I kept some sodium chlorate molten (around 270- 300°C) for two hours to make sodium perchlorate (for easy manufacture of perchloric acid without
distillation). The intended reaction was: 4NaClO3-----> 3NaClO4 + NaCl.
There was only extremely little oxygen evolution, and no measurable weight loss ( I used 35g and my scale has an accuracy of 0,1g).
After that time, I tested some of the substance for chlorate by adding HCl. It fizzed, became yellow and the test tube filled entirely with yellow
ClO2/Cl2 gas .
I tested the substance for chloride with silver nitrate solution (my used NaClO3 was very pure and chloride- free), and NO chloride could be detected.
WTF!? The test is so sensitive, it even detects the chloride ions in tap water!
So the NaClO3 did not decompose at all at the used temperatures.
Does anyone have a description of the process for NaClO3--->NaClO4 + NaCl disproportionation? It once was an industrial process.
All I need is the temperature and time needed for the process.
All the texts about disproportionation of chlorates mention KClO3 as the starting material. Does it only work with KClO3 (maybe because of the higher
melting point)? That would be bad, since KClO4 is useless for me, I need NaClO4 (it doesn't matter if it contains NaCl).12AX7 - 26-6-2005 at 16:29
How much O2 is lost in the process? (Depending on impurities I presume, especially transition metals.)
TimBromicAcid - 26-6-2005 at 19:17
ClO2 isn't terribly bad, I've accidently made large quantities of it with it only exploding once. Just don't put it under pressure of
course, any pressure. And from what I've read from many sources the method of mixing a chlorate with oxylic acid to produce ClO2 is much safer
then reacting aqueous chlorate with acid, oxylic acid looking more like an inorganic acid then an organic acid after all.
At one time I was pondering all of this and had to wonder weather there are some transition metal catalysts that could be added to speed up the
decomposition, however I figured they would also speed up the decomposition of the perchlorate formed, in addition there are some perchlorates that
are shock senstive (noteably the ones that chemist worry about forming in their fume hood exhaust and blowing up later) so I had to wonder about what
to use as a vessel.
The stability of moleten chlorates appears to be higher then expected, though not as stable as nitrates, maybe heating with hydroxides will speed up
this process, I have a paper on the subject (for the use as an oxidizing melt) but I don't have it handy at the moment.Lambda - 27-6-2005 at 00:38
Dear Garage Chemist, please wach out when mixing NaClO3 with a strong mineral acid. If your testtube would have had organic contaminats, specially
petroleum, then the ClO2 would have exploded. This may then result in expelling the acid contents into your eyes, or face were you may not have been
protected by your safety glasses. I have seen the results of such an accident, and it makes me sick to even think of it. In and old chemistry book
that I have, they reacted oxalic acid (weak acid) with NaClO3 to get ClO2. In the "Handbook of Inorganic Chemistry vol 1-3" is a much safer
procedure discribed. In this book you may allso find the procedure for Sodium Perchlorate and deffinitly Perchloric Acid (My books are nearly all
still packed in my removal boxes, this is very frustrating !).
The procedure that I will decribe here is extracted from a Dutch book on how to make Patasium Perchlorate.
REFFERENCE:
Chemische preparaten
Drs. L. P. Edel
Kluwer
Deventer - Antwerpen
(Page 65)
PROCEDURE:
Kaliumperchloraat (Dutch)
In een werkelijk schone en onbeschadigde porseleinen kroes met inhoud van 100 ml verhit men 50 g kaliumchloraat tot het juist begint te smelten. Dit
is bij ongeveer 400 graden celcius. Men houdt de smelt op deze temperatuur. Door het afsplitsen van zuurstof wordt de smelt taaier. Na 15 minuten is
de massa halfvast geworden en men laat afkoelen. Men overgiet het zout met 50 ml water en wrijft het zout fijn to er niets meer in oplossing gaat. Het
onopgeloste kaliumperchloraat wordt afgezogen en uit 200 ml heet water omgekristalliseert.
Potasium Perchlorate (Translation from Dutch)
In an absolutely clean and undamaged porcelein kroes with a volume of 100 ml, 50 g of Potasium Chlorate is heated to a point at wich it just starts to
melt. This is at about 400 degrees celcius. The melted salt is kept a this temperature. Because oxygen is evolved, the melted salt becomes stiffer.
After about 15 minutes the salt becomes semi-solid, and is then cooled down. 50 ml of water is then poured over this salt, and mixed to a fine
consistancy untill no more salt will dissolve anymore. The undissolved Potasium Perchlorate is sucked off, and recrystalised out of 200 ml hot water.
SORRY ABOUT THIS GUY'S !. This post was posted after Garage Chemist's first post, but was deleted instead of edited
(yes,...stupid me !).
You are right about that 12AX7, for Manganese dioxide or Iron oxide is often added to the Potasium chlorate melt inorder to give a swift and steady
stream of O2 gas in experiments were pure Oxygen is required. It then acts as a catalyst.
Yeah,...I know, Hydrogen peroxide may be cheaper to use for this kind of experiment.
[Edited on 27-6-2005 by Lambda]garage chemist - 27-6-2005 at 06:37
@ 12AX7: I said that there was no measurable weight loss (less than 0,1g with 35g of material), so nearly no oxygen was lost. But as you have seen, no
perchlorate was produced either.
@ BromicAcid: Metal compunds are a no-no in this reaction, as they only catalyse the evolution of oxygen from the chlorate and not the perchlorate
formation. I used a quartz crucible, a porcelain crucible can also be used, but no metal crucible, no matter what metal it consists of. (well, maybe
platinum is OK)
If the chlorate is self- made (by electrolysis), it must be cleaned from metal traces that may have gotten into it from electrode corrosion.
@ Lambda: The addition of HCl to the material was only meant as a test to find out if the material still contained chlorate. Neither perchlorate nor
chloride give any gas evolution with HCl, only chlorate does.
I also used less than 100mg of material.
Also note that mixing a chlorate with oxalic acid in solution doesn't give any reaction (I tried it). What you mean is the addition of conc.
H2SO4 to a mixture of oxalic acid and chlorate. The oxalic acid is decomposed to carbon dioxide and this dilutes the formed ClO2 to a less explosive
concentration.
So the decomposition of potassium chlorate needs about 400°C to work. I wonder if this temperature also works with sodium chlorate. As I have said
before, KClO4 is useless for me (I have some of it lying around, anyways).
What I need is a procedure for the reaction
4 NaClO3 ---> 3 NaClO4 + NaCl.
I'll heat my non- decomposed NaClO3 to 400°C for some time, to see if it yields any perchlorate.
I'll also weigh the material before and after the reaction, to see how much oxygen is evolved.
And of course I'll test for residual chlorate.
Who is the author of the mentioned "Handbook of Inorganic chemistry"?
Can I find it somewhere in the internet (FTP?) or is it only available as a real book?
[Edited on 27-6-2005 by garage chemist]
Production of CLO2 via Oxalic acid and Chlorate
Lambda - 27-6-2005 at 07:12
I am sorry Garage Chemist, for I was not clear on this matter. I still have this book somewhere packed in, and read it more than 25 years ago. But I
seem to recall that the Chlorate was mixed with Oxalic acid together with "SAND" as a protective medium to a fearly high temperature. So,
this was not done in solution. This big book (importaint things often look bigger to you when you were a kid) is more than 100 years old, and was my
first chemistry book. It was very much "Abra Ka Dabra" to me then. For this reason, it allso has special emotional value to me.
[Edited on 27-6-2005 by Lambda]garage chemist - 27-6-2005 at 07:52
I looked through my old downloaded PDFs and in "Chloratsprengstoffe" (downloaded by PAC) I found nearly the entire answer to my question.
First, sodium chlorate behaves the same as potassium chlorate on heating, it only melts at a much lower temperature, at which it doesn't
decompose noticeably.
Now, to the decomposition of both KClO3 and NaClO3, 400°C is the magic temperature, as stated in your preparation of KClO4.
However, the heating has to be continued over quite some time for a good yield to be obtained. A heating over the course of 26 hours was employed in
the early experiments!
This long heating period is not really necessary and also diminishes the yield somewhat due to subsequent decomposition of the perchlorate.
The maximum yield ever obtained was 60% Perchlorate, expressed as weight percent of the original chlorate.
About 8% oxygen is lost in the process.
Interestingly, if the temperature exceeds 420°C, the perchlorate decomposes into chlorate (which, in turn decomposes into oxygen and chloride).
So, one needs a good thermometer and a well adjustable burner or other heating source with good temp. control.
I'll try to heat my NaClO3 to about 400°C, my best thermometer only goes up to 350°C so I have to guess the correct temp.
Let's see what this produces.
My intended use for the resulting perchlorate/chloride/residual chlorate mix is to add an excess of HCl, which decomposes the residual chlorate (under
fume hood) and then boil the resulting mix. NaCl is very sparingly soluble in azeotropic HClO4 (only 1%) so it precipitates almost completely and I am
left with fairly pure HClO4 for manufacture of more exotic perchlorates like nitrosyl perchlorate and the like.
[Edited on 27-6-2005 by garage chemist]
Preparing Sodium Perchlorate
Lambda - 27-6-2005 at 11:20
Gerage Chemist, you may want to check out these very informative links on the production of Sodium Perchlorate:
Garage Chemist, do you actualy have a pdf version of "Die Chloratsprengstoffe" ?.
This book has been rewritten by Richard Escales: "Die Chloratsprengstoffe" originally printed in 1910. It has 208 pages.
I have been looking for some time now for this book, and have not yet succeded in finding a pdf copy of it. I will therefore very much appreciate it,
if you would be so kind as to make it available for download via Rapid Share.
Richard Escales: Ammonsalpetersprengstoffe.
Reprint from 1909.
Richard Escales, Alfred Stettbacher: Initialexplosivstoffe.
Reprint from 1917.
Richard Escales: Schwarzpulver und Sprengsalpeter.
Reprint from 1914.
Richard Escales: Nitroglyzerin und Dynamit.
Reprint from 1908.
[Edited on 27-6-2005 by Lambda]garage chemist - 6-7-2005 at 02:15
I've had some sucess with the thermal method now.
I heated my NaClO3 to the point where it was fizzing gently and about 40% of the surface was coated with a layer of tiny bubbles.
It was about 400- 430°C as I estimated with my 350°C thermometer.
It fizzed on and on for about 2 hours, and after that time, a flocculent (not really cristalline) precipitate occured, which slowly grew in amount,
until after about 3 hours nearly the entire melt was converted into a pudding- like sustance and heating was stopped because the lower part became
overheated due to lack of thermal convection in the pudding.
I wasn't able to weigh the end product because it had drawn a lot of water until the next morning (the pure NaClO3 had stayed perfectly dry in
the hot climate here!), indicating the formation of the very hygroscopic sodium perchlorate.
Adding some HCl to a bit of the substance still produced some yellowish gas but much less than with NaClO3.
I'll soon try to make HClO4 from it.
This is a good method for the production of perchlorates without the trouble of making/finding suitable anodes for a perchlorate cell.
Especially KClO4 is very easy to make this way, as the resulting KClO3/KClO4/KCl mix is easy to separate due to the large solubility differences of
these salts (solubility/temperature curves can be found on the sites mentioned by Lambda).
Residual chlorate can also be easily destroyed with sodium sulfite/bisulfite if the perchlorate will be used in pyro.
I'll upload "Chloratsprengstoffe" when I get around to it (I have it on CD-R).
EDIT: ATTENTION: I have uploaded "Die Chloratsprengstoffe" to Rapidshare!
The book is in german. It does not only contain lots of chlorate- based compositions, it also contains lots of methods for the PRODUCTION of chlorates
via several ways, like the electrolytical way and also the chemial way from alkali + chlorine.
Also covered is the production of perchlorates and compositions with them.
If you can read german, be sure to download this book!
[Edited on 6-7-2005 by garage chemist]
[Edited on 6-7-2005 by garage chemist]garage chemist - 6-7-2005 at 04:48
In a test I have boiled a spatula of the reaction mixture with a large excess of dilute HCl. A rather vigorous evolution of ClO2/Cl2 first started,
which then quickly stopped when all residual chlorate had been destroyed, to leave a slightly yellow solution. On further boiling, it became colorless
(ClO2 decomposes in boiling water).
It was boiled down, but it started bumping a lot and some of the liquid landed on my arm, which forced me to stop the experiment.
Some cubical crystals have crystallized from the liquid, they look like NaCl crystals (which is what one would expect from this procedure).
2ml of the liquid were added to 10ml dilute (3%) KCl solution, and soon some fine colorless crystals have precipitated out. They are pure potassium
perchlorate without a doubt, for they are not leaflets like KClO3 but rhombic like KClO4.
KClO3 would also never crystallize out from such a dilute solution.
The very low solubility of KClO4 in water is a good test of its identity.Scratch- - 6-7-2005 at 08:31
I cant read German and I dont know how to feed a PDF through a translator. So can you elaborate on the clorine + alkali => clorate. I have heard of
it, and know that it requires hot alkali solution unlike hypochlorates. I dont know however, the exact details such as yield and what reaction is
being carried out.garage chemist - 9-7-2005 at 13:09
@ Scratch: You know essentially all necessary things about this reaction. It's really not difficult.
Pass Chlorine into hot NaOH until no more is absorbed, let it stand for a few days and boil it for some time to destroy residual hypochlorite. Then
boil it down some amount (nothing should crystallise out- if it does, add some water). Then precipitate your chlorate as KClO3 with KCl solution.
A cheaper variant of this reaction is to feed the chlorine into hot calcium hydroxide suspension (Ca(OH)2 is much cheaper than NaOH) and then
precipitate KClO3 from this solution of calcium chlorate.
To the subject of this thread:
With 10g of the decomposed NaClO3, I performed the destruction of residual chlorate (using HCl) and precipitated KClO4 by adding KCl solution, cooling
and filtering. The KClO4 was washed with dist. water and recrystallized once to form larger crystals.
They were washed and dried.
The yield was a whopping 6,1g of pure KClO4, much higher than I had expected!
Boiling it with HCl gives absolutely no reaction and not even the slightest yellow color. So it is essentially chlorate- free.
The thermal decomposition of a chlorate has therefore been proved to be a good alternative perchlorate production process.
However, I won't follow this path any further, since I have constructed a neat perchlorate cell with an anode consisting of platinum wire (pics
will follow in a few days in another thread).
[Edited on 9-7-2005 by garage chemist]Zinc - 5-11-2006 at 12:28
At http://www.roguesci.org/megalomania/explo/perchlorates.html I have read that potassium perchlorate can be made by slowly adding 50 mL of
concentrated sulfuric acid to 2-5 g of potassium chlorate. The addition should be slow to avoid explosion. H2SO4 can also be replaced by nitric acid,
phosphoric acid and cromium trioxide.
Is that true? If it does work what should I do then (I assume cool the solution so that KClO4 can cristallyse)?unionised - 5-11-2006 at 13:11
I wouldn't try that unless you have remote sample handling equiment and blast screens.
Vogel's qualitative inorganic book gives that reaction as explosive and says not to try it with more than 0.1g of chlorate. On the other hand
perchorate is given as one of the products.AndersHoveland - 15-9-2011 at 10:28
Journal für praktische Chemie, Volume 23
Here is a translation of the article. Note that "chloric acid" could also alternatively refer to ClO2 in the earlier german language scientific
terminology so it is not known with certainty which exact compound the authors are refering to.
"With the intention to investigate the action of nitric acid upon potassium chlorate, a certain mass of the salt was mixed in a retort with a measured
quantity of acid and then the mixture was heated in a sand-bath. As soon as it became warm, chlorine and oxygen were formed but not as a compound, the
potassium chlorate slowly disappeared. The solution was evaporated to dryness, and then the remaining salt was found to be a mixture of potassium
perchlorate and nitrate, in a ratio of 3 equivalents of the latter to one of the first.
The action of nitric acid onto potassium chlorate differs than that of sulfuric acid on the same salt. Through nitric acid the salt is decomposed
smoothly as chlorine and oxygen form unbound, whereas through sulfuric acid, a dangerous exploding compound is formed (either chloric acid or chlorine
dioxide)
Therefore nitric acid should be used, since with this the operation can be done without a violent detonation, which occurs easily with sulfuric acid.
The action of nitric acid onto sodium chlorate is the same as by potassium chlorate. The liberated chlorine and the oxygen are in the form of a
mixture, and every 4 atoms of the salt will get 3 atoms of sodium nitrate and 1 atom of sodium perchlorate. Sodium perchlorate is very soluble and
crystallizes in rhomboides."
(The paper also deals with reaction of iodates and bromates with HNO3.)
So to disproportionate chlorate into perchlorate, use 70% nitric acid. Using sulfuric acid will likely result in explosions. This is likely because
the nitric acid can immediately oxidize the dangerous ClO2 as it is formed back to chlorate.
40% conc HNO3 is best to use.
If the nitric acid is very concentrated, it then needs to be heated while chlorate is slowly and gradually added in small portions, so that the NO2
can escape, otherwise no perchlorate will form.
phosphoric acid also decomposed chlorate to perchlorate
"The action of acids in the formation of potassium perchlorate from potassium chlorate" Hosmer Ward Stone
Some of you will no doubt be wondering about the reation of concentrated nitric acid with bromates. Such reactions only produce nitrates, while
liberating elemental bromine and oxygen, no perbromate is obtained.
"On the action of nitric acid on the chlorates, iodates, and bromates of Potassa and Soda" Fred Penny
[Edited on 15-9-2011 by AndersHoveland]dann2 - 9-11-2011 at 15:47
Some good info here on making Perk from Chlorate by heat
[Edited on 19-11-2011 by AndersHoveland]AndersHoveland - 30-6-2012 at 13:59
Here is a video showing the reaction of concentrated sulfuric acid with potassium chlorate. It becomes reddish-orange in color, begins liberating gas,
then makes a several small popping bursts that scatter some of the unreacted chlorate upwards. http://vimeo.com/37456393
The red color is probably from chloryl cations, ClO2+, that exist transiently in equilibrium. As soon as some water is added the red color completely
dissappears.
The bursts are likely little explosions from the unstable chlorine oxides produced. Just to emphasise, reacting concentrated sulfuric acid with
chlorate is extremely dangerous, and can result in violent explosion. A mix of nitric acid and concentrated sulfuric acid can, however, be used to
oxidize the chlorate to perchlorate with less danger, since the nitric acid immediately oxidizes the chlorine dioxide as fast as it forms.
[Edited on 30-6-2012 by AndersHoveland]AJKOER - 1-7-2012 at 07:42
Here is a new idea to prepare dry perchlorates, which is untested and requires some feedback, but otherwise may seemingly be doable. First prepare
Chlorine perchlorate, a pale greenish liquid which decomposes at room temperature, formula Cl2O4 or better ClOClO3. It is produced by the photolysis
of chlorine dioxide at room temperature with 436 nm of ultraviolet light:
"Abstract
OClO/O2/N2 mixtures were photolyzed in a temperature controlled 4201 reaction chamber at temperatures between 249 and 300 K and total pressures
between 0.5 and 1000 mbar. Initial OClO concentrations were in the range (1.7–5.7) · 10^15 molecule/cm3. Reaction mixtures were analyzed in situ
via long-path IR absorption using a Fourier-transform spectrometer. In some experiments product spectra were simultaneously monitored in the IR and
the UV. Depending on reaction conditions, the product IR spectra were dominated by absorption bands of Cl2O3 or Cl2O4 or a mixture of both. Evidence
is presented for the crucial role of O atoms in the Cl2O4 formation, suggesting either of the two mechanisms: (I) OClO + O + M → ClO3 + M
→ ClO3 + ClO + M → Cl2O4 + M, or (II) OClO + ClO + M → Cl2O3 + M, Cl2O3 + O + M → Cl2O4 + M. Both the weak temperature
dependence and the strong pressure dependence of the Cl2O4 yield support mechanism (I). In addition, Cl2O6 was detected as a minor product of OClO
photolysis under certain reaction conditions, both by its IR and UV absorption."
"Abstract
U.v. absorption spectra have been recorded during the low-intensity photolysis of chlorine dioxide, OClO, using a diode-array spectrometer. A
broad-band u.v. spectrum was observed which was favoured by low temperature and high OClO pressure. The absorption could be explained only in part by
the presence of ClO dimer, Cl2O2. Unequivocal assignment of the residual spectrum was not possible but it may be due to the chlorine perchlorate
molecule, ClOClO3, a recently discovered product of OClO photolysis."
Now, per Wiki Chlorine Perchlorate "is less stable than ClO2 and decomposes to O2, Cl2 and Cl2O6 at room temperature.
2 ClOClO3 → O2 + Cl2 + Cl2O6
Chlorine perchlorate reacts with metal chlorides forming anhydrous perchlorates:
which is a possible new path to the current thread topic of the production of perchlorates.
With respect to handling Chlorine dioxide per Wiki: "At gas phase concentrations greater than 30% volume in air at STP (more correctly: at partial
pressures above 10 kPa [7]), ClO2 may explosively decompose into chlorine and oxygen. The decomposition can be initiated by, for example, light, hot
spots, chemical reaction, or pressure shock. Thus, chlorine dioxide gas is never handled in concentrated form, but is almost always handled as a
dissolved gas in water in a concentration range of 0.5 to 10 grams per liter."
So having the means of producing the correct frequency of UV light to foster the reaction, noting the reaction temperatures, pressure and ClO2/O2/N2
inert gas concentration mentioned above (limiting the explosion hazard), may indeed provide a new path to dry perchlorate production. EDIT: The
following article states at room temperature that Cl2O4 is the major product of the photolysis of ClO2 with both a continuous wave (mercury lamp) and
pulsed (XeCl UV laser) light sources. Link: http://pubs.acs.org/doi/abs/10.1021/j100221a001
I also think it would be interesting in trying to dissolve ClO2 in an organic solvent (like CCl4) to form a dilute solution (under 15%), as per one
source (http://www.thesabrecompanies.com/science/chemistry.aspx ) ClO2 is highly soluble in solvents and oils as in water. Then, treat with the proper UV
exposure to create Cl2O4 and then add a suspension of say, dry CrO2Cl2, to create a perchlorate salt in a stainless steel vesel. My reading of
associated patents (see Patent 4012492 on "Synthesis of anhydrous metal perchlorates") of employing Chlorine Perchlorate in forming perchlorate salts
is that current known salts can be produced although with some new perchlorates (titanium tetraperchlorate, vanadium perchlorate, and
chromylperchlorate) can so be directly formed (link: http://www.google.com/patents/US4012492 ).
[Edited on 2-7-2012 by AJKOER]AJKOER - 5-7-2012 at 15:14
On precautions involving HClO4, please note the author's comments. In particular, the author notes that "For CCl4, HClO4 is insoluble in CCl4, and
gives upon shaking, a green emulsion, which discolors brown after several minutes welling up under formation of HCl and COCl2 (Vorländer, v.
Schilling, Lieb. Ann. 310 [1900] 374)" and further "So mixing something like CCl4 and HClO4 can cost one their lives if not wearing protective gear,
doing under fume hood,etc. It is dangerous to extrapolate so assuredly." mysteriusbhoice - 4-6-2020 at 14:26
yea tried making perchlorate with sodium chlorate using heat and I end up overshooting the temp, no clear indication like using potassium based salts
where the potassium perchlorate has a higher melting point than the chlorate and so the pastyness is even thicker.
I also think my sodium chlorate may not be as pure due to me probably pouring ferric chloride in that jar previously.
. 
