Sciencemadness Discussion Board - titanium chemistry

titanium chemistry

Jor - 23-6-2009 at 15:14

Yesterday I decided to do some chemistry with titanium.
I took a spatula of 99,9% titanium powder and added a few mL of reagent grade 37% hydrochloric acid.
Next day I have a very deep purple solution. This was diluted about 10 times with dilute HCl.

With this I did many experiments, including trying to reduce it with zinc to Ti(II) (this failed), oxidation with various oxidisers in acidic environment, action of various ligands on the solution.

But I was surprised in one experiment and had NO clue what happened:
To a few mL of purple Ti(III)-solution was added a spatula full of hydrazine sulfate. Nothing happened and I heated to dissolve most of the solid. Next the solution was made alkaline with 2M NaOH. At once the liquid became colorless and a white suspension was formed, in a fraction of a second. I know titanium is quickly oxidised by air in alkaline environment, but not this fast. Besides hydrazine base is strongly reducing. I do not recall if there was any bubbling, I forgot to write this down. so I will repeat this tomorrow.

As my mother screamed to come inside

, I decided to finish the experiment. As my current waste bottle is full, I had to destroy the hydrazine, so I added some sodium nitrite (there was slow gas evolution at this point, but not sure if there wasn't before I added NaNO2) and acidified. There was very vigorous bubbling and the solution was still white suspension. I flushed it away (it was very little).

So what has happened here. the white precitipate can't be Ti(III)hydroxide, as this is blue AFAIK, and it is not TiO2.nH2O as this can't be formed this fast. So either Ti(III) forms an insoluble complex with hydrazine, or hydrazine reduces Ti(III) tot Ti(II) followed by formation of Ti(II)-complex wich is insoluble, but I have no idea what this could be.

Anyways, sorry for the lack of useful information, but i didn't have the time at that point. I will do this again, but the I will check the stability in prolonged air contact, and do some other tests.

Anyone knows what happens? Any suggestions?

not_important - 23-6-2009 at 16:50

Check the dashed green lines in the diagram, the lower is the reduction of H2O to H2. In your case Ti(III) being oxidised to Ti(IV) while reducing water.

Titanium_in_water_porbiax_diagram.png - 8kB

[Edited on 24-6-2009 by not_important]

benzylchloride1 - 24-6-2009 at 10:49

Another interesting experiment that can be conducted with aqueous solutions of Ti (III) and Ti (IV) is to add hydrogen peroxide to the solution, forming an orange/red colored solution of the peroxo complex.

woelen - 24-6-2009 at 12:40

Jor, hydrazine is a remarkable compound. It also can act as an oxidizer in rare cases, just like H2O2 which can act as an oxidizer (normal mode), but also as reductur (rare cases).

Titanium(III) is a strong reductor, so I can imagine that hydrazine acts as oxidizer in this case (the hydrazine then is converted to ammonia).

Jor - 24-6-2009 at 14:36

Not important thank you for the diagram. I completely forgot to make one myself (I have that software).

Benzyl, I know about the peroxo-complex. I have done quite some exploration involving peroxo-complex, including isolation KVO2(O2).2H2O, K3CrO8, and many aqeous experiments.

Woelen, you might be right. Can you give me an example of an oxidation involving hydrazine as oxidiser?

I really need to some more experiments. I always thought Ti(OH)3 reduces water only slowly, but now I remember the liquid was quite hot (dissolving hydrazine sulfate), so this might made the reaction very fast. Also I am not sure if there was vigorous bubbling after adding it, as I spilled a few drops of NaOH, so that always draws my full attention, regardless of what happens else in my eyesight (so I don't 'save' other information at that point). So the results I gave you are pretty incomplete I must say, so i will perform some otehr experiments tomorrow probably, with more accurate observations.

What is an ideal way of destroying hydrazine? Nitrous acid right?

While we're at hydrazine, I have been considering buying hydrazine hydrate (100mL from Aldrich), but i am not sure what to expect of it's dangers. MSDS says it's toxic, carcinogenic and corrosive. Sounds familiar to many otehr chemicals right?
But I have heard stories about this chemicals that it is so evil. What makes this chemical so dangerous?

Lambda-Eyde - 24-6-2009 at 15:03

The anhydrous form is extremely unstable and dangerous to handle. An aqueous solution, on the other hand, is not that dangerous/evil.

DJF90 - 24-6-2009 at 16:21

Hydrazine also needs to be distilled under inert gas, due to explosion possibility. IIRC "anhydrous" hydrazine which is a little wet will attack glass (etching).

zed - 24-6-2009 at 16:45

I would suggest it is the N-N linkage. It isn't very common, and where it exists, toxicity and/or instability often follow.

In simple terms, molecules with that linkage....."Like"..."to unload" and become the very stable N2. If the hydrogen present can be reacted with O2 simultaneously, all the better. Lots of energy will be released in the process, and it will be released very quickly. Without a large volume of H20 acting as a heat sink, there may be an explosion.

That would explain explosiveness.

I have read a bit about the toxicity of such materials, but haven't yet found a universal explanation that encompasses all azides, azines, etc.. Suffice to say, such materials are highly reactive, and they react with and block critical enzyme systems in your body, rendering you dead.

You have aroused my curiosity. I will read further.

not_important - 24-6-2009 at 20:56

Jor, I think you need to take another look at that diagram. Ti(II) doesn't look to be stable in water without some unusual coordination stabilization.

I think that if you repeated your experiment, dividing the Ti(III) solution between two test tubes, then adding equal molar amounts of hydrazine sulfate to one and sodium sulfate to the other, and finally adding the NaOH, you'll see similar results in both tubes. Once the pH moves out of being strongly acid, Ti(III) can react with water.

As for hydrazine, this PDF http://www.ehponline.org/members/1985/064/64016.PDF suggests that in part its toxicity derives from free radical reactions. Hydrazine also damages DNA through direct and indirect means.

woelen - 26-6-2009 at 01:04

Quote: Originally posted by Jor

Woelen, you might be right. Can you give me an example of an oxidation involving hydrazine as oxidiser?

A documented example of hydrazine acting as oxidizer is with Ti(3+)

I by accident stumbled on it.

Look in the energetic materials thread on chlorine azide in the post just before my last post and download that document. At the end of page 120 this book mentions hydrazine as oxidizer when combined with Zn, Sn, Sn(2+) and Ti(3+).

Arthur Dent - 26-6-2011 at 08:29

Sorry to revive this antique thread, but I have come across a fairly good chunk of titanium (the top cover of a mac titanium powerbook. It appears like the real thing, a vaguely yellowish metal. So I tought that i'd use the most of it to make strips that i'll use as electrodes in my future electrochemistry experiments but I was interested in the chemistry of Titanium.

So I looked-up Brauer's and there are mentions of TiO2 and TiCl4, and other ceramic-like compounds like the nitride... but it appears that there are few interesting chemicals that can be synthetized without extreme temperatures and noxious gases. Following the OP's experiment, I put a few bits of metallic titanium in a test tube and dropped some HCl but nothing so far. No colorful salt or interesting, stable compound can be easily synthetized, reminds me of Germanium...

So i guess that I'd rather stick to using the metal for electrodes and that will be pretty much it... Unless someone has interesting recipes (that don't involve a kiln or streams of hot HCl or HF gas, LOL).

Robert

woelen - 26-6-2011 at 10:47

Put one of the strips in a solution of NaF in dilute HCl. If it is titanium it quickly dissolves, giving a green solution.

LanthanumK - 26-6-2011 at 11:23

Did you use concentrated HCl or just dilute stuff?

cyanureeves - 26-6-2011 at 11:32

lightly boil a piece in hcl acid like you do to make tin chloride and if its titanium it will turn purple. this stuff will boil and boil until the end of time and seem never to totally dissolve.even after turning the acid purple from the get go it will last forever. i ended up with titanium dioxide or carbonate when i tried to keep up with plante1999.the acid made the bottom of my pyrex rough and is still frosty looking.now i know why they coat this stuff and use as electrodes also you can do the annodizing thing by putting titanium annode and cathode in a solution and electrolyze. the negative side will turn rainbow color and eventually form a coat that wont respond to current.

woelen - 26-6-2011 at 12:44

If the titanium is in the form of sheets or lumps then using HCl only hardly is feasible. Dissolving the metal in dilute HCl takes sooooo looooong. Even in conc. HCl it will take days.

If you have NaF (or KF or NH4HF2) though, then just 10% HCl with some NaF added will do the job amazingly well. The titanium then dissolves quickly, bubbles of hydrogen being produced vigorously. In a few tens of minutes you can dissolve a few mm thichness of metal with a dilute HCl/NaF solution (or HCl/HF). The result of dissolving the metal then is green instead of purple.

http://woelen.homescience.net/science/chem/riddles/titanium+...

I'm not sure about the precise nature of the green compound. It most likely is some fluoro-complex of titanium(III) or a mixed chloro/fluoro complex of titanium(III).

Arthur Dent - 27-6-2011 at 05:28

Thanks guys for the excellent suggestions. Indeed this morning, the acid with the metal pieces was a very lovely shade of violet, and a bit of heat from my heat gun deepened the color. So I got myself a very nice chunk of titanium.

@ Woelen, I do not have NaF, but I do have a bottle of CaF2, would this work?

@ cyanureeves, indeed upon heating, the metal bubbled vigorously but hasn't decreased in size, so this purple stuff is TiCl3? I wonder if i'll be able to dehydrate to crystalline form, although I don't see the point since it's supposed to be very hygroscopic.

EDIT: I read the MSDS on TiCl3 and it seems hard to keep, they even suggest to store under an inert atmosphere. I'll fill a little vial with the solution and see if it changes over time...

Robert

[Edited on 27-6-2011 by Arthur Dent]

blogfast25 - 27-6-2011 at 07:28

Quote: Originally posted by woelen

If the titanium is in the form of sheets or lumps then using HCl only hardly is feasible. Dissolving the metal in dilute HCl takes sooooo looooong. Even in conc. HCl it will take days.

It's a bit faster than that. I dissolved Ti pieces (about one gram each, thermite metal) in 20 % HCl at reflux in a few hours. I also produced standardised Ti3+ solutions from a 99.9 % Ti bulk artefact in about the same amount of time. But to dissolve entire chunks in even sat. HCl would be impractical.

For Ti salts, I use TiO2 + NaHSO4 (fusion) to obtain TiOSO4 (titanyl sulphate), this can then be reduced with Al or Zn to Ti3+ if required.

Metacelsus - 7-3-2014 at 17:50

While trying to make titanyl sulfate, I have just reacted TiO₂ (99%, 10 grams) with excess sulfuric acid (98%, 20 ml) at 125 C. I got a viscous black sludge. Any ideas to what happened (besides that the seller of the titanium dioxide possibly ripped me off) or how I could recover the titanyl sulfate?

plante1999 - 7-3-2014 at 18:17

I think I have been asked... If the mixture turned black, you may want to look to the purity of your sulphuric acid, stabilizer are often organic, make a blank with only sulphuric acid. Then, if it does not turn black, you can still use physical properties of titanium dioxide to analyse it.

Bisulphate fusion is called for to dissolve titanium dioxide, or in-situe making titanium tetrachloride. You can test your solution for titanium using diluted hydrogen peroxide, it will get red if titanium is present.

As for recovering it, I would do test first. It is actually possible to do fusion with sulphuric acid, but you have to heat the sulphuric acid until it stops fuming, which takes time.

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

For people looking for my sources & references, it is my personal experience in the field.

jock88 - 8-3-2014 at 06:03

Quote: Originally posted by woelen

If the titanium is in the form of sheets or lumps then using HCl only hardly is feasible. Dissolving the metal in dilute HCl takes sooooo looooong. Even in conc. HCl it will take days.

If you have NaF (or KF or NH4HF2) though, then just 10% HCl with some NaF added will do the job amazingly well. The titanium then dissolves quickly, bubbles of hydrogen being produced vigorously. In a few tens of minutes you can dissolve a few mm thichness of metal with a dilute HCl/NaF solution (or HCl/HF). The result of dissolving the metal then is green instead of purple.

...

Would ammonium biflouride do instead of the NaF ?

blogfast25 - 8-3-2014 at 06:15

Quote: Originally posted by jock88

Quote: Originally posted by woelen

If you have NaF (or KF or NH4HF2) though, then just 10% HCl with some NaF added will do the job amazingly well. ...

Would ammonium biflouride do instead of the NaF ?

Which part of 'KF or NH4HF2' didn't you understand?

Titanium metal also dissolves fairly well in conc. hot H2SO4, although I've only ever used Ti powder for that.

jock88 - 8-3-2014 at 13:23

Sorry about that!

Grade 5 (contains some Aluminium(6%) and Vanadium(4%)) will dissolve more easily than grade 1, 2, 3 or 4 (only ever had grades 1, 2, and 5 myself).
You are going to get Al + V contamination with the grade 5 I would imagine.

Metacelsus - 8-3-2014 at 16:10

"Black crud" problem solved: I heated it like crazy for around an hour, and the reaction worked.

blogfast25 - 9-3-2014 at 10:50

Quote: Originally posted by Cheddite Cheese

"Black crud" problem solved: I heated it like crazy for around an hour, and the reaction worked.

You mean it dissolved? In that case you're fairly lucky: my anatase food grade TiO2 isn't even dented by 2 h of 200 C 98 % H2SO4. Only alkali fusion worked. Its solubility is strongly dependent on grade.

Your black crud may have been due to organic coatings that are sometimes applied to commercial TiO2, to improve certain characteristics in real life applications.

[Edited on 9-3-2014 by blogfast25]

Morgan - 9-3-2014 at 19:20

Could your anatase food grade titanium dioxide be coated with Al203 or SiO2 and if so in such a way as to result in the insolubility in sulfuric acid?

"Titanium dioxide (E 171, INS 171) is approved for use in food by the European Union, by the United States FDA and by the Codex Alimentarius of the FAO/WHO. The Joint WHO/FAO
Expert Committee of Food Additives (JECFA) evaluated titanium dioxide and allocated an Acceptable Daily Intake not specified (JECFA 1969). In addition to the safety evaluation,
JECFA established a set of purity criteria for titanium dioxide which do not differentiate between the anatase and rutile forms of titanium dioxide. In the European Union, titanium
dioxide (E171) is included in the list of approved colouring agents in Directive 94/36/EC. The purity criteria, however, mention explicitly that titanium dioxide essentially consists of the pure anatase form which may be coated with small amounts of alumina and/or silica to improve the technological properties of this product."
http://www.efsa.europa.eu/en/efsajournal/doc/163.pdf

"Titanium dioxide may be coated with small amounts of alumina and silica to improve technological properties. Such coatings can prevent possible reactions between the highly reactive surfaces of the extremely fine titanium dioxide crystals and the matrix in which the pigment is dispersed and they can improve the dispersion of the titanium dioxide in the host matrix (Kirk-Othmer, 2006)."
ftp://ftp.fao.org/ag/agn/jecfa/cta_tio2.pdf

blogfast25 - 10-3-2014 at 14:24

Morgan:

Both pdfs are interesting but I doubt if these minute additions have much effect on the bulk solubility of the anatase in conc. H2SO4.

Metacelsus - 11-3-2014 at 10:45

Most, but not all, of it dissolved. The insoluble stuff is a grey powder. The solution is very viscous; I will probably have to dilute it with water in order to filter it.

blogfast25 - 11-3-2014 at 13:39

CC:

Yes, concentrated solutions of titanyl sulphate are quite viscous. Over-dilution will cause the oxide to precipitate again (the basis of the 'sulphate process' for TiO2). Unless you have glass (or ceramic) filters, decantation may be better than filtering.

[Edited on 11-3-2014 by blogfast25]

Metacelsus - 11-3-2014 at 17:10

Quote: Originally posted by blogfast25

Over-dilution will cause the oxide to precipitate again (the basis of the 'sulphate process' for TiO2).
[Edited on 11-3-2014 by blogfast25]

That's good to know. No, I don't have anything other than paper filters. Decantation is going to be tricky.

About how much can I dilute it before the titanyl sulfate hydrolyzes?

blogfast25 - 12-3-2014 at 05:41

Quote: Originally posted by Cheddite Cheese

About how much can I dilute it before the titanyl sulfate hydrolyzes?

Not sure really, I know it also depends on temperature.

Take a ml of the original solution and dilute it 1:1. Then take 1 ml of that dilution and dilute it again 1:1. Carry on until cloudiness sets in. Use the dilutions to check your paper: it gets chewed up badly if the acid is too strong.