A cylindrically symmetrical bond with no nodal planes. All single covalent bonds are σ-bonds,
Single bonds consist of one σ-bond.
Double bonds consist of one σ-bond and one π-bond.
Triple bonds consist of one σ-bond and two π-bonds.
3.4 Sigma and Pi Bonds Summary
In valence-bond theory, we assume that bonds form when unpaired electrons in valence-shell atomic orbitals pair; the atomic orbitals overlap end to end to form σ-bonds or side by side to form π-bonds.
3.5 Electron Promotion and the Hybridization of Orbitals Summary
The promotion of electrons will occur if, overall, it leads to a lowering of energy by permitting the formation of more bonds. Hybrid orbitals are constructed on an atom to reproduce the electron arrangement characteristic of the experimentally determined shape of a molecule.
Hybrid orbitals formed from one s-orbital and three p-orbitals. 4 atomic orbitals are involved, and 4 hybrid orbitals are produced.
sp³ hybridized atoms.
Whenever an atom in a molecule has a tetrahedral electron arrangement, it is said to be sp³ hybridized.
Hybrid orbitals formed from one s-orbital and two p-orbitals. 3 atomic orbitals are involved, and 3 hybrid orbitals are produced.
sp² hybridized atoms.
Whenever an atom in a molecule has a square planar arrangement, it is said to be sp² hybridized.
sp hybrid orbitals
Hybrid orbitals formed from one s-orbital and one p-orbital. 2 atomic orbitals are involved, and 2 hybrid orbitals are produced. The resulting shape is linear.
sp³d hybrid orbitals
5 hybrid orbitals formed from 5 atomic orbitals to form a trigonal bipyramidal shape.
3.6 Other Common Types of Hybridization Summary
A hybridization scheme is adopted to match the electron arrangement of the molecule. Valence-shell expansion requires the use of d-orbitals.
3.7 Characteristics of Multiple Bonds Summary
Multiple bonds are formed when an atom forms a σ-bond by using an sp or sp² hybrid orbital and one or more π-bonds by using unhybridized p-orbitals. The side-by-side overlap that forms a π-bond makes a molecule resistant to twisting, results in bonds weaker than σ-bonds, and prevents atoms with large radii from forming multiple bonds.
One that tends to move into a magnetic field and has at least one unpaired electron.
A compound with too few valence electrons to be assigned a valid Lewis structure.
Molecular Orbital Theory
An improvement upon Lewis structures; uses atomic orbitals to form molecular orbitals, which can then be used to explain a number of behaviors.
3.8 The Limitations of Lewis's Theories Summary
Unlike Lewis's theory, molecular orbital theory can account for the paramagnetism of oxygen and the existence of electron-deficient compounds.
Molecular Orbitals are
Orbitals constructed from atomic orbitals that are spread throughout the entire molecule. Electrons that were localized on atoms or between pairs of atoms are delocalized in molecular orbital theory.
A molecular orbital for H₂
ψ = ψA1s + ψB1s
ψA1s = a 1s-orbital centered on atom A
ψB1s = a 1s-orbital centered on atom B
A combination of atomic orbitals that results in an overall lowering of energy; constructive interference between two atoms concentrates the bonding electrons between the nuclei more.
A combination of atomic orbitals that results in an overall rising of energy; destructive interference forms a nodal surface between the two atoms where the atomic orbitals cancel completely.
The ratio of molecular orbitals to atomic orbitals
N molecular orbitals can be constructed from N atomic orbitals.
Molecular Orbital Energy-Level Diagram
A diagram showing the relative energy levels of atomic orbitals, bonding orbitals, and antibonding orbitals.
3.9 Molecular Orbitals Summary
Molecular orbitals are built from linear combinations of atomic orbitals; when atomic orbitals interfere constructively, they give rise to bonding orbitals; then they interfere destructively, they give rise to antibonding orbitals. N atomic orbitals combine to give N molecular orbitals.
Molecular orbitals and the building-up principle
Molecular orbital theory follows the building-up principle when filling in orbitals with electrons.
1. Electrons are accommodated in the lowest-energy molecular orbital, then in orbitals of increasingly higher energy.
2. According to the Pauli exclusion principle, each molecular orbital can accommodate up to two electrons. If two electrons are present in one orbital, they must be paired.
3. If more than one molecular obital of the same energy is available, the electrons enter them singly and adopt parallel spins (Hund's rule).
Bond order = # of bonding electrons - # electrons of antibonding orbitals ÷ 2.
BO = (σ-σ*)/2
3.10 Electron Configurations of Diatomic Molecules Summary
The ground-state electron configuration of diatomic molecules are deduced by forming molecular orbitals from all the valence-shell atomic orbitals of the two atoms and adding the valence electrons to the molecular orbitals in order of increasing energy, in accord with the building-up principle.
Heteronuclear Diatomic Molecule
A diatomic molecule build from atoms of two different atoms. The bond of a heteronuclear diatomic molecule is polar.
The sharing of charge/electrons in a
Nonpolar Covalent Bond
The electron pair is shared equally between the two atoms.
The sharing of charge/electrons in a
Polar Covalent Bond
The atomic orbital belonging to the more electronegative atom has the lower energy and so it makes the larger contribution to the lowest energy molecular orbital; the contribution to the highest-energy (most antibonding) orbital is greater for the higher-energy atomic orbital, which belongs to the less electronegative atom.
3.11 Bonding in Heteronuclear Diatomic Molecules Summary
Bonding in heteronuclear diatomic molecules involves an unequal sharing of the bonding electrons. The more electronegative element contributes more strongly to the bonding orbitals, whereas the less electronegative element contributes more strongly to the antibonding orbitals.