bonding and antibonding orbitals
When two water waves meet, they can add up into a bigger wave or cancel into a flat patch. Electron clouds, being waves, do the same. Combine two atomic orbitals 'in step' and the electron density piles up between the nuclei; combine them 'out of step' and a gap of zero density opens between them. The first kind is a bonding orbital, the second an antibonding orbital.
Precisely, when two atomic orbitals combine they always make two molecular orbitals: a bonding orbital, lower in energy, with extra electron density between the nuclei that pulls them together; and an antibonding orbital, higher in energy, with a node (a surface of zero density) between the nuclei that lets them repel. Antibonding orbitals are often marked with a star, like σ* or π*.
This pair is the heart of molecular orbital theory. A molecule is held together when bonding orbitals hold more electrons than antibonding ones — that net excess is what makes a bond. The honest caveat is that antibonding orbitals are not just empty curiosities: pushing an electron up into one (for instance by absorbing light) weakens or breaks the bond, which underlies much of color and photochemistry.
He₂ does not exist as a stable molecule. Its four electrons fill both the bonding orbital and the antibonding orbital, so the two cancel out — there is no net bonding, and the helium atoms simply drift apart. H₂, with only two electrons, fills the bonding orbital alone and holds together.
Filled antibonding orbitals cancel filled bonding ones — that is why He₂ won't form.
An antibonding orbital is real, not merely 'the absence of bonding'. It is a genuine quantum state with its own energy and shape, and electrons really can occupy it. The antibonding level is also raised in energy by more than the bonding level is lowered, so filling both is slightly destabilizing overall.