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Author: Subject: Mass Oxygen Generation/Generators
MrTechGuy1995
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[*] posted on 9-3-2012 at 15:37
Mass Oxygen Generation/Generators


Well After buying an Oxygen Cylinder at my local hardware store. I was planning on using it for my NOx Generator, to increase production. For $8-9 it seemed to be a great deal at the time.

However.... The whole thing lasted a good 2 minutes, Not even constant flow, before it ran out Completely....
And it wasn't even worth the money in the end.

Knowing this, I'll need a decent way of generating Oxygen.


Moving to some thoughts....

These are the only ones I know from heart:
-Electrolysis
-H202 Decomp

How ever these I have tryed in the past, and they weren't as efficient as the flow from Cylinder. I'm at a true loss here as to how to properly generate O2 for a fair price.

Do you guys have any suggestions?

I'm going to post a link to my Apparatus for generating NOx some other time. After I find a suitable source of O2.
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weiming1998
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[*] posted on 9-3-2012 at 16:01


Fill a heat-resistant container with a bit of Ca(ClO)2, used to chlorinate pools. Heat that to slightly over 100 degrees celsius. It will rapidly decompose and generate oxygen. But if I was you, I'd use a small amount of it at one time, because large amounts will probably rupture the container as it decomposes very rapidly.

If you want a slower flow of oxygen and a safer way of generating it, then heat either a chlorate, a perchlorate, a nitrate or a permanganate in the presence of MnO2. The container required to be heated at higher temperatures though.
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[*] posted on 9-3-2012 at 16:20


Quote: Originally posted by weiming1998  
Fill a heat-resistant container with a bit of Ca(ClO)2, used to chlorinate pools. Heat that to slightly over 100 degrees celsius. It will rapidly decompose and generate oxygen. But if I was you, I'd use a small amount of it at one time, because large amounts will probably rupture the container as it decomposes very rapidly.



If that at all decomps into a small amount of Chlorine, The apparatus will certainly produce various Chlorine Oxides, which are know to be quite explosive.

I will certainly look into this.

[Edited on 10-3-2012 by MrTechGuy1995]
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[*] posted on 9-3-2012 at 16:21




Oxygen concentrators are available if that's an option. They are used by outpatients who need enriched air at home for lung problems, rather than constantly dragging cylinders around and having to maintain their charge. There are even portable versions.

The concentrators use zeolite (which is also used in kitty litter and washing powder) filters. The zeolite is able to separate the oxygen and nitrogen from the air via it's pores. The filter is regenerated by placing it under negative pressure. There is usually more than one filter, such that one is being regenerated whilst the other is under positive pressure. The reverse is also available, nitrogen concentrators.

The genuine (certified as medically sound) generators are expensive, as are the high purity nitrogen versions for labs.

However, China comes to the rescue. This one will supposedly give up to 90% O2.

Those disposable oxygen cylinders are not cheap per unit of gas. You pay a premium on the bottle going in the bin. If you are using enough to drain one in 2 minutes, you need to look at commercial (big welding) cylinders. As you're in the US, try calling around to see if the local places will refill privately owned cylinders at the yard. Buying the cylinder, rather than renting it, works out cheaper if you plan on keeping it for more than two or three years.

In the long run, I doubt you will be able to produce the oxygen yourself at a competitive price, relative to the big cylinders, once you factor in the equipment and effort required. If it were otherwise, shipyards and hospitals would be generating it on site, but they both have the cylinder people drop it off instead.

A big tank of LOX sat outside of a hospital in London:


This is from the Champs Flameoff (glassblowing) competition in the US. The big silver tanks are the oxygen. The answer to his question is, "lots". The rest is mainly bongs, but it's still very impressive. I expect a lot of these guys are as good as, or better than, scientific glass blowers (some of them may even be scientific glass blowers).
<iframe sandbox width="640" height="360" src="http://www.youtube.com/embed/ar3QofSFdBk#t=00m50s" frameborder="0" allowfullscreen></iframe>

[Edited on 10-3-2012 by peach]




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[*] posted on 9-3-2012 at 16:21


The cheapest source is oxygen bottle for welding. Initial costs are pretty high, but if you plan to use it a lot, it's the best and cheapest choice.
Electrolysis is the second cheapest, but requires sturdy electrodes, car battery charger (2 of them, preferably) and a system with a stopcock and water giving pressure to the developed gas, and of course drying tube. You can forget about carbon rods.
BTW some ozone is inevitable at great currents, but that's not important in your case.


peach, gotta love those hospital LOX supplies. :)

[Edited on 10-3-2012 by Endimion17]




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weiming1998
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[*] posted on 9-3-2012 at 17:14


Quote: Originally posted by MrTechGuy1995  
Quote: Originally posted by weiming1998  
Fill a heat-resistant container with a bit of Ca(ClO)2, used to chlorinate pools. Heat that to slightly over 100 degrees celsius. It will rapidly decompose and generate oxygen. But if I was you, I'd use a small amount of it at one time, because large amounts will probably rupture the container as it decomposes very rapidly.



If that at all decomps into a small amount of Chlorine, The apparatus will certainly produce various Chlorine Oxides, which are know to be quite explosive.

I will certainly look into this.

[Edited on 10-3-2012 by MrTechGuy1995]


If your Ca(ClO)2 and the container is dry(try put some in a dessicator bag just to make sure), then chlorine production will be minimal. Also, the reaction between chlorine and oxygen is not favoured anyway, as chlorine oxides decompose into chlorine and oxygen in the first place! So very small amounts of chlorine (from atmospheric humidity reacting with calcium hypochlorite) will generate hardly any chlorine oxides.

Edit: Ca(ClO)2 by itself will not decompose into chlorine. Only when a suitable acid is used (H2O counts), does it liberate Cl2.

[Edited on 10-3-2012 by weiming1998]
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[*] posted on 10-3-2012 at 05:22


Quote: Originally posted by MrTechGuy1995  
Well After buying an Oxygen Cylinder at my local hardware store. I was planning on using it for my NOx Generator, to increase production.


Your NOx generator? Care to explain?




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[*] posted on 10-3-2012 at 05:34


Quote: Originally posted by blogfast25  
Your NOx generator? Care to explain?


I think he means this: Air contains 78% N2 and 21% O2. When he makes NO2, there is not enough O2 for all the nitrogen to react. He wants higher yields.




Rest In Pieces!
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[*] posted on 10-3-2012 at 05:48


OK, you are going to need a lot of Ca(ClO)2 and the issue, in my opinion is safety over storage. The reaction I would worry about is between moist air containing CO2 and the Bleaching Powder (Ca(ClO)2.CaCl2.xH2O). The reaction of my concern:

Ca(ClO)2 + H2O + CO2 --> CaCO3 + Cl2O + H2O

Where I am in agreement with Wikipedia (http://en.wikipedia.org/wiki/Calcium_hypochlorite) as normally in solution:

Ca(ClO)2 + H2CO3 --> CaCO3 (s) + 2 HOCl

but with just moist air, I would speculate that very concentrated unstable HOCl forms and liberates some of its Cl2O. Normally an explosion hazard and more toxic than Cl2, but the diChlorine Mono-oxide may be converting in Chlorine (see http://onlinelibrary.wiley.com/doi/10.1111/j.1478-4408.1911.... )

On how this occurs, some opinion is as a direct reaction between HOCl and CaCl2.H2O forming Cl2, or I think a possible limited hydrolysis of the CaCl2.xH2O (increasing with heat) forming HCl and a reaction with the forming HOCl:

HCl + HOCl ---> Cl2 (g) + H2O

which adds water adding to the hydrolysis reaction and further increasing the Chlorine generation.

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

I would not heat the Ca(ClO)2, but instead try dripping an inexpensive source of H2O2 onto the Ca(ClO)2 as:

Ca(ClO)2 + 2 H2O2 --> CaCl2 + 2 H2O + 2 O2

[Edited on 10-3-2012 by AJKOER]
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[*] posted on 10-3-2012 at 05:59






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weiming1998
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[*] posted on 10-3-2012 at 06:31


Quote: Originally posted by AJKOER  
OK, you are going to need a lot of Ca(ClO)2 and the issue, in my opinion is safety over storage. The reaction I would worry about is between moist air containing CO2 and the Bleaching Powder (Ca(ClO)2.CaCl2.xH2O). The reaction of my concern:

Ca(ClO)2 + H2O + CO2 --> CaCO3 + Cl2O + H2O

Normally in solution:

Ca(ClO)2 + H2CO3 --> CaCO3 (s) + 2 HOCl

but with just moist air, very concentrated unstable HOCl forms and liberates some of its Cl2O (an explosion hazard and more toxic than Cl2).

There will also be some Chlorine. Some opinion is as a direct reaction between HOCl and CaCl2.H2O forming Cl2, or I think a possible limited hydrolysis of the CaCl2.xH2O (increasing with heat) forming HCl and a reaction with the forming HOCl:

HCl + HOCl ---> Cl2 (g) + H2O

which adds water adding to the hydrolysis reaction and further increasing the Chlorine generation.

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

I would not heat the Ca(ClO)2, but instead try dripping an inexpensive source of H2O2 onto the Ca(ClO)2 as:

Ca(ClO)2 + 2 H2O2 --> CaCl2 + 2 H2O + 2 O2


Reading Wikipedia's article on HOCl, transitional metal oxides can cause accelerated decomposition of HOCl http://en.wikipedia.org/wiki/Hypochlorous_acid
Some copper oxide sprinkled in is going to greatly reduce the chance of the formation of Cl2O, as it would probably be catalytically decomposed as well. It also provides an alterative source of oxygen, but the HCl gas liberated is going to react with the NOx when dissolved in water. So it is not good.

An interesting point is the Ca(ClO)2 is mixed with CuO (or other appropriate catalysts), and soaked in soda water, a self-decomposition reaction of the whole lot of Ca(ClO)2 might happen:
First Ca(ClO)2+H2CO3===>CaCO3+2HClO
Then 2HClO==catalyst===>2 HCl+O2
Then Ca(ClO)2+2HCl===>CaCl2+2HClO
and repeat in endless cycle, until all the Ca(ClO)2 is exhausted away.
If the HClO can decompose fast enough with the catalyst, then this reaction might be another "cold" way to generate O2, provided the hypochlorite/CuO mix is soaked in soda water. The CuO can even be recycled, by dissolving the final products in water, then filter to retrieve CuO.

But I do not know the decomposition rate of HClO boosted with catalyst, nor do I know of the reaction rate between HCl and HClO under conditions described. So, do you think this is possible under normal conditions?
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[*] posted on 10-3-2012 at 10:54


The presence of CuO in Chlorine water, I believe favors the production of HOCl:

[1] Cl2 + H2O <---> H(+) + Cl(-) + HOCl

[2] CuO + 2 HCl --> H2O + CuCl2

over the decomposition of HOCl. More HOCl per [1] means more Cl2. Now, add H2O2:

HOCl + H2O2 ---> H2O + O2 + HCl

to move reaction [1[ to the right and form O2.

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

Here is a reference from Mellor, "A comprehensive treatise on inorganic and theoretical chemistry", (Volume 2), page 46, on the action of CO2 on Bleaching powder:

"If the bleaching powder or the filtered soln. be treated with very dil. nitric, hydrochloric, or sulphuric acid, just sufficient to neutralize the free and combined lime, hypochlorous acid, HOC1, is formed : Ca(OCl)Cl+HCl=CaCl 2 +HOCl, and
the soln. smells of hypochlorous acid but not of chlorine. The hypochlorous acid can be separated by distillation. If an excess of acid be present, the hypochlorous acid is decomposed, forming water and chlorine : HOCl+HCl=H 2 0 + Cl 2 . Dry carbon dioxide has little or no action on dry bleaching powder, but with moist carbon dioxide at 70, most of the chlorine is removed, although this gas has no action on calcium chloride. According to R. L. Taylor, 13 the action of carbon dioxide on bleaching powder solid or soln. is like that of any other acid, for hypochlorous and hydrochloric acids are produced and these decompose one another with the evolution of chlorine."

LINK:
http://www.ebooksread.com/authors-eng/joseph-william-mellor/...



[Edited on 10-3-2012 by AJKOER]
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[*] posted on 10-3-2012 at 17:22


Quote: Originally posted by Endimion17  


Why use welding oxygen when you can use, erm... BLEACH, eh? :D

I guess some say it with flowers, others with bleach...




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[*] posted on 10-3-2012 at 17:41


Quote: Originally posted by AJKOER  
The presence of CuO in Chlorine water, I believe favors the production of HOCl:

[1] Cl2 + H2O <---> H(+) + Cl(-) + HOCl

[2] CuO + 2 HCl --> H2O + CuCl2

over the decomposition of HOCl. More HOCl per [1] means more Cl2. Now, add H2O2:

HOCl + H2O2 ---> H2O + O2 + HCl

to move reaction [1[ to the right and form O2.

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

Here is a reference from Mellor, "A comprehensive treatise on inorganic and theoretical chemistry", (Volume 2), page 46, on the action of CO2 on Bleaching powder:

"If the bleaching powder or the filtered soln. be treated with very dil. nitric, hydrochloric, or sulphuric acid, just sufficient to neutralize the free and combined lime, hypochlorous acid, HOC1, is formed : Ca(OCl)Cl+HCl=CaCl 2 +HOCl, and
the soln. smells of hypochlorous acid but not of chlorine. The hypochlorous acid can be separated by distillation. If an excess of acid be present, the hypochlorous acid is decomposed, forming water and chlorine : HOCl+HCl=H 2 0 + Cl 2 . Dry carbon dioxide has little or no action on dry bleaching powder, but with moist carbon dioxide at 70, most of the chlorine is removed, although this gas has no action on calcium chloride. According to R. L. Taylor, 13 the action of carbon dioxide on bleaching powder solid or soln. is like that of any other acid, for hypochlorous and hydrochloric acids are produced and these decompose one another with the evolution of chlorine."

LINK:
http://www.ebooksread.com/authors-eng/joseph-william-mellor/...



[Edited on 10-3-2012 by AJKOER]


Would CuO react with HOCl to form CuOCl, which would be very unstable? Or is HOCl too weak as an acid?
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[*] posted on 10-3-2012 at 18:17


Quote: Originally posted by blogfast25  
Why use welding oxygen when you can use, erm... BLEACH, eh? :D

I guess some say it with flowers, others with bleach...


I really meant no harm. As soon as I saw him talking about it, I LOLd as I remembered of the aliens guy and I just had to use the meme. :D

Actually, I like bleach, too. Sodium hypochlorite solution is one of the cheapest sources for oxydation at home. 5% NaOCl and 19% HCl is the leading source of homemade chlorine in my region. :D

[Edited on 11-3-2012 by Endimion17]




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[*] posted on 10-3-2012 at 21:52


Quote: Originally posted by blogfast25  
Quote: Originally posted by MrTechGuy1995  
Well After buying an Oxygen Cylinder at my local hardware store. I was planning on using it for my NOx Generator, to increase production.


Your NOx generator? Care to explain?




After my fail of an NOx generator in a video Tutorial of mine.
( I am quite embarrassed how I poorly planned out that video)
http://www.youtube.com/watch?v=xYrl8mPg8Ag

I remade my Generator, and it has improved, but I feel I can make room for more improvements.
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[*] posted on 11-3-2012 at 22:13


Quote: Originally posted by weiming1998  
Would CuO react with HOCl to form CuOCl, which would be very unstable? Or is HOCl too weak as an acid?


To hopefully answer your question here is a dated reference from "Hand-book of chemistry", Volume 5, by Leopold Gmelin and Henry Watts,page 442:

"Hypochlorite Of Cupric Oxide, or Cupric Hypociilorite.— This salt cannot be obtained by precipitating blue vitriol with hypochlorite of lime, inasmuch as a reaction is thereby produced similar to that with zinc (p. 31). —It is formed by dissolving cupric oxide in aqueous hypochlorous acid. The solution, when distilled, gives off hypochlorous acid and probably also free oxygen, and leaves a fine green oxychloride, which, when treated with excess of cupric oxide, gives off oxygen and chlorine, and is converted into an insoluble oxychloride. (Balard.)— Cupric oxide dissolves readily in chlorine water (Chenevix); the saturated solution contains an equal number of atoms of cupric oxide and chlorine [CuCl + CuO,Cl0], and decolorizes indigo after half an hour's boiling; when distilled, it leaves hypochlorons acid. (Balard.)"

LINK:
http://books.google.com/books?ei=ugP_TqKJB6bL0QGxsv3RAg&...

Also, "Modern inorganic chemistry", by Joseph William Mellor, page 271:

"R. Chenevix notes the ready solubility of cupric oxide in chlorine water, and P. Grouvelle found that the soln. obtained by passing chlorine into water with cupric oxide in suspension possessed bleaching properties, and these were retained even after the soln. had been boiled for a quarter of an hour. A. J. Balard found that the distillation of P. Grouvelle's liquor furnished some hypochlorous acid and a green oxychloride, 3CuO.CuCl2.4fl20, was formed in the retort. A. J. Balard prepared a soln. of cupric hypochlorite by dissolving cupric hydroxide in hypochlorous acid. It is also made by the action of cupric sulphate on calcium hypochlorite. A. J. Balard found that copper filings are partially dissolved by hypochlorous acid, the soln. after standing some time contains cupric chloride, and deposits a green pulverulent cupric oxychloride."

LINK:
http://books.google.com/books?id=pNNIAAAAIAAJ&q=hypochlo...

FYI, some possible reactions of interest with HOCl and copper per my notes:

HOCl + 2 Cu --> Cu2O + HCl

2 HOCl +2 Cu --> Cu(OH)2.CuCl2

4 HOCl + 2 Cu2O --> 2 CuCl2.Cu(OH)2 + O2 (g)

2 HOCl --Cu Catalyst--> 2 HCl + O2 (g)

so Copper (II) Oxygen Chloride (which forms green crystalline needles) and oxygen gas can be created. One source Mellor. "A comprehensive treatise on inorganic and theoretical chemistry", Volume 2, page 271 (link below), confirms CuCl2.Cu(OH)2 with no mention of oxygen.

http://books.google.com/books?pg=PA275&lpg=PA275&dq=...

For decomposing HOCl, Cobalt and Nickel are cited as catalyst along with direct sunlight. See "Modern inorganic chemistry", by Joseph William Mellor, page 284, at:

http://books.google.com/books?id=pNNIAAAAIAAJ&q=hypochlo...

However, avoid diffused sunlight which has been attributed to the formation of HClO3:

3 HClO --> 2 HCl + HClO3

But not particularly efficient reaction as the new HCl triggers an HOCl decomposition:

2 HCl + 2 HOCl --> 2 H2O + 2 Cl2

or, on net:

5 HOCl ---> 2 H2O + 2 Cl2 + HClO3

so assuming no other side reactions and an initial 100% disproportionation of the HOCl (unlikely), only a 20% of the HOCl is converted in HClO3. Nevertheless, one should not overlook the possible presence/creation of even small amounts of chlorates.


[Edited on 12-3-2012 by AJKOER]
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[*] posted on 24-3-2012 at 10:53


Quote: Originally posted by AJKOER  

However, avoid diffused sunlight which has been attributed to the formation of HClO3:

3 HClO --> 2 HCl + HClO3

But not particularly efficient reaction as the new HCl triggers an HOCl decomposition:

2 HCl + 2 HOCl --> 2 H2O + 2 Cl2

or, on net:

5 HOCl ---> 2 H2O + 2 Cl2 + HClO3

so assuming no other side reactions and an initial 100% disproportionation of the HOCl (unlikely), only a 20% of the HOCl is converted in HClO3. Nevertheless, one should not overlook the possible presence/creation of even small amounts of chlorates.


I have found a Patent reference were the author notes my precise reactions (see attached pdf page 4) including the net reaction:

5 HOCl ---> 2 H2O + 2 Cl2 + HClO3

Also noted, is that the reaction

3 HOCl --> HClO3 + 2 HCl

is actually (in accord with other hypochlorite/chlorate disproportionations) a multi-stage reaction:

HOCl + HOCl --> HClO2 + HCl

HClO2 + HOCl --> HClO3 + HCl

and, as expected, relatively slow as compared to:

HCl + HOCl <---> Cl2 + H2O

which explains the net reaction's creation of Cl2 at the expense of more HClO3. What is also interesting is how this relatively inefficient reaction for the production of HClO3 is moved to the right to be a viable commercial process.

(please see page 4 of this attachment):

Attachment: WO91-03421.pdf (534kB)
This file has been downloaded 310 times

[Edited on 24-3-2012 by AJKOER]
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[*] posted on 30-3-2015 at 10:17


Some of my prior dated references should be possibly be updated based on recent water research.

For example, per Liu, C., U. von Gunten, et al. (2012). "Enhanced Bromate Formation during Chlorination of Bromide-Containing Waters in the Presence of CuO: Catalytic Disproportionation of Hypobromous Acid." Environmental Science & Technology 46(20): 11054-11061. Link: http://www.researchgate.net/publication/230827707_Enhanced_B... . To quote from the abstract:

"Bromate (BrO3-) in drinking water is traditionally seen as an ozonation byproduct from the oxidation of bromide (Br-), and its formation during chlorination is usually not significant. This study shows enhanced bromate formation during chlorination of bromide-containing waters in the presence of cupric oxide (CuO). CuO was effective to catalyze hypochlorous acid (HOCl) or hypobromous acid (HOBr) decay (e.g., at least 10(4) times enhancement for HOBr at pH 8.6 by 0.2 g L-1 CuO). Significant halate concentrations were formed from a CuO-catalyzed hypohalite disproportionation pathway. For example, the chlorate concentration was 2.7 +/- 0.2 mu M (225.5 +/- 16.7 mu g L-1) after 90 min for HOCl (C-o = 37 mu M, 2.6 mg L-1 Cl-2) in the presence of 0.2 g L-1 CuO at pH 7.6, and the bromate concentration was 6.6 +/- 0.5 mu M (844.8 +/- 64 mu g L-1) after 180 min for HOBr (C-o = 35 mu M) in the presence of 0.2 g L-1 CuO at pH 8.6. The maximum halate formation was at pHs 7.6 and 8.6 For HOC or HOBr, respectively, which are close to their corresponding plc values. In a HOCl-Br--CuO system, BrO3- formation increases with increasing CuO doses and initial HOC and Br- concentrations. A molar conversion (Br- to BrO3-) of up to (90 +/- 1)% could be achieved in the HOCl-Br--CuO system because of recycling of Br- to HOBr by HOCl, whereas the maximum BrO3- yield in HOBr CuO is only 26%. Bromate formation is initiated by the formation of a complex between CuO and HOBr/OBr-, which then reacts with HOBr to generate bromite. Bromite is further oxidized to BrO3- by a second CuO-catalyzed process. These novel findings may have implications for bromate formation during chlorination of bromide-containing drinking waters in copper pipes."

To quote from the body of the work:

"Recently, we have shown that CuO can catalyze HOCl or HOBr disproportionation to produce chloride (Cl−) and chlorate (ClO3−) or Br−and bromate (BrO3−), respectively.26,27The latter pathway (eq 1) may lead to elevated BrO3− concentrations during chlorination of bromide-containing waters.....It was proposed that CuO can enhance the reactivity of halogen-containing oxidants (e.g., HOCl, HOBr, and ClO2), thereby accelerating the disproportionation.26−28"

Also:

"RESULTS AND DISCUSSION: Chlorate Formation from HOCl Decay in the Presence of CuO. In the absence of metal oxides, the depletion of HOCl mainly occurs through slow disproportionation (eq 7),27and oxygen generation (eq 8) is minor (less than 10%).22,28"

And also:

"In agreement with previous studies, significant chlorine depletion was observed at all pH values in the presence of CuO.18,20. The chlorine decay rates decrease as the pH increases from 6.6 to 9.6.... At pH 7.6, this rate constant will be even lower.27 Therefore, the decay rate of HOCl was enhanced about 108 times.... After 30 min, the increase of chlorate is concomitant to the chlorine decay. Only traces of chlorate were detected at pH 9.6. Chlorite is a possible intermediate in chlorate formation. However, chlorite concentrations were insignificant (below 0.02 μM), which is in agreement with studies performed in the absence of CuO.27,28 Furthermore, the adsorption of chlorate on CuO is insignificant"

And most interesting with respect to optimal pH for CuO catalyzed CuO chlorate formation:

"The fitted slopes (i.e., fractions of chlorate formation over the total HOCl decay) are (18 ± 1)%, (27 ± 1)%, (17 ± 1)%, and 1% for pHs 6.6, 7.6, 8.6, and 9.6, respectively. "

References cited:
(25) Shim, J. J.; Kim, J. G. Copper corrosion in potable water distribution systems: Influence of copper products on the corrosion behavior. Mater. Lett. 2004, 58, 2002−2006.
(26) Liu, C.; von Gunten, U.; Croué, J.-P. Enhanced bromate formation during chlorination of bromide-containing waters in the presence of CuO: Catalytic disproportionation of hypobromous acid. Environ. Sci. Technol. 2012, 46 (20), 11054−11061.
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Related work to quote:

"ABSTRACT: It has been shown previously that the disproportionation of halogen-containing oxidants (e.g., HOCl, HOBr, and ClO2) is enhanced by a CuO-catalyzed process. .....Similarly, the iodate concentrations decrease as a function of time in the CuO−HOCl−IO3−or CuO−HOBr−IO3− system, and the rates increase with decreasing pH (9.6−6.6) due to the enhanced reactivity of HOCl or HOBr in the presence of CuO. It could be demonstrated that iodate is oxidized to periodate by a CuO-activated hypohalous acid, which is adsorbed on the CuO surface. No periodate was found in solution.."

Link: http://www.researchgate.net/publication/266944672_Chlorinati...

[Edited on 31-3-2015 by AJKOER]
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