blue_vitriol
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KClO4 synthesis
Does anyone know a method for production of potassium perchlorate that is feasible for the home chemist? I have heard of syntheses using electrolysis
of KCl, but I am confused as to why it would not produce KOH and Cl2.
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12AX7
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Electrolysis does make KOH and Cl2, transiently, but the Cl2 dissolves in water and forms hypochlorite, which oxidizes further, and so on.
Chlorate can also be formed directly, at the electrode, without chlorine or hypochlorite.
A cheap graphite anode can be used. Few metals can be used, because they will be oxidized, which is handy if you wanted to oxidize the anode, but
doesn't get any chlorate.
Perchlorate can be made by very careful decomposition of a chlorate (dangerous, low yield), or by electrolysis with an expensive anode (lead dioxide,
platinum, or I think MMO).
All of this is covered in excruciating detail on the Technochemistry section.
Unless you're interested in the process itself, perchlorate is not generally a worthwhile endeavor for the ameteur -- buy it instead.
Tim
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blue_vitriol
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Does dissolved chlorine not produce HCl?
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elementcollector1
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It produces a mix of HCl and HOCl (this is the disassociation of H2O and Cl2).
Elements Collected:52/87
Latest Acquired: Cl
Next in Line: Nd
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blue_vitriol
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Thank you. If I am correct KClO3 is insoluble and will precipitate.
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hissingnoise
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KClO3 isn't insoluble but its solubility is such that it easily precipitates from a hot, saturated soln..
And NaCl is preferred as starting material because KClO3 precipitation onto electrodes reduces efficiency of reaction!
There's some relevant reading here, courtesy Alan Yates.
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AJKOER
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Here are some of my thoughts/references from a recent thread (link: https://www.sciencemadness.org/whisper/viewthread.php?tid=40... ) addressing the action of UV light on gases ClO2 as a path to dry perchlorates.
Quote: Originally posted by AJKOER  | 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:
2 ClO2 → ClOClO3
Source: See Wiki and references therein, link: http://en.wikipedia.org/wiki/Chlorine_perchlorate
Also, see "Chlorine Perchlorate Formation in the Gas Phase Photolysis of Chlorine Dioxide" by F. Zabel, link: http://onlinelibrary.wiley.com/doi/10.1002/bbpc.19910950809/...
"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."
See also "Novel ultraviolet product spectra in the photolysis of chlorine dioxide", link: http://pubs.rsc.org/en/content/articlelanding/1984/f1/f19848...
"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:
CrO2Cl2 + 2 ClOClO3 → 2 Cl2 + CrO2(ClO4)2
TiCl4 + 4 ClOClO3 → 4 Cl2 + Ti(ClO4)4 "
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 ). |
Also, the comments:
| Quote: Originally posted by AJKOER | As this thread involves the discussion of perchlorates, I can across a possibly valuable source. See https://sites.google.com/site/energeticscribble/perchloric-a...
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."
|
Also, with respect to electrochemical path to HClO4, I have previously noted that in the case of chloride free HOCl (dissolving Cl2O in water is one
preparation), the disproportionation reaction of HOCl proceeds to perchloric electrochemically. To quote "concentrated Cl-free HOCl can be oxidized
electrochemically to chloric and perchloric acids (97)." Page 554. The reference (97) is a patented process by World Pat. 9,114,614 (Oct. 17, 1991),
D. W. Crawford and co-workers (to Olin Corp.). See reference at "DICHLORINE MONOXIDE, HOCl, HYPOCHLORITES", Volume 8. Link: http://www.scribd.com/doc/30121142/Dichlorine-Monoxide-Hypoc...
With respect to the reaction of hot water, ClO2 and UV light, I would expect the creation of Cl2O4 and its decomposition (as noted above) into O2, Cl2
and Cl2O6 (or, structurally O2Cl-OClO3), the latter reacting with water as follows:
Cl2O6 + H2O --> HClO3 + HClO4
See: "Inorganic Chemistry For Undergraduates" by Gopalan, R., page 511.
Link:
http://books.google.com/books?id=Fs4zQ-hNTz8C&pg=PA511&a...
Per another source, I would expect the hydrolysis of Cl2O4 in cold water to proceed as follows:
Cl2O4 + H2O --> HOCl + HClO4
or, structurally:
Cl-OClO3 + HOH --> HO-Cl + H-OClO3
See Figure 1 in "Chlorine Oxoacids and Structure of Dichlorine Oxides", page 277.
Link: http://mdp.academia.edu/SandraQuiroga/Papers/1621127/Chlorin...
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AndersHoveland
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Warm electrolysis of salt water can produce chlorate, which is separated out and then electrolysed again. Hydroxyl radicals are able to oxidize the
chlorate anions to perchlorate. Because of the dynamics of the reaction, it typically requires high ammounts of electric current, with specialised
electrodes (discussed in the technochemistry section of this forum). Making perchlorate from electrolysis is certainly possible by the amateur home
chemist, and several members on this forum have done it.
Other methods of making perchlorate involve the careful heating of molten chlorate, and selective destruction of the contaminating chlorate by
reacting with hydrochloric acid. This is discussed in detail elsewhere in this forum.
Typically the chlorate/perchlorate is separated out from the chloride through its less soluble potassium salt, through fractional crystallization.
Ammonium perchlorate also has a very low solubility, and nearly any metal perchlorate can be derived from the ammonium salt by heating the metal
carbonate or hydroxide in solution to drive off the ammonia. (ammonium chlorate, however is dangerously unstable/hazardous and should be avoided)
| Quote: Originally posted by AJKOER | 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...
2 ClO2 → ClOClO3
|
I am not sure how practical your suggested route is. Chlorine perchlorate is extremely unstable. Unless it was dissolved in some solvent as it
was being formed, you would likely only get a few small droplets of the product. I am not sure if it would oxidize / react with carbon tetrachloride.
Perhaps it would react only gradually, or perhaps there would be an immediate explosion on contact. Pressurised or cryogenically liquified CF4 might
be the best solvent. SF6 is available in tanks for educational demonstration purposes, and is a surprisingly inert non-toxic gas.
You method might be suitable for making extreme anhydrous chlorine oxides or anhydrous perchloric acid, but for just making perchlorate salts, there
are other far easier and more practical methods.
[Edited on 2-1-2013 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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woelen
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| Quote: Originally posted by 12AX7 | | Perchlorate can be made by very careful decomposition of a chlorate (dangerous, low yield), or by electrolysis with an expensive anode (lead dioxide,
platinum, or I think MMO). |
With a PbO2-coated anode or with a platinum anode one can make perchlorates, not
with MMO. MMO is very well suited for making chlorate. Once you have a lot of chlorate, then PbO2 or Pt can be used to make perchlorate.
MMO does not make perchlorate, it simply leads to formation of oxygen and acid at the anode. PbO2 and Pt should not be used in a chloride-rich
solution, they are corroded by chloride.
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AJKOER
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Andershoveland:
I will agree for those who have invested in the equipment and knowledge base for electrolysis that it may be a better path to chlorates/perchlorates.
Yes, there are discussion/development points that need to be resolved for the Cl2O4 approach, but I would not consider the 'equipment' and 'knowledge
base' anything like that needed for electrolysis to be workable on a small scale. For example, in transparent container (use your imagination and
experiment), add KCl, spray it with water and shake the vessel to increase the surface area contact of the moist salt. Then, fill the vessel with ClO2
keeping the air/ClO2 mix under 30% to reduce explosion hazard (note, given the high solubility of the ClO2 in water, one may assume (but perhaps not
safely) that a little more than 30% is permissible. Then, leave the vessel in front of a Mercury lamp. That's it, simple and practical. Some possible
reactions include:
2 ClO2 --hv--> ClOClO3
ClOClO3 (g) + H2O --> HOCl + HClO4
KCl + ClOClO3 (g) --> Cl2 (g) + KClO4
2 ClOClO3 --> O2 + Cl2 + Cl2O6
Cl2O6 + H2O --> HClO3 + HClO4
HClO4 + KCl --> KClO4 + HCl
HCl + HOCl --> Cl2 (g) + H2O
Other solvent base approaches are also possible, and are most likely better for larger scales, but this is just an example of a relatively easy
synthesis for those with limited access to solvents other than water (note, an aqueous KCl solution saturated with ClO2 and exposed to a Mercury lamp
may also prove successful). I would also comment that small scale is not necessarily a bad starting point when it comes to the synthesis of
explosives.
To be honest, I have a wide range of feelings on possible yields. I guess, we will just have to wait until someone actually tries this or a similar
synthesis for feedback.
[Edited on 2-1-2013 by AJKOER]
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AndersHoveland
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Perchlorate can also made by heating chlorate with concentrated nitric acid, though oxygen and chlorine are also byproducts.
Penny, Lieb. Ann. 37 [1841] 204; J.pr. Ch. 23 [1841] 296.
I believe the reaction is best represented by the following equation:
(16) KClO3 + (12) HNO3 ==> (4) KClO4 + (12) KNO3 + (6) Cl2 + (6) H2O + (12) O2
By multiple evaporation of KClO3 with the regular conc. HNO3 or anhydrous HNO3, KClO4 in a yield of 30% is obtained, but the yield is almost zero if
fuming HNO3 is used, due to the nitrogen oxides which reduce the chloric acid before it can disproportionate.
(J.Am.Soc. 44[1922]143)
| Quote: |
"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."
Journal für praktische Chemie, Volume 23
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Boiling KClO3 with 85% pure H3PO4 until the yellow color has disappeared gives a 15% yield of KClO4. Boiling chromium trioxide with KClO3 gave an 11%
yield of KClO4. Not surprisingly, no KClO4 was obtained from treating 30% HClO3 with KClO3. (J.Am.Soc. 44[1922]143) (same reference again)
Concentrated sulfuric acid should not be added to chlorate, as a violent explosions often result. Careful execution of this procedure has, however,
been successful at producing perchlorate. (Sérullas. Ann. Chim. Phys. [2] 45 [1830] 272,273)
[Edited on 5-1-2013 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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AJKOER
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AndersHoveland:
Interesting and simple synthesis with a slightly less than good yield.
What's worry's me though is that does one lose one of the benefits of making perchlorates in the 1st place, namely increased stability. My concern
lies with the undesirable sensitivity (or not in terms of total explosive power) with the presence of KNO3. In particular per Wikipedia on KNO3: "it
can react explosively with reducing agents, but it is not explosive on its own" whereas for perchlorates again per Wiki: "The central chlorine in the
perchlorate anion is a closed shell atom and is well protected by the four oxygens. Hence, perchlorate reacts sluggishly. Most perchlorate compounds,
especially salts of electro-positive metals such as sodium perchlorate or potassium perchlorate, are slow to react unless heated. This property is
useful in many applications, such as flares, where the device should not explode, or even catch fire spontaneously.
Mixtures of perchlorates with organic compounds are more reactive. Although they do not usually catch fire or explode unless heated"
Please tell me I am wrong (a frequent occurrence).
[EDIT] I am wrong, the presence of KNO3 is not so much of an issue as it is very highly soluble in water while KClO4 is not, so removal and repeated
washing should make the KClO4 a safe product for mixtures.
[Edited on 6-1-2013 by AJKOER]
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hissingnoise
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The most effective oxidiser for conversion of chlorate to perchlorate is ozone . . .
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AndersHoveland
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Actually. . . ozone does not oxidize chlorate.
Yes, it is true that ozone can oxidize chlorine dioxide to Cl2O6, but this reaction apparently proceeds through a radical mechanism. This works much
better in the absence of water.
In water, the yield of perchlorate from treating ozone with chlorite is only 2.7%, the remainder only gets oxidized to chlorate.
"Perchlorate Formation by Ozone Oxidation of Aqueous Chlorine/Oxy-Chlorine Species: Role of ClxOy Radicals",
Balaji Rao, Todd A. Anderson, Aaron Redder, W. Andrew Jackson, Environ. Sci. Technol., 2010, 44 (8), pp 2961–2967
| Quote: |
"exposing aqueous solutions or Cl(-) coated sand or glass surfaces to O3 (0.96%) generated ClO4(-) with molar yields of 0.007
and 0.01% for aqueous Cl(-) solutions and 0.025 and 0.42% for Cl(-) coated sand and glass, respectively. Aqueous solutions of Cl(-) produced less
ClO4(-) than Cl(-) coated sand or glass as well as a higher ratio of ClO3(-) to ClO4(-). "
...
"a ClO2(-) solution exposed to O3 produced substantial molar yields of ClO4(-) (4% molar yield).
"Perchlorate production by ozone oxidation of chloride in aqueous and dry systems", N. Kang, W.A. Jackson, P.K. Dasgupta PK, T.A. Anderson
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In other words, only very small traces of perchlorate are formed by oxidizing chloride with ozone. The investigation did not find any traces of
perchlorate when chlorate was treated with ozone.
[Edited on 7-1-2013 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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hissingnoise
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| Quote: | Actually. . . ozone does not oxidize chlorate.
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Nooo shit? And your references?
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AndersHoveland
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| Quote: |
Solutions containing one gram of sodium chlorate and 2 cc. sulphuric acid (specific gravity 1.82) in 25 cc. water were used and were maintained at 100
°C by means of water bath. Ozonized oxygen produced by the apparatus described was allowed to bubble through these solutions for two hours. As in the
previous cases a slight oxidation of the chlorate was obtained.
The experiments with ozone show, therefore, that while ozone is capable of oxidizing acid solutions of chlorates, the oxidation is not at all
efficient. In all cases described the amount of ozone used was in excess of the theoretical amount necessary to oxidize all chlorate present. It may
be pointed out that while the amount oxidized is small it is nevertheless certain that perchlorate was produced since the micro-chemical test showed,
unquestionably, the presence of perchlorate in all solutions after treatment with ozone.
"Electrolytic Formation of Perchlorate", C. W. Bennett, E. L. Mack, Transactions of the Electrochemical Society, Volume 29, pp336-337
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While ozone is a stronger oxidizing agent in the presence of acid,
Ozone (aqueous neutral solution) 1.24v
Ozone (aqueous acidic solution) 2.08v
I suspect the difference in this reaction has more to do with the some of the chloric acid in solution disproportionating, or having some small
equilibrium with lower oxides of chlorine, through which the ozone can oxidize it to perchlorate.
[Edited on 6-1-2013 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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