Review for Chapter 10: Chemical Bonding II
1. The valence electrons, the electrons in the outermost occupied shell of an atom, are the
electrons involved in chemical bonding. Their arrangement determines the geometry of a
molecule.
2. The Valence-Shell Electron-Pair Repulsion (VSEPR) model is a simple model for
predicting molecular geometry. It is based on the assumption that valence-shell electron pairs
repel one another and tend to stay as far apart as possible.
3. To use the VSEPR model, the Lewis structure for the molecule is drawn and the number
of pairs of bonding electrons and lone pairs of electrons are determined. For molecules
where there are only bonding pairs of electrons (B) around the central atom (A), Table 10.1 can
be used to determine the arrangement of the electron pairs and the molecular geometry. For
molecules containing both bonding pairs (B) and lone pairs of electrons (E) around the central
atom (A), Table 10.2 is used to predict the arrangement of the electron pairs and the molecular
geometry.
4. Lone pairs of electrons repel other pairs of electrons more forcefully than bonding pairs do
and thus distort the bond angles from the ideal geometry.
5. A dipole moment is a measure of the charge separation in molecules containing atoms of
different electronegativities. For example, HF has a dipole moment because fluorine is more
electronegative than H and the electrons spend more time in the vicinity of the F nucleus than
near the H nucleus. This leads to a partial negative charge on the F end of the molecule and a
partial positive charge on the H end of the molecule. HF is an example of a polar molecule.
6. For molecules containing three or more atoms, the polarity of the bonds and the molecular
geometry determine whether there is a dipole moment. Even if polar bonds are present, the
molecule will not have a dipole moment if the bond moments cancel out due to a symmetrical
arrangement of the bonds around the central atom. If the electron pairs are not symmetrically
arranged around the central atom, however, the bond moments may add together to give a dipole
moment. For example, H2O has a bent shape and possesses a dipole moment because the bond
moments add together instead of canceling each other out as would be the case for a linear
molecule.
7. Polar molecules possessing dipole moments can be made to align using an electric field.
Nonpolar molecules are not affected by an electric field.
8. Valence bond theory is a quantum mechanical explanation for covalent bond formation.
Hybridized atomic orbitals are formed by the combination and rearrangement of orbitals from
the same atom. The number of hybridized orbitals is equal to the number of pure atomic orbitals
that are combined.
9. The hybrid orbitals and their shapes are summarized in Table 10.4. For linear molecules
in which there are two pairs of bonding electrons arranged around a central atom, sp hybrid
orbitals are used. For trigonal planar molecules in which three pairs of electrons are arranged
around a central atom, sp2 hybrid orbitals are used. For tetrahedral molecules in which four pairs
of electrons are arranged around a central atom, sp3 hybrid orbitals are used. For trigonal
bipyramidal molecules in which five pairs of electrons are arranged around a central atom, sp3d
hybrid orbitals are employed. The inclusion of the d orbital enables “expanded octets” to occur.
For octahedral molecules in which six pairs of electrons are arranged around a central atom,
sp3d2 hybrid orbitals are used.
10. Sigma bonds are covalent bonds formed by orbitals overlapping end-to-end, with the
electron density concentrated between the nuclei of the bonding atoms.
11. Pi bonds are covalent bonds formed by orbitals overlapping sideways, with the electron
density concentrated above and below the plane of the nuclei of the bonding atoms. Pi bonds can
be made using unhybridized p orbitals.
12. A carbon-carbon double bond consists of one sigma bond and one pi bond. A carboncarbon triple bond consists of one sigma bond and two pi bonds.
13. Molecular orbital theory is another quantum mechanical explanation for covalent bond
formation. Bonding is described in terms of the formation of new orbitals that are associated
with the molecule as a whole.
14. Two types of molecular orbitals are formed: bonding molecular orbitals and antibonding
molecular orbitals. A bonding molecular orbital is lower in energy and greater in stability than
the atomic orbitals from which it was formed. An antibonding molecular orbital is higher in
energy and lower in stability than the atomic orbitals from which it was formed.
15. Know how to draw molecular orbital diagrams for simple molecules as in Figure 10.26
and 10.27. The electrons fill the molecular orbitals from the lowest energy orbital up.
16. The bond order of a compound is calculated as follows:
bond order = ½ (number of electrons in bonding molecular orbitals – number of electrons in
antibonding molecular orbitals)
Molecules with a bond order of 0 are not stable and do not exist. Molecules with bond orders
greater than zero are stable and do exist.
17. The molecular orbital diagrams for Li2, B2, C2, N2, O2, and F2 are shown in Table 10.5.
Note that 2p orbitals are lower in energy than 2p orbitals for elements with an atomic number of
7 or less.
18. Molecular orbital theory helps explain why O2 shows paramagnetic properties. This
occurs because there are two unpaired electrons in the antibonding 2p* orbitals. Valence bond
theory was unable to explain this observation.
19. Delocalized molecular orbitals, in which electrons are shared among more than two atoms,
can be formed by the overlap of p orbitals on adjacent atoms. For example, in benzene (C6H6),
some electrons are shared among all six carbon atoms. Benzene is described as forming one set
of sigma bonds between the carbon atoms and another set of “partial” pi bonds in which
electrons are shared among all six carbons. Delocalized molecular orbitals are an alternative
approach to resonance structure in explaining observed molecular characteristics.