Chapter 8
Basic Concepts of Chemical Bonding
8.1 Chemical Bonds, Lewis Symbols,
and the Octet Rule
The forces that hold atoms or ions together in compounds are called chemical bonds.
Chemical bonds can be classified as ionic, covalent, or metallic.
An ionic bond results from the powerful electrostatic forces that exist between oppositely
charged ions. Ionic substances form readily between elements from the far left of the
periodic table (metals) and elements from the far right of the periodic table (nonmetals).
Metals tend to lose electrons to form positively charged cations, while nonmetals tend to
gain electrons to form negatively charged anions. Ionic substances are solids at room
temperature.
When atoms are similar in their tendencies to lose or gain electrons, they share electrons
to form a covalent bond. The most familiar examples of covalent bonds are found in the
interactions of nonmetals. Substances held together by covalent bonds can be solids,
liquids, or gases at room temperature.
Atoms in metallic solids such as copper, iron, and aluminum are held together by
metallic bonding. In this type of bonding, valence electrons of the metal atoms are free
to move throughout the metal solid.
The electrons lost or gained in ionic bond formation and shared in covalent bond
formation are the outermost or valence electrons. A Lewis symbol (also called an
electron-dot symbol) can be used to show the valence electrons of an atom or ion. The
number of valence electrons of any element is the same as the group number of the
element in the periodic table. For example, the Lewis symbol for oxygen, a member of
group 6A, shows 6 dots. The dots are placed around the symbol for the element; in the
case of oxygen, around the O. A maximum of eight dots can be placed around a symbol,
where each dot represents a valence electron.
Dots are placed above, below, to the left, and to the right of the element symbol. Each
position can accommodate two electrons, and electrons are not "paired" until each of the
four positions contains a single electron. Table 8.1 gives the Lewis symbols for elements
of the second period.
Atoms tend to lose, gain, or share enough electrons to achieve a noble-gas electron
configuration. Chlorine, for example, gains one electron to become isoelectronic with
argon. Sodium loses one electron to become isoelectronic with neon. When two chlorine
atoms bond to form a chlorine molecule, each achieves a noble-gas electron configuration
through sharing electrons in the covalent bond.
These observations are summarized in the octet rule. Atoms tend to gain, lose, or share
electrons until they are surrounded by eight valence electrons.
8.2
Ionic Bonding
Sodium chloride is an example of an ionic compound. It forms spontaneously and
exothermically when sodium metal is brought into contact with chlorine gas.
In the formation of sodium chloride, sodium and chlorine each achieve a noble-gas
configuration: sodium by loss of its one valence electron, and chlorine by gaining one
electron to complete its octet. The formation of this ionic solid is facilitated by
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the low ionization energy of sodium (496 kJ/mol)
chlorine's tremendous electron affinity (–349 kJ/mol)
the electrostatic attraction that pulls together resulting ions of opposite charge
A measure of the magnitude of the electrostatic attractive forces in ionic compounds is
the lattice energy. The lattice energy is defined as the energy required to convert a mole
of an ionic solid into its constituent ions in the gas phase. Although it is not possible to
measure the lattice energy directly, it is possible to calculate it with Hess's law, using
measurable quantities and the Born-Haber cycle.
Figure 8. 4. The Born-Haber cycle shows the energetic relationships
in the formation of ionic solids from the elements—in this case,
sodium chloride. The enthalpy of formation is equal to the sum of the
energies of several individual steps.
Differences in lattice energies can be understood in the context of Coulomb's law. The
magnitude of the lattice energy is proportional to the product of the charges, Q1 and Q2,
divided between the distance between them.
The constant k is 8.99 109 J-m/C2, and d is the distance between ion centers. Table 8.2
gives the lattice energies for some common ionic compounds.
The discussion of ions so far has been limited to main-group atomic ions. The ions in
ionic solids can also be transition-metal ions or polyatomic ions.
Transition-metal ions generally do not have a noble-gas electron configuration. Although
the octet rule can be useful, it does not apply to every situation. When a transition-metal
atom becomes an ion, the electrons lost first are those in the orbitals with the highest
principal quantum number. Note that these are not the last electrons assigned when
determining the electron configuration. A zinc atom loses two electrons to become Zn2+.
The electrons it loses come from the 4s orbital.
The atoms in polyatomic ions are held together by covalent bonds. A polyatomic ion acts
as a single charged unit when interacting with oppositely charged ions to form an ionic
solid.
8.3
Size of Ions
When an atom loses one or more electrons, the resulting cation is smaller than the
original, neutral atom. One reason for this is simply the removal of the outermost electron
shell. The loss of an electron from a sodium atom, for instance, changes an atom with its
outermost electrons in principal quantum level 3 to an ion with its outermost electrons in
principal quantum level 2.
When an atom gains one or more electrons, the resulting anion is larger than the original,
neutral atom. Even with the added electrons going into the same electron shell as the
existing valence electrons, the radius increases. This is caused by the increased
electrostatic repulsions between the valence electrons.
Figure 8.5 illustrates the sizes of main-group atoms and their corresponding ions. Note
that just as atomic size increases from top to bottom within a group, the sizes of ions
increase from top to bottom as well. Note also that for an isoelectronic series such as Li+,
Be2+, and B3+, the radius decreases from left to right. Coulomb's law explains this. For an
isoelectronic series, the species with the largest nuclear charge will have the smallest
radius. All other things being equal, the larger the positive charge attracting the sphere of
valence electrons, the more closely the valence electrons are drawn to the nucleus.
Figure 8.
5.
8.4
Covalent Bonding
The formation of a covalent bond involves the sharing of electrons between atoms.
Species held together by covalent bonds, such as molecules or polyatomic ions, can be
illustrated using Lewis structures. In a Lewis structure, electrons that are shared in a
covalent bond are represented by lines and unshared electrons are written as dots.
A single line denotes the sharing of two electrons in a single bond. Lewis structures can
also be used to show multiple bonding in a molecule or polyatomic ion. A double line
means that two pairs of electrons are shared in a double bond, and a triple line means
that three pairs of electrons are shared in a triple bond.
8.5 Bond Polarity and
Electronegativity
The bond between hydrogen atoms in a hydrogen molecule is a nonpolar covalent bond,
meaning that the electrons shared by the two atoms are shared equally. In most cases the
sharing of electrons between two atoms is not equal. Unequal sharing of electrons
constitutes a polar covalent bond. Atoms generally do not share electrons equally
because of differences in their abilities to attract electrons toward themselves. The ability
of an atom in a molecule to attract electrons to itself is called electronegativity.
Electronegativity is related to, but is not the same as, electron affinity. A difference in
electronegativity causes atoms in a molecule to have "partial charges," whereas ions have
discrete charges
Electronegativity increases from left to right across the periodic table and decreases from
top to bottom within a group. The polarity of a bond is determined by considering the
difference in electronegativities of the two bonding atoms.
The greater the difference in electronegativity, the greater the polarity of the bond. A
quantitative measure of the polarity of the bond in a diatomic molecule is the dipole
moment. When two charges of equal magnitude and opposite charge are separated by
some distance, a dipole is established. The dipole moment, m , is calculated as
where Q is the magnitude of the separated charges and r is the distance between them. If
charges of –1.60 10–19 C (the charge on an electron) and +1.60 10–19 C are separated
by a distance of 1.00 Å, the dipole moment is:
The last two terms in the calculation are for the purpose of canceling units and getting
dipole moment in units of debyes (D).
The dipole moment calculated in this way assumes that the charges on the atoms are +1
and –1 (in units of e, electronic charge). However, if two atoms truly had charges of +1
and –1, they would not be connected by a covalent bond. Instead, they would be ions and
would be held together by ionic bonding.
Dipole moments can be measured experimentally. With an experimentally measured
dipole moment and a bond length, we can determine the actual magnitude of "partial
charge" on the atoms held together by a polar covalent bond.
Table 8.3 presents the bond lengths and measured dipole moments of the hydrogen
halides.
When there is more than one bond in a molecule, the individual bond polarities contribute
to the overall polarity of the molecule.
8.6
Drawing Lewis Structures
Drawing Lewis structures is an important first step toward understanding bonding in
compounds. Use the following guidelines to practice drawing these structures.
1. Sum the valence electrons from all atoms. (Use the periodic table to help you do
this.) For an anion, add an electron to the total for each negative charge. For a
cation, subtract an electron for each positive charge. Don't worry about keeping
track of which electrons come from which atoms. Only the total number is
important.
2. Write the symbols for the atoms to show which atoms are attached to which, and
connect them with a single bond (a dash, representing two electrons). Chemical
formulas are often written in the order in which the atoms are connected in the
molecule or ion. When a central atom has a group of other atoms attached to it,
the central atom is usually written first. In other cases you may need more
information before you can draw the Lewis structure.
3. Complete the octets of the atoms bonded to the central atom. (Remember,
however, that hydrogen can have only two electrons.)
4. Place any leftover electrons on the central atom, even if doing so results in more
than an octet.
5. If there are not enough electrons to give the central atom an octet, try multiple
bonds. Use one or more of the unshared pairs of electrons in the atoms bonded to
the central atom to form double or triple bonds.
The skeleton for a Lewis structure is usually straightforward. Place in the center a unique
atom of lowest electronegativity. On occasion, though, the two criteria (uniqueness and
low electronegativity) do not give the same central atom. The N2O molecule is an
example of this. Should the oxygen be in the center of the skeleton because it is the
unique atom? Or should one of the nitrogen atoms be in the center because nitrogen is
less electronegative than oxygen? To solve this problem, we use a formal charge
analysis of the Lewis structures resulting from each possibility.
The formal charge of an atom is its number of valence electrons minus the number of
electrons associated with it in the Lewis structure. Associated electrons include all
nonbonding electrons and half of the bonding electrons.
8.7
Resonance Structures
It is sometimes possible to write more than one correct Lewis structure for a molecule.
When two or more correct Lewis structures differ only by placement of electrons, they
are called resonance structures.
8.8
Exceptions to the Octet Rule
As useful as it is, the octet rule is not always obeyed. The three main types of exceptions
are
1. Molecules with an odd number
of electrons
2. Molecules in which an atom has
less than an octet
3. Molecules in which an atom has
more than an octet
1. Molecules with an odd number of electrons
Molecules such as NO have an odd number of valence electrons and cannot obey the
octet rule. When distributing the electrons in such a molecule, remember to distribute
electrons first to the most electronegative element(s). Thus, oxygen has a complete octet
in NO, while nitrogen does not.
2. Molecules in which an atom has less than an octet
Another example is boron trifluoride, which is discussed in your textbook.
3. Molecules in which an atom has more than an octet
Large atoms, those in the third period and beyond, can accommodate more than an octet
of electrons. The sulfur atom in SF6 has 12 electrons around it. Like many elements,
sulfur can obey the octet rule. But it doesn't have to obey it. (There are six electrons
around each fluorine atom, but for the purpose of clarity, they are not shown in this
structure.)
8.9
Strengths of Covalent Bonds
The strength of a bond is measured by how much energy is required to break it. For
instance, it requires 436 kJ/mol to turn 1 mol of H2(g) into 2 mols of H(g). This process
requires the breaking of 1 mol of H–H bonds and the energy required to do it is called the
H–H bond enthalpy.
The H–H bond occurs only in diatomic hydrogen. Similarly, the F–F bond occurs only in
diatomic fluorine. These bond enthalpies are measured exactly. Most bond enthalpies are
averages of the strength of a particular bond that occurs in more than just one compound.
The C–H bond, for example, occurs in a wide variety of organic compounds. Table 8.4
gives the average bond enthalpies for a variety of bonds.
Note that all bond enthalpies are positive numbers because they correspond to H values
for breaking bonds. Breaking bonds is always an endothermic process. The bond
enthalpy for a particular bond is represented as D(bond type). For example, D(N–H) is
the average bond enthalpy (in kJ/mol) of the N–H bond.
It is possible to estimate the H of a reaction using Lewis structures for the reactants and
products of a reaction and using average bond enthalpies. Consider the reaction
where the net result of the reaction is the breaking of a C–H bond and a Cl–Cl bond and
the formation of a C–Cl bond and an H–Cl bond. It is easier to visualize this with Lewis
structures.
Figure 8.13. Illustration of the use of average bond enthalpies to
calculate Hrxn for the reaction above. Breaking the C–H and C–Cl
bonds produces a positive enthalpy change, H1, whereas making
the C–Cl and H–Cl bonds causes a negative enthalpy change, H2.
The values of H1 and H2 are estimated from the values in Table
8.4. From Hess's law, Hrxn = H1 + H2.
Note that bond length is inversely proportional to bond enthalpy. The greater the number
of electrons shared between two nuclei, the shorter and stronger the bond.