Chapter 6
Chemical Bonding
Introduction to Chemical Bonding
Atoms seldom exist as
independent particles in
nature--instead they are
combinations of atoms that
are held together by
chemical bonds.
A chemical bond is a mutual
electrical attraction between the
nuclei and valence electrons of
different atoms that binds the
atoms together.
By bonding, the atoms
decrease in potential energy,
thereby creating more stable
Types of chemical bonding:
Ionic bonding: chemical
bonding that results from the
electrical attraction between
large numbers of cations and
anions (metal and nonmetal)
 Atoms completely give up
electrons to other atoms
Covalent bondingchemical bonding that
results from the sharing
of electron pairs between
two atoms (nonmetal and
nonmetal)
Ionic or Covalent???
Bonding between atoms of
different elements is rarely purely
ionic or covalent. It usually falls
somewhere between, depending on
how strongly the atoms of each
element attract electrons.
Electronegativity- measure of an
atom’s ability to attract electrons
Remember…
Electronegativity increases across a
period and decreases down a group.
Which atom has the highest
electronegativity?
 Flourine
Which atom has the lowest?
 Francium
Ionic vs. Covalent
Bond Type
Electronegativity % ionic
difference
character
Nonpolar
covalent
0.0 to <0.4
Polar
covalent
0.4 to <1.7
Ionic
1.7 and greater
4% ionic
96 %
covalent
4-50% ionic
50%
covalent
50% ionic
50%
covalent
Practice
Elements
H and Cl
Li and Cl
Na and Cl
H and H
Electronegativit
y difference
Bond type
Nonpolar covalent bond- a
covalent bond in which the
bonding electrons are shared
equally by the bonding atoms,
resulting in a balanced
distribution of electrical
charge
Polar bond- bonds that have an
uneven distribution of charge(one
end of the molecule is more
electronegative than the others)
Polar covalent bond- a covalent
bond in which the bonded atoms
have an unequal attraction for the
shared electrons
Examples:
Nonpolar H-H
Polar H-Cl
What about HBr, CCl4, CO2,
NH3?
The composition of a compound
is given by its chemical formula.
 Chemical formula- indicates
the relative numbers of atoms
of each kind in a chemical
compound by using atomic
symbols and numerical
subscripts
Covalent Bonding Molecular Compounds
Many chemical compounds including most
of the chemicals that are in living things,
are composed of molecules
 Molecule- a neutral group of atoms that
are held together by covalent bonds
 Molecular compound- a chemical
compound whose simplest units are
molecules
Molecular formula- shows the
types and numbers of atoms
combined in a single molecule of a
molecular compound

Diatomic molecule- a molecule
containing only two atoms


examples:
H2
O2
Covalent Bond
Bond Length- the distance between two
bonded atoms at their minimum potential
energy, that is, the average distance
between two bonded atoms
Bond Energy- the energy required to break
a chemical bond and form neutral isolated
atoms (in kj/mol)
 Bond length and bond energies vary with
the types of atoms that have combined
Types of covalent bonds
Single

Hydrogen (H2)
Double

Ethene (C2H4)
Triple

Nitrogen (N2)
Octet Rule
Octet rule- chemical compounds tend
to form so that each atom, by gaining,
losing, or sharing electrons, has an
octet of electrons in its highest
occupied energy level
They want to be like the noble gases,
with their outer s and p orbitals
completely filled with 8 electrons.
Exceptions to the Octet Rule
Hydrogen-only needs 2 electrons
to be happy
Be can have only 4 electrons
B and Al can have only 6 electrons
rd
3 row or lower can sometimes
have more than 8 electrons (SF6,
PCl5)
Lewis Structures
Draw the following:
-IBr
-CH3Br
-C2HCl
-SiCl4
-F2O
Ionic Bonding and Ionic Compounds
Ionic compound- composed of
positive and negative ions that are
combined so that the numbers of
positive and negative charges are
equal
 The chemical formula of an ionic
compound represents the
simplest ratio of the compounds
combined ions
Formula unit- the simplest
collection of atoms from
which an ionic compound’s
formula can be established
The ratio depends on the
charges of the ions
combined

Characteristics of Ionic Bonding
In order for ions to minimize their potential
energy, they combine in an orderly
arrangement known as a crystal lattice.
Lattice energy- the energy released when
one mole of an ionic crystalline compound
is formed from gaseous ions
 Energy is released when crystals are
formed
Ionic vs. Molecular Compounds
Molecular (covalent)
Ionic
Low melting point
High melting point
Low boiling point
High boiling point
volatile
Non-volatile
Doesn’t conduct
electricity
Conducts electricity
in solution
Hard, brittle
Ionic Compounds
Remember, valence electrons are
transferred from one atom to another
Must have an element with low
electronegativity and an element with high
electronegativity
2 ions are formed-one positive and one
negative (cation, anion)
 They are attracted to each other by their
opposite charges
Electron-Dot Notation
Electron-dot notation- an electron
configuration in which only the valence
electrons of an atom of a particular
element are shown, indicated by dots
placed around the elements symbol
Examples:
F
H
N
Dot Diagrams for Ionic Compounds
Metals lose valence electrons and
have positive charges
Nonmetals gain enough electrons to
have 8 valence electrons and have a
negative charge
Alternate positive and negative ions in
the compound formula (if possible)
Practice
NaCl
MgF2
Ga2O3
CaO
Lewis Structures
Electron-dot notation can also be used to
represent molecules by combining the
notations of two individual atoms
H:H
F F
 Shared pair- a pair of electrons involved in
bonding
 Unshared pair (lone pair)- a pair of electrons
that is not involved in bonding
 You can change the shared pair to a dash
H - H
F-F
Lewis Structures- formulas in which
atomic symbols represent nuclei and innershell electrons, dot-pairs or dashes
between two atomic symbols represent
electron pairs in covalent bonds, and dots
adjacent to only one atomic symbol
represent unshared pairs
Structural formula- indicates what kind,
number, arrangement and bonds, but not
the unshared pairs of the atoms in a
molecule
Drawing Lewis Structures
1. Determine the type and number of atoms
in the molecule
2. Write the electron dot notation for each
type of atom in the molecule
3. Determine the total number of valence
electrons in the atoms to be combined
4. Arrange the atoms to form a skeleton
structure for the molecule (hint: if carbon
is present, it is always in the
middle…why?)
5. Add unshared pairs of electrons so that each
hydrogen atom shares a pair of electrons and
each other nonmetal is surrounded by 8
electrons.
6a. Count the electrons in the structure to be
sure that the number of valence electrons
used equals the number available
6b. If too many electrons have been used,
subtract one or more lone pairs until the total
number of valence electrons is correct. Then
move lone pairs until all have 8.
(this will require a double or triple bond)
Practice
NH3
H2S
CH4
H2O
Multiple Covalent Bonds
Some atoms can share more than one electron
pair
 Double bond- a covalent bond produced by
the sharing of 2 pairs of electrons between
two atoms
 Triple bond- a covalent bond produced by the
sharing of 3 pairs of electrons between two
atoms
 Examples: C2H4, C2H2, N2
More examples…
CH2O
CO2
HCN
Resonance Structures
Some molecules and ions cannot be
represented by a single Lewis structure
Resonance- refers to bonding in molecules
or ions that cannot be correctly
represented by a single Lewis structure
 Example: O3 (ozone), SO3
Polyatomic Ions
Polyatomic ions- a charged group of
covalently bonded atoms
 Polyatomic ions combine with ions of
opposite charge to form ionic
compounds
 The charge of the ion results from an
excess of electrons (- charge) or a
shortage of electrons (+ charge)
 Examples: NH4+, NO3-, SO4-2
Molecular Geometry
The properties of molecules
depend not only on the bonding of
the atoms but also on molecular
geometry
Molecular polarity- the uneven
distribution of molecular charge
VSEPR Theory
VSEPR (valence shell, electron-pair
repulsion) states that repulsion
between the sets of valence-level
electrons surrounding an atom
causes these sets to be oriented as
far as possible away from each
other
 Examples: BeF2, BF3, CH4
Linear
Trigonal
planar
Tetrahedral
Practice (use page 184 & 186)
Give the molecular geometry and bond
angles of the following:
 HI
 CBr4
 AlBr3
 CH2Cl2
VSEPR and Unshared Pairs
The lone pairs occupy space around
the central atom, but the actual shape
of the molecule is determined by the
positions of the atoms only
 Examples
NH3
H2O
ClO3
Hybridization
To explain how the orbitals of an atom
become rearranged when the atom
forms covalent bonds, a different
model is used
Hybridization- the mixing of 2 or more
atomic orbitals of similar energies of
the same atom to produce new orbitals
of equal energies
Hybrid orbitals- orbitals of equal
energy produced by the combination of
2 or more orbitals on the same atom
 Look at carbon, 2 of carbon’s valence
electrons occupy the 2s and 2
occupy the 2p. To achieve four
equivalent bonds, carbon’s 2s and 2p
orbitals hybridize
 They form 4 new identical orbitals
called sp3.
Metallic Bonding
Metals have unique properties
1. Excellent electrical properties
2. Shiny appearance
3. Malleability- the ability of a
substance to be hammered into thin
sheets
4. Ductility- the ability of a substance
to be pulled into wire
The Metallic Bond Model
The highest energy levels of most
metal atoms are occupied by very
few electrons
The vacant orbitals overlap
The electrons become delocalizedthey move freely in the empty
orbitals (called a sea of electrons)
The chemical bonding that results
from the attraction between metal
atoms and the surrounding sea of
electrons is metallic bonding.
Intermolecular Forces
The forces of attraction between molecules are
known as intermolecular forces
These forces vary in strength but are generally
weaker than bonds that join atoms in
molecules, ionic compounds, or metal atoms in
solid metals.
In other words, it generally takes far less energy
to separate molecules from one another than it
does to take molecules apart.
Dipole-Dipole Forces
The strongest intermolecular forces exist
between polar molecules and only polar
molecules
Dipole- created by equal but opposite charges
that are separated by a short distance
The direction of a dipole is from the dipole’s
positive pole to its negative pole
Polar covalent molecules act as little magnets,
they have positive ends and negative ends
which attract each other
A dipole is represented by an arrow
with a head pointing toward the
negative pole and a crossed tail
situated at the positive pole
 H Cl
 The forces of attraction between
polar molecules are known as dipoledipole forces
 Examples: NH3, CO2
Hydrogen Bonding
Hydrogen bonding- the intermolecular
force in which a hydrogen atom that is
bonded to a highly electronegative
atom is attracted to an unshared pair of
electrons of an electronegative atom in
a nearby molecule (usually N, O, or F)
Hydrogen bonding occurs between
water molecules
London Dispersion Forces
Even noble-gas atoms and molecules that
are nonpolar experience a weak
intermolecular attraction
London dispersion forces- the
intermolecular attractions resulting from
the constant motion of electrons and the
creation of an instantaneous dipole
London forces increase with
increasing atomic or molar mass- the
bigger the molecule, the stronger the
force!
These forces arise from the weak
attractive force of the electrons on one
molecule for the nuclei of another
molecule.
This is the only way the noble gases
stay together.