Sciencemadness Discussion Board

KClO4 from heating NaClO3

vanBassum - 23-3-2020 at 05:54


I found on this forum somewhere a topic where the production of NaClO4 was discussed. This was done by heating NaClO3 in a ceramic crucible where the NaClO3 decomposes into NaCl and NaClO4. I have tried this method and it seems to work reasonably well. After this reaction I dissolved the products in water and added some acid to destroy the leftover chlorate. Which is then neutralized with some sodium hydroxide. At this point I should be left with a PH nutral mixture of NaCl and NaClO4. Then I used KCl to percipitate the perchlorate out of solution and do the necessary washing steps. This leaves me with potassium perchlorate free of chlorates.

Now for the question:
When is de conversion from chlorate to perchlorate and chloride done?
Right now I am eyeballing this, when things start to solidify I stop.
Problem is, I don't want to stop early and waste chlorate and acid but I also don't want to destroy my perchlorate.


B(a)P - 23-3-2020 at 11:31

Chemplayer explains this well here

AJKOER - 1-4-2020 at 20:38

Heating chlorate is apparently a good path as other paths, for example, per Wiki on HClO3 (see

"Chloric acid is stable in cold aqueous solution up to a concentration of approximately 30%, and solution of up to 40% can be prepared by careful evaporation under reduced pressure. Above these concentrations, chloric acid solutions decompose to give a variety of products, for example:

8 HClO3 → 4 HClO4 + 2 H2O + 2 Cl2 + 3 O2
3 HClO3 → HClO4 + H2O + 2 ClO2 "

where I suspect the reaction could be inherently explosive with apparently a reduced yield of perchlorate (as compared to the heating of dry KClO3, also potentially explosive with any metal oxide/organic impurity presence).

I suspect a chemically safer and more direct approach is possible, with low yields only starting, for example, with a source of hydroxyl radicals (like the UV photolysis of aqueous N2O) acting on the chlorate ion in the presence of H+:

•OH + ClO3- = OH- + •ClO3

Source: See related Reaction [14] at Supplement Table S1 at: and click on Supplement F1 to download.

And, with more hydroxyl radicals possibly:

•ClO3 + •OH -> HOClO3 = HClO4

or, via the action of the hydroxyl radical on the hypochlorite ion, followed by interaction with the chlorate radical:

ClO- + •OH -> •ClO + OH-

Source: See Eq [48] at Supplement Table S1 at: and click on Supplement F1 to download.

•ClO + •ClO3 -> ClOClO3

2 ClOClO3 -> O2 + Cl2 + Cl2O6

Source: Wikipedia , where Cl2O6 behaves as [ClO2](+)[ClO4](-).

Interestingly, the action of UV light on chlorine water (basically starting with Cl2 & HOCl and converting it into a small amount of HClO3 and then to HClO4) can form at best a 1% solution of perchlorate (see ) as:

HOCl + hv (UV Light) --> •OH + •Cl

or, more perhaps from the assistance of a photocatalytic iron complex. A distance-related source: which notes the presence of chlorate and perchlorate on the surface of Mars.

More interestingly and complex, is the possible action of HOCl on aqueous Cuprous (and Ferrous?) chlorate (or, a Fenton-type reaction occurring in the presence of chlorate):

Cu(+)/Fe(2+) + HOCl/NaOCl --( pH > 5 & <9 )--> Cu(2+)/Fe(3+) + •OH + Cl-

Source: See comments and cited references at .

•OH + ClO3- = OH- + •ClO3 (Source: See above)

•ClO3 + •OH -> HClO4

Fe(2+) + Cu(2+) <--> Fe(3+) + Cu(+) (Metal Redox Couple)

where all metal ions are present in only small amounts, but still effective due to cycling of the reaction system (and possibly limiting an alternate reaction consuming chlorate by ferrous/cuprous, see

[Edit] Found a study (Second Progress Report for “Hypochlorite – An Assessment of Factors That Influence the Formation of Perchlorate and Other Contaminants” ) that apparently supports the impact of transition metal presence in both chlorate and perchlorate formation. To quote:

"• Transition metals rapidly decompose bleach at 20mg/L.
• Perchlorate is forming as hypochlorite is decomposing. However, if hypochlorite is quenched, there is no additional perchlorate formation.
• Transition metals may play an important role in perchlorate formation, though this will need to be tested by spiking bleach solutions with lower metal concentrations to reduce the decomposition of hypochlorite.
• Chlorate and hypochlorite both appear to be involved in the formation of perchlorate
• After 10 days of incubation, the perchlorate formation is slowed down probably due to negligible amounts of hypochlorite remaining. "

The above report comments imply perhaps one should replenish periodically a transition metal-rich mix having a chlorate presence with some pH-adjusted (pH > 5 & <9) NaOCl, employing CO2 or Citric acid or Acetic acid, combined with usual cited factors of elevated temperature (but not boiling), higher strength mixes, and more ionic salts like magnesium hypochlorite/MgCl2 (derived from NaOCl/NaCl by adding MgSO4 and freezing out the Na2SO4 hydrate) or employ Ca(OCl)2.

Here is a reference on the chemistry of oxychloride:

[Edited on 2-4-2020 by AJKOER]

[Edited on 2-4-2020 by AJKOER]

AJKOER - 3-4-2020 at 09:08

OK, found a reference which is a game-changer (like when one could potentially get a 100% conversion of chloride to perchlorate) for those with access to Boron doped diamond film electrodes (BDD). Apparently, great electrodes, but perhaps too good, as when applied to the electrolysis of hypochlorite, they contributed to some major unwanted byproduct formations, not just chlorate but perchlorate as well.

Per " Understanding Chlorite, Chlorate and Perchlorate Formation When Generating Hypochlorite Using Boron Doped Diamond Film Electrodes available to quote:

"Oxychlorine radicals (ClO•, ClO2•, ClO3•) were found to chemically adsorb to both secondary and tertiary carbon atoms on the BDD surface. These chemisorbed intermediates could react with hydroxyl radicals to regenerate the original chlorine oxyanion (ClO−, ClO2-, ClO3-), and produce ≡CO• and =C•HO sites on the BDD surface. The ≡C-O• and =C•HO sites also reacted with oxychlorine radicals to form chemisorbed intermediates, which could then be converted to higher oxidation states (ClO2, ClO3−, ClO4-) via reaction with hydroxyl radicals. "


"Batch experiments show complete conversion of chloride to perchlorate can be achieved with prolonged electrolysis times. Perchlorate production can be minimized in batch and flow-through systems by using low current densities, high mass transfer rates, and high concentrations of chloride (9, 11, 12). High mass transfer rates, driven by fluid convection, are hypothesized to affect the multistep reaction for perchlorate formation from chloride, as chloride ions are progressively oxidized to higher oxychlorine anions, as illustrated by:

Cl− -> OCl− -> ClO2− -> ClO3− -> ClO4

High rates of mass transfer near the surface result in low concentrations of intermediate products, so that complete chloride oxidation to perchlorate can be minimized. "

Cited reactions of interest include:

HOCl + HO• -> OCl• + H2O [2]

OCl• + HO• -> HClO2 [4]

HOCl2 + HO• -> ClO2• + H2O [13]

ClO2• + HO• -> HClO3 [15]

More on perchlorate formation:

"Production of perchlorate from chlorate has been presented previously in Azizi et al. (37). A brief summary of the most important reactions is contained here. Figure 6 shows the generation of ClO3• and HO• near the BDD anode surface, which subsequently combine, activationlessly, to form HClO4. Chlorate radical production occurs more readily than water oxidation, as it becomes activationless at 0.76 V/SHE. The combination of radicals is activationless, with a reaction energy ΔE = −133 kJ/mol."

Comments from a prior work "Mechanism of Perchlorate Formation on Boron-Doped Diamond Film Anodes" per, Azizi et al, at , to quote:

"Perchloric acid is then formed via the activationless homogeneous reaction between ClO(3)(•) and OH(•) in the diffuse layer next to the BDD surface. DFT simulations also indicate that the reduction of ClO(3)(•) can occur at radical sites on the BDD surface to form ClO(3)(-) and ClO(2), which limits the overall rate of ClO(4)(-) formation."

The above works suggest possible parallel investigations, not just in electrolysis experiments, but also in battery/electrochemical cells (for example, the "Bleach Battery") with different carbon-based electrodes.

[Edited on 3-4-2020 by AJKOER]

vanBassum - 6-4-2020 at 04:42

Wow, I need a moment to let this information sink in. My theoretical chemistry knowledge isn't too sharp, I don't have a degree in chemistry so everything I know came from reading and experimenting.
The BDD sound like the way to go, although I have no idea how to get a few of these electrodes. The great advantage of the thermal decomposition is the ease of the process and the lack of expensive or exotic tools.