CHEMICAL BONDING

What is a chemical bond?

• A chemical bond is a strong
electrostatic attraction between
the valence electrons of atoms.

• Valence electrons are those in

the highest energy level of an
atom.
 What is a valence orbital?
• To describe where electrons exist in the
atom, chemists created the concept of an
orbital.
• An orbital refers to a region of space
around an atom’s nucleus where an electron
may exist.
• A valence orbital is the outer orbital that
houses the valence electrons.
 A further look at valence electrons
• Valence electrons can be classified as
bonding electrons or lone pairs.
• Bonding electrons are single electrons
found in the valence orbital.
• Lone pairs are paired electrons found in the
valence orbital.
• Bonding electrons are those that generally
participate in bond formation.
• The number and occupancy of orbitals in
an element are determined by the
following theoretical rules:
1. Each orbital can hold a max of 2
electrons.
2. The 1st energy level has only one orbital
(max of 2 electrons).
3. The 2nd and 3rd energy level have 4 orbitals
each (max of 8 electrons per energy
level).
4. Electrons within an energy level occupy
each orbital singly before pairing up.
 The Octet or Duet Rule
• Noble gases have a full valence shell (2 e -
for He and 8 e- for all other noble gases)
and as a result are generally unreactive with
other elements because of this special
stability.
• The Duet Rule applies to the first energy
level – meaning for hydrogen only.
• The Octet Rule applies to all other atoms in
the other energy levels.
 Lewis Symbols
• Lewis symbols or electron-dot diagrams
consist of:
– a chemical symbol that represents the nucleus
and core electrons of an atom
– dots placed around the symbol to represent the
valence electrons.
 To construct a Lewis symbol:
• Write the atomic symbol for the atom – this
represents the nucleus and core electrons.
• Start by placing a single valence electron
(using a dot) on each side of the symbol
first – each one of the four orbitals must be
filled first.
• Pairing occurs with the 5th electron up to a
maximum of 8 electrons (octet rule).
• Note that some elements do not follow the
duet or octet rule i.e. B, Be, Al
 Draw Lewis symbols of :
 Elements of Period 3

 Give the number of bonding

electrons and lone pairs for each
 Types of Chemical Bonds
• There are 3 types of chemical bonds. Each type of
bond gives rise to distinctive physical properties
of the substance formed:
1. Ionic bonds
2. Covalent bonds
3. Metallic bonds
• The type of bond that atoms form depends on the
attraction for electrons of the atoms involved –
their electronegativity.
 What is Electronegativity?
• Electronegativity is a measure of an
atom’s attraction for electrons
• A high electronegativity means that the
element has a high attraction for
electrons
• In general, electronegativities increase
diagonally from the lower left (Cs) to
the upper right (F) of the periodic table
The difference in the electronegativity
value for 2 elements can describe the
nature of a bond:
Electronegativity Type of Bond Description
Difference
> or = 1.7 IONIC Transfer of electrons
b/w a metal and a
nonmetal to form ions.

< 1.7 COVALENT Electrons are shared

between unlike atoms.
 IONIC BONDING
• Ionic bonds are the result of electrostatic
attractions between cations and anions.
• When an ionic bond is formed the metal
atom transfers electrons to the nonmetal
atom so as to obtain a stable electron shell
structure of the nearest noble gas.
• Example:
• The + and - ions are
held together by a
strong electrostatic
force forming a 3-D
crystal lattice in which
each anion is
surrounded by a cation
and vice versa such
that the net charge is
zero.
Properties of ionic compounds include:

• solid at room temperature

• form electrolytic solutions when
aqueous or molten (solutions conduct
electricity)
• solids are hard
• high melting and boiling points
 Ionic Crystals
• Ionic compounds are abundant in nature.
• Soluble ionic compounds are present in both fresh
water and salt water.
• Ionic compounds with low solubility make up
most rocks and minerals.
• Examples
– table salt (NaCl + KI)
– antacids (Mg(OH)2, CaCO3)
– home cleaning products (NaOH)
– rust (Fe(OH)3)
– lime (CaO)
– tarnish formed on silver metal (Ag2S).
• Ionic compounds containing polyatomic ions, such
as Na2CO3, also have some covalent character to
them.
• The bonds between C and O atoms in the
polyatomic ion are covalently bonded but the
overall charge on the group causes it to act like a
monoatomic ion.
• Crystal lattices for such ionic compounds are more
complex, and the ranges of hardness and melting
and boiling points vary much more widely.
• Some of these compounds will decompose before
they get hot enough to melt because the covalent
bonds in the polyatomic ions begin to break before
the ionic bonding forces in the crystal are
overcome.
 Metallic Bonding
• Metallic bonds hold pure metals together.
i.e. Cu, Ag, Al
• Metallic bonds result from the electrostatic
attractions between the freely moving
electrons (delocalized electrons) of the
atoms making up the metal and the resulting
positively charged ions.
A model for a metallic bond:

• Each metal will consist of closely packed cations

surrounded by mobile electrons.
• This model also explains why metals are ductile and
malleable.
•When a metal is subjected to pressure, the metal cations
easily slide past one another, preventing shattering of the
metal.
 Properties of Metals
• Shiny
• Flexible
• Good conductors of electricity and heat
• Hardness varies from soft to hard
• Melting points vary from low to high as
well
 Covalent Bonding
• A covalent bond results from the electrostatic attraction
between the electrons of one atom and the nucleus of
another atom and vice versa.
• When 2 atoms approach each other, the nucleus of each
atom attracts the electrons of the other atom. The result is
a mutual sharing of electrons between the 2 nuclei.

• Depending on the number of pairs of electrons shared by 2

atoms, covalent bonds may be classified as single, double,
or triple bonds.
• A single covalent bond consists of a shared
pair of electrons between two atoms as in
the hydrogen molecule: H : H
• A double covalent bond consists of two
shared pairs of electrons between two atoms
as in the oxygen molecule: :O :: O:
• A triple covalent bond consists of three
shared pairs of electrons between two atoms
as in the nitrogen molecule: :N:::N:
Properties of Molecular Compounds

• Solids, liquids, or gases at room

temperature
• Low melting and boiling points
• Non-conductors of electricity both in the
pure form and in solution
• Solids are generally soft
 Non-polar Covalent Bonds
• In diatomic covalent compounds such as H2 or N2, where
the two atoms are identical, the electrons making up the
covalent bond will be shared equally between the two
atoms.
• Electronegativity difference is zero
• This type of covalent bond is called nonpolar.
 Polar covalent bonds
• In covalent compounds such as HF or CO, where the two
atoms have a different electronegativity, the electrons
making up the covalent bond will not be shared equally
between the two atoms.
• Instead, the more electronegative atom will tend to attract
the electrons closer to its nuclei, resulting in a polar
covalent bond:
Chemical Bond = competition for bonding electrons to gain a
stable octet (full outer energy level)
Atoms with equal EN = electrons shared equally = non-polar covalent
(EN difference 0 or < 0.4)
Atoms with unequal EN = polar covalent bond (EN difference < 1.7)
Atoms with unequal EN = ionic bond (EN difference >1.7)
• There are 2 ways of symbolizing polar
covalent bonds:
1. Using a delta symbol with a positive and
negative sign.

2. Drawing a bond dipole arrow with the tip

pointing the more electronegative atom.
H Cl
 Coordinate Covalent Bonds
• An atom with a lone pair of electrons, such as N or P,
provides BOTH electrons to an atom that has been
stripped of its electrons, such as a H+

• Example: formation of ammonium ion, (NH4+)

H+ + :NH3 [H NH3]+

• Example: formation of a hydronium ion, (H3O+)

 Network Covalent Bonds
• Network covalent solids
consist of a crystal where
all the atoms are bonded
to each other by a network
of covalent bonds
• In diamond, each carbon
is covalently bonded to 4
other carbon atoms.
• Other examples include
graphite and quartz (SiO2).
• Properties of network covalent solids
include:
– High melting and boiling points
– Insoluble in most solvents
– Form nonconducting solutions
– Form very hard solids

• The overall bonding in a network covalent

solid is very strong – stronger than most
ionic bonding.
 Other Covalent Networks of
Carbon
• Carbon is an extremely versatile atom in
terms of its bonding and structures.
• Carbon can bond to itself to form a variety
of pure carbon substances:
– 3-D tetrahedral arrangements (diamond)
– Layer of sheets (graphite)
– 60-atom spherical molecules (buckyballs)
– Long, thin tubes (carbon nanotubes)
Fullerene (buckyball)
Carbon Nanotube
Graphite
• Graphite is unlike most covalent crystals in that it readily
conducts electricity.
• It also acts as a lubricant.
• This is explained by its structure; hexagonal sheets of C
atoms.
• Within the planar sheets, the bonding is a covalent network
but between the sheets the bonding is relatively weak.
• The lubricating property of graphite can be explained by
the sliding of the sheets over each other.
• The electrical conductivity of graphite is explained by the
concept that each C atom only has 3 covalent bonds. The
unbonded 4th electron of each atom is free to move through
the space between the 2-D sheets of atoms.