I think in this question there is a need to provide another
reagent, it is impossible to tell wheter A is a oxidising agent
where there is only A involved
But in my opinion, since a oxidising agent is one that get
reduced, Fe2+ had a better chance of being the oxidising agent as
the oxidation state is +2 as compared to -2 in (SO42-).
However, can add sutible reagent to Iron (II) Sulphate and
isolate one product to the other, and test using aqueous potassium
dichromate and potassium ioidide, then, if one of the product,
perfrably the one with the Iron (II) ion is reduced and the other
is unchanged, Iron (II) is the reducing agent
correct me if i am wrong
>>> Fe2+ had a better chance of being the oxidising
agent as the oxidation state is +2 as compared to -2 in (SO42-).
<<<
The IONIC CHARGE of the dinegative sulfate(VI) ion is
2-. It's incorrect (and a common student error) to say "the O.S. of
a dinegative ion is -2". The O.S. must always refer to the
O.S. of an ELEMENT (or a specified atom of an element, if the
Kekule structure is drawn) within a specified compound.
Eg. O.S. of SULFUR in the following compounds are (to properly
understand this business of O.S. and it's relationship to formal
charges and electronegativity, draw the Kekule structures) :
sulfur dioxide hydrolyzed to sulfate(IV) ion SO3 2-
----> +4
sulfur trioxide hydrolyzed to sulfate(VI) ion SO4 2-
----> +6
thiosulfate ion S2O3 2- ----> +2 (average of), +4
and 0.
tetrathionate ion S4O6 2- ----> +2.5 (average of),
+5, 0, 0, +5.
peroxydisulfate(VI) ion S2O8 2- ----> +6
(6 oxygen atoms have O.S. of -2 each, the 2 oxygen atoms of the
peroxy group have O.S. of -1 each).