Oxidation related question

[Mathman23]

Hi All,

\mathrm{Fe^{3+}(aq) + SCN^{-} (aq) \rightleftharpoons FeSCN^{2+}(aq)}

I add some \mathrm{NaHSO_{3}(s)} to the above reaction !

This makes the above chemical solution sour, consequently in a sour solution \mathrm{HSO^{-}_{3}} will reduce Fe^{3+} to Fe^{2+}, meanwhile HSO_{3}^{-} is itself oxidated to SO^{2-}_{4}.

a/ The color of the original resolution changes then adding \mathrm{NaHSO_{3}(s)} to yellow.

b/ By adding \mathrm{NaHSO_{3}(s)} to the orginal solution the equilibrium direction changes, such that it runs from right to left.

c/ I need help writing the reaction between HSO^{-}_{3} and SO^{-}_{4}.

d/ Which influence does the adding of HSO^{-}_{3} have on the concentration of Fe^{3+}

I hope that there someone out there who is willing to assist me since this is my last post !

Sincerely
Fred

[chem_tr]

Hello,

Let me first do the redox thing in atomic level, since it's easier than dealing with a bunch of atoms.

We know that iron(III) is reduced to iron(II), and in turn, sulfite (S4+) is oxidized to sulfate (S6+):

S^{4+} \longrightarrow S^{6+}+2e^-
2Fe^{3+}+2e^-\longrightarrow 2Fe^{2+}
--------------------------------
S^{4+}+2Fe^{3+} \longrightarrow S^{6+}+2Fe^{2+}

In the real redox, you'll need to put some other correcting factors (water and hydroxide, I think) to balance additional oxygens.

Now that we've got rid of the redox reaction, we may look other things. When you add some sodium bisulfate to the medium, a redox reaction occurs; and so the thiocyanate complex changes from \displaystyle Fe(SCN)^{2+} to \displaystyle Fe(SCN)^+; this gives a totally different color. Iron(II) sulfate is also formed from the reaction, and the color may have also come from this one, so I don't think that the equilibrium shifts to the left, but I may be wrong; I don't insist on that.