fusso
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Why is selenic acid more oxidising than sulphuric?
Why is selenic acid more oxidising than sulphuric?
(I did utfse)
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woelen
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I do not know an explanation, but there is a trend in multiple columns of the periodic table.
If we look at P, S, Cl, then all of these can be brought into maximum oxidation states +5, +6, +7 relatively easily and the ions at these oxidation
states are quite inert in aqueous solution, even when the acids are present in aqueous solution at e.g. 60% they are nearly non-oxidizing.
If we look at As, Se, Br at oxidation states +5, +6 and +7, then all of them are strong oxidizers. E.g. As2O5 oxidizes HCl to Cl2, aqueous H2SeO4 at
higher concentrations can dissolve gold and HBrO4 is not even stable at higher concentrations. It also is much more difficult to bring these elements
to high oxidation states. Bromine only was obtained in oxidation state +7 fairly recently in the 1960's.
If we go down one more row, to Sb, Te, I, then we see that it is much easier again to go to the maximum oxidation state. Sb2O5 only is weakly
oxidizing (I know from experience, I have this chemical), H6TeO6 is not an oxidizing acid, H5IO6 is oxidizing, but quite stable, while HBrO4 is not
stable, unless present in low concentrations.
There is an anomaly in the row of As, Se, Br. I have read about quantummechanical explanations, but explaining it in simple language apparently is not
that easy. One needs to invoke fairly advanced concepts to explain this anomaly.
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clearly_not_atara
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It's a combination of the d-block contraction and the fact that larger atoms form weaker covalent bonds. Both of those effects destabilize Se(VI).
https://en.wikipedia.org/wiki/D-block_contraction
[Edited on 04-20-1969 by clearly_not_atara]
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fusso
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Quote: Originally posted by woelen  | If we look at As, Se, Br at oxidation states +5, +6 and +7, then all of them are strong oxidizers. E.g. As2O5 oxidizes HCl to Cl2, aqueous H2SeO4 at
higher concentrations can dissolve gold and HBrO4 is not even stable at higher concentrations. It also is much more difficult to bring these elements
to high oxidation states. Bromine only was obtained in oxidation state +7 fairly recently in the 1960's.
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If As(V) is so oxidising why can arsenite reduce iodoform to DIM?
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Bedlasky
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H6TeO6 have oxidizing properities.
There is one more way how to we find out if some compound have oxidation or reduction properties. Let see standard electrode potentials.
For acidic environment:
H2SO4 + 2H+ + 2e- <--> H2SO3 + H2O E° = -0,93 V
H2SeO4 + 2H+ + 2e- <--> H2SeO3 + H2O E° = 1,151 V
H6TeO6 + 2H+ + 2e- <--> H2TeO3 + 3H2O E° = 1,02 V
For basic environment:
(SO4)2- + H2O + 2e- <--> (SO3)2- + 2OH- E° = -0,93 V
(SeO4)2- + H2O + 2e- <--> (SeO3)2- + 2OH- E° = 0,05 V
For (TeO6)6- I didn't found data.
In acidic environment we see that sulfates are terrible oxidizers (but sulfites are very good reductors), but selenates and tellurates are very good
oxidizers (selenates are the strongest from this three).
In basic environment we see that sulfates are terrible oxidizers again and selenates aren't really good oxidizers.
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woelen
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My experience with H6TeO6 tells otherwise. Having a high redox potential is not always a guarantee for good oxidizer properties. Sometimes strong
oxidizers on paper are very sluggish in practice. There must be a mechanistic pathway which makes them reactive and fast.
@Fusso: I don't know the answer to your question. It may have to do with coordination of certain atoms to the ion, making it easier to react. As2O5,
when added to HBr forms bromine. On the other hand, As2O3 indeed can be used as reductor in some cases. Oxidizers can sometimes surprise you. H2O2 is
a strong oxidizer, but it also can act as a decent reductor in some very specific cases. It can e.g. oxidize ferrocyanide to ferricyanide, but at
different pH it can be used to reduce ferricyanide to ferrocyanide.
All in all, my conclusion, based on observations, is that redox-reactions are more complex and more unpredicatable than one would expect, solely based
on redox potentials. Redox potentials are a nice tool, but the environment in which they are valid is limited. Having some other seemingly unrelated
compounds around can drastically change the redox properties of a compound.
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woelen
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Thread Split
9-5-2019 at 23:41 |
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