Upsilon - 15-9-2015 at 18:35
Is there a such thing as ferrous nitrate (iron(ii) nitrate)? I can't seem to find any information on it anywhere. If not, then can someone explain why
it can't exist?
Texium - 15-9-2015 at 18:45
Iron(II) nitrate is not stable because the iron(II) ion is easily oxidized by the nitrate ion to iron(III). You will see that most of the time,
oxidizing anions like nitrate are only stable when paired with higher oxidation state cations.
Upsilon - 15-9-2015 at 21:12
Ok. What about manganese(ii) oxalate? Can't seem to find information on it anywhere. I was curious as to what would happen if I flooded manganese
dioxide with a solution of oxalic acid (since manganese dioxide oxidizes the chlorine in hydrochloric acid, surely it will do something similar to an
even weaker acid).
woelen - 15-9-2015 at 22:43
In principle manganese(II) oxalate can be made and it also will be stable. Many manganese(II) salts of weak acids, however, are very easily oxidized
by oxygen from the air, especially if the air is somewhat humid, or when the compound is not perfectly dry. You can see that yourself. Prepare a
solution of MnSO4 or MnCl2 and add a solution of a salt of a weak acid, which forms an insoluble salt of manganese(II) (e.g. hydroxide, carbonate,
phosphate, oxalate, borate, all of them will do). You at first get a pale pink/off white precipitate, but where the off-white material is in contact
with air (near the surface, material sticking to the glass) it soon turns brown, due to formation of basic salts of manganese(III) or even
manganese(IV). The manganese(II) is oxidized by oxygen from the air.
Manganese dioxide, flooded with a solution of oxalic acid will hardly react. The reaction is very slow. Maybe, if you allow the powder to sit in the
solution for many days and if you shake every now and then, it might dissolve with formation of manganese(II), but even then I expect only partial
reaction at best. If a lot of extra strong acid is added (e.g. H2SO4), then the MnO2 may dissolve somewhat faster.
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In principle, ferrous nitrate could exist. Nitrate ion, especially when no acid is available, can perfectly coexist with reducing ions at room
temperature. Only at elevated temperatures or at high concentrations of acid, the nitrate ion becomes oxidizing.
Making ferrous nitrate could be done by mixing a solution of lead nitrate and ferrous sulfate in the right proportions. The resulting insoluble lead
sulfate must be filtered and the remaining liquid must be allowed to evaporate.
In practice, this preparation will not be easy. It must be done with rigorous exclusion of air (otherwise you will have oxidation of iron(II) to
iron(III) and formation of basic iron(III) nitrate) and probably you also need a very strong drying agent. Most likely, ferrous nitrate will be very
hygroscopic.
[Edited on 16-9-15 by woelen]
ave369 - 16-9-2015 at 10:28
Reducing cation + oxidizing anion is generally not the best combination for a salt. That's the reason why ammonium ferrate does not exist: the anion
attacks the cation.
DraconicAcid - 16-9-2015 at 10:37
Nitrate ion is only strongly oxidizing in really acidic solution- there shouldn't really be any trouble in making iron(II) nitrate, apart from the
risk of oxidizing the iron(II) by exposure to air.
deltaH - 16-9-2015 at 10:51
Yes it exists and can be prepared by reacting iron filings with ferric nitrate and dilute nitric acid.
See attached ORNL document for details.
Attachment: THE USE OF FERROUS NITRATE AS A PLUTONIUM REDUCTANT.pdf (344kB)
This file has been downloaded 539 times
[Edited on 16-9-2015 by deltaH]
chornedsnorkack - 16-9-2015 at 12:58
Tin(II)nitrate is described as "reasonably stable" in dilute aqueous solutions, but always decays on concentration, so the solid nitrate is unknown.