I am a fish
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Beautiful Titanium-based Multilayer Precipitate.
I dissolved an excess of titanium granules in dilute hydrofluoric acid (generated in situ, by placing the titanium in 10% hydrochloric acid, and then
adding sodium fluoride solution). The reaction proceeded slowly, and at first, resulted is a green solution, presumably due to the presence of
[TiF<sub>6</sub>]<sup>3-</sup> ions. A day later, the colour had changed to purple, presumably due to a drop in hydrogen
fluoride concentration, causing the presence of uncomplexed titanium 3+ ions (strictly speaking
[Ti(H<sub>2</sub>O)<sub>6</sub>]<sup>3+</sup> ions).
I decanted the solution off, and started experimenting with coordination complexes. I noticed that basic compounds produced a thick, grainy and very
dark precipitate. (I first noticed this with ethylene diamine, and then produced exactly the same result with ammonia, and finally with potassium
hydroxide).
In order to investigate this precipitate further, I added the remainder of my titanium solution to an excess of potassium hydroxide solution, and left
the mixture to settle overnight.
The following day, I was confronted with the most beautiful precipitate I have ever seen. It was divided into four layers. The top layer was snowy
white, the second layer was a light grey-blue colour, the third layer was a deeper grey-blue colour and the bottom layer was a blue-black colour. It
should be noted that the the colours did not form a continuum, but instead were divided into four discrete layers.
I don't know what the precipitate is, and will investigate it further. Also, the author of Science Made Alive briefly mentions a similar precipitate, but makes no mention of layering.
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Pyridinium
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This sounds great. Do you have any pictures of it? I'd really love to see it.
Four layers, wow. I love chemistry.
Have you tried oxidation of Ti with 5-10% sulfuric acid and 3% H2O2? If it's free of fluoride, Fe, Ni, Cr, it will make a nice colored complex.
(can't remember offhand what the formula is).
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I am a fish
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Hydrofluoric acid oxidizes titanium to the Ti<sup>3+</sup> state. However, I think a mixture of sulphuric acid and hydrogen peroxide
would oxidize it further, to the Ti<sup>4+</sup> state. When I added hydrogen peroxide to my titanium(III) fluoride solution, it turned
bright red, which is the distinctive colour of Ti<sup>4+</sup> complexed with hydrogen peroxide.
Nevertheless, I will see what happens when I prepare the titanium solution in this manner. I am currently dissolving the titanium, but the reaction
is progressing very slowly, and so I have nothing yet to report.
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Lambda
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| Quote: | Originally posted by I am a fish
I dissolved an excess of titanium granules in dilute hydrofluoric acid (generated in situ, by placing the titanium in 10% hydrochloric acid, and then
adding sodium fluoride solution). |
I am a fish, personly I would rather keep the titanium granules. Titanium dioxide and aluminium trioxide are one of the most used white pigments.
Titanium dioxide, having a overwelming white appearance, is as pigment incredibly cheap. Hughe quantaties can be obtained via any paintstore that
mixes it's own pigments.
Could this not be a cheap resource to conduct your very interesting experiment with ?
Just adding my two cents. I to, look foreward in seeing pictures of your results.
Good luck !
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Pyridinium
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| Quote: | Originally posted by Lambda
I am a fish, personly I would rather keep the titanium granules. Titanium dioxide and aluminium trioxide are one of the most used white pigments.
Titanium dioxide, having a overwelming white appearance, is as pigment incredibly cheap. Hughe quantaties can be obtained via any paintstore that
mixes it's own pigments.
Could this not be a cheap resource to conduct your very interesting experiment with ?
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It depends on how the TiO2 was made. Precipitated TiO2 (from aqueous solutions) will dissolve in acid. TiO2 made by high temperature ignition will
not, at least not very well.
I don't know why there's a difference, except maybe there's an alpha and beta modification or something.
I am a fish, I looked up that H2O2 complex formula with sulfuric acid. It's TiO2(SO4)2(2-). Orange to red. Apparently you may also get some
H4TiO5.
It sounds as though at least one of your precipitate layers is a hydrogel, made by neutralizing an acidic Ti ion solution with excess alkali. I have
a 1940 reference here that says this gel has been analyzed with X-ray diffraction, and they say that, if it's allowed to age, the pattern is the
same as the anatase form of TiO2. Interesting!
The bottom layer may also have Ti(OH)3, slowly turning to TiO2. Just a guess though.
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12AX7
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| Quote: | Originally posted by Pyridinium
It depends on how the TiO2 was made. Precipitated TiO2 (from aqueous solutions) will dissolve in acid. TiO2 made by high temperature ignition will
not, at least not very well. |
Probably surface area, same thing with alumina and magnesia, big refractory products. Calcined (and dead burned, respectively) are pretty inert.
Tim
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I am a fish
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| Quote: | Originally posted by Pyridinium
Have you tried oxidation of Ti with 5-10% sulfuric acid and 3% H2O2? |
I added titanium to such a mixture (see above), and after a few days I obtained a deep red solution. However, on adding potassium hydroxide solution,
I merely obtained a plain white precipitate, which I assume is hydrated titanium dioxide.
This result in unsurprising, as the precipiation was from a Ti<sup>4+</sup> solution (rather than a Ti<sup>3+</sup> solution).
However, I will try reducing the solution with zinc first, and then see what precipitate I get.
It has also occured to me that Ti<sup>3+</sup> ions are oxidized by the oxygen in air. Therefore, my four layer precipitate may have
contained mixed oxidation states.
| Quote: | Originally posted by Lambda
I am a fish, personly I would rather keep the titanium granules. Titanium dioxide and aluminium trioxide are one of the most used white pigments.
Titanium dioxide, having a overwelming white appearance, is as pigment incredibly cheap. Hughe quantaties can be obtained via any paintstore that
mixes it's own pigments.
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This doesn't matter a great deal, as I usually work on a semi-microscale. However, I will look into using this route (see below).
| Quote: | Originally posted by Pyridinium
Precipitated TiO2 (from aqueous solutions) will dissolve in acid. TiO2 made by high temperature ignition will not, at least not very well.
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I will try dissolving pottery grade titanium dioxide in sulphuric acid. I will also try dissolving it in a mixture of sulphuric acid and hydrogen
peroxide. (Although the hydrogen peroxide is not needed as an oxidizer (as the titanium is already in a Ti<sup>4+</sup> state), it may
help dissolution, by forming a complex with the product.)
[Edited on 21-6-2005 by I am a fish]
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I am a fish
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Update:
The four layer precipitate has now become a two layer precipitate. The blue-black bottom layer has remained the same, whilst the rest of the
precipitate has turned white. This could be because the intermediate layers consisted of mixtures of the upper and lower components, which have now
fully seperated. However, if this was the case, I expect that a continuum between white and deep blue would have initially formed, rather than
discrete layers. Another possibility is that the intermediate layers have been oxidised into the same substance (probably
TiO<sub>2</sub> as the top layer.
| Quote: | | I will try dissolving pottery grade titanium dioxide in sulphuric acid. I will also try dissolving it in a mixture of sulphuric acid and hydrogen
peroxide. (Although the hydrogen peroxide is not needed as an oxidizer (as the titanium is already in a Ti<sup>4+</sup> state), it may
help dissolution, by forming a complex with the product.) |
I added pottery grade TiO<sub>2</sub> to 10% sulphuric acid about a week ago. Yesterday, when I added hydrogen peroxide to the mixture,
it turned a faint orange/yellow colour. I assume that this is due to low concentration of the strongly coloured Ti<sup>4+</sup> /
H<sub>2</sub>O<sub>2</sub> complex. This indicates that a reaction had occured, but only to a very small extent. Since
adding the H<sub>2</sub>O<sub>2</sub>, I have not noticed any increase in the colour's intensity, and so its presence
cannot be significantly accelerating the reaction.
| Quote: | | I will try reducing the solution with zinc first, and then see what precipitate I get. |
I added an excess of zinc powder to my Ti<sup>4+</sup> / H<sub>2</sub>O<sub>2</sub> solution. I decanted off a
liquid, which was the characterstic purple colour of Ti<sup>3+</sup>. I then added potassium hydroxide solution. However, instead of the
previous result, I obtained a dark grey metallic-looking sheet-like precipitate, which later faded to white.
Titanium chemistry is starting to baffle me. I will try the experiment with my original setup, to see if I can repeat the original result.
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woelen
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I'm new to scimad and found this topic. I'm also very interested in titanium chemistry and I found an even more interesting phenomenon.
Titanium (III) and titanium (IV) form a mixed-valence complex with multiple titanium ions in a single ion. The exact nature of the stuff, I did not
yet determine. I added some pictures on my site about the titanium experiment. See here.
I also did more experiments with the blue/black precipitate of titanium (III). It is oxidized by water, forming hydrogen gas, itself being converted
to white hydrous titanium (IV) oxide. It is most remarkable that a metal hydroxide forms hydrogen gas with water! Nice and detailed pictures of this
will follow in due time on my site.
The layered structure of the precipitate, I indeed did not observe, but what I sometimes did observe is that the precipitate is changing color from
dark to white slowly, starting at the surface of the liquid. This probably is due to oxidation by oxygen from the air.
[Edited on 29-8-2005 by woelen]
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chloric1
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I am fish, I am thinking some of your layers contained Ti+3 and these, like the ferrous ion, are much more suseptible to oxidation by air when in an
alkaline state. If you can reduce vanadium to the bivalent state add alkali and I am sure you would find simular results.
As far as titanium white goes, I think it shares the property with niobium and tantalum oxides in that a bisulfate fusion shall make them soluble.
Woelen, if I knew trivalent Ti was that beautiful I would have tried long ago! Where did I put my Titanium sheet !?!
[Edited on 8/30/2005 by chloric1]
[Edited on 8/30/2005 by chloric1]
Fellow molecular manipulator
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woelen
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This is a detailed picture of the precipitate of titanium (III) hydroxide, which reacts with water to form hydrogen gas:
http://81.207.88.128/science/chem/solutions/tiIIIoh.jpg
This is the same precipitate a little later. You can already see it turn lighter near the surface of the liquid.
http://81.207.88.128/science/chem/solutions/tiIIIoh-2.jpg
When it is shaken:
http://81.207.88.128/science/chem/solutions/tiIIIoh-3.jpg
When it is shaken in contact with air for a longer time:
http://81.207.88.128/science/chem/solutions/tiIVoh-2.jpg
The following also is nice. It is the complex of titanium (IV) with hydrogen peroxide at moderate concentration:
http://81.207.88.128/science/chem/solutions/tiIVh2o2.jpg
These pictures are part of a section in my website about many transition metals (all first row and many second and third row). I want to make an
overview of all colors of all oxidation states, all precipitates of hydroxides, many common complexes and all solutions of higher oxidation states.
This is a large job and it may take a few more months, but if it is finished, then I'll certainly let you know. Pictures from it will be
published sometimes earlier, when appropriate.
I most like titanium, vanadium and chromium. These are really interesting .
[Edited on 30-8-2005 by woelen]
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chloric1
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WOW!
Those are really great! I especially like how the Ti+3 hydroxide on shaking with air turns into the white dioxide.
Speaking of the dioxide I know that the dioxide mixed with carbon will react with dry chlorine gas at about 500 to 600 C, do you know what are the
conditions of the chlorine on the pure metal? Is it very dangerous? To me it would be impractical like heating aluminum in chlorine.
Fellow molecular manipulator
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I am a fish
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My original precipitate is now completely white. The two middle layers were presumably different mixed-valence complexes, which were both highly
susceptible to oxidation (hence their short lifespan). The bottom layer was presumably titanium(III) hydroxide, whislt the white top layer was always
titanium dioxide.
| Quote: | Originally posted by woelen
When it is shaken in contact with air for a longer time:
http://81.207.88.128/science/chem/solutions/tiIVoh-2.jpg
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The colours in this image are very similar to those of my original precipitate. The particles suspended in the liquid appear to have different
colours (although this could be an illusion caused by them being at different distances from the surface of the tube, and so being occluded by
different amounts), and so could be of differing chemical composition. I expect that my coloured layers formed from a mixture of very similar
composition to that shown in the image. If you had sealed the tube at this point, and had let the mixture settle, you may have ended up with the same
multi-layer phenomenon.
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