Sciencemadness Discussion Board

HClO3 Preparation

Kola - 7-3-2012 at 10:14

I'd like to know if there's any system in which HOCl can be oxidized to HClO3.
Also; i'm particularly interested in the concentration. What method can I use to prepare Conc HClO3 from HOCl?

disulfideprotein - 7-3-2012 at 10:24

Heat the HOCl and it will oxidise to HClO3 in this equation: 3HClO → HClO3 + 2 HCl Pleased be warned that hydrogen chloride gas is produced so where a respirator or do this under a fume hood. ALSO: IT IS AN EXTREMELY POWERFUL OXIDISER.


[Edited on 7-3-2012 by disulfideprotein]

Adas - 7-3-2012 at 10:56

This is bullshit. HClO3 above certain concentration decomposes into ClO2, which is an explosive gass and can be set off by sunlight. HClO3 explodes on contact with organics.

disulfideprotein - 7-3-2012 at 11:04

I will say that it is explosive in an edit but it does not always decompose when you make it, you could just make it at a dilution 30-40 percent is fine. Also it has to be further heated after the dilution passes this range. What I said is not bullshit but it is certainly not a safe way to make this, the safer way would be to use barium chlorate in this reaction Ba(ClO3)2 + H2SO4 → 2HClO3 + BaSO4

woelen - 7-3-2012 at 11:21

Making HClO3 from HOCl is utter nonsense. HOCl will decompose to mainly HCl, Cl2 and H2O when it is heated. HOCl is very unstable. Hypochlorite indeed can be converted to chlorate, but this reaction is very pH-sensitive. When all hypochlorite is present as ClO(-), then the reaction is slow and incomplete. When most hypochlorite is present as HOCl, then there will be self-catalysed decomposition to Cl2. Any HClO3 formed will also lead to formation of HCl, which very strongly reacts with HOCl to give Cl2 and H2O. The best pH is around 6, which is very weakly acidic.

If you really want to make a solution of HClO3, then indeed the way to go is reaction of dilute sulphuric acid with barium chlorate. Another option, which gives fairly pure HClO3 is mixing a hot solution of KClO3 with a hot solution of HClO4. On cooling down, KClO4 is formed. When the liquid is cooled to well below zero, then 99% or so of all KClO4 precipitates from the liquid. Use a slight excess of HClO4 and the resulting liquid will be almost free of any potassium.

At concentrations of 20% or so, a solution of HClO3 is quite stable and then it may be hot. At 35% to 40% the solution already is much less stable and then heating leads to disproportionation to HClO4, ClO2, Cl2 and H2O. You can easily see this, the liquid then turns yellow.

disulfideprotein - 7-3-2012 at 11:26

Very cool I knew it decomposed but not in self catalysed way upon heating.

Formatik - 9-3-2012 at 02:22

The (thermal) decomposition of hypochlorous acid can be used to make chloric acid. This is seen in the international patent attached below.

Chloric acid has the potential for explosion when combined with organic materials, especially since ClO2 is a decomposition product. The demonstration of its strong oxidizing power has been to let a paper soak in the concentrated acid and then allowing to evaporate, the paper inflames.

But chloric acid is a relatively tame acid, I wouldn't store it since it degrades slowly forming ClO2 and Cl2. But the fresh material is not going to just explode on you or something even when being vacuum concentrated (up to around 40%, it decomposes rapidly around at 50% and more).

If you are looking to get perchloric acid out of it, then this is known to form by boiling the concentrated acid to decomposition.

Attachment: WO91-03421.pdf (534kB)
This file has been downloaded 486 times

AndersHoveland - 9-3-2012 at 02:32

Actually, it is not impossible to make HClO3 from HOCl (with ozone!).
It is just very impractical because there are much better ways.

And HClO3 will oxidize HCl.

The usual method is reacting around 20% conc. H2SO4 solution with barium chlorate.

Agree completely with Woelen's post.

AirCowPeaCock - 9-3-2012 at 07:01

Anders, wouldn't it oxidize into the unstable chlorous acid (which may decompose significantly) before being oxidized to chloric acid? And couldn't the chloric acid be oxidized further to perchloric acid?

woelen - 9-3-2012 at 07:48

"wouldn't it oxidize into" --> very unclear, what is it? What compound is oxidized and what is the oxidizer and what is the reductor?

I have the impression that you are talking about a disproportionation reaction, in which the same compound is both an oxidizer and a reductor.

HOCl hardly decomposes to HClO2. The compound HClO2 is more stable than HOCl, but in the presence of Cl(-) it nearly quantitatively is converted to ClO2. I have done many experiments with chlorites and found that in the presence of chloride you get pure ClO2 and this reaction is fast. Without chloride the decomposition is much slower, even at low pH, but once the decomposition sets in, it speeds up more and more. This can be explained, because as soon as some chloride is formed from decomposition, then more chloride is formed from HClO2 (this disproportionates to ClO2 and Cl(-) and this reaction is catalyzed by Cl(-)).

Chloric acid is not easily oxidized to perchloric acid, unless you allow it to decompose at high concentration or in the presence of conc. H2SO4. Under such condition you have further disproportionation to HClO4, ClO2 and chlorine in lower oxidation states. This reaction is very dangerous though and easily leads to explosion. A nice experiment, demonstrating this is the following:
- Take 100 mg or so of solid KClO3
- Add a single large drop of conc. H2SO4
You will see formation of orange material. This is ClO2 at high concentration, dissolved in an acid mix. You also see some fumes, which are coming from HCl, escaping the mix and (invisibly) you get HClO4, mixed with the H2SO4. Any water formed in the reaction remains dissolved in the H2SO4 and HClO4.
Adding a small piece of paper leads to instant ignition of the paper and/or explosion of the droplet of acid mix and ClO2.

NEVER scale up the above experiment. It is safe with the suggested amounts, provided you have eye protection.

AirCowPeaCock - 9-3-2012 at 07:58

I was referring to the oxidation of HOCl with O3 that Anders mentioned. Such as HOCl + O3 -> HClO2 + O2 and then HClO2 + O3 -> HClO3 + O2. And I thought it would be relevant because HClO2 is unstable and it decompose to, would decompose by disproportionation to HOCl and HClO3, wouldn't it?

Formatik - 10-3-2012 at 21:11

The precipitation of potassium through using very simple fluoro acids has also been discussed on a German forum:,204084,.html#204084 But it could end up using quite a few reagents. The paper referenced in the post describes the nature of perchloric acid preparation from chloric acid, namely the obtained chloric acid solution is just boiled down to get the azeotropic acid after necessary filtrations, and is then distilled and further purified.

woelen - 12-3-2012 at 23:17

HClO2 mainly disproportionates to ClO2, H(+) and Cl(-). The Cl(-) formed in this reaction in turn catalyses further disproportionation. This disproportionation can go so fast that the liquid starts bubbling and bubbles of ClO2 escape from the liquid.

Equation: 5HClO2 --> 4ClO2 + 2H2O + H(+) + Cl(-)

ClO2(-), ClO2 and HClO2 are quite remarkable oxidizers. They in some way are sluggish in their reactions to form Cl2, much unlike hypochlorite.

HOCl + H(+) + Cl(-) gives almost complete reaction to H2O + Cl2
HClO2 + 3H(+) + 3Cl(-) gives only marginal amounts of 2H2O + 2Cl2, instead you get ClO2 and more chloride.

AndersHoveland - 13-3-2012 at 00:09

Quote: Originally posted by woelen  
HClO2 + 3H(+) + 3Cl(-) gives only marginal amounts of 2H2O + 2Cl2, instead you get ClO2 and more chloride.

It depends on how much acid and chloride ions are present.
If excess hydrochloric acid is added to a dilute solution of sodium chlorite NaClO2, essentially only chlorine will be produced.

These type of reactions involve complicated equilibriums, and can be very dependant on reactant ratios and pH.

(4)NaClO2 + (8)HCl --> (4)NaCl + (4)H2O + (3)Cl2 + (2)ClO2

(2)NaClO2 + (8)HCl --> NaCl + (2)NaCl + (4)H2O + (4)Cl2

Chlorine is interestingly a stronger oxidizer than chlorine dioxide, but ClO2 can oxidize HCl. This is not really so much of a paradox because the presence of additional hydrogen ions alters the reduction potentials of ClO2's oxidizing strength, similar to many other reactions.

(2)NaClO2 + Cl2 --> (2)NaCl + (2)ClO2

(2)ClO2 + (8)HCl --> (5)Cl2 + (4)H2O

In the absence of additional hydrogen ions, chlorine dioxide (0.96 v) is a weaker oxidizer than chlorine (1.36 v).

[Edited on 13-3-2012 by AndersHoveland]

platedish29 - 17-10-2012 at 00:29

Quote: Originally posted by disulfideprotein  
I will say that it is explosive in an edit but it does not always decompose when you make it, you could just make it at a dilution 30-40 percent is fine. Also it has to be further heated after the dilution passes this range. What I said is not bullshit but it is certainly not a safe way to make this, the safer way would be to use barium chlorate in this reaction Ba(ClO3)2 + H2SO4 → 2HClO3 + BaSO4

Why Barium? Why not simply Calcium?

Just making ammonium chlorate could be the most suitable form to work with (ammonia is displaced after another base is added to ammonium chlorate).


wrote by Woelen:
Making HClO3 from HOCl is utter nonsense. HOCl will decompose to mainly HCl, Cl2 and H2O when it is heated. HOCl is very unstable. Hypochlorite indeed can be converted to chlorate, but this reaction is very pH-sensitive

Hows that pH senstive for the tradional disproportionation of bleach too? I mean, the reaction
3 NaClO --> 2NaCl + NaClO3 is somehow affected by pH changes?
Hardware sodium hypochlorite is very basic and contains a lot of NaCl, probably to preserve it. Also, the sodium chlorate extract seems to be stabilized all the way through the final dry salt because of this excess hydroxide they put in.

AJKOER - 17-10-2012 at 10:17

Here is a route to chlorate from HOCl via Silver. To quote from "A comprehensive treatise on inorganic and theoretical chemistry", Volume 2, by Joseph William Mellor, page 271:

"L. N. Vauquelin found that when chlorine acts on silver oxide diffused in water, a mixture of silver chloride and chlorate is formed, and it was accordingly supposed that these same products are obtained when chlorine acts on the salts of silver. The products observed by L. N. Vauquelin were shown by A. J. Balard to be end products, being preceded by the formation of silver hypochlorite. According to A. J. Balard, finely divided silver immediately decomposes hypochlorous acid with the evolution of oxygen, and the formation of silver chloride. Again, if an alkali hypochlorite be treated with silver nitrate, or if silver oxide, Ag20, suspended in water, be treated with chlorine, much heat is developed and silver chloride and silver peroxide are precipitated while a liquid with bleaching properties is formed. The liquid is very unstable, and decomposes in a few minutes with the separation of silver chloride and the formation of a soln. of silver chlorate which does not bleach. If an excess of chlorine be employed, all the silver is precipitated and a soln. of hypochlorous and chloric acids remains. If an alkali hydroxide be added to the bleaching liquid, oxygen is evolved, and a mixture of silver chloride and peroxide is precipitated. Similar results are obtained if an aq. soln. of silver chlorate, nitrate, or acetate be employed except that the corresponding acid is liberated. J. S. Stas has •shown that probably no chloric acid or silver chlorate is formed in the primary reaction:


If the silver oxide or carbonate be in excess, the silver oxide gradually forms silver hypochlorite,

2HOCl + Ag20 =2AgOCl+H20

which is readily soluble, and the soln. is stable so long as it is shaken with the excess of silver oxide present. This salt is partially decomposed on standing in darkness, and completely decomposed at 60° into silver chloride and chlorate:


and the latter remains in soln. in the alkaline liquid. J. S. Stas found no signs of the formation of perchloric acid. F. Raschig prepared silver hypochlorite by the action of alkaline sodium hypochlorite on silver nitrate, and also by adding a soln. of silver nitrate to sodium azide, NaN3, or to a soln. of chloroazide in sodium hydroxide."


So the formation of Silver Chlorate is apparently not slow and one can actually visibly observe the formation of the AgCl. I suspect that AgClO3 should be quickly be employed in a reaction, or dried per the directions, as the aqueous solution is most likely not stable. In fact, Mellor (cited previously), page 340, notes that AgClO3 decomposes into AgCl and O2 in the presence of HCl, HNO3 and even Acetic acid. Chlorine also reacts forming AgCl, HClO3 and O2.

Also per another source "The principles of chemistry", Volume 2, by Dmitrïi Ivanovich Mendelieev, the bottom of page 403, to quote:

"This is how he describes the phenomenon which then takes place: if silver oxide or carbonate be suspended in water, and an excess of water saturated with chlorine be added, then all the silver is converted into chloride, just as is the case with oxide or carbonate of mercury, and the water then contains, besides the excess of chlorine, only pure hypochlorous acid without the least trace of chloric or chlorous acid. If a stream of chlorine be passed into water containing an excess of silver oxide or silver carbonate, while the liquid is continually shaken, then the reaction is the same as the preceding; silver chloride and hypochlorous acid are formed. But this acid does not long remain in a free state; it gradually acts on the silver oxide and gives silver hypochlorite. If, after some time, the current of chlorine is stopped but the shaking is continued, then the liquid loses its characteristic odour of hypochlorous acid, while preserving its energetic decolorising property, because the silver hypochlorite which is formed is easily soluble in water. In the presence of an excess of silver oxide this salt can be kept for several days without decomposition, but it is exceedingly unstable when there is not an excess of silver oxide or carbonate present. So long as the solution of silver hypochlorite is shaken up with the silver oxide, it preserves its transparency and bleaching property, but directly it is allowed to stand, and the silver oxide settles, it becomes rapidly cloudy and deposits large flakes of silver chloride, so that the black silver oxide which hud settled becomes covered with the white precipitate. The liquid then loses its bleaching properties, and contains silver chlorato iu solution,...."

And from the bottom of page 404:

" In this manner the reactions whiih are consecutively accomplished may be expressed by the equations:

6Cl2 + 3Ag2O + 3H2O = 6AgCl + 6HCl0; 6HCl0 + 8Ag20 = 3 H20 + 6AgCIO;
6AgClO = 4AgCl + 2AgCl03.

Hence, Slas gives the following method for the preparation of silver chlorate. A slow current of chlorine is caused to act on oxide of silver, suspended in water, which is put into a state of continuous movement. The shaking is continued after the supply of chlorine has been stopped, in order that the free hypochlorous acid should pass into silver hypochlorite, and the resultant solution of the hypochlorite is drawn off from the sediment of the excess of silver oxide. This solution decomposes spontaneously into silver chloride and chlorate. The pure silver chlorate, AgClO3, does not change under the action of light. The salt is made ready for further use by drying it in dry air at 150°. It is necessary during drying to prevent the access of any organic matter; this is done by filtering the air through cotton wool, and passing it over a layer of red-hot copper oxide."


So either the presence of free Chlorine in the bleach (which effectively lows the pH by forming HCl and also HOCl), or by the direct addition of a weak acid (like CO2 or vinegar) forming HOCl, which reacts directly with Ag2O or AgCO3 to create the highly unstable silver hypochlorite. Then, AgClO quickly disproportionates into AgCl and AgClO3 unless there is an excess of Ag2O or AgCO3 to promote the stability of the Silver hypochlorite.

I had added NaClO to Silver acetate dissolved in Acetic acid. The bleach and weak acid formed HOCl, which then reacted with Silver acetate to form a white precipitate of AgCl. The solution also bubbled for a few weeks (O2 from the decomposition of AgClO3).

Note, only 1/3 of the Silver metal is converted into AgClO3, so this is an expensive route, but you may have some Silver literally laying around.

platedish29 - 17-10-2012 at 12:16

It is too dangerous a procedure! Before preparing chloric acid ones must first ask experienced people about it!

Dilute Chloric acid Preparation

AJKOER - 18-10-2012 at 18:36

Per Mellor (see page 300 at ), originally R. Bottger treated aqueous NaClO3 with Oxalic acid (H2C2O4) to form sparingly soluble Na2C2O4 and HClO3:

2 NaClO3 + H2C2O4 <--> Na2C2O4 (s) + 2 HClO3

Note, one can prepare NaClO3 by reacting Hypochlorous (HOCl) with an excess of NaOCl. The reaction is cited in Mellor also (see page 299) with excess NaOCl in a slightly acidified solution at 70 C:

NaOCl + 2 HOCl <--> NaClO3 + 2 HCl

The created HCl reacts with the excess Sodium hypochlorite driving the reaction to the right as:

HCl + NaOCl --> NaCl + HOCl

produces more HOCl that is consumed until all the NaOCl is gone. More recent work actually indicates that the reaction is two steps with the intermediate rate determining step involving the formation of NaClO2:

NaOCl + HOCl <--> NaClO2 + HCl
NaClO2 + HOCl <--> NaClO3 + HCl

However, I would avoid using an excess of NaOCl in the current context as the intent is to subsequently form a stable HClO3 (discussed more below).

Hypochlorous acid can be prepared in many ways including treating NaOCl with a weak acid (like acetic) and distilling off half of the solution which consists of most of the volatile HOCl. Pure NaOCl (free of NaCl found in commercial chlorine bleach) can be prepared by add NaOH to the Hypochlorous acid. Note, if we have NaCl present then the reaction with the Oxalic acid produces HCl as well and, as noted by Mellor and others, Chloric acid is not stable in the presence of Cl2 or HOCl or HCl (see, for example, Mellor referenced above page 288 and more recent work "Kinetics and Mechanism of the Decomposition of Chlorous Acid" from J. Phys. Chem. A 2003, 107, pages 6966-6973 at ), which is the reason for avoiding excess NaOCl.

In addition, one should avoid an excess of H2C2O4 as it can also reduce the newly formed HClO3 (see Mellor page 310).

[Edited on 19-10-2012 by AJKOER]

AJKOER - 17-2-2013 at 19:53

Here is an interesting comment from the "Kinetics and Mechanism of the Decomposition of Chlorous Acid" from J. Phys. Chem. A 2003, 107, pages 6966-6973 at ), to quote:

"Several groups of investigators[5-7] have found
that in the absence of chloride ion the stoichiometry of the
decomposition of chlorous acid is given by reaction A:

4HClO2 --> 2ClO2 + ClO3- + Cl- + 2H+ + H2O (A)

The stoichiometry of the decomposition of chlorous acid in the
presence of chloride ion is given by reaction B:

5HClO2 --> 4ClO2 + Cl- + H+ + 2H2O (B) "

Also, to quote:

"Earlier studies,[9,18] in agreement with our present results, have also found the formation of more chlorate than predicted from reaction A. Reaction C

3HClO2 --> 2ClO3- + Cl- + 3H+ (C)

also plays a role in determining the stoichiometry at higher HClO2 concentrations."

Now, I have seen the following stated reaction (see ) which is apparently appropriate in high HClO2 concentrations in the absence of chlorides, per the prior quoted research:

6 ClO2 + 3 H2O → 3 HClO2 + 3 HClO3 → HCl + 5 HClO3

So if one places ClO2 gas in contact with water for a sufficient period, the reaction may move with time to the right to produce significant Chloric acid. However, there is a safety issue to quote from Wikipedia ( ):

"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. Its solubility increases at lower temperatures: it is thus common to use chilled water (5 °C or 41 °F) when storing at concentrations above 3 grams per liter."

This suggests to me, to avoid a ClO2 explosion, and still prepare strong HClO3, that the synthesis should be repeated employing only safe levels of dilute ClO2 in air (that is, using the final HClO3 and water solution in place of the starting water) until the HClO3 concentration approaches at most 40%. Upon completion of each run, cautiously adding a little Ag2O as:

1/2 Ag2O + HCl --> AgCl (s) + 1/2 H2O

and a chloride-free strong solution of Chloric acid could be produced. I would prefer to have a small excess of Ag2O as any AgClO3 formed would decompose in weeks forming a white precipitate of AgCl and O2 gas bubbles, while any HCl present would contribute to the instability of the Chloric acid itself. Again per Wikipedia, this is important as:

HClO3 + HCl → HClO2 + HOCl

HClO3 + HClO2 → 2 ClO2 + Cl2 + 2 H2O

HOCl + HCl → Cl2 + H2O

Note, this synthesis assumes the prior generation of ClO2 (discussed in at least one prior thread and Wikipedia).

[EDIT] An alternate version of this synthesis would be to add Ag2O suspension to the water immediately. Reactions:

4 ClO2 + 2 H2O → 2 HClO2 + 2 HClO3

Ag2O + 2 HClO2 --> 2 AgClO2 (s) + H2O

Or, on net:

4 ClO2 + H2O + Ag2O --> 2 HClO3 + 2 AgClO2 (s)

The advantage of this approach is that the reaction should move more quickly to the right. Another is that the Chlorite formation can used to generate ClO2 for more HClO3 production as follows:

2 AgClO2 + 2 HCl + NaOCl → 2 ClO2 + NaCl + 2 AgCl + H2O

Note, any Silver chloride formed can also be recycle to Ag2O upon treatment with NaOH/Sugar. This means that other than the start-up chemicals (a chlorate or chlorite salt and silver oxide), the only ongoing consumables include water, HCl, NaOCl, NaOH and sugar for Chloric acid production.

The disadvantage of this proposed synthesis is the formation of Silver chlorite itself, which forms yellow unstable explosive crystals, which are impact sensitive, cannot be finely grounded and decomposing energetically (exploding) at 105° C. Note, this product may be formed inadvertently anyway from the 1st synthesis which may not have moved completely to the right. Also, both synthesis suffer from the employment of the toxic and explosive (if not properly diluted and handled) Chlorine dioxide.

[Edited on 18-2-2013 by AJKOER]