Sciencemadness Discussion Board - Processing chlorate cell electrolyte

Processing chlorate cell electrolyte

frogfot - 1-4-2003 at 02:23

Hi, i got a really annoying problem. After preparing chlorate electrolytically i found it hard to remove carbon particles..
My procidure was simple, first i filtered electrolyte through a piece of compressed cotton, this gave a clear green solution. But with a flashlight test one could see the beam through the liquid, which meant that small particles was still present.
After 30 min boiling solution got gradually orange to black. Any clue why this color change happened?
After filtering this i ubtained an orange liquid. Flashlight test showed the same thing as first time though light beam was much weaker..

Is there any way one could filter electrolyte to remove all particles? I heard many documents mention diatomeous earth, is it worth bying? (though i havn't looked for it yet)
Or maby one can leave the electrolyte for a month, maby all particles will settle down?

BASF - 1-4-2003 at 12:30

Can´t help you in terms of your color change, but:

diatomeeous earth = "kieselguhr"= "infusorial earth" = "terra silicea".
It is not exactly cheap, but not really expensive.

But as you suggested, i am also sure that time would do its thing and the particles would settle.

I may be wrong, but personally i don´t think such a small amount of carbon in the final product would increase its sensitivity considerably.

Organikum - 1-4-2003 at 13:33

Quote:

diatomeeous earth = "kieselguhr"= "infusorial earth" = "terra silicea"

aka "celite". Aquarium or brewery supply.

but not perlite!

vulture - 1-4-2003 at 13:36

Did you add any dichromate by any chance?

odd color change

blazter - 1-4-2003 at 15:35

I tried this a looong time ago with NaCl, then converting it to the potassium salt. I too got the bleach green color and tons of carbon crap. For me it more or less resolved overnight and a simple filtering through medium lab filter paper did it just fine. Upon boiling, It remained rather green and I only got a few impure crystals. Never the less when ground with sugar they burnt with a long purple flame which looked interesting.

PHILOU Zrealone - 1-4-2003 at 18:05

Purple flame?With Na?

NaClO3 converted to KClO3

blazter - 1-4-2003 at 21:18

I used KCl to convert the NaClO3 to KClO3 which is MUCH less soluable and would've hopefully precipitated first. As i found it isnt that cut and dry, and some contaminants precipated along with it. But yes I beleive it was the K that gave the purpleish flame.

frogfot - 2-4-2003 at 03:32

I didn't used cromates.. but maby cathode is responcible for this, it was usuall iron. Recently i thoat that cell chamber could be responcible for strange color, they're made of polystyrene, and it should react with chlorine. Cells was transparent and had some forry spots after operation.

Well, i forgot about how expencive things are in my country.. BASF, youre maby right, afterall, product looked slightly grey, this shouldn't be a problem

Blazter, that mustve been some high quality KClO3 you've made. I heard that color from potassium is jammed very easily by traces of Na, how many times have you recrysallised it?

Actually it wasn't that pure

blazter - 2-4-2003 at 12:13

I sort of got impatient with the cell I was running and the anode was getting eaten to nothing pretty fast. IIRC I ran it for something like 60amp hours. Even then it was obvious that it wasn't a very pure product. I came to that conclusion because actual flame tests with the material showed a strong sodium color, though with filtered with cobalt glass some of the flame could still be seen, indicating potassium ions. The crystals also didnt form particularly well, they were sort of pyrimidal or something like that. When I made chlorate from boiling down bleach I assume I got much higher purity crystals because they formed in leaflets, which is characteristic of pure chlorate. On top of this, the pyrotechnic mix I made with it wasn't particularly fast burning. This is a strong contrast with the chlorate from bleach which burns almost as fast as flash.
Hope ya have good luck with your project, and you may need to run your cell again to convert more to chlorate. Remember you can recycle the electrolyte as many times as needed so nothing is wasted!

BTW, for the cell I made was a 200ml glass beaker with a spiral of 14 guage copper wire lining the inside of the beaker.

Just an update

frogfot - 11-4-2003 at 12:46

Well, i got totally 217g of product from above named electrolyte. After second recrystallisation i got 181g snow white product:
http://www.geocities.com/frogfot/stuff/chlorate.jpg

It seems that such big mass loss indicates that lots of KCl was present, about 20g. Btw, I used vacuum filtration, it makes wonders.

I have "modified" the processing procidure a bit, i boil it before filtering, this way it can't destroy cotton filter.
Procidure goes like that:

First, pH in crude electrolyte is adjusted to 6, and then it is boiled for 30 minutes. Then it's chilled and filtered through compressed cotton. After it's ready for extraction of product. I found that there are no difference between using pure KCl and a mineral salt (a mix of NaCl and KCl 40:60), well, it gives same mass after first recrystallization.

It seems that used procidure is sufficient to remove those small annoying coal particles (and strange colors). Btw, now im pretty sure that strange color was from polystyrene cell. Next cell will be made of polypropylene and chromates will be used.

chlorate

tryptamine - 26-5-2003 at 23:24

Tips.

Your yield is going to be mostly salt if you don't use chromates, there is a reason that industrial plants use them, they prevent your product from getting reduced as soon as it is made.

You need to run the reaction at 80-90 degrees celcius, and after you are done running current through it you need to let it sit and react at this temperature for a while.

Adjust the pH to 6 using bleach before your run, bleach is an intermediate.

Pure chlorate crystals are not leaflets, all the chlorate I've ever seen(think thousands of tonnes) is like salt, some of the bigger crystals (I saw some today) are sorta rhomboid.

Carbon is old school, new school is stainless and platinum coated titanium. If you must use carbon you still need chromates, to get it out filter through celite(pool store).

A high currrent and low voltage is used, the cell banks in an industrial plant are at about 3.5 V per cell, and well over 100,000 amperes.

Anything else pertaining to chlorate manufacture? For the non totse members, I work at a chemical plant where one of the main products is chlorate.

Marvin - 29-5-2003 at 12:41

ph6 is acid, bleach is a base. Possibly you were indending to say ph9 ish? I think mild base is best, but if it gets too high the carbon rod erosion rate suffers, as its OH- catalysed. The ph of the solution should be maintained by adding dilute HCl as needed during the electroylsis.

perhaps

tryptamine - 30-5-2003 at 17:59

Perhaps I should have said that in the industrial plant they use hypochlorite to adjust the pH. They also use HCl.

frogfot - 6-6-2003 at 00:23

Soz, this is kinda late reply but i couldn't let it be..

Chromates is used to rise efficiency with 10-20%, this is alot for an industry but not in home cell. However i ubtained some potassium chromate and will use it in next cell.

"Your yield is going to be mostly salt if..."
Do you mean i will yield NaCl after a recrystallisation? Impossible, then it wouldn't precipitate with temperature fall.

80-90*C is inaproppriate for carbon electrodes, they will corrode at great rate. Recommended temperature is 40*C.

Why addjust pH with bleach? This simply wouldn't be economical. HCl do the trick for me, and most factories

2kg sodium hypochlorite cost 40$ here.

Carbon is used in many chlorate plants even today, since it's cheap. Titanium would cost here same as gold....

I got one question. I heard that graphite plants have a continious flow of electrolyte through a hypochlorite converting chamber, where it's heated to 90*C. How big current per volume requiers this? I mean, if my new cell will operate at 8A/liter, do i need to stop electrolysis and heat electrolyte from time to time?

Chlorate!

Theoretic - 30-6-2003 at 03:58

I suggest putting in some AN or any ammonium salt in the electrolysis vessel - equimolar quantity of NH4+ to ClO3-. These form a nicely insoluble precipitate.
It's OK as long as you're happy with the explosive NH4ClO3, which doesn't give the flame a colour. :cool:

Theoretic - 30-6-2003 at 07:03

Can anyone tell me, what apart from pH, influences whether Cl2 or ClO3- is produced (I know that chlorine disproportionates faster in alkaline solutions)?

rikkitikkitavi - 30-6-2003 at 09:03

actually, chlorine doesnt disproportionates in alkaline solution, it reacts with hydroxide ions forming hypochlorite ions, this reaction goes faster the higher conc of hydroxide it is.

Hypochlorite disproportionates to chlorine and chlorate, maximum speed is at pH 6-7 because it is a reaction between hypochlorous acid (HClO) and hypochlorite ion. Higher pH , there are virtually none acid, lower the acid decomposes into Cl2 and Cl-. higher temperature also affect the reaction positively with a maximum around 70-80 C.

/rickard

Theoretic - 1-7-2003 at 06:46

"actually, chlorine doesnt disproportionates in alkaline solution, it reacts with hydroxide ions forming hypochlorite ions, this reaction goes faster the higher conc of hydroxide it is."

Well, the reaction between chlorine and OH- to form ClO- and Cl- is called disproportionation :mad:

.
Also chlorine could directly disproportionate to ClO3- and Cl-, as it can from ClO-.

vulture - 1-7-2003 at 08:52

Well, the reaction between chlorine and OH- to form ClO- and Cl- is called disproportionation

Not true. A disproportionation reaction is a reaction where ONE compound reacts with itself to form 2 compounds with different oxidation states of a central atom. In hypochlorite to chlorate, some of the chlorite gets oxidized to chlorate, will some chlorite will be reduced to chloride.

Disproportionation

Theoretic - 3-7-2003 at 04:33

Not, true!
A reaction where a compond or an element react with another compound or itself, where a certain element gets both oxidized and reduced is called disproportionation.
In CHEMISTRY BOOKS the reaction of Cl2 with alkalis is called disproportionation, so is that of sulfur.

markx - 7-8-2003 at 02:23

All the fine tuning (pH control by acid or base, adding dichromates or other compounds to inhibit cathodic reduction of ClO3- etc.) is not a must when it comes to chlorate production on a domestic scale!
Believe me I have produced several kg of KClO3 by electrochemical means and after trying all those methods to improve productivity I can say with absolute confidence that it's not worth the trouble! The effect
of those additives on the final amount produced can only yield some results on an industrial scale at home the difference is so small that there is no need to add anything into the solution besides the chloride.
I used a concentrated solution (at room temperature that is) of fertilizer grade KCl in regular tap water. Tap water contains calcium compounds which inhibit the ClO3- reduction at the cathode just like dichromates do.
There is no need to use NaCl solution and then convert it to KClO3. Smart books say the production rate is higher in NaCl but in the end it's unnoticeable and there's a ton of trouble and more things can go wrong. So just keep it simple and use KCl to begin with!
I took 'bout 5L of straight KCl solution a 5*5*10cm graphite anode and a 5*10cm stainless cathode (works great, doesn't corrode!!) and ran 20A of current through this setup for at least 12 hours.
The reaction temperature is around 40-50C and is kept by the current and the cooling system. Just tipped the bottom of the 5L cell into running water (a creek that is!!), works great he-he

The temp is really not critical as long as it's over 40C.
As exp shows this system is rather foolproof as long as one keeps the temp over 40C the current over 3A/per liter of electrolyte and the additives out of the cell!
There really isn't much sense in prolonging the
reaction time over 24 hours. Taken the same amount of KCl and current that is! After the 12 hours I get about 100-150g of KClO3 and it's enough for my needs. Don't take these numbers as an exact reference but about 8 cycles yield a kg of KClO3 so the single cycle must produce about that amount.

When it comes to filtering there's only one solution to get all the fine crap out of the solution - glass wool! I use standard insulation wool (the yellow stuff in the walls to keep warmth in house) it costs nothing and filters like a mother!! I usually collect the dirty results of 4 to 5 cycles and
then hot filter them once and the final product is white as snow and a real eye-candy!

Damn I've been blabbering on....
Well so long!

Theoretic - 8-8-2003 at 05:58

Well... if anyone wants to try NaClO3 for a change? More oxygen per gram and a "golden flame".

markx - 13-8-2003 at 00:28

Yup! And it's hygroscopic and readily soluble in water, meaning that you will have alotta trouble with separation from chlorides and drying.
But it's soubility makes it a perfect starting material for producing other chlorates that have a low solubility (metathesis reactions) BaClO3 or KClO3...
When producing KClO3 it's better to start with KCl and not mess around with the metathesis! It will save time and trouble and no residual Na ions will be introduced into the final product.

CommonScientist - 4-3-2004 at 11:21

Could ytou hand some KCLO3 by a string or wire into your electrolyte to form one big crystal? The electrolyte would be the mother liquor and the KCLO3 would be the seed crystal. It should work, but I cant back it up cbecause I havnt tried it before.

I belive there is a section on wouter's page that explains why you shoudnt start out with a KCL brine solution as your electrolyte.

axehandle - 4-3-2004 at 11:51

Sorry for the late reply, countrymate, but is a Pt anode out of the question? You could use normal 18-10 stainless as the cathode (use a baking bowl and a hacksaw...) and a platinum wire as the anode. I ordered one 3 weeks ago, it should arrive soon. A 300mm long wire with a diametre of 0.8mm. Cost: about 1500SEK. But it will outlive me and you, and you won't have problems with carbon particles.

Saerynide - 22-3-2004 at 06:26

Im curious. If Cl2 and H2 explode when exposed to UV light, then would a chlorate cell explode if it was made of glass or anyother UV penetratable material?

t_Pyro - 25-3-2004 at 10:17

Unless the cell is absolutely air-tight, there shouldn't be any problem. Hydrogen, being the lightest gas, diffuses through even the smallest of openings pretty fast. Whatever small quantity of hydrogen gas that does combine with chlorine will immediately get dissolved in the water again. Also, glass is not transparent to UV light.

Saerynide - 25-3-2004 at 11:08

Glass is not transparent to UV light??

Then why do glasses need anti-UVray coatings and how did I get sun-burned through the window on a 6 hour bus ride?

t_Pyro - 26-3-2004 at 02:50

No, glass is not enitrely transparent to UV light.
Regarding sunburns, you might be wondering <a href="http://physics.about.com/cs/light/a/010803.htm"> "Can you get Sunburned Thorugh Glass?"</a> That should answer all your questions.

[Edited on 26-3-2004 by t_Pyro]

needle - 3-6-2004 at 08:18

Many writes of the dichromate-addition in the cell to raise efficiancy. But what does it do? I've both read that it prevents the reduction of ClO3 to ClO, and that it itself it reduced to Cr(III) as a protecting layer around the cathode. Is this true?

Could one add another, less dangerous, chemical with the same effect? Perhaps KMnO4?

Thanks.

evilgecko - 6-1-2005 at 19:32

I know this is an old topic but I have got one hell of an annoying chlorate cell in my shed

. It firstly consisted on two 100mm long (immersed) carbon rods of 8mm diametre about 15mm apart in a concentrated salt solution. The current was only 0.8A so I added 7 more electrodes, forming a circle, each one 15mm away from the two nearest. Each one has two of the opposing charge beside it eg cathode-anode-cathode all the way around the circle. The current with this setup was 3A but after 5mins dropped back down to around 0.8A. I could have 7A going through this withough succseding the current density of 43mA/cm2. Any suggestions to why the current is so low and how to increase the current?

hodges - 7-1-2005 at 17:42

I'm not sure why this would reduce the current but the approach you are using for series cells is not likely to work very well. There should only be two electrodes in each cell; connect cells in separate containers in series to run multiple cells at once. Otherwise the electrodes being in the same cell are going to short each other out.

Also, two carbon rods is not ideal for a cell. You want a carbon rod for the anode, but for the cathode you want a large piece of metal arranged so that the chlorine bubbles from the anode rise and mix with the solution next to the cathode. Otherwise the chlorine will escape and you will end up making sodium hydroxide.

evilgecko - 7-1-2005 at 19:24

Well I found the problem with the current, it was to do with the connections between the power supply and the electrodes. I had used clothes pegs with aluminium foil wrapped around them and the wire inside. They just gripped onto the electrode. I discovered this when I though that the problem was the voltage, and increased it to 12V. As you would expect these got very hot and they started melting. So I made some connections similar to the ones for 12V car batteries out of a cat food tin. With these suddenly my current jumped to 7A. I lowered the voltage back down to 5V and now it is working at 4A with 700mL of electrolyite. I heard somewhere that optimum is 2A per 100mL. Anway I'm designing my 5th cell at the moment, using a carbon rod and a stainless steel spoon. If I drill a 12mm hole on the spoon, then I can bend it so that the carbon rod goes through the middle of it, so rising chlorine will hit the spoon and get trapped by its conical shape for a bit. Is this a good idea? Also the cell will be taller and thinner, so that there is only 500mL electrolyte.

evilgecko - 9-1-2005 at 01:01

One question, can you use 10A 240V fuses on 5V and still have it blow at 10A? Or is it on power, so that i would need 0.2A 240V fuse to blow when 5V reaches 10A. Or is it something compeltely different which means I'll have to find a 5V fuse. I would of done a practicle to find this out but then I realised all my fuse wire had gone into making electrical ignitors.

hodges - 9-1-2005 at 14:13

10A 240V fuse will blow just as well at 5V as at 240V. It will blow at its rated amperage. You can use a fuse at a lower voltage than its rating, but not at a higher voltage (since a higher voltage might arc across the gap left when the fuse blows).

evilgecko - 9-1-2005 at 15:26

Thanks, hodges. Now I can leave my chlorate cell on all night.