Covalent bonds formed by combining atomic orbitals and are regions of space where electrons are most likely to be formed with the calculation of bond order = (number of electrons involved in bonding - number of electrons involved in antibonding) / 2
For MOs with bonding, energy level is lower than initial energy and for MOs with antibonding (eg lone pairs), energy is higher than initial energy.
Bond order of 0 means molecule is unstable and the bond order > 1 represents the number of covalent bonds formed eg 2 = double bond
For depiction on diagrams, consider relative electronegativities of the atoms involved in bonding. Atoms with higher electronegativities would have MOs (bonding) lying closer in energy to the AOs (electrons spend more time with more electronegative atoms), and confers a partial ionic charge (dipole) on the molecule. For ionic compounds, the energy levels of the respective AO are too wide apart, hence not feasible for MO formation, thus exist as ions, instead of a covalent molecule. In addition, the filling of the s and p electrons are drawn from the lowest MO (can be sigma or pi orbital, depending on Ar with atoms having Ar lower than O having the pi orbitals at the lowest energy [more stable & lesser dispersion forces] ) and continued using Hund's rule and Aufbau principle until all the orbitals are fully filled