Sciencemadness Discussion Board

Perchlorate chemistry

12AX7 - 20-1-2008 at 08:45

From Technochemistry;
Quote:
Originally posted by chloric1
Quote:
Originally posted by DerAlte
Yes, not with...

Have no use for perchlorate myself. It's rather boring, apart from pyrotechnics. However, your efforts to produce the MMO, etc, Anodes are very impressive.
Der Alte


Oh man perchorates boring?! You let me down man.:( Ammonium perchlorate is easily converted to perchloric acid for which metal perchlorates can be made. I had a nice nickel deposit from a nickel perchlorate solution I electroyzed with no buffers or additives.

Don't forget potassium perchlorate's role in Self-Propagated High Temperatures reactions. This was mostly researched by the Russians so finding english papers takes a little sleuthing work.

Perchlorate occupies roles that chlorate is unable to do to the inherent dangers of instability.


How much aqueous chemistry is there for perchlorate? For instance, chlorate becomes chloric acid at very low pH (pKa ~ 1), which is unstable, making an active oxidizer, and disproportionating to ClO2 or Cl2 depending on conditions. But the perchlorate ion is much more stable, acting about like iodide, without the redox chemistry.

Does perchlorate serve in fusion oxidation type reactions? I suppose that's my interest, so I should get another pound of Cr2O3 and test if I can make K2Cr2O7 with KClO4.

Tim

DerAlte - 20-1-2008 at 12:47

Good idea to repost the above here, 12AX7. It was lost and almost OT where it was. Here it might stimulate some discussion. You pose the valid question - How much aqueous chemistry is there for perchlorate?

This is the valid point. All perchlorates (bar K,Rb,Cs) are very soluble (AFAIK) - hence no intersting separations are possible. It is useless for oxidations unless the kinetic activation barrier is overcome - not possible at aqueous solution temperatures, but done in pyrotechnical use. Not so for chlorates, as you say.

The very stability of the perchlorate ion and the high solubility make it boring, Chloric! Admittedly, as perchloric acid might be interesting to make; it is a very strong oxidant when anhydrous. Also those intrigued by Energetic Materials may find organic perchlorates interesting. but not for me as a 99% inorganic dabbler.

12AX7, I think there's a good chance you could make K2Cr2O7 the way you suggest, if you add KOH to 'mop up' the Cl- produced with K+ ion. After all it works with MnO2 or Mn2O3 to produce manganate this way, and manganate is far less stable than dichromate. At least you ought to get chromate, which is even more stable and fairly easy to turn into dichromate.. Give it a try if you've got Cr2O3.\

Regards,

Der Alte

BromicAcid - 20-1-2008 at 13:15

I've seen molten oxidations with chlorate and hydroxide but I am not sure if I have ever seen the use of perchlorate though I can figure no reason why such a reaction would not work, it would decompose all the same in the end.

Regarding the solubility of various perchlorates, sodium perchlorate is fairly soluble but potassium chlorate is not something I would call very soluble by a long shot (J.T. Baker states 1.5 g / 100 g H<sub>2</sub>O in their MSDS).

12AX7 - 20-1-2008 at 15:27

Fairly he says! Come to think of it, the perchlorates of Na and K may have the largest difference in solubility out of any pair of salts of Na and K with the same anion.

How about chemistry relating to its behavior as an ion? It's known to substitute for iodide, e.g. in the thyroid, inducing hypothyroidism, presumably because of a similar ionic size. What other potentially interesting properties does it have?

Tim

garage chemist - 20-1-2008 at 15:45

Lithium perchlorate is a very interesting compound. It is soluble in a variety of aprotic organic solvents, like THF and diethyl ether, those solutions conduct electricity and most likely deposit lithium at the cathode.
It is used in lithium batteries as an electrolyte salt for a reason.

If you can make HClO4 (either via the NaClO4/HCl method on dann2's website, or from NH4ClO4 and HNO3/HCl or by ion exchange from any soluble perchlorate) LiClO4 would be an interesting project. I want to try that myself when I have the time.
The most difficult part will be the drying though- anhydrous LiClO4 is extremely hygroscopic and binds water very tightly, you will need to heat it under vacuum to drive off the water of hydration from the trihydrate.

chloric1 - 20-1-2008 at 16:10

Quote:
Originally posted by garage chemist
If you can make HClO4 (either via the NaClO4/HCl method on dann2's website, or from NH4ClO4 and HNO3/HCl or by ion exchange from any soluble perchlorate) LiClO4 would be an interesting project. I want to try that myself when I have the time.
The most difficult part will be the drying though- anhydrous LiClO4 is extremely hygroscopic and binds water very tightly, you will need to heat it under vacuum to drive off the water of hydration from the trihydrate.


The HCl/HNO3 process for oxidizing the ammonium ion in ammonium perchlorate is straight forward. Just mix and boil, boil until no more HCl,Cl2,NOx evolves. It is not really an aqua regia process, because we are actually using nitrous acid as the oxidizer. So an equimolar mix of 1:1 HCl/HNO3 is of order here.


Lithium perchlorate should easily be prepared in a solid state reaction between lithium carbonate with ammonium perchlorate. Of coarse considering the lithium perchlorate can withstand the hydrolsis and sublimation temperature of ammonium carbonate which I am sure it can.

Perchlorate is a good inert spectator ion the produces VERY soluble salts which can be very usefull when solution with high metal values is needed. Besides, it is a heck of alot cheaper than trifluoromethanesulfonic acid or whatever.

garage chemist - 20-1-2008 at 16:46

Good point about NH4ClO4 and Li2CO3.
The only problem I can see is Li2CO3s low solubility (less than 1g/100ml), which even decreases with increasing temperature.
One would have to find out oneself how long it takes for Li2CO3 to dissolve in a boiling NH4ClO4 solution.

For the actual preparation of LiClO4 by this method one would use Li2CO3 in excess, boil until no more NH3 and CO2 are evolved, and filter hot (to minimize the solubility of Li2CO3).
Then concentrate and dehydrate the solution. Any Li2CO3 in the product will be left behind upon dissolution in ether.

(EDIT: I see you meant a dry reaction between NH4ClO4 and Li2CO3. I have never tried something like that, but I have produced KNO3 from AN and K2CO3 in solution myself many times and can say for sure that NH3 and CO2 are rapidly and completely expelled from a solution containing ammonium and carbonate ions upon boiling).

[Edited on 21-1-2008 by garage chemist]

chloric1 - 20-1-2008 at 16:56

Quote:
Originally posted by garage chemist

The only problem I can see is Li2CO3s low solubility (less than 1g/100ml), which even decreases with increasing temperature.
One would have to find out oneself how long it takes for Li2CO3 to dissolve in a boiling NH4ClO4 solution.


Ehh, not exactly what I meant. I said solid state for a reason. I meant physically mixing the powders and heating together. If it does not go right away, maybe a DROP or two of water can be added to the mix.

woelen - 20-1-2008 at 23:54

Quote:
Originally posted by Der Alte

Good idea to repost the above here, 12AX7. It was lost and almost OT where it was. Here it might stimulate some discussion. You pose the valid question - How much aqueous chemistry is there for perchlorate?

I have perchloric acid 60% and I did a few tests with it, already some time ago. I noticed that this acid, even at 60% concentration is remarkably inert. It only is a strong acid, but not a strong oxidizer.

I put in some solid NaI, in the 60% acid. No reaction at all, the solid dissolves.
Next, I carefully heated the liquid, with the NaI in it, until it started boiling/simmering a little Only after a few minutes of simmering, the solution turns pale yellow, noting more.

I have a similar experience with Na2SO3 added to the acid. By boiling, you drive off SO2 and NaClO4 remains in solution. There is no, or only very little, redox activity.

So, at concentrations up to at least 60%, the acid is remarkably inert and not that interesting. It might be interesting in situations, where one wants a strong acid, without coordinating and redox properties. Dilute HClO4 can be regarded as a nearly ideal acid, which is acid only and nothing else. All other acids I know have at least one other property besides their acidity, such as being easily reduced, being easily oxidized, its anion being able to coordinate strongly to metal ions.

I also have NaClO4.H2O. In aqueous solution, this is one of the dullest chemicals I have. It is even less reactive than NaCl or Na2SO4! Of course, at high temperature, in pyrotechnic mixtures, things become very different.

garage chemist - 21-1-2008 at 23:46

The 60% HClO4 can actually be a vigorous oxidizer under the right circumstances.
When you let a few drops of it soak into a charred piece of wood (like out of a wood fire) and light it, it burns very fiercely. The fire can be sustained by adding more HClO4. It works even better with 72% HClO4, I have been told.
The same can be done with vacuum-concentrated H2O2 (around 70%).

not_important - 22-1-2008 at 09:37

Quote:
Originally posted by garage chemist
Lithium perchlorate is a very interesting compound. It is soluble in a variety of aprotic organic solvents, like THF and diethyl ether, those solutions conduct electricity and most likely deposit lithium at the cathode.
It is used in lithium batteries as an electrolyte salt for a reason.

...


Then there is silver perchlorate, which also just about can be listed as "soluble". In water at 25 C 557 g dissolve in 100 cc of water, 5.3 g in 100 cc aniline, 26.4 g in pyridine, 5.3 in benzene, and 101 in toluene; also soluble in alcohols and polyols, chlorobenzene, acetonitrile, dichloromethane, ether, and nitromethane. The aromatic solvents tend to form solid complexes, "solvent of crystallisation", which generally are explosive which discourages study of them.

It's stable up to 486 C, beating out AgNO3 which is stable to 460 C. It's also one of the least light sensitive silver salts.

The_Davster - 22-1-2008 at 21:30

Working with silver perchlorate in anything but aqueous can be very frustrating as a result of how deliquescent it is. A huge problem if you are trying to maintain anhydrous or semi-anhydrous conditions. The stuff does not store nicely outside a dessicator.
Quote:
Originally posted by garage chemist
The 60% HClO4 can actually be a vigorous oxidizer under the right circumstances.
When you let a few drops of it soak into a charred piece of wood (like out of a wood fire) and light it, it burns very fiercely. The fire can be sustained by adding more HClO4. It works even better with 72% HClO4, I have been told.
The same can be done with vacuum-concentrated H2O2 (around 70%).

I have seen this demonstration done before, but have not gotten around to trying it myself. It surprises me and actually makes me nervous about my small bottle of 70%. It is one of the few chemicals for which I have 3 levels of containment, the bottle, a plastic bucket, then another plastic bucket.



[Edited on 22-1-2008 by The_Davster]

woelen - 22-1-2008 at 23:38

Quote:
Originally posted by garage chemist
The 60% HClO4 can actually be a vigorous oxidizer under the right circumstances.
When you let a few drops of it soak into a charred piece of wood (like out of a wood fire) and light it, it burns very fiercely. The fire can be sustained by adding more HClO4. It works even better with 72% HClO4, I have been told.
The same can be done with vacuum-concentrated H2O2 (around 70%).

That sounds like a cool experiment. This is something I definitely will try. I'll use a little wood stick from a match, which I will char by keeping it above a flame, but without really burning it. Once the little stick is soaked with the acid, I'll light it. At such a small scale, I do not expect any dangerous situations.

But again, this is reactivity of perchlorate at high temperatures, not at the liquid water temperature range.

JohnWW - 23-1-2008 at 05:13

That reminds me - now that the cation ClF6+, as the salt of SbF6-, has been prepared by reaction of ClF5 with [KrF][SbF6] (in which the cation decomposes to Kr and F+), I wonder if [ClF6][ClO4], containing two Cl(VII) atoms in the cation and anion, could be prepared metathetically. It would be an extremely powerful explosive, but reasonably safe to handle.

woelen - 23-6-2014 at 03:45

I came across a very nice article about perchloric acid. It is one of the few articles which gives real safety info and does not try to scare the shit out of its readers. Perchloric acid is demonized in many publications, while it only under very specific conditions is as dangerous as many people believe. Read this little article, it puts everything in perspective:

http://www.msdshazcom.com/MSDS/G/gfschemicals/perchloricacid...

[Edited on 23-6-14 by woelen]

Brain&Force - 23-6-2014 at 08:43

The fume hood I was working in had this warning right next to the safety procedures:
"DO NOT USE PERCHLORIC ACID IN THIS UNIT."
Not like I could have any, this was at school.

jock88 - 23-6-2014 at 14:06


Is this available to anyone?

http://books.google.ie/books/about/Perchloric_acid.html?id=G...

kmno4 - 23-6-2014 at 23:19

Quote: Originally posted by jock88  

Is this available to anyone?

http://books.google.ie/books/about/Perchloric_acid.html?id=G...

Yes, available for everyone and for free from HathiTrust Library :P