Sciencemadness Discussion Board

Periodic acid -- properties and experiments

woelen - 7-4-2008 at 09:55

Although there is some information about this compound on sciencemadness, I think it is appropriate to create a dedicated thread about the properties of this compound. Information about this compound is sparse. The only thing I could find is that it can be used to cleave -C(OH)-C(OH)- groups and that it is a strong oxidizer.

woelen - 7-4-2008 at 09:56

As a start, I can give some nice experiments I did.

1) A solution of H5IO6 gives a white precipitate of KIO4 when a small amount of a solution of KOH is added. When more KOH is added, the precipitate redissolves again. When dilute nitric acid is added, then a white compact crystalline precipitate is formed again.
This is quite interesting. It looks like being amphoteric. With a little amount of base, the IO4(-) ion seems to be formed, which gives a precipitate with K(+) ions. At higher pH, IO6(..) based ions are formed?
The KIO4 can be separated very easily. Just rinsing with distilled water, decanting the liquid, having the liquid absorbed in tissue paper and then drying. The KIO4 is an energetic material which burns fast with metal powder, and it sort of explodes with red P.

2) A mix of slightly crunched H5IO6 (30 mg) with powdered red P (7 mg) has the interesting property that it is self-igniting with a time delay. The self ignition occurs when the mix is put in a pile. Just wait and at a certain moment it ignites. When the powdered mix is spread over a large area, then it does not ignite.
The reaction between red P and H5IO6 is very fast, like flash powder. The mix does not seem very friction sensitive, I could not ignite it by crunching/rubbing the mix.

3) The compound H5IO6 reacts with conc. HCl, giving loads of Cl2 and yellow ICl3, which remains in solution.

I did not try the cleaving method of di-ols. Does this mean that sugars like sorbitol (which contains OH groups on each C-atom) are completely broken down? Under which conditions do such reactions occur?

pantone159 - 7-4-2008 at 10:35

Checking some of my books...

Organic Chemistry, Brown and Foote, 3rd ed:

The mechanism for splitting glycols involves first forming a five-membered cyclic periodic ester, and electron redistribution then gives two C=O groups. Glycols which cannot form such an intermediate (i.e. steric reasons) cannot undergo such a reaction.

This reaction is said to be useful in analyzing sugars, by analyzing the amount of periodic acid used and the products. Also, the periodic acid can similarly attack a-hydroxyketones and a-hydroxyaldehydes.

Macroscale and Microscale Organic Experiments, Williamson, 2nd ed., has the following test for vicinal glycols:
Prepare periodic reagent by dissolving 0.25 g H5IO6 in 50 ml H2O.
To 2 ml of this reagent, add 1 drop (no more) of conc HNO3 and shake. Then add one drop or a small crystal of the unknown. Shake for 15 sec and add 1-2 drops of 0.3 M AgNO3 (aq). Instantaneous formation of a white ppt is a positive test.

This test gives positive results with a-hydroxy aldehydes, a-hydroxy ketones, a-hydroxy acids, a-amino alcohols, 1,2-diketones, as well as for the 1,2-glycols.

EDIT: Also, Solomons, 4th ed, states that lead tetraacetate brings about similar cleavage reactions. The lead compound works well in organic solvents, while periodic acid works well in aqueous solutions, so the two are complementary.

The periodic reaction usually has quantitative yields.

EDIT2: I believe that sorbitol would indeed be broken down completely, the end carbons would become formaldehyde and the inner ones formic acid, so 1 mol sorbitol -> 2 mol formaldehyde + 4 mol formic acid.


[Edited on 7-4-2008 by pantone159]

[Edited on 7-4-2008 by pantone159]

Ammonium periodate

woelen - 12-4-2008 at 04:57

Today I made ammonium periodate. It is made very easily. If you make a solution of H5IO6 in a little water and then add some ammonia, then a compact white crystalline precipitate is formed. This precipitate does not dissolve when much more ammonia is added, not even in 25% ammonia.

I let this precipitate dry above a heat radiator (50 C or so). A fine non-hygroscopic white powder is obtained. This compound is a remarkably powerful explosive.

I put a very small amount in a thick-walled test tube, unconfined, and then heated the test tube above a small alcohol burner. The result is quite impressive. At a certain point in time, the solid explodes with a remarkably loud report and iodine vapor is produced at once. I made a video of this reaction, a small write-up on this will follow tonight or tomorrow.

What most surprises me is how easily this compound can be made from periodic acid. Apparently, periodates belong to a class of anions, which hardly have any soluble compounds. Even the potassium and ammonium salts are almost completely insoluble in water. Only the meta-periodates are like this. When NaOH is added, then the solid dissolves and ortho-periodates go into solution.

Zinc - 12-4-2008 at 05:17

Very interesting:P

I wish I have some periodic acid. Does anyone know where could it be obtained OTC or do I have to order it from some chem. company?

And could organic periodates be made (I know perchlorates can and they are very explosive)?

sparkgap - 12-4-2008 at 09:35

@woelen: If you want to test the polyol-cleaving property of periodic acid, even sucrose will suffice. :D

@Zinc: AFAIK, periodic acid cannot be had OTC.

sparky (~_~)

woelen - 12-4-2008 at 12:26

Here follows the write-up on the preparation and explosive properties of ammonium periodate:

http://woelen.homescience.net/science/chem/exps/ammonium_per...

I'll try the cleavage reaction with sugar. I expect bubbles of CO2, first cleaving and then further oxidation to CO2. If not, then I at least should smell formaldehyde and/or formic acid. Both chemicals have strong and distinct smells.

[Edited on 9-5-11 by woelen]

chloric1 - 12-4-2008 at 13:51

Periodic acid can be quite expensive. If I recall, a certain ebay seller sells little 1 or 4 oz bottles along with iodides and few other chems. His products are concerned with printing.

$62.99 for just25 grams JEEZ! Just search ebay for "iodide" and you will find the store.

[Edited on 4/12/2008 by chloric1]

garage chemist - 12-4-2008 at 15:13

Periodates are remarkably easy to make, much easier than perchlorates.
Check out Brauer, there's a detailed synthesis of NaIO4 in there in three steps from iodine.
The first one is HNO3-catalyzed oxidation of iodine with NaClO3, forming NaIO3 solution, second one production of sodium periodate by oxidation of the iodate with chlorine in NaOH solution.
It is precipitated as the sparingly soluble salt Na3H2IO6, and in the last step this is reacted with HNO3 to form the easily soluble NaIO4 which is isolated by fractional crystallization after concentration according to the procedure.
From NaIO4, ammonium periodate can be easily prepared without a doubt.

Making free crystalline periodic acid requires the barium salt Ba3(H2IO6)2, which is reacted with an excess of conc. HNO3 in which Ba(NO3)2 is insoluble. The procedure is given in Brauer as well.

Everyone interested in periodate preparation check out those procedures in Brauer, they are very nice laboratory procedures that surely are a pleasure to do and don't require any exotic chems!

Woelen, very interesting work with the ammonium periodate. I am surprised that this salt is explosive by simple heating, in contrast to ammonium perchlorate, which only decomposes quite silently with a yellow flame when heated.

woelen - 13-4-2008 at 07:00

I also tried to make hydrazinium periodate, but this is not possible. As soon as H5IO6 is brought in contact with hydrazine, or one of its salts, then an EXTREMELY violent reaction occurs. I have never seen a reaction between two aqueous solutions, which is so violent and hot, that iodine escapes from the liquid as vapor.

This is a remarkable experiment in its own. If one has H5IO6 and hydrazine (or one of its salts) then make a solution of H5IO6 in a small amount of water, make a solution of hydrazine or its salt (e.g. N2H4.2HCl) and then add the two solutions to each other. This results in the formation of a thick plume of purple iodine vapor!

YT2095 - 13-4-2008 at 07:23

sounds like wet chemistry version of Rocket fuel, almost Messerschmitt worthy :)

Axt - 13-4-2008 at 10:49

You can produce hexaiodobenzene (see attachment). Then attempt to <a href="http://www.sciencemadness.org/talk/viewthread.php?tid=4766">oxidise to iodyl</a>, though I did do a lot of searching and never did I come across reference to iodyl groups ortho to each other.

Or see if it changes colour with temperature http://www.chemicalforums.com/index.php?topic=19644.0

Somewhat novel, it forms copper(III) periodate complexes as well.

Attachment: Periodination of Benzene.pdf (284kB)
This file has been downloaded 2314 times

woelen - 13-4-2008 at 11:56

This periodination of benzene looks very interesting. Unfortunately, benzene is VERY difficult to obtain in NL, it is one of the chemicals which is banned completely. I do have toluene, what would you expect with toluene, under the conditions, as described in the article.

Thanks anyway for all this new info. A whole new world of iodine chemistry opens up to me :), this old thread was buried too deep, so I have not seen it before. I especially like the benzene with the ICl2 groep on it. That really is interesting chemistry!

Axt - 14-4-2008 at 12:56

Yep, I've attached the follow up article which includes pentaiodotoluene amoungst others.

Oh and I'm an idiot, I posted myself in the iodoxy thread that o-diiodylbenzene had been prepared, so yes theres an example of two iodyl moieties orto to each other. The highest "iodylated" benzene i've seen is 1,3,5-triiodylbenzene, and was only mentioned as an "exotic initiator of the future" :/ No reference nor indication that it has been made.

[Edited on 15-4-2008 by Axt]

Attachment: Direct Aromatic Periodination.pdf (440kB)
This file has been downloaded 1422 times

UnintentionalChaos - 14-4-2008 at 14:41

Quote:

Somewhat novel, it forms copper(III) periodate complexes as well.


Attached a paper explaining somewhat vaguely this process.

Oxidation of Cu (II) by persulfate or hypochlorite in alkaline medium in the presence of periodate.

Attachment: Copper (III).pdf (157kB)
This file has been downloaded 1384 times

woelen - 4-5-2008 at 08:07

@garage chemist: I have done the preparation of KIO4, based on a slightly modified procedure, as described in the reference you gave. This really works nice. I have a lot of KIO3, and now I can easily make KIO4. I already had some, but only a very limited amount (it is soooo expensive). Making the sodium salt also is easy for me, I have a lot of NaI and the periodate can directly be made from a boiling solution of NaI, NaOH through which Cl2 is bubbled.

A write-up with pictures will follow on the synthesis of KIO4.

woelen - 5-5-2008 at 12:49

Here follows the write-up of the synth of KIO4:

http://woelen.homescience.net/science/chem/exps/KIO4_synth/i...

Have fun and try it yourself, it really is easy.

[Edited on 9-5-11 by woelen]

woelen - 18-5-2008 at 11:31

Quote:
Originally posted by woelen
I also tried to make hydrazinium periodate, but this is not possible. As soon as H5IO6 is brought in contact with hydrazine, or one of its salts, then an EXTREMELY violent reaction occurs. I have never seen a reaction between two aqueous solutions, which is so violent and hot, that iodine escapes from the liquid as vapor.

This is a remarkable experiment in its own. If one has H5IO6 and hydrazine (or one of its salts) then make a solution of H5IO6 in a small amount of water, make a solution of hydrazine or its salt (e.g. N2H4.2HCl) and then add the two solutions to each other. This results in the formation of a thick plume of purple iodine vapor!

I did some further experimenting on this and under no condition, periodate and hydrazine can coexist. Not at high pH as free base, nor at low pH as hydrazinium ion. In all cases there is a violent reaction. This reaction is fun in its own. I made a small webpage of this:

http://woelen.homescience.net/science/chem/exps/hydrazine_pe...

[Edited on 9-5-11 by woelen]

Jor - 18-5-2008 at 11:50

Very nice. I'm gonna make some KIO4 really soon! I have about 100g KIO3. I Also have hydrazine sulphate.

I'm gonna try this with potassium bromate. Should be really violent!

Formatik - 18-5-2008 at 12:21

Quote:
Originally posted by woelen I did some further experimenting on this and under no condition, periodate and hydrazine can coexist. Not at high pH as free base, nor at low pH as hydrazinium ion. In all cases there is a violent reaction. This reaction is fun in its own. I made a small webpage of this:

http://woelen.homescience.net/science/chem/exps/hydrazine_pe...


How about in liquid nitrogen, it still reacts violently?

Edit(woelen): Made link work again.

[Edited on 30-7-16 by woelen]

JohnWW - 18-5-2008 at 13:41

Quote:
Originally posted by JorI'm gonna make some KIO4 really soon! I have about 100g KIO3. I Also have hydrazine sulphate. I'm gonna try this with potassium bromate.
I presume you are going to try to make perbromate. Perbromates are very difficult to prepare, being very strong oxidants, and were first obtained in any quantity only relatively recently, unlike perchlorates and periodates. Before that, they were known only in small quantities through the decay of a short-lived Se isotope as selenate, which is itself oxidizing.

Jor - 19-5-2008 at 03:03

I know perbromates are very hard to make.
I ment that I was going to substitute a preiodate by a bromate in the experiment on woelen's site.

I always wondered why it is so impossible to make perbromate without F2 or XeF2 though. Things like ferrate(VI), manganese heptoxide, persulphate, bismuthate, cerium(IV) etc. should be able to do the job in theory right? Because they are stronger oxidisers than perbromate. I can get all of these chems (except ferrate, but it i possible to prepare barium ferrate >>> see versuchschemie), so i will try it some day.

Why is it so hard to make? Is it the Br going through oxidation state 6+ before achieving 7+? Maybe 6+ is so unstable that it will decompose immidiately? I have no idea...

Jor - 19-5-2008 at 10:51

I mixed some a small amount of potassium bromate with a small amount of potassium permanganate on a petri dish. To this a large drop of conc. sulphuric acid was added. Something was being oxidised, because the solution went brown. This must have been the bromate being oxidised. But yes, as expected, there was a lot of bubbling. A saw a brown tint; indicates bromine. But there was also a lot of colorless gass : oxygen?
What exactly is happening here?

ShadowWarrior4444 - 19-5-2008 at 10:58

Quote:
Originally posted by Jor
I mixed some a small amount of potassium bromate with a small amount of potassium permanganate on a petri dish. To this a large drop of conc. sulphuric acid was added. Something was being oxidised, because the solution went brown. This must have been the bromate being oxidised. But yes, as expected, there was a lot of bubbling. A saw a brown tint; indicates bromine. But there was also a lot of colorless gass : oxygen?
What exactly is happening here?


Perhaps the production of elemental bromine with the evolution of oxygen. Did you detect any characteristic scents?
Quote:
I can get all of these chems (except ferrate, but it i possible to prepare barium ferrate >>> see versuchschemie), so i will try it some day.


Potassium ferrate is quite easy to prepare, and is used in 'Super-Iron' batteries. It is usually stabilized by the hypochlorite ion.

[Edited on 5-19-2008 by ShadowWarrior4444]

Jor - 19-5-2008 at 11:53

According to wiki, perbromic acid rapidly decomposes to bromic acid and oxygen, with the evolution of bromine vapor.

I read in a ACS-journal that bromate is not appreciably oxidised by persulphate at 100C. If my English is correct, this would mean that some bromate is oxidised. Worth a try , if you do it for a long time, the boiling.

JohnWW - 19-5-2008 at 15:34

There is an article on the preparation of KBrO4 in Inorganic Syntheses vol.13 (2007), described at:

http://doi.wiley.com/10.1002/9780470132449.ch1 or
http://www3.interscience.wiley.com/cgi-bin/summary/114037735... or http://www3.interscience.wiley.com/cgi-bin/booktext/11403773... (1,618 Kb) .

Someone with access to a library with a Wiley subscription, please download it, and post it here.
PS Someone has posted this article, from Inorganic Syntheses vol.13 (1972), in the References section. Thanks.

[Edited on 21-5-08 by JohnWW]

Jor - 27-5-2008 at 02:23

For some reason I don't have access to the References section anymore. I had the password from vulture, but it doesnn't work anymore. Has the pass changed?

Time to do some experiments with periodic acid! Just recieved 25 grams, together with 500mL ether, 25 gram mercuric oxide, 100g ammonium dichromate, 100g potassium dichromate, 250ml 0,1M silver nitrate, 100ml ethyl acetate, 25 phenolphtalein, 250g zinc chloride and 50g malonic acid for only 32 EURO, including shipping. Good buy IMO!

Jor - 29-5-2008 at 11:30

I was wondering, can perbromates be prepared by dry reaction of high oxides like NiO2 or AgO with potassium bromate? Like a pyrotechnic mixture.
Or maybe with ferrates or bismuthates, but not in water, but dry powders.

[Edited on 29-5-2008 by Jor]

AndersHoveland - 7-5-2011 at 15:16

Quote: Originally posted by BromicAcid  

I've always wanted to make perbromic acid, and aside from Xenon Difluoride this was the only way.


Surely there must be some other creative combination of regents which can oxidize bromate to perbromate? What about the slow addition of peroxysulfuric acid to a mixture of permanganese heptoxide and bromate in perfluorooctanesulfonyl fluoride solvent? Supposedly highly reactive atomic oxygen is an intermediate in the decomposition of Mn2O7. Alternatively peroxydisulfuryl difluoride (FSO2OOSO2F) likely would work. Consider this, the compound actually has an equilibrium with the SO3F radical (fluorosulfonate without the extra electron)!
reference: http://www.sciencedirect.com/science?_ob=ArticleURL&_udi...

Supposedly, one of the strongest oxidation reactions that can be done with reagents stable in aqueous solution is acidifying a ferrate(VI) salt (yes that is correct +6 oxidation state for iron). This is more oxidizing than acidified permanganate, which makes one wonder why permanganates are so commonly used for oxidation reactions when iron is much more commonly available.
http://www.youtube.com/watch?v=pUvdETUQPuo
The redox potential of ferrate is 2.2 volts in acid and 0.7v in base. Note that chlorine is 1.4v and fluorine is 2.87v.

There was also someone in Chemical Forums who posted a topic about whether ferrate(IV) could oxidize bromate into perbromate.

Some additional information that "BromicAcid" may find of interest:

The acids HClO3 and HBrO3, unlike HIO3 , cannot be isolated from their aqueous solutions, and attempts to concentrate them brings about their decomposition. For example, in the case of HClO3 the following occurs:
(3)HClO3 ➝ HClO4 + (2)O2 + Cl2 + H2O

Unlike chloric acid, unfortunately, HBrO4 is not formed from the decomposition of bromic acid.

[Edited on 7-5-2011 by AndersHoveland]

AndersHoveland - 8-5-2011 at 16:40

(moderator: please do not merge this post with the one above. Too much time has elapsed and the forum would not allow me to edit the old one. If the posts are merged, the link included in this post will then not work)

Neither persulfate, nor ozone is capable of oxidizing bromate. However from electrolytic experiments, hydroxyl radicals appear to be able to form perbromate. Good results have been shown with conductive-diamond coated electrodes. In order to form any perbromate high current densities are required, greater than 300A per square meter. Generation of perbromate is favored at 20-30C.
“Perbromate was easier formed starting the electrolysis from bromide and not from bromate. Bromide electrolysis is a new and promising method for synthesizing perbromate that was recently confirmed.” Saez, C. et al.,J. Appl. Electrochem., Online publication. March 5, 2010

Potassium Ferrate (K2FeO4) has only been commercially available very recently. "Ferratec and Electrosysthesis announce availability of the powerful chemical oxidizer ferrate by the kilogram. Potassium ferrate(VI) (K2FeO4) has many applications, including use as a biocide, as a powerful oxidizer in organic synthesis and as a water treatment compound. The 95-percent pure product contains no chlorides or other halogen impurities, is shelf stable and environmentally friendly."

I made a convenient list of reduction potentials here:
https://sites.google.com/site/ecpreparation/ferrate-vi

Is it possible to make AgFeO4 silver(II) ferrate ? Perhaps with finely powdered K2FeO4 and AgF in pentafluoropyridine solvent? No doubt it would immediately hydrolyze on contact with water, but it could potentially be a very strong oxidizer.

[Edited on 9-5-2011 by AndersHoveland]

AndersHoveland - 14-9-2011 at 13:25

Not sure if this is true, since the article is very old:

"...an aqueous solution of perbromic acid is easily obtained by the action of bromine on the hydrate of perchloric acid dissolved in water, and that this solution, when neutralized with caustic potash, deposits crystals of potassium bromate."
Journal of the Chemical Society, Volume 27
On the Perbromates, M. M. Pattison Muir

This would indeed be very surprising since perchloric acid is virtually inert in terms of its oxidizing reactivity below 70% concentration at room temperature, and bromates are extremely difficult to oxidize to perbromates- typically fluorine (bubbled into aqueous solution) or hydroxyl radicals are required.

woelen - 14-9-2011 at 22:55

Indeed, this article must be utter nonsense! Perchloric acid is rather inert (in terms of oxidation, not in terms of acidiity) in concentrations up to 70%. I personally tried this by adding NaI to hot 70% HClO4 and even Na2SO3. In the latter case just SO2 escapes and in the case of NaI nothing happens at all, the liquid only turns very pale yellow, but this most likely is iodide oxidized by oxygen from the air. Bromide certainly will not react al all.

Anhydrous HClO4 is another matter. That stuff is extremely energetic and will react violently or even explosively with the chemicals mentioned above. But with bromide I do not expect any perbromate.

AndersHoveland - 15-9-2011 at 10:47

It is, however, plausible that highly concentrated HClO4 (over 70% concentration), might oxidize bromine to perbromate. Anhydrous HClO4 can oxidize CCl4 at room temperature after a short time, whereas CCl4 can be used as a solvent for ozone. It can thus be inferred that anhydrous HClO4 is a stronger oxidizer than ozone.

The monohydrate of perchloric acid, HClO4·H2O, contains 84.6% perchloric acid, so would qualify as the reactive form.

A study of the oxidation potential of perchloric acid (in the concentration range 70–80 %) showed a probable redox value of 2.0–2.1 V or higher.
"The role of 70–80% perchloric acid as oxygen donor and the oxidation potentials made available"
G. Frederick Smith

For comparison, the redox potential of ferrate (while being acidified) is 2.2V, and that for hydroxyl radicals (which are known to be able to oxidize bromate to perbromate) are 2.8V. Ozone has a potential of 2.07V.

"When cold, perchloric acid solutions are non-oxidizing at any acid concentration below 73%. The oxidation potential of hot concentrated perchloric acid is [roughly around] 2.0 V"
John R. Long, GFS Chemicals Inc.

It is interesting that the 73% reactivity threshold is so close to the azeotropic 72.4% concentration (the highest concentration that can be achieved by boiling out the water, the azeotropic solution boiling at 203C). The 72.4% concentration is close to a dihydrate composition.

I do not know if 98% concentrated HClO4 would be more oxidizing (have a higher redox value) than 80%. It may be possible, but I doubt this is the case, since molecules of HClO4 likely exist in equilibrium in the monohydrate.

the monohydrate of perchloric acid is a solid which melts at 50°C, whereas both the azeotropic concentration and the anhydrous form are liquids, even below 0°C.

[Edited on 15-9-2011 by AndersHoveland]

woelen - 15-9-2011 at 22:41

I have no possibility to try out reactions with more than 70% HClO4 (which is very hard to obtain and also hard to keep around safely for a longer time), but I do not believe that reaction of anhydrous HClO4 with bromine or bromide will give any perbromate. I expect that the reaction products just will be decomposition products, mainly oxygen, chlorine, chlorine oxides and water. The bromide/bromine may be oxidized to all kinds of oxo-species of bromine which in the heat of the reaction also will decompose to bromine and oxygen.

AndersHoveland - 16-9-2011 at 12:15

one source states that the oxidizing potential of acidified perbromate is 1.763 V.
Perbromic acid is unstable and quickly decomposes. The oxidizing potential might be significantly higher if KBrO4 is added into pyrosulfuric acid (also known as superconcentrated sulfuric acid, oleum, H2S2O7), perhaps more so than perchloric acid.
This would no doubt be a very violent reaction.

Apparently bromate can be oxidized by hypobromite to perbromate, at a highly alkaline 12.5pH at 40degC.
BrO[-] + BrO3[-] --> BrO4[-] + Br[-]

"Two New Methods of Synthesis for the Perbromate Ion"
Aleksey N. Pisarenko, Robert Young, Oscar Quiones, Brett J. Vanderford, and Douglas B. Mawhinney
http://pubs.acs.org/doi/abs/10.1021/ic201329q

It should be mentioned that this paper came from a somwhat less reputable research institution, so one might be inclined to hold some reservations and be a little sceptical, in light of the well known unusual difficulty of oxidizing bromate.

AndersHoveland - 2-11-2012 at 23:54

Required redox potential for oxidizing bromate to perbromate

A redox value of -1.763 V is given for the following reaction:

BrO4[-] + 2H[+] + 2e[-] --> BrO3[-] + H2O

G. K. Johnson, P.N. Smith, E.H. Appleman, W.N. Hubbard, Inorganic Chemistry, 9, 119 (1970).

This value is not that high, but unfortunately neither bromate nor perbromate likes acidic conditions and it will just decompose, so that excludes most of the typical oxidizers.

Just to give some idea of the problem, I will include these values:
Ozone (aqueous acidic solution) 2.08v
Ozone (aqueous neutral solution) 1.24v
Ferrate (acidified) 2.20v
Ferrate (alkaline) 0.72v

Boiling a solution of persulfate with bromate might be enough to oxidize the bromate to perbromate, if this -1.763 V is actually correct. Boiling the solution causes the persulfate to break into more powerful persulfate radicals (2.6v).


Possibility of oxidizing bromate using hydrogen peroxide

I initially had the idea that perhaps an alkaline solution of hydrogen peroxide might be able to work here, since all sorts of powerful oxidizing species are formed during the base-catalysed decomposition of hydrogen peroxide. But then I remembered a paper that mentioned that hydrogen peroxide was unable to oxidize chlorate.

Hydrogen peroxide failed to oxidize chlorate to perchlorate under alkaline, neutral, or acidic conditions, although minute traces of perchlorate did form under acidic conditions.
“Electrolytic Formation of Perchlorate” C. W. Bennett, E.L. Mack. Chemical Engineer, volume 23, p206
Quote:

A boiling solution of sodium peroxide failed to oxidize chlorate to perchlorate.
A solution containing one gram sodium chlorate and 1cc ammonium hydroxide (specific gravity 0.90) in 15 cc hydrogen peroxide (30%) was boiled for 30 minutes. Analysis of the solution failed to show any traces of perchlorate, thus showing that alkaline hydrogen peroxide is not a sufficiently powerful oxidizer to convert chlorates to perchlorates.


Still, I would not be so quick to dismiss the possibility of using alkaline hydrogen peroxide as an oxidizer. Perhaps whatever reactive species formed reacted much faster with the NH4OH before it had a chance to oxidize any of the chlorate? This experiment should be repeated using NaOH instead, without boiling, and with the hydrogen peroxide given plenty of time to fully decompose under the catalytic action of the base. And then also with boiling.

If the conditions can be optimised to get alkaline H2O2 to be able to oxidize chlorate (if that is possible), it might similarly be able oxidize bromate to perbromate. Not an easy thing to do, but there are some very reactive intermediaries (superoxide anions and hydrogen trioxide) during the base-catalysed decomposition of H2O2.
For comparison, the reduction value of acidified hydrogen peroxide is 1.78v, a value slightly greater than what it should take to oxidize bromate, but again if it were not for the fact that this requires acidic conditions and so would decompose the bromate, and any perbromate that may form. (Bromate and perbromate are much less stable than chlorate/perchlorate/iodate/periodate, and under acidic conditions decompose into elemental bromine releasing bubbles of oxygen).

As far as I know, the only two methods that have been successful in oxidizing bromate to perbromate in the literature are with fluorine, and electrolytic oxidation (presumably through formation of hydroxyl radicals).

Fluorine 2.87v
Hydroxyl radical 2.80v

[Edited on 3-11-2012 by AndersHoveland]

AndersHoveland - 6-1-2013 at 07:29

Apparently ozone cannot directly oxidize chlorate, though it can oxidize chlorite, ClO2-, with a small ammount of perchlorate forming. In water, the yield of perchlorate from treating chlorite with ozone is only 2.7%, the remainder only gets oxidized to chlorate.
"Perchlorate Formation by Ozone Oxidation of Aqueous Chlorine/Oxy-Chlorine Species: Role of ClxOy Radicals",
Balaji Rao, Environ. Sci. Technol., 2010, 44 (8), pp 2961–2967

I remember reading in the literature somewhere that ozone was not successful at oxidizing bromate. This would hardly be surprising, as bromate would be even more difficult to oxidize than chlorate. Probably the investigators conducting the experiment never read the above article, if they had they might have tried reacting bromite with ozone, to see if any traces of perbromates form. (for potentially much higher yields, bromine dioxide dissolved in liquified CF4, reacted with ozone, might work well I think)

Fowler and Grant found that on heating chlorate with silver oxide that the chlorate was completely converted to perchlorate without loss of oxygen, metallic silver also forming.
J. Chem. Soc. 57, 272 (year 1890)

This might similarly work with bromate also, although potassium perbromate does decompose at around 275 °C.

Very recently a new method of making perbromate was discovered. Hypobromite can oxidize bromate to perbromate (40 °C, pH 12.5), the reaction was carried out over a 13 day period.
"Two New Methods of Synthesis for the Perbromate Ion", Aleksey N. Pisarenko, Robert Young, Oscar Quiñones, Brett J. Vanderford, and Douglas B. Mawhinney, Inorg. Chem., 2011, 50 (18), pp 8691–8693

Hypobromite can be easily formed by reacting bromine with base, but it is unstable, and at room temperature decomposes to bromate and bromide after around 20 minutes.

[Edited on 6-1-2013 by AndersHoveland]