How is the oxidation of #"ferrous ion"# to #"ferric ion"# by #"permanganate ion"# represented?

1 Answer
May 16, 2017

Answer:

We write out the individual redox processes to get:

#MnO_4^(-) + 5Fe^(2+) +8H^(+)rarr Mn^(2+) + 5Fe^(3+) +4H_2O(l)#

Explanation:

#"Oxidation reaction (i):"#

#Fe^(2+) rarr Fe^(3+) + e^-#

#"Reduction reaction (ii):"#

#MnO_4^(-) + 8H^(+) + 5e^(-) rarr Mn^(2+) + 4H_2O(l)#

For both #(i)# and #(ii)#, #"charge is balanced"# and #"mass is balanced"#, as indeed they must be if we purport to represent chemical reality. The final redox reaction adds #5xx(i)+(ii)# so that electrons, conceptual particles, do not appear in the final redox equation:

#MnO_4^(-) + 5Fe^(2+) +8H^(+)rarr Mn^(2+) + 5Fe^(3+) +4H_2O(l)#

Which is balanced (is it?) with respect to mass and charge. The reaction has a built in indicator in that #MnO_4^-# in intensely purple, whereas the reduction product #Mn^(2+)# is almost colourless. An endpoint could be vizualized.