As two atoms approach each other, their electron orbitals begin to overlap. This overlap forms a molecular bond between the two atoms with its own molecular orbital shape. These orbitals follow the Pauli exclusion principle in the same way as atomic orbitals. No two electrons in an orbital can have the same quantum state. If the original atoms contain electrons where a bond would violate the rules, the electron will populate the higher energy antibonding orbital. Antibonding orbitals are denoted by an asterisk symbol next to the associated type of molecular orbital. σ* is the antibonding orbital associated with sigma orbitals and π* orbitals are antibonding pi orbitals. When speaking of these orbitals, the word 'star' is often added to the end of the orbital name: σ* = sigma-star.
Hydrogen atoms have a single 1s electron. The 1s orbital has room for 2 electrons, a spin "up" electron and a spin "down" electron. If a hydrogen atom contains an extra electron, forming a H- ion, the 1s orbital is filled.
If a H atom and H- ion approach each other, a sigma bond will form between the two atoms. Each atom will contribute an electron to the bond filling the lower energy σ bond. The extra electron will fill a higher energy state to avoid interacting with the other two electrons. This higher energy orbital is called the antibonding orbital. In this case, the orbital is a σ* antibonding orbital.
See the picture for the energy profile of the bonds formed between the H and H- atoms.