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woelen
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It can be done without, but the yield will be lower. When dichromate is present, then at the cathode, water is reduced to hydrogen and hydroxide ions
are formed. When no dichromate is present, then the bromate is reduced back to bromide to a fairly large extent, once a certain concentration is
reached. I did the experiment, however, without dichromate as well, and the yield still was acceptable. Also, I think that with a carbon cathode this
effect of back-reduction of bromate to bromide is less severe than with a metal cathode.
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Zinc
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Be careful when working with bromates because they are carcinogenic.
http://en.wikipedia.org/wiki/Bromate
(Yay! 100 posts!)
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woelen
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I now finally have platinum wire and I can only say WOW about the difference with graphite rods.
I used a big titanium cathode and a platinum anode (9 inches of 0.3 mm wire + 4 inches of 0.4 mm wire). The solution remains perfectly clear, and most
important, the chromium of the dichromate is not converted to the +3 oxidation state. I did some more tests with the graphite, and I suspect that with
a graphite anode, the chromium first is converted to some peroxo-compounds, and that in turn is converted to chromium in the +3 oxidation state. The
KBrO3 I obtained with the platinum anode is really white, while the KBrO3 with the graphite anode is light green.
I have noticed, that chromium (III) is amazingly difficult to separate from bromate. Even repeated recrystallization does not remove chromium (III),
while even without recrystallization I had no chromium (VI) in my rinsed product (although it still contained bromide, but that is easily removed by a
single recrystallization). Apparently, chromium (III) co-crystallizes with bromate, or maybe it even is coordinated to bromate, making separation very
difficult.
Right now, I have two batches, one light green sample and one purely white sample. I will make a web-page on this subject in due time, and then I will
post the link over here.
The only drawback of using platinum is the lower efficiency. With the same number of ampere-hours, I get less KBrO3 with platinum. I don't know why.
Maybe because chromium (III) increases the process, and prevents back-reduction of bromate at the cathode better than chromium (VI)? Another reason
could be the much higher current density at the very thin wires, compared to the thick graphite rods, resulting in formation of oxygen, instead of
bromine. I could not look very well at the anode, because of the vigorous bubbling at the cathode.
[Edited on 19-11-06 by woelen]
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woelen
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Here is the promised web page:
http://woelen.homescience.net/science/chem/exps/KBrO3_synth/...
EDIT: Edited URL, such that the link works again
[Edited on 7-11-12 by woelen]
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pyrochem
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Can permanganate be used instead of dichromate to prevent bromate reduction?
If it's only partially reduced, it could form manganate, which might be hard to remove.
[Edited on 30-11-2006 by pyrochem]
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not_important
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Woelen, a thought on cleaning up the bromate.
Make a small amount of Al(OH)3 by adding KOH or K2CO3 solution to the solution of an aluminium salt. Some Al foil dissolved in HCl sould be OK. Wash
the ppt with water once or twice, then dissolve in a small amount of concentratated KOH solution.
After dissolving the crude KBrO3 in hot water, add the potassium aluminate solution, then blow some air through it or bubble some CO2 into it. The
aluminate becomes K2CO3 and Al(OH)3, in the slightly alkaline solution Cr(III) should follow the Al(OH)3 down so both can be filtered off. It might
even help trap some of the particulate carbon. Might need to gently boil it a bit to get the Al(OH)3 into a nice filtratable form.
The amount of aluminium need should be small in comparison wit the amount of KBrO3, but should be much more than the amount of dichromate added.
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YT2095
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does anyone know if LiBrO3 is hygroscopic or not?
my LiBr arrived this morning, and I plan on setting up a small bromate cell to do 50g or so.
however I may have to rethink if it`s like Copper Chlorate (that stuff never crystalises if made wet).
the solubility data shows that the Bromate is much more soluble than the Bromide, and that`s all the data I have.
Edit: another thought occured to me, since LiOH will exist during electrolysis, will it be safe in a Glass cell?
[Edited on 12-5-2007 by YT2095]
\"In a world full of wonders mankind has managed to invent boredom\" - Death
Twinkies don\'t have a shelf life. They have a half-life! -Caine (a friend of mine)
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woelen
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Making LiBrO3 in this way will be very hard. I'm almost 100% sure that crystallizing the LiBrO3 will be a real pain, it will be extremely hygroscopic.
LiBr already is extremely hygroscopic, LiBrO3 probably is more so.
I studied the properties of bromates, because I wanted to do some flame color experiments with them, and I found the only bromates, which are easy to
prepare are
KBrO3
Ba(BrO3)2
RbBrO3
CsBrO3
AgBrO3 can easily be made by adding a solution of AgNO3 to a solution of KBrO3. It is insoluble in water and precipitates as a coarse, easy to filter
precipitate.
NaBrO3 can be made in solution, but crystallizing this is hard. Crystallizing that also is not that interesting. However, you could make a solution of
NaBrO3 (plus remains of NaBr) and use that to make the other bromates, mentioned above (except the silver salt, because that will be contaminated with
the bromide). Getting rid of the Na(+) ions is relatively easy, because of the large difference of solubilities of NaBrO3 and the other bromates in
cold water. Maybe one, but certainly two recrystallizations from distilled water will remove the Na(+) and make it suitable for flame color
experiments.
One bromate, which is borderline is Sr(BrO3)2. All others are very hygroscopic or even deliquescent solids, which are really hard to prepare in a
reasonably pure solid state.
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JohnWW
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Have you examined their explosive properties? The best bromate for that purpose would probably be NH4BrO3, if it could be made. However, like
chlorates, they may be too unstable for practical or commercial use.
Also, having made bromates, how about trying to prepare perbromates? Similarly to perchlorates and periodates in comprison with chlorates and iodates,
they would be more stable than bromates, but more strongly oxidizing than both perchlorates and periodates. But they are supposed to be difficult to
prepare, and were first made only a few years ago, as I understand. Electrolysis of an alkaline solution of a bromate at low temperature would be the
most likely method. Other possible methods could involve direct reaction of bromate with atomic oxygen.
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dann2
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Hello,
There is an artical on Making Bromates in the
"Further reading section" on this page:
http://www.geocities.com/CapeCanaveral/Campus/5361/basechem....
Dann2
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YT2095
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| Quote: | Originally posted by JohnWW
how about trying to prepare perbromates? |
perbromates are not possible to make without the use of radioactive isotopes.
\"In a world full of wonders mankind has managed to invent boredom\" - Death
Twinkies don\'t have a shelf life. They have a half-life! -Caine (a friend of mine)
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Nerro
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I have 75g of superfine silver wire lying around. If I were to coil the wire into a 6 or 12 fold wire could I use it as an elektrode in this synth? I
have no idea how Ag might stand up to the surroundings in this case...
#261501 +(11351)- [X]
the \"bishop\" came to our church today
he was a fucken impostor
never once moved diagonally
courtesy of bash
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Zinc
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A little off topic but is AgBrO3 explosive? I have heared that AgClO3 is but I don't know if that is true. Is it?
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dann2
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Silver wire
| Quote: | Originally posted by Nerro
I have 75g of superfine silver wire lying around. If I were to coil the wire into a 6 or 12 fold wire could I use it as an elektrode in this synth? I
have no idea how Ag might stand up to the surroundings in this case... |
Silver will erode away.
Dann2
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woelen
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Silver wire indeed is useless for this synth. You need either graphite, or even better, platinum.
Making perbromates indeed is very difficult. I also asked this question some time ago, but now I did some research on it myself. The first preparation
involved a radioactive selenate, which emitted a beta-particle to give perbromate:
Se*O4(2-) ---> BrO4(-) + e(-), here Se* is some radioactive isotope of Se.
Later preparations, however were chemical only. I read that the current mode of preparing this chemical is dissolving KBrO3 in a dilute solution of
KOH and bubbling fluorine through this, or adding XeF2 to this. Both are totally out of reach of the home chemist and require real good lab apparatus
and expensive chemicals, not available for the general public. So, having perbromate probably will remain a nice dream for me.
I made some AgBrO3, and this chemical is not explosive. It is easy to prepare from AgNO3 and KBrO3. Dry AgBrO3 actually is quite unreactive, but with
finely powdered metals like Al or Mg it reacts violently, when ignited. I did the experiment, and it is nice to see formation of small silver
droplets. This, however, also is quite dangerous. You must not think of having a small silver droplet on your hand, that must be really painful.
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Antwain
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Massive problems....
I was running my platinum electrode as shown in the platinum wire thread. The current was 1.0A and the cathode was titanium. The surface area of the
anode was at least 5cm^2.
Bubbles formed on the cathode as expected, but falling from the anode was a solution with a brown tinge. This gathered at the bottom. I left it
running for a while hoping that this was some stupidly small amount of iron that had been transfered from the hammering, but then I noticeed that the
platinum wire which was encased in the coil was starting to show through. I left it a little longer, and I am convinced that the erosion continued,
exposing a larger patch of the wire.
Why the fuck is my platinum dissolving? That electrode cost me a lot of money and now it appears to be completely useless. I should add that there was
~0.5g of potassium dichromate in the solution as well.
[Edited on 4-11-2007 by Antwain]
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Antwain
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JohnWW
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| Quote: | Originally posted by woelen
Making perbromates indeed is very difficult. I also asked this question some time ago, but now I did some research on it myself. The first preparation
involved a radioactive selenate, which emitted a beta-particle to give perbromate:
Se*O4(2-) ---> BrO4(-) + e(-), here Se* is some radioactive isotope of Se.
Later preparations, however were chemical only. I read that the current mode of preparing this chemical is dissolving KBrO3 in a dilute solution of
KOH and bubbling fluorine through this, or adding XeF2 to this.(cut) |
Have there been any successful attempts at preparing perbromates electrolytically, like perchlorates?
It has been theorized that argon tetroxide, ArO4, could be similarly made, by the radioactive decay to a stable argon isotope (either Ar-36, 38, or
40) by beta-emission of a neutron-rich isotope of chlorine (it would have to be Cl-38 or Cl-40, because Cl-36, often used as a tracer, is too
long-lived and much of it decays to S-36), as the isoelectronic perchlorate anion, ClO4-. I wonder if this has ever been actually attempted. (Only
ArF2 and HArF have ever been reported to be made by ordinary chemical means.)
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garage chemist
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Antwain, thats just the bromine that is forming on the anode- think about the mechanism of bromate formation by electrolysis, bromine formation at the
anode is the first step! Your anode most likely isnt corroding.
Also, why do you keep your electrodes so far apart? Thats counterproductive since it increases the resistive losses in the cell and therefore
undesirable heat production. Also, the hydroxide ions from the cathode are supposed to react with the bromine from the anode, so put your electrodes
close together!
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Antwain
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@ garage chemist- are you sure that it is my imagination? ie. that a Pt anode will not dissolve under these conditions. I don't want to put them too
close together because it is not that stable and I don't want a short. Would a stirrer bar be good enough?
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garage chemist
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Yes, stirring is certainly helpful. If your anode is pure platinum or even better Pt/Ir it wont corrode, I have made bromate myself with 99% Pt wire.
If your electrodes are not very mechanically stable then have something like 2cm distance between them, but just not how they are now! Have you read
about industrial electrolysis setups? In aqueous electrolyses there is something like 3-5mm (yes, millimeters!) between the electrodes to cut down on
resistive losses, in almost all electrochemical processes, like chloralkali, chlorate and perchlorate synthesis etc...
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Antwain
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Well the plan was to have the glass casings in thermometer fittings, but I broke one a few months ago and didn't realise that my other one was the
wrong quickfit size. By putting these into a 3 necked flask the positions could have been carefully adjusted, but right now it is the old bottle with
holes drilled in the lid trick.
The Pt coin is .9995 platinum, don't know what the impurity is but I would assume precious metal. Ok then, I will give it another go when I have the
time and see what happens. Now that I know for certain that bromine is made the colour doesn't seem so scary. In hindsight this was stupid, but I
somehow mistakenly thought that the entire reaction would take place at the anode so that no bromine would be formed. The dissolved bromine was
sinking to the bottom, but was depleted around the cathode, so that is a good sign too. I was just really worried about losing my electrode - even if
you neglect construction time, that electrode cost me 8 hours of my job (which is boring) and I really didn't want to have to separate the platinum
from solution because bromide is a precious substance for me too at the moment.
On that note, I did try it at several different currents, but my home made water cooled resistor failed dismally, so I had to use a low enough current
that a power resistor could dissipate the heat. I will do it again and aim for a decent current now that I know that brown stuff is not platinum.
Incidently, what do platinum halides/hexahaloplatinic salts look like? my merk does not describe colour. Wiki gives the colours of both PtCl2 and
PtCl4 as dark brown. I may have heard this before, because I was worried by the colour in my electrolysis.
PS. I just remembered. I once dissolved palladium in aqua regia to a deep brown solution. I believe platinum behaves similarly.
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woelen
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Garage chemist is totally right. The brown material indeed is bromine, formed at the anode. It is mixed with hydroxide from the cathode and that
combination in turn results in formation of bromate (through intermediate hypobromite).
I only had thin wires, while you have a nice surface area, so you even have a much better setup than I had. I did not see any corrosion on my anodes
at all, so I expect that in your situation things should be the same. If you do have excessive corrosion, then your anode is not made of platinum.
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Antwain
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Thank you both for setting my mind at ease. Left to my own devices I would probably not have powered it up again.
In this case I have physically damaged my electrode (slightly). I put the warning out.... heating platinum does make it soft and more malleable (this
I already knew, my wire bent much more easily, to the point of nearly worrying me). Also, beating platinum metal even a bit can cause it to tear if it
is thin.
Hopefully it has enough structural integrity left to work, but I will have to be careful manipulating it.
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Antwain
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Ok, I decided to do it now (at midnight). It seems to be working well enough. The bromine is being pulled into the updraught caused by all the
hydrogen and there is no discoloration forming at the bottom. It is running at 3.4A.
Heres the crappy part. I used 120g of KBr which coincidently is 1.0mol. When I did the calculations for a 6 electron oxidation I arrived at the most
astonishing figure of 160Ah, at 100% efficiency. I knew that electrolysis was slow, but thats just taking the piss. That means running it for 2 days.
Anyone who has done electrolytic reactions before is probably familiar with this huge number of columbs per mole, but I have never done it
quantitatively, to get a product.
I suppose that it kind of helps to demonstrate that we shouldn't be surprised that molar quantities of reducing and oxidising agents can explode with
a large release of energy.
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