Eddygp
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Redox reaction explained incorrectly?
So I was in one of these boring chemistry demonstrations just after the students had learned how redox reactions work. The teacher explains that he is
going to mix a solution of copper(II) sulphate with aluminium foil to yield aluminium(III) sulphate and copper metal, in the presence of chloride
ions.
Up to this point, everything seems rather normal and straightforward. However, when he mixes the substances in the solution, I notice that, while
copper is evidently being generated, gas bubbles are also formed, with that characteristic smell of SO2. I therefore realised that the
reaction must not be
2Al + 3Cu2+ ---> 2Al3+ + 3Cu
but most likely is something similar to the following one:
8H+ + 2Al + 5Cu2+ + 2SO42+ ---> 2Al3+ + 5Cu + 2SO2 + 4H2O
(notice that the 8H+ imply that this is in an acidic solution)
(also, notice that the 3 "missing" sulphate ions are spectator ions)
This would imply that the aluminium is capable of reducing sulphur(VI) to sulphur(IV)... but I find no other possible explanation for the evolution of
the sulphur dioxide gas.
Are there other common redox reactions in which the baccalaureate-level explanations and the practice have significant inconsistencies?
Edit: the formats do not seem to work :/
[Edited on 5-4-2015 by Eddygp]
there may be bugs in gfind
[ˌɛdidʒiˈpiː] IPA pronunciation for my Username |
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blogfast25
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The gas generated is almost certainly mainly hydrogen, not SO<sub>2</sub>. But hydrogen generated when metals dissolve often has a
peculiar smell, that could easily be mistaken for a sulphurous smell.
'On paper', aluminium can indeed reduce sulphate ions to sulphite ions in acid conditions (reduction potential = 0.2 V). In reality that doesn't seem
to happen, for instance Al metal with a sodium sulphate solution will not yield any reaction. In watery medium, sulphates are incredibly hard to
reduce and have no oxidising properties at all. Hence also use of sulphuric acid as a non-oxidising acid.
The reduction of cupric sulphate by Al in the absence of chloride ions has also been discussed elsewhere on this forum and is quite slow,
without any reaction of the Al with the sulphate anions.
[Edited on 5-4-2015 by blogfast25]
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Metacelsus
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Quote: Originally posted by Eddygp  |
8H+ + 2Al + 5Cu2+ + 2SO42+ ---> 2Al3+ + 5Cu + 2SO2 + 4H2O
(notice that the 8H+ imply that this is in an acidic solution)
(also, notice that the 3 "missing" sulphate ions are spectator ions)
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Sulfate is negatively charged, not positively. Thus, the second equation is balanced wrong.
The situation you proposed is equivalent to:
2 Al + 3 Cu2+ ---> 2 Al3+ + 3 Cu
and
2 Al + 12 H+ + 3 SO42- = 3 SO2 + 6 H2O + 2 Al3+
happening at the same time.
Let me find my reduction potential tables, and see if the reaction is spontaneous.
Edit: Ninja'd! Apparently the reaction is spontaneous, but the kinetics are unfavorable.
[Edited on 5-4-2015 by Cheddite Cheese]
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blogfast25
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| Quote: Originally posted by Cheddite Cheese | Let me find my reduction potential tables, and see if the reaction is spontaneous.
Edit: Ninja'd! Apparently the reaction is spontaneous, but the kinetics are unfavorable.
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'Spontaneous' is a very misleading term here. Aqueous sulphates are basically impossible to reduce, so the reduction reaction is hardly 'spontaneous'.
If the cell potential of a redox reaction Ecell = Eox + Ered > 0 that means that ΔGreaction < 0 and the reaction is thermodynamically
favourable. But thermodynamics makes no pronouncements about kinetics. Here, the reduction of aqueous sulphates doesn't happen, depite the reaction
being thermodynamically favourable.
Al can reduce sulphates in non-aqueous conditions, for instance stoichiometric mixtures of calcium sulphate and aluminium powder burn very hot,
according:
CaSO4 + 8/3 Al === > CaS + 4/3 Al2O3 (S reduced from +6 to -2)
But the enthalpies and Free Energy involved are very different in that case.
[Edited on 5-4-2015 by blogfast25]
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