For Cl, and O, there are 7, and 6 valence electrons respectively associated with the neutral atoms.
For hypochlorite ion, Cl-O^-, we have to distribute 7+6+1 electrons in the Lewis structure. There are thus 7 electron pairs. One of these electron pairs is conceived to form the Cl-O bond, and so around each chlorine and each oxygen atom there are 3 lone pairs of electrons. Because the bonding pair of electron is shared, i.e. one electron is claimed by Cl, and one by O, this means that the chlorine atom owns 7 valence electrons, and is thus formally neutral, and the oxygen atom also owns 7 valence electrons, and thus has a FORMAL negative charge.
That is oxygen, Z=8, has 7 valence electrons, and 2 inner core electrons, and thus 9 electrons in total. Given this electronic formalism, the oxygen centre is formally negative, and our Lewis structure certainly represents this.
And now for chlorate, ClO_3^-. We have 7+6+6+6+1 valence electrons, 26 electrons, and 13 electron pairs to distribute. A Lewis stucture of (O=)_2ddotCl(-O^-) is reasonable, and I think, correctly accounts for the charge. Chlorine is neutral, and the singly bound oxygen has a negative charge. Of course, this charge is distributed to the other oxygen centres by resonance.
For completeness, we should consider perchlorate, ClO_4^-, where chlorine has its max Group 7 oxidation number of +VII, Here, we have 7+4xx6+1=32 valence electrons, and a Lewis structure of (""^(-)O-)Cl(=O)_3, again with the charge FORMALLY residing on an oxygen atom....which we CONCEIVE to be singly bound to chlorine...
