General Chemistry CHM115

The Molecular Nature of Matter and Change

Fall semester 2020

1
9-1
 Chapter 9

Models of Chemical Bonding

9-2
 Models of Chemical Bonding

9.1 Atomic Properties and Chemical Bonds

9.3 The Covalent Bonding Model

9.5 Between the Extremes:

Electronegativity and Bond Polarity

9.6 Introduction to metallic bonding

9-3
 Models of Chemical Bonding

9.1 Atomic Properties and Chemical Bonds

9.2 The Ionic Bonding Model

9.3 The Covalent Bonding Model

9.5 Between the Extremes:

Electronegativity and Bond Polarity

9.6 An Introduction to Metallic Bonding

9-4
 Figure 9.1 A comparison of metals and nonmetals.

9-5
 Lewis Electron-Dot Symbols
To draw the Lewis symbol for any main-group element:
• Note the 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 5 and therefore has 5 valence electrons.

•• • • •
•N• or • N• or • N or N•
••

••
• •• • •

9-6
 Figure 9.4
Lewis electron-dot symbols for elements in Periods 2 and 3.

9-7
 Lewis Symbols and Bonding

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: duet rule).

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.

9-8
 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.

9-9
 Figure 9.2 Three models of chemical bonding.

9-10
 Models of Chemical Bonding

9.1 Atomic Properties and Chemical Bonds

9.2 The Ionic Bonding Model

9.3 The Covalent Bonding Model

9.5 Between the Extremes:

Electronegativity and Bond Polarity

9.6 An Introduction to Metallic Bonding

9-11
 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 atom(s).

9-12
 Figure 9.5 Three ways to depict electron transfer in the formation
of Li+ and F–.

Electron configurations: Li 1s22s1 + F 1s22s22p5 → 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
••

••
••
•• ••

9-13
 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•

9-14
 Sample Problem 9.1

Na ↑
3s 3p + O ↑↓ ↑↓ ↑ ↑
Na ↑ 2s 2p

3s 3p

2Na+ + O2– ↑↓ ↑↓ ↑↓ ↑↓
2s 2p

The formula is Na2O

9-15
 Properties of Ionic Compounds

• Ionic compounds tend to be hard, rigid, with high melting

points, form ion pairs when vaporize. “Strong attraction
between the ions”
• Ionic compounds crack upon hammering.
• 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.

9-16
 Table 9.1 Melting and Boiling Points of Some Ionic Compounds

Compound mp (ºC) bp (ºC)

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

9-17
 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.

9-18
 Figure 9.9 Why ionic compounds crack.

9-19
 Figure 9.10 Electrical conductance and ion mobility.

Solid ionic Molten ionic Ionic compound

compound compound dissolved in water
9-20
 Models of Chemical Bonding

9.1 Atomic Properties and Chemical Bonds

9.2 The Ionic Bonding Model

9.3 The Covalent Bonding Model

9.5 Between the Extremes:

Electronegativity and Bond Polarity

9.6 An Introduction to Metallic Bonding

9-21
 Figure 9.12 Covalent bond formation in H2.

9-22
 Bonding Pairs and Lone Pairs
in Covalent Compounds
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
••
••
••

••

••

•• •• •• ••

9-23
 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.

9-24
 Table 9.2 Average Bond Energies (kJ/mol) and Bond Lengths (pm)

9-25
 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 (stronger bond)

Bond length increases down a group in the periodic table

“as the atomic size increases” and decreases across the
period “as the atomic size decreases” . Bond energy shows
the opposite trend.

9-26
 Table 9.3 The Relation of Bond Order, Bond Length, and
Bond Energy

9-27
 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

9-28
 Sample Problem 9.3 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.

9-29
 Sample Problem 9.3

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

9-30
 Figure 9.15
Strong forces within molecules and weak forces between them.

9-31
 Models of Chemical Bonding

9.1 Atomic Properties and Chemical Bonds

9.2 The Ionic Bonding Model

9.3 The Covalent Bonding Model

9.5 Between the Extremes:

Electronegativity and Bond Polarity

9.6 An Introduction to Metallic Bonding

9-32
 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.

9-33
Figure 9.20 Bonding between the models.

Polar covalent bonds are much

more common than either pure
ionic or pure covalent bonds.

9-34
 Polar Covalent Bonds

The unequal sharing of electrons is the origin of the polar

covalent bond.
This bond can be depicted by a polar arrow. The head of
the arrow points to the more electronegative element.

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

symbols.
δ+ δ–

9-35
 Figure 9.23 Electron density distributions in H2, F2, and HF.

In HF, the electron density shifts from H to F.

The H–F bond has partial ionic character.

9-36
 Trends in Electronegativity

In general electronegativity decreases down a group as

atomic size increases.

In general electronegativity increases across a period

as atomic size decreases.

The most electronegative element is fluorine.

Nonmetals tend to be more electronegative than metals.

9-37
 Figure 9.21 The Pauling electronegativity (EN) scale.

9-38
 Figure 9.22 Electronegativity and atomic size.

9-39
 Electronegativity and Bond Polarity

The higher the electronegativity difference

(Δ EN), the higher the bond polarity
and the higher the ionic character of
the bond. This means higher bond energy,
higher melting point and electrical
conductivity of the compound.

Figure 9.24
ΔEN ranges for classifying the partial ionic
character of bonds.

9-40
 Figure 9.26 Electron density distributions in bonds of the
Period 3 chlorides.

There is a steady increase in electron sharing from left to right.

9-41
 Figure 9.27 Properties of the Period 3 chlorides.
Copyright © The McGraw-Hill Companies, Inc. Permission required for reproduction or display.

As DEN decreases, melting point and electrical conductivity decrease because

the bond type changes from ionic to polar covalent to nonpolar covalent.

9-42
 Sample Problem 9.5 Determining Bond Polarity from EN Values

PROBLEM: (a) Use a polar arrow to indicate the polarity of each

bond: N–H, F–N, I–Cl.
(b) Rank the following bonds in order of increasing
polarity: H–N, H–O, H–C.
PLAN: (a) We use Figure 9.21 to find the EN values for each
element. The polar arrow points toward the more
electronegative element.
(b) The greater the DEN between the atoms, the more polar
the bond.
SOLUTION: (a) The EN values are:
N = 3.0, H = 2.1; F = 4.0; I = 2.5, Cl = 3.0

N–H F–N I–Cl

9-43
 Sample Problem 9.5

(b) The EN values are:

N = 3.0, H = 2.1; O = 3.5; C = 2.5

DEN for H–N = 3.0 – 2.1 = 0.9

DEN for H–O = 3.5 – 2.1 = 1.4
DEN for H–C = 2.5 – 2.1 = 0.4

In order of polarity:
H-C < H-N < H-O

9-44
 Models of Chemical Bonding

9.1 Atomic Properties and Chemical Bonds

9.2 The Ionic Bonding Model

9.3 The Covalent Bonding Model

9.5 Between the Extremes:

Electronegativity and Bond Polarity

9.6 An Introduction to Metallic Bonding

9-45
 Metallic Bonding
“Electron Sea Model”

The electron sea model of metallic bonding proposes that:

• All metal atoms in the sample give up 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.
• The most familiar example of the Metallic bonding is the
“Alloys” that are mixtures of different metals in specific
ratio.

9-46
9-47
 Properties of Metals
• Metals are generally solids with moderate to high melting
points and much higher boiling points.
• Strong attraction between the cations and the electron sea.
• Metals can be shaped without breaking.
• The electron sea prevents repulsion between cations but rather it
allows the metal ions to slide past each other and end up in a new
position, so metals dent and bent contrary to ionic compounds.
• 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.

9-48
 Table 9.5 Melting and Boiling Points of Some Metals

Element mp (ºC) bp (ºC)

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

9-49
 Figure 9.28
Melting points of the Group 1A(1) and Group 2A(2) metals.

Metals of group 2 are divalent, so stronger attraction

between the cation and the electron sea.

9-50
 Figure 9.29 Why metals dent and bend rather than crack.

9-51