Ch. 6
What is a chemical bond?

mutual electrical attraction between the nuclei and
valence electrons of different atoms that bind the
atoms together

Why don’t noble gases do this?
 Already have filled s and p orbitals
 stable octet: 8 valence e- (or 2 if, you’re helium)
Atoms that don’t have a
stable octet are more
reactive
 Key Point #1: By forming bonds with each other, most
atoms reduce their potential energy, becoming more
stable.
 This is a chemical change! All chemical changes
involve energy!
What types of bonds can be formed?
Metallic Bonding
•
In a metal, the empty orbitals in the atoms’ outer
energy levels overlap

Delocalized Electron: outer electron that does
not belong to any one atom but can move
freely through the metal’s network of empty
atomic orbitals.

sea of electrons: mobile electrons around the
metal atoms, which are packed together in a
crystal lattice.

metallic bonding: chemical bonding that results
from the attraction between metal atoms and
surrounding sea of electrons

Key Point: In metallic bonding, valence
electrons move freely throughout a network of
metal atoms.
Unique Characteristics of Metals
Metals have many unique properties because of
their sea of electrons

• Malleability: ability of a substance to be
hammered or beaten into thin sheets
• Ductility: ability of a substance to be pulled
into a thin wire
• Why? atoms can slide past one another along a
plane without breaking bonds
 Luster: shiny appearance
• Why? Absorb a wide range of light frequencies,
many orbitals separated by small energy
differences
 Conductivity
• Thermal: ability to conduct heat
• Electrical: ability to conduct electricity
• Why? Electrons move easily through network of
empty orbitals
Metallic Bonding Strength

The strength of metallic bonding is determined
by the enthalpy of vaporization:
• the amount of energy required to vaporize (turn
into a gas) 1 mol of a metal

In general, the strength of the metallic bond
INCREASES moving left to right across the
periodic table.
• Soft metals (less dense) metals  harder (more
dense) metals toward right
Properties of Metals: Malleability and Ductility
Properties of Metals: Surface Appearance
Properties of Metals: Electrical and Thermal
Conductivity
Types of Bonds
 What type of bonds can be formed?
 Ionic bond
 Covalent bond
○ Nonpolar covalent
○ Polar covalent
 Ionic bonding: bonds that result from electrical
attractions between cations and anions
 1 atom losses electrons
 1 atom gains electrons
 Covalent bonding: sharing of electrons between 2 or
more atoms

Key Point 2: Rarely is bonding between atoms purely
ionic or purely covalent.
 Instead, it usually falls somewhere between the two
extremes. Why?

Key Point 3: The extent of ionic or covalent bonding
between two atoms can be estimated by calculating
the difference in each elements’ electronegativity.
Covalent Bonding
 Large difference in E.N.: bond has more ionic
character
 Small difference in E.N: bond has more covalent
character
Types of Covalent Bonds
 Non-polar covalent bonding: both electrons equally
shared between atoms

Polar covalent bonding: unequal attraction for the
shared electrons
6.1 Practice Worksheet
Part 1
The property of electronegativity, which is the measure of
an atom’s ability to attract electrons, can be used to
predict the degree to which the bonding between atoms
of two elements is ionic or covalent.
The greater the electronegativity difference, the more
ionic the bonding is.

If the calculated electronegative difference is…
 > 1.7 : ionic bond is formed
 > 0.3 , < 1.7 : polar-covalent bond
 0 – 0.3 : non-polar covalent bond
Increasing difference in electronegativity
Nonpolar
Covalent
share e-
Polar Covalent
partial transfer of e-
Ionic
transfer e-
Elements
Mg to Cl
H to O
C to Cl
N to H
C to S
K to F
Na to Cl
H to H
Electronegativity Electronegativity
Difference
Element Element
Bond Type
1
2
1.2
2.1
3.0
1.8
3.5
1.4
Polar covalent
2.5
3.0
.5
Polar covalent
3.0
2.1
.9
2.5
2.5
.8
4.0
3.2
Ionic
.9
3.0
2.1
Ionic
2.1
2.1
0
Ionic
Polar covalent
0 Non-polar covalent
Non-polar covalent
Polyatomic Ions

It is also possible if a compound contains polyatomic
ions, for both types of bonding to be present.
 Monatomic Ions: Fe2+ , Na+, Cl Polyatomic Ions: PO43-, NH4+ , NO-1
 Groups of atoms are bonded covalent together, but
because of few or more than expected valence
electrons they have an overall charge (so they can
also bond ionically with other ions)
 Ex: Ca2+ and SO42-  CaSO4 (metal & diff. nonmetals)
Classify the following as ionic, covalent, or both
Ionic
1. CaCl2 = __________
(metal & nonmetal)
Both
5. BaSO4 = ___________
(metal & diff. nonmetals)
Covalent
2. CO2 = __________
(nonmetal & nonmetal)
Covalent
6. H2O = ____________
(nonmetal & nonmetal)
Ionic
3. MgO = __________
(metal & nonmetal)
Covalent
7. SO3 = ___________
(nonmetal & nonmetal)
Covalent
4. HCl = ___________
(nonmetal & nonmetal)
Both
8. AlPO4 = ___________
(metal & diff. nonmetals)
Section 6.2
What is a molecule?

Neutral group of atoms that are held together by
covalent bonds.
 Chemical formula: indicates the relative numbers of
atoms of each kind in a chemical compound by using
atomic symbols and numerical subscripts.
Formation of Covalent Bonds

The electrons of one atom and protons of the other
atom attract each another.

The two nuclei and two electrons repel each other.
These two forces cancel out to
form a covalent bond at a length
where the potential energy is
at a minimum.

Bond Length vs. Bond Energy
 Bond length (pm): distance between two bonded
atoms at their minimum potential energy
 Bond energy (kJ/mol): energy required to break a
chemical bond and form neutral isolated atoms.
○ Breaking bonds: absorbs (requires) energy
○ Forming bonds: releases energy

Key Point: As you increase the number of bonds
between 2 atoms the bond energy increases, while
the bond length decreases. This is an inverse
relationship.
Bond Energies & Bond Lengths
 A. How many electrons are shared in a
 single bond:
 double bond:
 triple bond:
 B. Which bond is shorter? C – C or C = C
 C. Which bond requires more energy to break?
 In addition to finding an ideal bond length, atoms
also lower their potential energy by achieving a
stable octet of 8 valence electrons
Bond Energies & Bond Lengths
 A. How many electrons are shared in a
 single bond: 2 e double bond: 4 e triple bond: 6e B. Which bond is shorter? C – C or C = C
 C. Which bond requires more energy to break? =
 In addition to finding an ideal bond length, atoms
also lower their potential energy by achieving a
stable octet of 8 valence electrons
Octet Rule

Octet Rule: chemical compounds tend to form so that
each atom has an octet of e-’s in its highest occupied
energy level
 Exceptions to the octet rule:
 Atoms that cannot fit eight electrons
 Atoms that can fit more than eight electrons
Hydrogen: 2eBoron: 6ePhosphorus, Sulfur, & Xenon: expanded
valence, more than 8e-
How can we represent molecules?

Lewis Structures: formulas in which atomic symbols
represent nuclei and inner shells, which are surrounded
by dot-pairs/dashes represent valence electrons
Chapter 6.5
VSEPR THEORY

Lewis Structures are 2D but we live in a 3D world!
 molecular geometry: the three-dimensional arrangement
of a molecule’s atoms

What do those 3D structure/shapes look like??
 Follow the Valance Shell Electron Pair Repulsion
Theory or VSEPR
○ Repulsion between the sets of valence electrons
surrounding an atom causes them to be oriented as
far away from each other as possible
Why use VSEPR Theory?

Key Point: VSERP Theory is used to predict the shape
of molecules based on the fact that electron pairs
strongly repel each other.

Following VSEPR allows us to predict bond polarity:
 uneven distribution of electrons
AB2 – Linear
Central
atom
Atoms/group of atoms
attached to central atom
Atoms bonded to
central atom (B)
Number of Lone Pairs
on central atom (E)
Bond Angle
2
0
180˚
Other Linear Geometries

The shape of two atoms bonded together is not
given in the chart.
 Ex: F2
 What is the only possible shape a binary
compound can have?
○ LINEAR!
AB2E1 – Bent
Atoms bonded to Number of Lone
central atom (B)
Pairs (E)
2
1
Bond Angle
<120˚
What happens to the bond angle
between atoms as you increase
the number of “lone pair electrons”
on the central atom?
Bond angles decrease!
AB2E2 – Bent
Atoms bonded to Number of Lone
central atom (B)
Pairs (E)
2
2
Bond Angle
104.5˚
AB3 – Trigonal Planar
Atoms bonded to Number of Lone
central atom (B)
Pairs (E)
3
0
Bond Angle
120˚
Shape is often
associated with atoms
that break octet rule,
but doesn’t have to be
AB3E1 – Trigonal Pyramidal
Atoms bonded to Number of Lone
central atom (B)
Pairs (E)
3
1
Bond Angle
107˚
AB4 – Tetrahedral
Atoms bonded to Number of Lone
central atom (B)
Pairs (E)
4
0
Bond Angle
109.5˚
Predicting Molecular Geometry
1. Draw Lewis structure for molecule.
VSEPR theory: if any lone pairs of electrons are
found on the central atom, these electrons decrease
the bond angles of atoms attached to it.
2. Draw a revised Lewis structure to show more
accurate geometry
O
S
AB2E
bent
F
O
F
S
F
AB4E
F
tetrahedral
Predicting Molecular Polarity
3. To indicate the polarity of the bonds, we use this
symbol: __________________ , which always points
toward the more electronegative element.
electron rich
region
electron poor
region
H
F
e- poor e- rich
H
+
d
4. When multiple bonds are found in a molecule, we
must identify polarity of each bond.
F
d
5. Observe the overall polarity of the molecule. Think of
it as “tug-of-war” for valence electrons between the
various atoms.
Non-polar covalent molecules: If the atom is
symmetrical and all atoms have an equal pull on
electrons
Polar covalent molecules: If the atom is not
symmetrical and/or the atoms do not all have an
equal chance of winning the tug of war for
electrons
Intermolecular forces:
attractive forces between molecules.
Intramolecular forces:
attractive forces within a molecule (the bonds)
Intermolecular
Forces
Intramolecular
Forces (bond)
Intramolecular
Forces
intermolecular forces are much weaker than
Strength of IMF
strongest
weakest
 Hydrogen
Bond
 Dipole – Dipole
 Induced Dipole
 London Dispersion
Forces
Dipoles
 What
is a dipole?
 A polar molecule
 Uneven sharing of electrons so there is a
separation of charge
electron poor
region
H
electron rich
region
F
e- poor
H
F
d+ d-
Dipole-Dipole Forces
 Attraction
between two polar molecules
—
+
—
+
Hydrogen Bonding
Special type of Dipole – Dipole
 Attraction between:
Hydrogen & Nitrogen/Oxygen/Fluorine

Induced Dipole

Attraction between one polar and one nonpolar
molecule
Electrons
shift toward
positive end
of dipole
—
—
+
+
—
+
London Dispersion Forces

Attraction between two nonpolar molecules
Electrons
become
uneven and
form a dipole
—
+
—
+
What does IMF effect?
 Viscosity
 Surface
Tension
 Boiling Point
Boiling Point

Point at which liquid particles escape the
surface of the liquid into the gas phase
Stronger IMF  Higher Boiling Point
Surface Tension
result of an imbalance of forces at the surface
of a liquid.
Stronger IMF  Higher Surface Tension
Viscosity

Measures a fluid’s resistance to flow
Stronger IMF  Higher Viscosity