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manvstaco
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Quick question on a potassium chlorate cell
Ok so I will be making a decent power supply out of an old transformer like this video here -- > http://www.youtube.com/watch?v=Yr8HKlX4VnU , getting tied of the pc power supply and how slow things move......... I am wondering....
I have heard of people using carbon rods for both the anode and a cathode if this is the case can I use a low voltage high amp A.C. to power my
cell since my electrodes are the same material? is this possible or am i dreaming? I dont want to go buy diodes to convert it to D.C.  I.e. 5 vlts @ 200 amps A.C.
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hissingnoise
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Graphite cathodes and anodes are fine for electrolysis but you need direct current for it to work. . .
I suggest you read up on the subject before going ahead!
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manvstaco
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Yeah I generally use 6 volts dc with 15 amps of juice, Carbon for the positive and Stainless steel for the negative with great results just throwing
ideas out..
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hissingnoise
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This board is awash with info on chlorate/perchlorate cells. . .
You just need to UTFSE!
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manvstaco
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yeah most cells dont use the same materials for both electrodes requiring a D.C. , but if they are both the same would A.C. work is all I am
wondering, I have been tweaker style on the search engines for weeks about this and have been reading this website since 03 but didnt feel I had a
legitimate question till now..
I have a good understanding of chemistry and was in demolitions with the military for a 4 year term, please dont curse at me it hurt my feelings ahaha
just joking
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gregxy
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Its an interesting question. Using AC would save the cost and losses
in the diode. But I suspect many of the things that are oxidized at the anode
would be reduced before they get a chance to react when the anode becomes
the cathode 60 times per second resulting in low efficiency and excessive heating.
Note that chromates are added to the solution specifically to prevent losses due to
reduction of chlorate at metal cathodes.
Diodes are cheap and don't wear out, and metal cathodes are easier to come by
than graphite so I doubt that you would gain much. Industrial processes use DC
with MMO anodes and stainless steel cathodes.
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hissingnoise
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| Quote: | | But I suspect many of the things that are oxidized at the anode would be reduced before they get a chance to react when the anode becomes the
cathode 60 times per second resulting in low efficiency and excessive heating. |
There is no overall reaction using AC as each cycle cancels the preceding one.
My cathode and anode were graphite because I didn't want metal ions in the electrolyte.
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manvstaco
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Thank you for all your answers guys you have been helpful! I just rigged up 3 300 watt computer supplies to my cell 5 volts each rated at 20 amps, on
a parallel circuit, 14 electrodes total. Its got a 5 gallon bucket hot to the touch lol, i have to change out 7 3/8ths carbon rods a foot long every
other day, the cathodes are ss butter knives.. I have maybe 32thousanths resistance through my cell how can I tell how many amps its pulling through
it? they could do 60 amps but im thinking maybe im only hitting like 40? any idea?
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densest
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There are a couple of ways to check the current... the simplest, but most expensive, is to buy a snap-on current meter. They're down to $30 or less.
Be sure to get one that reads DC amps, because some only read AC. I assume you have a voltmeter.
Another way is to measure the voltage drop in a measured distance along one of the power wires. The resistance is known if you know the wire gauge.
Assuming it's #18, that's 0.0210 ohms/meter, so each millivolt per meter = 48mA. For 10cm, each mV is 4.8A,
multiplied by the number of wires in parallel.
How did you measure the cell resistance?
Is the voltage between electrodes 5V, 12V, ???
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quicksilver
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You really have to get a clamp meter to be sure. Especially with a computer PSU as it will switch on you as the solution changes via natural
alteration in solution composition & evaporation.Computer CPU's are tricky. They can pump out good heavy amperage for a few hours then drop for
reasons only they know....
I have used graphite rods quite a bit. And they are fine but the way to do it is to use a 5gallon bucket and a LARGE level of solution! You are going
to have to clean up the final product: that's simply a fact. SO why clean a little amount several times when you can clean ONE 4 pound yield????
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pjig
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Computer CPU's..... I have one from the late 90's will it be enough to run a cell of any size( like a 5gal) or do I need to go to a bigger power unit?
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gregxy
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If you do use multiple supplies keep the outputs separate, for example
if you use 12 rods tie 4 rods to each of the 3 supplies. Otherwise they may
not evenly share the current.
It may be less hassel to run at a lower current and just wait longer rather than
have to keep changing the rods. I seem to remember 100ma/cm^2 as a good
current density for graphite (Its been a long time so I'm not sure).
There used to be a guy on Ebay selling graphite plates that were good for this
kind of thing.
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pjig
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Any way of getting a good source of graphite plates or rods? In the industrial world.? I thought I saw a post a while back showing a cheap source of
graphite.... not sure if it was on this forum though..
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watson.fawkes
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Quote: Originally posted by pjig  | | Any way of getting a good source of graphite plates or rods? In the industrial world.? I thought I saw a post a while back showing a cheap source of
graphite.... not sure if it was on this forum though.. |
Always check McMaster-Carr: http://www.mcmaster.com/#graphite/.
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quicksilver
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Naturally I used scarfing rods from an industrial welder. I got them for next to nothing and they are quite thick (about 7/10 of an inch). I treated
them so they wouldn't break down so easily (they don't break down for a LONG time anyway).
I agree on using more than one PSU for several electrodes: it's a sensible thing to do. You CAN get several anyway if you start hunting; they are
mostly free. EXCEPT if you want to find a 1000 watt unit.....that may be a bit harder to find.
I was VERY lucky in getting some really professional power supplies but I HAVE used a computer PSU and it was fine and still works well after several
huge setups. Most of the time I end up with a beige KCLO3 from my starting material: KCLO water softener. I would rather spend the money and get KCLO3
than go through the additional steps to go from NaCLO3, etc. With the beige KCLO3, I simply re-crystallize and my final product is reagent pure.
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pjig
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Awesome! how did you treat your graphite, if you don't mind sharing...
I would like to scale up to a 5gal sys too. But am unsure of the needed power and run time with that source, to obtain a completed reaction. Any ideas
using the computer source of power, or maybe microwave source( might take a bit of work to convert)... I am aiming at simple, easy setup at a low
cost. ( I like the welding rod idea, and stainless knifes as stated above).
Now off to find 2 more computer power supplies. ( if that is the cheapest route).
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manvstaco
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pjig watch the video I originally posted it will teach you a lot.. Also an update.. I treated my carbon cutting rods with veggie oil and havent had
to replace them going on 3 days now, with little sign of wear, give it a shot and tell me how that works. I didnt want to spend 25 bucks for liinseed
oil
I have a new setup also, 3 liter glass cookie jar with the 3 computer power supplies running together, giant bbq spatula with 2 12" treated carbon
cutting rods... combinded 64 amps total output, I was expecting it to boil over but soo far so good. keeping the temp right around 140 F
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pjig
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Do you have a vid. of the running cell? I liked the rewinding idea, it looks much easier than I thought it would be...
Thanks for the tip on rod treatment too....
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The WiZard is In
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Chlorates And Perchlorates: Their Production
I didn't find this on la form so....
Chlorates And Perchlorates: Their Production
BY JAMES FINCKBONE
American Pyrotechnist Fireworks News
Volume 7, Number 6 June, 1974
The current drives by pressured bureaurocrats [sic] and safety authorities to
restrict the flow of chemicals with a potential for abuse to the public,
especially to hobbyists (e.g., the Child Protection and Hazardous Substances
Acts of 1966 and 1967), have substantially achieved their goal, with most
bulk-breakers losing too much money to continue selling chemicals to such
customers. Because of this dwindling supply of certain chemicals, particularly
oxidizers, a few chemists among the deprived hobbyists are considering a
do-it-yourself approach, compounding their own oxidizers from naturally
occurring and relatively abundant elements, not unlike the efforts during our
Revolutionary War to collect nitrates from every household's privy. But the
most important high-energy oxidizers for the pyrotechnist, the chlorates and
perchlorates, are not found in nature. Here we consider some of the
problems encountered by the hobbyist attempting to produce these chemicals
on a limited scale and with limited equipment. The writer does not, however,
claim the present article to be the final word on the subject by a long shot,
and readers are strongly urged to forward their own observations and
suggestions for printing in future issues.
The two classes of agents in question did not become widely available until
the advent of the Industrial Revolution, although they had previously been
made by several European chemists by heating red lead (Pb304) with mixes
of prepared chlorine and a base. Today the chloralkali industry produces
chlorine and sodium hydroxide in abundance, using an electrolytic process
and ocean brine. The chloride ion in the brine is converted to the chlorate ion
by the application of electrical energy (noted in the following formulas as ''f",
designating faradays, unit quantities of electricity), and finally to the
perchlorate ion. A metal donor added to the solution causes a precipitate to
form; e.g., barium carbonate will precipitate out barium chlorate. The overall
reaction is:
Cl - + 3(H20) + 6f -> Cl03-+ 3H2(gas)
Stepwise, the conversion is:
1 NaCl + H20 + 1f -> NaOH + 1/2(H2) + 1/2(CI2) -1 -0*
2. 1/2(CI2) + H20 + 1f -> NaClO + 1/2(H2) 0-+1
3. NaClO + 2(H20) + 2f -> NaCl03 + 2(H2) +1 -+5
4. NaCl03 + H20 + 1f NaCl04 + H2 +5 -+7
(*oxidation number of the chlorine)
There is an unstable intermediate between steps 2 and 3: the chlorite ion,
C102-.
Considering for the moment only the chemistry of the electrolytic chlorate
cell, free chlorine gas is initially formed at the anode (the positively-charged
electrode) as soon as current flows. Simultaneously, at the
negatively-charged cathode, hydrogen gas and hydroxyl ions are forming.
(Metallic sodium is also formed but is unstable in water and reduces its
hydrogen.) If the electrodes are not separated by a diaphragm, the dissolved
chlorine gas in the electrolyte diffuses toward the hydroxyl ions and
vice-versa, forming the hypochlorite, hypochlorous acid (HCIO) and the
chloride ion (which is later oxidized). Since the oxidation voltage for con-
verting hypochlorite to chlorate is the same as for conversion from chloride to
hypochlorite, chlorate soon appears (the oxygen being derived from the
water, as confirmed by tagged-isotope study), and more hydrogen gas is
evolved from the electrolyte. At the same time, the competing reaction of
water electrolysis is releasing oxygen gas at the anode and hydrogen at the
cathode, so the reaction efficiency is reduced to about 60 to 80 percent of
theoretical. The final step, formation of the perchlorate ion, occurs only when
all chloride has been removed or reacted, as its potential is of a higher value.
The acidity (pH) of the solution will influence other side-reactions, such as the
evolution of chlorine dioxide.
Industrially, chlorates and perchlorates are made in multi-cell tanks of a
corrosion-resistant steel alloy which serve as the cathode, with dozens to
hundreds of rod- or bar-shaped anodes connected to a bus-bar and
projecting vertically from the tank cover into the electrolyte. Vents are
provided to remove hydrogen and other evolved gases and, along with the
electrode connections, deck the tank cover. Solutions are kept cool by means
of a series of steel pipes running horizontally through the tank, as electrical
resistance can easily send temperatures up to above 100o C if not checked.
Vents on the bottom are used to empty and fill the tank.
Concentrated brine (200-300 grams of sodium chloride to a liter of water)
is introduced into the cell, along with a small amount of sodium or potassium
dichromate (1-7 gms./liter) to hinder reduction of the chlorate ion and protect
the cathode against corrosion. Dilute hydrochloric acid may also be added to
set the pH of the solution at 6 to 7. There is a potential drop of 3 to 4.5 volts
across each anode and a current often as high as 10,000 to 30,000 amperes
per cell. [Assuming a 100 cell tank with 3.75 volts across the anodes and
cathode and 20,000 amperes flowing in each leg as average values, Ohm's
law would show a power consumption of 75 kilowatts, which would rather
tend to discourage the amateur experimenter! --Ed. (Van)] Temperature is
held at or below 35oC, as graphite anodes are most commonly used. If
solution temperature gets above 45oC, such anodes are seriously corroded
and oxidized away to carbon dioxide (catalyzed by the hydroxyl ion). Graphite
anodes are widely used because of low cost and high efficiency-; but
sometimes, more frequently in the past few years, anodes of steel coated
with lead dioxide (Pb02) or magnetite (Fe304) have been employed, since they
can be used at higher temperatures.
As solution level drops, more brine is added, until about 500 to 800
gms./liter have been added and 500 to 600 gms. of sodium chlorate have
been formed. Then the electrolyte is pumped out, boiled, and treated with re-
ducing agents like thiosulfate ion (the "hypo" used in photographic
darkrooms) to destroy remaining hypochlorite ions, after which the solution is
filtered to remove the chloride ion. Next, evaporation to a smaller volume and
chilling to about 0oC separates the chlorate crystals. These may be bagged
for use elsewhere, converted by precipitation to another chlorate salt, or put
into a cell not unlike the chlorate cell and further oxidized to form the
perchlorate ion, which represents the highest positive state into which
chlorine can enter. Perhaps half of the bagged chlorate may ultimately be
converted back to chlorine gas at its usage point, as that gas is used
extensively to bleach textiles and wheat flour.
In the production of perchlorates, the sodium salt is again preferred over its
potassium counterpart, as it is much more water-soluble, allowing it to be
more easily separated from sodium chloride, and permitting the cell to be run
longer between drainings. Again, dichromate is added, and the pH is held at
6-7. Voltages of 5 to 6.5 volts per anode are the most common. Here the
anodes are usually of platinum or of lead dioxide coated steel, not graphite,
for reasons of efficiency. Temperature is usually 40-60oC, and the cell is run
until a concentration of 600-800 gms. of sodium perchlorate per liter is
achieved. The electrolyte is now treated to remove dichromate and chlorate
(barium carbonate is useful), and the perchlorates of ammonium or
potassium are usually formed by double decomposition.
On a small-scale basis, successful chlorate/perchlorate production is a
real challenge, and this chemist was largely a failure at the effort. The
requirements of a high-amperage, low-voltage power source largely rule out
batteries on any continuous basis. The common 6-volt battery charger is the
most convenient power source but suffers from lack of variability. The large
types used by garages can deliver an output of 50-100 amperes. The
problem of varying the voltage might be solved by use of a rheostat hooked
in series with the positive or negative output lead, but amperage would
(suffer.
The container, whether a beaker, battery jar or other vessel, must possess
several practical qualities: it must be capable of being cooled, either with ice
packs or, preferably, a continuous stream; it must be resistant to corrosion (if
not used as the cathode), with glass or steel perhaps best qualifying for
continuing use; the top should be able to support connections to the anode
(and cathode) properly and vent the evolved gases without exposing the
operator to their often poisonous fumes (if used indoors). If a separate
cathode is used, platinum, platinized metal or stainless steel (such as a large
spoon) would be best. The anode cannot be of copper, steel or other base
metal, because of the unbelievably corrosive action of chlorine and oxyacids
attacking it. So it must be of platinum, platinized metal, graphite for chlorate
production only) or lead-dioxide-coated graphite or metal. Graphite anodes
can be made from the central element of large 1.5-volt dry cells, with the
added convenience that they already have a terminal attached. The local (or
mail-order) rockhound dealer can furnish platinized anodes, but platinum
ones will cost from $30 to $50 each. Stainless steel or graphite elements can
be plated with lead dioxide by making them -the anodes in a bath of lead
nitrate. A copper or steel cathode is used, and a touch of copper sulfate or
nitrate is added to the electrolyte (see Ref. 2 below).
The clips to the electrodes must be shielded from the fumes or covered
with Saran-Wrap or other plastic to protect them from corrosion. The
electrolyte used for small-scale production is similar to that described above
for industrial use, although the salt concentration may be less. If lead dioxide
anodes are used, chromate or dichromate should be omitted. If the anodes
are made of graphite and the electrolyte temperature should rise above
30-35oC, the solution will be filled with colloidal graphite. The evolution of
chlorine gas and chlorine dioxide can be minimized by continual stirring and
low temperature, but it may present a health problem indoors and should be
dealt with thoughtfully. Of course, smoking around an electrolysis apparatus
is a no-no!
On the basis of knowing the voltage and amperage between the electrodes,
one can calculate when the conversion to chlorate is reasonably complete,
using the relationship of electrical units to molecular weight. One faraday (1
mole of electrons) = 96,500 coulombs; one coulomb = one ampere flowing for
one second; conversion to chlorate requires 6 faradays (-1 to +5 oxidation
number) per mole (107 grams).
Some "hypo" from the local camera supply shop can serve to destroy
excess hypochlorite after the remaining table salt crystals have been
separated away. The nitrates of barium and potassium could be used to pre-
cipitate out most of the available chlorate from a cold solution if it were not to
be carried into perchlorate production. Morton's Salt Substitute could supply
potassium chloride, but it is more expensive per pound than pyro-grade
nitrates. A slight saving in electricity and time could also be effected by the
use of commercially available bleach solutions, which are 5 or 10 percent by
weight of sodium hypochlorite, but the cost of the bleach itself would tend to
offset these savings. A more concentrated hypochlorite can also be made by
reaction of household washing soda (sodium carbonate decahydrate) with
powdered 70% chlorine pool bleaches, e.g., HTH brand. Solutions of alkaline
earth chlorides or hypochlorites tend to electrolize less efficently.
Overall, this writer found the biggest drawback of home production of
chlorates and perchlorates to be economics rather than chemistry, time or
safety factors. At today's outrageous prices for electricity delivered to the
home, often 25 cent to 32 cent per kilowatt-hour, and considering that, at 3.6
volts, one pound of sodium chlorate requires about 2.4 kilowatt-hours to
produce, the current retail prices don't seem quite so bad. And it takes about
3.6 to 4 kilowatt-hours to produce a pound of perchlorate, so the cost is even
more, in contrast to the usual market price of this chemical. The home
production of chlorates and perchlorates should perhaps be considered a
desperate last resort for the day when all legitimate and practical sources of
the commercial products may be cut off. It's really much more fun making the
compositions than the chemicals!
REFERENCES FOR FURTHER STUDY:
1. Kirk & Othmer (Editors) ENCYCLOPEDIA OF CHEMICAL
TECHNOLOGY (Second Edition), interscience Publishers, New York, 1963,
Volume 8, "Chlorine Oxides". (an excellent guide to this topic)
2. Hampel, Clifford A. (Editor) ENCYCLOPEDIA OF
ELECTROCHEMISTRY, Reinhold Publishing Co., New York, 1964. (much
practical information on lead dioxide electrode manufacturing, chlorates,
chlorine manufacturing and perchlorates)
3. Mantell, C. L., INDUSTRIAL ELECTROCHEMISTRY, McGraw Hill, New
York, 1931. (although dated, the book is filled with all kinds of useful
information on the manufacture of oxidizers by electrolytic means)
4. Milazzo, Giulio, ELECTROCHEMISTRY: THEORETICAL PRINCIPLES
AND PRACTICAL APPLICATIONS, Elsevier, New York, 1963. (probably the
best and clearest guide to the chemical and practical considerations of cell
variables; should be available in public libraries)
5. Schumacher, J. C., PERCHLORATES, Reinhold, New York 1960. (a
thorough, monograph-like treatment, with much information on electrolytic
cells, hardware and processes, etc.)
[Editor's note: we are always happy to print such well researched and
excellently-written articles as the one above. We hope to have comments
from readers and possible suggestions as to other means of side track in the
oxidizer shortage than Jim's "brute-force" method.]
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pjig
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Is it really too costly to produce your own chlorate v.s. buying it ?
I was under the impression that the cost was a bit less. Most people can get bulk starting materials, but the cost of electricity is the big cost
factor. SO ..., what does that come out to, say on a lb for lb cost basis then?
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Laboratory of Liptakov
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anodes on NaClO3 and 4
| Quote: Originally posted by The WiZard is In |
The home
production of chlorates and perchlorates should perhaps be considered a
desperate last resort for the day when all legitimate and practical sources of
the commercial products may be cut off. |
OK, Judgment Day has come. It's time to start the electrodes.
Development of primarily - secondary substances CHP (2015) Lithex (2022) Brightelite (2023) Nitrocelite and KC primer (2024)
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metalresearcher
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Rather easy.
I have a 5V 40A power supply ordered at amazon.de for $30. And I use an MMO anode.
Works very nice, I ran the cell (600ml) at 16A and it got 90 C.
I used 'lo salt' which consists of 2 parts KCl and 1 part NaCl. Filtering the raw crystallized KClO3 followed by rinsing it with ice cold water
yileded Na-free KClO3 as the flame color was lilac. Even the slightest trace of Na salt (even 0.1%) result in yellow discoloration of the flame.
https://www.youtube.com/watch?v=IlijSQsZ0mI
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Laboratory of Liptakov
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KClO3
Really pure KClO3. Purple flame color. Good work. Congratulation...
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Laboratory of Liptakov
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anodes
1) And still question: multi metal oxide (MMO) anode has how final surface ?
2) Important is good experience with really use anodes. Because e-bay can be fake, false... your recommend link for buy please ?
Thanks.........LL
Development of primarily - secondary substances CHP (2015) Lithex (2022) Brightelite (2023) Nitrocelite and KC primer (2024)
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MrHomeScientist
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Just out of curiosity, what are people making such large quantities of chlorates for? Resale? Rocketry?
I made a few small batches for particular experiments (and the flaming gummy bear demo) but don't have much use for it beyond that.
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