Review for Chapter 9: Chemical Bonding I
1. A Lewis dot symbol consists of the symbol for an element surrounded by one dot for each valence electron found
in an atom of the element. Lewis dot structures are used for representative elements and noble gases but not for
transition metals, lanthanides, or actinides.
2. An ionic bond is formed by the electrostatic forces of attraction between positive and negative ions. An ionic
compound is formed by a large network of ions in which the positive and negative charges are balanced. The net
attractive forces are maximized in the structure of the compound.
3. The elements that are most likely to form ionic compounds have low ionization energies (such as alkali metals
and alkaline earth metals, which form cations) or high electron affinities (such as the halogens and oxygen, which
form anions).
4. Coulomb’s law states that the potential energy E between two ions is directly proportional to the product of their
charges (Qcation x Qanion) divided by the distance of separation between them, r:
Ek
QcationQanion
r
5. The lattice energy is the energy required to completely separate one mole of a solid ionic compound into
gaseous ions. The higher the lattice energy, the more stable the compound. It can be calculated using the BornHaber cycle, which is based on Hess’s law (refer to Section 9.3).
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6. A covalent bond is a bond in which two electrons (one pair) are shared by two atoms. In multiple covalent
bonds (double or triple bonds), two or three pairs of electrons are shared by two atoms. Some covalently-bonded
atoms also have lone pairs of electrons (pairs of valence electrons that are not involved in bonding) around them.
7. Bond length is the distance between the nuclei of two covalently-bonded atoms in a molecule. Bond length
decreases as bonding increases from a single to a double to a triple bond.
8. Ionic and covalent compounds show significantly different physical properties due to the difference in their
bonding. The electrostatic forces holding ions together in an ionic compound are usually very strong, leading to
compounds that are solids at room temperature and have high melting points. Many ionic compounds are soluble in
water and their aqueous solutions conduct electricity. Covalent compounds are usually gases, liquids, or lowmelting solids. Many covalent solids are insoluble in water. If they do dissolve, they do not dissociate so they do
not generally conduct electricity because no charged particles are present in solution.
9. The electrons in some covalent bonds are not shared equally between two atoms, giving rise to polar covalent
bonds in which the electrons spend more time in the area around one atom than the other. The bond in HF is an
example of a polar covalent bond. This type of bond is intermediate between ionic and pure covalent bonds.
10. Electronegativity is the ability of an atom to attract toward itself the electrons in a chemical bond. Figure 9.5
gives the relative electronegativities of common elements. Fluorine has the highest electronegativity, with a value
of 4.0. Electronegativity generally increases from left to right within a period of the periodic table.
Electronegativity generally decreases with increasing atomic number within a group of the periodic table.
11. An ionic bond tends to form when the electronegativity difference between the two bonding atoms is 2.0 or
more. A pure covalent bond has an electronegativity difference of 0. Polar covalent bonds correspond to
electronegativity differences greater than 0 but less than 2.
12. A Lewis structure is a representation of covalent bonding in which shared electron pairs are shown either as
lines or as pairs of dots between two atoms, and lone pairs are shown as pairs of dots on individual atoms. Only
valence electrons are shown in a Lewis structure.
13. The octet rule states that atoms tend to form bonds until they are surrounded by eight valence electrons.
14. Know how to write Lewis structures for molecules and ions! Use a four-step procedure as follows:
a. Write the chemical skeleton structure for the molecule using chemical symbols and placing bonded atoms next to
one another. In general, the least electronegative atom occupies the central position.
b. Total up the number of valence electrons contributed by each atom in the molecule. If the molecule has a charge,
be sure to account for it as well in the electron total.
c. Draw a single covalent bond between the central atom and each of the surrounding atoms. Complete the octets of
the atoms bonded to the central atom. Place any remaining electrons from the total on the central atom.
d. Check the central atom to see if the octet is complete. If not, try adding double or triple bonds between the
surrounding atoms and the central atom using lone pairs from the surrounding atoms until the central atom’s octet is
complete.
15. There are three categories of exceptions to the octet rule: incomplete octets, odd-electron molecules, and
expanded octets.
16. Incomplete octets occur for some elements in certain compounds where structures are formed with fewer than
8 electrons around the central atom. Examples include H (2 electrons), Be (4 electrons), B (6 electrons) and Al (6
electrons).
17. Odd-electron molecules occur for molecules in which the total number of electrons is odd rather than even.
Examples: NO, NO2. Unpaired electrons result in structures called radicals that tend to be highly reactive.
18. Atoms of elements in and beyond the third period of the periodic table form some compounds in which there are
more than 8 electrons around the central atom. These elements have 3d orbitals that can be used in bonding and
allow the atoms to have “expanded octets”, i.e., more than 8 electrons around the central atom. Some examples
are SF4, BrF3, XeF2, SF6, PF5, IF5, and XeF4.
19. A formal charge on an atom is the electrical charge difference between the number of valence electrons
contributed by an atom and the number of electrons assigned to that atom in a Lewis structure. The number of
electrons assigned to an atom in a Lewis structure includes all of the lone pair electrons on the atom plus one-half of
the number of shared electrons around the atom. For molecules, the sum of all of the formal charges must add up to
zero. For cations, the sum of the formal charges must equal a positive charge and, for anions, a negative charge.
20. Sometimes there is more than one acceptable Lewis structure for a given species. The most plausible Lewis
structure can be selected using the following guidelines:
a. For molecules, a Lewis structure in which there are no formal charges is preferable to one in which formal
charges are present.
b. Lewis structures with small formal charges (+1, -1) are preferable to structures with large formal charges (+2, +3,
and/or -2, -3, etc.)
c. Among Lewis structures having similar distributions of formal charges, the most likely structure is the one in
which the negative formal charges are placed on the more electronegative atoms.
21. A resonance structure is one of two or more Lewis structures for a single molecule that cannot be represented
accurately by only one Lewis structure.
Example: Ozone (O3)
O=O-O  O-O=O
22. Bond enthalpy is the enthalpy change required to break a specific bond in 1 mole of gaseous molecules. It is a
measure of the stability of a molecule. Different kinds of bonds have different bond enthalpies as summarized in
Table 9.4. Triple bonds have higher bond enthalpies than double bonds, which are greater than single bonds.
23. The enthalpy of a reaction in the gas phase can be estimated using the equation:
∆H° =  Bond enthalpies (reactant molecules) –  Bond enthalpies (product molecules)
using the bond enthalpies in Table 9.4. Know how to work problems like Example 9.13 and 9.14.