Sciencemadness Discussion Board

bromate synthesis

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IodineForLunch - 2-10-2002 at 13:13

Madscientist and I were talking, and the discussion turned to bromates. We came up with various possible ways of preparing bromates, but we're not sure which ones might work, and which ones won't. Here is what we've thought of:

1. React elemental bromine with nitric acid to yeild bromic acid, water, and ntirogen oxides. Then add a hydroxide to precipitate the corresponding bromate According to madscientist, the first step would be endothermic, but I don't know what form of energy input it would require (that is, if it works at all).

2. Electrolysis of aqueous potassium bromide. I think this is how it might be done in industry. In room-temp water (~25 degrees) KBr is 5 times as soluble as KBrO3.

3. Oxidation of metallic bromides or hydrobromic acid by Mn2O7.

4. Oxidation of hydrobromic acid by hydrogen peroxide, then adding hydroxide to precipitate the corresponding bromate.

Any thoughts?

David Hansen

madscientist - 2-10-2002 at 13:22

The only process I have any real confidence in is electrolysis of an aqueous solution of a bromide.

Oxidation of bromine with nitric acid sounds unlikely. HBr is oxidized by hydrogen peroxide, yielding elemental bromine and water (of course, a small amount of the bromine reacts with the water). Any non-electrolytic method of oxidizing bromine is going to be difficult to perform.

Still interesting to think about... I'd like to know more about non-electrolytic methods. I do think that a solution of HOBr (prepared by dissolving elemental bromine in water or hydrogen peroxide) could yield HBrO3 when heated.

IodineForLunch - 2-10-2002 at 15:21

How bout this three-step reaction?

1. React elemental bromine with hydrogen peroxide to form hypobromous acid.

3Br2 + 3H2O2 --> 6HOBr

2. Heat HOBr solution to initiate decomposition of hypobromous acid into hydrobromic acid and bromic acid.

6HOBr --heat--> 4HBr + 2HBrO3

3. Add potassium hydroxide in order to precipitate potassium bromate.

2HBrO3 + 2KOH --> 2KBrO3 + 2H2O

Net reaction:
3Br2 + 3H2O2 + 2KOH --> 2KBrO3 + 4HBr + 2H2O

Increasing the amount of H2O2 used may lead to a higher bromate yield:

3Br2 + 6H2O2 + 3KOH --> 3KBrO3 + 3HBr + 6H2O

The only problem I can think of is formation of undesired potassium bromide upon addition of the hydroxide - however, potassium bromide is 4-5 times as soluble as potassium bromate, and will likely stay in solution while the potassium bromate precipitates out.

I'd like to hear more possible methods.

David Hansen

IodineForLunch - 2-10-2002 at 15:23

However, I'm sure that PHILOU Zrealone will come along and prove us wrong! :P

J/K, no harm intended.

David Hansen

vulture - 3-10-2002 at 07:53

IIRC, the order of halogens is reversed when it comes to oxygen compounds. To be short, this comes down that bromine will replace chlorine in chlorate.

IodineForLunch - 10-10-2002 at 14:19

Or...

K2CO3 + 5H2O2 + Br2 ----> 2KBrO3 + 5H2O + CO2

I found that in Hawley's condensed chemical dictionary.

David Hansen

BromicAcid - 3-12-2003 at 09:52

You can react a hot potassium hydroxide solution with bromine to produce potassium bromate, potassium bromide, and water.

Br2 + KOH + H2O ------> KBrO3 + H2O + KBr

The KBr and KBrO3 can be seperated from each other by fractional crystalization.

blip - 3-12-2003 at 21:54

Why do you not reduce your coefficients down to the smallest possible, IodineForLunch? I must be missing something. For example:
Quote:
3Br2 + 3H2O2 --> 6HOBr

can be reduced down to Br<sub>2</sub> + H<sub>2</sub>O<sub>2</sub> <s>&nbsp;&nbsp;></s> 2HOBr

KABOOOM(pyrojustforfun) - 5-12-2003 at 20:47

<blockquote>quote:<hr>Br2 + KOH + H2O ------> KBrO3 + H2O + KBr <hr></blockquote>that's unbalanced.
6KOH + 3Br<sub>2</sub> <s>&nbsp;&nbsp;&nbsp;></s> KBrO<sub>3</sub> + 5KBr + 3H<sub>2</sub>O

woelen - 12-10-2006 at 01:17

*BUMP*

I did the synthesis of KBrO3 by means of electrolysis. There is a nice synth on www.versuchschemie.de by garage chemist:
http://www.versuchschemie.de/htopic,6656,kaliumbromat.html.
He used platinum wire as anode. I had to use a thick graphite rod as anode. I also used such a graphite rod as cathode.

I dissolved as much as possible of KBr in 35 ml of water.
I also added 20 mg of K2Cr2O7 and dissolved this.
The solution now is pale yellow.

I dipped the electrodes in the liquid and used a current of 4 A for the electrolysis (using a 13.8 V power source, with suitable series resistor in order to regulate current and obtain a fairly constant current, independent of precise chemical composition of the electrolyte).

The evolution of hydrogen worked very well, and it also remained good for all time I did the electrolysis. The effect, described by garage chemist, that bromate is reduced back to bromide, did not occur. I did not add a drop of soap in order to suppress spraying. For that purpose, I used a plastic cap, with tight holes in it for the electrodes. Any spraying only wetted the bottom side of the cap and the liquid runs back into the beaker.

The liquid became very hot (probably well over 70 C, I really could not stand touching the long/tall glass in which the solution was without gloves).

I stopped electrolysis after 3 hours. After that period the liquid has turned deep green and there was some black soot on the bottom from the graphite anode, but this was only a very little bit. There hardly is any erosion of the anode. That's good.

I decanted the hot clear liquid from the black soot and hence got rid of the black stuff, and what remains is a hot clear green liquid with a brown/yellow tinge. On slow cooling down, fairly large crystals of KBrO3 settled. I was nicely surprised by the good yield.

Now comes the problem: My crystals are not nice white, but they are light green with a yellow/brown tinge. I think they have chromium (III) in them, together with remains of dichromate. I carefully rinsed the crystals with cold distilled water, but that does not help. The KBrO3 is a potent oxidizer, so that is OK, but I do not like its color, it should be white.

I could of course recrystallize, but that would introduce quite a large loss of KBrO3. Why do I get green/brown crystals and garage chemist obtains snowwhite crystals, also directly from his first batch, without the need to recrystallize? He even used more K2Cr2O7 than I did, relative to the amount of KBr he used. Also, his K2Cr2O7 did not become green and apparently is not reduced to chromium (III). My liquid was dark green at the end, his still is yellow.

[Edited on 12-10-06 by woelen]

woelen - 12-10-2006 at 13:09

I did recrystallize the material and took the loss for granted. The bromate now is almost white. Only a very vague green color remains, which is only noticeable, if the KBrO3 is put next to a purely white solid.

But still, I do not understand how the pure material can be achieved at once. I did another experiment, now with 1 A current, but the result is the same. The solid I obtain is off-white and yellow or green, depending on how long I perform the electrolysis. I simply cannot prevent contamination of the crystals with chromium :(.

I also did a test for bromide content in my sample. With respect to bromide it is very clean. If I add a spatula of this to 2 M H2SO4, then no yellow color at all can be observed, and no smell of bromine can be observed. As soon as I add a spatula of KBr as well, the liquid becomes turbid and rusty brown, due to very finely dispersed droplets of Br2. After a few minutes I have a big drop of Br2 at the bottom.

So, I regard my making of KBrO3 quite succesful, the only flaw remaining some traces of Cr (III) and possibly Cr(VI).

[Edited on 13-10-06 by woelen]

not_important - 12-10-2006 at 22:53

Interesting. Are you willing to sacrifice a few grams of bromate to try crystallizing with controlled amounts of Cr(VI) or Cr(III) to see if it is only the +3 state that is causing the problem? It may be that the carbon electrodes cause too much reduction of the hexavalent chromium.

woelen - 13-10-2006 at 01:56

I have quite some KBrO3 now, I now can make this routinely if I wish, approximately 10 grams per evening. I do not want electrolysis cells running overnight, hence the slow progress, but it does not bother me. I only need gram-quantities.

I have an analytic balance with 0.1 mg resolution, so I can do nice precise experiments.

I can do an experiment with 500 mg of KBrO3 in 3 ml of water and 5 mg of K2Cr2O7 added, and see how yellow my crystals of KBrO3 are after recrystallization.

With Cr(III) things become less satisfactory. I do have CrCl3.6H2O and KCr(SO4)2.12H2O, but none of these provide a good model for the situation in the electrolysis cell. The chloride and sulfate definitely will interfere, they both strongly coordinate to chromium (III) ions, when the solution is heated (try it, dissolve some chrome alum in water, and then heat, you'll see a color change from violet to deep green).

I'll do the experiment with KBrO3/K2Cr2O7 this weekend. If the resulting crystals are still nice white, then we know that the Cr(III) is the problem.

--------------------------------------------------------------------------

Yet another thing, could it be possible to make KBrO4 from the KBrO3? I would love to experiment with that compound. I read something about the need of fluorine, but aren't there any other methods, which are more easily accessible for the home chemist? What would be the solubility of KBrO4, would it be easy to isolate? If there is analogy with KClO4 then I expect it would be easy.

[Edited on 13-10-06 by woelen]

Purification of Potassium Bromate

Chemist514 - 13-10-2006 at 06:34

Hello!

Not sure if this is in anyway usefull to you peeps, but its from The Purification of Laboratory Chemicals 5th edition...

Potassium bromate [7758-01-21 M 167.0, m 350°(dec at 370°), d 3.27. Crystd from distilled
H20(2mL/g) between 100° and Oo. To remove bromide contamination, a 5% soln in distilled H20, cooled to
loo, has been bubbled with gaseous chlorine for 2h, then filtered and extracted with reagent grade CCl4 until
colourless and odourless. After evaporating the aqueous phase to about half its volume, it was cooled again
slowly to about loo. The crystalline KBrO3 was separated, washed with 95% EtOH and vacuum dried [Boyd,
Cobble and Wexler J Am Chem Soc 74 237 19521. Another way to remove Br- ions was by stirring several
times in MeOH and then dried at 150° [Field and Boyd J Phys Chem 89 3767 19851.

Chemist514 - 13-10-2006 at 06:38

Opps, lemme correct a couple things that got messed up...

I should read "between 100 and 0 degrees"
and "then cooled slowly to 10 degrees"

not_important - 13-10-2006 at 08:29

Quote:
Originally posted by woelen
(snip)

Yet another thing, could it be possible to make KBrO4 from the KBrO3? I would love to experiment with that compound. I read something about the need of fluorine, but aren't there any other methods, which are more easily accessible for the home chemist? What would be the solubility of KBrO4, would it be easy to isolate? If there is analogy with KClO4 then I expect it would be easy.

[Edited on 13-10-06 by woelen]


I've never read of any other method than using F2, XeF2, or possibly higher xenon fluorides. JACS 90, 1900 (1968), Inorg Chem 8, 223 (1969), and Inorg Synthesis 13,1 (1972) are the references that I have, along with a scribbled note about some patents using F2.


This might have more, it's pay-to-see for me
http://www.turpion.org/php/paper.phtml?journal_id=rc&pap...

You could try for BrO3F, which attacks almost everything including teflon. 8-)

woelen - 13-10-2006 at 13:39

Quote:
You could try for BrO3F, which attacks almost everything including teflon. 8-)


I'll first stick to the KBrO3/K2Cr2O7 recrystallisation experiment and post an update if I have results (after this weekend).

Next, I'll have a look at the attic and see if I still have a few pounds of XeF4 or BrO3F lying around. I'll post an update on this as well, but that could take a _slightly_ longer time than next weekend :D.

Al Koholic - 13-10-2006 at 18:17

You know woelen...perhaps the soap is having an effect on the precipitation of the bromate crystals. You should add a drop to your next batch to see what happens.

garage chemist - 14-10-2006 at 05:12

I have once tried to make perbromate by electrolysis of bromate, because I read in Brauer that perbromates could be prepared in this way.

However, electrolysis of a pure NaBrO3 (NaBrO3 was from a chemical supplier) solution with platinum anode did only yield oxygen at the anode and nothing else.

Recently I had a look into a book on inorganic chemistry (Earnshaw, Chemistry of the elements) and read that perbromate was prepared in 1% yield by electrolysis of lithium bromate solution. So this is not a viable way to get perbromates.

The only way to obtain useable quantities of perbromate is to oxidise a strongly alkaline bromate solution with fluorine, as done in Brauer. This is of course much cheaper than using xenon compounds, but also less convenient.

woelen - 15-10-2006 at 06:06

I did the experiment of recrystallizing KBrO3 with 1% K2CrO7. I had 660 mg KBrO3 and 6.4 mg K2Cr2O7, dissolved in appr. 3 ml hot H2O. This solution has quite a strong yellow color. On cooling down, almost all KBrO3 separates from this solution. When it is rinsed, then it is totally white!! This really is amazing.

I now had much more dichromate, relative to the amount of KBrO3, then in the electrolysis experiment. So, the chromium (III) must be the problem. The color also confirms this, the solid, obtained from the electrolysis cell is light green.

not_important - 15-10-2006 at 07:23

So it would seem to be that the graphite electrodes reduce enough Cr(VI) to Cr(III) to give you problems. Going to be a problem removing it, except possibly with ion exchange or co-precipitation.

woelen - 15-10-2006 at 11:44

I used a graphite cathode as well. Garage chemist used a copper wire. I used a graphite cathode, because I did not want any extra metal contamination besides the chromium. But probably this was not the right decision.

garage chemist - 15-10-2006 at 12:06

I think that the graphite anode is the cause of your contamination. It gives off colloidal partivcles which can adsorb metal ions, like chromium(III) and be impossible to remove by filtration.

A copper or iron cathode does not introduce any extra metal ions.
The metal is completely protected against corrosion by the negative charge.

I'm sure that using a platinum anode and copper cathode will solve your problem.

woelen - 15-10-2006 at 12:55

Quote:
I'm sure that using a platinum anode and copper cathode will solve your problem.

Yes, I also am sure about that. I ordered some platinum wire (5 inch of 0.4 mm Pt/Ir alloy, 95/5, and 10 inch of 0.3 mm of 99.9% platinum) and when it arrives, I'll try again. But still, the graphite anode method also is quite useful. The batch of KBrO3 I now have is VERY pale green after a single recrystallization. Without a pure white reference (e.g. white table salt) besides it, you don't notice the color. For the great majority of experiments, this KBrO3 is perfectly suitable. It has no dichromate in it, it only is chromium (III). I did the very sensitive acid/H2O2 test on dichromate, and this test was negative.
I like the graphite method better, because it is accessible for more people. If I share such a method, then I like it if other people also can make such an interesting compound without too much problems.

I'll forget about the making of KBrO4. The F2-method of course is out of reach of home chemists and other homelab methods indeed seem not to exist.


[Edited on 15-10-06 by woelen]

Zinc - 24-10-2006 at 05:25

Quote:
Originally posted by woelen
I dissolved as much as possible of KBr in 35 ml of water.
I also added 20 mg of K2Cr2O7 and dissolved this.
The solution now is pale yellow.


Is K2Cr2O7 required or can the electrolysis be performed with out it?

woelen - 24-10-2006 at 07:32

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.

Zinc - 30-10-2006 at 12:43

Be careful when working with bromates because they are carcinogenic.
http://en.wikipedia.org/wiki/Bromate

(Yay! 100 posts!)

woelen - 19-11-2006 at 08:05

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]

woelen - 23-11-2006 at 02:48

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]

pyrochem - 29-11-2006 at 22:10

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]

not_important - 29-11-2006 at 23:08

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.

YT2095 - 12-5-2007 at 02:13

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]

woelen - 12-5-2007 at 13:11

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.

JohnWW - 12-5-2007 at 14:07

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.

dann2 - 12-5-2007 at 17:02

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

YT2095 - 13-5-2007 at 00:17

Quote:
Originally posted by JohnWW
how about trying to prepare perbromates?


perbromates are not possible to make without the use of radioactive isotopes.

Nerro - 13-5-2007 at 01:19

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...

Zinc - 13-5-2007 at 01:55

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?

Silver wire

dann2 - 13-5-2007 at 05:51

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

woelen - 13-5-2007 at 10:23

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.

Antwain - 3-11-2007 at 18:50

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]

disolving platinum.JPG - 36kB

Antwain - 3-11-2007 at 18:51

:mad:

disolving platinum1.JPG - 38kB

JohnWW - 4-11-2007 at 00:23

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.)

garage chemist - 4-11-2007 at 00:30

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!

Antwain - 4-11-2007 at 00:53

@ 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?

garage chemist - 4-11-2007 at 02:32

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...

Antwain - 4-11-2007 at 03:04

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.

woelen - 4-11-2007 at 03:35

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.

Antwain - 4-11-2007 at 04:19

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.

Antwain - 4-11-2007 at 07:22

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.

woelen - 4-11-2007 at 08:37

Yes, electrolysis is a slow process. You could have known if you had looked at the quantities I showed on my webpage ;).

I doubt that you can convert the 120 g of KBr quantitatively into KBrO3 in one run. I would first let it run for one day or so, and collect the KBrO3. Also have a frequent look at your cell. The anode tends to be clogged by crystals of KBrO3 somewhat later in the process and this may completely bring the process to a standstill. Every now and then, you may need to cleanup the anode, scraping off crystals of KBrO3.

If you have been running your cell for some time, then let the liquid cool down in the fridge, such that you have maximum yield of KBrO3. The liquid simply can be decanted (don't throw it away, it can be used with more KBr and water to make a new cell) and the solid KBrO3 must be rinsed with ICE-COLD water, and then recrystallized. You'll see that it can be made amazingly pure with only a single recrystallization. Dissolve the crystals in as little as possible of hot - near boiling - water and then let cool down agian. Finally put it in the fridge. This gives a yield of 90% or so.
The remaining 10% in the liquid can be retrieved as well, with impurity, but this is useful for the more raw pyrotechnic experiments, or for making bromine.

12AX7 - 4-11-2007 at 08:38

Electrolysis in quantity easily runs days. Is that a problem?

I have some H2PtCl6. The solid powder is orange, while the solution ranges from dilute yellowish green (much like dilute alkaline chromate) to yellow.

Tim

Fleaker - 4-11-2007 at 10:23

The solution is very orange red, dilute it is yellow.

Palladium solution in nitric acid or aqua regia actually a very very potent yellow colour that looks brown. I will show you all what I mean by this shortly.

Antwain - 4-11-2007 at 10:46

Thanks Tim. No, its not a problem.... but I reserve the right to consider it a pain in the bum :P I want it now!!! :D

I was concerned initially about the dissolution of the anode as I said before but it appears to be sitting pretty 3.5hrs into the process. There is brown scum on the surface but I think that is a titanium compound. The outside of some of the the Ti wire was inadvertently oxidised during the glassing process. It went all multicoloured like many metal when they are heated in air and some of it may have oxidised more than that. It didn't matter because I ran 8 wires into the solution with a greater surface area by a factor of at least 2 than the anode. 10ft of the wire, including postage from England, cost far less than a tenth of the platinum so I didn't care about wasting it. Anyhow, its clean now and there is black scum so Im guessing that the Ti wire is to blame.

@wolen- Firstly, I read your page but was more concerned with procedure than numbers. And I knew that it was bad, I had the number 10000 in my head for columbs per mole (hey, its the right order of magnitude). What I didn't do mentally was multiply it by 6 electrons and realise that dividing by 3600 was still a lot.

No need to worry, I'm pretty good when it comes to recrys. Ive had way too much practice. I'm also obsessive compulsive about purity, so it will have at least 2 recrystallisations, probably one each for each batch isolated and another for the whole lot all added together.

Also- and in one sense this speaks highly of my cell- it is stone cold. I added a bunch of ganged small resistors to push the current up to 3.6A but it is not going to make that much of a difference. I wish it was running hot because then I could isolate a fair bit of bromate by cooling. It is sitting on a hotplate so if necessary I will add a water bath when crystals start to appear.

Is there any disadvantage to running a cell like this hot? Will it significantly increase the wear on the anode? And IIRC, it will increase the current, so I will have to be careful of runaway.... Is this right?

[Edited on 5-11-2007 by Antwain]

Antwain - 4-11-2007 at 10:49

Quote:
Originally posted by Fleaker
The solution is very orange red, dilute it is yellow.

Palladium solution in nitric acid or aqua regia actually a very very potent yellow colour that looks brown. I will show you all what I mean by this shortly.


Yes, I have made it. Well its brown by definition because it looks brown. And I assure you it stains stuff brown. But yes, when diluted it becomes yellow and remains coloured to a very high dilution.

chloric1 - 4-11-2007 at 11:11

Quote:
Originally posted by Antwain

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.


Absolutely right! When I burn my chlorate impregnated filter papers after thorough drying, the energy release is terrific. Twenty to 30 grams of oxidizer being reduced in less than a second. The heat is intense as I have seen NaCl vapor condense on room temperature surfaces!:o:o



The first virgin run of my 2 liter sodium chlorate cell assuming 50% was 14 days! I only have 6 amps to play with at this writting. A 12 amp setup is in testing. I boiled the solution to half volume and reran another 7 days to get 600 or 700 grams of sodium chlorate. With 10 or 12 amps I could complete the first run in a week and secondary run in 4 days. But still this takes a while.

[Edited on 11/4/2007 by chloric1]

Antwain - 4-11-2007 at 11:33

At the moment I am runing the shameful battery charger, rated for 6 amps. But I just grabbed a bare transformer my father had custom wound (why I do not know :o) in the 70s. It is supposedly good for 2* 20A @8v but probably not parallel bridgeable. I will need to get a case for it, a fuse and power cord, but this seems like it is up to it is up to any challenge I could reasonably set for it. And the diode bridges attached to it :o. If god needed diode bridges, these are the ones he would use ;), they are huge. Now I just need to build a variable resistor that can dissipate a gazillion watts.

12AX7 - 4-11-2007 at 20:34

Bah, you haven't seen anything 'til you get your hands on hockey puck rectifiers. :D Four kiloamperes, anyone?...

Conductivity of the solution will probably about double by the 60C range. My cell runs about 30A when the solution is changed, 60A when it's reached equilibrium around 60C or whatever it sits at. Which is less these days, as the anode is down to about 3/4" dia..

Burning NaClO3 is indeed quite impressive. It contains much more oxygen than the potassium salt, and let's not forget the wildly luminous sodium d-lines giving that intense yellow firey look to it. I've got a cardboard drying tray that's encrusted with NaClO3. One of these days, I'm going to take it into the street and light it. With a piece of fuse, mind you. I have two KClO3 trays, but I know they won't be nearly as bright or fast.

Tim

Eclectic - 5-11-2007 at 10:29

Occasionally a nice 200A, 12V adjustable Lambda power supply will turn up on Ebay. :D


Good for electroplating, electroforming, and electrowinning also. (answer for Woelen V)

[Edited on 11-5-2007 by Eclectic]

woelen - 5-11-2007 at 10:34

Why all those high currents? If you don't have high currents, then use lower currents in cells, connected in series. You easily can take 24 volts and add 4 cells to it. In many cases that is more efficient.

I, however, just am patient. I only need small amounts, and I then can live perfectly with 2 A.

Antwain - 5-11-2007 at 15:27

After crystals came out I heated it a bit and then 7 hrs later more crystals appeared. So I took a first crop and its a bit ugly, definitely a candidate for 2 recrystallisations. Next time I think I will just let it run, because although the anode was covered in crystals the current had not dropped by more than 0.1A if that. I added some more dichmorate because I don't know if there was enough there. Interestingly bromine (the bad smell not visible amounts) was liberated but it shouldn't have been because the solution was alkaline... or is this a case of dichromate (bichromate) acting as an acid? It was about another 1/2 a gram added to only 20mL of solution so that may have changed the pH.

Something dodgy is going on too... for some reason the electrolysis seems to increase the solubility of (presumably) the bromate. When I stopped the current and heated the solution quite hot the net result was MORE precipitation, although it was actually a lot of precipitation followed by some dissolution on heating. Does anyone have an explanation for that? I have stopped mechanical stirring and only dredge it up sometimes so it is probably heterogeneous in the jar.

Also, based on what happened after my latest dichromate addition, I think the brown/black muck is a chromium compound, probably a hydrated oxide. There is too much for it to be dissolution of either electrode. How does dichromate work if it doesn't stay in solution?

Also, will the electrodes suffer damage if the bromide concentration becomes too low?(I believe so but just checking). And if so then how will I know when this happens. Bubbles at the anode will be hard to see. Will the current suddenly change?

Incidentally, the current hasn't changed by more than ~5% across a wide temperature range. In fact less, it is staying between 4.5 and 4.8A. Always.

chloric1 - 5-11-2007 at 15:30

@tim

That sounds like a lot of fun. For goodness sakes give a good distance. I have not seen many oxidizers that burn with such vigor. The low end roaring sound sounds like detonation is imminent.

Antwain - 5-11-2007 at 15:48

Actually, I may have found a good solution to the crystals problem. The electrodes are now together at one side of the top of the jar and a very slow convection is taking place because of the hydrogen evolution. As such the bottom of the jar is cooler (feels about 10*C) so crystals are now growing on the bottom. Before, the bottom half was in heated water so the top was cooler.

Antwain - 5-11-2007 at 16:05

I know that this thread is 'bromate synthesis', but considering some of the off topic stuff so far this is more closely related.

How powerful an oxidiser is bromate? I have seen a reference somewhere comparing it to permanganate. The standard reduction potentials never tell the story, because concentrations and especially pH differences change the character of redox reaction quite a lot.

In acid solution, is bromate or chlorate the more powerful oxidiser?

12AX7 - 5-11-2007 at 18:46

Quote:
Originally posted by Antwain
After crystals came out I heated it a bit and then 7 hrs later more crystals appeared. So I took a first crop and its a bit ugly, definitely a candidate for 2 recrystallisations.


What's ugly about it, particulates or just green? Particulates are hot filtered, green can be neutralized (H2O2, if it doesn't react with bromate) or lost to the recrystallizing wash.
Quote:
Next time I think I will just let it run, because although the anode was covered in crystals the current had not dropped by more than 0.1A if that. I added some more dichmorate because I don't know if there was enough there. Interestingly bromine (the bad smell not visible amounts) was liberated but it shouldn't have been because the solution was alkaline... or is this a case of dichromate (bichromate) acting as an acid?


Regardless of pH, my cell continuously gives off chlorine fumes. Even though the solution may become alkaline as a result, it only takes a few ppm (if even that) to smell. There is plenty of hypohalide present to form a small equilibrium concentration of dissolved Cl2 regardless of pH, and plenty of gas bubbling through (i.e., the hydrogen) to sparge it out.

After all, bleach smells noticably "chlorine-ey", despite its pH of 10 or so.

I wouldn't worry about pH, it will always force itself back up as a result of halogen release. After crystallizing, I add acid to the liquor until the chromate changes color (pH ~ 5), or until too much chlorine gas (pH 7-8) is being produced from the hypochlorite that hasn't decomposed.
Quote:
When I stopped the current and heated the solution quite hot the net result was MORE precipitation, although it was actually a lot of precipitation followed by some dissolution on heating. Does anyone have an explanation for that?


Hypobromite decomposing from heat?
Quote:
Also, based on what happened after my latest dichromate addition, I think the brown/black muck is a chromium compound, probably a hydrated oxide. There is too much for it to be dissolution of either electrode. How does dichromate work if it doesn't stay in solution?


Is the solution, after filtering, green? If not, then it's probably some chromium compound. I don't know what it would be to be black.

When electrolyzing a sodium chlorate solution with platinized titanium anode and copper cathode, I noticed a large amount of brown to black material forming, and often being attracted to the cathode, either as a result of deposition or electrostatic attraction. (Interesting thing that attraction, I've had particulates in certain solutions which were slowly attracted to metals placed in the solution. Electrostatic collection without all the kilovolts!) It turns out the copper cathode is actually attacked by reactive species in solution, regardless of its negative potential. Since the solution is so reactive, copper sponge doesn't form on the cathode, only more black crud, making it appear to agglomerate.
Quote:
Incidentally, the current hasn't changed by more than ~5% across a wide temperature range. In fact less, it is staying between 4.5 and 4.8A. Always.


What voltage?

Tim

Antwain - 5-11-2007 at 20:28

@Tim- 3.08v across the electrodes plus 4m of 15A rated wire, I can't measure the voltage across the cell directly.

Yes, it may just have been cold enough before heating for hypobromite to exist.

There is very negligible bromine coming off now, but it stank really well when I heated the dichromate solution.

I took a picture of the gunk but my camera wouldn't focus. It is blackish and solid. If I didn't know better I would suspect it was carbon.

[Edited on 6-11-2007 by Antwain]

woelen - 6-11-2007 at 00:03

Antwain, do not add so much dichromate. You really only need to add a pinch to your solution. In my experiment, I only used 50 mg or so for 40 ml of liquid. Dichromate is reduced at the cathode, giving green chromium(III) species, and these give a green color to your product. If you have so much dichromate, then you get a lot of fine particles, which may be very hard to remove (too small for filtering, and they act as nucleation sites for crystal formation).

KBrO3 in aqueous solution is a potent oxidizer, not because of its really high redox potential, but because it acts fast. Just for fun, add a little pinch of solid KBrO3 (may be wet or impure, does not matter) to concentrated HCl. A vigorous reaction starts, in which a mix of Cl2 and Br2 is produced. Unfortunately, bromate is not a clean oxidizer. It can be reduced to bromine, or to bromide, depending on whether it is present in excess amount or not. It only works at low pH. At high pH, bromate is not capable of oxidizing much. I would say, that compared with chlorate in aqueous solution, bromate is more powerful. Not really stronger, but acting faster.

In organic reactions, bromate is not the oxidizer of choice, because of bromine formation. An important unwanted side reaction may be bromination of compounds, and not only oxidation. E.g. toluene, suspended in a H2SO4/KBrO3/water mix will be oxidized to benzoic acid, but one get a funny sweet smell as well, most likely this is some brominated compound of benzene or toluene. Not what you want. With KMnO4 you only get benzoic acid.

Bromate also is a very powerful (and dangerous!) oxidizer in dry mixes, compositions based on bromate are more sensitive than comparable compositions, based on chlorate. So, for real production work it is not suitable, simply too dangerous and too high a risk of unwanted ignition. On Usenet:rec.pyrotechnics, it was compared to chlorate as follows: "Playing with chlorates is like playing with a deadly poisonous snake. Playing with bromates is like playing with a deadly poisonous snake, which is slightly pissed off.".

I use the bromate mainly for making bromine. It is a fantastic source for making bromine. Simply mixing it with a 5-fold molar amount of KBr or NaBr and adding this to 20% sulphuric acid makes the bromine drop out and simply pipetting is suitable for obtaining already 75% yield. For the remaining 25% you'll need a distillation setup. I do not store bromine, I make it when I want to experiment with it.

If you want to protect your anode, then simply stop electrolysis, when 75% to 80% of the calculated time for full conversion has passed. In this way, you certainly will not perform electrolysis in a cell, which hardly has any bromide. The liquor then can be kept for future cells. No need to throw it away, it will be a nice starting point for a new cell. I indeed would be careful with electrolysis when hardly any bromide is left. There will be much more oxidative strain on the anode in that case, because oxygen must be formed, instead of bromine.

Finally, I would not add any acids or whatever stuff to your cell for purification. Just collect the solid matter and recrystallize. Acid also adds contamination.

Antwain - 6-11-2007 at 03:40

Yes, the dichromate was a bad idea. Well, you learn these things. My solution is a mess :(, well it can be cleaned.

So as I mentioned in another topic solution got up my anode and started to corrode it. I sealed it with wax but this is temporary at best. I have put the connecting wire through hell it is way TOO flexible for my liking. The wax has come off a bit and has clearly reacted, a bit, but I can deal with that; there isn't much and it can be gotten out.

Now I find out that my cathode is leaking like a sieve too. At least the solution in that isn't causing damage in there, just putting a few bubbles up the glass tubing. I cant see why.... well I can, it was because in each case I sealed multiple strands of wire into the glass. But for the cathode I melted it so long and hard that the copper oxidised a lot and the oxide dissolved in the glass. You can't get more forceful than that and it still wasn't enough to actually seal it.

I think that after this I am about to give up on sealing stuff in glass, it just isn't happening for me. At least the platinum seems to be doing ok.

Antwain - 6-11-2007 at 07:44

Oh well, may as well give the blow by blow. Even if noone wants to see it now it may help someone in the future if I am thorough.

After about 12-18hrs, maybe 6hrs room temp and 12hrs at ~30-40*C, the solution started depositing crystals again. They are a pain to get out so I decided to freeze it and then return it for more electrolysis. The first 6 hours when nothing deposited was because I did a dirty thing and added the bromide/bromate I made as described in the bromine thread a while back, maybe 30-45g in the 5:1 ratio, then got out much of the bromate because common ion stopped it from dissolving (if it has KBr in it I will get that on the recrys). The solution volume also increased a bit as stuff was washed into the jar.

Anyhow, when crystals started appearing at room temp I gave it the heat with waterbath treatment. And just now it started to crystallise warm so I poured it out and iced it. I don't know the density of my product as there is still liquid in it, but I would guess 15-20g came out. Pretty poor by current, but at least its working. The following picture shows the 2 crops I have obtained thus far. You can clearly see the new and old stuff because the new stuff is still very yellow, but last time the bromate settled out and left yellow solution when it warmed up so hopefully it will again. Also there is a large amount of brown/black gunk now thanks to the dichromate, but it does settle, so worst case I will be doing a decant and have a lot of water to boil down.

electrolysis - bromate-a.JPG - 22kB

chloric1 - 6-11-2007 at 18:43

Quote:
Originally posted by woelen
I would say, that compared with chlorate in aqueous solution, bromate is more powerful. Not really stronger, but acting faster.


@woelin-the percieved activity of the bromate over the chlorate is probably due the decreased stability of bromine oxides. Most chlorate oxidations seem to go like:chloric acid,chlorine dioxide,chlorine/chloride.

Antwain - 10-11-2007 at 03:11

Ok several updates. Much of this will seem obvious to the experienced electrochemist, but I am pulling myself up by the bootlaces here and trying to detail everything.

Don't use a steel lid on your jar, it will corrode and run into the solution. It can be separated with only a loss of time, but I am sure it was hurting current efficiency. I am using a plastic lidded jar now and its going good.

Don't place the cathode under the anode. I don't know if Pt adsorbed hydrogen was screwing things up, but its going better now. There will still be lots of bubbles to react with Pt catalyst but its not being blasted with them.. I only did this, btw, to keep the electrodes close in my original bad-geometry setup.

Next to each other is good and off the bottom works well too. The bromine sinks and there is a distinct layer stopping a mm below the bottom of the cathode. This is also good because it keeps the electrodes slightly warmed by their own power and precipitation happens at the bottom.

I estimate previous current effiency to have been <10% and maybe <5%. Just from the precipitate now seen after several hours, it is clearly much higher. 5-10 fold higher.

Antwain - 10-11-2007 at 13:10

2 questions. Firstly, there are some bubbles forming on the anode. Much less than the cathode, by like 20 fold or more. There was still orange dense fluid pouring off it at a good rate. Does this mean that my current density is too high or that I am out of bromide, or could it be either?

This question is kind of in unison with the last. I have a very bad feeling about my cell. Does KBrO3 dissolve with a very noticeable cooling of the solution, like KBr? If not, then I am very concerned that I have been isolating KBr from my cell :(. Sometimes it seems like there are two different solids there sometimes just one. It should not be possible to be isolating KBr, based on the current volume of solution and the quantity of KBr used initially, regardless of whether the stuff I have pulled out so far is KBr or KBrO3. But for something which isn't soluble in cold water it is acting like something which is very soluble, not so much in terms of how much dissolves (the crystals are too small to know how much exactly I have) but 5g/100mL into 20mL shouldn't be able to produce this much cooling. Or should it?

Actually, one piece of additional information. Despite having gotten out a reasonable amount of solid. Should a MAXIMUM of 100g of KBr, dissolved in 180mL of solution be able to crash out enough solid to half fill the beaker with slurry when cooled in an ice bath?

Ok, so I panicked. I just killed 2 birds with one stone, doing a solubility test while recrystallising all the product I have so far. It must not have been packed as tight as I thought, because it was ~70-80mL of what seemed like near dry crystals, but at density>3 that is impossible. Say it was at least 100g. It would not dissolve in 200mL of water despite best efforts... As it dissolved from the bottom it was snowing down from the top which was being evaporatively cooled by stirring. It dissolved with difficulty in 300mL of boiling water. Definitely not KBr then.

[Edited on 11-11-2007 by Antwain]

[Edited on 11-11-2007 by Antwain]

[Edited on 11-11-2007 by Antwain]

woelen - 10-11-2007 at 14:47

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).

chloric1 - 10-11-2007 at 17:01

Quote:
Originally posted by woelen

I use the bromate mainly for making bromine. It is a fantastic source for making bromine. Simply mixing it with a 5-fold molar amount of KBr or NaBr and adding this to 20% sulphuric acid makes the bromine drop out and simply pipetting is suitable for obtaining already 75% yield. For the remaining 25% you'll need a distillation setup. I do not store bromine, I make it when I want to experiment with it.




Woelin- I did the math and it looks great. I assume your 20% sulfuric acid is cool as not to vaporize your bromine. Question is when you make up your bromate/bromide mix in water how much water do you need? Or do you use the crystaline salts into the acid? Is this reaction exotheric to any great extent?
I once made bromine by adding 30% H2O2 to concentrated HBr. The was a slight delay and then a fast reaction occured vaporizing most of the bromine sending me running for higher ground.

woelen - 11-11-2007 at 04:00

I use the solid salts for making the bromine. I mix solid KBrO3 and KBr (or NaBr) in a molar ratio of 1 : 5. I use a slight excess of bromate, in order to avoid the problem of bromine, dissolving very well in solutions, which contain bromide.

As a start, try the following for making 10 grams of bromine, of which appr. 7.5 grams can simply be pipetted from below the aqueous layer.

Take 3.5 grams of KBrO3 and 12 grams of KBr (or 10 grams of NaBr), and finely grind the chemicals and mix them thoroughly. No need to fear any reaction, as long as no acid is added.

Take 40 ml of cool 20% H2SO4. Add the solid mix, 2 gram at a time and stir after each addition. Soon, you'll see bromine dropping out. When you have added all solid, then cap the container and carefully shake, in order to have all salt mix dissolved. Lots of bromine drop out. After 30 minutes or so, you'll have a clear liquid (bright red/orange) with some white crap (mostly KHSO4/K2SO4) and a big blob of bromine at the bottom. With a long pasteur pippette you simply suck up this bromine. Yield: 75% to 80%, appr. 2.5 ml. The remaining bromine is dissolved in the red/orange liquid. You also should keep this liquid, bromine water also is quite interesting for many experiments. If you want to go to the max, then you can distill the bromine water to get the remaining 20% of bromine. I, however, do not take that effort, I simply keep the bromine water.


You can scale up this procedure, the reaction is only slightly exothermic. Use 2 gram lots for each addition. So, on scaling up, you'll need more time for making the bromine, but it works equally well.

Be careful with this, LOTS of bromine vapor are created, and when swirling the solution, keep it capped, or covered with a piece of glass!

garage chemist - 11-11-2007 at 06:01

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.

chloric1 - 11-11-2007 at 06:52

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.

garage chemist - 11-11-2007 at 08:25

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.

Antwain - 13-11-2007 at 21:20

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

Antwain - 13-11-2007 at 21:24

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?

Antwain - 13-11-2007 at 21:40

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%

woelen - 13-11-2007 at 23:26

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).

sodium bromate

chloric1 - 14-11-2007 at 15:25

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.

Klute - 22-5-2008 at 17:10

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.

woelen - 23-5-2008 at 01:58

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.

Klute - 23-5-2008 at 02:46

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.

12AX7 - 23-5-2008 at 05:00

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

woelen - 23-5-2008 at 11:35

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]

Klute - 23-5-2008 at 13:13

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 :)

dann2 - 23-5-2008 at 15:32

Perhaps this has been posted before somewhere

Attachment: bromates.pdf (1.1MB)
This file has been downloaded 756 times

Polverone - 23-5-2008 at 16:27

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.

MagicJigPipe - 23-5-2008 at 18:24

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]

Klute - 24-5-2008 at 09:47

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.

MagicJigPipe - 24-5-2008 at 11:02

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.

woelen - 24-5-2008 at 13:03

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 :P . 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.

MagicJigPipe - 24-5-2008 at 17:56

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?

12AX7 - 24-5-2008 at 20:28

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

not_important - 24-5-2008 at 21:07

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.

JohnWW - 25-5-2008 at 02:47

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.

dann2 - 25-5-2008 at 13:15

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 815 times

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