Hybrid Orbitals in Chemistry

The process of combining one s orbital and one p orbital, resulting in two equivalent hybrid orbitals oriented in a straight line (180° angle).

The process of moving an electron from a paired orbital to an empty orbital within the same energy level, as part of hybridization.

The resulting hybrid orbitals from sp hybridization, each containing equal amounts of s and p characteristics.

Beryllium atoms, in their neutral state, have all paired electrons in their outermost shell. They therefore undergo hybridization to achieve more stable bonding.

The combination of orbitals of the same quantum number is called hybridization. A hybrid orbital is an orbital that results from the hybridization process.

The angle between the two sp hybrid orbitals, which is 180°. This leads to linear molecular geometry.

In sp hybridization, the two remaining p orbitals (p_y and p_z) don't participate in the process and stay unoccupied.

The shape of sp hybrid orbitals, similar to p orbitals with one lobe being larger than the other. This facilitates stronger and more directional bonding.

In ammonia (NH3), the nitrogen atom forms three N-H bonds and has one lone pair of electrons. The four orbitals involved in bonding and the lone pair are sp3 hybridized, resulting in a tetrahedral arrangement. The lone pair exerts a slightly greater repulsive force, pushing the N-H bonds away from the top of the pyramid, resulting in slightly smaller H-N-H bond angles (107.3°) compared to a perfect tetrahedral structure (109.5°).

In water (H2O), the oxygen atom is sp3 hybridized, with two of the hybrid orbitals forming bonds with the two hydrogen atoms, and the other two orbitals containing the lone pairs. This arrangement results in a bent molecular shape with an angle of approximately 104.5°.

A type of covalent bond formed by the direct overlap (end-to-end) of atomic orbitals along the internuclear axis. This type of overlap is also called 'head-on' or 'axial' overlap.

A type of covalent bond formed by the sideways overlap of atomic orbitals above and below the internuclear axis. This overlap results in a bond with electron density concentrated above and below the plane of the sigma bond.

The sideways overlap of unhybridized p-orbitals on adjacent carbon atoms leads to the formation of pi bonds. Pi bonds are weaker than sigma bonds and contribute to the double or triple bonds in molecules.

The head-on overlap of hybridized orbitals, or an s orbital and a hybridized orbital, on adjacent atoms leads to the formation of sigma bonds. These bonds are strong and are the primary bonds in a molecule.

The molecule has a tetrahedral arrangement with four sp3 hybrid orbitals around the central atom. Each sp3 orbital forms a sigma bond with a hydrogen atom.

The nitrogen atom uses sp3 hybridization to form a trigonal pyramidal structure. Three of the sp3 orbitals form sigma bonds with hydrogen atoms, and the fourth sp3 orbital contains a lone pair of electrons.

The oxygen atom uses sp3 hybridization to form a bent structure. Two of the sp3 orbitals form sigma bonds with hydrogen atoms, and the remaining two sp3 orbitals contain lone pairs of electrons.

Boron trifluoride (BF3) has sp2 hybridization, resulting in a trigonal planar geometry. The boron atom forms three sigma bonds with fluorine atoms.

The strength of a covalent bond is determined by the amount of overlap between the atomic orbitals involved.

Sigma bonds are stronger than Pi bonds because of the greater extent of overlapping.

The formation of a Sigma bond allows free rotation around the bond axis, while Pi bonds restrict rotation.

Sigma bonds can exist alone, while Pi bonds always form in addition to a Sigma bond.

Sigma bonds are formed from the overlap of s and p atomic orbitals, while Pi bonds are formed from the overlap of only p orbitals.

The overlap of atomic orbitals forms molecular orbitals and covalent bonds, where electrons are shared between atoms.

One of the ways sigma bonds can occur in a molecule, where the s-orbital of one atom overlaps with the s-orbital of another atom, leading to the formation of an H2 molecule.

A type of sigma bond formed when the s-orbital of one atom overlaps with the half-filled p-orbital of another atom, forming molecules like HF and HCl.

Another type of sigma bond formed when the p-orbitals of two atoms overlap directly along the internuclear axis, leading to the formation of molecules such as F2, Cl2, and Br2.

The presence of a pi bond, which is stronger and shorter than a single bond, leads to a shorter bond length in ethylene compared to molecules with single carbon-carbon bonds. This is because the pi bond contributes to the overall bond strength.

The presence of a pibond means two atoms are bonded by a double bond instead of a single bond. It is also called a double bond.

The pibond allows for free rotation around the bond axis in molecules like ethane, unlike restricted rotation around the double bond in ethylene.