Chapter 9

Models of Chemical Bonding

9-2
 Types of Chemical Bonding

Ionic bonding involves the transfer of electrons and is

usually observed when a metal bonds to a nonmetal.

Covalent bonding involves the sharing of electrons and is

usually observed when a nonmetal bonds to a nonmetal.

Metallic bonding involves electron pooling and occurs

when a metal bonds to another metal.

9-5
 Figure 9.2 Three models of chemical bonding.

9-6
 Lewis Electron-Dot Symbols

To draw the Lewis symbol for any main-group element:

• Note the A-group number, which gives the number of
valence electrons.
• Place one dot at a time on each of the four sides of the
element symbol.
• Keep adding dots, pairing them, until all are used up.

Example:
Nitrogen, N, is in Group 5A and therefore has 5 valence electrons.

•• • • •
•N• or • N• or • N or N•
••

••
• •• • •

9-8
 Lewis Symbols and Bonding

For a metal, the total number of dots in the Lewis symbol

is the number of electrons the atom loses to form a cation.

For a nonmetal, the number of unpaired dots equals

- the number of electrons the atom gains to form an anion
- or the number it shares to form covalent bonds.

The octet rule states that when atoms bond, they lose,
gain, or share electrons to attain a filled outer level of 8
electrons (or 2, for H and Li).

9-9
 Figure 9.4
Lewis electron-dot symbols for elements in Periods 2 and 3.

9-10
 The Ionic Bonding Model

An ionic bond is formed when a metal transfers electrons

to a nonmetal to form ions, which attract each other to
give a solid compound.

The total number of electrons lost by the metal atom(s)

equals the total number of electrons gained by the
nonmetal atoms.

9-11
 Figure 9.5 Three ways to depict electron transfer in the formation
of Li+ and F-.

Electron configurations Li 1s22s1 + F 1s22p5 → Li+ 1s2 + F- 1s22s22p6

Orbital diagrams

Li ↑↓ ↑ Li+ ↑↓
+ 1s 2s 2p 1s 2s 2p

F ↑↓ ↑↓ ↑↓ ↑↓ ↑ F- ↑↓ ↑↓ ↑↓ ↑↓ ↑↓
1s 2s 2p 1s 2s 2p
Lewis electron-dot symbols

•• •• -
Li• •F Li+ + F
••

••
••
•• ••

9-12
 Sample Problem 9.1 Depicting Ion Formation

PROBLEM: Use partial orbital diagrams and Lewis symbols to depict

the formation of Na+ and O2− ions from the atoms, and
determine the formula of the compound formed.

PLAN: Draw orbital diagrams and Lewis symbols for Na and O

atoms. To attain filled outer levels, Na loses one electron and
O gains two. Two Na atoms are needed for each O atom so
that the number of electrons lost equals the number of
electrons gained.

SOLUTION:
Na• •• •• 2-
•O + O

••
••
2Na+
••

• ••
Na•

9-13
 Sample Problem 9.1

Na ↑
3s 3p + O ↑↓ ↑↓ ↑ ↑
Na ↑ 2s 2p

3s 3p

2Na+ + O2- ↑↓ ↑↓ ↑↓ ↑↓
2s 2p

The formula is Na2O

9-14
 Periodic Trends in Lattice Energy

Lattice energy is the energy required to separate 1 mol of

an ionic solid into gaseous ions.
Lattice energy is a measure of the strength of the ionic bond.

Coloumb’s Law

charge A x charge B
Electrostatic energy 
distance

cation charge x anion charge

Electrostatic energy  cation radius + anion radius  DHolattice

9-17
 Periodic Trends in Lattice Energy

Lattice energy is affected by ionic size and ionic charge.

As ionic size increases, lattice energy decreases.

Lattice energy therefore decreases down a group on the
periodic table.

As ionic charge increases, lattice energy increases.

9-18
 Properties of Ionic Compounds

• Ionic compounds tend to be hard, rigid, and brittle, with

high melting points.
• Ionic compounds do not conduct electricity in the solid
state.
– In the solid state, the ions are fixed in place in the lattice and do
not move.
• Ionic compounds conduct electricity when melted or
dissolved.
– In the liquid state or in solution, the ions are free to move and
carry a current.

9-20
 Figure 9.9 Why ionic compounds crack.

9-21
 Figure 9.10 Electrical conductance and ion mobility.

Solid ionic Molten ionic Ionic compound

compound compound dissolved in water
9-22
 Table 9.1 Melting and Boiling Points of Some Ionic Compounds

Compound mp (°C) bp (°C)

CsBr 636 1300

NaI 661 1304
MgCl2 714 1412

KBr 734 1435

CaCl2 782 >1600
NaCl 801 1413
LiF 845 1676
KF 858 1505
MgO 2852 3600

9-23
 Figure 9.11 Ion pairs formed when an ionic compound vaporizes.

Interionic attractions are so strong that when an ionic compound

is vaporized, ion pairs are formed.

9-24
 Bonding Pairs and Lone Pairs

Atoms share electrons to achieve a full outer level of

electrons. The shared electrons are called a shared pair
or bonding pair.

The shared pair is represented as a pair of dots or a line:

••
H H or H–H

An outer-level electron pair that is not involved in

bonding is called a lone pair, or unshared pair.

•• •• •• ••
F F or F–F
••
••
••

••

••

•• •• •• ••

9-27
 Properties of a Covalent Bond

The bond order is the number of electron pairs being

shared by a given pair of atoms.
A single bond consists of one bonding pair and has a bond order of 1.

The bond energy (BE) is the energy needed to

overcome the attraction between the nuclei and the
shared electrons. The stronger the bond the higher the
bond energy.

The bond length is the distance between the nuclei of

the bonded atoms.

9-28
 Trends in bond order, energy, and length

For a given pair of atoms, a higher bond order results in a

shorter bond length and higher bond energy.

For a given pair of atoms, a shorter bond is a stronger bond.

Bond length increases down a group in the periodic table

and decreases across the period.
Bond energy shows the opposite trend.

9-29
 Table 9.2 Average Bond Energies (kJ/mol) and Bond Lengths (pm)

9-30
 Table 9.3 The Relation of Bond Order, Bond Length, and
Bond Energy

9-31
 Figure 9.14 Bond length and covalent radius.

Internuclear distance Covalent Internuclear distance Covalent

(bond length) radius (bond length) radius

72 pm 114 pm

Internuclear distance Covalent Internuclear distance Covalent

(bond length) radius (bond length) radius

100 pm 133 pm

9-32
 Sample Problem 9.2 Comparing Bond Length and Bond Strength

PROBLEM: Using the periodic table, but not Tables 9.2 or 9.3, rank
the bonds in each set in order of decreasing bond length
and decreasing bond strength:
(a) S–F, S–Br, S–Cl (b) C=O, C–O, CΞO

PLAN: (a) S is singly bonded to three different halogen atoms, so the

bond order is the same. Bond length increases and bond
strength decreases as the atomic radius of the halogen
increases.
(b) The same two atoms are bonded in each case, but the
bond orders differ. Bond strength increases and bond
length decreases as bond order increases.

9-33
 Sample Problem 9.2

SOLUTION:

(a) Atomic size increases going down a group, so F < Cl < Br.

Bond length: S–Br > S–Cl > S–F

Bond strength: S–F > S–Cl > S–Br

(b) By ranking the bond orders, we get

Bond length: C–O > C=O > CΞO

Bond strength: CΞO > C=O > C–O

9-34
 Figure 9.15
Strong forces within molecules and weak forces between them.

9-35
 Figure 9.16 Covalent bonds of network covalent solids:
quartz and diamond.

9-36
 Electronegativity and Bond Polarity

A covalent bond in which the shared electron pair is not

shared equally, but remains closer to one atom than the
other, is a polar covalent bond.

The ability of an atom in a covalent bond to attract the

shared electron pair is called its electronegativity.

Unequal sharing of electrons causes the more

electronegative atom of the bond to be partially negative
and the less electronegative atom to be partially positive.

9-48
Figure 9.20 Bonding between the models.

Polar covalent bonds are much

more common than either pure
ionic or pure covalent bonds.

9-49
 Trends in Electronegativity

The most electronegative element is fluorine.

In general electronegativity decreases down a group as

atomic size increases.

In general electronegativity increases across a period

as atomic size decreases.

Nonmetals are more electronegative than metals.

9-51
 Figure 9.22 Electronegativity and atomic size.

9-52
 Depicting Polar Bonds

The unequal sharing of electrons can be depicted by a

polar arrow. The head of the arrow points to the more
electronegative element.

A polar bond can also be marked using δ+ and δ- symbols.

9-55
 Metallic Bonding

The electron sea model of metallic bonding proposes that:

• All metal atoms in the sample contribute their valence

electrons to form a delocalized electron “sea”.
• The metal “ions” (nuclei with core electrons) lie in an
orderly array within this mobile sea.
• All the atoms in the sample share the electrons.
• The metal is held together by the attraction between the
metal “cations” and the “sea” of valence electrons.

9-63
 Properties of Metals

• Metals are generally solids with moderate to high melting

points and much higher boiling points.
– Melting points decrease down a group and increase across a
period.
• Metals can be shaped without breaking.
– The electron sea allows the metal ions to slide past each other.
• Metals are good conductors of electricity in both the solid
and liquid states.
– The electron sea is mobile in both phases.
• Metals are good conductors of heat.

9-64
 Table 9.5 Melting and Boiling Points of Some Metals

Element mp (°C) bp (°C)

Lithium (Li) 180 1347

Tin (Sn) 232 2623

Aluminum (Al) 660 2467

Barium (Ba) 727 1850

Silver (Ag) 961 2155

Copper (Cu) 1083 2570

Uranium (U) 1130 3930

9-65
 Figure 9.28

Melting points of the Group 1A(1) and Group 2A(2) metals.

9-66
 Figure 9.29 Why metals dent and bend rather than crack.

9-67