Chapter 8
Basic Concepts of
Chemical Bonding
Chemical Bond
• A chemical bond occurs
between atoms or ions when
they are strongly attracted to
each other
Types of Bonds
• ionic bond – electrostatic forces
between ions of opposite
charges (usually a metal and
nonmetal)
Types of Bonds
• ionic bond
Types of Bonds
• covalent bond – results from
sharing of electrons between
two atoms (usually 2 or more
nonmetals)
Types of Bonds
• metallic bond- found in metals
like copper, iron, aluminum
• array of positive ions immersed
in a sea of delocalized valence
electrons
Lewis Symbols
• consist of the chemical symbol
for the element plus a dot for
each valence electron
Lewis Symbols
Octet Rule
• states that atoms tend to gain,
lose, or share electrons until
they are surrounded by 8
valence electrons
Ionic Bonding
• Example of formation of ionic
bond
Ionic Bonding
Ionic Bonding
• Lattice Energy is the energy
required to completely separate
a mole of solid ionic compound
into its ions
• (breaking bonds is an endothermic process)
Ionic Bonding
• The magnitude of lattice energy
depends on the charge of ions,
their sizes, and their
arrangement
• Eel = k Q1Q2/ d
Ionic Bonding
Covalent Bonding
• Shared pairs of electrons bind
atoms together
• The attraction of the nucleus to
the electrons overcomes the
repulsion of the electrons
Covalent Bonding
• Lewis structures show each
electron pair shared between
atoms as a line and unshared
electrons as dots
Bond Polarity
• describes the sharing of
electrons between atoms
Bond Polarity
• A nonpolar covalent bond is
one in which electrons are
shared equally between 2 atoms
Bond Polarity
• A polar covalent bond is where
one of the atoms exerts a
greater attraction for the
bonding electrons than the
other.
Bond Polarity
• if the difference in the attraction
between two atoms is large
enough an ionic bond occurs
Electronegativity
• ability of an atom in a molecule
to attract electrons to itself
• depends on the ionization
energy and electron affinity
Electronegativity
• Based on Pauling’s Scale
Electronegativity
• Consider F2, HF, and LiF
Polar Molecules
• Not only can individual bonds
be classified as polar but so can
an entire molecule
Polar molecules
• Polar molecules have a positive
end and a negative end which
accounts for many properties of
substances
Polar Molecules
• In polar molecules a dipole is
established based on the
separation of charge in the
molecule
Polar Molecules
• A dipole moment is the
measure of the magnitude of the
dipole
• u=Qr , Q= value of charge, r =
distance
Polar Molecules
• it is measured in debyes (D) a
unit that is equal to 3.34 x 10-30
Coulomb – meters (C-m)
Polar Molecules
Drawing Lewis Structures
• 1. Sum the total valence
electrons
• 2. Write the symbols for the
atoms to show which atoms are
attached to which and connect
them with a single bond
Drawing Lewis Structures
• A. Often they are written in the
order of which they are attached
• B. When the central atom has a
group of atoms bonded to it, the
central atom is written first
• C. usually the central atom is less
electronegative
Drawing Lewis Structures
• 3. Complete the octet of the atoms
bonded to the central atom
• 4. Place left over electrons on the
central atom, even if it results in more
than an octet
• 5.If there are not enough electrons to
give the central atom an octet try
multiple bonds
Practice 8.6 – 8.8
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Draw the Lewis structure for:
PCl3
CH2Cl2
HCN
NO+
C2H4
BrO3–
ClO2–
PO43 –
Formal Charge
• CO2
Formal Charge
• Used when more than one
structure can be drawn for a
molecule
Formal Charge
• equals the number of valence
electrons in the atom minus the
number of unshared electrons
minus half the bonding
electrons
Formal Charge
• As a general rule, formal
charges of 0 are preferred and
any negative charge should
reside on the more
electronegative atom
Sample Exercise 8.9 Lewis
Structures and Formal Charges
• The following are three possible Lewis structures
for the thiocyanate ion, NCS–:
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• (a) Determine the formal charges of the atoms in
each structure. (b) Which Lewis structure is the
• preferred one
Resonance Structures
• When the placement of atoms in
a Lewis structure is the same
and the placement of electrons
is different, we use resonance
structures
Bond Length
• as the number of bonds between
2 atoms increases, the bond
grows shorter and shorter and
stronger and stronger
Bond Length
Resonance Structures
• Examples O3, NO3-, C6H6
Sample Exercise 8.10 Resonance
Structures
Which is predicted to have the shorter sulfur–
oxygen bonds, SO3or SO32–?
Exceptions to the octet rule
• Molecules with odd numbers of
electrons – NO
Exceptions to the octet rule
• Molecules in which an atom has
less than an octet – BF3, and
usually it reacts with a molecule
with unshared electrons like
NH3
Exceptions to the octet rule
• More than an octet – PCl5, SF6,
XeF4
Sample Exercise 8.11 Lewis
Structures for an Ion with an
Expanded Valence Shell
Draw the Lewis structure for ICl4– and XeF2
Strengths of Covalent Bonds
• Bond enthalpy is the enthalpy
change DH for the breaking of a
particular bond
Strengths of Covalent Bonds
Strengths of Covalent Bonds
DHrxn = Sbond enthalpies
broken - Sbond enthalpies
formed
Strengths of Covalent Bonds
Sample Exercise 8.12 Using
Average Bond Enthalpies
Using Table 8.4, estimate ΔH for the following reaction
Sample Integrative Exercise
Putting Concepts Together
• Phosgene, a substance used in poisonous gas warfare
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during World War I.
Phosgene has the following elemental composition:
12.14% C, 16.17% O, and 71.69% Cl by mass. Its molar
mass is 98.9 g/mol.
(a) Determine the molecular formula of this compound.
(b) Draw three Lewis structures for the molecule that
satisfy the octet rule for each atom. (The Cl and O atoms
bond to C.)
(c) Using formal charges, determine which Lewis structure
is the most important one.
(d) Using average bond enthalpies, estimate ΔH for the
formation of gaseous phosgene from CO(g) and Cl2(g).
Review
•SO2
Review
•CO2
Review
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+
NO
Review
•ICl2
-
Review
•Br2
Review
•BCl3
Review
•CO3
2-
Review
•NO2
-
Review
•SO4
2-
Review
•XeF4
Review
•ClO2