Sciencemadness Discussion Board - 2 Synthesis problems

2 Synthesis problems

Chemistry Alchemist - 25-10-2011 at 02:42

1st Synthesis
Synthesis of Sodium Thiosulfate

My first experiment i am doing is Burning Sulfur to produce Sulfur dioxide then pushing that through a solution of Sodium Hydroxide to form Sodium Sulfite which i would then dissolve more sulfur in with that to produce Sodium Thiosulfate

S + O2 = SO2
SO2 + 2NaOH = Na2SO3
S + Na2SO3 = Na2S2O3

Ive done the first main step to this (dissolving the SO2 into the sodium hydroxide) after the sulfur candle burnt out, the solution turned a slight gray colour, i let it sit for abit to settle out and there was a gray precipitate with bits of white precipitate aswell... would would this be? and how can i separate any Sodium Hydroxide from the sodium sulfite so i know how much Sulfur to add? is one soluble in a solvent leaving the other behind? or could i just continue on to the next step? would the Sodium hydroxide if any is left interfere int he rest of the procedure?

2nd Synthesis
Chlorination of hot Sodium Hydroxide

In this synthesis more then one chemical reaction is taking place at the same time... im using impure Manganese dioxide for this so it is hard to weigh out the correct amount to get the other amounts right... So im adding HCl to the manganese dioxide to form Chlorine gas which is then passed through a hot solution of Sodium Hydroxide, this forms a solution of Sodium Chloride and Sodium Chlorate aswell as water.

2MnO2 + 8HCl = 2MnCl2 + 3Cl2 + 4H2O
6NaOH + 3Cl2 = NaClO3 + 5NaCl + 3H2O

Now the solution has turned a noticeable yellow colour which im guessing is from dissolved Chlorine in the water but that shouldn't be too much of a issue...

I don't want Sodium Chlorate due to its hard to separate from the chloride so i wanna make Potassium Chlorate instead
so my question is should i boil/evaporate the solution until crystals appear b4 adding the potassium chloride?? then i just poor of the solution, discard the crystals and then make a saturated solution of Potassium Chloride and transfer the solution in to the Sodium Chlorate with out transferring the KCl crystals that hadn't dissolved... this will replace the sodium chlorate with the chloride to form potassium chlorate and sodium chloride... then i just stick the solution into the freezer, only around 3 grams of potassium chlorate is soluble at 0 degrees C so i just filter and wash the chlorate

SmashGlass - 25-10-2011 at 03:21

2nd synthesis - It would probably be better to collect the crystallized sodium
perchlorate first, even if it is contaminated with some NaCl.
Then do the reaction with KCl. You will only get small quantities, but it's
better then nothing I guess.

I'm not sure whether your method would be an improvement as you are
creating even more NaCl and it might also salt out.
Experimentation is always the ultimate proof.

An even better way to make it would be from KOH using your first method.
Wood ash contains reasonable quantities of potassium. Hydroxide and carbonate.
There are plenty of sources on-line for the conversion of K2CO3 to to KOH.

It might work better in vegemite jars too!

Good luck and happy tinkering.

Chemistry Alchemist - 25-10-2011 at 03:41

When does Perchlorate come into this? i didnt think it made perchlorate, ill let it evaporate a bit first b4 i add the potassium chloride, i have some potassium hydroxide (finally) so i could aways give it a go, i was gonna make Potassium nitrate out of it but Potassium Chlorate is a better oxidizer

ps: i dont like vegemite

bbartlog - 25-10-2011 at 07:44

Perchlorate does not come into this. Further, the initial product of bubbling chlorine into NaOH is not sodium chlorate, but sodium hypochlorite (and sodium chloride as you suggested). You've made bleach! Now, over time, it's true that sodium hypochlorite disproportionates into NaCl and NaClO3 (months, maybe years at RT... boiling would help). If you had started with KOH, then the disproportionation into KCl and KClO3 would be pretty much immediate.

Chemistry Alchemist - 25-10-2011 at 07:54

Bubbling Chlorine into Sodium Hydroxide could cause 2 things, A cold solution will yield Sodium hypochlorite while a hot solution will yield sodium chloride and sodium chlorate... the whole time i was bubbling chlorine, the solution was hot... wouldnt that mean sodium chlorate has formed?

TheNaKLaB - 25-10-2011 at 17:04

Would the gray colour formed during your synthesis of Sodium Thiosulfate be caused from the sulfur-paper wick? The wick would contaminate the solution with Carbon and other materials, wouldn't it? It really depends on how long you burnt the wick for aswell.

Chemistry Alchemist - 25-10-2011 at 19:15

this time i used just foil to hold the sullfur instead so it wouldnt be from the wick... and carbon would be black, not really gray :S

Oops! My bad

SmashGlass - 26-10-2011 at 00:02

sorry about writing perchlorate.
-I had a brain fart-
chlorate is correct.

Apologies

woelen - 26-10-2011 at 00:14

Quote: Originally posted by bbartlog

Now, over time, it's true that sodium hypochlorite disproportionates into NaCl and NaClO3 (months, maybe years at RT... boiling would help). If you had started with KOH, then the disproportionation into KCl and KClO3 would be pretty much immediate.

A solution of KOH does not have different disproportionation properties than a solution of NaOH. In both cases there is slow disproportionation of hypochlorite to chlorate and chloride at room temperature and faster disproportionation when the solution is heated.
The speed of disproportionation can be further increased if the pH is adjusted, such that most of the hypochlorite exists as HOCl instead of free OCl(-) ion. In commercial chlorate electrolysis cells the pH is constantly adjusted and this is done in order to increase the speed of disproportionation.

Chemistry Alchemist - 26-10-2011 at 00:50

What way would i need to adjust the disproportionate? More acidic or basic? and how much faster would it disproportionate?

TheNaKLaB - 27-10-2011 at 16:53

Quote: Originally posted by Chemistry Alchemist

this time i used just foil to hold the sullfur instead so it wouldnt be from the wick... and carbon would be black, not really gray :S

If it is a small amount of carbon contamination, the solution would turn grey. It seems like the only logical explanation

Chemistry Alchemist - 27-10-2011 at 19:35

So just filter it off and continue with the synthesis... would it matter if i add too much sulfur dioxide to the solution just to make sure all the sodium hydroxide reacts?

TheNaKLaB - 27-10-2011 at 22:01

Yeah filter it off. I would make sure that all the Sodium Hydroxide has reacted with the Sulfur Dioxide.

woelen - 27-10-2011 at 22:28

If too much SO2 is passed into the liquid, then you obtain a solution of sodium bisulfite and in such a solution you hardly will be able to have any sulphur dissolved. It is best to pass in SO2 until the liquid reeks of SO2 (the liquid then contains bisulfite) and then to this liquid add NaOH-solution dropwise until the pH of the liquid is slightly basic (e.g. pH equal to 9 or 10). Use pH paper for testing this.

Keep in mind that the dissolving of sulphur in a solution of Na2SO3 also is a slow process. You will need heating of the liquid and long stirring.

Chemistry Alchemist - 12-11-2011 at 00:10

So i think my Potassium Chlorate synthesis may of failed but i did get a slight precipitate so ill check that out later.

For now i took my sodium sulfite solution and filted it, i then added sulfur to the liquid and heated, it started to dissolve and the solution turned a deep read... does that mean it was dissolving into he solution? i took it off heat after about 30 min of strong heating, filtered the insoluble sulfur that didnt dissolve and now i have a deep red solution... how to i find out if Sodium thiosulfate is in? do i boil down the solution until crystals appear then put it in the freezer to squeeze out thiosulfate for a kinda pure product and then just let it evaporate to get more crystals? is the solution even ment to be deep red?

Chemistry Alchemist - 16-11-2011 at 01:51

I boiled down the Thiosulfate solution untill crystals appeared, i then took off heat and let it cool, alot of crystals formed so i began to filter (the crystals are badly impure... they are a yellowish brown colour) i then went to test the filtered solution to see if it contains any thiosulfate, so i added HCl to it... solution turned white and then a pale yellow with lots of bubbling and the distinct smell of rotten eggs. would this indicate the crystals are in fact thiosulfate?

AJKOER - 17-11-2011 at 06:44

On the preparation of KClO3, first note that the disproportionation reaction proceeds on HClO as well:

3 HClO --> 2 HCl + HClO3 [1]

Boiling is not necessary, I would heat to around 70 C for an hour (following the suggested process for NaClO3 production from NaClO). Now add K2CO3, for example, and your done.

2 HClO3 + K2CO3 -->2 KClO3 + H2O + CO2 (g)

As dilute HClO solutions are reported to be more stable, the more concentrated the HClO I would suspect the more rapid the formation of HClO3. There is also a secondary decomposition reaction (general under 20% with 80% per reaction [1], but light and certain catalysts can accelerate):

2 HClO --> O2 + 2 HCl

Note, HClO can be made by adding a weak acid like vinegar (HAc), or a very dilute mineral acid, to NaClO:

NaClO + HAc --> NaAc + HClO

Then, distill off half of the solution and you will also double the concentration of the pure HClO (most of the Cl2O comes over early in the distillation).

Note, I also believe (?) I read that HClO3 itself will further disproportionate into HClO4 (I will see if I can find the reference).

Do all work in a well ventilated area!

hissingnoise - 17-11-2011 at 07:50

Chlorates are produced by disproportionation of hypochlorite salts!
It's neither desirable nor necessary to liberate acid!
High purity NaClO3 disproportionates slowly to perchlorate and chloride when the salt is fused under controlled conditions.
Dust and organic material must be excluded, otherwise runaway decomposition will occur . . .

AndersHoveland - 19-11-2011 at 14:32

Quote: Originally posted by AJKOER

3 HClO --> 2 HCl + HClO3

Actually, a solution of hypochlorous acid will not disproportionate into chloric acid. The reaction (actually an equilibrium) is:

(8)HOCl <==> (4)H2O + (2)ClO2 + (3)Cl2

These types of reactions can be very confusing, because of the many different equilibriums, which can shift depending on pH and reactant ratios. In fact, in this reaction there will be some chloric acid in equilibrium, but its presence will not be favorable.

(2)H2O + (4)ClO2 <==> HOCl + (3)HClO3

There are many different ways this equilibrium could be written, and despite the different equations appearing with completely different reactants and products, they are all essentially describing the same reaction equilibrium.

Quote: Originally posted by AJKOER

Note, I also believe (?) I read that HClO3 itself will further disproportionate into HClO4

This is actually true. The chloric acid must be fairly concentrated (>40%) for this to procede.

Concentrated acid can disproportionate chlorate salts into perchlorate. [nitric acid should be used, not sulfuric acid which could result in an explosive reaction]
see: http://www.sciencemadness.org/talk/viewthread.php?tid=4077#p...

Quote:

Journal für praktische Chemie, Volume 23

(16)KClO3 + (12)HNO3 --> (4)KClO4 + (12)KNO3 + (6)Cl2 + (13)O2 + (6)H2O

[Edited on 19-11-2011 by AndersHoveland]

Formatik - 21-11-2011 at 00:24

Boiling aq. commercial NaClO to get get chlorate works only with difficultly. I've done it many times. I've even boiled to dryness (only in glass of course, metal gets attacked) many times and hypochlorite still lingers especially because of the hydroxide that's added to stabilize the ClO-. This hydroxide concentrates upon boiling.

I've mixed concentrated solutions of this ClO- filtrate with saturated KCl. The filtrate is very corrosive to filter paper and eats through it and weakens its structure. Even after washing formed KClO3 (crystallized from the freezer) with water several times, the filter paper shows strong weakening and crumbling. Because of strong hypochlorite presence, chlorate reduces in yield because of the many necessary water washes.

Bottom line: despite boiling, a significant amount of the hypochlorite does not decompose, and so the yields tend to be even more miserable.

A way to likely remedy this is to let the boiled mixture sit in a warm place (away from sunlight, I think I read somewhere in Gmelin sunlight decomposes ClO- to NaCl and O2) several days or weeks until it decomposes on its own. The indicator of decomposition here would be obvious loss of hypochlorite odor.

I would forget about attempting to neutralize the acid, it could need a very skilled balance, and hours in the lab of determination.

The preparation of hypochlorous acid from a chloride of lime of known concentration of hypochlorite and calcium oxide/hydroxide and very dilute mineral acids was brought up in this following hypochlorous acid thread in the first attached reference: http://www.sciencemadness.org/talk/viewthread.php?tid=12869 Too much dilute acid and the hypochlorite converted entierly to Cl2 with variations between HClO and Cl2 depending on the amount, with conversion entierly to HClO also possible.

woelen - 21-11-2011 at 01:16

AJKOER is right to some extent. If you want to make chlorate at a reasonable reaction rate from hypochlorites, then _some_ acid must be present. Strongly alkaline ClO(-) hardly disproportionates. In commercial electrolysis cells great effort is put in controlling the pH of the solution, such that part of the hypochlorite is present as HOCl and part as ClO(-). In a home-made electrolysis cell this is difficult and pH usually is too high. In an electrolysis cell this leads to loss of 1/3 of current efficiency, because no disproportionation takes place in solution and all ClO(-) must be oxidized at the anode to ClO3(-). At the right pH (around 6), the ClO(-) disproportionates in solution and no additional oxidation is needed at the anode.

If no electrolytic process is used, but simply heating of hypochlorite solutions, then one should adjest the pH to around 6 and then start heating. Too high a pH is not desirable, as that will lead to loss of Cl2 (and possibly some ClO2 as well). So, in practice the method of making chlorate from solutions of hypochlorite is less straightforward than many people think it is. A good way to adjust pH is to use boric acid. Dilute hydrochloric acid can also be used, but that should be added dropwise under vigorous stirring and realtime pH monitoring.

Chemistry Alchemist - 21-11-2011 at 21:24

Could we possibly talk about the purification of my Sodium Thiosulfate? i know i have made it but it is still a deep yellow colour... is there anyway of purifying it?

ScienceSquirrel - 22-11-2011 at 03:56

Dissolve it in water, boil with a little charcoal, filter and recrystallise.

Chemistry Alchemist - 22-11-2011 at 07:35

The only charcoal I have is from burn wood... Would this work?

ScienceSquirrel - 22-11-2011 at 07:59

Quote: Originally posted by Chemistry Alchemist

The only charcoal I have is from burn wood... Would this work?

I don't know.
Years ago we had decolourising charcoal which was finely ground willow charcoal, I think. It got the gunk out of all sorts of things.

Neil - 23-11-2011 at 06:08

Not likely, burnt charcoal losses most of its density and has its pore structure destroyed my marauding oxygen. Just pop some wood into a tin can, cover the top of the can with tin foil and have a small hole poked through the foil.

Set it next to a fire or set burning charcoal around it, as soon as it stops venting cover the hole with a bit of ash and let it cool; instant charcoal.

If you soak the wood in an activating agent before you destructively distil it; you get a better performance out of the end charcoal.

I've used home-made charcoal to soak the yellow lemon scented satan spit out of cleaning ammonia with great success.

Chemistry Alchemist - 23-11-2011 at 06:13

in other words activated carbon would be the simple explanation? when some of the water evaporates, a black precipitate forms... i dont know what it is tho...

bbartlog - 2-12-2011 at 18:37

I decided to have a try at producing potassium chlorate from commercial bleach (sodium hypochlorite).

First, I took 1240g of 6% bleach (containing one mole of sodium hypochlorite). I added 75g of sodium bicarbonate, reasoning that this would lower the pH somewhat by turning the NaOH used to stabilize the bleach into Na2CO3 (even after all of the HCO3 was driven out of solution). This I put into a glass coffee pot, covered with a watch glass (to slow evaporation, initially), and heated on a hot plate so that the temperature stayed in the range of 88-95C. I left it this way for 18 hours.

At that point the solution had lost any trace of yellowish color. I then removed the watch glass and turned up the heat to produce a rolling boil. After another five hours, the liquid volume had been reduced to about 400ml. I let the solution cool to -5C (it gets cold in my shed at night). A large mass of crystals precipitated - two different types, presumably Na2CO3 hydrate and NaCl, were apparent.
The liquid was decanted from the crystals. Only about 150ml of solution (sp gr 1.33) was obtained.
To this I added 100ml of distilled water and 25g (about 1/3 mole) of KCl. This was heated in a 500ml beaker and stirred to speed dissolution of the KCl. Once a clear solution resulted (at about 50C) I set the beaker aside, covered, to cool.
Once again this was left to cool overnightat about -5C. In the morning I decanted the liquid from the crystals that had formed, and dried them between paper towels (probably a stupid choice for KClO3, come to think of it). 20g of crystals were obtained (49% of theory). Losses are probably due to incomplete decomposition, liquid trapped in the initial mass of precipitated crystals, and KClO3 still dissolved in the solution.
No tests for purity have yet been done.
If I do this again I will probably try to use HCl to adjust the pH. I shied away from it this time, having chlorinated myself previously, but the Na2CO3 complicates things needlessly here.

AJKOER - 2-12-2011 at 21:24

First, recall that

NaClO + NaHCO3 --> Na2CO3 + HClO

or adding a weak acid (Boric acid has been mentioned) to NaClO will also liberate HClO.

Now, much of the discussion on pH control, in my opinion, could be viewed as relating to HClO production. The reason is based on the following important reactions:

NaClO + HClO --> NaClO2 + HCl

NaClO2 + HClO --> NaClO3 + HCl

Also, an important catalyst to create Chlorate is the concentration of the solution. Hence, only incidental to the act of concentrating is the prolonged boiling although not optimal from a temperature perspective.

----------------------

As a final comment, on the disputed disproportionation of concentrated HClO into HClO3 yields an equilibrium involving HClO + HClO3 (on the right) and some ClO2 (on the left), the reaction I cited is best expressed as:

3 HClO --> 2 HCl (g) + HClO3

See Wikipedia on Chloric acid "Another method is the heating of hypochlorous acid, of which products include chloric acid and hydrogen chloride:

3HClO → HClO3 + 2 HCl "

where the reference is most likely only correct (thanks) as to the products upon distillation of HClO (but I would add HClO and Cl2O to the distillation products).

In the context of aqueous solution, I agree there can be some ClO2 as:

6 ClO2 + 3 H2O <==> 5 HClO3 + HCl

where the presence of even a small amount of HCl (relative to HClO3) appears to form an equilibrium creating ClO2.

[Edited on 3-12-2011 by AJKOER]

[Edited on 3-12-2011 by AJKOER]

[Edited on 3-12-2011 by AJKOER]

bbartlog - 3-12-2011 at 05:48

Quote:

NaClO + NaHCO3 --> Na2CO3 + HClO

Not really... this being a solution we have to think about ion concentrations rather than compounds. And once you heat the solution, excess CO2 (HCO3-) will be driven off as CO2. I added the NaHCO3 reasoning that by turning all of the NaOH into Na2CO3 I would drop the pH to just a little over 11, where before it was likely around 12. pH 7 would of course be better.

Quote:

Hence, only incidental to the act of concentrating is the prolonged boiling

I don't agree. Both concentration and temperature increase the rate, but temperature does so more dramatically. See here: http://www.omegachem.com.au/docs/mega_handbook.pdf (they are trying to *avoid* chlorate formation but have some very nice tables and references). You can see in the chart they give for rate constants of decomposition that increasing concentration from 5% to 16% speeds up the reaction 4-5x, while increasing the temperature from 15C to 55C (to say nothing of 90C) speeds it up about 150x.

AJKOER - 3-12-2011 at 06:18

bbartlog: I believe we agree as I cited (per my first post) the ideal temperature (per one source from recollection) at around 70 C.

My concern is that increasing the temperature to boiling, while concentrating, could lose HClO necessary to chlorate formation.

AJKOER - 5-12-2011 at 12:14

While my discussion has focused on heating HClO solutions, I believe some may find very interesting what is reported in the literature with respect to very dilute solutions and the power of diffused sunlight and my actual observations last winter.

From "A treatise on chemistry", Volume 1 By Henry Enfield Roscoe, Carl Schorlemmer, page 192:

"Saturated chlorine water gives off chlorine freely on exposure to the air, and bleaches organic colouring matters. When exposed to direct sunlight it is, if sufficiently dilute, gradually converted into hydrochloric acid with evolution of oxygen:

Cl2 + 2H20 = 4HCl + 02.

It has been proposed to employ this reaction in measuring the chemical action of light, but the decomposition is not sufficiently regular for this purpose ; thus Pedler (1) has shown that a solution containing 1 molecule of chlorine to 64 of water undergoes no appreciable alteration during two months' exposure to tropical sunlight, whilst more dilute solutions undergo more or less decomposition, as shown in the following table

Mols. H20 for 1 mol. Cl2 / Percentage of Cl2 acting on water.

64 no action
88 29%
130 46%
140 29%
412 78%

In the case of more dilute solutions, the reaction in sunlight appears to take place almost completely in accordance with the above equation, except in so far as small quantities of chloric acid are formed. In diffused daylight, however, a considerable quantity of the latter acid is obtained, so that in this case the reactions are probably those put forward by Popper (2):

Cl2 + H20 = HCl + HClO

8HC10 = 2HCl03 + 6HCl + 02

Under certain conditions, however, sunlight brings about the reverse change, causing the formation of free chlorine from a mixture of hydrogen chloride and oxygen (see p. 200).

1 Journ. Chem. Soc. 1890, 57, 613. 1 Annalen, 1885, 227, 161 "

Now, last winter I left out in the sun (partially open to the air) some fresh dilute HClO in a thick transparent glass flower vase (diffused light?), which I further re-diluted (actually intending to discard hence the diluting to save the pumbling!). After two weeks, I noticed that the solution developed a much stronger chlorine like smell (so much for my effort to dilute!). Passing NH3 near the top produced a cloud of NH4Cl. Upon discarding the solution down the drain in an old shower, I noticed that where I splashed some on the shower floor, an intense bleaching action occurred (HClO3?). I now suspect that the diffused sunlight on the dilute HClO produced HCl and HClO3, the latter acid being so strong as to account for the smell and bleaching action. The cold temperature and dilution may have assisted in keeping gases dissolved in the solution and preserving the HClO to be acted upon by the sunlight.

Note, as the vase was covered, but not sealed, some CO2 may have dissolved into the solution. The above authors also report that when Cl2 is mixed with other gases (including CO2), the volume of Cl2 that is dissolved increases. Thus, this also is a possible factor in producing the results. I will more rigorously repeat the whole experiment this winter to confirm my observations.

[Edited on 5-12-2011 by AJKOER]

[Edited on 6-12-2011 by AJKOER]

woelen - 6-12-2011 at 00:35

Keep in mind that formation of HCl from decomposition of HClO causes formation of Cl2. In solutions, containing only HOCl the only way of forming Cl2 is by means of a complicated disproportionation reaction. In the presence of HCl, the formation of Cl2, however, is fast and easy:

HCl + HOCl --> Cl2 + H2O

This reaction is used for making Cl2 by adding hydrochloric acid to hypochlorites.

I once added dilute HNO3 to pure Ca(ClO)2. When this is done, then a pale green solution is obtained and hardly any gas is produced. When dilute HCl of similar concentration is added, then immediately there is bubbling and copious amounts of Cl2 are produced.

So, your observation can perfectly be explained. HOCl decomposes to HCl and O2. The resulting HCl then further reacts with HOCl to form Cl2. This explains the much stronger smell of Cl2.

HClO3 hardly is involved in the bleaching reactions and smell. Solutions of HClO3 are odorless and have no appreciable bleaching capabilities. They also are much more acidic than solutions of HOCl.

AJKOER - 24-12-2011 at 15:25

Quote: Originally posted by AndersHoveland

Quote: Originally posted by AJKOER

3 HClO --> 2 HCl + HClO3

Actually, a solution of hypochlorous acid will not disproportionate into chloric acid. The reaction (actually an equilibrium) is:

(8)HOCl <==> (4)H2O + (2)ClO2 + (3)Cl2

Actually, from a practical point of view AndersHoveland may be correct, as per cited source below, page 554, "A parallel pathway involving HOCl instead of ClO- is a 1000 times slower" referring to the multi-stage disproportionation reaction forming chlorate.

Also, same source, on page 553, "Photolysis of aqueous HOCl is also initiated by formation HO and Cl radicals that undergo a series of further reactions producing hydrochloric and chloric acids and oxygen (87)." Note, this appears to follow the reaction I previously posted attributed to Popper:

8HCl0 = 2HCl03 + 6HCl + 02

The reference (87) is: "A. J. Allmand, P. W. Cunliffe, and R. E. W. Maddison, J. Chem. Soc., 822 (1925); 655(1927); A. J. Allmand and W. W. Webb, Z., Phys. Chem. 131, 189 (1928); L. Bonnet, Rev. Gen. Mater. Color. 39, 29 (1935); K. W. Young and A. J. Allmand, Can. J. Res. 27B, 381 (1949); M. W. Lister, Can. J. Chem. 30, 879 (1952). "

However, it turns out perhaps most interestingly that in the case of chloride free HOCl, the disproportionation reaction of HOCl proceeds to even 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.).

Another interesting comment, same source, by the author is "Hypochlorite ion is oxidized to chlorate by ozone (142)." Page 559.

REFERENCE: "DICHLORINE MONOXIDE, HOCl, HYPOCHLORITES", Volume 8.

LINK
http://www.questscan.com/?tmp=redir_bho_bing&prt=Qstscan...

[Edited on 24-12-2011 by AJKOER]

AJKOER - 24-4-2012 at 04:59

Wrong Link:

REFERENCE: "DICHLORINE MONOXIDE, HOCl, HYPOCHLORITES", Volume 8

http://www.scribd.com/doc/30121142/Dichlorine-Monoxide-Hypoc...

or, do a google search on: DICHLORINE MONOXIDE scribd

New users may have to register with scribd which I would highly recommend given its free and the high quality of the research provided.