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Author: Subject: Trying to understand the electrolytic chlorate cell
A Halogenated Substance
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[*] posted on 19-3-2017 at 15:35
Trying to understand the electrolytic chlorate cell


Earlier this day, after watching two separate videos on potassium chlorate production, I was wondering why the electrolysis method was so much more efficient than the 'boil the bleach' method. I went and began researching about electrolysis to try understand it more (though this only helped slightly).

I learned about standard potential and found a reference graph for the basics of it:
http://hyperphysics.phy-astr.gsu.edu/hbase/Tables/electpot.h...

After working some of the equations out on paper to compare them of that of formation of various oxychlorions, this is my understanding...

The conversion of chloride to hypochlorite has a standard potential of -0.90 volts but hypochlorite to chlorite has a standard potential of -0.59 volts. In that way, bleach seems to be used as a bypass to skip the chloride to hypochlorite step. Then, chlorite is converted to chlorate (with a standard potential of -0.35 volts).

I suspect the 'boiling' method is ineffective because these reactions aren't spontaneous so it is difficult to set the reaction in favor of adding oxygen atoms to the chlorine based anion.

However, there are some things I'm a little confused about. On the graph from the source above, there is an additional equation of chlorate converting to perchlorate with a standard potential of -0.17 volts. This seems like a low demand compared to the other reactions. Why doesn't the chlorate get converted to perchlorate in an electrolytic chlorate cell? Is it because the standard potential isn't so low that it converts back to chlorate under standard conditions? Can a bleach eletrolyte also be used to produce chlorite and perchlorate efficiently instead of chlorate?

I'm just trying to ensure my understanding of this is correct as I would like to try it out for myself. Any comments or corrections are appreciated!

Edit: fixed spelling error

[Edited on 3-20-2017 by A Halogenated Substance]




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JJay
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[*] posted on 19-3-2017 at 16:08


I was looking at that recently to see if I could find a way to easily make potassium hypochlorite, and it looks like you've researched the topic more thoroughly than I did. I thought chlorate production was so efficient mainly because potassium hypochlorite easily oxidizes to potassium chlorate, but I didn't look at the standard potentials... that does seem to explain why industrial hypochlorite production usually involves the use of a hydroxide and chlorine.



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Zyklon-A
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[*] posted on 21-3-2017 at 09:36


It probably has more to do with price when small scale batches are desired consistantly. The boiling method is quite efficient, but it's not cheaper unless you're already making hypochlorite with products from chloralkili.
Chlorite is an intermediate oxidation state, not stable under cell conditions and I doubt it could be isolated effectively. Perchlorate is easily made with higher voltage and durable anodes. Kinetics are likely the reason perchlorate is harder to get than the SP would indicate. For example hypochorite -> chlorite has a greater - potential but the kinetics require so little activation energy that it occurs spontaneously above 60°C. Unlike Cl+1, Cl+5 is quite stable and therefore any reaction, regardless of thermodynamics, has to overcome that energy well.
Keeping chlorate molten for about an hour is sufficient energy for it to disproportionate into perchlorate and chloride, with passable yields.
No perchlorate will not be reduced to chlorate under these conditions or any others that I'm aware of. I seem to recall that perchlorate generally won't form if chloride ions are present (during electrolysis).
Thermodynamics and kinetics go hand in hand, both are required for any practical purposes.
[Edit lol let's pretend I put the oxidation # on the other side of the + symbol.:P]

[Edited on 21-3-2017 by Zyklon-A]
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