A single bond with all remaining electrons gives an unfavorable electron arrangement.
Form an O=O double bond.
Move one lone pair into the bond.
Leave two lone pairs on each oxygen.
All 12 electrons are used and both formal charges are zero.
Electron check
The double bond contains 4 electrons and the four lone pairs contain 8, totaling 12.
Model limitation
The simple Lewis structure is useful for electron counting, but it does not explain oxygen's observed paramagnetism. Molecular orbital theory is needed for that behavior.
Common mistakes
Drawing only a single bond.
Showing three lone pairs on each oxygen after adding a double bond.
Using the Lewis diagram as proof that O2 is diamagnetic.