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fusso
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CrO5 isn't an acid.
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teodor
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I believe that happens with soluble chromates only if the solution is alkaline. In acidic media H2O2 reduces chromium (VI) to chromium (III). In
neutral solution H2O2 doesn't react with BaCrO4 at all (and I think the same should be true for BaCrO4 in alkaline media because this salt is
insoluble).
My BaCrO4 samples atr lemon-yellow, exactly the same colour you use in the font to write the word "Barium".
Copper (II) and Iron (II ?) chromates are quite different, they are brown. But their solution in water has the same yellow colour as other chromates.
[Edited on 9-7-2019 by teodor]
[Edited on 9-7-2019 by teodor]
[Edited on 9-7-2019 by teodor]
[Edited on 9-7-2019 by teodor]
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teodor
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Also I think complex formation, as with HCl, HNO3 prevents oxidation Cr(III) to Cr(VI) when the solution changes pH to alkaline, but it is not the
case when I used H3PO4 - when I tried to add KOH to the solution to precipitate Cr(OH)3 it was oxidized back to chromate. But not sure, should try
this again.
(To put it in other words, an insoluble compound forms before oxidation - possible there is some performance from that hydroxo...quarions, which could
participate in electron transfers - according to example between Fe(III) and Cr(II) https://www.britannica.com/science/oxidation-reduction-react...)
[Edited on 9-7-2019 by teodor]
[Edited on 9-7-2019 by teodor]
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teodor
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Once I've got a chromate by oxidizing Cr3+ with Ca(OCl)2. Oxidation of HCl by Cr6+ possible will give HOCl again. So, I think you have to allow to
escape some Cl2 and that will drive the equilibrium to Cr3+ side.
Or possible I am wrong, there is no such thing as equilibrium in the oxidation-reduction, but in that case I can't explain that.
[Edited on 9-7-2019 by teodor]
[Edited on 9-7-2019 by teodor]
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teodor
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Quote: Originally posted by teodor  |
I believe that happens with soluble chromates only if the solution is alkaline. In acidic media H2O2 reduces chromium (VI) to chromium (III).
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There is something to add, I red that in Treadwell , Analytical Chemistry, Vol. 1, see under "Chromates":
"In acid solutions the chromic cation is the most stable condition for chromium, but in alkaline solutions the chromate anion is the more stable
condition. In acid solutions, therefore, it is easy to reduce a chromate to chromic salt and in alkaline solutions it is easy to oxidize a chromic
salt to chromate".
(In 4th english edition it is page 137).
Then
"Oxidation in acid solution may be effected by boiling with very energetic oxidizing agents such as conc. nitric acid and potassium chlorate, sodium
bismuthate (or bismuth tetroxide), or potassium permanganate. ... the chloride should not be used, or any other salt of which the anion is capable of
oxidation ...
So, you are right, H2O2 usually works in the opposite direction - oxidizer in acidic solution and reducer in alkaline, but not in the case when Cr
ions are present, their presence changes this.
Also, Treadwell describes the formation of perchromic acid H3CrO8 with help of H2O2 in _alkaline solution_. In neutral solution it will be H3CrO7. And
also in some circumstances H7CrO10.
[Edited on 9-7-2019 by teodor]
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Bedlasky
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| Quote: Originally posted by teodor |
I believe that happens with soluble chromates only if the solution is alkaline. In acidic media H2O2 reduces chromium (VI) to chromium (III). In
neutral solution H2O2 doesn't react with BaCrO4 at all (and I think the same should be true for BaCrO4 in alkaline media because this salt is
insoluble). |
Dichromate reacts with hydrogen peroxide in acidic media to form unstable oxide-peroxide CrO5. In aqueous solution CrO5 quickly decomposes in to
oxygen and Cr3+. But in organic solvent like diethyl ether, ethyl acetate or amyl alcohol it is much more stable. See this video.
| Quote: Originally posted by teodor | | Also, Treadwell describes the formation of perchromic acid H3CrO8 with help of H2O2 in _alkaline solution_. In neutral solution it will be H3CrO7. And
also in some circumstances H7CrO10. |
In alkaline media it doesn't exist peroxochromic acid but peroxochromate(V) - CrO8(3-) (better forumla is [Cr(O2)4]3-).
Woelen have on his site few articles about chromium-peroxide complexes in acidic and also in alkaline media:
CrO5
K3CrO8 synthesis
K3CrO8 properties
[Cr(NH3)3(O2)2]
| Quote: Originally posted by teodor | Once I've got a chromate by oxidizing Cr3+ with Ca(OCl)2. Oxidation of HCl by Cr6+ possible will give HOCl again. So, I think you have to allow to
escape some Cl2 and that will drive the equilibrium to Cr3+ side.
Or possible I am wrong, there is no such thing as equilibrium in the oxidation-reduction, but in that case I can't explain that.
|
Commercial Ca(OCl)2 contains some CaCl2 and Ca(OH)2 so its solutions are basic.
Cr(3+) salts are oxidized by hypochlorite or chlorine in alkaline media.
If dichromate oxidize in acidic solution HCl, chlorine is released and it doesn't reoxidize Cr3+.
| Quote: Originally posted by teodor | | "Oxidation in acid solution may be effected by boiling with very energetic oxidizing agents such as conc. nitric acid and potassium chlorate, sodium
bismuthate (or bismuth tetroxide), or potassium permanganate. ... the chloride should not be used, or any other salt of which the anion is capable of
oxidation ...[Edited on 9-7-2019 by teodor] |
I don't confirm this but I doubt that permanganate oxidize Cr3+ in to dichromate in acidic media. But maybe I am wrong. Often are Cr3+ salts oxidized
in alkaline conditions.
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