Consider the redox reaction between hydrogen peroxide and chlorine dioxide to fo
ID: 530275 • Letter: C
Question
Consider the redox reaction between hydrogen peroxide and chlorine dioxide to form chlorite and oxygen: H_2 O_2 (aq) + ClO_2 (g) rightarrow ClO_2^- (aq) + O_2 (g) Assign oxidation states to each element in this reaction. Which element has been oxidized? O_2 Which element has been reduced? Cl Write the oxidation and reduction half-cell reactions. Oxidation: H_2 O_2 (aq) rightarrow O_2 (g) + H_2 + 2e Reduction: C|O_2 (g) + e rightarrow ClO_2^- (aq) Use your answers in #2 to write a complete and balanced equation for this reaction. If this reaction were performed in an electrochemical cell, calculate the standard cell potential. Is this reaction spontaneous or nonspontaneous?Explanation / Answer
Ans 1.
The oxidation states of elements in reactants are
H - +2
O - -2
Cl - +4
O- -2
In products
Cl - +3
O - -2
O2 - 0
So as you may see the oxidation state of oxygen has been increased from -2 in reactants to 0 in products hence it is oxidised.
The oxidation state of chlorine is reduced from +4 in reactants to +3 in products hence it has been reduced.