Structure and Properties Help

Hybridization and Bonding

A carbon atom must provide four equal-energy orbitals in order to form four equivalent bonds, as in methane, CH₄. It is assumed that the four equivalent orbitals are formed by blending the 2s and the three 2p AO's to give four new hybrid orbitals, called sp³ orbitals. The larger lobe, the "head," having most of the electron density, overlaps with an orbital of its bonding mate to form the bond. The smaller lobe, the "tail," is often omitted when depicting hybrid orbitals. However, at times the "tail" plays an important role in an organic reaction.

The s and p orbitals of carbon can hybridize in ways other than sp³, as shown below. Repulsion between pairs of electrons causes these hybrid orbitals to have the maximum bond angles. The sp² and sp hybrid orbitals induce geometries about the C's as shown below.

Head-to-head overlap of AO's gives a sigma (s) bond. The bond angles (angles between s-bonds) at sp³ carbons are 109.5°, leading to a tetrahedral geometry, The bond angles at sp² carbons are 120°, leading to a trigonal planar geometry, and the bond angles at sp carbons are 180°, leading to a linear geometry. The imaginary line joining the nuclei of the bonding atoms is the bond axis, whose length is the bond length.

Two parallel p orbitals overlap side-by-side to form a pi (π) bond. The bond axis lies in a nodal plane (plane of zero electronic density). Single bonds are σ bonds. A double bond is one σ and one π bond. A triple bond is one σ and two π bonds.

Electronegativity and Polarity

The electronegativity of an atom is its tendency to attract bonding electrons toward itself. The higher the electronegativity, the more strongly the atom attracts and holds electrons. A nonpolar covalent bond exists between atoms having a very small or zero difference in electronegativity. A few relative electronegativities are

F (4.0) > O (3.5) > Cl, N (3.0) > Br (2.8) > S, C, I (2.5) > H (2.1)

A bond formed by atoms of dissimilar electronegativities is called polar due to partial charge separation. The more electronegative element of a covalent bond is relatively negative in charge, while the less electronegative element is relatively positive. The symbols δ+ and δ– represent partial charges in polar bonds. These partial charges should not be confused with ionic charges. Polar bonds are indicated by ; the arrow points toward the more electronegative atom.

The vector sum of all individual bond moments gives the net dipole moment of the molecule. H₂O has polar bonds. Since the molecule has a bent shape, the dipoles of the bonds do not cancel and the molecule has a net dipole moment.

Resonance and Delocalized π Electrons

Resonance theory describes species for which a single structure does not adequately describe the species' properties. As an example, consider the cation on the next page (called the allyl cation):

A comparison of the calculated and observed bond lengths shows that the 2 C–C bonds are the same length. Neither resonance structure alone can explain this similarity in bond length. When resonance structures form a resonance hybrid, we obtain a structure consistent with the observed bond length. The resonance hybrid has some double-bond character between the central carbon and both outside carbons. This state of affairs is described by the non-Lewis structure in which dotted lines stand for the partial bonds in which there are delocalized π electrons in an extended π bond created from overlap of p orbitals on each atom. The symbol ↔ denotes resonance, not equilibrium.

The hybrid is more stable than any single resonance structure. The more nearly equal in energy the contributing structures, the greater the resonance energy. When contributing structures have dissimilar energies, the hybrid looks most like the lowest-energy structure. Contributing structures (a) differ only in positions of electrons (atomic nuclei must have the same positions); (b) must have the same number of paired electrons; and (c) must not place more than 8 electrons on any second period atom. Relative energies of contributing structures are assessed by the following rules.

1. Structures with the greatest number of covalent bonds are most stable. However, for second-period elements (C, O, N) the octet rule must be observed.
2. With few exceptions, structures with the least charge separation are most stable.
3. If all structures have formal charge, the most stable (lowest energy) one has – on the more electronegative atom and + on the more electropositive atom.
4. Structures with like formal charges on adjacent atoms have very high energies.
5. Resonance structures with electron-deficient, positively charged atoms have very high energy, and are usually ignored.

Practice problems for these concepts can be found at: Structure and Properties Practice Problems