Chapter 6
Chemical Bonding
Section 1:
 Introduction to
chemical bonding
Introduction to chemical bonding
What is a chemical bond???
A mutual electrical attraction
between the nuclei and valence
electrons of different atoms
that binds the atoms together
Introduction to chemical bonding
 Why do atoms bond?
They are working to achieve
more stable arrangements where
the bonded atoms will have lower
potential energy than they do
when existing as individual
atoms- increase stability.
Introduction to chemical bonding
 Types of Chemical Bonding:
1. Ionic – an electrical attraction that
forms between cations (+) and
anions (-)
2. Covalent – are formed when
electrons are shared between atoms
3. Metallic – formed by many atoms
sharing many electrons
Introduction to chemical bonding
However….
 Bonds
are never purely
covalent or purely ionic.
 The degree of ionic-ness or
covalent-ness depends on
property of electronegativity.
Degree of Ionic/Covalent
Character in Chemical Bonds
Ionic
100%
50%
Polar-Covalent
5%
Nonpolar-Covalent
0%
Introduction to chemical bonding
 Recall what electronegativity is:
The ability or degree of
attraction that an atom has to
electrons that are within a
bonded compound.
(see page 161)
Introduction to chemical bonding
 To determine the degree of
ionic-ness or covalent-ness you
must take each of the
electronegativities for the
elements in the compound and
subtract them.
Introduction to chemical bonding
 If difference is 0-0.4 = nonpolar
covalent
 If difference is 0.5 – 1.7 = polar
covalent
 Above 1.8 = Ionic
Ionic/Covalent Character Due to
Electronegativity Differences
3.3
Ionic
1.7
100%
50%
Polar-Covalent
0.4
0
5%
Nonpolar-Covalent
0%
Introduction to chemical bonding
 Sulfur + Hydrogen
2.5 - 2.1 = 0.4
NonPolar Covalent
 Sulfur + Cesium
2.5 - 0.7 = 1.8
Ionic
 Sulfur + Chlorine
2.5 – 3.0 = 0.5
Polar Covalent
Introduction to chemical bonding
 In general however…
If bonding elements are on
opposite sides of the periodic
table (metal with a nonmetal)
then they tend to be ionic.
If elements are close together
(nonmetal to nonmetal), then
they tend to be covalent.
Section 2:
 Covalent Bonding &
Molecular Compounds
Covalent Bonding
 What is a molecule?
A neutral group of atoms that are
held together by covalent bonds.
 May be different atoms such as
H2O or C6H12O6
 May be the same atoms such as O2
Covalent Bonding
 Molecular compounds are made of
molecules ….. Not ions!
 We represent covalent or molecular
compounds by chemical formulas
that show numbers of atoms of
each kind of element in the
compound. CH4 - methane
Covalent Bonding
 Diatomic molecules are those
elements that exist in pairs of like
atoms that are bonded together.
 There are 7 diatomic molecules:
H2 N2 O2 F2 Cl2 I2 Br2
 Big 7
Covalent Bonding
Formation of a covalent bond:
 When atoms are far apart they do
not attract – potential energy is
zero.
 As they come closer the electrons
are attracted to protons but
electrons and electrons repel – but
e- to p attraction is stronger!
Covalent Bonding
 The electron clouds of the
bonded atoms are overlapped
and form a “bond length.”
Covalent Bonding
 Energy is released when these
atoms join together with a bond.
 Energy must be added to separate
these atoms into neutral isolated
atoms – called bond energies.
 Bond energy is expressed in
kilojoules per mole.
Covalent Bonding
 Octet Rule – Atoms will either gain,
lose, or share electrons so that their
outer energy levels will contain eight
electrons (H is an exception since it
can only have 2 in the outer level).
 These electrons that are being
gained, lost, or shared are
represented by using the electron
dot diagrams.
Examples of
electron dot notations
 1 valence electron
 3 valence electrons
X
X
 5 valence electrons
 7 valance electrons
X
X
Covalent Bonding
 Shared electron pairs and unshared
pairs:
Cl:ClShared pair
Unshared pairs
Covalent Bonding
 These electron dot representations
are called Lewis structures.
 Dots represent the valence
electrons
Covalent Bonding
 Lewis structures can also be
represented using structural
formulas.
 Dashes indicate bonds of shared
electrons (unshared e- are not shown
Cl - Cl
 One pair (2 e-) is shared here.
Steps To Drawing Lewis
Structures
 Calculate the number of valance
electrons.
 Arrange atoms.
 Compare number of electrons used with
number of electrons available.
 Check octet rule.
 Change dots to dashes where
appropriate.
Covalent Bonding
Lewis structure for ammonia
(NH3)
Covalent Bonding
Practice:
 Draw
Lewis structure for
methane CH4
 Ammonia NH3
 Hydrogen Sulfide H2S
 Phosphorus trifluoride PF3
More Guidelines
 H and halogen atoms usually bond to
only one other atom in a molecule and
are usually on the outside or end of a
molecule (each only need 1 electron to
form stable octet and electronegativity)
More Guidelines
 The atom with the smallest electro-
negativity is often the central atom
 When a molecule contains more atoms
of 1 element than the other, these atoms
often surround the central atom
Covalent Bonding
Some atoms can form multiple
bonds – especially C, O, & N.
Double bonds are bonds that
share 2 pair of electrons
C=C means C::C
Triple bonds share 3 pair
C≡C means C:::C
Covalent Bonding
Resonance:
Some substances cannot be
drawn correctly with Lewis
structure diagrams
Some electrons share time with
other atoms – ex. Ozone – O3
Covalent Bonding
Electrons in ozone may be
represented as: O = O–O
Other times it may be
represented as O–O=O
Actually these structures are
shared – electrons “resonate”
(go back & forth) between them
Section 3:
 Ionic Bonding and
Ionic Compounds
Section 3:
Ionic Bonding & Compounds
 Ionic compounds are formed
of positive and negative ions
 When combined these
charges equal zero
Ex: Na = 1+
0 charge
Cl = 1-
Section 3:
Ionic Bonding & Compounds
 The electrostatic attraction
between ions
Opposites attract! But why???
Paula Abdul said so!
Section 3:
Ionic Bonding & Compounds
Coulomb’s Law
E = (2.31X10-19 J•nm) (Q1 Q2/ r)
Q1 = charge of cation
Q2 = charge of anion
r = distance between ions
What is the sign on E? Why?
Section 3:
Ionic Bonding & Compounds
 Crystals in ionic compounds
exist in orderly arrangements
known as a crystal lattice.
Section 3:
Ionic Bonding & Compounds
 Ionic substances are usually
solids
 Ionic solids are generally
crystalline in shape
 An ionic compound is a 3-D
network of + and – ions that
are attracted to each other
Section 3:
Ionic Bonding & Compounds
 Ionic substances are not
referred to as “molecules”
 Ionic substances are
referred to as “formula
units”
 A formula unit is the
simplest ratio of the ions
that are bonded together.
Section 3:
Ionic Bonding & Compounds
 The ratio of ions depends on
the charges.
 What would result when Fcombines with Ca2+?
CaF2
Section 3:
Ionic Bonding & Compounds
 When ions are written using
electron dot structures the
dots are written and
symbols for their charges.
 Na.  Na+
 Cl

-
Compared to molecular
compounds, ionic compounds:
 Have very strong attractions
 Are hard, but brittle
 Have higher melting points and
boiling points
 When dissolved or in the molten
state they will conduct
electricity
Polyatomic Ions:
 A group of atoms covalently




bonded together but with a
charge.
Sulfate
SO42Carbonate CO32Nitrate
NO3Ammonium
NH4+
Section 4:
 Metallic Bonding
Metallic Bonding
Metals are excellent electrical
conductors in the solid state.
This is due to highly mobile
valence electrons that travel
from atom to atom.
e-
Metallic Bonding
Generally metals have either 1
or 2 s electrons
p orbitals are vacant
Many are filling in the d level
Electrons become delocalized
and move between atoms (sea
of electrons)
Metallic Bonding
A metallic bond is the mutual
sharing of many electrons
among many atoms.
Metallic Properties
 High electrical conductivity
 High thermal conductivity
 High luster
 Malleable (can be hammered or
pressed into shape)
 Ductile (capable of being drawn or
extruded through small openings to
produce a wire)
Metallic Bond Strength
Varies with nuclear charge and
number of electrons shared.
High bond strengths result in
high heats of vaporization (when
metals are changed into
gaseous phase)
Section 5:
 Molecular Geometry
Molecular geometry…
A molecule’s properties depend
on bonding of atoms, but also
the molecular geometry.
Molecular geometry…
 Is the three dimensional
arrangement of a
molecule’s atoms in space.
VSEPR Theory
 Valence Shell Electron Pair
Repulsion
 Electrons around a nucleus
repel each other to be as far
away from each other as
possible.
VSEPR Theory
 Types of e- Pairs


Bonding pairs - form bonds
Lone pairs - nonbonding e-
Lone pairs repel
more strongly than
bonding pairs!!!
 Draw the Lewis Diagram.
 Tally up e- pairs on central atom.

double/triple bonds = ONE pair
 Shape is determined by the # of bonding
pairs and lone pairs.
Know the common shapes
& their bond angles!
Common Molecular Shapes
2 total
2 bond
0 lone
BeH2
LINEAR
180°
Common Molecular Shapes
3 total
3 bond
0 lone
BF3
TRIGONAL PLANAR
120°
Common Molecular Shapes
4 total
4 bond
0 lone
CH4
TETRAHEDRAL
109.5°
Common Molecular Shapes
4 total
3 bond
1 lone
NH3
TRIGONAL PYRAMIDAL
107°
Common Molecular Shapes
4 total
2 bond
2 lone
H2O
BENT
104.5°
Examples
 PF3
4 total
3 bond
1 lone
F P F
F
TRIGONAL
PYRAMIDAL
107°
Examples
 CO2
2 total
2 bond
0 lone
O C O
LINEAR
180°
Hybridization
 Explains how atom’s orbitals
become rearranged to form
covalent bonds.
 Hybridization is the mixing of 2 or
more orbitals of similar energies on
the same atom to produce new
orbitals of equal energies.
Hybridization
 Methane (CH4) is an example of
hybridization:
 Carbon’s normal configuration is
2s22p2
 In methane all the electrons in the
2nd energy level become equal in
energy and is referred to as sp3
Intermolecular Forces
Intermolecular forces are attractive forces between molecules.
Intramolecular forces hold atoms together in a molecule.
Intermolecular vs Intramolecular
•
41 kJ to vaporize 1 mole of water (inter)
•
930 kJ to break all O-H bonds in 1 mole of water (intra)
“Measure” of intermolecular force
Generally,
boiling point
intermolecular
melting point
forces are much
weaker than
intramolecular
forces.
11.2
Intermolecular Forces:
 Strong IM forces exist in polar
molecules.
 Polar molecules act as tiny
“dipoles” (equal & opposite
charges separated by short
distances)
Types of Intermolecular Forces
Dipole Forces
Attractive forces between polar molecules
Orientation of Polar Molecules in a Solid
11.2
Types of IMF
 Dipole-Dipole Forces
-
+
View animation online.
Types of Intermolecular Forces
Dipole Forces
Attractive forces between an ion and a polar molecule
Ion-Dipole Interaction
11.2
Intermolecular Forces:
 Another IM force is Hydrogen
bonding.
 Is the strongest type of dipoledipole force
 Explains high boiling points of
H-containing substances such
as water and ammonia
Intermolecular Forces:
 In hydrogen bonding, a
hydrogen atom is attracted to
an unshared pair of electrons
of an electronegative atom in a
nearby molecule.
Types of Intermolecular Forces
Hydrogen Bond (strongest)
The hydrogen bond is a special dipole-dipole interaction
between the hydrogen atom in a polar N-H, O-H, or F-H bond
and an electronegative O, N, or F atom. IT IS NOT A BOND.
A
H…B
or
A
H…A
A & B are N, O, or F
11.2
Types of IMF
 Hydrogen Bonding
Intermolecular Forces:
London/ Dispersion/VanDerWaals
forces:
 Are very weak bonds
 Occur due to the fact that since
electrons are in constant motion
that briefly there are moments
where electrons are unevenly
distributed and thus the molecule
briefly has a charged area.
Types of IMF
 London Dispersion Forces
View animation online.