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Rosco Bodine
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| Quote: | Originally posted by Xenoid
Yeah, I have some potassium phosphate, and I think I can get some Sb and Bi chemicals but at the moment I just want to see if the anode holds
together!
We can try adding all the fancy chemicals at a later stage.
I'll use the 2nd anode on some NaCl solution assuming the first one holds together, and makes perchlorate. 
Xenoid |
It should make perchlorate , but you have to hold your mouth right for it to work 
Bismuth fishing sinkers are in the stores . You could make a nitrate from that . It's only a catalyst , makes the efficiency go up a bit .
The phosphate probably ought to go in for either case .
It possibly protects the anode and maybe increases the
efficiency similarly as does persulfate or fluoride for the PbO2 anode or it could be just a wetting agent , I don't know .
[Edited on 26-6-2007 by Rosco Bodine]
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Xenoid
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Shedding and Shredding
Sorry folks, total and utter failure. 
Obviously I didn't hold my mouth the right way.
Set up a nice little 800 ml cell, specially designed for testing anodes. Anode - cathode separation (centre to centre) was 40 mm. Cathode was Ti.
130mm of the anode was in the electrolyte with a surface area of about 39 cm2. I used my lab power supply which has a max. output of 4 Amps. Because
of this current limitation I only used about 77 mA/cm2 for total current of 3 Amps. The lab supply was operated in constant current mode and for 3
Amps a voltage of 5.4 V was required. Oxygen (and some ozone) were evolved from the anode and there was a whispy pink swirl in the solution for a few
minutes. After about 10 minutes, black flakes (some as large as 5mm x 5mm) started to drop off. 
I left the cell running for about another 5 mins, then turned off the current and lifted the anode out, it had slightly pink droplets on the end.
I put the anode back and let the cell run for an hour, by which time there was a pile of MnO2 flakes under the anode, the electrolyte had turned
brownish and black carbon particles were floating around!
At this point I turned off the power, interestingly however the anode continued to evolve ?oxygen and shed material for about another hour, even when
the electrodes were shorted out! Not sure what was going on at this point, perhaps I have discovered a material which is capable of converting
chlorate into chloride and oygen - that must have heaps of practical applications - I could become rich!! 
I still think the Dip 'n Bake technique holds a lot of promise, perhaps not on gouging rods however! I still havn't finished the rod that I was just
surface coating - it seemed to be forming a much smoother coat, so I'll try that next, and then move onto some other material......
Xenoid
[Edited on 27-6-2007 by Xenoid]
[Edited on 27-6-2007 by Xenoid]
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Rosco Bodine
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It is surprising the coating crapped out that quick .
I am wondering if it is sloughing off loosely adhered
oxide or is it the carbon disintegrating or both .
The erosion may slow down as it goes if it is just losing
outer layers and gets down to more adhered material.
Maybe let it run for awhile longer to make sure .
pH could be an issue too . I was thinking about the phosphate and they are commonly used as buffers ,
so that could be what the phosphate was all about .
When you short the electrodes , you have created a shorted battery of sorts . Leave it open circuit when it is off .
Try baking the second one hotter and longer before you try it . Maybe a finish bake for an hour at 400C or much longer at a lower temperature . The
pink swirling could mean you still have some , too much manganese nitrate that didn't change to dioxide . That possibly could be the whole problem .
It could be the sulfate impurity from the original material synthesis too . Know you gotta love hearing that .....but it is possible . It is stated
in one of the baked
beta manganese dioxide patents that sulfate impurity can kill
the integrity of the layer due to side reactions favoring lower
oxides ....and that could also be the source of the pink swirl .
Even at the best , not all of the nitrate converts to the dioxide , but it is a slightly substoichiometric dioxide with
some monoxide which gives it the added conductivity of a mixed oxide . If the proportional conversion on pyrolysis
is off ....then the required composition is not produced .
The cobalt should work better as it is specifically described effective on graphite , but it would seem like even the manganese should have worked
better than it did .
The whole idea of the "activated graphite" is ambiguous
since there are no definite specific examples for the graphite substrate being used at the higher current densities of a perchlorate cell , and on
reviewing the patent , the duriron
substrates fared a lot better than graphite and all of the test anodes were being used at a very low current density compared with what a perchlorate
cell would encounter .
There may have to be an indigenous metallic oxide on the surface of a substrate to provide a transitional bonding layer
for a baked on oxide , and with carbon there is not that oxide interface which would provide adhesion . To use graphite , maybe a thin film of some
intermediate layer which will stick to carbon better is necessary , similarly as is the case for some oxide layers on titanium where only a few
materials can be applied directly to the titanium to form an adherent and conductive interface . Or perhaps some doped
or mixed spinel is required for good adhesion to the graphite , that is that manganese dioxide alone won't work ,
but some mixed oxide may be required for a graphite substrate .
It may be that for the current densities required of a perchlorate anode that only a cold process electrodeposition
can provide an adherent and conductive and impervious oxide layer , and that any baked on coating simply has a porosity that makes it vulnerable .
[Edited on 27-6-2007 by Rosco Bodine]
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dann2
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| Quote: | Originally posted by Xenoid
Sorry folks, total and utter failure.
Obviously I didn't hold my mouth the right way.
Set up a nice little 800 ml cell, specially designed for testing anodes. Anode - cathode separation (centre to centre) was 40 mm. Cathode was Ti.
130mm of the anode was in the electrolyte with a surface area of about 39 cm2. I used my lab power supply which has a max. output of 4 Amps. Because
of this current limitation I only used about 77 mA/cm2 for total current of 3 Amps. The lab supply was operated in constant current mode and for 3
Amps a voltage of 5.4 V was required. Oxygen (and some ozone) were evolved from the anode and there was a whispy pink swirl in the solution for a few
minutes. After about 10 minutes, black flakes (some as large as 5mm x 5mm) started to drop off.
I left the cell running for about another 5 mins, then turned off the current and lifted the anode out, it had slightly pink droplets on the end.
I put the anode back and let the cell run for an hour, by which time there was a pile of MnO2 flakes under the anode, the electrolyte had turned
brownish and black carbon particles were floating around!
At this point I turned off the power, interestingly however the anode continued to evolve ?oxygen and shed material for about another hour, even when
the electrodes were shorted out! Not sure what was going on at this point, perhaps I have discovered a material which is capable of converting
chlorate into chloride and oygen - that must have heaps of practical applications - I could become rich!!
I still think the Dip 'n Bake technique holds a lot of promise, perhaps not on gouging rods however! I still havn't finished the rod that I was just
surface coating - it seemed to be forming a much smoother coat, so I'll try that next, and then move onto some other material......
Xenoid
[Edited on 27-6-2007 by Xenoid]
[Edited on 27-6-2007 by Xenoid] |
Hello Xenoid,
Commiserations.
Not even a big black mess, but a pink one. They say a change is as good as a rest.
What compound gives a pink colour?
Dann2
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Rosco Bodine
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The pink is evidence of Mn++ , and Mn++ shouldn't be there . It is evidence that not all of the (II) Mn nitrate
was completely decomposed by the pyrolysis .....
not hot enough for long enough to do the job of
oxidizing it to the (IV) Mn of MnO2 .
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12AX7
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I would be prone to suspect Mn(III)...Mn(II) is a really weak color.
What of Mn2O3? It's the most stable one I think. You can probably melt it in a graphite crucible. It should be pretty easy to pyrolyze your coating
to it at a thousand degrees or so. If nothing else, the heat would be good for it. I get the feeling you really just need some diffusion bonding to
get it to stick together nice.
Tim
[Edited on 6-27-2007 by 12AX7]
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Rosco Bodine
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Maybe an initial electrodeposition of MnO2 followed by the baked on coatings would have better adhesion .
After further reading ...
It's not too surprising that the baked MnO2 would encounter adhesion problems on the graphite ,
as baked MnO2 doesn't stick to titanium either , even with
the potential there for formation of a mixed oxide bond ,
the MnO2 alone won't attach securely . If a cobalt
oxide is added then it does stick , but the cobalt oxide
spinel is the glue that makes it work , so it is more than
just a dopant for the MnO2 . It possibly requires a
surface oxide interface for the baked cobalt to stick .
And only the electrodeposited cobalt suboxide is actually
reported to stick to graphite . Adding a baked layer
of the ZnCo2O4 spinel on top of that might work , or
the heat might destroy the adhesion of the cold formed layer . The ZnCo2O4 baked on spinel might stick to graphite as an initial coating , if it
doesn't rely upon bonding with only an oxides layer . Anyway the , bimetal spinel of zinc and cobalt appears to be very good for a baked coating on
titanium , and it might go onto graphite okay alone . Or it might work mixed with a fibrous alumina binder for reenforcement . There may be some
different specific conductive spinel type compound which has a better adhesion for graphite .
There's plenty of room for experiments with graphite .
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Xenoid
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| Quote: | Originally posted by Rosco Bodine
Maybe an initial electrodeposition of MnO2 followed by the baked on coatings would have better adhesion .
After further reading ...
It's not too surprising that the baked MnO2 would encounter adhesion problems on the graphite ,
as baked MnO2 doesn't stick to titanium either , even with
the potential there for formation of a mixed oxide bond ,
the MnO2 alone won't attach securely . |
That was my reasoning for doing the vacuum impregnation / baking cycles, I figured MnO2 crystallites forming in the pore spaces would have a good
mechanical bond to the carbon rod, and the surface coating would be bonded to these. Sort of like having "fine penetrating roots" to use a gardening
analogy! Also, because the gouging rods are so crappy (full of holes and pores) this might actually be an advantage over high quality graphite.
I had previously tried plating MnO2 at 5 mA/cm2 using the recipe on Dann2's pages (MnSO4 in dil. H2SO4) but it did not adhere, in fact it didn't even
seem to form!
I am currently giving the 2 remaining rods extra coats at 320oC for 30 mins.
Don't hold your breath.......
Xenoid
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Rosco Bodine
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I don't know if it would be of any use in this application , but there is a conductive bimetal manganese chromite spinel that can be formed in situ I
am pretty sure from manganese nitrate and ammonium dichromate .....followed by heating to ignition of the dried intermediate which undergoes an
exothermic decomposition to leave a glowing hot residue of the combined oxides . It may just be a good way of blowing up the porous carbon rods when
you want to try something radical 
When I was reading about this , it struck me that it might be possible to use this reaction as a heat source in some sort of mixture of oxides which
might form a melt from its
own heat of reaction .....and possibly making a conductive ceramic anode would be as simple as pouring the granular precursor in a mold and baking it
up to the ignition temperature .....or maybe using some thermite as a primer and lighting the fuse to produce a self casting
ceramic anode
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Beaters
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A couple of questions:
1) Anyone ever try the type of carbon rods used in arc lights (search lights, some cinema projectors)? I have a few of them. I'm not sure how they
were made but they are quite hard, about 1/4 inch diameter and maybe 20 cm long.
2) Any thoughts concerning the merits of various types of power supplies for small scale chlorate / perc production regardless of the choice of
anode(s)?
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Xenoid
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| Quote: | Originally posted by Beaters
A couple of questions:
1) Anyone ever try the type of carbon rods used in arc lights (search lights, some cinema projectors)? I have a few of them. I'm not sure how they
were made but they are quite hard, about 1/4 inch diameter and maybe 20 cm long.
2) Any thoughts concerning the merits of various types of power supplies for small scale chlorate / perc production regardless of the choice of
anode(s)? |
I guess they would work (for chlorate), but they may have fillers to harden them or produce more light.
For a first experimental cell, carbon anodes from a lantern battery (zinc-carbon not alkaline!) are a good choice, so you can learn about all the
problems of running a chlorate cell. I searched around all the local recycling centres, pulling the old batteries out of the lanterns. These carbon
rods are already impregnated with a waxy material. They are a bit short so you will need a fairly low plastic food container or similar, holding about
500 - 600mls. My first cell had four and later eight rods in a circular pattern, surrounded by a stainless steel cathode. The carbon rods were
inserted through 6mm rubber grommets and the electrical connections were made using fuse clips or tool clips, if the brass tops are OK you can solder
directly to them. Run the cell at about 2 Amps. Make sure you have an ammeter (the 10Amp setting on a small cheap digital multimeter is fine)
otherwise you won't have a clue whats going on with cell. The best bet for a small cell like this is a small variable lab power supply otherwise a
6volt transformer with a bridge rectifier, run of a variac, but you probably don't have a variac! A small cell like this only requires about 3.5 volts
to produce a current of 2 Amps, so if you hook it up to a 5 volt computer power supply it will probably vaporize. You will need a high wattage resistor to drop the voltage. (Use Ohm's law). You can use nichrome wire from an old
heater or toaster, or carbon rods wired in series, ( a stripped gouging rod has a resistance of about .1 - .15 ohms).
Most importantly, read all the websites dealing with home chlorate production.
[Edited on 29-6-2007 by Xenoid]
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Beaters
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Thanks for the feed-back. Nice photo.
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Xenoid
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MnO2 coated gouging rods - the final indignity!!!
Greetings,
Well I gave my remaining two gouging rods extra coats at higher temperatures - taking on board Rosco's advice.
I won't bother going into details, suffice to say it was a large number, the final coats were baked for 30 mins. at 320oC.
For the electrolysis, I used 500g/litre pure NaClO3, with a Ti cathode.
I tested the impregnated rod first, starting at a cell current of 1 Amp (25mA/cm2) voltage required was 3.8V, Hmmm... looked good, (It had that
characteristic low oxygen production and small bubbles, like my Pt anodes). Let it run for about 20 mins no flaking!
Turned the current up to 2 Amps (51mA/cm2), Voltage 4.5V. Solution had a slight pink tinge. I let it run for 40 mins, still no flaking, but solution
was now distinctly pink!
Turned up the current to 3 Amps / 4.9V ( about 75mA/cm2) solution now rosy pink! One hour later and the anode was shedding.
I then connected the rod that was just coated (not impregnated) it had a much smoother surface. I ran it at 2 Amps / 4.1 Volts but after 6 hours it
too had started shedding, by this time the electrolyte was a rich, rosy bergundy colour, very pretty!
Whilst they were working, the anodes seemed to behave in a similar fashion to Pt. With a tripling of the current there was very little increase in
oxygen production, indicating that the current was going into perchlorate production, (greater efficiency at higher current densities). Does that
sound right?
I can understand why the anodes are failing mechanically, and shedding flakes of MnO2 and carbon, but why am I getting so much Mn++ in solution, I
don't believe it is all coming from unconverted Mn(NO3)2.
WARNING - Anyone messing with MnO2 anodes, it is a real chore to get rid of the Mn from your chlorate, the solution keeps precipitating brown and
black crud during the concentration and crystallisation phases, despite much filtering and removal of the bulk of the Mn by precipitating with
Na2CO3!!
Xenoid
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Rosco Bodine
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Hmmm...I'm starting to wonder if your pink might actually be formation of *per*manganate ???? Maybe if you go above a certain voltage , it is like a
breaking point ??
You follow what I mean , you may be overdriving the anode and actually oxidizing the anode itself instead
of what you are trying to do the work on that is in the electrolyte .
It could be that the electrode gap being wide is causing you to use too much voltage to get good current ,
pushing the voltage up to a point that the potential is actually attacking the anode coating from underneath ,
at the carbon to MnO2 interface ? ...or at any rate
wherever it is occuring , maybe the overvoltage is what is killing the coating and you have to run the gap much closer to keep down the voltage much
lower .
[Edited on 1-7-2007 by Rosco Bodine]
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Xenoid
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| Quote: | Originally posted by Rosco Bodine
Hmmm...I'm starting to wonder if your pink might actually be formation of *per*manganate ???? Maybe if you go above a certain voltage , it is like a
breaking point ??
[Edited on 1-7-2007 by Rosco Bodine] |
So if there is sodium permanganate in the solution, and I add NaCl, it will be oxidised to NaClO3 and MnO2 will be precipitated, according to my
chemistry book.
Xenoid
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Xenoid
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Permanganate!
Yeah!
Just tried this, made the solution alkaline by adding NaOH, and added NaCl, the distinctly pinky-purple colour has disappeared, and browny/black crud
MnO2 has precipitated!
So where to from here, shall I make another anode and try it at a lower voltage - what is the limiting voltage!
Xenoid
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MadHatter
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Anodes
http://cgi.ebay.com/Potassium-Perchlorate-Machine_W0QQitemZ2...
This eBay seller claims the machine will produce KClO4 but is offering COPPER anodes.
I can't believe it would work.
A couple of sellers are offering platinum-coated anodes for $95-100 each. Too rich for
my blood.
I'm not sure if we'll ever find the 'perfect' anode. I made some more 'perc' recently.
I had a thin coating of PbO2 over the gouging rod this time around but even that was
torn up by the electrolysis. The perchlorate cell has to be one of the most hostile
environments there is in electrochemistry. Pinholes in the PbO2 layer are probably the
cause of this destruction. It's like a shark smelling blood after the 1st bite, and the
feeding frenzy occurs.
BTW, a snapshot of the gouging rods I use before peeling the copper plating.
[Edited on 2007/7/1 by MadHatter]
From opening of NCIS New Orleans - It goes a BOOM ! BOOM ! BOOM ! MUHAHAHAHAHAHAHA !
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Rosco Bodine
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limiting overvoltage found charted in CRC !
@Xenoid
This relates precisely to something brought up about
ripple voltage and electrode potentials causing mixed products in another thread , and it seems the anode material itself can certainly be affected
also . Never really thought about this much , but it does make sense .
Thinking back on it , there is an engineering parameter
concerning a sort of "anode depassivation voltage" ???...but I can't remember the correct/exact term for it or even where I read it .....I think it
was in one of the cathodic protection patents where the parameter was mentioned for duriron anodes that perform excellent up to a certain
"depassivation?? voltage" which is the limit
for noncorroding operation .
I have not a single Electrochemistry book to reference
and find what is the minimum ( anode ) oxidation potential on a chart for the formation of permanganate ,
but evidently that's going to be the defining limit for
operation of an MnO2 anode coating to keep it from entering the reaction itself .
You might have to close the electrode gap down to 3 mm or so , where there is a risk of the thing shorting out from any flaking , so you will
absolutely require some sort of active current limiting PS to babysit the process .
As an interesting side note ...MnO2 can also be formed
as a product from Mn(NO3)2 and permanganate ,
and IIRC the MnO2 first forms as an unstable hydrated sol / gel system and then precipitates a nanocrystalline
*beta* MnO2 in a similar way as can magnetite be gotten
as a nanocrystalline form by chemical means . These
processes are generally applied to precipitation on
carriers for use as catalysts following calcination to sinter
the carrier to the carried particles of catalyst , and are also
processes for ink and paint pigments ....but might have some adaptive use to electrode manufacture as well .
The only problem I can see is that with using Mn(NO3)2
and an alkali permangante is the KNO3 byproduct could
cause ignition of the carbon on baking , so another Mn
salt like an acetate would probably be a better choice if
it reacts the same way with permanganate . This might
eliminate a lot of baking time and achieve a more complete conversion at a smaller grain size for the MnO2 .
Never tried it and never read anything specifically describing this for anode application ....but haven't really searched for it either , so it might
already be a known
technique , or something for a logical experiment .
For the cathode , here's a thought . Maybe get or make some plastic washers as a spacer for a coaxial arrangement
of bare wires spaced parallel to the center anode , put about
eight of them around it . Make the holes in the spacer washer for the wires a little small so it is a press fit to hold
them like a birdcage around the anode , very closely gapped
like 2-3 mm . Or you might use PVC fittings as a spacer for
some sort of metal tube as the cathode with the anode spaced inside it with three point support at 120 degree spacing to keep the anode centered
coaxially top and bottom .
Edit : CRC gives the following reaction potentials :
ClO4 + 2H + 2e <----> ClO3 + H2O 1.19 volts
MnO4 + 4H + 3e <----> MnO2 + 2H2O 1.679 volts
I interpret that as meaning if you go one half volt too high ,
your MnO2 anode is toast .
Your overvoltage limit is 0.489 volts ( above the minimum 1.19 volts for perchlorate production ) so you only have ~ 0.4
volts allowable increase beyond where perchlorate production begins , before you start also destroying the anode .
[Edited on 1-7-2007 by Rosco Bodine]
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Xenoid
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Damnation!
Now the black/brown precipitate has settled, the solution is actually a bluish/green!!!
Manganate I assume, how do I get rid of this!!
Re. above voltages, the cell starts to conduct at about 2.4V, thats with a 40mm electrode spacing, so maybe 10 - 20mm spacing would be OK, but at
quite low currents. Less than 5mm and the electrode would be just about in the hydrogen stream coming of the cathode.
Xenoid
[Edited on 1-7-2007 by Xenoid]
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Rosco Bodine
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chlorine evolution potential
You need to get some circulation flow going , to strip that plume of hydrogen down to a thinner section so you can close up the gap for an MnO2 anode
...as a wide gap is intolerable given the narrow voltage gradient between
perchlorate production and anode destruction . This is simply an inescapable materials limitation which is going
to apply to MnO2 being used as the anode material in a perchlorate cell .
And as if the precise voltage for ClO3 to ClO4 is not exacting enough .....let's see what else applies as a possible complication . If I understand
this correctly , the plot is thickened by possible unwanted cathode reactions which can reverse the direction of desired reactions if the cell voltage
and pH isn't just right .
This is interesting , there are a couple of possible
chlorine discharge potentials listed in CRC that are *intermediate* values between the two above , which
I will show again for location in the range and it appears
that the both the ClO3 and ClO4 ion can actually be reduced again all the way back to chlorine in an undivided cell where such ions are exposed to the
cathode . It would seem an acidic pH would favor this undesired result , as some alkalinity would immediately react with the molecular or ionic
chlorine and keep it in the reaction . However that might also
encourage the undesired reaction so this could be only minimzed by careful voltage control and a near neutral pH ,
in an undivided cell .
ClO4 + 2H + 2e <----> ClO3 + H2O 1.19 volts
ClO4 + 8H + 7e <----> 1/2Cl2 + 4H2O 1.34 volts
ClO3 + 6H + 5e <----> 1/2Cl2 + 3H2O 1.47 volts
MnO4 + 4H + 3e <----> MnO2 + 2H2O 1.679 volts
Concerning your evidence of some Mn++ residue that would make sense , because following the oxidation of the MnO2
to MnO4 at 1.679 volts , MnO4 itself is further reduced
according to :
MnO4 + 8H + 5e <----> (Mn++) + 4H2O 1.491 volts
It is a certainty you would find Mn++ compounds in the electrolyte after disintegration of an MnO2 anode , and the
permanganate was something that existed only as a transient intermediate species during the MnO2 decomposition .
Electrochemistry is one of my lesser areas of understanding ,
so if I have gotten this wrong or misunderstand the values
interpretation , anyone feel free to speak up .
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DerAlte
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@ Rosco - electrochemistry is understood by academics, but electrolyses only by artists! Don't belittle yourself: quote " Electrochemistry is one of
my lesser areas of understanding..." I'ms ure you have a better grip on the theory than most. Being a retired EE, I am always fascinated by
electolytic methods, whose power can exceed any chemical means. Think of fluorine or Li,K, Cs, Rb etc. it's the only way to go.
DerAlte
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Rosco Bodine
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The voltages given above are not absolute cell voltages
like you would read across the power supply leads with a voltmeter , but are difference voltages compared to a
"standard hydrogen electrode" which is a deliberately polarized platinum electrode being operated in a 1 M HCl
electrolyte as a laboratory "standard" and how standard it is seems to depend on whose laboratory is making and reading it ......since every printed table of redox potentials seems to vary from the next
. However
the tables in CRC were all done by the same person in the same lab using the same equipment , so their relative
values show something useful , in terms of the "spread"
of measured voltages which have detectable effects .
The "absolute voltage" of the standard hydrogen electrode is not so absolute from one lab to another ,
except as the reference used for experiments in that
particular setup . But I have seen typical values for
that absolute voltage ranging from 4.2 volts to 4.8 volts ,
and while experimental conditions relate to one molar electrolytes , obviously the concentration differing and
the electrode spacing would affect the actual cell operating voltage and the actual discharge potentials ,
as I understand it anyway . So these charted values
for electrode potentials are ambiguous values which
do not directly translate to a cell of different design
and different content .
Neither is the catalytic effect of electrode materials
which may be specific and preferential for certain reactions addressed by such tables , where the actual production cell may be using electrode
materials and
process catalysts in the electrolyte also which can shift
the standard potentials several tenths of a volt in
a favorable direction for selectivity of the desired product .
So considering these things , must be done while "reading" such tables . What I get from such
reading of the table above is that the desired cell operating point would be found by slowly increasing the voltage until chlorine evolution is barely
detectable as a byproduct and that would correspond to reaching an
electrode potential for your cell , which is parallel to
the second reaction above . Regardless of the absolute numbers , you know then that the voltage reading for your cell and your electrolyte has
reached a level where
some of the perchlorate being produced at the anode
is being reduced again at the cathode as evidenced by the chlorine being evolved .
This is a point where certain electrolyte additives can form a film on the cathode which obstructs contact of the electrolyte with reactive ionic
hydrogen which reduces
the perchlorate as a competing reaction , favoring the
ionic hydrogen instead combining to form molecular hydrogen and escaping as bubbles unstead of reducing
the just formed perchlorate which is desired to remain intact . This additive could be dichromate and magnesium chloride in very small amount , like
0.1 g per liter , and there are specific combinations of different materials used depending upon what is the electrode materials used in the
particular cell , materials which basically form a conductive semi-solid permeable membrane film on electrodes and steer the reaction towards the
desired products and/or interfere with production of the undesired byproducts . Such additives
emulate the function of a divided cell by forming a porous
barrier layer on the surface of the electrode , interfering
with the surface chemistry which would otherwise occur there on the cathode particularly , by shielding the nascent
ionic hydrogen from contact with the electrolyte just long enough for it to spontaneously combine with itself to form much less reactive H2 .
[Edited on 2-7-2007 by Rosco Bodine]
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Xenoid
National Hazard
Posts: 775
Registered: 14-6-2007
Location: Springs Junction, New Zealand
Member Is Offline
Mood: Comfortably Numb
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| Quote: | Originally posted by MadHatter
A couple of sellers are offering platinum-coated anodes for $95-100 each. Too rich for
my blood.
[Edited on 2007/7/1 by MadHatter] |
I just ordered a 60mm x 90mm platinised titanium electrode from Palloys;
http://www.palloys.com.au/category7_1.htm
It was about $81 Australian, but I just got an email to say that the price had gone up and the new price was A$97 + A$5 postage (their supplier had
increased the price).
Looks like there are too many people out there making perchlorate.....
They have a smaller one for about A$50, not sure if the price has gone up on this one, or not.
Xenoid
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garage chemist
chemical wizard
Posts: 1803
Registered: 16-8-2004
Location: Germany
Member Is Offline
Mood: No Mood
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Have you tried the thermal disproportionation of chlorate at least once before spending so much money on a perchlorate anode?
Today, another member on the german forum has tried my method of thermal chlorate disproportionation and found it to work very well. He used KClO3
from a graphite anode chlorate cell which was still wet and dirty.
The thermal disproportionation tolerates any amount of chloride contamination and seems to be less sensitive to heavy metals as I previously thought
(there was slight MnO2 contamination in the KClO3).
You can even recycle the KCl byproduct by boiling down the filtrate from the residual chlorate destruction step.
I keep repeating that the thermal disproportionation works well and gives reasonable yields without the need for pure chlorate and any anodes.
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Eclectic
National Hazard
Posts: 899
Registered: 14-11-2004
Member Is Offline
Mood: Obsessive
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It sort of depends on whether your primary interest is in the journey, or the destination, grasshopper.
May the electromotive force be with you.
[Edited on 7-2-2007 by Eclectic]
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