Sciencemadness Discussion Board - Perchloric acid preparation

Perchloric acid preparation

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plasma - 20-5-2002 at 18:43

Perchloric acid How can I make perchloric acid (HClO4) from hydrochloric acid (HCl) ?

I think the reaction would be like this:
HCl + H2SO4 -> HClO4 + 2H + S

Is this equation right (I don't really think so) ?

madscientist - 20-5-2002 at 18:43

Unfortunately, production of chlorine oxides is far more complex than that. It is possible to easily make perchloric acid from a perchlorate salt and sulfuric acid.

Coen - 20-5-2002 at 18:44

You should see that this reaction is thermodynamically not right already.

You can also make HClO4 using a chlorate + H2SO4 and then heating it in the air. Partially oxidized to HClO4 then.

plasma - 20-5-2002 at 18:44

What about this ?

36% 64% 32% 35% 33%
3NaOH + 3Cl2 -> NaClO3 + 2NaCl + 3HCl
warm

70% 30% 56% 44%
2NaClO3 + (COOH)2 -> 2NaHCO3 + 2ClO2
90-100 C

94% 6% 70% 24% 6%
4ClO2 + H2O -> 2HClO4 + Cl2 + O

Is there any easier way to get chlorine gas than by electrolysis ?

madscientist - 20-5-2002 at 18:45

Yes, there are much easier ways. Hydrochloric acid will react with various strong oxidizers to liberate chlorine gas. I recommend using either hydrogen peroxide, potassium permanganate, or manganese dioxide.

A good way to prepare sodium chlorate is the electrolysis of a hot aqueous solution of sodium chloride. Electrolysis of a cold solution of sodium chloride will yield sodium hypochlorite. Electrolysis of a cold solution of sodium chlorate will yield sodium perchlorate.

Chlorine gas will hydrolize to form hydrochloric acid and hypochlorous acid (HOCl).

Cl2 + H2O --> HCl + HOCl

Therefore, you could prepare an equamolar mixture of sodium chloride and sodium hypochlorite by bubbling chlorine gas through an aqueous solution of sodium hydroxide.

I believe the following reaction would occur:

2NaClO3 + (COOH)2 --> (COONa)2 + 2ClO2 + O2

DO NOT mix sodium chlorate with concentrated sulfuric acid, the reaction is explosive (just a general warning to any readers).

I believe the following reaction would occur:

2ClO2 + H2O --> HClO3 + HClO2

Both of those acids are unstable; I believe both of them have never been isolated in pure form.

8HClO2 --> 4H2O + 4ClO2 + 2Cl2O + O2
4HClO3 --> 2H2O + 4ClO2 + O2

plasma - 20-5-2002 at 18:45

Thank you for your response
I will have to disagree with you

; chlorine dioxide mixed with hot water decomposes to perchloric acid and gives if chlorine and oxygen like this.

4ClO2 + H20 -> 2HClO4 + Cl2 + O

I got the theory confirmed by this site http://www.xrefer.com/entry.jsp?xrefid=486546&secid=.-

madscientist - 20-5-2002 at 18:46

Intriguing, thanks for that link!

I still think this reaction would occur:

2ClO2 + H2O --> HClO3 + HClO2

But, only in cold water. Clearly you are correct that ClO2 reacts with hot water to form HClO4 while liberating chlorine and oxygen gas; makes sense now that I think about it. This is what I think is occuring (I'll try to do some research on this tonight):

2ClO2 + H2O --> HClO3 + HClO2

(when the solution of HClO3 and HClO2 is heated, the HClO2, which is an unstable, powerful oxidizing agent, decomposes, oxidizing the HClO3 to HClO4)
2HClO3 + 2HClO2 --> 2HClO4 + H2O + Cl2 + O2

Therefore bubbling ClO2 through hot water should yield HClO4 while liberating chlorine and oxygen gas.

I'm thinking this is the equation for the reaction of moist (COOH)2 and NaClO3 at fairly high temperatures (90C-100C):

2NaClO3 + (COOH)2 --> Na2CO3 + 2ClO2 + H2O + CO2

madscientist - 20-5-2002 at 18:46

Oh, by the way, you can produce chlorates by heating an aqueous solution of hypochlorites.

3NaOCl ---(heat)---> NaClO3 + 2NaCl

Therefore bubbling Cl2 through a hot aqueous solution of NaOH would be a working method of preparing NaClO3.

plasma - 20-5-2002 at 18:46

The hypochlorite method.
NaClO3 could be extracted by filtering, but I wonder; If 209g NaClO3 and 36g NaCl is mixed with 100ml water, does the NaCl solve at all ? If it doesn't, aquos sodiumhypochlorite could then be boiled to the salts formed. The right amount of water could then be added to the salts (NaClO3 & NaCl), NaCl would not solve and could easily be filtered out. Is any of this correct ?

Soulability :
NaClO3 -> 209g/100ml @ 20 C
NaCl -> 36g/100ml @ 20 C

madscientist - 20-5-2002 at 18:47

I really can't think of a simple solution to the inert-gas problem. Most of the time, the gas is going to get heated to autoignition temperature; however it will not combust if there is no oxygen present. Try boiling hexane and using the hexane vapors as an inert gas; higher autoignition temperature. I have produced CS2 which has an autoignition temperature of around 90C by heating sulfur and carbon in a pyrex flask at temperatures more like 500C; never have had any trouble. Mostly, just make sure you're using a flask, and do it outside (it case the gas being used does ignite). If it ignites while outside, it really doesn't matter.

I just looked up the solubility of sodium chlorate here: http://www.jtbaker.com/msds/S3314.htm according to them, 100g NaClO3 dissolves in 100g H2O at 20C.

If you placed 100g NaClO3 and 36g NaCl in 100mL of water, I believe about 73.5g NaClO3 and about 9.5g NaCl would be dissolved.

Therefore, the best method for preparing high-purity NaClO3 is the electrolysis method. However, if you just plan on using the NaClO3 for preparing ClO2, then high levels of contamination with NaCl is not to be worried about.

A little accident

plasma - 19-9-2002 at 01:21

I tried making perchloric acid with my "professional"

equipment.
I used a three necked flask, a 70 degree adapter and a allin condenser. I measured out the right amount of oxalic acid and put it in the threee- necked flask. I heated the oxalic acid to it melted (90 - 80 degrees celsius). Then i added portions of Sodium chlorate (wich was made from NaClO decomp. and was very polluted). A yellowish-green gas developed at a great speed. I bubbled the gas trough cold water, and after five min. the water had turned yellow. Just as I leaved the lab to get some food, a explosion occured. I heard a farly high report, but no damage was done.

What could have been the reason for the explosion, ClO2 explosion or to high HClO4 conc. ?

Thanks

vulture - 19-9-2002 at 04:33

Explosive decomposition of ClO2....

plasma - 19-9-2002 at 09:53

What is the decomposition temp. of ClO2

vulture - 19-9-2002 at 10:25

Chlorine oxides decompose violently in contact with organic matter. Since you used oxalic acid....

Forgot something

vulture - 19-9-2002 at 10:26

To add, there is no temperature listed at which it explodes but it is very unpredictable. Treat like gaseous NI3.NH3

plasma - 19-9-2002 at 10:58

I doubt it was the contact with oxalic acid that made the chlorine dioxide explode. This is where I got the facts

Maybe the collecting flask was a bit dirty ?

Psycho - 19-9-2002 at 16:57

Some Accidents Involving Perchloric Acid

1. Explosions may occur when 72% perchloric acid is used to determine chromium in steel, apparently due to the
formation of mixtures of perchloric acid vapor and hydrogen. These vapor mixtures can be exploded by the catalytic
action of steelparticles.2
2. Two workers are reported to have dried 11,000 samples of alkali-washed hydro-carbon gas with magnesium
Perchlorate over a period of 7 years without accident. However, one sample containing butyl fluoride caused a
purplediscoloration of the magnesium Perchlorate, with the subsequent explosion of the latter.2
3 . A worker using magnesium Perchlorate to dry argon reported an explosion andwarned that warming and
contact with oxidizable substances should be avoided.2
4. An explosion was reported when anhydrous magnesium Perchlorate used indrying unsaturated hydrocarbons was
heated to 220'C.2
5 . An explosive reaction takes place between perchloric acid and bismuth or certainof its alloys, especially during
electrolytic polishing.2,5
6. Several explosions reported as having occurred during the determination ofpotassium as the Perchlorate are
probably attributable to heating in the presenceof concentrated perchloric acid and traces of alcohol. An incident
in a French laboratory is typical: an experienced worker in the course of a separation ofsodium and potassium
removed a platinum crucible containing a few decigramsof material and continued the heating on a small gas flame.
An explosion pulverized the crucible, a piece of platinum entering the eye of the chemist.7
7. A violent explosion took place in an exhaust duct from a laboratory hood inwhich perchloric acid solution was being
fumed over a gas plate. It blew outwindows, bulged the exterior walls, lifted the roof, and extensively
damagedequipment and supplies. Some time prior to the explosion, the hood had beenused for the analysis of
miscellaneous materials. The explosion apparentlyoriginated in deposits of perchloric acid and organic material in the
hood andduct.8
8. A chemist was drying alcohol off a small anode over a Bunsen burner in a hoodreserved for tests involving
perchloric acid. An explosion tore the exhaust ductfrom the hood, bent a portion of the ductwork near the fan,
and blew out many panes of window glass.8
9. An employee dropped a 7-pound (3.2 kg) bottle of perchloric acid solution ona concrete floor. The liquid was
taken up with sawdust and placed in a covered, metal waste can. Four hours later, a light explosion blew open
the hinged cover of the can. A flash fire opened three sprinklers which promptly extinguished the fire.8
10. A 7-pound bottle of perchloric acid solution broke while an employee wasunpacking a case containing three bottles.
The spilled acid instantly set the wood floor on fire, but it was put out quickly with a soda-acid extinguisher.8
11. At a malleable iron foundry, perchloric acid had been used for about 4 years inthe laboratory for the determination
of the silicon contents of iron samples. Acast iron, wash-sink drain at the bench used for this purpose had
corroded and the leaking acid had soaked into the wood flooring, which was later ignited whilea lead joint was
being poured. This fire was extinguished and part of the woodflooring was removed. Later in the day, at a point
slightly removed from the location of the first fire, a similar fire occurred when hot lead was again spilled.This time
the fire flashed with explosive violence into the exhaust hood and stackabove the workbench. Laboratory
equipment and records were wetted downextensively and damaged.
12. A stone table of a fume hood was patched with a glycerin cement and severalyears later, when the hood was being
removed, the table exploded when a worker struck the stone with a chisel. The hood had been used for digestions
with perchloric acid and, presumably, acid spills had not been properly cleaned up.9
13. A conventional chemical hood normally used for other chemical reactions, including distillation and ashing of organic
materials, was also used during the same time for perchloric acid digestion. During a routine ashing procedure, the
hot gases went up the 12-inch tubular transit exhaust duct and one of a series of explosions occurred that tore the
duct apart at several angles and on the horizontal runs.9
14. During routine maintenance involving partial dismantling of the exhaust blower on a perchloric acid ventilating
system, a detonation followed a light blow with a hammer on a chisel held against the fan at or near the seal
between the rear cover plate and the fan casing. The intensity of the explosion was such that it was heard 4 miles
away. Of the three employees in the vicinity, one sustained face lacerations and slight eye injury; the second
suffered loss of four fingers on one hand and possible loss of sight in one eye; the third was fatally injured with the
6-inch chisel entering below his left nostril and embedded in the brain.9
15. A 6-pound (2.7 kilograms) bottle of perchloric acid broke and ran over a fairly large area of a wooden laboratory
floor. It was cleaned up, but some ran down over wooden joists. Several years later a bottle of sulfuric acid was
spilled in this same location and fire broke out immediately in the floor and the joists.9
16. A chemist reached for a body of perchloric acid stored on a window sill above a steam radiator. The bottle struck
the radiator, broke, and the acid flowed over the hot coils. Within a few minutes the wooden floor beneath the
radiator burst into flame.9
17. An explosion occurred when an attempt was made to destroy benzyl celluloses by boiling with perchloric acid.11
18. An explosion occurred as anhydrous perchloric acid was being prepared via sulfuric acid dehydration and extraction
with methylene chloride when a stopper was removed from the separatory flask.14
19. A rat carcass was dissolved in nitric acid, the fat skimmed off, and perchloric acid added. The mixture was
heated to dryness and touched, setting off an explosion that cracked the fume hood and nearly blew out the sash.16
20. A perchlorate-doped polyacetylene film was prepared and stored under argon in a seated vessel. Two weeks later,
the film detonated when the vessel top was being removed. Earlier safety testing failed to show any reaction to flame
or impact.17
21. Perchlorate-doped polyacetylene samples combusted violently in the oxygen atmosphere of a Schoniger flask.18
22. An explosion occurred in a fume hood upon ether drying of a second crop of crystals of hexaminechromium (111)
perchloratethatwere washed with absolute ethanol and anhydrous diethylether. Following aspiration of the ether
wash, the ether damp filter cake was agitated with a glass stirring rod and the mass detonated.19
23. Perchlorate-doped, highly conducting polythiophene Pt-CIO, exhibits excellent ambient stability, but the film should
not be heated above 1000C. Touching an extremely dry Pt-CIO, film (kept in a desiccator over P20,) with
tweezers might cause an explosion of the film.20
24. Some samples of rare earth organic fluoride were re-ashed with perchloric, sulfuric, and nitric acid in I -liter
beakers. One of the beakers started foaming, turned yellow, and then exploded. The surface of the hot plate
was bent downward, and the imprint of the beaker was left in the metal surface of the hot plate by the force of the
explosion.1
25. Drying an acetonitrile adduct of neodymium Perchlorate at SO'C in vacuumapparently produces a compound that
can detonate on mechanical contact

Any of those explain the explosion

Psycho - 19-9-2002 at 17:05

I hope this helps:

http://www.geocities.com/CapeCanaveral/Campus/5361/chlorate/...

rikkitikkitavi - 19-9-2002 at 22:05

still, ClO2 is the most widely used bleaching chemical in the pulp and paperindustry. In sweden alone several tens of thousands of tonnes of NaClO3 is manufacured everyyear for the sole purpose of making ClO2.

www.ekachemicals.com

so it can be handled safely in conctact with organic material, it is just that it is probably very difficult to do this in a lab.

/rickard

daryl - 21-9-2002 at 14:30

I was recently researching perchloric acid, it appears to be a bastard of a chemical. Not only does it dissolve experimenters, it can spill onto wooden floorboards and either catch fire right now or wait for a few months/years and form a coating that is friction sensitive and burn/explode later. It also seems to create merry hell with fume cupboards.

I would like to hear from a person who regularly uses perchloric acid in their normal work.

IodineForLunch - 1-10-2002 at 16:29

How about reacting a stoichemical amount of potassium perchlorate and conc. sulfuric acid, and distilling the resulting perchloric acid to tilt the equilibrium in the favorable direction? HClO4 boils at 39 degrees; concentrated H2SO4 boils at 290 degrees. Little bit of a difference there.

Although, I don't know if perchloric acid would enjoy being distilled...

David Hansen

BLAST_X - 1-10-2002 at 22:31

mmmmm,

ever heared somebody a big booom
by heating HClO4 on the open air or by
contact with draining chemicals?

Higher a concentration 50%, it`s a bad thing to cooking a soup with >75 C.

You can prepare >70% HClO4 by dropping up H2SO4 to KClO4 in a vacuum distillation-device and following all reakting, distil very careful with waterbad under vacuum !

vulture - 2-10-2002 at 08:00

The boiling point of perchloric acid is 120,5C.

rikkitikkitavi - 2-10-2002 at 08:01

HClO4 formes a aezotrope with water at 71 %, BP = 203 C.

Conentrated (> 90 %) can not be heated to its boiling point at atmospheric pressure, decompostion starts at 100 c or so.
decomposition starts with the acid changing colour to the redish , and very soon after that explodes.

If handling there must be a possibility to dump everything into water without a second of delay!!!!

Even at room temperature this can happen.

Less than 71 % it is generally considered safe, given standard precautions for a strong oxidizing acid forming explosive mixture with most organic materials..

It is a excellent reagence for wet-ashing filters , but needs much care in the process.

Industrially it is made by excess HCl and NaClO4 , NaCl preciptates and is filtered off, whereafter the HCl is removed by distillation. Left is NaClO4 dissolved in concentrated HClO4 wich is separated by vacuum distillation.

/rickard

BLAST_X - 3-10-2002 at 02:40

I`d read any time ago,
the bp of HClO4 is a this densitys:

1,53 g/cm3 =~ 160 C
1,68 g/cm3 => 198 C
?

Theoretic - 23-7-2003 at 07:32

I read that:
1)Chlorine dioxide dissolves in water without reacting unless
2)When the water is hot it disproportionates into chlorous and chloric acid or
3)On sunlight or UV the solution disproportionates into HCl and HClO4.
4)HClO2 decomposes into chlorate, chloride and ClO2
Also, only from a single (however hard have I tried to find out more) source I read of thic STRANGE reaction:
Cl2+NaNO3=>NaClO2+NOCl
I suggest it continues with:
5NaClO2+4HCl=>4ClO2+5NaCl+2H2O
8ClO2+4H2O(sunlight)=>3HCl+5HClO4

BromicAcid - 16-10-2003 at 11:51

I found this out today and it just makes me agitated that I didn't figure it out on my own earlier. I remember back in chemistry lab when we would take saturated solutions of sodium chloride, run a little HCl gas through them and precipiate out the sodium chloride with the common ion effect. Today I came across.

NaClO4(aq) + HCl(aq) -----> NaCl (s) + HClO4(aq)

I could post more details if someone wants but appartently after filtering off the NaCl and heating to about 100C a 50% solution of HClO4 is acheived and the HCl volitizes off, the resulting solution gives no reaction with AgNO3 sounds insanely simple to me.

Yes that works

chloric1 - 19-10-2003 at 06:07

A saturated NaClO4 solution with excess 35% HCl does work as indicated in my perchlorate book. But there may be trace NaClO4 in the perchloric acid and the book recommended vacuum distillation of the 72%HClO4 azetrope for purity. You do have to heat the acid to about 200C to drive off the excess HCl. Preferrably distill the 20%HCl azetrope for further sythesis to eleminate waste and recycle

[Edited on 10/19/2003 by chloric1]

[Edited on 10/19/2003 by chloric1]

Extraction of perchloric acid with DCM

Mephisto - 19-10-2003 at 07:40

This might be interesting:

Quote:

quoted from U.S. patent 3,981,975
Anhydrous perchloric acid has been prepared by extraction of a perchloric acid-oleum solution with methylene chloride. See the article by P. H. Plesch et al., "Chemistry and Industry," September, 1971, pp. 1043 and 1044.

Has someone here access to this article?

vulture - 19-10-2003 at 12:43

IIRC, that method was abandoned because the mixture of DCM and nearly anhydrous HClO4 was highly explosive.

the timeless - 3-11-2003 at 02:43

On the previous site of this topic there is the auto-oxidising -method of producing NaClO3:

3NaClO-> NaClO3+2NaCl

I have known this method for ages but now when I saw it I thought that I could ask if you know a good method of either removing the NaCl from the solution or extracting the NaClO3 from it. My methods are either dangerous or ineffective, for example causing the NaClO3 to decompose.
So, do you know any better ways?

[Edited on 3-11-2003 by the timeless]

BromicAcid - 3-11-2003 at 08:15

To seperate out the chlorate add some potassium chloride. The chlorate will precipitate out as the potassium salt. Solubility of potassium chlorate is 7.1g/100ml H2O and in boiling water it is 57g/100ml H2O. Compare to 79g/100ml cold H2O and 230g/100ml boling water for sodium chlorate. The potassium salt will crystalize out preferedly over the sodium chlorate the potassium chloride and the sodium chloride.

Even easier

PrimoPyro - 3-11-2003 at 10:34

If you actually want the NaClO3 as opposed to KClO3, or youdon't have KCl on hand, just add lots more NaCl to a clean solution of NaCl/NaClO3 and water, and fractionally crystallize out the chlorate.

The excess of chloride ions reduce the solubility of sodium chlorate. Add salt to a warm (not boiling) solution and quickly chill it in the freezer, filter the crystals, re-warm it, rechill it (no need to add more salt) and repeat. Don't add somuch salt that you create a supersaturated solution at cold temps, but it will dissolve at warm temps,or you will contaminate your NaClO3 with precipitated NaCl.

PrimoPyro

So, has anyone here made perchloric acid by the classical method?

BromicAcid - 17-3-2004 at 19:38

Right now I have 5 lbs of KClO4 just sitting around, it's very pure and I have absolutely no use for it, my pyro days ended about 3 years ago, but there is no way in hell I'm just going to throw it away (although I could always eBay it). So I was thinking about perchloric acid.

Well, perchloric acid from KClO4 and H2SO4 requires high vacuum right? And oil bath I would assume, something like 70% H2SO4, actually, I've got all the conditions on hand but I was wondering if anyone has actually attempted to make it this way or has succeded, I've heard that even by this method, with an oil bath and careful heating it still explodes often, also, how much vacuum do I need if I decide to make an attempt?

BromicAcid - 2-5-2004 at 10:11

I decided that this might help, I've scanned in some interesting information regarding different methods of production of perchloric acid at my web site:

http://members.aol.com/bromicacid/perchlorate/perchlorate.ht...

So if you want some good information head on over.

BromicAcid - 9-10-2004 at 13:19

I attempted the classical method of perchloric acid production today. All my glassware was presoaked in concentrated HNO3 for a period of about 2 weeks then washed with hot distilled water to ensure cleanliness. All chemicals used were pure.

I took a 14/20 distillation apparatus and filled it with water and dropped a stir bar into it, and closed it. I applied vacuum the suction of my vacuum pump and it boiled rapidly at room temperature and pulled the liquid over into the receiver. Excelent.

Knowing that I had a powerful vacuum that would allow me to distill my perchloric well below atmospheric pressure and the 210 C where it decomposes redily I charged the 100 ml reaction flask with 12 g of KClO4 and 25 ml of 33% H2SO4. A stirbar was added and the apparatus connected together. It was held in a bath of water saturated with CaCl2 which gives it a high boiling point and the vacuum connection lead into an erlemyer full of a saturated NaOH solution and finally to the vacuum pump. Stirring was begun and the vacuum turned on.

Bad Move.

As you just read, I was using an aqueous solution of NaOH for the trap, as soon as it got under vacuum it foamed and frothed and boiled into the vacuum pump. Hummm... bad idea. Anyway, I replaced the liquid in the trap with a test tube containing solid NaOH which I placed in the erlemyer and put the tube to lead through that. One problem fixed.

After vacuum was applied this second time heating was begun and the mix kind of sat there. But I found out that the pipe I had leading into the solid NaOH touched the bottom too throughly and it had cut off the vacuum to the flask, found out too late is more like it. Because when I had found out the flask was already at 90 C and when I moved the tube by pulling from above the vacuum went though the vessel and flash boiled the perchlorate powder though the vessel and defeated the purpose of everything.

So I cleaned everything out with distilled water again and recharged the flask. Then heating was begun again. About 60 C the perchlorate powder at the bottom started to chaotically float from top to bottom. About 70 some bubbles came to the top and about 73 I got some distillate coming over. The flask was being magnetically stirred but I blame my air conditioner compressor. It had fits of low vacuum and high vacuum and it would occasionally almost pull over the reaction mixture. I turned down the hot plate to keep the mixture at about 80 C and distillate kept coming over. There was no problems at all. Well, except that my receiving flask had a crack developing from the chipped lip that it had. But suddenly the vacuum shot up again and caused the mix to boil fiercely and pulled over my KClO4 solid into the collection flask.

Worrying about the crack developing in the flask I shut down everything. I'm going to clean up in about 10 minutes.

The point of all this being. HClO4 distillation from a reaction mixture of H2SO4/KClO4 seems quite easy if you have a reliable vacuum source. I really want to try this again in the future, there was no strange reactivity in the reaction flask, it seemed incredibly tame compared to distilling HNO3 with the red vapors and everything. Just be sure to keep it below 150 C which is where HClO4 starts to decompose noticeably, use clean good reagents, and no organics at all! Soaking your glassware in 70% HNO3 seems like a good idea. Distilling anhydrous perchloric calls for greasing your glass joints with a KClO4/95% H2SO4 mixture but for the dihydrate (the azeotrope of about 70% HClO4) that distills over here I had no noticeable reactivity at all with teflon tape.

budullewraagh - 9-10-2004 at 13:44

chlorine peroxide will explode on contact with light. do not make it unless you have a death wish. if you breathe it in you will get pulmonary edema, also

BromicAcid - 9-10-2004 at 14:35

Where does Cl2O2 come into this?

The purpose of my post was to illustrate the relative safety of distilling HClO4 under reduced pressure. Significant decomposition only occurs above 150 C and I was distilling about 80 C which is well below the boiling point. People have distilled at atmospheric 210 C and lived to tell about it, it's just that decomposition occurs in upwards of 15% of the substance and there is a chance of explosion, usually chlorine dioxide is generated in this case, never heard of chlorine peroxide being formed.

However, what worried me more then the temperature or awkward byproducts was the inclusion of small pieces of organic matter in my reaction system, that would have probably given me a somewhat bad day.

vulture - 10-10-2004 at 07:47

Quote:

I had no noticeable reactivity at all with teflon tape

Why should teflon not be compatible with perchloric acid?

BromicAcid - 10-10-2004 at 08:12

Although logically I didn't think it, there was a part of me that said perchloric acid would react with almost anything. Therefore it wasn't a surprize, just a sort of relief when the teflon tape didn't start bubbling and turning black, although I was worried about impurities.

Theoretic - 11-10-2004 at 06:58

You could react ammonium perchlorate with nitrous acid to destroy the ammonium ion and leave perchloric acid behind.
Although the rising acidity may start to force the nitrous acid out of solution, this can be solved by slowly distilling N2O3 into the solution, which reversibly dissolves to form HNO2, which then reacts, as much is made as can be dissolved. In this way, I think a 100% yield can be achieved
The nitric/hydrochloric acid procedure in the PDF seems very nice, especially as you use 34 moles of NH4ClO4 and get 36 moles of HClO4 (with 99% yield, so still more than "possible"

.
Another procedure is to take NH4ClO4 (the poor thing

) and bubble chlorine through it. This will make HClO4, HCl and NCl3. On the bottom of the vessel there should be a catalyst that decomposes NCl3 preferably rapidly, but non-explosively. This will make chlorine, which will react with more NH4ClO4, and N2. When you're done with bubbling chlorine and decomposing NCl3, boil to remove the HCl. Though hazardous, I think this is a procedure that's at least as good as the others.

budullewraagh - 11-10-2004 at 07:15

Quote:

Where does Cl2O2 come into this?

Cl2O2 has nothing to do with this, but chlorine peroxide (ClO2) has everything to do with this, as it is incredibly unstable

BromicAcid - 11-10-2004 at 07:25

ClO2 is chlorine dioxide, Cl2O2 is chlorine peroxide.

Theoretic - 12-10-2004 at 06:47

Also, aqua regia could be reacted with NH4ClO4 (this one is different from the oxidation mentioned in the PDF). The hydrochloric acid actually serves as a catalyst:

NH4ClO4 <=> NH4+ + ClO4-
HNO3 + 3HCl <=> Cl2 + NOCl + 2H2O
NH4+ + Cl2 <=> NH2Cl + HCl + H+
H+ + ClO4 <=> HClO4
NOCl <=> NO+ + Cl-
NH2Cl + NO+ <=> NONHCl + H+
NONHCl <=>NONCl- + H+
NONCl- => N2O + Cl-
2H+ + 2Cl- <=> 2HCl

Thus all your hydrochloric returns to you, so it functions as a catalyst.

Perchloric acid preparation

rocketscience - 21-11-2004 at 13:20

I was wondering how I could hclo4 from naclo4 with out using hcl

If h2so4 works how could I react it safely with no explosions

and if it forms nahso4 will that decompose with heat to make a sulfur oxside

darkflame89 - 22-11-2004 at 00:55

Is it possible that chlorine dioxide be contained in a gel? You see, the supermarkets near my place sell a type of stuff that removes odour and dehumidifies the place. I have container of such stuff at home. It claims to contain ClO2, and besides the gel appears green in colour.

Is this possible? Isn't ClO2 dangerous?

Reverend Necroticus Rex - 22-11-2004 at 18:37

ClO2 is very unstable, and in concentration, explodes on contact with light.

You might have a hypochlorite perhaps? I have never smelled ClO2, but hypochlorites smell utterly foul, like chlorine bleach.

garage chemist - 23-7-2005 at 14:35

I just came back from my first attempt to distill dilute perchloric acid. A scary experience.

In the boiling flask, there were placed 14g finely ground KClO4 and 5,6ml 96% H2SO4 diluted with 10ml water (more water than needed to make azeotropic HClO4). The joints of the apparatus were greased with conc. H2SO4.
It was distilled at atmospheric pressure first. About 7ml of distillate came over at 100°C, which was assumed to be water and was discarded. Then the temperature went up a lot, to about 180°C, where white fumes were produced in the distillation flask. However, I saw that noticeable evolution of incondensable gas started (presumably oxygen from decomposition of the HClO4), so I let it cool and connected my aspirator to the apparatus.
I turned on the water for the aspirator and heated the mixture slightly, more white fumes were produced which condensed nicely into the receiver. Then I took a look at the condenser- and saw a deposit of white material which could not possibly be KClO4 because the solution didnt splash.

I remembered that HClO4 forms a crystalline monohydrate. "If the condenser contains the monohydrate, what is the boiling flask going to contain?" I asked myself.

I pulled away the burner at once and backed away from the apparatus to let it cool down. I disconnected the aspirator hose very carefully and turned off the water. After cooling down, I very carefully dismantled the apparatus, wearing heavy gloves and face protection.
The liquid which wetted the stillhead from the inside fumed strongly in air, even though it was at room temperature! 72% HClO4 doesn't fume in air.

I let the apparatus air out until it didn't fume any more and thoroughly washed it.

However, I don't know what to do with the fuming boiling flask which contains anhydrous HClO4 with NaHSO4. Anhydrous HClO4 explodes from the heat of dilution when water is added.
I'm letting it stand open overnight, in hope that air humidity will dilute it safely until tomorrow.

I have no idea why the monohydrate and then anhydrous HClO4 was formed. Maybe the NaHSO4 retained the water?
Perhaps I should use an excess of KClO4?

BTW, the 3ml of distillate which came over in vacuum was HClO4 of quite a high concentration. It was not oily like the 72% HClO4, but it precipitated large amounts of KClO4 when a few drops of it were added to dilute KCl solution.

Now, what should I do? I want to make HClO4 from KClO4 without having to deal with the anhydrous acid. Are there really no procedures for this?

neutrino - 23-7-2005 at 16:03

Would Bromic's method (upthread) be feasible here?

Lambda - 23-7-2005 at 16:10

Garage chemist, "Brauer" has a good prep. for HClO4. Maybe by adding Anhydrouse Calcium chloride (CaCl2) you will be able to dilute your distilation flask content, for it will attract moisture from the air. You then can easily sepperate CaSO4 and solution of HCl, CaCl2 and Ca(ClO4)2. By adding KCl solution to the glasswool filtrate, you will be able to recover KClO4 by precipitation.

BromicAcid - 23-7-2005 at 17:17

There is always a steam distillation, that would give your perchloric acid at a lower temperature and with nearly zero risk of what just happened to you. However sulfuric acid will come over too though considerably less then the amount of perchloric that will come over. So a second distillation would be necessary, not to mention that your perchloric will be dilluted.

As for what happened with the perchloric monohydrate, shouldn't it have been a liquid considering the temperature of the inside walls of the glass? Strange that you got anything like that considering the hygroscopicity of the monohydrate and the extreme difficulty in obtaining the anhydrous form or anything greater then the monohydrate (involving distilling the dihydrate with 6x weight of 99%+ H2SO4

. Very strange indeed.

garage chemist - 24-7-2005 at 03:15

Yes, it's strange, but I was kinda expecting something like that. KClO4 is too insuluble to react fast enough, and I'm just boiling away the water of the H2SO4 first and then I have conc. H2SO4 with KClO4 and I get HClO4 of a dangerously high concentration which will clog the condenser as the crystalline monohydrate.

I am aware of the procedure in "Brauer" and I take every measurement to PREVENT that the distillation goes like described there. It yields the anhydrous HClO4! The receiver needs to be cooled to -40°C in order to minimize the explosion risk.

Ammonium Perchlorate

chloric1 - 24-7-2005 at 11:02

Several years ago I made about 100 to 150ml of Aqua regia and then added ammonium perchlorate. You get a soup of nitrous acid, oxides, and chlorine. The solution is black cherry colored and I let it set undesturbed for the night to eliminate risk for explosions. After the solution is clear or off yellow, it is safe to distill off the excess nitrous-hydrochloric-chlorine solution and the 72 % constant boiling acid is in the flask. The file I am attaching should elighten you all to the theories on the mechanism for this reaction. Mind you, my explorations where 7 years before I obtained this file so the mounts of nitric and hydrochloric acid I used where quite excessive. This proceedure would definately be a cheaper as far as reagents are concerned.

Attachment: The preparation of Perchloric acid.pdf (431kB)
This file has been downloaded 1418 times

garage chemist - 24-7-2005 at 13:55

Thanks for the instructions, but I have no access to ammonium perchlorate.
I want to use KClO4 as the starting material!

Perchloric Acid By Electrolysis - Patent

jpsmith123 - 27-8-2005 at 05:21

Just in case anyone's interested, here's a patent I found regarding the production of HClO4 by electrolytic oxidation of HCl.

Regards,
Joe

Attachment: 1271633.pdf (294kB)
This file has been downloaded 1154 times

Taaie-Neuskoek - 27-8-2005 at 07:42

If one has BaCl2 ánd a working chlorate cell, it is pretty staight forward:

BaCl2 electrolyses --> Ba(ClO4)2, Ba(ClO4) + dil . H2SO4 --> BaSO4 (s) + HClO4.

From Mg(ClO4)2, which have on hand: Mg(ClO4)2 + Ba(OH)2 --> Ba(ClO4)2 + Mg(OH)2(s)
Via bariumperchlorate it looks rather easy, but the high toxicity of barium is a point to consider.

12AX7 - 27-8-2005 at 10:17

Explain this to me, you are writing in a perchloric acid thread that you are concerned about barium toxicity? I'm missing something here! LOL

Tim

Taaie-Neuskoek - 27-8-2005 at 13:55

You're left with a waste of Ba, which is toxic, even though the sulphate make it a lot safer due to it's insolubility.
Anyway, you're handling all the way very soluble barium salt's, which you have to filter, purify, etc... BaCl2 is considered as T by Merck, I don't know about the perchlorate. I do have a bottle of HClO4 60%, so I am not that interested in making a lot of the stuff, but if I can find some time and can make some Ba(OH)2 I might give it a shot to make a few ml.
Furthermore did I read that people were vacuum distilling perchloric acid, which is not advisable according to this site:
http://www-safety.deas.harvard.edu/advise/PerchloricAcid.htm...
I quote:

Quote:

Do not distill perchloric acid in a vacuum, because the unstable anhydride may be formed and cause a spontaneous explosion. Protect vacuum sources from perchloric acid/perchlorate contamination. Vacuum pumps should be thoroughly flushed and refilled with Kel-F or Fluorolube.

[Edited on 28-8-2005 by Taaie-Neuskoek]

garage chemist - 3-9-2005 at 11:40

Now I have a very good perchlorate cell running with a platinum anode (0,5mm wire, 25cm long, coiled up, was really expensive but it was well worth it!) and store- bought 45% sodium chlorate solution (liquid weedkiller from France, sold in 5L canisters) with added dichromate. It produces a sodium perchlorate solution of about 520g/L concentration, with minor impurities like NaCl, Na2SO4 and Na2SO3 from the destruction of residual chlorate with bisulfite.

I have successfully produced KClO4 (very good yield, over 90%) and NH4ClO4 (acceptable yield, ca. 60%) from this cell liquor.

This cell provides me with a never- ending supply of large amounts of KClO4 and NH4ClO4. The platinum wire was a very good investment. I live near to France and can get the necessary chloride- free (very important! Don't use homemade NaClO3 solution!) NaClO3 solution easily.

Today I tried the production of HClO4 directly from NaClO4 and HCl by those instructions: http://www.geocities.com/CapeCanaveral/Campus/5361/chlorate/...

I used 50ml NaClO4 solution and 125ml HCl. I used a porous glass suction filter to filter the NaCl, which worked beautifully.

It went as described. I concentrated the HClO4 in a still and used a scrubber to get rid of the HCl offgas. When the volume was at about 25ml, the temperature shot up to 209°C and the liquid became yellow. I had concentrated HClO4 at this moment.
But on cooling, an enormous amount of NaClO4 crystallized out. The crystals filled the entire liquid. Apparently my 520g/L NaClO4 solution wasn't concentrated enough and too much water was left in the solution, which had the effect of higher NaCl solubility.

I vacuum distilled it with a boiling capillary and got 17ml oily liquid looking exactly like conc. H2SO4, being maybe a bit more viscous. I measured the density as 1,68 and was very pleased when I looked up the literature value for 70% HClO4.

This procedure is a good pathway to HClO4, but one needs a really concentrated NaClO4 solution and also HCl of really 37% concentration, otherwise vacuum distillation of the crude HClO4 is imperative because of the enormous amounts of crystallizing NaClO4 (this also lowers the yield). Maybe one should boil down the NaClO4 solution completely and add that to the HCl in order to have a higher HCl concentration in the liquid and hence a lower solubility of NaCl.

Also, having to use 125ml of conc. HCl in order to get 17ml HClO4 is rather uneconomic.

A possibility would also be to boil down my NaClO4 solution and vacuum distill that with dilute H2SO4. The high solubility of NaClO4 should effect a complete reaction.

BromicAcid - 3-9-2005 at 11:52

Excelent to finally hear some sucess in this thread, maybe next time consider bubbling HCl gas into your solution of hydrochloric acid and sodium perchlorate to help eek out some more of that sodium as sodium chloride.

garage chemist - 3-9-2005 at 13:06

That's a good idea, and would perhaps allow for less conc. HCl being used and therefore save time at the evaporation step.

However, if I succeed in directly distilling HClO4 from NaClO4 and dilute H2SO4, then I'm done with the research on HClO4 production since this would be the most economic and cost- effective way.

jpsmith123 - 4-9-2005 at 09:47

I'm glad to hear of your success, garage chemist.

I'm wondering, what current are you running in your 0.5 mm platinum wire?

I have two pieces of pt wire; one is rather thick, maybe about 1 mm diameter and about 60 cm long, while the other is 0.5 mm x 30 cm, similar to yours. If my PbO2 electrode effort ultimately fails, I'm just going to end up doing a two stage process using some of this pt wire for the perchlorate part.

(BTW there is a guy selling 0.5 mm x 30 mm platinum wire on ebay for $34.00 USD http://cgi.ebay.com/PLATINUM-WIRE-12-inches-99-9-PURE_W0QQit...).

Quote:

Originally posted by garage chemist
Now I have a very good perchlorate cell running with a platinum anode (0,5mm wire, 25cm long, coiled up, was really expensive but it was well worth it!) and store- bought 45% sodium chlorate solution (liquid weedkiller from France, sold in 5L canisters) with added dichromate. It produces a sodium perchlorate solution of about 520g/L concentration, with minor impurities like NaCl, Na2SO4 and Na2SO3 from the destruction of residual chlorate with bisulfite...

garage chemist - 4-9-2005 at 10:27

My Pt anode can be used with a current of up to 3 Ampere, though I'm using 2 Ampere to limit heating of the cell (the efficiency drops sharply at over 40°C). Cathode is a coil of thick copper wire. The cell volume is 250ml, and it is standing in a tub of cooling water.
It is stirred magnetically. Mixing of the electrolyte is very important in perchlorate cells, and the mixing from the hydrogen at the cathode is NOT sufficient.
With 2 amperes and a 45% NaClO3 solution and 1g dichromate, the cell is done in 48 hours of non- stop running. The cell gasses start smelling like ozone in the last few hours strongly, this indicates that mostly water is electrolyzed and a lot of oxygen is evolving at the anode.
The used NaClO3 solution must be chloride- free (test it with AgNO3 solution!), this is extremely important and failure to use absolutely chloride- free solutions will result in severe corrosion of the Pt anode. Tap water already contains too much chloride.

I once runned the cell without dichromate, and the Pt got etched noticeably because the nascent hydrogen at the cathode reduced a small portion of the chlorate to chloride. Dichromate prevents this by forming a diaphragm of hydrated Cr2O3 around the cathode. Also, dont make the cathode of too large a surface.

Running a perchlorate cell is considerably more difficult than a chlorate cell. A lot of factors are important.
However, pH doesn't need to be adjusted.

The done cell liquor is strongly acidified with HCl and a spatula of Na2S2O5 is added. This turns the liquid green by reduction of the dichromate. If on boiling the solution still smells like SO2, the chlorate is also destroyed. If necessary, add a bit more Na2S2O5 until it smells like SO2.

Then basify with NaOH solution, this will precipitate Cr2O3 and the above solution is clear and pure. It contains 520g/L NaClO4 and no chlorate.

jpsmith123 - 4-9-2005 at 13:00

Thanks for the info.

I think I've read that persulfate works as well as dichromate to prevent reduction of chlorate, do you know anything about that?

Anyway, I'm wondering how practical it would be to produce NaClO4 by thermal decomposition of NaClO3, rather than by electrolysis?

If NaClO3 can be bought rather easily and cheaply, maybe it would be somewhat less troublesome to make NaClO4 by heating it and then purifying the NaClO4 by recrystallization?

If I remember correctly, NaClO3 decomposes at about 300 degrees C, whereas NaClO4 decomposes around 500 degrees C, so all you'd have to do is hold the temperature somewhere between, say, 300 to 400 degrees for a little while.

If I've done the math correctly, 100 grams of NaClO3 would provide a theoretical yield of about 86 grams of NaClO4, which should be fairly easy to seperate because of the large difference in solubilities of the the products.

What do you think about that?

garage chemist - 4-9-2005 at 14:06

Look at this thread:
https://sciencemadness.org/talk/viewthread.php?tid=4077

The method is very time -and especially propane- consuming. Also quite some NaClO3 is left undecomposed, which has to be destroyed first with strong HCl, then with Na2S2O5, which would give a very impure solution due to the large amount of chlorate to be destroyed.
But it works beautifully for preparation of KClO4.

Getting the right temperature (exactly 400°C) is difficult without an electronic thermometer and requires a lot of experience.

JoPann - 10-9-2005 at 12:37

hi

Iv'e just made some HClO4 in dissolving ClO2 into water (but I used a mix of ClO2 and CO2, which is safer).

After this step I put the mix of acids (HCl, HClO2, HClO3, HClO4) out into the sunlight for 2 days, to let the HClO2 and HClO3 reakt to HCl and HClO4.

Then I enriched the HClO4 in distilling off the HCl under vacuum, when the temperature began to rise i stopped heating, because i don't want anhydrous perchloric acid.

I tested the final produkt with AgNO3 for Cl^-, and some was laft, but not much.

Then i neutralized it and calculated a concentration of 0,8mol/l (enough for me)

does anybody have suggestions for this kind of procedure?

king regards

[Edited on 10-9-2005 by JoPann]

[Edited on 11-9-2005 by JoPann]

concentartion density table

SAM4CH - 6-3-2006 at 11:22

I need a table of relations of concentration of perchloric acid and its Specific Gravity "like any aqueous solution"!!?

[Edited on 6-3-2006 by SAM4CH]

mick - 6-3-2006 at 13:44

A quick coment,
As an organic chemist by trade
Organic perchlorates, organic azides and organic nitrates should be treated with respect.

Take care
mick

SAM4CH - 14-3-2006 at 03:21

Can I reach 90-95% yield in preparing Ammonium perchlorate from concentrated solution of sodium perchlorate "250g/L" at 70-80º C by using ammonium chloride, I read the yield will not be over 70%, but can we use an precipitating agent "e.g. commen ion, organic solvent,...etc" which can ppt. NaCl and can not precipitate AP!!!!!!!!!!?

[Edited on 14-3-2006 by SAM4CH]

Madandcrazy - 15-3-2006 at 08:38

What will you prepared from,
Ammonim perchlorate from perchloric acid or
perchloric acid from ammonim perchlorate ?

Maybe perchloric acid from sodium perchlorate and HCl
and than ammonim perchlorate by NH4OH (NH3)
HCl and sodium perchlorate, by eletrolysis

, or by
perchloric acid and sodium perchlorate.

Is the ammonium perchlorate staple and prepared by electrolysis ?

[Edited on 15-3-2006 by Madandcrazy]

SAM4CH - 15-3-2006 at 10:35

I am searching for the best method which can give largest yield, I have sodium perchlorate from electrolysis of NaCl and then I like to convert it to pure ammonium perchlorate, my dream is having yield more than 85%.

Ammonium perchlorate from ammonium chloride by electrolysis is not in my planning because it is very dangerous method which can produce unstable material "ammonium chlorate".!!?

praseodym - 16-3-2006 at 02:58

The reason you gave for not using the electrolysis of ammonium chloride is that there is a possibility of the production of ammonium chlorate which you think is too unstable. However, isnt ammonium perchlorate itself a rather unstable compound since all perchlorates are powerful oxidising agents? Furthermore, the ammonium chlorate can be converted to ammonium perchlorate simply by heating.

woelen - 16-3-2006 at 05:26

Electrolysis of NH4Cl definitely does not give NH4ClO3. The Cl2, formed at the anode, mainly forms NCl3 with ammonium ion. NCl3 is an exceedingly dangerous chemical. It is very unstable and looking bad at it may be sufficient to let it explode

.

Suppose you had some ammonium chlorate, then heating of this definitely causes the material to violently decompose or explode.

The ammonium ion cannot be simply thought of as an alternative for potassium ion or sodium ion in these cases. The ammonium ion is too good a reductor to be compatible with Cl2 and ClO3(-).

So, summarizing: NH4Cl is useless for making chlorate and it is useless for making perchlorate.

JohnWW - 16-3-2006 at 06:33

Woelen is correct. The Cl produced at the anode would initially form chloro-substituted ammonium salts of chloramines (NH2Cl, NHCl2), which are less basic than NH3, with HCl as byproduct, and finally NCl3, which is dangerously explosive without warning. The French chemist Dulong who discovered it in the 19th century was seriously injured by it.

It would be safer to try electrolysis of a solution of an alkali metal chloride. This results in the formation of a strongly alkaline solution of hypochlorite (being the industrial process for household bleach), and further electrolysis of this, with the liberated Cl2 being entrained in solution, would result in chlorate and perchlorate.

YT2095 - 27-7-2006 at 00:02

does anyone know the solubility of Ammonium perchlorate as compared to potassium chloride?
I have the KCl listed but not the NH4ClO4.

in a wet reaction I have NH4Cl and KClO4 dissolved together (only mg quantities), and there are crystals forming already as the soln evaporates slowly.

My guess is that the NH4ClO4 is least soluble, as the crystals (although very tiny) don`t resemble KCl.
can anyone confirm this?

woelen - 27-7-2006 at 02:14

I don't have figures for solubility of NH4ClO4, but as far as I remember, this stuff is MUCH more soluble than KClO4. So, in your solution of KClO4 and NH4Cl, I think that you just again get crystals of KClO4.

Once you have KClO4, it is very hard to make any other perchlorate of it. KClO4 is one of the least soluble salts of perchlorate. For this reason I have purchased NaClO4 as my primary perchlorate salt, that leaves me the possibility to make any other perchlorate salt. Finally, when the perchlorate is not used anymore in aqueous solution, then I add KCl just to precipitate the waste perchlorate and recrystallizing the KClO4 gives my perchlorate a second life, but this time not anymore for aqueous chemistry.

YT2095 - 27-7-2006 at 02:29

so the the Cl won`t pull the K then?
leaving my ammonium perchlorate, as it would if it was a Sodium salt.

Damn! :mad:

thnx anyway

Engager - 26-8-2006 at 17:32

Solubility in g/100ml - y axis, temperature in celsius - x axis. For NH4ClO4:

[img]http://www.ruspyro.net/Chlorates/NH4ClO4(Sol).bmp[/img]

For KClO4:

[img]http://www.ruspyro.net/Chlorates/KClO4(Sol).bmp[/img]

[Edited on 27-8-2006 by Engager]

hashashan - 6-2-2008 at 04:28

Guys, what will happen if one attempts to make HClO4 out of KClO4 and H2SO4 distillation
suppose, we boil KClO4 and H2SO4(diluted of course) in order to make some K2SO4 and HClO4, then we filter all the solids(K2SO4, and unreacted KClO4) and then distill WITHOUT vacum at 203 degrees.... what will happen?
will the HClO4 decopmose? will the anhydrous acid come out .. what will happen?

BTW I can filter only with glass filters right?

Fleaker - 6-2-2008 at 10:45

Yes only glass, keep away from any organics whether you're dealing with dilute or concentrated. It is important to note that while the conc. (90% +) acid can cause an instant explosion, dilute solutions carry the risk of forming unstable compounds that make explode later on (sometimes years) due to friction, shock, or heat.

I would not attempt a distillation at its boiling point...that is asking for an explosion. If you're after the anhydrous acid, I'm sure that Brauer has a safe, effective Preparation for it. If you don't have Brauer's excellent text, it is available in the forum library.

hashashan - 6-2-2008 at 11:41

Dude, i dont need the anhydrous ... I am after the azetropic acid.
I was told that It is possible to distill it at 203 without any problems .. I just dont understand why everyone does that under vacum

Formatik - 21-8-2008 at 20:11

Quote:

Originally posted by chloric1
Several years ago I made about 100 to 150ml of Aqua regia and then added ammonium perchlorate. You get a soup of nitrous acid, oxides, and chlorine. The solution is black cherry colored and I let it set undesturbed for the night to eliminate risk for explosions. After the solution is clear or off yellow, it is safe to distill off the excess nitrous-hydrochloric-chlorine solution and the 72 % constant boiling acid is in the flask. The file I am attaching should elighten you all to the theories on the mechanism for this reaction. Mind you, my explorations where 7 years before I obtained this file so the mounts of nitric and hydrochloric acid I used where quite excessive. This proceedure would definately be a cheaper as far as reagents are concerned.

I've used a method similar to this described in Gmelin. They say the yield is 99.7 to 99.8%. I adjusted it to use subconcentrated HNO3 (Gmelin says to use 68 to 70% HNO3) and scaled it down by 10 times:

50 g NH4ClO4 is put into 60 mL H2O and then mixed. Then added 51.5 g 54% HNO3 and mixed in a 400 mL beaker. These are then heated to boiling on a hot plate and eventually all of the AP dissolved.

Then rapidly a significant portion of 12.4 g 31% HCl in 38 mL H2O is added while the mixture is still boiling, and then later portionwise is the rest added.

After a bit of boiling there was a significant Cl2 odor (done outside), but I didn't see any nitrogen oxides or strong coloring in the solution. Then later it started boiling less, and then a white powder began precipitating here is where heating was stopped. A lot more ammonium salts deposit as the solution cools. Here I thought something must be wrong.

So, the liquid was siphoned off with a few crystals coming over in the process. Now after boiling this liquid in the beaker on a hot plate, an extremely massive amount of very thick, odorless white fumes began to show (HClO4) and heating stopped here. Then after sitting several hours the liquid was still cloudy from the solids and the volume is just under 30 mL. Letting this oily liquid stand overnight, there was finally separation possible and it was clear and colorless and could be siphoned off, where the density was about 1.68 g/cc (71.5%). Likley the formation of solids may have been avoided by use of higher strength acid and containing nitrogen oxides.

Quote:

Originally posted by garage chemist Today I tried the production of HClO4 directly from NaClO4 and HCl by those instructions: http://www.geocities.com/CapeCanaveral/Campus/5361/chlorate/...

Gmelin also mentions this method and says it doesn't work with KClO4, and gives a low yield of 80% using Ba(ClO4)2, where NaClO4 gives about 95% yield. Other formation and preparations covered are: KClO4 and H2SiF6, aqueous HClO3, Cl2 and O2 or O3, Cl2 water, O3 and hypochlorite, oxidation of chlorate, and formation from ClO2, distillation of KClO4 with H2SO4, and vacuum distillation of the same, and distillation of HClO3.

Quote:

Originally posted by hashashan
Dude, i dont need the anhydrous ... I am after the azetropic acid.
I was told that It is possible to distill it at 203 without any problems .. I just dont understand why everyone does that under vacum

72% HClO4 decomposes partly when distilled in atmospheric pressure, so it should be distilled under reduced pressure (the decomposition is about 10% at 760 mm Hg, so the highest should be 200 mm Hg according to Gmelin, preferably at 15 to 20 mm, but even this acid has a weak odor of chlorine oxides).

[Edited on 22-8-2008 by Schockwave]

dann2 - 10-11-2008 at 08:00

Hello,

Since there is no Ammonium Perchlorae thread I reckoned I would post this here.

Can Nafion (reg. trade mark) be had by the common Joe (or indeed the common Mary if she was so inclined!).

Dann2

Attachment: amm_perk.pdf (635kB)
This file has been downloaded 2104 times

dann2 - 17-11-2008 at 20:28

Hi,

Came accross this while scouring the net (as one does) and decided to plonk it up here. There are some rather 'Heath Robinson' looking apparati in it.
Have a good bedtime read!!

Dann2

Attachment: Kalichevsky_va_1924.pdf (1.6MB)
This file has been downloaded 835 times

watson.fawkes - 18-11-2008 at 07:32

Quote:

Originally posted by dann2
Can Nafion (reg. trade mark) be had by the common Joe (or indeed the common Mary if she was so inclined!).

Yeah, but it's pricey. I seem to remember that there's a project using that membrane in Fuel Cell Projects for the Evil Genius. Here's an internet seller of the membrane. A 15 cm square piece goes for 60 - 80 USD, depending on thickness. But there are more materials problems than just that.

Quote:

From A novel electrochemical process for the production of ammonium perchlorate
The 20 amp, cathode-membrane sandwich perchloric acid cell with Hastelloy cathode and platinum coated niobium anode was operated at current densitities ranging up to 10 kA m-2.

Got a small-scale source for Hastelloy or platinum-coated niobium? I suspect that anode is even more expensive.

12AX7 - 18-11-2008 at 10:46

McMaster sells hastelloy hardware, but no sheet, surprisingly. Washers up to 2" o.d.

Tim

watson.fawkes - 18-11-2008 at 13:10

Here's the Haynes International page on their corrosion resistant alloys. It's the C series that's relevant to oxidizing acids. C-2000 is probably the right one to use, as its specifically designed for chloride resistance.

JohnWW - 18-11-2008 at 14:55

See section 23 of Perrys Chemical Engineers Handbook. It has a good tabulation of corrosion-resistant alloys and other materials and their compositions, and their specific corrosion resistances or otherwise. From these tables, of the Hastelloys, only Hastelloy B, which is about 61% Ni, 28% Mo, 5% Fe, with small amounts of Mn and Si, seems to be fully satisfactory against chloride; but it is "not recommended" for use with nitric acid (which may also be the case with chloric and perchloric acids which are similarly oxidizing acids, but the data for them is not given). However, the other Hastelloys (A and C), Ni-o-nel, and 14% Si iron are quite satisfactory against oxidizing acids.

chloric1 - 18-11-2008 at 15:07

Quote:

Originally posted by Formatik

Quote:

Originally posted by garage chemist Today I tried the production of HClO4 directly from NaClO4 and HCl by those instructions: http://www.geocities.com/CapeCanaveral/Campus/5361/chlorate/...

Quote:

Well I had solid ammonium perchlorate in the beaker, added 68%, then 31.45% HCl. The solution was dark colored because the nascent chlorine formed from the HCl started attacking the ammonium ion as well as forming NOCl and NOx. When the ammonium perchlorate crystals disappeared the next day, if figured ammonium was destroyed leaving nitrogen oxides, chlorine, excess HCl, hydronium and perchlorate ions. In fact I had this as I boiled said mixture until excessing smoking stopped and I had a viscous clear excessively caustic odorless liquid. I dropped some on concrete and reaction commensed only slightly faster than 98% sulfuric acid on concrete. Some more detailed analysis would have been better here, but I know this reaction works.

S.C. Wack - 18-11-2008 at 17:25

Quote:

Originally posted by Formatik

I've used a method similar to this described in Gmelin.
...
Gmelin also mentions this method and says it doesn't work with KClO4, and gives a low yield of 80% using Ba(ClO4)2, where NaClO4 gives about 95% yield.

These references might be pretty easy to obtain for those interested in reading them, so it would be nice to know what Gmelin says they are. I'm guessing
http://dx.doi.org/10.1021/ja02212a006
and
http://dx.doi.org/10.1021/ja01919a004

Formatik - 18-11-2008 at 18:24

Quote:

Originally posted by S.C. Wack

Quote:

The NaClO4 method cited is that from Kreider in: Z. anorg. Ch. 9 [1895] 342; but also used by others in Z. anorg. Ch. 66 [1910] 244 and JACS 32 [1910] 4. Non-applicability of KClO4 and low yield from Ba(ClO4)2 is: JACS 32 [1910] 66 and Ch. Ztg. 34 [1910] 1317.

watson.fawkes - 19-11-2008 at 06:59

Quote:

Originally posted by JohnWW
of the Hastelloys, only Hastelloy B, [...] fully satisfactory against chloride; but it is "not recommended" for use with nitric acid (which may also be the case with chloric and perchloric acids which are similarly oxidizing acids, but the data for them is not given).

The Haynes site states clearly that the Hastelloy B series is resistant against non-oxidizing acids and that the Hastelloy C series against oxidizing acids. So I did just say "chloride", when I should have said "chloride with oxidizing acids".

Having said all this, the Haynes site has a link on their home page to get metallurgical advice. Anybody consider actually doing this should just contact them rather than trusting me.

Formatik - 21-11-2008 at 14:31

Quote:

Originally posted by jpsmith123
Just in case anyone's interested, here's a patent I found regarding the production of HClO4 by electrolytic oxidation of HCl.

Regards,
Joe

This would be a nice way to prepare HClO4, if one didn't have to use highly dilute solutions. Gmelin also talks about it in Cl 6, 218: electrolysis of a dilute (0.1 n-) aq. solution of HCl using an anode of pure Pt, and a cathode of copper, 50% of the used HCl converts to HClO4. The larger part which is not converted escapes as Cl2. With increasing concentration of the HCl the conversion of HCl -> HClO4 reduces drastically, using a 0.5 n-solution, only 10% is oxidized to HClO4, where the remaining HCl (in n-HCl almost completely) escapes as Cl2. However, the content of HClO3 increases with increasing HCl.

Might as well just electrolyse a -Cl or -ClO3 salt for the perchlorate salt and then use that to form the acid.

Formatik - 5-4-2009 at 16:31

Another preparation could be through oxalic acid. Using a calcium chlorate solution and oxalic acid solution to form HClO3. The calcium oxalate precipitate then removed. Calcium oxalate is around as insoluble as BaSO4. But oxalates are already more obtainable than barium salts. Lieb. Ann. 57 [1846] 138 details a method using solutions of NaClO3 and saturated oxalic acid frozen with a freezing mixture to precipitate most sodium oxalate. Though the acid made there is impure. Then decomposing the HClO3.

PHILOU Zrealone - 6-4-2009 at 07:20

Quote: Originally posted by Formatik

Another preparation could be through oxalic acid. Using a calcium chlorate solution and oxalic acid solution to form HClO3. The calcium oxalate precipitate then removed. Calcium oxalate is around as insoluble as BaSO4. But oxalates are already more obtainable than barium salts. Lieb. Ann. 57 [1846] 138 details a method using solutions of NaClO3 and saturated oxalic acid frozen with a freezing mixture to precipitate most sodium oxalate. Though the acid made there is impure. Then decomposing the HClO3.

No crossed oxydoredox between strongly opxydant HOClO2 and reducer HO2C-CO2H?

In books they propose exces HCl (30-35%) to admix with NaClO3...
So the maximum concentration of HClO3 remains in the range where it is stable (<30%), NaCl precipitates out because it is less soluble than NaClO3 or NaClO4...then boiling this down produces the desired NaClO4/HClO4...

Formatik - 6-4-2009 at 14:30

Quote: Originally posted by PHILOU Zrealone

No crossed oxydoredox between strongly opxydant HOClO2 and reducer HO2C-CO2H?

I don't think they react since no mention was made of it, but possibly on boiling or slowly on prolonged standing. Oxalic acid seems to be fairly resistant to some oxidizing acids (recall its simple preparation out of sugar, starch, etc. and boiling aq. HNO3 with evolution of copious nitrogen oxides until they cease).

Quote:

In books they propose exces HCl (30-35%) to admix with NaClO3...
So the maximum concentration of HClO3 remains in the range where it is stable (<30%), NaCl precipitates out because it is less soluble than NaClO3 or NaClO4...then boiling this down produces the desired NaClO4/HClO4...

That won't work. HClO3 reacts with HCl to form ClO2, Cl2. With more HCl you get more Cl2. Cl2 is favored with more acid and higher temperature. You will be basically left with NaCl after everything has boiled away.

Also note that boiling NaClO4 with conc. HCl (without filtering first the NaCl) will leave you with sodium perchlorate as the HCl as boiled off since boiling chlorides with aq. HClO4 leaves behind the perchlorate salt. But as discussed above, there is process starting from NaClO4 (either solid or aqueous) and conc. HCl to get the HClO4 in very high yield.

chief - 6-4-2009 at 16:35

What about Ba(ClO4)2 with H2SO4 ? This would ppt. very (completely) unsoluble BaSO4 and leave only HClO4 in the liquid phase, possible to be filtered through glasswool. Quite a simple one, is it not ?

Formatik - 6-4-2009 at 17:09

Quote: Originally posted by chief

What about Ba(ClO4)2 with H2SO4 ? This would ppt. very (completely) unsoluble BaSO4 and leave only HClO4 in the liquid phase, possible to be filtered through glasswool. Quite a simple one, is it not ?

More obtainable HCl will work as mentioned above, though there the yield is lower than if using NaClO4. Barium perchlorate is very soluble and it will work with H2SO4.

chief - 7-4-2009 at 05:13

But HCl then will yield BaCl2 ; H2SO4 on the contrary will yield _insoluble_ BaSO4, and the liquid phase will be clean HClO4 (as clean as weighted in, thereby with trace-amounts of the excess-reagent, either H2SO4 or Ba(ClO4)2 )

Formatik - 6-5-2009 at 18:22

Quote: Originally posted by Formatik

... Another preparation could be through oxalic acid. Using a calcium chlorate solution and oxalic acid solution to form HClO3. The calcium oxalate precipitate then removed. ...

I've done a probe experiment with a small amount, and this works. But I didn't get a smooth decomposition of HClO3 on boiling, either because of impurities in the chlorate or unstoichiometry. I took some aq. Ca(ClO3)2 (which was impure with a bit CaCl2) and mixed it with a clear conc. aq. oxalic acid solution. I was relying on a precipitate for stoichiometry, but it didn't appear right after mixing. So I let it sit in the cold, in dark area for several hours (longer than necessary actually, about 2 days) in a 50mL flask, lightly stoppered (the rubber cork was wrapped in PE plastic to prevent potential violent interaction with ClO2).

Later there was a fine white precipitate and the solution was now pale yellow. The solution had now the smell of ClO2, and much more so than Cl2. And after filtering, the white precipitate remained insoluble in an excess of water. After some research I also found out Ca(ClO3)2 can be obtained pure from CaCl2 or Ca(ClO2)2 by using acetone, in which it's soluble (USP2075179). Serullas (Pogg. Ann. 21 [1831] 165) had described the distillation of aq. HClO3 forming HClO4: at first the watery part comes over, then viscous, colorless liquid appears and with greater heat distills over from the retort, where chlorine and oxygen are given off at the same time.

[Edited on 7-5-2009 by Formatik]

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