Because there is less electron density between the two nuclei, antibonding orbitals have a greater energy. To draw an electron away from a nucleus, energy is required. As a result, electrons in an antibonding orbit have a higher energy level when they spend less time between the two nuclei. Bonds are formed when two atoms join together. The resulting pair of electrons has lower energy than those in individual atoms because it isn't necessary for them to be close to each other. Thus, bonds form when two electrons in different atoms come together and share their negative charge.
Bonds can be thought of as pairs of electrons that want to be far apart from each other. Because electrons are orbiting around nuclei, they will try to avoid being too close to them. Therefore, electrons in bonding orbits are closer together than they would be if there were no bond present. In contrast, electrons in anti-bonding orbits are more separated than they would be if there were no bond present.
The strength of a chemical bond is measured by its binding energy. This means that the stronger the bond, the higher its energy will be. For example, consider hydrogen gas. It consists of one single particle in an orbital with zero energy because there is no potential energy due to the nucleus-electron attraction. Since this orbital is not filled, the bond must be considered non-bonded.
Antibonding orbitals occur as a result of out-of-phase orbital overlap, which produces disruptive interference. Because atomic orbitals are conserved, they frequently develop alongside bonding orbitals. For example, 2s and 2p electrons form pairs with opposite spins in atoms like helium, neon, and argon. These pairs of electrons constitute the 1s and 3d subshells, which are filled before any other subshell is full.
Bonding and antibonding combinations of orbitals can be thought of as "quasi-levels" of energy that alternate in sign each time an atom bonds or refuses to bond. In general, bonding interactions increase the stability of molecules and materials while antibonding interactions do not necessarily destabilize molecules but they may cause problems for certain processes such as chemical reactions or measurements. Bonding and antibonding interactions can also influence the properties of solids such as conductivity. Electrical conductors such as gold have only spacial bonding between their electrons while insulators such as quartz have only spacial antibonding between their electrons.
In general, bonding interactions are more stable than antibonding interactions. However, antibonding interactions can become more stable if one considers the total energy of the system instead of just the energies of the individual particles.
The antibonding orbital has a greater energy than both the bonding and hydrogen 1s orbitals due to the reduction in electron density between the nuclei. Bonding molecular orbitals, in general, contain less energy than either of their parent atomic orbitals. The 1s orbital of carbon is an example of this phenomenon. The 1s orbital can be thought of as made up of two half-orbitals that overlap with each other at the nucleus where they combine to form a single bond with three electrons. Thus, it contains less energy than its atomic counterpart which has only one electron in its 1s orbital.
An antibonding orbital will always have more energy than a bonding orbital because there is less overlap between the nuclei and their valence shells in the antibonding case than in the bonding case. If we think about the 1s orbital of carbon again, it can be said to have less overlap between the nuclei and their valence shell than does the 2s orbital since there are two electrons in the 1s orbital but only one electron in the 2s orbital. The 1s orbital therefore has more energy than the 2s orbital.
This difference in energy is what allows us to identify which orbitals are antibonding and which are not. If an orbital has less energy than either of its parent atomic orbitals then it is antibonding.
Antibonding orbitals occur as a result of out-of-phase orbital overlap, which is referred to as destructive interference. This restricts the locations in which electrons may exist, increasing electron repulsion and making antibonding orbitals more energetic than the comparable bonding orbital. The energy of an antibonding orbital is not fixed: it may be higher or lower than that of the corresponding bonding orbital.
Because antibonding orbitals are less stable than bonding orbitals, molecules often prefer to increase stability by forming bonds with other atoms or molecules. For example, carbon monoxide has an antibonding orbital that is significantly below the energy of the corresponding bonding orbital; thus, it is strongly bound to iron (which has a major contribution from p-orbitals), but only weakly bound to oxygen (whose p-orbitals do not overlap with those of CO). In contrast, both bonding and anti-bonding orbitals of hydrogen exist at similar energies, so it can easily share its electrons with other elements; for example, water consists of two hydrogens bonded to one oxygen atom. However, antibonding orbitals of certain elements may become occupied due to the presence of another element with which they cannot form a bond; for example, the antibonding orbital of nitrogen is filled with 2p electrons because there are no other options available for occupation due to spin-orbital coupling.
Antibonding orbitals play an important role in many chemical reactions.
Bonding molecular orbitals contain less energy and hence more stability than antibonding molecular orbitals. This is because the electrons in bonding orbits are closer together than those in anti-bonding orbits, so they can attract each other's positive charges (the nuclei) more strongly.
Antibonding and bonding molecular orbitals can be either symmetric or asymmetric. Symmetric antibonding and bonding molecular orbitals will have equal numbers of electrons in them. Asymmetric antibonding and bonding molecular orbitals will have different numbers of electrons in them. The difference in number between the two types of molecules is called the oxidation state. Oxidation states indicate the degree to which elements are oxidized; higher values mean more oxidation. For example, carbon in its pure form is in a zero oxidation state. Because it has no electrons lost nor gained, its electron configuration is 2s^2 2p^2. In oxygen gas at room temperature, there are four electrons in the 1s orbital, two electrons in the 2s orbital, and one electron in each of the 2p orbitals. Thus, even though oxygen has a greater atomic mass than carbon, it is considered an electropositive element because its valence shell is filled with electrons.