Hybridization

Single bonds are

always sigma bonds.

This process of hybridization gives the same number of

hybrid orbitals as the number of atomic orbitals involved—but these orbitals have the same energy (degenerate), are symmetrically arranged around the atom, and are more directional. We say that the atom undergoes sp3 hybridization or is sp3 hybridized.

Double bonds always consist of

one sigma bond and one pi bond.

You cannot have a

stand-alone pi bond

When atoms share more than one pair of electrons,

the additional pairs are in pi bonds. The centers of charge density in a pi bond lie above and below the bond axis.

Whenever a set of equivalent tetrahedral orbitals is "required" by an atom because it has 4 pairs of electrons that

"need a home" according to VSEPR, the atom adopts a set of sp3 orbitals to achieve minimum energy (ie, a more stable arrangement for the molecule). In other words... if VSEPR indicates that an atom "needs" a tetradral electron group geometry, it adopts sp3 hybridization.

sp hybridization

The mixing of the s orbital with a p orbital yields 2 molecular orbitals with equal energy capable of forming a pair of sigma bonding orbitals oriented at 180° angles (providing linear orbital geometry). The leftover unhybridized p orbitals are perpendicular to the bonding axis of the sp hybrid.

sp2 hybridization

The mixing of the s orbital with two p orbitals yields 3 orbitals with equal energy capable of forming a set of sigma bonding orbitals oriented in a plane at 120° angles (providing trigonal planar orbital geometry).

When atoms join together to form molecules

(except H), their outer atomic orbitals interact with each other to produce hybrid orbitals.

Pi (л) bond

A covalent bond in which electron density is concentrated above and below the line joining the bonded atoms (the internuclear axis).

Sigma (б) bond

a covalent bond in which electron density is concentrated along the internuclear axis.

Hybridization

is the mixing of native atomic orbitals to form special degenerate orbitals for bonding.In other words, when an atom bonds, the atomic orbitals involved in forming the sigma bonds or accommodating the lone pairs of electrons, interact with each other to form an equal numberof highly directional hybrid orbitals of equal energy.

Delocalization

A pi bond results from the side-ways interaction of unhybridized p-orbitals on two atoms—each containing 1 electron. (i.e., the electrons are localized) In many molecules, we cannot adequately describe the bonding as being entirely localized. This situation arises particularly in molecules that have pi bonds and have more than one resonance structure.

sp3 hybridization

The mixing of the s orbital with all three of the p orbitals yields 4 orbitals with equal energy capable of forming 4 identical bonds oriented in a tetrahedron.

Carbon and hydrogen

form methane, a perfectly tetrahedral molecule:

Why do atoms hybridize?

hybridized orbitals are lower in energy compared to their separated (atomic orbitals) unhybridized counterparts. more stable compounds when hybridization occurs. major parts of the hybridized orbitals, or the frontal lobes, overlap better than the lobes of unhybridized orbitals. This leads to better bonding.

The electrons in sigma bonds are

localized in the region between two bonded atoms and do not make a significant contribution to the bonding between any other two atoms.

When a double covalent bond forms between atoms, the two bonds are

not identical. The first bond, a sigma bond, is formed by head-on overlap between atomic or bonding (hybrid) orbitals along the internuclear axis (an imaginary line joining the two nuclei). The electron density is concentrated directly between the nuclei, along the internuclear axis. The second bond, a pi bond, is formed by the sideways interaction (overlap) of electrons in p-orbitals at right angles to the internuclear axis. The electron density is concentrated on either side of the internuclear axis between the two nuclei.

Every pair of bonded atoms shares

one or more pairs of electrons. a.In every bond, at least one pair of electrons is localized in space between the atoms in a sigma bond. b.The appropriate set of hybrid orbitals used to form the sigma bonds in an atom and its neighbors is determined by the observed geometry of the molecule. (Use VSPR for this!)

Triple bonds always consist of

one sigma bond and two pi bonds.

The leftover unhybridized p orbital is

perpendicular to bonding plane of the sp2 hybrids.

Molecules with two or more resonance structures can have

pi bonds that extend over more than two bonded atoms. Electrons in pi bonds that extend over more than two atoms are said to be delocalized.