Sciencemadness Discussion Board

Titanium IV Perchlorate Ti(ClO4)4

DubaiAmateurRocketry - 15-6-2013 at 06:12

I recently got interested into Titanium IV perchlorate.

In the preview of this report, (I wish some one could purchase this and give it to us for free)
http://link.springer.com/article/10.1007/BF00922105

It stated that Ti(ClO4)4 decompose into TiO2 and the rest chlorine and oxygen, this is such a high amount of oxdizier released !!

Only 17.9% of this compound is turned into solid upon decomposition, and the rest is oxidizers, even the TiO2 could react with aluminum !

It contains more oxidizer by weight than AP and denser. However its reported that it explode at 130 degree, but i dont think it would be that dangerous, since decomposition is endothermic.

I also found a place to purchase this, extremely expensive though.

[Edited on 15-6-2013 by DubaiAmateurRocketry]

Fantasma4500 - 15-6-2013 at 06:26

interesting, i think what approach you would want would be dilute HClO4 + Ti as Ti even reacts with HCl of what ive heard

whats the price for say.. 100g Ti(ClO4)4?? sounds expensive, not only by the idea that titanium compounds are pricey but tetraperchlorate..??

DubaiAmateurRocketry - 15-6-2013 at 06:29

Quote: Originally posted by Antiswat  
interesting, i think what approach you would want would be dilute HClO4 + Ti as Ti even reacts with HCl of what ive heard

whats the price for say.. 100g Ti(ClO4)4?? sounds expensive, not only by the idea that titanium compounds are pricey but tetraperchlorate..??


In your dreams bro, Titanium Tetra perchlorate reacts with water.

Making it, you need a solution of 98% Anhydrous perchloric acid and 2% Cl2O7 as catalyst. Which is why it is so expensive.

The price for 250 grams of TP is almost one dollar per gram..

[Edited on 15-6-2013 by DubaiAmateurRocketry]

Trotsky - 15-6-2013 at 08:38

That's actually pretty cheap. I mean, comparably. Try buying gaboxadol or anything cool for that price.

Someone should tell the coke dealers that a dollar per gram is expensive ;)

If it's as good as it looks a few grams will pack a powerful punch. I'd worry about ddt if you're thinking of it as a propellant. That won't work well

DubaiAmateurRocketry - 15-6-2013 at 08:59

Quote: Originally posted by Trotsky  
That's actually pretty cheap. I mean, comparably. Try buying gaboxadol or anything cool for that price.

Someone should tell the coke dealers that a dollar per gram is expensive ;)

If it's as good as it looks a few grams will pack a powerful punch. I'd worry about ddt if you're thinking of it as a propellant. That won't work well


It might be, I will try purchase around 50 gram first, It is not reactive at all, if it had a reactivity NFPA it would be maximum 2.

Also the decomposition composition is..

Ti(ClO4)4 + Heat = TiO2 + 7O2 + 2Cl2

Although it first makes some chlorine oxides, it will furthur decompose.

[Edited on 15-6-2013 by DubaiAmateurRocketry]

hyfalcon - 15-6-2013 at 10:02

Quote: Originally posted by DubaiAmateurRocketry  
Quote: Originally posted by Antiswat  
interesting, i think what approach you would want would be dilute HClO4 + Ti as Ti even reacts with HCl of what ive heard

whats the price for say.. 100g Ti(ClO4)4?? sounds expensive, not only by the idea that titanium compounds are pricey but tetraperchlorate..??


In your dreams bro, Titanium Tetra perchlorate reacts with water.

Making it, you need a solution of 98% Anhydrous perchloric acid and 2% Cl2O7 as catalyst. Which is why it is so expensive.

The price for 250 grams of TP is almost one dollar per gram..

[Edited on 15-6-2013 by DubaiAmateurRocketry]



Correct me if I'm wrong, but 98% perchloric acid is explosive in and of itself. Anything over 70% and you're asking for an explosion.

Trotsky - 15-6-2013 at 11:17

Does anyone know if you could form phenyl titanium triperchlorate? Or perhaps trinitro phenyl titanium triperchlorate? Or perhaps some other organic substitutent to make for a more neutral OB, remove the need for an additional fuel source.

franklyn - 15-6-2013 at 12:17

Bretherick's Handbook of Reactive Chemical Hazards Vol 1
[ 4 1 7 0 ] Titanium Tetraperchlorate
Kirk Othmer, 4th. Edn., 1996, Vol. 18, 161
The compound sublimes from 70 °C and explodes on heating to 130 °C
at atmospheric pressure. US Patent 3157464
Described as insensitive to heat or shock, this powerful oxidant explodes
on contact with Diethyl ether , and ignites with Formamide , or Dimethylformamide.


Studies of Complex Perchlorates

www.dtic.mil/dtic/tr/fulltext/u2/430803.pdf

www.dtic.mil/dtic/tr/fulltext/u2/607947.pdf

www.dtic.mil/dtic/tr/fulltext/u2/634105.pdf

___________________________________


Another Oxidizer of Interest

The reaction between zinc and liquid dinitrogen tetroxide involves the evolution of
nitric oxide, and the formation of a compound Insoluble In the tetroxide havlng the
empirical formula Zn(NO3)2 • 2N2O4. Tha action of heat on this compound, and its
reactions with organIc solvents ( particularly ethyl ether ) are considered in the light
of possible structures for the compound. Heating at 100 º results In the removal of
the two moIocuies of dinltrogen tetroxide, leaving zinc nitrate in the pure anhydrous
form. This method of preparation is simpler than methods previously described.
Zinc nitrate is stable at 100 º, but decomposes slowly in the temperature range
100 - 240 º, and rapidly above 240 º.

For this abstract and first page , of part VII , the first cited below.
http://pubs.rsc.org/en/content/articlelanding/1951/jr10.1039...

The Liquid Dinitrogen Tetroxide Solvent System

Part VII. Products of reaction of zinc with liquid dinitrogen tetroxide
J. Chem. Soc. , 1951 , 2829 - 2833

Part VIII. Rates, and possible mechanisms, of reaction of zinc with liquid dinitrogen tetroxide
J. Chem. Soc. , 1951 , 2833 - 2838

Part IX. Products, rates, and possible mechanisms of the reaction of zinc with liquid dinitrogen trioxide–tetroxide mixtures
J. Chem. Soc. , 1951 , 2838 - 2843


This whole series of papers related to N2O4 are cited here , if you have access _
www.sciencemadness.org/talk/viewthread.php?tid=19098&pag...

____________________________________


Ti(ClO4)4
As you point out produces 17.9 % solid ( TiO2 )
9 mols of gas and one mol of TiO2 or 10 %

LiClO4
produces just 14 % solid ( Li2O )
9 mols of gas and 2 mols of Li2O or 18.2 %

Zn(NO3)2 • 2N2O4
produces 21.8 % solid ( ZnO )
21 mols of gas and 2 mols of ZnO or 8.7 %

The density of the final pyrotechnic mixture will
also factor significantly in the performance derived.

.

[Edited on 16-6-2013 by franklyn]

AndersHoveland - 15-6-2013 at 15:29

Quote:

Nonsolvated titanium perchlorate was produced by the action of TiCl4 on anhydrous perchloric acid (containing up to 2% Cl2O7). Titanium perchlorate possesses an appreciable intrinsic vapor pressure and can be repeatedly sublimed at 70 °C under vacuum. The properties of sublimed Ti(ClO4)4 differ from the properties of the starting material.

It was established by the method of differential thermal analysis that Ti(ClO4)4 melts with partial decomposition at 85–95 °C, and above 110°C under vacuum it decomposes to TiO2, O2, and oxides of chlorine. The heating of Ti(ClO4)4 at atmospheric pressure ends in explosive decomposition at ∼130 °C.

Ti(ClO4)4 dissolves vigorously in water, it dissolves with decomposition in nitromethane and acetonitrile - the solution turns yellow. It is insoluble in CCl4 but above 25 °C it reacts, forming a yellow solution.

Titanium perchlorate is a colorless hygroscopic crystalline substance. ... The density of Ti(ClO4)4 ... is equal to 2.35 g/cm3.

"Volatile titanium perchlorate", V. P. Babaeva, V. Ya. Rosolovskii,
Bulletin of the Academy of Sciences of the USSR, Division of chemical science, November 1974, Volume 23, Issue 11, pp 2330-2334


Quote: Originally posted by Trotsky  
Does anyone know if you could form phenyl titanium triperchlorate? Or perhaps trinitro phenyl titanium triperchlorate? Or perhaps some other organic substitutent to make for a more neutral OB, remove the need for an additional fuel source.

It seems that Ti(ClO4)4 spontaneously reacts with any organic compound (much like anhydrous Al(ClO4)3 does). I do not think it would be any less oxidizing if it just had 3 perchlorate groups.

However, it might be possible a Lewis adduct could be formed. US patent 2,363,684 describes an aduct (6:1) between DMSO and Al(ClO4)3, though it can be noted that the reaction did not involve the anhydrous form of aluminum perchlorate, but rather the DMSO was used to displace the water of hydration. Anhydrous Al(ClO4)3 immediately oxidizes most organic solvents, though it did react to form 1:3 adducts with both nitromethane and tetrahydrofuran.
("Solvation, Solvolysis, and Complexing of Anhydrous Aluminum Perchlorate in Anhydrous Media", Z.K. Nikirina, V. Ya. Rosolovskii, Izvestiya Akademii Nauk SSSR, Seriya Khimicheskaya, No. 6, pp. 1223–1227, June, 1980.)

Another possibility is mixed azide-perchlorate complexes. Both Ti(ClO4)62‒  and Ti(N3)62‒  exist.

Though consideration would be needed, because anhydrous Ti(ClO4)4 would likely immediately oxidize any free azide ions (though once the azide groups are part of the complex they will no longer be vulnerable to oxidation, as they will be electron donating).

And I also doubt Ti(N3)4 would be a strong enough Lewis acid to form an adduct with perchlorate ions, so the resulting complex ion would probably have to contain at least 3-4 perchlorate groups.

[Edited on 16-6-2013 by AndersHoveland]

DubaiAmateurRocketry - 16-6-2013 at 00:38

Quote: Originally posted by franklyn  
Bretherick's Handbook of Reactive Chemical Hazards Vol 1
[ 4 1 7 0 ] Titanium Tetraperchlorate
Kirk Othmer, 4th. Edn., 1996, Vol. 18, 161
The compound sublimes from 70 °C and explodes on heating to 130 °C
at atmospheric pressure. US Patent 3157464
Described as insensitive to heat or shock, this powerful oxidant explodes
on contact with Diethyl ether , and ignites with Formamide , or Dimethylformamide.


Studies of Complex Perchlorates

www.dtic.mil/dtic/tr/fulltext/u2/430803.pdf

www.dtic.mil/dtic/tr/fulltext/u2/607947.pdf

www.dtic.mil/dtic/tr/fulltext/u2/634105.pdf

___________________________________


Another Oxidizer of Interest

The reaction between zinc and liquid dinitrogen tetroxide involves the evolution of
nitric oxide, and the formation of a compound Insoluble In the tetroxide havlng the
empirical formula Zn(NO3)2 • 2N2O4. Tha action of heat on this compound, and its
reactions with organIc solvents ( particularly ethyl ether ) are considered in the light
of possible structures for the compound. Heating at 100 º results In the removal of
the two moIocuies of dinltrogen tetroxide, leaving zinc nitrate in the pure anhydrous
form. This method of preparation is simpler than methods previously described.
Zinc nitrate is stable at 100 º, but decomposes slowly in the temperature range
100 - 240 º, and rapidly above 240 º.

For this abstract and first page , of part VII , the first cited below.
http://pubs.rsc.org/en/content/articlelanding/1951/jr10.1039...

The Liquid Dinitrogen Tetroxide Solvent System

Part VII. Products of reaction of zinc with liquid dinitrogen tetroxide
J. Chem. Soc. , 1951 , 2829 - 2833

Part VIII. Rates, and possible mechanisms, of reaction of zinc with liquid dinitrogen tetroxide
J. Chem. Soc. , 1951 , 2833 - 2838

Part IX. Products, rates, and possible mechanisms of the reaction of zinc with liquid dinitrogen trioxide–tetroxide mixtures
J. Chem. Soc. , 1951 , 2838 - 2843


This whole series of papers related to N2O4 are cited here , if you have access _
www.sciencemadness.org/talk/viewthread.php?tid=19098&pag...

____________________________________


Ti(ClO4)4
As you point out produces 17.9 % solid ( TiO2 )
9 mols of gas and one mol of TiO2 or 10 %

LiClO4
produces just 14 % solid ( Li2O )
9 mols of gas and 2 mols of Li2O or 18.2 %

Zn(NO3)2 • 2N2O4
produces 21.8 % solid ( ZnO )
21 mols of gas and 2 mols of ZnO or 8.7 %

The density of the final pyrotechnic mixture will
also factor significantly in the performance derived.

.

[Edited on 16-6-2013 by franklyn]
\


Are you sure Lithium Perchlorate decomposes to Li2O instead of LiCl ? , because wiki says LiCl. Which is 40% Solid.

I just checked some reported most did not state or stated LiCl.

Titanium perchlorate is some times dangerous with some types of liquid though.

[Edited on 16-6-2013 by DubaiAmateurRocketry]

[Edited on 16-6-2013 by DubaiAmateurRocketry]

[Edited on 16-6-2013 by DubaiAmateurRocketry]

AndersHoveland - 16-6-2013 at 08:31

Quote: Originally posted by DubaiAmateurRocketry  

Are you sure Lithium Perchlorate decomposes to Li2O instead of LiCl ? , because wiki says LiCl.

I would suspect a mixture of byproducts is produced upon initial reaction. That is the problem with some of these energetic propellants, the byproducts themselves can absorb energy or be slow to fully react and release all their energy. However, as the byproducts cool away from the initial reaction zone, the end product would indeed tend to be LiCl, though below a certain activation temperature Li2O, Cl2, and H2O can no longer combine, and I am sure some HCl manages to escape also.

The decomposition of some of these compounds is not always stoichiometric.

With magnesium perchlorate, I would expect much higher ratios of MgO to form, rather than the chloride.

DubaiAmateurRocketry - 16-6-2013 at 10:23

Quote: Originally posted by AndersHoveland  
Quote: Originally posted by DubaiAmateurRocketry  

Are you sure Lithium Perchlorate decomposes to Li2O instead of LiCl ? , because wiki says LiCl.

I would suspect a mixture of byproducts is produced upon initial reaction. That is the problem with some of these energetic propellants, the byproducts themselves can absorb energy or be slow to fully react and release all their energy. However, as the byproducts cool away from the initial reaction zone, the end product would indeed tend to be LiCl, though below a certain activation temperature Li2O, Cl2, and H2O can no longer combine, and I am sure some HCl manages to escape also.

The decomposition of some of these compounds is not always stoichiometric.

With magnesium perchlorate, I would expect much higher ratios of MgO to form, rather than the chloride.


Oh yes i do agree that LP does form some Li2O and decompositions are not really stoichiometric, but i was just trying to make sure the main product is LiCl, although some traces of Li2O may be found.

Fantasma4500 - 16-6-2013 at 10:43

this thread surely grew quickly haha...

98% is do-able infact, but i would only handle that in drop-amount treating it as liquid NI3, or in other words NCl3..
theres a video of anh. HClO4 reacting with paper tho
explodes on contact..

but i expected a much much higher price really..

octanitrocubane has a price that exceeds the price of gold per gramme before the 30% gold price drop just recently

now a 100g charge of that would be very expensive, despite that it exceeds 10000m/s VoD
in any possible way you wouldnt want this for rocket propellants anyways

with the price of this perchlorate, i suppose you want this as a catalyst more than main oxidizer..?

You are right — too late to edit now though

franklyn - 16-6-2013 at 17:25

I just assumed 4 LiClO4 => 2 Li2O + 7 O2 + 2 Cl2

It actually is LiClO4 => LiCl + 2 O2

produces 39.8 % solid ( LiCl )
2 mols of gas and 1 mols of LiCl or 33.3 %


The higher heat of formation determines the product

ΔfH solid LiCl = −409 kJ/mol

ΔfH solid Li2O = −300 kJ/mol


http://books.google.com/books?id=13sMAQAAIAAJ&pg=PA16&am...

Lithium: Theoretical Studies and Practical Applications
http://books.google.com/books/download/Lithium.pdf?id=13sMAQ...

.

virgilius1979 - 17-6-2013 at 09:06

What would be the result of reaction between Ti(ClO4)4 and H2O ?
Because I was thinking at the first time one could synthesize Ti(ClO4) by electrolysing concentrated HClO4 with Ti electrodes.

a bit off topic question: Did anyone produced some AgClO4 ? Do you also need anhydrous HClO4 ?

DraconicAcid - 17-6-2013 at 09:08

Quote: Originally posted by virgilius1979  
What would be the result of reaction between Ti(ClO4)4 and H2O ?
Because I was thinking at the first time one could synthesize Ti(ClO4) by electrolysing concentrated HClO4 with Ti electrodes.

Probably TiO2 and HClO4. Perchlorate is a very poor ligand, and titanium is very oxophilic.

DubaiAmateurRocketry - 17-6-2013 at 12:08

Quote: Originally posted by DraconicAcid  
Quote: Originally posted by virgilius1979  
What would be the result of reaction between Ti(ClO4)4 and H2O ?
Because I was thinking at the first time one could synthesize Ti(ClO4) by electrolysing concentrated HClO4 with Ti electrodes.

Probably TiO2 and HClO4. Perchlorate is a very poor ligand, and titanium is very oxophilic.


The link i gave already stated the best route to it.

TiCl4 and anhydrous HClO4, with few drops of Cl2O7, containing at least 2% of the solution, could help give a good yeild of TiCl4O16. Make more there are 4 HClO4 at least for each TiClO4, and add some extra would be better.

Reaction must been done under -10 degree, a bit lower is ok.

Heating of TiCl4O16 releases large amounts of gas, could lead to explosions, although the decomposition is endothermic.

AndersHoveland - 18-6-2013 at 04:25

For those who are not aware of the dangers, anhydrous HClO4 is not stable in storage. In fact, if left in a jar it gradually decomposes, increasing in temperature until there is a runaway decomposition and it finally explodes.
Quote:
Pure perchloric acid is unstable. If left standing at room temperature for between 10 and 30 days, at some point the temperature of the acid will spontaneously rise to 90 °C, at which point there will be an explosion.


72% perchloric acid is completely stable, however. It is just that the chemical properties of perchloric acid change markedly over 72% concentration. We can think of this concentration threshold as actually being hydroxonium perchlorate, H3OClO4.

[Edited on 18-6-2013 by AndersHoveland]

DubaiAmateurRocketry - 18-6-2013 at 05:32

Quote: Originally posted by AndersHoveland  
For those who are not aware of the dangers, anhydrous HClO4 is not stable in storage. In fact, if left in a jar it gradually decomposes, increasing in temperature until there is a runaway decomposition and it finally explodes.
Quote:
Pure perchloric acid is unstable. If left standing at room temperature for between 10 and 30 days, at some point the temperature of the acid will spontaneously rise to 90 °C, at which point there will be an explosion.


72% perchloric acid is completely stable, however. It is just that the chemical properties of perchloric acid change markedly over 72% concentration. We can think of this concentration threshold as actually being hydroxonium perchlorate, H3OClO4.

[Edited on 18-6-2013 by AndersHoveland]



Titanium perchlorate react with water, so it have to be prepared at temperature below -10 degree. -15 would guarantee no accidents, since that's how the Russian team performed and synthesized it

Can you distillate water from 72% HClO4 ?

[Edited on 18-6-2013 by DubaiAmateurRocketry]

woelen - 18-6-2013 at 06:18

Please use full words and not slang like 'u' where you should put 'you'.
It makes u l00k l1ke a k3wl.

Back to topic, you cannot distill off the water from 72% HClO4. This is the azeotrope of perchloric acid with water and when you boil this liquid, then water and acid escape in exactly a 28 to 72 ratio.
If you want to concentrate perchloric acid beyond 72%, then you need a strong drying agent to bind the water. Concentrated sulphuric acid in large excess amount does the job. P4O10 also does the job. I, however, STRONGLY advice against trying this. Distilling anhydrous perchloric acid is insanely dangerous and leads to violent explosions very easily. In professional labs this usually is not allowed at all, or if people need anhydrous perchloric acid, then they have to work in pairs, using blast shields and personal protection gear. You cannot store anhydrous perchloric acid. It explodes after appr. 1 month of storage if kept cool, it explodes earlier when kept at room temperature or when not entirely pure.
Acid at 72% concentration is the highest stuff which can be obtained commercially. Anything above 75% is dangerous to have around. It is really remarkable to see how much the properties of perchloric acid change with increasing concentration. At 70% it is actually amazingly inert and not even capable of oxidizing fairly strongly reducing stuff like Na2SO3 or KI, while anhydrous acid explodes on contact with nearly any reductor (e.g. dust, people, animals, plants :o).

@AndersHoveland: The 72% acid is the azeotrope, it is not hydronium perchlorate. The 72% acid does not have some well-defined stoichiometry, it just is a mix of water and acid. Hydronium perchlorate, H3O(+)/ClO4(-) is a solid and it contains appr. 85% by weight HClO4.



[Edited on 18-6-13 by woelen]

watson.fawkes - 18-6-2013 at 09:43

Quote: Originally posted by woelen  
In professional labs this usually is not allowed at all, or if people need anhydrous perchloric acid, then they have to work in pairs, using blast shields and personal protection gear.
Not only that, there's a specific class of fume hoods for working with perchloric acid. They have integral washdown systems on the interior walls.

Nitro-esteban - 20-6-2013 at 13:36

What would happen if titanium were reacted with 60% HClO4?

woelen - 20-6-2013 at 22:51

Titanium most likely does not react with 60% HClO4. Or, if it does, then I expect formation of purple titanium(III) ions in solution. On boiling down, the titanium is oxidized by air and hydrous TiO2 will separate from the liquid.

DubaiAmateurRocketry - 21-6-2013 at 02:12

Guys I purchased the paper, here it is :)

Page 1




[Edited on 21-6-2013 by DubaiAmateurRocketry]

DubaiAmateurRocketry - 21-6-2013 at 02:14








[Edited on 21-6-2013 by DubaiAmateurRocketry]

DubaiAmateurRocketry - 21-6-2013 at 02:16

I uploaded the report the Russian team posted that i bought, its in the comment before :)

[Edited on 21-6-2013 by DubaiAmateurRocketry]

caterpillar - 21-6-2013 at 02:27

Quote: Originally posted by woelen  
Anything above 75% is dangerous to have around. It is really remarkable to see how much the properties of perchloric acid change with increasing concentration. At 70% it is actually amazingly inert and not even capable of oxidizing fairly strongly reducing stuff like Na2SO3 or KI, while anhydrous acid explodes on contact with nearly any reductor (e.g. dust, people, animals, plants :o).
[Edited on 18-6-13 by woelen]


My dad was a chemist. He studied reactions like Fe++ -> Fe+++, Cu+ -> Cu++ and so on. Once a girl in his lab put some iron powder into HClO4, wanting to prepare Fe(ClO4)2. I do not know, which concentration this acid has- surely not above 70%, but explosion occurred just when this girl arrived to my dad to report that all is correct and reaction is going well...

DubaiAmateurRocketry - 21-6-2013 at 02:43

Quote: Originally posted by caterpillar  
Quote: Originally posted by woelen  
Anything above 75% is dangerous to have around. It is really remarkable to see how much the properties of perchloric acid change with increasing concentration. At 70% it is actually amazingly inert and not even capable of oxidizing fairly strongly reducing stuff like Na2SO3 or KI, while anhydrous acid explodes on contact with nearly any reductor (e.g. dust, people, animals, plants :o).
[Edited on 18-6-13 by woelen]


My dad was a chemist. He studied reactions like Fe++ -> Fe+++, Cu+ -> Cu++ and so on. Once a girl in his lab put some iron powder into HClO4, wanting to prepare Fe(ClO4)2. I do not know, which concentration this acid has- surely not above 70%, but explosion occurred just when this girl arrived to my dad to report that all is correct and reaction is going well...


Some perchlorates react with water.

woelen - 21-6-2013 at 06:12

Even with plain acids like 30% HCl you can get very violent reactions (nearly explosive), especially if the amounts are scaled up. You need to provide more info about the experiment, especially the scale at which it was performed.

I do not believe that adding a small spatula full of iron powder, added to e.g. 60% HClO4 will lead to an explosion. If, however, a tablespoon full of iron is added to e.g. 50 ml of 50% HClO4, then I can imagine that the reaction at first starts gently, but slowly it becomes more and more violent and finally you get a severe runaway. If you want an idea of what can happen, add some coarse granules of Al to 30% HCl. At first nothing happens, but after a few minutes the reaction becomes extremely violent. Only use test tube quantities for this experiment!
Here, simple acidity does the job already, I'm not talking about strongly oxidizing anion-reactions.

What I wanted to say with my previous post is that HClO4 is amazingly inert besides its acidity when the concentration remains below appr. 70%. It just reacts like a strong acid and not like a strong oxidizer such as HNO3 at similar concentration.

@DubaiAmateurRocketry: The fact that some perchlorates react with water does not mean anything in this context. These perchlorates would not even be formed under the conditions, discussed here. What you mean is that some metal salts easily are hydrolysed. This is not unique to perchlorate, it is unique to the metal-ion. Ti(ClO4)4 certainly cannot exist in water, it reacts to hydrous TiO2 and HClO4. But a similar thing is true for any Ti(IV) compound. E.g. TiCl4 reacts violently with water to form TiO2 and HCl and Ti(NO3)4 reacts to TiO2 and HNO3. There are many metal ions which show this behavior. Some of these ions hydrolyse completely in neutral solution (e.g. Ti(4+), Zr(4+), Sn(4+), Sb(3+), Bi(3+), Mn(3+)), others hydrolyse partly (e.g. Fe(3+), Al(3+), Cr(3+), Ce(4+), Sn(2+), Pb(2+)).



[Edited on 21-6-13 by woelen]

caterpillar - 21-6-2013 at 22:13

Unfortunately, I cannot add details above what I already told. My dad is dead and therefore I have no way to ask him, what really happened. One wise boy wanted to get dry (CH3)3PO4- ordinary trimetilphosfate, but he used Mg(ClO4)2. In this case at least it is clear, why this mixture exploded.

AndersHoveland - 21-6-2013 at 22:20

Quote: Originally posted by caterpillar  
My dad was a chemist. He studied reactions like Fe++ -> Fe+++, Cu+ -> Cu++ and so on. Once a girl in his lab put some iron powder into HClO4, wanting to prepare Fe(ClO4)2. I do not know, which concentration this acid has- surely not above 70%, but explosion occurred just when this girl arrived to my dad to report that all is correct and reaction is going well...

I think this probably had to do with the perchlorate not being very stable in the presence of Fe+2 ions while the concentrated solution was being heated. Once the activation energy was overcome, the Fe+2 was oxidized. So this particular problem would probably not be an issue with other salts. Although added caution never hurts.

With Ti(ClO4)4 hydrolyzing in water, I am not sure what the maximum concentration of perchloric acid that can result, but I suspect pure anhydrous HClO4 is not obtainable. Or in other words, that the hydrolysis is not completely stoichiometric. And similarly, I also doubt Ti(ClO4)4 would hydrolyze in ~90% perchloric acid, though perhaps an especially acidic adduct could form, H2Ti(ClO4)6.

I am sure that dropping a piece of Ti(ClO4)4 into water would be an extremely violent reaction, but whether it could actually explode, I am not sure. Because it is possible that the concentrated perchloric acid generated on hydrolysis could detonate from the intense heat of hydration. I am not sure what the minimum concentration is that can detonate (without any other reducing agent present), but from what I have read it seems to be around 85%.

[Edited on 22-6-2013 by AndersHoveland]

caterpillar - 21-6-2013 at 22:21

Quote: Originally posted by woelen  


What I wanted to say with my previous post is that HClO4 is amazingly inert besides its acidity when the concentration remains below appr. 70%. It just reacts like a strong acid and not like a strong oxidizer such as HNO3 at similar concentration.

[Edited on 21-6-13 by woelen]


You are right, talking about COLD HClO4. But hot acid (near its boiling point) is strong oxidizer- and even if reaction starts with cold acid, temperature may easily go up.

AndersHoveland - 21-6-2013 at 23:05

Quote: Originally posted by woelen  
Titanium most likely does not react with 60% HClO4.

If we look at other non-oxidizing acids, titanium tends to be very resistant to attack at dilute concentrations, but at higher acid concentrations or heated, the attack is more rapid. I think if the acid concentration is high enough, TiO2+ begins to dissolve in solution.