AX2E2 indicates two bonding pairs and two lone pairs, resulting in a bent shape.

F2 has a bond order of 1, while C2 has a bond order of 2.

He2 has a bond order of 0, as it has equal bonding and antibonding electrons.

The molecular orbital diagram of F2 is similar to that of O2, with the same energy level arrangement.

O2 and F2 both have a bond order of 2.

The nitrogen atom in NH3 is sp3 hybridized, forming three bonds and one lone pair.

The nitrogen atom in NH3 is sp3 hybridized, forming three N-H bonds and one lone pair.

In O2, the π2p orbitals are the highest energy orbitals.

The energy levels of molecular orbitals can vary depending on the specific atoms and their electronegativities.

In N2, the 2s orbitals (σ2s and σ*2s) are filled before the 2p orbitals.

π molecular orbitals can accommodate a maximum of 2 electrons, similar to all molecular orbitals.

The 2s molecular orbitals are lower in energy than the 2p molecular orbitals.

The σ2s orbital is lower in energy than the π2p and σ2p orbitals in homonuclear diatomic molecules.

The highest energy orbital in O2 is σ*2p, which is an antibonding orbital.

In O2, the highest occupied molecular orbital (HOMO) is π2p.

s and s orbitals can combine to form a sigma bond.

Sigma bonds are formed by the head-on overlap of s and p orbitals.

In PCl5, phosphorus exhibits dsp3 hybridization, forming five P-Cl bonds.

In C2H2, the carbon atoms are sp hybridized, forming a triple bond between them.

In PCl5, the phosphorus atom is sp3d hybridized, allowing it to form five P-Cl bonds.

H2 has a bond order of 1, indicating a single bond between the two hydrogen atoms.

O2 is paramagnetic due to the presence of two unpaired electrons in its molecular orbitals.

O2 has two unpaired electrons in its π* orbitals, making it paramagnetic.

O2 is paramagnetic due to the presence of two unpaired electrons in its molecular orbitals.

O2 is paramagnetic due to the presence of two unpaired electrons in its molecular orbitals.

N2 has a triple bond, with a bond order of 3 according to molecular orbital theory.

O2 has unpaired electrons in its molecular orbital configuration, making it paramagnetic.

In methane, the bond angle is approximately 109.5 degrees due to its tetrahedral geometry.

In a tetrahedral geometry, the bond angles are approximately 109.5 degrees.

In a trigonal planar geometry, the bond angles are 120 degrees.