Also please provide a reference to the... "the classic wet chemistry
analysis where Cd is used to reduce HCl04 in a dilute aqueous solution
at RT. I will bet a lot of money that doesn't happen.

On Tue, 24 Oct 2006 17:23:58 GMT, "hanson" <hanson@quick.net> wrote:

,;"Unknown" <dwilkins@unitelc.com> wrote in message
,;news:hgurj2lfecknrk9spcapnrr1p4pqelg2ch@4ax.com...
,;> "tadchem@comcast.net"
,;> <tadchem@comcast.net> wrote:
,;>>,;Protoman wrote:
,;>>,;Fe(OH)2+HClO4 -> Fe(ClO4)3+H2O
,;>>,;> How is the neutralization of perchloric acid with ferrous
hydroxide
,;>>,;> unbalanceable?
,;>>,;
,;> <tadchem@comcast.net> wrote:
,;>>,;As written, the proposed reaction cannot occur
,;>>,;The *balanced* redox reaction would also have different products:
,;>>,;2*Fe(OH)2 + 7*HClO4 -> 2*Fe(ClO4)3 + HClO3 + 5*H2O

How can you have HCl03 as a product. What do you suppose HCl03 would
do when it encountered some of that divalent iron. Are you suggesting
that the HCl04 reacts so fast that none of the produced HCl03 ever
sees the divalent iron???

,;>>,;Notice that the acid-base reaction occurs simultaneously.
,;>>,;If my intent was to produce ferric perchlorate, I would slowly add
the
,;>>,;ferrous hydroxide (again with stirring) to an excess of perchloric
,;>>,;acid. This would insure that the iron was oxidized as fast as it
,;>>,;was added.
,;>
,;[Wilkins]
,;> You better learn some perchloric acid chemistry before pontificating.
,;> Perchloric acid is an oxidizing agent ONLY when hot AND concentrated.
,;> You can pour room temperature 72% perchloric acid onto your hand, wash
,;> your hands with it, and then rinse with water. If you have scratches
,;> it will sting a bit but your skin is undamaged.
,;>
,;> I did this demo many times when the "safety engineers" with your
,;> knowledge of perchloric acid chemistry went balistic when I ordered 50
,;> pound quantities of concentrated perchloric acid. At the conclusion I
,;> would ask them "now lets see you do that with concentrated sulfuric
,;> acid".
,;>
,;> If you add aqueous ferrous iron to concentrate HCl04 you will end up
,;> with divalent iron. No oxidation will take place and yes you can pour
,;> water into concentrated HCl04 unlike concentrated H2S04. Fe(Cl04)2 is
,;> stable in room temperature concentrated (~72%) HCl04.
,;>
,;[hanson]
,;ahahahaha..... macho, macho!. Hey, I buy your gig here!
,;With reservations. I am not so sure that Fe(ClO4)2 is that
,;stable at RT, because if conditions exists where Fe2+ can
,;liberate nascent H from the solution (Fe2+ + H+ --> Fe3+ H_o)
,;then ClO4 will reduce, as is demonstrated in the classic wet
,;chemistry analysis where Cd is used to reduce HClO4 in
,;a dilute aqueous solution at RT.

Please provide a reference. I suspect that you are confusing this with
another reaction where the Fe+2 is complexed and the energy is
provided by a light source.

Also please provide a reference to the... "the classic wet chemistry
analysis where Cd is used to reduce HCl04 in a dilute aqueous solution
at RT. I will bet a lot of money that doesn't happen.

For one of my sources see Harvey Diehl "Quantitative Analysis
ISBN 0-914902-02-4 page 91

I quote:

"When cold, or when either hot or cold at concentrations less than 50
per cent HCl04, perchloric acid is not an oxidizing agent: hot and
concentrated , however, the dihydrate is one of the most powerful of
all oxidizing agents."

( For your info conc. HCl04 is very close to the dihydrate)

On page 92..."Metallic iron is dissolved by perchloric acid dihydrate
at room temperature , or by hot dilute perchloric acid with the
evolution of hydrogen and the formation of ferrous perchlorate, a
striking illustration of the absence of any oxidizing power. At higher
concentrations (and boiling temperatures), ferric perchlorate is
formed."

I knew Harvey (Michigan student of Willard) and he was one damned good
analytical chemist and teacher (Iowa State). He also was a consultant
to the GF Smith Chemical Co.

,;>
,;[Wilkins]
,;> For the rest of you in this thread Cl03-1 is not a product of the
,;> reaction in the oxidation of divalent iron with hot concentrated HCl04
,;> so none of you have balanced the oxidation equation correctly.
,;>
,;[hanson]
,;... AHAHAHAHA.. your tone of supreme authority is heart rendering!
,;But who was saying anything about "hot concentrated HCl04"
,;besides you? ....Now show us YOUR correct equations...
,;(1) for hot and cold, ... (2) for dilute and conc....(3) for the other
,;combinations of (1 & 2) conditions...

Perhaps the reason why I was the only one that introduced "hot and
concentrated" was because apparently I am the only one here that knows
if it ain't hot and concentrated there ain't no oxidation. If you want
RT here is the balanced equation.

Fe(0H)2 + 2 HCl04---> Fe(Cl04)2 + 2 H20

You want what happens in hot concentrated read some of GF Smith's
publications. He made a career teaching at the U. of Illinois and a
fortune manufacturing analytical reagents. Most (if not all) of the
reagent grade perchloric acid was produced by the GF Smith Chemical
Company. Hooker produced most of the commercial grade.

,;>
,;OTOH I 'll buy unconditionally your advice that hot and conc. HClO4
,;is a powerful oxidizer. This was proven and demonstrated in LA in
,;~ 1953, when the lab guys mixed Glacial Acetic Acid with conc.
,;HClO4 and heated it to 200F and got from it mirror-like gleaming
,;surfaces when they dipped or electro polished Aluminum parts in it.
,;This superjuice was rushed into immediate production in a tank that
,;contained some 3000-5000 lbs of the stuff. Shortly after start up
,;of production, it, 3 city blocks and 18 people disappeared in a bang!

You are missing some of the facts. That was a mixture of HCl04
dihydrate, glacial acetic acid, AND acetic anhydride. That mixture was
used for years as an electropolishing solution. One of the largest
uses was for deburring cash register gears. Yes the amount of solution
was large and most of a city block moved out. What you failed to
mention is that it was a large scale electropolishing operation which
generated a lot of heat AND that the cooling coils contained an
organic fluid. When the cooling coils opened and the organics were
pumped into the pot all hell broke loose. Not surprising. My
recollection is that the circulating coolant was a glycol. Does make
beautiful stainless steel or aluminum mirrors though if you know what
you are doing.

This is discussed at length in one of GF Smith's monographs. There is
a lot of info available on perchloric acid chemistry from monographs
produced by GFS Chemical Co. Most of his Ph. D. students were exposed
to a lot of perchloric acid chemistry. I was one of those students. It
is a remarkable chemical but probably has more miss information
printed about it than any other regent as shown in this thread. I
suspect the misinformation arises because of its dualistic properties
which few people understand and frankly most instructors are afraid of
it so all they teach is "its dangerous stay away from it".

,;>
,;This is a great thread, reminiscent of the discussions about SR/GR
,;where the rubber never really meets the road for practical reasons,
,;yet theoretical arguments say that it should do so .... ahahahaha...
,;Thanks for the laughs, guys.... ahahahaha.... ahahahanson

One of the few threads in this newsgroup now days that has any
chemistry.