CHEMICAL
BONDING AND
MOLECULAR
STRUCTURE.
FROM : EXCEED
EVERYTHING
IMPORTANT TOPICS :-Ionic vs. covalent bonding
• Valence electron & Lewis dot structures

octet vs. non-octet

resonance structures

formal charges
• VSEPR - predicting shapes of molecules
• Bond properties

polarity, bond order, bond strength
•
CHEMICAL BONDS



Atoms or ions are held together in molecules or
compounds by chemical bonds.
The type and number of electrons in the outer
electronic shells of atoms or ions are instrumental
in how atoms react with each other to form stable
chemical bonds.
Over the last 150 years scientists developed
several theories to explain why and how elements
combine with each other.
Bonding in Chemistry


Central theme in chemistry: Why and How atoms
attach together
This will help us understand how to:
1.
2.
3.
Predict the shapes of molecules.
Predict properties of substances.
Design and build molecules with particular sets of chemical
and physical properties.
CHEMICAL BONDS
Two of the most common substance on our
dining table are salt and granulated sugar
NaCl
C12H22O11
The properties of substances are determined in large part
by the chemical bonds that hold their atoms together
Chemical Bonds
All chemical reactions involve breaking of some
bonds and formation of new ones which yield new
products with different properties.
Bonding Theories
Lewis
bond Theory
Valence Bond Theory
Molecular Orbital Theory
Gilbert Newton Lewis
Lewis Bonding Theory
 Atoms
ONLY come together to produce a more
stable electron configuration.
 Atoms bond together by either transferring or
sharing electrons.
 Many of atoms like to have 8 electrons in their
outer shell.
 Octet
rule.
 There are some exceptions to this rule—the key to
remember is to try to get an electron configuration
like a noble gas. Li and Be try to achieve the He
electron arrangement.
Lewis Symbols of Atoms
 Uses
symbol of element to represent nucleus
and inner electrons.
 Uses dots around the symbol to represent
valence electrons.
 Puts
one electron on each side first, then pair.
 Remember
that elements in the same
group have the same number of valence
electrons; therefore, their Lewis dot
symbols will look alike.
Li•
Be•
•
•B•
•
•
•C•
•
••
•N•
•
••
•O:
•
••
:F:
•
••
:Ne:
••
Valence electrons
Practice to write the Lewis symbol for
Arsenic
 As
is in group 15 (5), therefore it has 5 valence
electrons.

 As 

Examples for Lewis representation
of some chemical bonds
••
••
••
••
••
••
H
••
•O
••
•
F
O
•H
••
•O
••
•
F
H
••
••
F
••
H• • O
••
••
•
••
•F
••
••
••
••
F•
••
••
F
••
O ••
•• O
O
O
Example:
Write the Lewis structure
of CO2.
Information:
Given: CO2
Find: Lewis structure
Solution Map: formula → skeletal →
electron distribution → Lewis
Total number of valence electrons = 6 + 4 + 6 = 16
 Actually 24 electrons needed for completing the octet of each
atom
 Thus 24 - 16 = 8 electrons are shared.
 Since two electrons make a bond, the molecule should have 4
bonds.
 The remaining 8 electrons are lone pair electrons.

..
O
..
C
..
O
..
Example NO3─
-
Write skeletal structure.

N is central because it is the most
metallic.
Count valence electrons.
N=5
O3 = 3 x 6 = 18
(-) = 1
Total = 24 e-
TYPES OF CHEMICAL BONDS
Ionic
bonds
Covalent
Metallic
bonds
bonds
Ionic compounds consist of a cation and an anion
• the formula is always the same as the empirical formula
• the sum of the charges on the cation and anion in each formula
unit
must equal zero. Lewis bonding theory is able to
explain ionic bonds very well.
The ionic compound NaCl
Ionic bonding
Ionic
substances are formed when an
atom that loses electrons relatively
easily react with an atom that has a
high affinity for electrons.
ex. metal-nonmetal compound
Chemical Bonds
Ionic bonds are formed by the attraction of oppositely
charged ions.
Ionic Bonds
 Metal
to nonmetal.
 Metal loses electrons to form cation.
 Nonmetal gains electrons to form anion.
 The electronegativity between the metal and
the nonmetal must be > than 2.
 Ionic bond results from + to − attraction.
 Larger
charge = stronger attraction.
 Smaller ion = stronger attraction.
 Lewis
theory allows us to predict the correct
formulas of ionic compounds.
Ions that pack as spheres in a very regular
pattern are called crystalline substances .
Formation of an Ionic Solid

1. Sublimation of the solid metal
M(s) → M(g)

[endothermic]
2. Ionization of the metal atoms
M(g) →M+(g) + e- [endothermic]

3. Dissociation of the nonmetal
1/2X2(g) → X(g) [endothermic]
Electron affinity
of F
Dissociation of F2
Ionization of Li
Sublimation of Li
Formation
of solid
What is Lattice Energy?

A measure of the energy contained in
the crystal lattice of a compound, equal
to the energy that would be released if
the component ions were brought
together from infinity.
More Gains and Losses
• Can elements lose
or gain more than
one electron?
• The element magnesium, Mg, in Group 2
can lose two electron and element oxygen
in Group 6 can gain two electrons to form
stable Nobel gas configurations. The ions
can come together to form a crystal
structure.
Relative sizes of some ions and their
parent atoms.
Structure of ionic crystals
Different types of
crystals are formed
depending on the
ionic radii and the
charge of the ions
involved.
Convalent Bonds—Sharing
• Some atoms are unlikely to lose or gain
electrons because the number of electrons
in their outer levels makes this difficult.
• Consider the Lewis dot structure of carbon
.
. C. .
C+4 + 4e-
• The alternative is sharing electrons.
Covalent Bonds
 Often
found between two nonmetals.
 Typical of molecular species.
 Atoms bonded together to form molecules.
 Strong
attraction.
 Atoms
share pairs of electrons to attain octets.
 Molecules generally weakly attracted to each other.
 Observed
physical properties of molecular substance due to
these attractions.
Covalent Bonding
Electron
are shared by nuclei
The Convalent Bond
• Shared electrons are attracted to the nuclei
of both atoms.
• They move back and forth between the
outer energy levels of each atom in the
covalent bond.
• So, each atom has a stable outer energy
level some of the time.
The formation of a bond between
two atoms.
An electron density plot for the H2
molecule shows that the shared electrons
occupy a volume equally distributed over
BOTH Hydrogen atoms.
Electron Density for the H2 molecule
Chemical Bonds
Covalent bonds form when atoms share 2 or more
valence electrons.
Covalent bond strength depends on the number of
electron pairs shared by the atoms.
single
bond
<
double
bond
<
triple
bond
Examples of Convalent Bond
• The neutral particle is formed when atoms
share electrons is called a molecule
Single Covalent Bonds
 Two

atoms share one pair of electrons.
2 electrons.
 One
atom may have more than one single
bond.
••
F
••
••
F
H•
H
•H
O H
••
F
••
••
F
••
••
••
••
• F
••
••
••
•O
••
••
•
F •
••
••
••
••
Double Covalent Bond
 Two
4
atoms sharing two pairs of electrons.
electrons.
 Shorter
and stronger than single bond.
••
•O
••
•
•
••
•O
••
O ••
•• O
O
O
Chemical Bonds
Bond Polarity
 Bonding
between unlike atoms results in unequal sharing
of the electrons.


One atom pulls the electrons in the bond closer to its side.
One end of the bond has larger electron density than the
other.
 The

result is bond polarity.
The end with the larger electron density gets a partial
negative charge and the end that is electron deficient gets a
partial positive charge.
d+ H •• Cl d-
Nonpolar and polar covalent bonds
Probability representations of the
electron sharing in HF.
Trends in electronegativity across a
period and down a group
In practice no bond is totally ionic. There will
always be a small amount of electron sharing.
Percent ionic character of chemical bonds
as a function of electronegativity
difference
Bond Polarity and Dipole
Moments
Dipole Moment
μ=QR
Q: center of charge of
magnitude
R: distance
Dipole Moment of HF
1D=3.336×10-30 coulomb meter
μ=(1.6×10-19 C)(9.17×10-11 m)=1.47×10-29
=4.4 D for fully ionic
Measured dipole moment=1.83 D
1.83×3.336×10-30=δ(9.17×10-11)
δ=6.66×10-20
Ionic character=1.83/4.4=41.6%
Polar Molecules and Electric Field
Molecular Geometry
 Molecules
are three-dimensional objects.
 We often describe the shape of a molecule with terms that
relate to geometric figures.
 These geometric figures have characteristic “corners” that
indicate the positions of the surrounding atoms with the central
atom in the center of the figure.
 The geometric figures also have characteristic angles that we
call bond angles.
Valence Shell Electron Pair
Repulsion (VSEPR Model)
 It
is used to predict the geometries of molecules
formed from nonmetals.
 Postulate: the structure around a given atom is
determined principally by minimizing electron
pair repulsion.
 The bonding and nonbonding pairs should be
positioned as far apart as possible.
Predicting a VSEPR Structure




Draw Lewis structure.
Put pairs as far apart as possible.
Determine positions of atoms
from the way electron pairs are
shared.
Determine the name of molecular
structure from positions of the
atoms.
For non-metals compounds, four pairs of
electrons around a given atom prefer to
form a tetrahedral geometry
 Draw
the Lewis structure
 Count the pairs of electrons and arrange them to minimize
repulsions
 Determine the positions of the atoms
 Name the molecular structure
Lone
pairs require more space than
bonding pair.
The bonding pairs are increasingly
squeezed together as the number of lone
pairs increases.
 The
bonding pair is shared between two
nuclei; and the electrons can be close to
either nucleus.
 A lone pair is localized on only one nucleus,
so both electrons are close to that nucleus
only.
Molecular Geometries
Electron Pairs
Practice drawing these shapes below
Linear
TP
Tetra
TBP
Octa
Polarity of Molecules

In order for a molecule to be polar it must:
1. Have polar bonds.


Electronegativity difference—theory.
Bond dipole moments—measured.
2. Have an unsymmetrical shape.


Vector addition.
Polarity effects the intermolecular forces of attraction.
Molecule Polarity
The O—C bond is polar. The bonding
electrons are pulled equally toward both O
ends of the molecule. The net result is a
nonpolar molecule.
Molecule Polarity
The H—O bond is polar. Both sets
of bonding electrons are pulled
toward the O end of the molecule.
The net result is a polar molecule.
Water molecule behaves as if it had a
positive and negative end.
The Covalent Chemical Bond
Bond Energies

Bond breaking requires energy (endothermic).

Bond formation releases energy (exothermic).

DH = SD(bonds broken) - SD(bonds formed)
energy required
energy released
Bond Energies
Covalent Bond Energies and
Chemical Reactions
H2+F2→2HF
ΔH=ΣD (bonds broken)-ΣD (bonds formed)
ΔH=DH-H+DF-F-2DH-F=1×432+1×154-2×565
=-544 kJ
Bond Energy of CH4
Experimental result : 1652 kJ/mol
C(g)+4H(g) →CH4(g) + 1652 kJ/mol
An average C-H bond energy per mole
of C-H bond: 1652/4=413 (kJ/mol)
Metallic Bonding
 The
model of metallic bonding
can be used to explain the
properties of metals.
 The luster, malleability, ductility,
and electrical and thermal
conductivity are all related to the
mobility of the electrons in the
solid.
 The strength of the metallic bond
varies, depending on the charge
and size of the cations, so the
melting points and DHfusion of
metals vary as well.
IONIC COMPOUNDS vs METALS
BREAKING INORGANIC MATERIAL
SLIP PLANES
ALLOY vs PURE METAL