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Author: Subject: KCLO3 by way of H2O + KCL
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shocked.gif posted on 1-7-2003 at 06:55


How, apart from making the solution alkaline at the start or waiting for it to turn alkaline, can one reduce the amount of Cl2 escaping?:o
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[*] posted on 13-7-2003 at 20:15


Increasing anode surface.
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[*] posted on 26-7-2003 at 21:09
another thread on chlorate


i thinks its called
sodium and potassium chlorate. was recent
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[*] posted on 7-8-2003 at 04:04


Put the anode at the bottom of the cell and the cathode just beneath the surface of the electrolyte!
That creates a reaction zone in the middle of the cell as the alkali floats down and chlorine rises up.
That works great for eliminating the escape of Cl2.
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[*] posted on 8-8-2003 at 03:05


Duh! Good idea. Amazing how our minds can be restricted by always seeing those chemistry textbook pictures where the electrodes are positioned upright.:o

CCl4 does form with electrolysis of molten chlorides because of the high temperatures and inert atmosphere used.

CF4 also forms to some extent when making Al from molten kryolithe.

But I doubt it will form in aqeous solution.

BTW, CCl4 isn't that bad. It's a victim of the well known chemofear phenomenon. It has been used in lab fire extinguishers in the past and those lab assistants didn't have a particulary high cancer toll.

[Edited on 8-8-2003 by vulture]




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[*] posted on 24-10-2004 at 06:22
KClO3 by Pt, Ti, KCl, H2O and electricity


http://species8472.dyndns.org/chem/KClO3/index.html

Thought I'd share some porn pictures of my <b>successful</b> KClO<SUB>3</SUB> cell using Pt and Ti with a saturated KCl solution...




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[*] posted on 24-10-2004 at 07:17


Those look like good yields. I take it that you've had a lack of success with the PbO<sub>2</sub> electrodes?
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[*] posted on 24-10-2004 at 07:46


Quote:

Those look like good yields. I take it that you've had a lack of success with the PbO2 electrodes?

Not at all, not at all, I'm still trying. I just wanted some ultra-pure (not contaminated by PbO<SUB>2</SUB>;) KClO<SUB>3</SUB> for use in a couple of percussion sensitive mixtures I'm very keen to try...




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[*] posted on 24-10-2004 at 15:39
KClO3


I couldn't resist responding to this thread given past experience in making chlorates and
perchlorates.

Vulture, mon ami, I'm glad you're responding to this thread. Your knowledge and
experiences are much needed here. Chlorates are indeed very powerful oxidizers.
Bad oxidizers ? No way ! Along the lines of your response, chlorates mixed with P or S
are unpredictable and dangerous. 30 years ago I blew off my eyebrows with KClO3 and
P - known as Armstrong's mixture. Most experienced pyros are familiar with that mix.

Frogfot, I've never detected the formation of CCl4 when making chlorates/perchlorates -
just that irritating chlorine gas. This is with gouging rods from a welding supply shop.

Rikkitickitavi, the salts you mentioned provide some cathodic protection but also
increase the efficiency of the cell. In the past, I've used potassium dicromate for this
purpose, but the fluorides and persulphates work much better for perchlorate production.

Markx, what you posted is obvious. The problem is finding an anode that won't be
corroded by the brine when completely submerged. I would LOVE to have an anode
with that capability !

Aaron-V2.0, manganese dioxide will work. but how will you get it into a usable shape ?

MadScientist, ClO2 was used a few years ago to kill the anthrax spores in the offices
of a few U.S. Senators. I don't know how the exterminators produced it but I'm sure
it would be interesting to the rest of us to produce this gas without the risk of explosion.

Last but not least, Axehandle, are you sure those are crystals of KClO3 ? In my own,
experience, KClO3 crystals look like cactus needles. From your picture, it looks more
like the rhombic KClO4 crystals. Just an observation !

Anyway, just a few observations and comments. I wish everybody the luck and
success in their scientific endeavours ! If it appears that I'm talking out of my ass,
PLEASE, don't hesitate to flame me ! I'd rather be be flamed than have something
send me to an early grave ! Thanks Again !




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[*] posted on 24-10-2004 at 19:20


I'm kind of new to electrolysis, so I'm going to ask a few basic questions on this thread. Basically I can take a 9v battery hook up the negative wire to a PbO2 anode, and a positive wire up to a graphite cathode, stick both of them in a solution of K and Cl ions and presto I have Hydrogen and KCLO3?



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[*] posted on 24-10-2004 at 19:36


That would work, but it would take a very long amount of time. The amperage emitted by a 9v battery(I am assuming you mean one of those smal .5"x1"x2" sqare ones) is not enough to produce chlorate at a reasonable rate.
ne=IT/F, where ne is moles of electrons, I is amps, T is time, and F is faraday's constant, 96500 C/mol
So the higher the amperage, the more moles of electrons produced and thus the more chlorate/perchlorate is produced.




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[*] posted on 24-10-2004 at 19:47


Don't forget, most importantly, the reaction vessel has to PREVENT THE NASCENT CHLORINE FROM ESCAPING - because of the DESIRED reaction of the freshly produced chlorine with the freshly produced KOH!, forming KOCl, which disproportionates eventually to KClO3 and KCl (where KCl is electrolysed to KOH and Cl2 once again).
That is the most important rule. If you let the chlorine gas escape, then you will PREVENT the electrolysis from performing in the desired manner!

[Edited on 25-10-2004 by chemoleo]




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[*] posted on 24-10-2004 at 22:50
Escaping chlorine


Chemoleo, I couldn't have said better myself ! The more the chlorine escapes, the less the
chlorate/perchlorate produced ! Also the brine becomes more alkaline(rising PH) as the Cl2
escapes.

In a previous post, someone talked about keeping the anode completely submerged.
DUH ! We would have done this already if it was possible ! So mark x, a little more practical
application and less theory would go a long way here !




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[*] posted on 25-10-2004 at 07:25


Quote:

Last but not least, Axehandle, are you sure those are crystals of KClO3 ? In my own,
experience, KClO3 crystals look like cactus needles. From your picture, it looks more
like the rhombic KClO4 crystals. Just an observation !

No, I'm not sure at all. Should be impossible though since I made it straight from KCl and the crystals precititated/crystallized out at the bottom of the vessel. Btw, the crystals in the pictures are really lumps of smaller crystals. Some of them are rod-shaped, some of indeterminable shape. I'll have to look closer, or perhaps try to decompose any KClO<SUB>3</SUB> with HCl, leaving the KClO<SUB>4</SUB>, if any... hmm.

Quote:

I'm kind of new to electrolysis, so I'm going to ask a few basic questions on this thread. Basically I can take a 9v battery hook up the negative wire to a PbO2 anode, and a positive wire up to a graphite cathode, stick both of them in a solution of K and Cl ions and presto I have Hydrogen and KCLO3?

No, if you hook up the negative wire to the PbO<SUB>2</SUB> anode, you'll end upp dissolving the anode. Try the positive wire :).

Leaving your probable typo aside, a battery wouldn't last 10 minutes. You need amps upon amps for days even with a small 1 liter cell such as mine.

Quote:

Don't forget, most importantly, the reaction vessel has to PREVENT THE NASCENT CHLORINE FROM ESCAPING - because of the DESIRED reaction of the freshly produced chlorine with the freshly produced KOH!, forming KOCl, which disproportionates eventually to KClO3 and KCl (where KCl is electrolysed to KOH and Cl2 once again).
That is the most important rule. If you let the chlorine gas escape, then you will PREVENT the electrolysis from performing in the desired manner!

I've tried to make the escape vents (actually, electrode holes in the lid) as small as possible. Still, I must have at least a small opening, otherwise the cell would go boooom...




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[*] posted on 25-10-2004 at 11:18


What about that ignition coil driver lab on power labs? Would that be a good source for a high amp, direct current flow?



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[*] posted on 25-10-2004 at 11:37


Quote:

What about that ignition coil driver lab on power labs? Would that be a good source for a high amp, direct current flow?

Certainly not! An ignition coil delivers a an extremely high voltage at an extremely low amperage. Ideal power sources for a (per)chlorate cell are an ATX PSU, a car battery charger or for a huge cell, a DC welding transformer.




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[*] posted on 25-10-2004 at 14:30


So for making my PbO2 anode I just connect a piece of lead to a positive chare and put it in a bath of strong sulfuric acid. Maybe use copper as my cathode? All so I saw someone bitching about thier KClO3 being contaiminated with PbO2. couldn't you simply filter the PbO2 out of the KClO3 solution? Oh ya so could I use copper and PbO2 instead of graphite and PbO2 in my chlorate cell?

[Edited on 25-10-2004 by tom haggen]




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[*] posted on 30-10-2004 at 17:57
Chlorates/Perchlorates


tom haggen, definitely lower voltage. Axehandle, you said the some of the crystals
were "rod" shaped. Probably the chlorate crystals. The rhombic crystals are ususally
perchlorate. Easy to separate due to the solubility differences.

Axehandle, do you have access to HCl and, more importantly, indigo carmine ?

Good job on your response to tom haggen's post ! I remember seeing a website on a
German company who produced sodium chlorate directly from a salt mine. It appears
they sent the brine up from the mine directly to chlorate cells constructed of concrete and
using graphite. The "push" was less than 5 volts although the amperage was higher
than hell. Talk about efficient !


[Edited on 31-10-2004 by MadHatter]




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[*] posted on 1-11-2004 at 20:05


Mad Hatter I fail to see why you cannot completly submerge a PbO2 anode in the solution. Perhaps it is my own limited knowledge, but it seems to me that if you set up a cell where the anode and cathode are protruding into the cell from the sides, anode below and cathode above, as long as those whole are sealed with some no-conductive sealent, the cell should be fine. The only other problem with that set up is the excess of corrosion at the point where the electrodes enter the cell, and in my experiance they break off. I'm thinking this could be fixed by slanting the ends of them so that the ends are closer together than from the begining of the cell. That would make it "easier" for the electrolysis to occur at the ends and would corrode it there, rather than at the junction. I don't know, maybe that thinking doesn't hold.


On PbO2 electrodes, I just recently cracked open an old car battery(used). Fron what I've read about them the fail becuase the anode and cathode are coated by PbSO4, stopping the elctrolysis. Yet when I ran current through them in an NaCl solution, the pleasent smell of chlorine filled the room.

I am however, having problems telling the Pb from the PbO2. I've fully removed one of the 6 cells. In it i found the electrode side by side, but one of them was covered in a removable plastic sheath. One of them has a thin white layer on it, and I'm thinking that this would be the PbO2 as the SO4 ion would be attracted to the positive lead ions coming off of it. Both are on fairly maleable, are

I don't know, I feel like I'm missing something in that logic, but my tired mind cannot grasp it at the moment.

Any idea's are welcome, its interesting that something so cool can be made out of such easily available and not so expensive materiel.




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[*] posted on 3-11-2004 at 17:22


Quote:

Mad Hatter I fail to see why you cannot completly submerge a PbO2 anode in the solution. Perhaps it is my own limited knowledge, but it seems to me that if you set up a cell where the anode and cathode are protruding into the cell from the sides, anode below and cathode above, as long as those whole are sealed with some no-conductive sealent, the cell should be fine. The only other problem with that set up is the excess of corrosion at the point where the electrodes enter the cell, and in my experiance they break off. I'm thinking this could be fixed by slanting the ends of them so that the ends are closer together than from the begining of the cell. That would make it "easier" for the electrolysis to occur at the ends and would corrode it there, rather than at the junction. I don't know, maybe that thinking doesn't hold.
...

Intractable, IMHO --- it's nigh impossible to find a sealant that would 1) be resistant to the Cl<SUB>2</SUB> and anodic oxidation as well as 2) would provide tight enough fit to stop the brine crawling out along the electrode by capillary action and 3) would adhere to both the PbO<SUB>2</SUB> layer on the electrode as well as the hole in the cell wall.

I could be wrong though.




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[*] posted on 3-11-2004 at 21:18


glass to metal seal



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[*] posted on 6-11-2004 at 13:33
Sealants and Slant Angles


Axehandle, excellent reply to Magius ! Every sealant I've checked the MSDS on so far
shows reactivity hazards with OXIDIZERS. Even the water-resistant silicone sealant has
this problem.

Magius, the slant angle does work. I place the 2 rods in parallel separated by 4 inches. The
cathode is on the top. This way the Cl2 rises to meet the hydroxide forming at the top.
Still, some of the Cl2 escapes.

The other idea I've considered is to cover the top end of the anode with a non-conductive
material so that the exposed end would remain completely submerged. I wouldn't
get the full surface area of the anode as a result but if I could retain most of the Cl2
in the brine I should get more efficient cell even if it means using 2 anode rods.

I saw a seller on eBay selling sheets of graphite. Might make an interesting cathode.




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[*] posted on 8-11-2004 at 21:25


In my recent experiments with the preparation of chlorates I turned to electrolysis. I simply used a beaker with two graphite electrodes and a saturated NaCl solution. So after around a day of fizzing it finally stopped. So I took the solution out of the beaker and into a filter. This was than placed in a stainless steel bowl and boiled down on an open flame. My girlfriend came around a little unexpected and I got distracted. When I finally came back I noticed that the White powder was turning black and reddish brown. I let it cool and scooped some of this powder out and mixed it roughly with sugar and one with charcoal neither burnt. So I have come to ask is there away to see if I have made chlorate?

I also noticed something weird. There is a thread about making your HCl by electrolysis a solution of NaCl to get H2 at one electrode and Cl2 at the other. Then burning them to get gas and bubbling it through water.

Electrolysing to get chlorate should be as follows

NaCl + 3H2O -> NaClO3 + 3H2

The only gas that should be involved is H2 but I noticed one electrode bubbled greatly and the other just bubbled a tiny bit. Why?

So under what conditions do each reaction occur.
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[*] posted on 9-11-2004 at 15:29


I see several problems here, Mark.

1. The reaction should not "stop fizzing". Were you using a battery that may have went dead? It takes a lot of current to produce just a little bit of chlorate from chloride, so if you were using a battery anything smaller than a car battery that is one problem.

2. Chlorine and sodium hydroxide are produced as intermediate products at the anode and cathode respectively. You must provide a way to continuously mix the two or else the chlorine will escape and you will not get the reaction you want.

3. Generally the temperature must be somewhat elevated above room temperature for the desired reaction to occur. If you have enough current flowing, usually this will take care of itself. But if you are using a small battery, it will not.

4. In order to separate the chlorate from the remaining chemicals (chloride and hypochlorite), you must add a potatium salt and chill the resulting solution. Potassium chlorate has a very low solubility at low temperatures so it will precipitate out. If you just evaporate everything you are going to get a mixture of products and the main ones are not likely to be your chlorates.

Hope that helps.
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[*] posted on 9-11-2004 at 22:40


I am using a 10 v 1 Amp dc current. It’s an adopter from something not sure what though.

Well this project started as way to get chlorine for organic hypochlorites. But not much happened so I decided might as well get some chlorate from it.

The temperature was at around 40 degrees most time.

Could I use KNO3 in the separation process?


Also thanks for the help!
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