1,000 C using a different medium into CCl2 and HCl (see http://pubs.acs.org/doi/abs/10.1021/jp971723g?journalCode=jp... ), so this appears be a higher temperature path than CCl4. One must also be aware of the potential presence of water vapor (as decompositions products apparently could include HCl acting on the TiO2) on cooling given the affinity of TiCl4 for hydrolysis. Also, the preparation using Chloroform could similarly result in the creation of COCl2 upon cooling of the exit gases and collecting TiCl4, so scrub gases into water and/or ammonia:

COCl2 + H2O → CO2 + 2 HCl

COCl2 + 4 NH3 → CO(NH2)2 + 2 NH4Cl

Source: http://en.wikipedia.org/wiki/COCl2

[EDIT] Interestingly, the low temperature interaction of TiO2 and CHCl3 in the presence of sunlight has been studied (see http://proj3.sinica.edu.tw/~chem/servxx6/files/paper_7390_12... ) as a path to destroy environmental contaminants, such as halogenated hydrocarbons. More interestingly, on page 572, the author does state, to quote:

"In the thermal decomposition of CHCl3 on P25-TiO2
generating CHCl intermediate, two surface Ti-O bonds
break and two Ti-Cl bonds form as shown in Scheme II.
Therefore the process forming CO(a) involves breakage of
one C-H bond, three C-Cl bonds, and two Ti-O bonds and
formation of three Ti-Cl bonds, one CO bond, and one
O-H (as suggested by the H2O formation)."

where P25-TiO2 is a commercially available TiO2 powder described on page 567. The author's comment appears to offer some support for my speculated reaction on the action of CHCl3 on TiO2 that I presented above.

For those willing to experiment without accurate knowledge of reaction paths and potential products under varying reaction conditions, this is one preparation that should give one second thoughts.


[Edited on 29-9-2013 by AJKOER]
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[*] posted on 30-9-2013 at 05:44


In various sources, I found that TiO2 dissolves in concenrtated H2SO4, forming the green Ti(SO4)2. In non-concentrated H2SO4, it is said to form TiOSO4. One source mentioning this reaction is Mellor, but various others exist.

Has anyone tried isolating the green sulfate, and converting it to TiCl4 in a stream of HCl gas?

(Due to Ti's affinity for oxygen, wet HCl seems no option for a low temperature process.)
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[*] posted on 30-9-2013 at 12:45


Quote: Originally posted by Bezaleel  
In various sources, I found that TiO2 dissolves in concenrtated H2SO4, forming the green Ti(SO4)2. In non-concentrated H2SO4, it is said to form TiOSO4. One source mentioning this reaction is Mellor, but various others exist.

Has anyone tried isolating the green sulfate, and converting it to TiCl4 in a stream of HCl gas?

(Due to Ti's affinity for oxygen, wet HCl seems no option for a low temperature process.)


I've dissolved calcined (commercial) TiO2 in concentrated H2SO4 plus heat and time, many times. It dissolves to a yellowish solution of TiOSO4 (titanyl sulphate). The main ore of titanium, Ilmenite (FeTiO3), does dissolve in concentrated H2SO4 to a green solution, due to FeSO4 being formed too.

In less concentrated acids, only fairly freshly precipitated "Ti(OH)4.nH2O" dissolves, not the commercial pigment form.

AFAIK, the sulphate is always as titanyl sulphate, not Ti(SO4)2. If you have credible literature references to the contrary, let me know.

TiOSO4 is offered by Sigma as an unspecified hydrate so it's possible to isolate it.




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[*] posted on 30-9-2013 at 20:12


Regardless of political affiliation, due to the government shutdown, I wouldn't feel it appropriate to use the fume hoods for personal use, so I will temporarily be relegated to reactions which can be performed outside with personal protective equipment.

I am interested in setting up a stoichiometric amount of potassium pyrosulfate, NaCl, and pigment TiO2 in a test tube on a small propane burner. If melting the pyrosulfate within the glassware is sufficient to produce gaseous TiCl4 and HCl, I could perform rudimentary tests such as litmus paper and observation of smoke produced. Any alternatives or suggestions?
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[*] posted on 1-10-2013 at 08:10


Even if the action by HCl is successful, is separation of TiCl4 from the created Sulfuric acid and water an issue? I am assuming the reaction may proceed (speculation) as follows:

TiOSO4 + 4 HCl --?--> TiCl4 + H2SO4 + H2O
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[*] posted on 1-10-2013 at 08:37


Quote: Originally posted by Bezaleel  
In various sources, I found that TiO2 dissolves in concenrtated H2SO4, forming the green Ti(SO4)2.


Also, green Ti(IV) compounds are unlikely. The electron configuration of Ti(0) is [Ar] 3d2 4s2. In a Ti (IV) compound the Ti atom has lost the 3d2 4s2 electrons. Colourful D-block element compounds arise from electrons in unfilled d orbitals. There are 5 types of d sub orbitals, all with slightly different energy levels. Electrons in these unfilled orbitals can absorb VIS, ending up in higher energy level sub d orbitals, the VIS absorption giving rise to colour.

The only coloured Ti(IV) compound I know of is a red peroxo complex, a simple test tube indicator of Ti compounds.

Ti(III) however does have one lone 3d electron and Ti(III) compounds are very colourful.

[Edited on 1-10-2013 by blogfast25]




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[*] posted on 1-10-2013 at 09:20


Quote: Originally posted by blogfast25  
Also, green Ti(IV) compounds are unlikely.
Titanium disulfate is reported as a green solid. The first link that came up with a search: http://hazmap.nlm.nih.gov/category-details?id=7454&table=copytblagents.
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