Write an equation for the disproportionation of chlorate(I) ions and show the oxidatiWatch
chlorate(i) ions and show the
oxidation states of chlorine in the products. Thanks
The easiest way to write any disproportionation is to write single electron reactions for the reduction and the oxidation and then add the two equations to each other.
Assuming it disproportionates to chloride ions and chlorate ions, you have the following half equations.
Reduction of ClO(-) into Cl(-)
0.5 ClO(-) + H(+) + e(-) --> 0.5 Cl(-) +0.5H2O
Oxidation of ClO(-) into ClO3(-)
0.25ClO(-) + 0.5 H2O --> 0.25ClO3(-) + H(+) + e(-)
Now add the two equations, and everything simplifies nicely with the hydrogen ions, electrons and waters cancelling out.
0.75 ClO(-) --> 0.5 Cl(-) + 0.25 ClO3(-)
This can also be written with integer coefficients of 3, 2 and 1 respectively.
EDIT: Just saw you asked for oxidation states too. ClO(-) has chlorine in oxidation state +1, Cl(-) has chlorine in oxidation state -1, ClO3(-) has chlorine in oxidation state +5. You can work this out as the oxidation state of oxygen in all non-peroxide compounds is (-2), and the sum of...
oxidation state of chlorine + ( (-2) * number of oxygen atoms) = overall charge on the ion.
This is why the reaction is a disproportionation, because 0.75 moles of chlorine at oxidation state 1 is simultaneously reduced to 0.5 moles of chloride ions at oxidation state -1 and oxidised to 0.25 moles of chlorate (V) ions at oxidation state +5.
Hope this helped!
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