The table below shows the main oxidation states of the Group elements. The major features of the redox chemistry of elements in this group are:
Like oxygen, fluorine only has negative oxidation states.
The most stable state for all elements is -1.
All of the elements and the compounds containing halogen in positive oxidation states are oxidants.
Despite being the most stable state for iodine, I– is a reasonable reductant in aqueous solution.
F
Cl
Br
I
+7 (ClO4–)
+7
+7
+5 (ClO3–)
+5
+5
+3 (ClO2–)
+1 (ClO–)
+1
+1
0 (F2)
0(Cl2)
0 (Br2)
0 (I2)
–1* (F–)
–1* (Cl–)
–1* (Br–)
–1* (I–)
With the exception of fluorine, halogens and their oxoanions are susceptible to autooxidation-reduction (disproportionation).
Autooxidation-reduction (disproportionation) is a reaction in which atoms of the same element act as both the oxidant and the reductant. Therefore in order to undergo auto-oxidation-reduction reaction, an atom of an element must be in an intermediate oxidation state.
For example Cl2 (oxidation state 0) reacts with water to to form HOCl (Cl oxidation state +1) and HCl (Cl oxidation state -1).