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garage chemist
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You can also simply replace bromate with chlorate for the bromine production, use 6 mol instead of 5 mol bromide, and use the correct amount of H2SO4:
ClO3- + 6 Br- + 6H+ -----> Cl- + 3 Br2 + 3H2O
The reaction requires some warming and time, in contrast to the bromate + bromide comproportionation which happens almost instantaneously in cold
solutions.
This is even more convenient for bromine production if you can get chlorate easily, but the correct stochiometry is important, otherwise you will have
chlorine impurity in your bromine.
Since bromate is easier for the amateur to produce than chlorate, this method is best suited for those who can simply buy pure chlorates, like me.
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chloric1
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Thanks woelin, you helped me alot. I knew there was a simple trick. When I wrote to you, I already was getting an idea that a solid mix was to be
used simply becuase too much water would prevent bromine from precipitating.
@garage chemist- I am able to make my own chlorate of good quality. I have yet to try bromate making. I may consider your method but I have 50lbs of
NaBr so bromate has 2 advantages for me; no foriegn halogens involves, and getting the most use from my bromide. I believe your synthesis is quite
satisfactory. If any chlorine impurity should be taken into the bromine, then a simple agitation with pure potassium or sodium bromide with the
elemental bromine should suffice.
Fellow molecular manipulator
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garage chemist
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Yes, it would be advantageous yield-wise to use a small excess of chlorate to convert all of the bromide in the first step, and purify the crude
bromine from chlorine impurity in a second step by agitation with an alkali bromide solution.
But any economic preparation of bromine from bromide will invariably use distillation to separate bromine from water, so the bromate method with
mechanic separation of the bromine phase is better for small-scale bromine preparation where yield is less important than convenience and simplicity.
I have put some thoughts into a method on how to extract the bromine from alkali bromides in large amounts- say, 1kg of NaBr per batch (giving about
250ml bromine).
I would go about it like this: put bromide into 2L three-neck flask with distillation bridge and dropping funnel, add the required amount of NaClO3
plus very small excess as an aqueous solution (or KClO3 in solid form), add enough water to make everything into a thin stirrable slurry, heat to
80°C and add conc. HCl (much cheaper than H2SO4) dropwise under magnetic stirring.
As the temperature is above the boiling point of bromine, it distills off as it is formed and is obtained as the distillate together with some water.
The crude bromine is then purified from chlorine and possible chlorine dioxide by stirring with NaBr solution followed by distillation, separated from
co-distilled water via separatory funnel, dried by shaking with conc H2SO4 and redistilled in a dry distillation setup.
The residue in the 2L flask will be almost pure NaCl solution/slurry if NaBr, NaClO3 and HCl were used.
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Antwain
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| Quote: | Originally posted by woelen
Antwain, actually, testing KBrO3 purity is very easy. Just take some of the crystals and add them to just 1 ml of 10% H2SO4 (not HCl, use H2SO4). If
the solution remains colorless, then you have really pure KBrO3, if the solution turns ligt yellow, then you have some KBr in your KBrO3. If the
solution turns orange and you see even some vapor of Br2, then you have very impure KBrO3. Also, try your solid, mixing it with some powdered S and/or
C, and ignite. If you have KBrO3, then you'll definitely notice, even if it is impure  .
If you see a steady stream of bubbles of oxygen at your anode, then it is best to stop. The concentration of bromide then has dropped considerably.
Let the contents of the cell cool down in the fridge in order to get most of your KBrO3 out of it, and keep the liquid for further production of KBrO3
(just add new KBr and maybe some water). |
After some fiddling, it is nowhere near done. No bubbles and plenty of bromine coming off the anode. I think my current effeiency is still shit, but
hey, it was a first attempt. At least it is making bromate.
I have isolated 56.3g so far. I just finished my last test at uni and so went down to try testing the KBrO3.... from memory. So anyway I used
98%H2SO4 and I got bromine, did i ever. I will try it again very shortly
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Antwain
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| Quote: | Originally posted by garage chemist
Yes, it would be advantageous yield-wise to use a small excess of chlorate to convert all of the bromide in the first step, and purify the crude
bromine from chlorine impurity in a second step by agitation with an alkali bromide solution.
But any economic preparation of bromine from bromide will invariably use distillation to separate bromine from water, so the bromate method with
mechanic separation of the bromine phase is better for small-scale bromine preparation where yield is less important than convenience and simplicity.
I have put some thoughts into a method on how to extract the bromine from alkali bromides in large amounts- say, 1kg of NaBr per batch (giving about
250ml bromine).
I would go about it like this: put bromide into 2L three-neck flask with distillation bridge and dropping funnel, add the required amount of NaClO3
plus very small excess as an aqueous solution (or KClO3 in solid form), add enough water to make everything into a thin stirrable slurry, heat to
80°C and add conc. HCl (much cheaper than H2SO4) dropwise under magnetic stirring.
As the temperature is above the boiling point of bromine, it distills off as it is formed and is obtained as the distillate together with some water.
The crude bromine is then purified from chlorine and possible chlorine dioxide by stirring with NaBr solution followed by distillation, separated from
co-distilled water via separatory funnel, dried by shaking with conc H2SO4 and redistilled in a dry distillation setup.
The residue in the 2L flask will be almost pure NaCl solution/slurry if NaBr, NaClO3 and HCl were used. |
You don;t think that HCl would distill across under those conditions?
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Antwain
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Test completed. The powder turned very slightly yellow on entering the solution, but after a quick swirl the solution was clear as far as I could
tell. If there is any bromide left it is much less than 1%
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woelen
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Antwain, well done! I have some NaBrO3 (commercial sample) and that gives a clearly visible light yellow color, when dissolved in dilute H2SO4. My
home-made KBrO3 is better than that NaBrO3, and yours also.
If the solution remains near colorless and you just see a hint of yellow color, then the amount of KBr in it is much less than 1%. A 1% contamination
would make things even orange! You have nice and pure KBrO3 now.
Just try how sensitive this reaction is. Take a spatula of your KBrO3 and dissolve this in some acid. Then add just a few crystals (1 mm3 or so)) of
KBr and you'll get a nice yellow/orange color. If you have done that test, then you will be even more convinced that your KBrO3 is pure (at least with
respect to bromide contamination).
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chloric1
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sodium bromate
I have some (maybe 100 grams) of commercial sodium bromate that I was going to make bromine with via your suggestion. After I used this I will
probably want to make my own. I think sodium bromate might even be easier to obtain pure than the chlorate. I say this based on Merck's claim that
it taked 2.5 parts cold water to dissolve 1 part sodium bromate. Sodium bromide solubility is more like that of sodium chloride and is relatively
constant.
Sodium bromate would be easier to use to make slightly soluble bromates such as the lead and barium bromates.
Fellow molecular manipulator
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Klute
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I'd like to share my results with this electrolysis, thanks to Woelen's excellent write-up and help. I have also noticed a few difference from
Woelen's results when using graphite rods.
I used a Al 901A ELC variable voltage generator, rated at 1A at 1V, 4A at 15V, with a good multimeter for the current, and a cheap one for mesuring
the voltage between the two electrodes (just to follow the resistance of the cell).
I used two graphite rods from a large 4R25 battery, very easy to recover: after opening the case, the battery contained 10 small zinc cells, each
with a brass "plug" on top. Simply taking out the plug with pliers gave the graphite electrode very easily. I kept the MnO2 slurry for a rainy day.
The electrodes were thoroughly washed with dilute H2So4, disted water, then used to electrolyse a brine solution for 10-15min, as advised by Woelen.
A little very fine carbon particlues floated on the surface or stayed at the bottom.
35g of KBr were dissolved in ~60mL, obtaining a saturated solution. To this was added 1mL of a dilute K2Cr2O7 solution, of unknow concentration
(enough to give a light yellow color to the solution).
The electrodes were then fitted in the two holes made in a plastic lid. It was pretty delicate to have them stay in place with the crocodiles in
place, I finished by using some wooden "pliers" (don't know the name, the things you use to attach your laundry to the line outside), and a elastic to
hold evrything in place. This worked great.
I then turned the power supply on, at maximum setting at first, then turned it down until the current limitation diode went off. This offered a
current of roughly 1.9A. I couldn't go other that without setting the current limitation on. The power supply displayed 5V, and there was ~2.8V across
the cell.
H2 evolution was vigorous but contained, and Br2 formation was immediatly seen:
I decided on throwing a stir bar in it, as the bromine seemed to stay at the bottom, the H2 generation not mixing everything that much.
Electrolysis was continued for several hours, periodically adapting the voltage as current increased or diminished. The current was maintained
between 1.6 and 1.9A, and the voltage across the cell slowly increased.
Strange enough, the cell didn't heat up. Not more than 25°C anyway. Concerned that this would slow down the decomposition of hypobromous acid to
bromate, i decided on heating the cell slight on a hotplate. It was kept very warm to the touch (~50°C).
Afetr 4h, some crystals started appearing:
This slowly continued, but at one point I realized crystals were forming on the electrodes:
As i didn't want them to diminish electrode surface, I cut the current, and scrapped them off with a spatula. I realized that heating the cell with
the hotplate meant that the electrodes were the coldest par of the cell, and that naturally the broamte would cristallize there, so I stopped heating.
Although the cell cooled down pretty much, it remained warm to the touch even several hours after.
The electrolysis was continued for over 10hours, at which point the current had decreased pretty much (down to 1.4A), and there was other 3.45V across
the cell. Current was stopped, and the electrodes removed after scrapping any remaining crysatls off. Pretty surprisingly, the solution was dark
yellow, and not green as Woelen's solution was. Also, there was only a very minimal amount of carbon particlues in the solution. Once the electrodes
were placed in dH2O, it was seen only the anode was very slightly degraded, leaving a very small amount of fine carbon particules on the bottom.
(The solution looks much darker than it really is, because of the light i suppose)
The stoppered cell was placed in the fridge for 2 dasy, and the dark yellow solution decanted into a bottle for re-use.
30mL of dH2o were added to the solids, and shaken up. The off white solids quickly decanted, leaving a very slightly green/yellow solution. Only a
very slight amount of fine carbon particules are present. The solution is still in the fridge.
The two major difference with Woelen's results are that there seems to be no reduction of the chromate, and that the electrodes are in very good
condition after 10h of electrolysis; there is also hardly any heating up of the cell. Maybe this is du to a lower current, although I think woelen
passed 2A, which isn't much more. Maybe the lectrodes are different, but I doubt it considering we used the same type of batteries to salvage them. I
haven't filtered the KBrO3 yet, so can't really compare yields though.
I did this electrolysis just as a first experiment, i guess I will use the bromate to make Br2 one day, but for the moment I have,n't got much use to
it.
I intend on trying to prepare pinacol by electrolysis of acetone in a compartimented cell, with the same electrodes. I might have a very high
resisatnce though, considering that the electrodes will be pretty far apart, so I'm considering making some lead electrodes with a high surface, but
thin enough to be able to fit into the joint of this beauty:
In any case, a big thank you to Woelen for sharing his results, for his excellent write up and unvaluable help, I really appreciate.
\"You can battle with a demon, you can embrace a demon; what the hell can you do with a fucking spiritual computer?\"
-Alice Parr
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woelen
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Klute, very nice write-up and good results. Also nice to see that the main results are similar, but at details there are differences.
I think that the most important difference is in the graphite rods. I used rods from an old used and depleted battery, did you also do that, or are
your rods from a new battery? This might be an important difference. Also good to see that you had much less anode corrosion. I obtained quite some
fine crap.
The green color I obtained, could be the result of reduction of the chromium, but it could also be due to the mix of black particles and yellow color
of chromium.
The final results are quite similar though and that is the most important. Nice to see that more people are making KBrO3 and that this chemical
becomes more and more common among home experimenters.
Just for fun, KBrO3 makes awesome black powder like stuff. Mix it with some S and C. Be careful, the mix is VERY energetic and also very sensitive.
Don't grind S/KBrO3 mixes! You will be surprised to see how fast such a mix burns.
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Klute
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I used a very old battery! It had been depleted for several years, and was laying outside... but still in good shape. The fact that it's a 4R25 would
suggest it's the same electrodes that are used, no?
I have yet to find a good oxidation procedure using KBrO3, but have been busy elsewhere, and I'm sure it cna be sued for lots of other things! Thanks
to you for spreading this reagent!
BTW, do you think peroxodisulfates could be produced in a similar fashion? On the opposite of bromates, they are cheap and readibly available, but if
one constructed a rather large cell, this could be another use for it.
\"You can battle with a demon, you can embrace a demon; what the hell can you do with a fucking spiritual computer?\"
-Alice Parr
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12AX7
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Sulfate oxidizes graphite, swelling and destroying the anode.
Every time I electrolyze a bromide solution, my graphite swells and degrades substantially. Too much voltage? Current seems reasonable, and I can
run much more current (and presumably at more voltage, I haven't measured) in a chlorate cell of identical dimensions (including anode and cathode).
Tim
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woelen
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Apparently there is a lot of difference in anode quality. I am also surprised by the difference between Klute's results and mine. I really had much
more carbon crap than Klute. This is worth some investigation, but I think it has to do with the compactness/hardness of the graphite rods. I have
been reading my synth in detail again and the only main difference is that I used a lower volume (only 40 ml with 15 grams of KBr). So, in my
situation I had approximately the same heat as Klute, but in only half the volume. That at least explains why my solution becomes warmer than Klute's.
The higher temperature is good for bromate formation from hypobromite, but it might cause much faster erosion of an anode. But this only is
speculation and I cannot back it up with research results.
Making peroxosulfates from sulfates is not as easy as making bromates, because this requires fairly precisely specified conditions. If the conditions
are not right, then only oxygen and acid are produced at the anode. I understood that the temperature must be low and the current density must be very
high, so you need very thin electrodes and large currents. This requires platinum wire, graphite will not withstand high currents. Graphite should not
have a current density of more than 100 mA per square cm, even better is 50 mA per square cm.
[Edited on 23-5-08 by woelen]
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Klute
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I could always send you a couple electrodes and vice versa and see if we get the same results this time 
Too bad for the peroxodisulfates.. I can't wait to try the pinacol synth though. Do you think fishing lead is suitable for such an electrolysis? I've
actually have a hard time finding good quality lead, and don't feel like cutting a car battery open, considering I wounded myself not that long ago,
and am just starting to get used to having my left hand back, I don't want to take any risk chopping a finger off
\"You can battle with a demon, you can embrace a demon; what the hell can you do with a fucking spiritual computer?\"
-Alice Parr
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dann2
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Perhaps this has been posted before somewhere
Attachment: bromates.pdf (1.1MB)
This file has been downloaded 757 times
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Polverone
Now celebrating 21 years of madness
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Can you repost that attachment as a file in scipics or using mihd.net, or something like that? Normally I can successfully download attachments from
the board using Opera or Curl if Firefox chokes, but I can't open this one no matter how I try to download.
PGP Key and corresponding e-mail address
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MagicJigPipe
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Klute, some bullets contain high purity lead. Just check the package. They are usually available at gun stores that sell reloading supplies. Some
even sell pure lead blocks for melting and casting one's own bullets.
That's where I get my lead if and when I need it.
[Edited on 5-23-2008 by MagicJigPipe]
"There must be no barriers to freedom of inquiry ... There is no place for dogma in science. The scientist is free, and must be free to ask any
question, to doubt any assertion, to seek for any evidence, to correct any errors. ... We know that the only way to avoid error is to detect it and
that the only way to detect it is to be free to inquire. And we know that as long as men are free to ask what they must, free to say what they think,
free to think what they will, freedom can never be lost, and science can never regress." -J. Robert Oppenheimer
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Klute
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We don't have the same laws in europe concerning amunition
I can't just walk in and ask for ammo without showing some kind a license or club membership card. Especially if the guy realize I don't know
anything in amunition or weapons in general
I've read alot of people referring to "roofing" lead, is lead still used in construction materials? I thought it was bad for one's health to live in
a house with lead.
I've tested my cell with the two same graphite electrodes and a brine solution, no current passing... the multimeter says 5200 Ohms resistance
between the two.... Which gives me a few mA of current at 15V... I hope lead electrodes with a slighter bigger surface will pass some current at
least, I can't have them too big or they will not go through the 14/23 joints... Hopefully dilute H2SO4 will also be more conductive.
These kind of cells are usually used for electrocrisatllizations, at a few mA current for days to obtain a few millimols of complexes, not really
for large current electrolysis.
\"You can battle with a demon, you can embrace a demon; what the hell can you do with a fucking spiritual computer?\"
-Alice Parr
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MagicJigPipe
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IMO, metallic lead isn't extremely hazardous as long as you don't use it as toothpaste or hand lotion. It's still widely used as fishing weights
where it's main purpose is to be tossed in the water. It's still used in bullets, which some people handle on a daily basis.
I would be 0% afraid to live in a house whose wood was replaced with lead. Up until 40 years ago it was still used to transport water! It's just
that it's not very soluble and it doesn't "get around" like mercury. I mean, the only way people were harmed by lead paint is if they ate it or
continuously breathed the dust. If an adult eats paint, perhaps they deserve lead poisoning.
Anyway, I wouldn't be shocked if you could still find it as "roofing lead", however, I don't know about Europe. They seem to be slightly more
paranoid about stuff like that.
Also, I forgot (if I knew) that you were in Europe. My mistake. If it weren't so heavy I'd send you some but you know how dense things mix with
shipping.
12AX7, every time I electrolyze anything with graphite it seems to slowly (sometimes very rapidly) disintegrate the graphite. I hate it.
That's why I hardly ever do anything with electrolysis because I hate cleaning the fine carbon "dust" out of the solution. Platinum is way too
expensive for me right now.
I suppose I'm somewhat of an idealist as I think electrolysis should go smoothly without all that contamination. I guess that's just unrealistic,
though. I suppose I might be a little OCD in that respect. I dislike using grease on my joints because of the contamination it might cause.
"There must be no barriers to freedom of inquiry ... There is no place for dogma in science. The scientist is free, and must be free to ask any
question, to doubt any assertion, to seek for any evidence, to correct any errors. ... We know that the only way to avoid error is to detect it and
that the only way to detect it is to be free to inquire. And we know that as long as men are free to ask what they must, free to say what they think,
free to think what they will, freedom can never be lost, and science can never regress." -J. Robert Oppenheimer
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woelen
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Also in Europe, lead still is used in buildings. The house, which I live in is built in 2002, and it contains quite a lot of lead-slabs. This is a
nice source of lead.
We don't have gun-stores over here, not in France and not in NL  . That source
of lead does not exist over here.
I also have the feeling that graphite disintegrates too fast when doing electrolysis. I have some platinum electrodes and I really like them. I,
however, sometimes do electrolysis with graphite, especially if I want to share results with others. I know that not everyone has platinum electrodes
and then it is good to be able to make a write-up which is based on graphite. That's why I also made some bromate with the graphite rods.
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MagicJigPipe
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What kind of cruel, cosmic joke is the universe trying to play on us chemists by making one of the most useful metals also the most rare and/or
diificult to obtain and subsequently expensive? I can't wait until a decent sized asteroid full of platinum crashes into the moon (or an unpopulated
place on Earth. Even better!) making mining efforts worthwile while, at the same time, making platinum about as valuble as silver or perhaps even
copper.
It's too bad platinum seems to even be rare in the rest of the solar system (especially in asteroids).
I know this is OT but have we ever discovered any significant (or any at all) amounts of platinum in our solar system? If we have I haven't heard of
it.
Also, what's the best source of platinum wire? Would platinum-plated wire suffice in most electrolysis situations?
"There must be no barriers to freedom of inquiry ... There is no place for dogma in science. The scientist is free, and must be free to ask any
question, to doubt any assertion, to seek for any evidence, to correct any errors. ... We know that the only way to avoid error is to detect it and
that the only way to detect it is to be free to inquire. And we know that as long as men are free to ask what they must, free to say what they think,
free to think what they will, freedom can never be lost, and science can never regress." -J. Robert Oppenheimer
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12AX7
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Platinum is a heavy metal, so it stands to reason much of it is dissolved in the earth's metallic core. Probably the same is true of metallic
asteroids (you may have heard the K-T boundary layer -- from the blast that killed the dinosaurs -- is "rich" in iridium, although I don't know how
much). Uranium likewise is probably deep in the earth, either dissolved in metallic phase or in heavy rocks deep in the mantle.
Zone refining, with solar ovens, has been proposed to purify asteroids in space. The more valuable elements (like platinum) would be economical to
send to Earth.
Tim
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not_important
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| Quote: | Originally posted by 12AX7
Platinum is a heavy metal, so it stands to reason much of it is dissolved in the earth's metallic core. Probably the same is true of metallic
asteroids (you may have heard the K-T boundary layer -- from the blast that killed the dinosaurs -- is "rich" in iridium, although I don't know how
much). Uranium likewise is probably deep in the earth, either dissolved in metallic phase or in heavy rocks deep in the mantle.
...
Tim |
Platinum metals, iron group, manganese and gold, but not uranium. The elements likely to be in the core are siderophiles, plus some sulfur, carbon,
and hydrogen; all those trapped by their solutions in the siderophiles.
Uranium, thorium, the lanthanides, the alkali and alkaline earthe metals are lithophiles; ending up in stony minerals. Along with the halogens and
oxygen, they concentrate in the crust. The radioactive elements likely have an extra push towards the crust due to their intrinsic heat generation.
The rest of the metals and semi-metals are chalcophiles, named from their ready combining with sulfur; they concentrate in the outer core and mantle.
There's a bit of overlap, the iron group shows up with chalcophiles, where oxygen is present iron also ends up with the lithophiles. Some elements
have unexpected distributions, usually attributed to their formation through decay of moderately long lived isotopes of elements concentrating in one
of the other families; tungsten formed from (192)Hf - a lithophile.
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JohnWW
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This is getting rather off-topic for this thread, but I remember reading somewhere that there is more gold (and probably platinum group metals and
silver) under the sea (where the crust consists of basalt and in places ultra-basic rocks like peridotite) than has ever been mined on land. So if or
when mining deep ocean floors, or even the Eath's mantle or core, ever becomes economic and practical, there should be price drops for those metals.
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dann2
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Hello,
The file I posted is corrupted. Not the boards fault.
ELECTROLYTIC PRODUCTION OF BROMATES
JOURNAL OF THE ELECTROCHEMICAL SOCIETY
July 1957 Vol. 104, No. 7
Attachment: bromates.pdf (383kB)
This file has been downloaded 816 times
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