VladimirLem
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KCLO4 from H2O2 ?
Hi
Does KCLO3 really reacting with H2O2 to KCLO4 ?
KCLO3+H2O2 -> KCLO4+H2O
(Potassiumchloarte+hydrogen-peroxide -> potassiumperchlorate+water)
Some "Chemical Equations Balanced"-Pages say yes, but i read some posts in other boards, that this doesnt work?
And google found nothing similar...
And, if that works, how many 20% H2O2 solution would be needet at, lets say, 10gramms KCLO3 ?
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quicksilver
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Frankly your question belongs in General Chemistry.
A balanced equation does not always discern a synthesis (but often it can). Also watch your capital letters, it's: KClO3 + H2O2 = KClO4 + H2O. The
page was designed for understanding the balancing of equations not developing a synthesis per se'. That page is a useful tool. Some years back a
fellow here wrote a small program (in Delphi I think) that was a "stand alone" balance of equations tool".
That an equation appears to balance does not highlight the idiosyncrasies (purity, percentages, application methods, etc) for what a synthesis would
entail in many cases.
Dinge sind nicht immer das, was sie scheinen.
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AndersHoveland
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(First, let me state that I have not actually done this reaction.)
No, there would not be any reaction between potassium chlorate and hydrogen peroxide unless the solution was acidified, in which case the hydrogen
peroxide would act like a reducing agent, with oxygen gas being liberated while leaving behind potassium chloride.
KClO3 + (3)H2O2 --> KCl + (3)H2O + (3)H2O
If you want to oxidize chlorate to perchlorate, you could use potassium ferrate, K2FeO4, then slightly acidify the solution. The ferrate may be easily
formed by reacting iron sulfate with a good portion of concentrated sodium hydroxide solution, then bubbling in chlorine. The solution will turn a
brownish-burgundy purple, indicating the formation of ferrate. Unfortunately the solution will still contain chloride ions that you will want to get
rid of before you acidify the solution to make the ferrate act as a powerful oxidizer. Usually either potassium or calcium ferrate can be precipitate
out, leaving the chloride ions behind still dissolved in solution. One note: ferrate immediately decomposes to oxygen if the solution is not kept
alkaline, so the acid should only be added after the alkaline ferrate is mixed with the chlorate. Chloride ions are not desirable (so do not use
hydrochloric acid) because ferrate which is being acidified is a very powerful oxidizer, and would oxidize the chloride ions instead of the chlorate.
http://library.sciencemadness.org/library/ferrates.html
https://sites.google.com/site/ecpreparation/ferrate-vi
Of course, alternatively chlorate can be made to disproportionate to perchlorate by careful heating. Read the section titled "Potassium Perchlorate"
here
https://sites.google.com/site/energeticchemical/time-tracker
(toward the middle of the page)
[Edited on 9-7-2011 by AndersHoveland]
I'm not saying let's go kill all the stupid people...I'm just saying lets remove all the warning labels and let the problem sort itself out.
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VladimirLem
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thanks, AndersHoveland
the last link is great 
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hissingnoise
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| Quote: | Ammonium Perchlorate
To a boiling solution that contains as much sodium perchlorate and ammonium chloride as will dissolve, allow this solution to cool to room
temperature, then cool to around 15degC. Crystals of Ammonium Perchlorate will start precipitating out of solution. NH4ClO4 is only about 10% as
soluble as NaClO4. |
Mixing cold solutions of NH<sub>4</sub>NO<sub>3</sub> and NaClO<sub>4</sub> produces an immediate precipitate of
reasonably pure NH<sub>4</sub>ClO<sub>4</sub>!
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franklyn
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Low pH promotes reaction with materials in H2O2 solution which
are susceptable to acids , dissolving Cu, Ag, Hg.
H2SO4 + H2O2 is more oxidative than Nitric acid.
High pH promotes reaction with other materials in solution, since
H2O2 is decomposed by basic substances.
Salts which are near pH neutral or acid are not reactive in H2O2.
An interesting exception is the ClO3- ion which neutral or alkaline is
not reactive. ClO4- ion is not reactive at any pH.
Mg(ClO4)2 is used as a drying agent when concentrating H2O2.
.
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VladimirLem
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i've got another short question...
many people say, that you should use HCL at the chloratecell to get a low ph-value...( a few drops all few hours)
but, wouldn't HCL reduce (destroy) the produced KCLO3 or (if it is made with NaCl) NaClO3 ?
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hissingnoise
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I used to add a little HCl every two, or so, days just to correct alkaline-drift and if I overdid the addition a lacrymatory gas was produced . . .
A chlorine oxide, I assumed!
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AndersHoveland
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"ClO4- ion is not reactive at any pH"
This is not entirely true. Although perchlorate salts are typically very inert at room temperature, the salts become reactive oxidizers under
extremely acidic conditions, such as the use of concentrated sulfuric acid. I suspect that 90% concentrated perchloric acid would violently react (if
not explode) with hydrogen peroxide to liberate oxygen and chlorine. Indeed, HClO4 concentrations over 74% can oxidize hydrogen chloride, whereas 70%
concentrated HClO4 only reacts with zinc to give off hydrogen gas, with none of the perchlorate anions being reduced.
"you should use HCL at the chloratecell to get a low ph-value...( a few drops all few hours),
but wouldn't HCL reduce (destroy) the produced KCLO3 ?"
Although HCl reacts with KClO3, it is actually an equilibrium reaction, so if there is only a very miniscule quantity of HCl, there will effectively
be no degradation of the chlorate. These types of equilibrium reactions are typically strongly pH dependent.
"I used to add a little HCl every two, or so, days just to correct alkaline-drift and if I overdid the addition a lacrymatory gas was produced . . . A
chlorine oxide, I assumed!"
The particular reaction of HCl and KClO3 only produces chlorine gas. The chemistry of preparing chlorine oxides is surprisingly complicated. This
topic has been discussed before.
http://www.sciencemadness.org/talk/viewthread.php?tid=16378
Chlorine dioxide can be ignited and is capable of self-sustaining decomposition with a pale grey flame. At higher concentration it explodes when
ignited (it sometimes even self-ignites, for no apparent reason).
I'm not saying let's go kill all the stupid people...I'm just saying lets remove all the warning labels and let the problem sort itself out.
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AndersHoveland
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oxidation of chlorate with hydrogen peroxide
Some research that should completely put to rest any questions about the reaction between potassium perchlorate and hydrogen peroxide:
Chlorates, when heated to moderate temperatures, disproportionate into perchlorate and chloride.
4NaClO4 --> NaCl + 3NaClO4
Fowler and Grant found that on heating chlorate with silver oxide that the chlorate was completely converted to perchlorate without loss of oxygen,
metallic silver also forming.
J. Chem. Soc. 57, 272 (year 1890)
A solution containing one gram sodium chlorate, 1 cc sulfuric acid (specific gravity 1.82), and one gram potassium permanganate in 100 cc water was
boiled for 30 minutes. The solution showed no perchlorate present.
Hydrogen peroxide failed to oxidize chlorate to perchlorate under alkaline, neutral, or acidic conditions, although minute traces of perchlorate did
form under acidic conditions.
“Electrolytic Formation of Perchlorate” C. W. Bennett, E.L. Mack. Chemical Engineer, volume 23, p206
Alkaline
A boiling solution of sodium peroxide failed to oxidize chlorate to perchlorate.
A solution containing one gram sodium chlorate and 1cc ammonium hydroxide (specific gravity 0.90) in 15 cc hydrogen peroxide (30%) was boiled for 30
minutes. Analysis of the solution failed to show any traces of perchlorate, thus showing that alkaline hydrogen peroxide is not a sufficiently
powerful oxidizer to convert chlorates to perchlorates.
Oxidizing Power of Alkaline Hydrogen Peroxide
Although alkaline solutions of peroxide failed to oxidize chlorate, it is nevertheless worth mentioning the high oxidizing power of such solutions.
Although alkaline peroxides (such as CaO2) are stable in the absence of water, hydogen peroxide slowly decomposes in aqueous alkaline solution. A
mixture of hydrogen peroxide and ammonium hydroxide (in a 1:3 ratio) acts as a reactive oxidizer, which can attack organic compounds and elemental
carbon. The reaction rate is negligible at room temperature, but when heated to 60°C the reaction becomes vigorous and self-sustaining. Such
solutions are sometimes known as "base piranha". With a 1:1:5 volume ratio of NH4OH, H2O2, and H2O, respectively, the half-life times of peroxide were
4 hours at 50°C and 40 minutes at 80°C. "Reaction of Ozone and H2O2 in NH4OH Solutions and Their Reaction with Silicon Wafers" Japanese Journal
Applied Physics. 43 (2004) pp. 3335-3339. Magnesium hydroxide inhibits the formation or reactive radicals in alkaline solutions of hydrogen peroxide,
interrupting the free radical chain reactions by catching the superoxide anion radicals. Zeronian SH & Inglesby MK (1995) "Bleaching of cellulose
by hydrogen peroxide". Cellulose 2: 265-272.
Nuetral
One gram of sodium chlorate was dissolved in 15 cc hydrogen peroxide (30%) and the solution evaporated to dryness on the water bath. The residue was
dissolved in a second 15 cc portion of hydrogen peroxide and again brought to dryness. After dissolving in water, and analysis was performed. This
experiment showed that chlorate in neutral solutions is not oxidized to perchlorate by 30% hydrogen peroxide.
Acidified
One gram sodium chlorate was dissolved in 25 cc of hydrogen peroxide (30%) which had previously been acidified wih 1 cc sulfuric acid (specific
gravity 1.82). The solution was boiled for one hour. Soon after the solution had reached he boiling point a yellow gas was evolved. This was at first
thought to be chlorine but more careful examination showed it to be a mixture of chlorine dioxide and chlorine. Analysis showed small traces of
perchlorate had formed. This experiment showed that chlorate, through the action of acidic solutions of hydrogen peroxide, is largely converted to
chloride. A considerable amount of chlorine and chlorine dioxide is evolved at the same time. Acidic solutions of 3% hydrogen peroxide also were shown
to reduce chlorate to chloride.
It has been suggested that in the above reaction the intermediate formation of small amounts of hydrogen chloride interferes with the reaction,
catalytically causing decomposition of the chlorate. Chlorine is known to react with hydrogen peroxide to form hydrochloric acid and oxygen gas. The
hydrochloric acid thus formed would attack the remaining chlorate, the products of the reaction being chlorine and chlorine dioxide, the chlorine then
reacting with more hydrogen peroxide to again form hydrogen chloride.
Reaction of concentrated sulfuric acid with sodium chlorate did not produce any perchlorate, but it has been reported by other sources that
perchlorate is indeed produced. This may be due to different acid concentrations and ratios of reactants.
The reaction between solutions of chloric acid (HClO3) and hydrogen peroxide does not have any appreciable reaction rate until a temperatures above
70degC. (note that perchlorate is not a reaction product in the decomposition reaction, although it may be likely that traces are formed). experiments
conducted by Sand, published in Zelt phys. Chem.,50, 465 (year 1904)
[Edited on 26-7-2011 by AndersHoveland]
I'm not saying let's go kill all the stupid people...I'm just saying lets remove all the warning labels and let the problem sort itself out.
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Globey
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".....Acrylic fabric basically a deadly gas that has been polymerized......."
Indeed, it is, The one overlooked the most however, is (PFIB) perfluoroisobutylene, This results from the pyrolysis of polytef,
(Teflon), and is roughly an order of magnitude more toxic that phosgeene, Plus, it is odorless, and the insidious nature of it's effect
don;t show themselves immediately, (1-4 hours). It's a real big nasty.
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The WiZard is In
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Noted in passing
Accession Number : AD0017221
Title : RESEARCH ON THE EXPLORATION OF METHODS TO
PRODUCE CHLORATES AND PERCHLORATES BY MEANS OTHER
THAN ELECTROLYTIC. PART 2
Descriptive Note : Interim research rept. 1 Jun-30 Nov 1952
Corporate Author : MATHIESON CHEMICAL CORP BALTIMORE MD
Personal Author(s) : Naughton, J. M.
Handle / proxy Url : http://handle.dtic.mil/100.2/AD017221
Report Date : 30 NOV 1952
Pagination or Media Count : 26
Descriptors : *PERCHLORATES, CHEMICAL REACTORS,
OZONE, SODIUM COMPOUNDS, CHLORATES, PRODUCTION,
POTASSIUM COMPOUNDS
Subject Categories : INDUSTRIAL CHEMISTRY AND CHEMICAL PROCESSING
------------
Accession Number : AD0014822
Title : RESEARCH ON THE EXPLORATION OF METHODS TO PRODUCE CHLORATES AND PERCHLORATES BY MEANS OTHER THAN ELECTROLYTIC. PART 1
Descriptive Note : Interim research rept. 1 Jun-30 Nov 1952
Corporate Author : MATHIESON CHEMICAL CORP BALTIMORE MD
Personal Author(s) : Jaszka, D. J.
Handle / proxy Url : http://handle.dtic.mil/100.2/AD014822
Report Date : 30 NOV 1952
Pagination or Media Count : 25
Descriptors : *SYNTHESIS(CHEMISTRY), *PRODUCTION, *PERCHLORATES, DIOXIDES, CHLORATES, THEORY, LEAD COMPOUNDS
Subject Categories : INDUSTRIAL CHEMISTRY AND CHEMICAL PROCESSING
----------
Accession Number : AD0016814
Title : RESEARCH ON THE EXPLORATION OF METHODS TO
PRODUCE CHLORATES AND PERCHLORATES BY MEANS OTHER
THAN ELECTROLYTIC
Descriptive Note : Interim research rept.
Corporate Author : OLIN MATHIESON CHEMICAL CORP BALTIMORE MD
Personal Author(s) : Dexter, T. H. ; Naughton, J. M.
Handle / proxy Url : http://handle.dtic.mil/100.2/AD016814
Report Date : MAY 1952
Pagination or Media Count : 42
Descriptors : *PRODUCTION, *PERCHLORATES, *CHLORATES, CHEMICAL REACTORS
Subject Categories : INDUSTRIAL CHEMISTRY AND CHEMICAL PROCESSING
Distribution Statement : APPROVED FOR PUBLIC RELEASE
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AndersHoveland
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Just want to mention that perchlorates can be prepared, in certain instances, by boiling chlorates with highly concentrated acid, but this can be very
dangerous.
References:
"The action of acids in the formation of potassium perchlorate from potassium chlorate" Hosmer Ward Stone
Thesis by Zick, University of Wisconsin, 1917.
Journal für praktische Chemie, Volume 23
See the thread "Decomposition of sodium chlorate",
http://www.sciencemadness.org/talk/viewthread.php?tid=4077
[Edited on 15-12-2011 by AndersHoveland]
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neptunium
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simply heating (gently!!) potassium chlorate will generate the perchlorate and the chloride much the same way sodium chlorate is made from bleach .
separating the chloride from the perchlorate is just a question of solubility in cold water ....
im just saying
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AndersHoveland
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Quote: Originally posted by neptunium  | simply heating (gently!!) potassium chlorate will generate the perchlorate and the chloride much the same way sodium chlorate is made from bleach .
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In the instance of disproportionating hypochlorite to chlorate, the solution only needs to reach a boiling heat. But to decompose chlorate to
perchlorate, it requires much higher temperatures, just above the melting point of the chlorate salt (m.p. of KClO3 is 356 °C ). At
this temperature, all the water would have already boiled away.
Simply boiling a solution of potassium chlorate in water is not going to cause any chemical changes.
Care must be taken to ensure that the temperature of the molten chlorate does not become too hot, as this could result in violent decomposition of the
chlorate and/or decomposition of the perchlorate that forms. The potassium chlorate is typically held just above its melting point for several
minutes, and the heat is removed if any bubbles appear. The potassium perchlorate that begins to form has a higher melting point, so
eventually after gradual and cautious heating, the molten salt may begin to solidify.
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Pulverulescent
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I sense a deja vu --- ozone is the only oxidiser aggressive enough to oxidise chlorate to perchlorate!
P
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AndersHoveland
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Bubbling elemental fluorine into a solution of chlorate and water will also oxidize the chlorate to perchlorate, but the reaction is not efficient.
KClO3 + H2O + F2 --> KClO4 + 2 HF
Mixed concentrated nitric and sulfuric acids can also oxidize chlorate under some conditions.
The disproportionation of chlorate using concentrated sulfuric acid will also result in the formation of some perchlorate, although this is not an
efficient reaction, and there is severe danger of explosion.
Potassium ferrate, K2FeO4, would also no doubt be able to oxidize chlorate. Iron can be easily oxidized to ferrate when the pH is high enough, but the
resulting ferrate becomes extremely oxidizing when the solution is made slightly acidic. (not that H2O2 will not work because it immediately reacts to
reduce ferrate to liberate oxygen)
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Neil
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ah finally a use for all my F2...
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AndersHoveland
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I sense sarcasm.
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