From: REMOVE_THISdwilkins@means.net (Don Wilkins)
Newsgroups: sci.chem
Subject: Re: "perchloric acid showed signs of crystallization"
Date: Wed, 27 May 1998 12:25:39 GMT

On Fri, 22 May 1998 17:17:03 -0700, Uncle Al <UncleAl0@ix.netcom.com>
wrote:

>Rebecca M. Chamberlin wrote:
>>
>> Hey gang,
>>
>> The article below appeared in our newsbulletin today.  Can anyone come up
>> with an explanation for how a bottle of perchloric acid could "show signs
>> of crystallization" -- and how it could subsequently be taken out to and
>> "safely burned"?

The article doesn't give enough information. One would like to know if
it was in the original container but even that would not assure one of
the contents.

It is contaminated with something. Perchloric acid will not crystalize
under those conditions unless someone had made the monohydrate and I
doubt that they did that. If it in fact is perchloric acid then
ammonium salts are a good guess. Prudence probably required the bomb
squad. Dilution and burning probably put their minds at ease. Dilute
perchloric won't burn but it will be dispersed enough so that their
liability coverage is safe.

For those who are beginning in the lab it is a good lesson on why any
solution which is made up should have the preparer's initials, the
date, and appropriate entries in a permanent lab notebook. With that
information one can at least have a reasonable guess at the contents.

I was a GF Smith graduate student at the time he got interested in
producing anhydrous perchloric acid and the monohydrate. We snuck into
one of the quantitative analysis labs on weekends do run the preps.
The head of the chemistry department would have gone ballistic if he
had know what we were making and Smith didn't want any interruptions.

We made the anhydrous by vacuum distillation from fuming sulfuric acid
and prepared the monohydrate by mixing equivalent amounts of the
anhydrous perchloric acid and the dihydrate. The anhydrous is not
stable.

The monohydrate (nominally 85%) crystallizes when you mix the two
acids. It reacts violently with most organic materials and is in fact
the reason he wanted to make it. Once the prep was worked out the
process was transferred to the GF Smith Chemical Company in Columbus.

The major use was by oil drilling companies. They would get the
ampoules down a well and then break them. There would be a major
explosion if there was oil present and apparently that was desirable
(the explosion that is).

At one time the monohydrate was listed in his catalog but I note that
it is no longer listed. I presume it was removed because of liability
concerns.


>>
>> "Legacy material storage problem prompts evacuation
>>
>> Nine people were evacuated from Wing 2 of the Chemical and Metallurgical
>> Research Building
>> Thursday following the discovery of two containers of chemically unstable
>> liquids. Workers Thursday morning spotted two partially filled
>> one-half-liter glass bottles of perchloric acid that showed signs of
>> crystallization, which makes the chemical unstable.  Thursday evening,
>> Laboratory and Los Alamos County emergency responders removed the two
>> bottles and took them to Technical Area 49 where they were safely burned
>> and destroyed. The perchloric acid had been moved to CMR's Wing 2 along
>> with numerous other legacy material in preparation for recycling or
>> disposal.  CMR staff will segregate the legacy chemicals as needed and
>> dispose of any other unstable chemicals that might be found before workers
>> are allowed to return to normal activities in Wing 2."
>
>What is the composition at which you get hydronium perchlorate, and is
>it a solid?

85% ought to be close enough for this kind of work and yes at room
temperature it is a solid.

>"Safely burned?"  I'd attach a small explosive charge or get a
>sharpshooter and break the bottles into a pit of crushed limestone or
>marble.  When the fizzing subsides, send the residue to a cement kiln.
>Or hold a bonfire and observe at a goodly distance until the two pops
>are heard.  The stuff really packs a punch.

From: REMOVE_THISdwilkins@means.net (Don Wilkins)
Newsgroups: sci.chem
Subject: Re: "perchloric acid showed signs of crystallization"
Date: Wed, 27 May 1998 12:25:41 GMT

On Sat, 23 May 1998 13:30:19 +12, B.Hamilton@irl.cri.nz (Bruce
Hamilton) wrote:

>In article <rmchamberlin-ya023680002205981722320001@news.lanl.gov>
> rmchamberlin@lanl.gov (Rebecca M. Chamberlin) writes:
>
>>The article below appeared in our news bulletin today.  Can anyone come up
>>with an explanation for how a bottle of perchloric acid could "show signs
>>of crystallization" -- and how it could subsequently be taken out to and
>>"safely burned"?
>
>Wow! you may have missed a real Guy Fawkes treat. I'd suspect that it may have
>been either 72.5% perchloric acid ( constant boiling 203C, approximately
>equivalent to the dihydrate) or anhydrous ( less likely ), rather than a
>dilute aqueous solution, but even the 72.5% constant boiling aqueous
>solution may become unstable over long time periods - if not properly stored.

I can guarantee it is not anhydrous that one is not stable on long
term storage.

The dihydrate can be stored indefinitely. The 72.5% acid is close to
the dihydrate in composition. Cold (room temperature) it is not an
oxidizing agent. This can be demonstrated by pouring some in the palm
of your hand and performing the hand washing experiment demonstrated
many times by Professor Smith at the U of I. If you have some small
cuts or scratches you will locate them by the sting. You finish up by
running water over your hands for the rinse.

In the demo the professor always asked someone if they would care to
try that with concentrated sulfuric acid. I did this demo for the
"safety expert" who challenged my order for perchloric acid. I will
give the guy credit because he did try it and then OKd my order.


>Most laboratories that use perchloric acid solutions  - especially
>for digestions - have special, dedicated fume cupboards that have
>water flushed walls and flues to facilitate the washing of perchlorate
>deposits from the flues. Depending on the form of the perchlorate,
>it can spontaneously explode when subjected to mild shock ( such
>as a vibrating flue ).

Yes this is true but the problem usually occurs because the perchloric
digestions have moved some concentrated perchloric acid up into the
duct system where it condenses and waits patiently for some organic
material to arrive. Even then the two can sit there calmly waiting for
a spark (hood fan motor) or other source of ignition. I have seen
ventilation fans that I couldn't lift put through the wall of the lab
(not my lab).

>
>Naturally the acid, and some deposits, can also react explosively  with
>many common organic materials, and the acid and some metal peroxides
>are particularly hazardous with combustible materials. The mixtures
>can be more sensitive and dangerous than some traditional explosives.
>One way to dispose of it is to allow it to be diluted with water, but it
>can also be carefully burned after dilution.

A dilute solution sure isn't going to burn but in an incinerator it
may react with some of the combustibles or it may vaporize up the
chimney.

>Consequently users of perchloric acid have strict guidelines, as many
>users have been harmed after explosions because digestion was allowed
>to dry out.

The explosion usually comes before it dries if one is decomposing
organic material. You get to recognize that characteristic chocolate
brown color and know instinctively that it is time to run. Don't
assume that just because it hasn't dried it won't explode and because
of hood construction a great deal of the force is directed out in your
direction (at lest until the wall goes).

>The most obvious demonstration of the power of perchloric acid
>occurred in 1947, when a acetic anhydride / perchloric acid mixture in a Los
>Angles factory exploded, killing 15 and injuring 400.

This was an electro-polishing facility. As I recall most of a city
block sized building was removed and there was a substantial hole
left. There was some evidence that the mixture was incorrect (I
believe the ratio of HCl04 and acetic anhydride was reversed and they
omitted the acetic acid)  that error coupled with some coolant pipes
containing organic material which ruptured leveled the place.

The correct mixture which also contains acetic acid was popular for
de-burring operations back in those old days when they made mechanical
cash registers and mechanical calculators. They dumped all of the
little wheels and gears into a barrel plating apparatus with the parts
as the anode in order to remove any burrs.

>There have been many
>other fatalities and injurys caused by perchloric acid - including poor
>unfortunates that have been dismantling laboratories where perchloric
>acid was used previously.

Or some poor unfortunate inherits a lab where unbeknownst to the new
occupant the previous person was a bit careless with the use of
perchloric acid.

Or some  concentrated perchloric was spilled on some of those old
wooden lab floors. Now when soaked in perchloric and dried a piece of
wood goes puff when ignited. This soaked floor just sat there waiting
for the dropped cigarette. This did not result in an explosion but the
flames got big fast. I have always thought that perchloric acid would
be great for stump removal.

>Aqueous solutions of perchloric acid only develop oxidising power with heat
>or concentration increases above 70%. Soutions that are ambient or below, and
>less than 70% are not considered to have oxidising ability, but as the
>concentration increases above 70% the oxidising ability increases dramatically.

If you are saying that the acid must be above 70% in order to be an
oxidant then that is wrong.

>The 72.5% acid reacts as a strong, non-oxidising acid at ambient temperatures,
>but around 160C it is both an extremely strong and active oxidising agent and
>a strong dehydrating reagent.

At room temperature you can wash your hands with 72.5% perchloric acid
BUT

This is where the problem occurs and why people get into trouble. The
oxidizing action of boiling perchloric acid solutions begins around
55% acid at about 150 degrees and becomes progressively stronger as
the acid concentration increases (boiling off water) and the
temperature rises.

Anyone who starts an organic digestion with 72.5% perchloric and
expects to keep it in the container when slowly raising the
temperature is asking for trouble and will find it sooner or later.


>The monohydrate ( 85% acid strength and a
>solid ), will even react with rubber, which the 72.5% acid doesn't do.  AFAIK,
>nobody today sells perchloric acid solutions at higher concentrations than the
>72.5% aqueous solution.
>
>Anhydrous perchloric acid is extremely unstable at even roon temperatures,
>and very few laboratories allow staff to use anhydrous perchloric acid these
>days, and those that do should require that it be disposed of within several
>days. This is because strong perchloric acid solutions can decompose
>spontaneously with explosive power - even after storage at room temperature,
>and one indication that the solution is becoming unstable is the appearance of
>turbidity. Anhydrous and monohydrate perchloric acid solutions can also slowly
>darken due to the accumulation of chlorine dioxide, and often explode about a
>month after preparation - if stored at ambient temperature.

Anhydrous perchloric acid is an oily liquid which is not stable and
should not be stored. The monohydrate however is a white crystalline
solid which melts at 50 degrees. It is quite stable but very
hygroscopic. At one time it was available in sealed ampoules from the
GF Smith Chemical Co. and could be shipped by common carrier.

It is no longer listed and I doubt if they would make it for anyone
but customers who they know have the appropriate experience to handle
that material.

>Inorganic chemists and metallurgists soon learn to respect perchloric acid,
>as the literature is replete with horrific examples of tragedy caused by
>somebody else's careless.  From memory, the CRC Handbook of L0aboratory
>Safety devotes a whole chapter to perchloric acid and perchlorates if you
>want more detail about the hazards.
>

From: B.Hamilton@irl.cri.nz (Bruce Hamilton)
Newsgroups: sci.chem
Subject: Re: "perchloric acid showed signs of crystallization"
Date: Sat, 30 May 1998 10:48:11 GMT

rmchamberlin@lanl.gov (Rebecca M. Chamberlin) wrote:

>I later heard, from a resident of the building that was evacuated, that the
>bottles in question were 4% perchloric acid.

I can't comment on hearsay, but chemists with pension plans that they
want to utilise, soon learn not to trust labels they didn't write on
containers that have been previously opened...

>"...other chemicals have been found, which are shock sensitive or reactive,
>such as perchloric acid, ethers, azides, and dioxane, that have either
>a) exceeded their shelf life, or
>b) crystallized and are, therefore, unstable...
>
>Requirements:  .Quantities of shock-sensitive materials must be kept to a
>minimum by maintaining precise inventory with the rate of use.  Strict
>inventory control must be maintained to ensure disposal of chemicals that
>tend to form unstable materials with age, such as ethers, and materials
>that become dangerous when they become dehydrated, such as perchloric and
>picric acids.
>
>Shock-sensitive materials must be stored in a cool, dry area, protected
>from heat and shock.  Refrigeration is recommended.  During storage, the
>materials must be segregated from incompatible materials, including
>flammables and corrosives."

>So.  Aside from the insanity of suggesting that ether must be segregated
>from flammables, and perchloric acid must be segregated from corrosives,

There is nothing insane about either of those suggestions, as
most segregation protocols would separate potentially shock
sensitive chemicals from other hazard groups, including flammable
and corrosive. There are procedures for assigning hazard priorities
to chemicals that have properties that comply with several hazard
classifications.

The formation of peroxides in ethers can result in explosions,
thus they fall into the "shock sensitive" category and should
be separated from the other flammables if peroxides are suspected.
The reason for separating perchloric acid is because anhydrous acid
( >85% ), or non-hermetic containers can also become shock sensitive
with time, and can also be flammable.

>my questions are:
>
>1) Is it possible for perchloric acid (say, 70%) to become dehydrated by
>prolonged storage at room temperature?

Not usually without a dessicating agent. But containers that are not
hermetic could be exposed to other chemicals, and form perchorates
that are shock sensitive. As the history of chemicals discovered in
labs is seldom documented, it's better to be safe than sorry. Never
trust labels on opened containers.

>2) At what dilution level (if any) do the hazards of perchoric acid become
>negligible?

Perchoric acid solutions can react to form perchlorates, which are
shock sensitive, and if the solution then evaporates... In other words,
perchloric acid has significant potential risk, depending on the volume,
concentration, and associated materials. Solutions can also react
explosively with a range of other materials... There is no general rule,
so cautious people assume the worst case...

>3) What are the comparative hazards of perchloric acid versus (solid)
>perchlorate salts?  (i.e. did they omit something important?)

As above, the hazards depend on the specific concentration, volume,
and container for the acid, and the individual perchlorate for the
solids. It's not a good idea to generalise, but to identify the
hazards specific to the compound, such as in Sax, MSDS, Bretherick,
etc.

>4) Can anyone suggest any good RATIONAL references on the safe handling of
>these materials?

I've previously pointed towards the chapter in the CRC Handbook of
Chemical Safety. I've just checked my 2nd edition, and there is a
whole section devoted to " Handling Perchloric Acid and Perchlorates"
p.265-276. It also points to an article in Chem. Eng. Prog. v.62
p.109-114 (1966) on "safe handling of perchloric acid' - which I
haven't read. There may be also be more information in a later
edition of tthe CRC Handbook....

Bretherick's Handbook of Reactive Chemical Hazards also has an
extensive section on Perchloric Acid, and points to an article
in the J.Chem.Ed. that details safe handling procedures ( v.49,
p.A463 (1972) - which I haven't read ).

>I am interested in sources that are aimed at a bachelor's degree
>level chemist, NOT at the ES&H "experts".  (A little
>knowledge is a dangerous thing!)

Those "experts" usually have to clean up a whole range of
chemicals, usually poorly-labelled, left hiding in unsuitable
cupboards by thoughtless "chemists". As my earlier message
noted, the use of perchloric acid in laboratories has been
punctuated by the death/injury of innocent workmen who have
been dismantling laboratories where perchloric acid was used.
Old, bold chemists aren't common.

    Bruce Hamilton

From: B.Hamilton@irl.cri.nz (Bruce Hamilton)
Newsgroups: sci.chem
Subject: Re: "perchloric acid showed signs of crystallization"
Date: Sat, 30 May 1998 10:48:17 GMT

REMOVE_THISdwilkins@means.net (Don Wilkins) wrote:

>On Sat, 23 May 1998 13:30:19 +12, B.Hamilton@irl.cri.nz (Bruce
>Hamilton) wrote:
...
>>Wow! you may have missed a real Guy Fawkes treat. I'd suspect that it may have
>>been either 72.5% perchloric acid ( constant boiling 203C, approximately
>>equivalent to the dihydrate) or anhydrous ( less likely ), rather than a
>>dilute aqueous solution, but even the 72.5% constant boiling aqueous
>>solution may become unstable over long time periods - if not properly stored.
>
>I can guarantee it is not anhydrous that one is not stable on long
>term storage.

My memory was that there was considerable debate about whether
anhydrous acid ( usually considered to be >85% ) instability
could be related to impurities, and that highly pure anhydrous
perchloric acid may be more stable.

  " Decomposition of Anhydrous Perchloric Acid.
  ... The explosive reputation of HClO4 itself is probably
  exaggerated, at least for very pure samples; however the acid
  ages badly, and the presence of impurities, especially Cl2O7
  greatly facilitates decomposition. " [1]

Perhaps there is later information that I haven't seen...

>>Most laboratories that use perchloric acid solutions  - especially
>>for digestions - have special, dedicated fume cupboards that have
>>water flushed walls and flues to facilitate the washing of perchlorate
>>deposits from the flues. Depending on the form of the perchlorate,
>>it can spontaneously explode when subjected to mild shock ( such
>>as a vibrating flue ).
>
>Yes this is true but the problem usually occurs because the perchloric
>digestions have moved some concentrated perchloric acid up into the
>duct system where it condenses and waits patiently for some organic
>material to arrive.

Well, " Dismantling an exhaust ventilation system suspected of
contamination with perchlorates " starts with the assumption that
shock-sensitive perchlorates are present, rather than perchloric
acid, a fairly justifiable assumption, given some of the accidents
reported [2].

>>Aqueous solutions of perchloric acid only develop oxidising power with heat
>>or concentration increases above 70%. Soutions that are ambient or below,
>>and less than 70% are not considered to have oxidising ability, but as the
>>concentration increases above 70% the oxidising ability increases
>>dramatically.
>>
>>The 72.5% acid reacts as a strong, non-oxidising acid at ambient temperatures,
>>but around 160C it is both an extremely strong and active oxidising agent and
>>a strong dehydrating reagent.
>
>If you are saying that the acid must be above 70% in order to be an
>oxidant then that is wrong.

" Cold perchloric acid, 70% or weaker, is not considered to have significant
  oxiding power. The oxidizing power, however, increases rapidly as the
  concentration increases above 70%. " [2].
" Although the 70-72% acid of commerce behaves when cold as a very strong,
  but non-oxidising acid,  it becomes an extreme oxidant and powerful
  dehydrator at elevated temperatures (160C) or when anhydrous." [3]

I admit that my memory had forgotten the "significant", but I clearly
stated that the temperature was at ambient or below, and the above
seem to consider cold 70% acid as non-oxidising....

     Bruce Hamilton

[1] Comprehensive Inorganic Chemistry - 1st ed.
    Pergamon Press ISBN 0-08-017275-X (1973)
    p.1446

[2] CRC Handbook of Laboratory Safety - 2nd ed.
    p.266,273.

[3] Brethericks Handbook of Reactive Chemical Hazards - 5th ed.
    Butterworth-Heinemann ISBN 0 7506 1557 5  (1995)
    p.1246

From: REMOVE_THISdwilkins@means.net (Don Wilkins)
Newsgroups: sci.chem
Subject: Re: Concentrating Azeotropic HCl
Date: Mon, 16 Oct 2000 09:34:44 -0500

On Sun, 15 Oct 2000 21:21:47 +0200, "chris.lee" <chris.lee@infonie.fr>
wrote:

>,;The composition of azeotropic mixtures is pressure dependent. You can make
>,;really conc HCl by evaporating the dilute stuff in a rotary evaporator. It's
>,;worse than fuming acid. I don't know where I read it, but as far as I
>,;remember you end up with the monohydrate of HCl.

I don't know where you read it either but it must have been in a
science fiction novel. I don't know where you get your information but
it sure is different than the stuff I learned.

pressure(in mm)------concentration of HCL(%)

780----20.173
770----20.197
760----20.221
750----20.245
740----20.269
730----20.293

Doesn't look very promising to make the monohydrate (66+%) does it?

>,;Incidentally, you can make anhydrous perchloric acid the same way. I don't
>,;advise you to try it as anything organic that happens to be present becomes
>,;rapidly inorganic.

This also is just plain wrong.

You can not make anhydrous perchloric acid by vacuum distillation. The
dihydrate (~72%) is the azeotropic concentration. I was a student of
Professor GF Smith and worked with him when he developed a procedure
for making anhydrous perchloric acid and for making the monohydrate.
You can't even make the monohydrate by vacuum distillation.

We started with the dihydrate, added it to fuming sulfuric acid and
distilled the anhydrous perchloric acid from that mixture. It is not a
procedure for the faint of heart or for those who don't understand the
properties of perchloric acids. Make a mistake and you can lose the
building.

You make the monohydrate by preparing the anhydrous acid and mixing it
with an equivalent amount of the dihydrate. The monohydrate is a solid
at room temperature and reasonably stable (but very reactive). The
anhydrous acid can be kept refrigerated for a short period of time but
is inherently unstable. The latter is a hazardous material to have
even when there are no reducing agents present (even when kept
refrigerated).

Professor Smith owned the G. Frederick Smith Chemical Company which
was/is the largest producer of perchloric acid in the USA. He and I
were making the anhydrous acid in a quant lab on the third floor of
the Chem Annex on weekends because he was concerned with making the
stuff at the factory in Columbus by people who were not trained
chemists.

The monohydrate was sold to oil drilling companies. There was no
market for the anhydrous acid. AFAIK neither the anhydrous acid or the
monohydrate are currently available commercially.

Believe me when I tell you that if all you had to do to prepare
anhydrous perchloric acid was distill it in a rotary evaporator the
Professor would have done it. He knew more about perchloric acid and
the perchlorates than any one else in the country.