I have a question about purifying NaClO3 from NaCl. To maintain a high current and reduce anode wear in an electrolytic cell, I would stop the
electrolysis when the NaCl level drops to ~10g/100g H2O, with ~50g NaClO3/100ml H2O. Without water, this is 17% NaCl and 83% NaClO3 by weight.
I just studied how to interpret mutual solubility graphs. According to this mutual solubility diagram
I can purify it by,
1. heating the 1 Cl to 5 ClO3 mixture to favor ClO3 solubility. At 100C, a mutually saturated solution would have 63% ClO3 and 4.5% Cl. Assume
initially both salts are completely dissolved. Since the intitial mixture is heavier in Cl than the mutual saturation point, it would soon cross from
the phase where everything is in solution to the phase where NaCl starts crystallizing out. This would continue until it reaches the mutual saturation
point, when both would crystallize.
2. removing the cystalized NaCl and NaO3 from #1. The remaining solution is 63% ClO3 and 4.5% Cl. Cooling the solution would favor a higher Cl to ClO3
ratio. Say I cool it to 20C. Then according to the diagram, 43% would be dissolved ClO3, 20% would be solid ClO3, and 4.5% as dissolved Cl.
3. Collect the pure NaClO3 crystals from #2. The solution would now be 54% ClO3 and 5.6% Cl. This 9.6 ratio is lower than the peak ratio of 14:1 in a
100c solution, so we can now repeat steps 1-3 to purify the remaining fraction.
Is my understanding correct? If it is, I think this is an aweful lot of work to do to get 100% NaClO3 since the yield from each fraction is < 20%.
Is there a better way?
I could tolerate a few % Cl. I was thinking of simplifying it to just #1, except at a high pressure and thus higher boiling point which could bump the
ClO3 to Cl ratio to beyond 14 ! Will it work?
[Edited on 26-1-2015 by UncleJoe1985]blogfast25 - 26-1-2015 at 16:46
Is my understanding correct? If it is, I think this is an aweful lot of work to do to get 100% NaClO3 since the yield from each fraction is < 20%.
Is there a better way?
[Edited on 26-1-2015 by UncleJoe1985]
I haven't followed your reasoning fully and don't know the reliability (or origin) of your data.
One thing's for sure: when two components have similar temperature/temperature dependencies then fractionated crystallisation is a pain in the
backside, requiring a cascaded design.
Consider separating the chlorate from the chloride by converting the former to the very poorly soluble (at 0 C) KClO3.
KClO3 has a very significant solubility/temperature dependence and thus is suitable for thermal recrystallization, post chloride/chlorate separation.
[Edited on 27-1-2015 by blogfast25]UncleJoe1985 - 27-1-2015 at 13:17
I should've done more research before asking. I've found an excellent solution that's much simpler.
My interpretation of the mutual solubility diagram seems to be a little off (i.e. in a system where not everything dissolves, the way to find the %
dissolved isn't to trace a line to the saturation curve parallel to an axis, but at an angle?)
But the principles are the same as in my earlier solution:
1. Use high temperature to favor ClO3 solubility over Cl. Then use this to build up the ClO3 concentration. No extra labor required - it happens in
the chlorate cell already.
2. Chill to low temperature. This lowers the ClO3 to Cl ratio at saturation, causing the extra ClO3 in #1 to precipitate out. Also add additional NaCl
to kick more ClO3 out of solution (intuitively makes sense, but what's this principle called?)
The solution I found is very simple because it's all done in situ without needing to take electrolyte out of the cell! Conceptually it's still a
fractional process since you can't take all the ClO3 out of solution in 1 step, but you don't need a bunch of physically cascaded stages (you would
if you want to make it run in a continuous manner). The remaining fraction is processed by the cell in-place by continuing the electrolysis.
I'm a bit hesitant about adding NaCl to kick out more NaClO3. It could double the yield of ClO3 from 20g/100g H2O to 40g/100g H2O, but if I add too
much, then the precipitate will end up with NaCl too.
To know if I added the right amount, I will need a simple test for measuring the ClO3 fraction.
I was thinking about using a flame to decompose the chlorate to chloride and measure the difference in weight. The smallest Cl fraction I care about
is 1%, so it seems a scale with 3 figure precision will be good enough. Has any one else tried this method or know of something better? jock88 - 27-1-2015 at 15:02
If you want solid sodium chlorate the best was is to keep adding salt solution to your sodium chlorate cell at a close to the amount of NaCl that is
being converted and sodium chlorate solid will start to appear on the cell bottom once the liquor has got 'full' of chlorate. This may take longer
than you think if you are impatent.Molecular Manipulations - 27-1-2015 at 15:10
I keep adding sodium chloride as the cell runs until sodium chlorate precipitates. Then I cool it to 0°C filter as many crystals as possible. Then
run the solution again and re-saturate with sodium chloride.
This works because sodium chloride's solubility only decreases about one gram/100mL from 100°C to 0°C.UncleJoe1985 - 27-1-2015 at 15:46
Quote:
keep adding sodium chloride as the cell runs until sodium chlorate precipitates
Right, you want to add enough so that it reaches the mutual saturation point (roughly 12% by weight at 20C).
But I would need to know how much Cl I already have in the solution so that I don't over add. Is there an easy way to know that? I was thinking of
using the cell current as an indicator, but have no experience. Does anyone have a rough idea of the cell current as a function of Cl concentration?
Quote:
How to interprete mutual sol. diagrams is here
That's the one I studied. But I'm still a bit confused on how to get the % dissolved when there's a solid phase present.
[Edited on 27-1-2015 by UncleJoe1985]Molecular Manipulations - 27-1-2015 at 15:53
Just pour the solution in a beaker, while boiling add salt until no more dissolves, filter off the excess and put the solution back in the cell.jock88 - 28-1-2015 at 08:57
I actually give you the wrong link but you may have the file anyways.
It's attached just to be sure.
The easiest way to go IMO is to keep adding NaCl solution.
You should have some clue as to current efficiency. About 50% of no pH control and about 80 or so if using pH control.
You can tell how much NaCl used up from the current efficiency + charged passed into cell (current by time).
Wait until NaCl is (roughly) 140 grams per 100ml solution.
Add water or perhaps dilute NaCl solution until this point. Start adding NaCl solution to keep cell topped up. Keep this going and solid chlorate will
start to appear on bottom.
You need to have evaporation from cell to allow this to happen. Mind the dreaded rust-everything-in-sight mist that may come from the cell.
The Chlorate will be at the point of saturation (because it is at a high concentration) and it is it what will keep coming out of solution. The NaCl
is not close to its point of saturation.
Cooling the cell etc only upsets things and there is no point as you will not make a single molecule more chlorate by doing this. It only placates the
impatient!!!!!!!!
Think of the cell as a long oil pipeline going from Alaska to new York. You have to pump for weeks before you see a single drop of oil at the end
(bottom of cell). Thats the way it is.
You can actually get sodium chlorate to come out of solution by adding solid salt to the solution and stirring and stirring until the salt actually
'substitutes' for the chlorate in the saturation picture buts its a pain to do with beakers / bucket cells etc. Its done in industry etc.
If you need to see how much Chloride set up a titration. It is dead easy to do. Seems like a lot of bother but once you have the solutions made up its
a few minutes job.
Attachment: mut_sol.zip (279kB) This file has been downloaded 535 times
[Edited on 28-1-2015 by jock88]blogfast25 - 28-1-2015 at 12:40
If you need to see how much Chloride set up a titration. It is dead easy to do. Seems like a lot of bother but once you have the solutions made up its
a few minutes job.
Mohr's titration of chloride with silver nitrate, use potassium chromate as indicator. Buffer sample to close to pH = 7.
Consider also density as a measure of chloride + chlorate content.
[Edited on 28-1-2015 by blogfast25]jock88 - 28-1-2015 at 15:53
How do you buffer the sample to close to pH 7blogfast25 - 28-1-2015 at 16:25
There's a ton of recipes for pH 7 - 8 buffers. Just make sure they don't contain any anions that interfere with Ag<sup>+</sup>, like
phosphates, halides and such like.
But it all depends on the pH of the electrolyte (NaCl/NaClO3 solution). If that solution is fairly neutral, you're good to go, without buffering.
If you do need to buffer, say you titrate 10.0 ml of the electrolyte then add 25 ml of a 3 M pH 7 - 8 buffer to fix pH at the needed level.
Bear in mind also the following. Your electrolyte solution is likely to be far too concentrated for direct titration. So you will need to dilute it to
about 0.05 to 0.1 M in Cl<sup>-</sup>. That dilution alone may bring the sample solution into the right pH range.
What is the expected range of NaCl (say in %) of the electrolyte solution during the operation of a chlorate cell?
