Pyrovus
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ClO4 !?
I ran across this passage in an old chemistry book:
"If silver perchlorate is treated with bromine, silver bromide and free perchlorate radical, ClO4 result: 2AgClO4 + Br2 -> 2AgBr + 2ClO4. This
substance also has an odd number of valence electrons, and is extremely reactive and unstable."
It doesn't go into any further detail, and I haven't heard this compound mentioned anywhere else. This certainally sounds like an
interesting substance - if it hangs around for more than a few seconds, that is, but presumably it does last for a macroscopic timescale as the
authors of the book felt it worth mentioning.
One thing I find baffling is how bromine is able to oxidise the perchlorate ion. Surely the perchlorate ion would oxidise the bromine: 2ClO4- +Br2
-> 2BrO4- + Cl2?
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vulture
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I think this is connected to the fact that perbromate couldn't be isolated except by radiologic transmutation of radioactive perselenate.
One shouldn't accept or resort to the mutilation of science to appease the mentally impaired.
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BromicAcid
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| Quote: |
2AgClO4 + Br2 -> 2AgBr + 2ClO4
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That would put chlorine in a +8 state and I don't think it would appreciate that. I would guess that the chlorine tetroxide would dimerize like
so:
2ClO4 ----> Cl2O7 + 1/2O2
Being that dichlorine heptoxide is significantly more stable. Maybe they detected the ClO4 by the Cl2O7 produced and figured it was an intermediate,
I've got some fairly detailed books at my library on chlorine oxides so I can look this one up.
| Quote: |
Surely the perchlorate ion would oxidize the bromine: 2ClO4- +Br2 -> 2BrO4- + Cl2
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The perbromate ion has too high a reduction potential.
BrO4- (aq) + 2H+ (aq) + 2e- ---> BrO3- (aq) + H2O (l) E = +1.74
"By contrast the reduction potential for perchlorate is +1.23 V; and for periodate +1.64V. Hence only extremely strong oxidizing agents such as
xenon difluoride and difluorine are capable of oxidizing bromate to perbromate..." Descriptive Inorganic Chemistry 3rd ed
It's funny to note that perbromic acid is surprisingly stable, it can be concentrated to 55% solution and it's salts are stable. It is a
very powerful oxidizing acid.
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GeneralChemTutor
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Bromine is oxidized in this reaction.
[Edited on 24-2-2004 by GeneralChemTutor]
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BromicAcid
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| Quote: |
Bromine is oxidized in this reaction.
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Which reaction? You mean in the initial: 2AgClO4 + Br2 -> 2AgBr + 2ClO4 where it's reduced? Or are you saying that it would be oxidized in
the reaction and not preform as indicated in the above reaction?
Or are you refering to my comments about a reduction potential? If so then reverse the reaction to get the -E oxidation potential. The perbromate
ion is very difficult to make, this reaction will not do it, otherwise it would have been isolated prior to 1968, of course the driving force is the
insolubility of the silver bromide produced so the reaction itself does not oxidize the bromine. Although once it starts bromine in solution will be
oxidized to oxoacids of bromine but not the perbromate, maybe some bromine monoxide as well as the other oxides of bromine are also very unstable.
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Theoretic
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Probably ClO4 does exist for a significant length of time, either as a free radical or its dimer, Cl2O8 (diperchloryl peroxide -
"(ClO3+)2O2--" . It would probably follow the trend which is started by
chlorine dioxide ClO2 (does not dimerise) and continued by chlorine trioxide ClO3 (dimerises 99% to Cl2O6, exists in equilibrium with it). This means
that ClO4 would probably be present to a smallextent along witth the dimer Cl2O8
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GeneralChemTutor
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I was looking at the wrong equation...read through the problem too quickly. My apologies.
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BromicAcid
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Chlorine Tetroxide
Copied from "Supplement to Mellor's Complete Treatise on Inorganic and Theoretical Chemistry" Supplement II Part I copyright 1956:
| Quote: |
Chlorine Tetroxide
It has been claimed that the oxide ClO4 is produced by the reaction of silver perchlorate with iodine in
anhydrous ether, the concentrations of the reactants being about 1%:
2AgClO4 + I2 ----> 2AgI + 2ClO4
The presence of a small amount of unstable iodo-compound, not volatile in ether, and probably dimeric, (ClO4)2. It was believed to react with water
according to the following equation:
2ClO4 + H2O -----> HClO4 + HO.ClO4
It has also been regarded as an intermediate in the reaction between perchloric acid and fluorine. The following facts are submitted as evidence that
the solutions in ether contain chlorine tetroxide rather then perchloric acid: (a) the solution liberated iodine from hydroiodic acid, thogh not in
equivalent amount; (b) with zinc and magnesium metals, perchlorates were produced but no hydrogen was evolved. (The reaction could thus be explained
in terms of a single electron trasfer process: with iron, tin, and copper, perchlorates of the metals in more then one valency state were produced,
and with cadmium, bismuth, and silver, reaction was slow and incomplete). (c) The reactions of solutions of perchloric acid and of ethyl perchlorate
differed from those of the chlorine tetroxide solutions.
The evidence for the existence of this oxide was in no way conclusive, and an alternative interpretation of the experimental data has been proposed
involving the intermediate formation of iodine perchlorate.:
AgClO4 + I2 -----> AgI + IClO4
followed by iodination of ether to give an unstable iodo-ether:
C2H5.O.C2H5 + IClO4 -----> C2H5.O.C2H4I + HClO4
The iodo-ether slowly reacts with more silver perchlorate:
AgClO4 + C2H5.O.C2H4I -----> AgI + C2H5.O.C2H4.ClO4
The overall reaction therefore leads to the formation of equivalent amounts of perchloric acid and ether perchlorate. On this basis, the oxidizing
power of the solution should vanish towards the end of the reaction, as the iodine perchlorate is consumed in iodinating the solvent. This has been
found to be the case, although the obeservation has been contested. The second and third stages are difficult to confirm because of the instability
of the iodinated ether, but receive support from later work. When silver perchlorate and iodine react in the presence of an aromatic compound C6H5X
(e.g. chloro- or nitro-benzene) the reaction is exclusively:
C6H5X + I2 + AgClO4 -----> C6H5XI + AgI + HClO4
and there is no evidence of the formation of chlorine tetroxide. The aromatic iodo-compound is not affected by excess of silver perchlorate. However
alkyl iodides do react:
C2H5I + AgClO4 -----> C2H5ClO4 + AgI
When iodine and silver perchlorate react in ether solution, the initial precipitation corresponds to 1 mol. of silver iodide per mol. of iodine.
Reaction with the solvent is an essential step in the formation of the second molecule of silver iodide. This precipitation is slow and during the
process the oxidizing power of the solution dissappers. The original postulate that chlorine tetroxide was formed was based on the assumption that
ether behaved as an inert solvent: since this appers to be untrue there is no apparent justification for supposing that chlorine tetroxide is formed
in the reaction between iodine and silver perchlorate.
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Still though I would like to find some more recent sources, preticularly from the re-emergence of inorganic chemistry in the 70's.
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BromicAcid
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From Preparative Inorganic Reactions Vol 1
A newer chemistry source from the 1970's on this reaction, exactly what I asked for above:
| Quote: | Halogen Perchlorates
Recent work in nonaqueous solvents has shown that at -85C. in ethanol, iodine and silver perchlorate react to give AgI and iodine(I) perchlorate,
IOClO3. Earlier work, mostly at room temperature, in ether was complicated by iodination of the solvent and from it the formation of ClO4 was
incorrectly posculated. The reaction of bromine and silver perchlorate in aromatic solvents has been studied and bromination of the solvent was
observed. The work did not state that BrOClO3 might have been the active reagent. Since IOClO3 was not isolated from the alcohol solvent, FOClO3 has
the distinction of being the only halogen (I) perchlorate to have been isolated. |
No mention of chlorine perchlorate though.
[Edited on 12/16/2004 by BromicAcid]
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JohnWW
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The ClO4 mentioned earlier in this thread would be a peroxide, in which one of the oxygens is single-bonded to the Cl and carries an unpaired
electron. The molecule would be resonance-stabilized by delocalization of this single bond and unpaired electron around the 4 Os. It would exist in
equilibrium with a dimer, Cl2O8, containing a peroxide bridge.
The most likely method by which it could be produced would probably be by electrolysis of a perchlorate solution at low temperatures. It would be a
fairly volatile compound, and highly explosive.
The claimed reaction 2AgClO4 + I2 ----> 2AgI + 2ClO4 in Mellor is MOST unlikely, especially as it involves reduction of I2 to I-. A more likely
reaction would be oxidation of I2 to iodite(III), iodonium(III), iodate(V), or periodate(VII), probably a mixture of these, with some of the Ag+ being
precipitated as AgCl.
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