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Author: Subject: Bad Science? MgCl2 (aq) + 3 O2 --> Mg(ClO3)2
AJKOER
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[*] posted on 17-3-2014 at 08:27
Bad Science? MgCl2 (aq) + 3 O2 --> Mg(ClO3)2


The equation:

MgCl2 (aq) + 3 O2 --> Mg(ClO3)2 (aq)

seems to be popular on education sites (see, for example, http://www.mychemistryclass.net/Files/Interactive%20Notebook... and https://www.google.com/url?sa=t&rct=j&q=&esrc=s&... ), but I suspect that the chemistry is misreported. There is some historical foundation, for example, see "The Manufacture of Sulphuric Acid and Alkali: Ammonia-soda, various ..." by Georg Lunge, an online googlebook. To quote from page 669:

"4. Magnesia Bleach-Liquor.
This can be made by decomposing a solution of bleachingpowder with Epsom salts and decanting from the gypsum formed. It has been proposed for bleaching by Claussen, and again by Ramsay; it is said to bleach more rapidly, and not to turn straw, flax, hemp, etc., brown, as it is free from lime; but it suffers decomposition more quickly than chloride of lime. The separated magnesia does not at all injure the fabrics (Bolley and Jokisch, Schwciz. polyt. Zeit., 1866, p. 120). Such a liquor was patented by Oliver Grantham, Sinnock, and Leverson (B. P. 2351 of 1861), without adding anything new. F. Hodges has shown that a liquor of the same kind (prepared from ordinary bleaching-powder and kieserite) can bleach linen fabrics without exposure on grass, if they have previously been steeped in a hot solution of sodium carbonate.
Balard prepared a bleaching-liquor by dissolving magnesia in aqueous hypochlorous acid; Grouvelle by passing chlorine into magnesia suspended in water.
None of these solutions had been properly investigated until Lunge and Landolt (Chem. Ind., 1855, p. 340) undertook this task. They tried first to prepare a bleaching-magnesia, analogous to bleaching-powder, by treating solid magnesium hydroxide with chlorine, either in the dry or in the moist state; but they did not obtain anything fit for use, viz., a. mass with only 0-15 available and 4-3 chloride-chlorine.
Now chlorine was passed into a milk of- magnesium hydroxide and water, at temperatures between o° and 100°. Even at 0°, together with magnesium hypochlorite, much chlorate was formed, more than corresponding to half of the chlorine entering into the reaction. At 15° a little more chlorate was formed, together with much hypochlorate, some of which was changed into chloride, with evolution of oxygen. In both solutions the hypochlorite is easily converted into chlorate, not merely by heating to 50°, but even by prolonged agitation by a current of air at ordinary temperatures. At 70° C, from the first mostly chlorate was formed, with a little chloride, produced by loss of oxygen. Hence magnesium hypochlorite in statu noscendi does not possess much stability and is easily transformed into chlorate."

Link: http://books.google.com/books?id=FnrTAAAAMAAJ&pg=PA670&a...

I do not dispute the formation of chlorate that occurs upon passing air into a hot concentration of Mg(OH)2 and Chlorine, or Mg(ClO)2. But, the more likely chemistry in my opinion, is not represented by the promulgated equation above.

In my humble opinion, it's not O2 but the CO2 in the air moving the reaction. My take:

CO2 + H2O = H2CO3

Mg(ClO)2 + H2CO3 --> MgCO3 (s) + 2 HOCl

or the soluble Magnesium bicarbonate, followed by:

Mg(ClO)2 + 2 HOCl --> Mg(ClO2)2 + 2 HCl

and the newly formed HCl attacking any formed carbonate recreating CO2 and more HOCl, and finally forming the chlorate:

Mg(ClO2)2 + 2 HOCl --> Mg(ClO3)2 + 2 HCl

Note, in my suggested path, CO2 acts in the role of a catalyst and the reason the reaction proceeds with Magnesium hypochlorite is its ascribed lack of stability.

Now, with respect to the reaction proceeding with just MgCl2 and not Mg(ClO)2, a weaker path may occur (I am doubtful) via the steps:

MgCl2 + 2 H2O = Mg(OH)2 + 2 HCl

2 HCl (dilute only) + air/ozone = 2 HOCl

the above reaction forming some Hypochlorous acid being based on the reported action of H2O2 (and monoatomic oxygen) on dilute HCl by Watts in Dictionary of Chemistry.

In any event, the reaction:

MgCl2 (aq) + 3 O2 --> Mg(ClO3)2

is at best, a complex series of reactions not accurately represented, in my opinion, by the above reaction.

However, if I am off base here, would someone please cite some research to the contrary.

Thanks.

[Edited on 17-3-2014 by AJKOER]
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[*] posted on 17-3-2014 at 08:39


Sounds reasonable, you should test your hypothesis. It would be quite easy, first test with pure oxygen, then with a mixture of carbon dioxide and oxygen. If your hypothesis is correct, the pure oxygen won't work.



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[*] posted on 17-3-2014 at 10:31


Quote: Originally posted by AJKOER  
The equation:
MgCl2 + 2 H2O = Mg(OH)2 + 2 HCl
[Edited on 17-3-2014 by AJKOER]


In what condition is the above happening?
If it is in liquid water, then the equation basically turns into

H2O <=> OH- + H+

In a acidic solutions like HCl, the number of OH- is driven down. The equilibrium is to the left side.
Mg2+ and Cl- are spectator ions.




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[*] posted on 17-3-2014 at 10:35


MgCl2: xH2O decompose to HCl and Mg(OH)2 or MgO, not sure what temperature this occurs though.



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[*] posted on 17-3-2014 at 12:24


Quote: Originally posted by AJKOER  
The equation:

MgCl2 (aq) + 3 O2 --> Mg(ClO3)2 (aq)

seems to be popular on education sites (see, for example, http://www.mychemistryclass.net/Files/Interactive%20Notebook... and https://www.google.com/url?sa=t&rct=j&q=&esrc=s&... ), but I suspect that the chemistry is misreported.
[Edited on 17-3-2014 by AJKOER]


C'mon AJ, these sources provide A-level students with examples of how to balance equations. No where does it imply these reactions are actually feasible.

Hell, one of them is the decomposition of Cs<sub>2</sub>CO<sub>3</sub>, ROFLOLPM...

Talk about barking up the wrong tree...




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[*] posted on 17-3-2014 at 13:41


What does "ROFLOLPM" mean?



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[*] posted on 17-3-2014 at 14:04


Rolling on the floor, laughing out loud, p*ss*ng myself.



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[*] posted on 17-3-2014 at 15:07


Lol that's so funny. I knew the first part, but not the "pissing myself" part.
I'll be using that sometime, I'm sure....




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[*] posted on 17-3-2014 at 17:09


Continuing on this cesium carbonate tangent:

The thermal decomposition of cesium carbonate apparently (under certain conditions) produces an n-type semiconductor of cesium superoxide doped with cesium suboxides.

http://yylab.seas.ucla.edu/papers/AFM%20Cs2CO3.pdf

About the original topic:

When you have experimental evidence, I will believe it.

[Edited on 18-3-2014 by Cheddite Cheese]




As below, so above.
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[*] posted on 17-3-2014 at 18:25


Actually, looking over my equations on how CO2 acts as a catalyst, there exists proof already. It is the very reason Chlorine bleach appears to have a short shelf life once opened. The reason being after the bottle is opened, any CO2 exposure can, with time, convert your Bleach to chlorate and chloride. This is why some bleach brands add NaOH, to capture CO2! Once this is exhausted, the Bleach will start to lose its hypochlorite. Oxygen exposure is not the issue as the Bleach had O2 exposure from the time it was created.

Now that I have inadverently explain one of the major weakness of Chlorine bleach, consumers should avoid over paying for it and chemist can strive to make it a better product.
----------------------------------------------------

Now, as to whether chemist should alert educators on bogus reactions is like requiring, in an English composition class, that one should use real words. To remain quiet is a disservice to the student and the science.

[Edited on 18-3-2014 by AJKOER]
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[*] posted on 18-3-2014 at 02:08


Quote: Originally posted by AJKOER  

Now, as to whether chemist should alert educators on bogus reactions is like requiring, in an English composition class, that one should use real words. To remain quiet is a disservice to the student and the science.

[Edited on 18-3-2014 by AJKOER]


No, it f*cking well isn't, you tw*t.

The 'words' used are REAL: it just so happens that some reactions aren't thermodynamically favourable. Not that you would know that, going by the reams of nonsense you've been posting.

Take the 'offending' reaction, for instance. It can be pointed out to students at some later date that the reaction from left to right probably isn't thermodynamically favourable but that the reaction from right to left is. I don't think you understand education processes very well.

The concept that educators would have to listen to... erm... you, is too preposterous to entertain. When will you finally take up knitting as a hobby?


[Edited on 18-3-2014 by blogfast25]




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[*] posted on 18-3-2014 at 03:14


In my second life in a parallel universe I make Mg(ClO3)2 routinely from MgCl2 and O2. I just open my container with MgCl2 and leave it standing overnight. The next morning I have Mg(ClO3)2 and I can make some quick and loud bangs in the streets again.



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[*] posted on 18-3-2014 at 07:39


Woelen:

I find your comment very interesting if you are referring to the action of air on MgCl2.nH20 crystals. That is, possibly some usual (and unexpected by me) surface chemistry taking place. The general lack of discussion in the literature is then due to the absence of true understanding of the underlying chemistry itself.

There is some truth in Blogfast comment relating to energy considerations, but surface chemistry (which can rework reaction paths energy requirements per my studies of atmospheric reactions), I find fascinating.

An important aspect may relate even as to how the MgCl2 was prepared. For example, per my research, methods include, to quote:

"FROM HYDROCHLORIC ACID ON MAGNESIUM OXIDE OR HYDROXIDE, ESPECIALLY THE LATTER WHEN PRECIPITATED FROM SEAWATER OR BRINES (GREAT SALT LAKE)

... PREPARED FROM MAGNESIUM AMMONIUM CHLORIDE HEXAHYDRATE IN PRESENCE OF HCL...

Molten magnesium chloride can be formed by the direct carbochlorination of magnesium oxide obtained from the calcination of magnesium carbonate ores or magnesium hydroxide.

Magnesium chloride can be produced in large quantities from (1) carnallite or the end brines of the potash industry; (2) magnesium hydroxide precipitated from seawater; (3) by chlorination of magnesium oxide from various sources in the presence of carbon or carbonaceous materials"

where the method may introduce impurities impacting lattice structure, catalytic and chemical properties of the resulting salt. For example, if created by evaporation/heating of a concentrated MgCl2 solution, some MgO and Mg(OH)Cl could be present as well. Also, some HCl is known to formed on hydrolysis of MgCl2 (see Mellor, page 244 at
http://books.google.com/books?id=1iQ7AQAAMAAJ&pg=PA244&a... ).

[Edited on 18-3-2014 by AJKOER]
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[*] posted on 18-3-2014 at 13:22


Quote: Originally posted by AJKOER  
Actually, looking over my equations on how CO2 acts as a catalyst, there exists proof already. It is the very reason Chlorine bleach appears to have a short shelf life once opened. The reason being after the bottle is opened, any CO2 exposure can, with time, convert your Bleach to chlorate and chloride. This is why some bleach brands add NaOH, to capture CO2! Once this is exhausted, the Bleach will start to lose its hypochlorite. Oxygen exposure is not the issue as the Bleach had O2 exposure from the time it was created.
[Edited on 18-3-2014 by AJKOER]


I doubt it. you wrote

Mg(ClO)2 + H2CO3 --> MgCO3 (s) + 2 HOCl
or the soluble Magnesium bicarbonate, followed by:
Mg(ClO)2 + 2 HOCl --> Mg(ClO2)2 + 2 HCl

I doubt that the 2nd equation takes place.

For bleach,
NaOH is present in bleach to avoid the acidity problem caused by CO2. Acidity causes hyprochlorite to decompose and forms free Cl2 which evaporates away.
In the case of aqueous Mg(ClO)2, it is going to be the same situation and the same applies to any soluble hypochlorite salt.




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[*] posted on 18-3-2014 at 14:16


Here is a link to "White's Handbook...", note Equations 9.4 and 9.5 on pages 463 and 464 (or pdf index pages 505 and 506) at http://f3.tiera.ru/3/Chemistry/Material%20Science/AppliedChe...

which are the equations governing the formation of chlorates as I gave previously.

The oxidation of the chlorite to chlorate is reputedly the slow (rate determining step) reaction. Interestingly, even electrolysis is largely governed by these reactions and in certain setups one will find holding tanks to give the reaction some time to complete as a means to increase chlorate yield!

[Edited on 18-3-2014 by AJKOER]
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[*] posted on 18-3-2014 at 17:32


Quote: Originally posted by AJKOER  
Here is a link to "White's Handbook...", note Equations 9.4 and 9.5 on pages 463 and 464 (or pdf index pages 505 and 506) at http://f3.tiera.ru/3/Chemistry/Material%20Science/AppliedChe...

which are the equations governing the formation of chlorates as I gave previously.

The oxidation of the chlorite to chlorate is reputedly the slow (rate determining step) reaction. Interestingly, even electrolysis is largely governed by these reactions and in certain setups one will find holding tanks to give the reaction some time to complete as a means to increase chlorate yield!

[Edited on 18-3-2014 by AJKOER]



Alright, that's good to know.
The text states this
OCl- + OCl- => ClO2- + Cl-
OClO2 + OCl => ClO3- + Cl-

while you wrote all of this
CO2 + H2O = H2CO3
Mg(ClO)2 + H2CO3 --> MgCO3 (s) + 2 HOCl
Mg(ClO)2 + 2 HOCl --> Mg(ClO2)2 + 2 HCl
Mg(ClO2)2 + 2 HOCl --> Mg(ClO3)2 + 2 HCl

The text is talking about how a hypochlorite solution degrades.
You are suggesting that CO2 acts as a catalyst which I don't know if it is true or not, but it has no bearing on the initial equation.

The initial equation suggests chlorine going from -1 to +5 which looks like none sense to me.
MgCl2 (aq) + 3 O2 --> Mg(ClO3)2 (aq)

And this
MgCl2 + 2 H2O = Mg(OH)2 + 2 HCl
2 HCl (dilute only) + air/ozone = 2 HOCl

I think that we can get rid of the first equation.
It turns into
2 Cl- (dilute only) + air/ozone = 2 OCl-

Unless if you think the presence of ions like Mg2+ and H+ serve some role.




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[*] posted on 18-3-2014 at 18:57


OK, in words, first, you miss the reaction where the Carbonic acid, a weak acid, attacks the hypochlorite, a salt of the even weaker Hypochlorous acid. A carbonate is formed along with HOCl which starts the reaction chain to chlorate and HCl. The latter is neutralized by either the recently formed carbonate (regenerating the CO2) or even Mg(ClO)2 (producing more HOCl which could react with free HCl forming Cl2).

Bottom line, the CO2 remains in a sealed system and moves the reaction slowly forward (hence, a catalyst).

Now, the reaction:

MgCl2 + 3 O2 --> Mg(ClO3)2

if valid for the solid salt exposed to air appears to be largely unexplained or attributed to other compounds/impurities or even special physical surface conditions.

[Edited on 19-3-2014 by AJKOER]
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[*] posted on 19-3-2014 at 03:30


Quote: Originally posted by AJKOER  

if valid for the solid salt exposed to air appears to be largely unexplained or attributed to other compounds/impurities or even special physical surface conditions.


If you say so. I don't know. I imagine that solid phase properties are going to be different.
It seemed like you were talking about solutions so I would get rid of spectator ions.

This equation
MgCl2 + 2 H2O = Mg(OH)2 + 2 HCl

is useless. It becomes
Cl- + H2O <=> OH- + HCl
but HCl isn't going to last long. It will quickly encounter some H2O that would cause dissociation
HCl => H+ + Cl-

and the H+ will encounter an OH-
H+ + OH- => H2O




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[*] posted on 19-3-2014 at 04:40


In concurrence with your comment, the hydrolysis of MgCl2 is described per my Mellor reference as "slight" as compared to CaCl2, which Mellor cites as "insignificant". However, the hydrolysis is significant if the reaction conditions are such (like boiling off the voltatile HCl) to drive the reaction to the right.

Interesting, the hydrolysis may also produce some Mg(OH)Cl. The latter in the presence of O2 and a metallic impurity (the source of the MgCl2 is possibly brine after all) may (speculation) even create an electrochemical cell forming Cl2. The action of moist Chlorine on MgO or Mg(OH)2 is a known preparation path to Mg(ClO3)2.

This thread is possibly a good educational tool as I have touched on many areas in chemistry seeking some insight as to how/if the topic's main reaction could proceed.

[Edited on 19-3-2014 by AJKOER]
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[*] posted on 19-3-2014 at 05:32


@vmelkon:

Don't feed the trolls.




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[*] posted on 19-3-2014 at 06:59


@blogfast25: Do not derail the thread in a personal quarrel. If you don't like AJKOER's posts, then please do not respond to his posts.

@vmelkon and AJKOER: Formation of HOCl from humid MgCl2 does not occur at any practical rate. It does not even occur at much more easily hydrolyzed chlorides like FeCl3.6H2O and AlCl3.6H2O. What does happen though is formation of free HCl. You can actualle smell it. The formation of the free HCl occurs in the compound itself, e.g.

FeCl3.6H2O --> Fe(OH)Cl2.5H2O + HCl

This kind of decomposition reactions occurs more easily when the chloride is heated.

I do not think, however, that this kind of reactions occurs much with MgCl2.2H2O, because in this compound, the metal-aqua complex is less acidic.

In the past I have done experiments with hydrated metal nitrates and metal chlorides and I found that quite a few can be decomposed to free acid and some basic nitrate or chloride, but earth alkali salts usually are not easily decomposed. E.g. I heated Ca(NO3)2.4H2O and CaCl2.2H2O and no or only very small amounts of free acid were produced. The same with BaCl2.2H2O and Sr(NO3)2.4H2O. I never purchased MgCl2 or one of its hydrated salts, because it is not very interesting, so I never tried this with the magnesium salt.

Hypochlorites are AJKOER's favorite and they appear over and over again in his posts, but in practice their formation only occurs in the presence of free chlorine or during electrolysis of chlorides. Air and HCl do not give free chlorine, nor HOCl in any appreciable amount and chloride without acid is even less reactive towards oxygen, so forget about that. Quantum mechanically speaking every reaction can occur (e.g. a particle which must overcome a huge potential wall still has a non-zero chance of crossing that wall), but the chance of reaction for a particular atom can be so low, that in practice the reaction never ever will be observed, or at most for one atom every few seconds or so and at that rate it will remain undetected with every, even the most sensitive, analytical methods.




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[*] posted on 19-3-2014 at 08:01


Well, I'm relatively new and don't look at each forum here so perhaps I missed all the other hypochlorite discussions.

Quote: Originally posted by woelen  
Air and HCl do not give free chlorine, nor HOCl in any appreciable amount


I was thinking about this.
Why doesn't O2 oxidize Cl-?
Is it because the bonds between O2 are difficult to break?
Is it because O needs 2 electrons and the second electron is difficult to pull away from the Cl-?
A combination of those two cases?

(I'm not talking about rare occurrences. I am talking about mass conversion.)

[Edited on 19-3-2014 by vmelkon]




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[*] posted on 19-3-2014 at 13:17


Quote: Originally posted by vmelkon  
....
I was thinking about this.
Why doesn't O2 oxidize Cl-?
Is it because the bonds between O2 are difficult to break?
Is it because O needs 2 electrons and the second electron is difficult to pull away from the Cl-?
A combination of those two cases?

(I'm not talking about rare occurrences. I am talking about mass conversion.)


OK, first O2 can be used to oxidize HCl to Cl2 in the presence of a catalyst (like CuCl2, or CuCl2 + NaCl) at a temperature of 450 C (see "Thermochemistry of the Deacon Process" at http://pubs.acs.org/doi/abs/10.1021/j100016a065 ). For recent improvements on the Deacon Process employing CuO.CuCl2, for example, see http://cfpub.epa.gov/ncer_abstracts/index.cfm/fuseaction/dis... .

Absence a catalyst, compounds like H2O2 (or acidified peroxides?) can convert concentrated HCl to a solution of effectively Chlorine water (note, Cl2 + H2O <--> HCl + HOCl) and, per a secondary reaction, evolve O2 (previously discussed on SM, see http://www.sciencemadness.org/talk/viewthread.php?tid=14490&... ). I previously cited Watt's Dictionary of Chemistry on the action of H2O2 on dilute HCl forming HOCl. This reaction is cited by Watt's as one of several preparations for Hypochlorous acid. To quote from "Watts' Dictionary of chemistry", page 16:

"—6. Addition of H2O2 Aq (containing 2.45 p.c. H2O2) to a large excess of Cl Aq produces HClOAq, according to Fairley (B. A. 1874, 57); if much H202 is added, the HClO Aq is decomposed forming HCl Aq, H20, and evolving O." Link: http://books.google.com/books?pg=PA13&dq=Watt+preparatio...

I also alluded to the action of ozone on HCl, to quote one source:

"The halide acid HCl has been shown to be effective in destroying ozone"

Source: http://www.geology.sdsu.edu/how_volcanoes_work/climate_effec...

And, per another source (see http://nix.nasa.gov/search.jsp?R=19770046873&qs=N%3D4294... ), to quote:

O3 + HCl --> O + O2 + HCl

where the HCl can subsequently be broken up by the monoatomic oxygen radical in a photochemical setting as follows:

HCl + O --> OH + Cl
HCl + O --> ClO + H
HCl + OH --> H2O + Cl
ClO + O --> Cl + O2
Cl + Cl --> Cl2

Sources: See Table I at http://books.google.com/books?id=BEImtW7ZR3oC&pg=PA257&a... and http://pubs.rsc.org/en/Content/ArticleLanding/2011/CP/c0cp02...

[Edited on 19-3-2014 by AJKOER]
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[*] posted on 20-3-2014 at 18:43


I came across an interesting article "OXIDATION OF CHLORIDE TO PERCHLORATE UNDER AMBIENT MARS CONDITIONS", by B. L. Carrier* and S. P. Kounaves, Department of Chemistry, Tufts University, Medford, MA, 02155, link: http://www.hou.usra.edu/meetings/lpsc2014/pdf/2570.pdf . The point of the article is to explain the significant amount of perchlorate formation on Mars from chloride. An important quote, in my opinion, is:

"One possible pathway is the heterogeneous reaction of soil chlorides with atmospherically produced oxidants or oxidants generated photochemically at the surface. An ongoing source of perchlorate formation would indicate the likely presence of other oxychlorine species during the oxidation of chloride to perchlorate, such as ClO-, ClO2- and ClO2(g) as well as other possible radicals such as ●OCl, ●Cl, or ●OH. The presence of these intermediates has implications for the survivability of organics on the surface due to their high reduction potentials. The current research aims to investigate the formation pathway for perchlorate on mineral surfaces under current Mars conditions. "

Also, a new thought based on an observation from the interaction/relationship between copper hypochlorite and its decomposition into copper oxychloride and oxygen back here on earth:

Cu(OH)2 + 2 HOCl = Cu(ClO)2 + 2 H2O
Cu(ClO)2 + 2 H2O + Cu(OH)2 = CuCl2 .Cu(OH)2 + 2 H2O + O2

Reference: "A comprehensive treatise on inorganic and theoretical chemistry", Volume 2, by Joseph William Mellor, page 271. To quote:

"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.CuCl2.4H20, 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."

Now, replacing Mg for Cu, a possibly similar reaction:

Mg(OH)2 + Mg(ClO)2 + 2 H2O = 2 Mg(OH)Cl + 2 H2O + O2

So, if this hypothetical reaction can be reversed(?), Magnesium chlorate formation could ensue. The catalyst could be, per the article cited above, via photochemically induced radicals, a heterogeneous mix, possibly from impurities in the MgCl2 itself, or even from bacteria (for example, see
http://link.springer.com/article/10.1007%2Fs11771-011-0856-6 ). The role of the catalyst could also be to lower the thermal barrier for the oxidation with O2 of HCl (a product of the limited hydrolysis of MgCl2 in addition to Mg(OH)Cl ). Arguments based on the particular properties of MgCl2 with regard to hydrolysis products, instability of Mg(ClO)2 with respect to disproportionation to chlorate, and potential impurities of the MgCl2 extracted from brine, build a stronger argument IMHO for the apparently indirect action on MgCl2 of oxygen allegedly unique to this magnesium salt.

As such, my opinion remains that the equation implying the direct oxidation of MgCl2 by oxygen to chlorate is misleading at best.

[EDIT] Interestingly, in support of an indirect path for the action of oxygen on MgCl2, per Bretherick’s Handbookof Reactive Chemical Hazards, SixthEdition, Vol 1, page 561, the only comment on MgCl2 was, to quote:

"A large steel evaporator used for magnesium chloride solution was shutdown for maintenance. During maintenance operations a fatality occurred from atmospheric oxygen deficiency inside the evaporator. It was found later that the oxygen content in the evaporator fell from the normal 21% to about 1% in under 24h, and this was confirmed in laboratory tests.This was attributed to very rapid rusting of the steel under warm humid conditions in the presence of traces of magnesium chloride[1]. Further work shows that other salts (calciumbromide, calciumchloride, magnesium sulfate, potassium chloride) behave similarly, and that presence of scale is a contributory factor[2]. Magnetitescale (Fe3O4) on mild steel increases the depletion rate by a factor of 10, while the rust formed during the corrosion has little effect[3]."

The rapidity with which oxygen was depleted suggests possibly galvanic corrosion, an electrochemical reaction. The products of such a reaction may provide catalyst for the reaction under question.

[Edited on 21-3-2014 by AJKOER]
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AJKOER
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[*] posted on 22-3-2014 at 06:59


A word of warning, I prepared MgCl2 from adding MgSO4 to aqueous NaClO:

MgSO4 + 2 NaClO = Na2SO4 + Mg(ClO)2

On serious cooling, Na2SO4(H2O)10, Glauber's salt, separates out, leaving largely Magnesium hypochlorite and possibly some chlorate as well given the propensity of the Mg(ClO)2 to undergoing disproportionation.

Next, add H2O2, to form aqueous MgCl2 (if one accepts the claim that agitation with air on Mg(ClO)2 produces chlorate, then the action of active oxygen released in this reaction, may similarly create some Mg(ClO3)2 as well):

Mg(ClO)2 + 2 H2O2 ----) MgCl2 + 2 H2O + O2 (g)

I let the solution in the open to evaporate overnight. Then, I proceeded to boil down the remainder in a stainless pan (to test the reaction on Fe). When most of the solution is boiled away, one starts to see a white salt in solution, MgO or possibly Mg(OH)Cl (and/or some Mg(OH)ClO3?). Then, the steel is seriously attacked. Per the hydrolysis previously detailed by Mellor, I suspect HCl, but in the presence of a concentrated and highly ionic MgCl2, the 'activity level' (or acid strength) of the HCl is likely enhanced explaining the attack on the shiny stainless steel. There could also be a small amount of HClO3 formed. Continuing with reaction, a light brown salt forms on the pan and in solution. Thank goodness it was an older pan, pretty much disfigured.

[Edited on 23-3-2014 by AJKOER]
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