Chapter 11 Theories of Covalent Bonding

Flashcards based on lecture notes about covalent bonding theories, valence bond theory, and molecular orbital theory.

Valence Bond (VB) Theory

A quantum mechanical theory that explains chemical bonding. It posits that when atoms approach each other, their orbitals overlap, and if this overlap is energetically favorable, a chemical bond forms.

Bond Length

The distance between two atoms where the repulsive and attractive forces are balanced, resulting in minimal potential energy. This separation helps determine bond strengths.

Hybrid Orbitals

New atomic orbitals formed from the 'mixing' of two or more nonequivalent atomic orbitals of the same atom. These orbitals account for molecular geometries predicted by the VSEPR model.

sp3 Hybrid Orbitals

Hybrid orbitals formed from one 2s orbital and three 2p orbitals, resulting in four identical hybrid orbitals oriented at angles of 109.5° with respect to each other, as seen in the carbon atom in CH4.

sp Hybrid Atomic Orbitals

Hybrid orbitals formed by the mixing of one s orbital and one p orbital, resulting in two hybrid orbitals. Occurs in molecules of type AX2 with linear geometry.

sp2 Hybrid Atomic Orbitals

Hybrid orbitals formed from one s orbital and two p orbitals. Occurs in molecules of type AX3 or AX2E, with a trigonal planar electron-group arrangement.

sp3d Hybrid Atomic Orbitals

Hybrid orbitals formed from the mixing of an s orbital, three p orbitals, and one d orbital, resulting in five hybrid orbitals with a trigonal bipyramidal arrangement.

sp3d2 Hybrid Atomic Orbitals

Hybrid orbitals formed from mixing one s orbital, three p orbitals, and two d orbitals, and which will have an octahedral arrangement.

Sigma Bond (σ bond)

A bond formed by the end-to-end overlap of orbitals, involving the sharing of two electrons. Example: formed when an sp2 hybrid orbital from each carbon atom overlaps with an sp2 hybrid orbital from the other carbon atom.

Pi Bond (π bond)

A bond formed by the side-to-side overlap of orbitals. The overlap occurs above and below the plane defined by the rest of the molecule. Example: unhybridized p orbital of each carbon overlapping with the unhybridized p orbital of the other carbon.

Delocalized Electrons

Electrons that are shared by three or more atoms and are not restricted to the region of just two atoms. This sharing leads to bonding between single and double bonds.

Molecular Orbitals (MO)

Orbitals associated with the entire molecule rather than with a specific atom. A molecule is considered a collection of nuclei with associated molecular orbitals.

Bonding Molecular Orbital

A molecular orbital that has a lower energy than either of the atomic orbitals from which it formed. The lower energy signifies that the molecule is in a more stable state than having separate atoms.

Antibonding Molecular Orbital

A molecular orbital that has a higher energy than either of the atomic orbitals from which it formed. Occupancy of this orbital reduces the stability of the molecule.

Sigma (σ) Molecular Orbital

A molecular orbital in which the electron density is concentrated directly along the bond axis (a line through the two nuclei).

Pi (π) Molecular Orbitals

Molecular orbitals formed when two p atomic orbitals combine in a side-to-side fashion; the electron density is concentrated above and below the bond axis.

Bond Order (BO)

A number expressing the relative stability of a bond by comparing the number of electrons occupying bonding orbitals to the number of electrons in antibonding orbitals: BO = 1/2 (# bonding e- - # antibonding e-).