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JohnWW
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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]
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Jor
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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!
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Jor
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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]
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AndersHoveland
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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]
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AndersHoveland
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(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]
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AndersHoveland
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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.
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woelen
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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.
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AndersHoveland
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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]
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woelen
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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.
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AndersHoveland
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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.
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AndersHoveland
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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
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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.
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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]
I'm not saying let's go kill all the stupid people...I'm just saying lets remove all the warning labels and let the problem sort itself out.
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AndersHoveland
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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]
I'm not saying let's go kill all the stupid people...I'm just saying lets remove all the warning labels and let the problem sort itself out.
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