Chemical bonding and molecular structure

• Know about valence electrons
• Learn various types of bond and bond parameters
• Lewis Structure
• Understand polar and covalent characters of
covalent and ionic bonds
• Concept of hybridisation
• Study of VSEPR and Molecular Orbital Theory

1. Atomicity of a gas: The number of atoms present in the molecule of a
gas is called its atomicity.
2. Bond dipole moment ( µ ).:A covalent bond between two atoms of
different elements is called a polar covalent bond . A polar bond is
partly covalent bond and partly ionic. The percentage of ionicity in a
covalent bond is called percentage ionic character in that bond . The
ionic character in a bond is expressed in terms of bond dipole moment
( µ ).
3. BORN-HABER CYCLE: This thermochemical cycle was devised by Born
and Haber in 1919. It relates the lattice energy of a crystalline
substance to other thermochemical data. The Born-Haber cycle is the
application of Hess's law to the enthalpy of formation of an ionic solid
at 298 K.

4. Chemical bond : The chemical force which keeps the atoms in any
molecule together is commonly described as a chemical bond.
5. Chemical compounds :Compounds are generally called chemical
compounds because they are formed due to the chemical
combination of the combining element.
6. Covalent bond : The Bond formed by Mutual sharing of electrons
between the combining atoms of the same or different elements is
called covalent bonds.
7. Double covalent bond :The bond formed between two atoms due to
the sharing of two electron-pairs is called a double covalent bond or
simply a double bond. It is denoted by two small horizontal lines (=)
drawn between the two atoms, e.g., O = O, O = C = O etc.
8. Electronegative or nonmetallic character :The tendency of an element
to accept electrons to form an anion is called its non metallic or
electronegative character.
9. ELECTRONEGATIVITY :The relative tendency of an atom in a molecule
to attract a shared pair of electrons towards itself is termed its
electronegativity.

10. Electronic configuration:The distribution of electrons amongst various
energy levels of a atom is termed its electronic configuration
11. HYBRIDISATION :The process of mixing of the atomic orbitals to form
new hybrid orbitals is called hybridisation.
12. Hybrid orbitals :According to the concept of hybridisation, certain
atomic orbitals of nearly the same energy undergo mixing to produce
equal number of new orbitals. The new orbitals so obtained are called
hybrid orbitals.
13. Hybridisation in carbon :Carbon shows sp 3
hybridisation in alkanes, sp
2
hybridisation in alkenes and sp hybridisation in alkynes.
14. HYDROGEN BOND :The bond between the hydrogen atom of one
molecule and a more electronegative atom of the same or another
molecule is called hydrogen bond.
15. Ionic (or Electrovalent) bond : An ionic (or electrovalent) bond is
formed by a complete transfer of one or more electrons from the atom
of a metal to that of a non-metal.
16. LATTICE ENERGY (Lattice Enthalpy) :The strength of binding forces in
solids is described by the term lattice enthalpy ( ∆ L
H ) (earlier the
term lattice energy was used). The molar enthalpy change
accompanying the complete separation of the constituent particles
that composed of the solid (such as ions for ionic solids and molecules
for molecular solids) under standard conditions is called lattice
enthalpy ( ∆ H° ). The lattice enthalpy is a positive quantity.

17. Lewis Formula (or Electronic Formula) of a Compound :The formula showing
the mode of electron-sharing between different atoms in the molecule of a
compound is called its electronic formula or Lewis formula.
18. Metallic crystals : In metallic crystals, the valence electrons of all the atoms
form a pool of mobile electrons. The nuclei with their inner electrons (called
Kernels) are embedded into this pool of free electrons. Thus, the constituent
particles in a metallic crystal are the positive kernels in a pool of electrons.
19. NON-POLAR COVALENT BOND :When a covalent bond is formed between
two atoms of the same element, the electrons are shared equally between
the two atoms. In other words, the shared electron-pair will lie exactly
midway between the two atoms. The resulting molecule will be electrically
symmetrical, i.e ., centre of the negative charge coincides with the centre of
the positive charge. This type of covalent bond is described as a non-polar
covalent bond. The bonds in the molecules H 2
, O 2
, Cl 2
etc., are non-polar
covalent bonds.
20. OCTET RULE : According to this theory, the atoms tend to adjust the
arrangement of their electrons in such a way that they ( except H and
He ) achieve eight electrons in their outermost shell. This is known as
the octet rule .
21. Pi ( π ) Bond : A covalent bond formed between the two atoms due
to the sideways overlap of their p -orbitals is called a pi ( π ) bond.

22. POLAR COVALENT BOND :When a covalent bond is formed between two
atoms of different elements, the bonding pair of electrons does not lie
exactly midway between the two atoms. In fact, it lies more towards the
atom which has more affinity for electrons. The atom with higher affinity
for electrons, thus, develops a slight negative charge, and the atom with
lesser affinity for electrons a slight positive charge. Such molecules are
called polar molecules. The covalent bond between two unlike atoms
which differ in their affinities for electrons is said to be a polar covalent
bond.
23. RESONANCE : When a molecule is represented by a number of electronic
structures such that none of them can exactly describe all the properties
of the molecule, but each structure has a contribution to it, then the
molecule is termed as a resonance hybrid of all these structures. Such
structures are called resonance structures and such a phenomenon is
called resonance.
24. Resonance hybrid :When a molecule is represented by a number of
electronic structures such that none of them can exactly describe all the
properties of the molecule, but each structure has a contribution to it,
then the molecule is termed as a resonance hybrid of all these structures.

25. Sigma ( σ ) Bond : A covalent bond formed due to the overlap of orbitals of
the two atoms along the line joining the two nuclei (orbital axis) is called
sigma ( σ ) bond.
26. Single Covalent Bond : A covalent bond formed by mutual sharing of one pair
of electrons is called a single covalent bond, or simply a single bond. A single
covalent bond is represented by a small line (−) between the two atoms.
27. Triple covalent bond : Bond formed due to the sharing of three electron-
pairs is called a triple covalent bond or simply a triple bond.
28. Valence electrons : Valence is one of the most important chemical property
of the elements. The chemical behaviour of an element depends upon the
number of electrons in the outermost shell of its atom. The electrons present
in the outermost shell are called valence electrons. The electrons in the
outermost shell are called valence electrons because the electrons in the
outermost shell determine the valence of an element .
29. Valency : The combining capacity of an atom of an element is described in
terms of its valency. It may be defined as,
The number of hydrogen or chlorine or double the number of oxygen atoms
which combine with one atom of the element is termed its valency.
It may also be defined as,
The number of electrons which an atom loses or gains or shares with other

• The electrons that are in the highest(outermost) energy level
• That level is also called the valence shell of the atom they are held
most loosely
• The number of valence electrons in an atom determines:
• The properties of the atom
• The way that atom will bond chemically
•As a rule, the fewer electrons in the valence shell, the more
reactive the element

• When an atom has eight electrons in the valence shell, it is stable
• Atoms usually react in a way that makes each atom more stable
• There are two ways this can happen:
•The number of valence electrons increases to eight
•Loosely held valence electrons are given up

• Atomic number = number of Electrons
• Electrons vary in the amount of energy
they possess, and they occur at certain
energy levels or electron shells.
• Electron shells determine how an atom
behaves when it encounters other atoms

• Electrons are placed in shells according to
rules:
• The 1st shell can hold up to two electrons,
and each shell thereafter can hold up to 8
electrons.

Octet Rule = atoms tend to gain, lose or share electrons so
as to have 8 electrons
C would like to
N would like to
O would like to
Gain 4 electrons
Gain 3 electrons
Gain 2 electrons

Group # Group Name # of valence electrons
1 Alkali Metals 1
2 Alkaline Earth Metals 2
3-12 Transition Metals 1 or 2
13 Boron Group 3
14 Carbon Group 4
15 Nitrogen Group 5
16 Oxygen Group 6
17 Halogens 7
18 Noble Gases 8

•The octet rule is a simple chemical rule of thumb
•Octet Rule says atoms with 8 electrons in their outer shell areatoms with 8 electrons in their outer shell are
stablestable
•Atoms tend to combine in such a way that they each have eight
electrons in their valence shells, giving them the same electron
configuration as a noble gas
• The rule applies to the main-group elements, especially carbon,
nitrogen, oxygen, the halogens, and also to metals such as sodium
or magnesium
•In simple terms, molecules or ions tend to be most stable when
the outermost electron shells of their constituent atoms contain 8
electrons

Doesn’t allow for
• H, He or Li
[stable with 2 e-
in their outer shells] -
Duet Rule
• Transition elements - 18 electron rule
• BF3 which only has 6 e-
in its outer shell

• named after Gilbert N. Lewis, who introduced it in his 1916
• also known as Lewis dot diagrams, electron dot diagrams,
"Lewis Dot formula" Lewis dot structures, and electron dot
structures)
• are diagrams that show the bonding between atoms of
a molecule and the lone pairs of electrons that may exist in the
molecule.
• can be drawn for any covalently bonded molecule, as well
as coordination compounds.

Symbols of atoms with dots to represent the valence-shell
electrons
1 2 13 14 15 16 17 18
H•
He:
• • • • • • • • • • •
Li•
Be• •
B • •
C • •
N • •
O •
: F •
:Ne :
• • • • • • • •
• • • • • • • • • • •
Na•
Mg• •
Al• •
Si• •
P• •
S•
:Cl •
:Ar :
• • • • • • • •

• A bond results from the attraction of nuclei
for electrons
– All atoms trying to achieve a stable octet
• IN OTHER WORDS
– the p+
in one nucleus are attracted to the e- of
another atom
• Electronegativity

• It is an exothermic process
Energy
released
E
N
E
R
G
Y
Reactants
Products

• Endothermic reaction
– energy must be put into the bond in order to
break it
E
N
E
R
G
Y
Reactants
Products
Energy
Absorbed

• Strong, STABLE bonds require lots of energy
to be formed or broken
• weak bonds require little E

Double bondDouble bond
Single bondSingle bond
TripleTriple
bondbond
AcrylonitrileAcrylonitrile

Fractional bond orders in resonance structures.
Consider NO2
-
bondsO—N2
bondsNOinpairs-e3
=orderBond
Bond order =
Total # of e- pairs used for a type of bond
Total # of bonds of that type
The N—O bond order = 1.5The N—O bond order = 1.5
O O O O
N
••
••
••
••
••
••••••
••
••
••
••
••
N

Bond order is proportional to two important
bond properties:
(a) bond strength
(b) bond length
745 kJ745 kJ
414 kJ414 kJ
110 pm110 pm
123 pm123 pm

• Bond length is the distance between
the nuclei of two bonded atoms.

Bond length depends
on bond order.
Bond distances measuredBond distances measured
using CAChe software. Inusing CAChe software. In
Angstrom units where 1 A =Angstrom units where 1 A =
1010-2-2
pm.pm.
Bond distances measuredBond distances measured
using CAChe software. Inusing CAChe software. In
Angstrom units where 1 A =Angstrom units where 1 A =
1010-2-2
pm.pm.

Using Bond Energies
Estimate the energy of the reaction
H—H + Cl—Cl ----> 2 H—Cl
Net energy = H∆ rxn =
= energy required to break bonds
- energy evolved when bonds are made
H—H = 436 kJ/molH—H = 436 kJ/mol
Cl—Cl = 242 kJ/molCl—Cl = 242 kJ/mol
H—Cl = 432 kJ/molH—Cl = 432 kJ/mol

Estimate the energy of the reaction
H—H + Cl—Cl ----> 2 H—ClH—H = 436 kJ/molH—H = 436 kJ/mol
Sum of H-H + Cl-Cl bond energies = 436 kJ +Sum of H-H + Cl-Cl bond energies = 436 kJ +
242 kJ = +678 kJ242 kJ = +678 kJ
Using Bond Energies
2 mol H-Cl bond energies = 864 kJ2 mol H-Cl bond energies = 864 kJ
Net = ∆H = +678 kJ - 864 kJ = -186 kJNet = ∆H = +678 kJ - 864 kJ = -186 kJ

Why do ionic compounds dissolve inWhy do ionic compounds dissolve in
water?water?
Boiling point = 100 ˚CBoiling point = 100 ˚C
Boiling point = -161 ˚CBoiling point = -161 ˚C
Why do water and methaneWhy do water and methane
differ so much in theirdiffer so much in their
boiling points?boiling points?

HCl isHCl is POLARPOLAR because itbecause it
has a positive end and ahas a positive end and a
negative end.negative end.
Cl has a greater share inCl has a greater share in
bonding electrons thanbonding electrons than
does H.does H.
Cl has slight negative chargeCl has slight negative charge (-(-δδ)) and H has slightand H has slight
positive chargepositive charge (+(+ δδ))
H Cl
••
••
+δ -δ
••H Cl
••
••
+δ -δ
••

Due to the bond polarity, the H—Cl bondDue to the bond polarity, the H—Cl bond
energy is GREATER than expectedenergy is GREATER than expected
for a “pure” covalent bond.for a “pure” covalent bond.
BONDBOND ENERGYENERGY
““pure” bondpure” bond 339 kJ/mol calc’d339 kJ/mol calc’d
real bondreal bond 432 kJ/mol measured432 kJ/mol measured
BONDBOND ENERGYENERGY
““pure” bondpure” bond 339 kJ/mol calc’d339 kJ/mol calc’d
real bondreal bond 432 kJ/mol measured432 kJ/mol measured
Difference = 92 kJ. This difference is proportional to the difference inDifference = 92 kJ. This difference is proportional to the difference in
ELECTRONEGATIVITYELECTRONEGATIVITY,, χχ..
Difference = 92 kJ. This difference is proportional to the difference inDifference = 92 kJ. This difference is proportional to the difference in
ELECTRONEGATIVITYELECTRONEGATIVITY,, χχ..
H Cl
••
••
+δ -δ
••

• Metallic bonding
– Occurs between like atoms of a metal in the
free state
– Valence e- are mobile (move freely among all
metal atoms)
– Positive ions in a sea of electrons
• Metallic characteristics
– High mp temps, ductile, malleable, shiny
– Hard substances
– Good conductors of heat and electricity as (s) and (l)

• electrons are transferred between valence
shells of atoms
• ionic compounds are
made of ions
• ionic compounds are called Salts or Crystals
NOT MOLECULES

• Always formed between metals and non-
metals
[METALS ]+
[NON-METALS ]
-
Lost e-
Gained e-

• Electronegativity difference > 2.0
– Look up e-neg of the atoms in the bond and
subtract
NaCl
CaCl2
• Compounds with polyatomic ions
NaNO3

Ionic Bond, A Sea of Electrons

The formations of ionic bond governed by the following factors:
1.Ionization energy:
• Formation of ionic bond metal atom loses electron to form cation
• Energy required for this equal to ionization energy
• Alkali metals have lowest ionization energy, thus have more tendency
to form cation
2.Electron gain enthalpy:
• Electron released in the formation of cation are to be accepted by the
other atom taking part in the ionic bond formation
• Electron accepting tendencies depend on upon the electron gain
enthalpy
• Defined as energy released when isolated gaseous atom takes up an
electron to form anion.
• Greater the negative enthalpy, easier the formation of anion

3. Lattice energy:
• Combination of oppositely charged ions to form ionic crystal, with
release of energy is referred as lattice energy
• Higher value of lattice energy, greater will be the stability of
compound
• Magnitude of lattice energy gives idea about the strength of interionic
forces
• Size of ions:
• In case of similar ions inter-nuclear distance is lesser due to which
inter-ionic attraction is greater and hence the magnitude of lattice
energy will be larger
• Charge on the ions:
• Ions have higher charge exerts stronger forces of attraction and hence
larger amount of energy is released. Thus value of lattice energy is
higher

• Ionic compound exist in solid state
• The network of ions have a definite geometric pattern which
depends on the size and charge of ions
• Posses high melting and boiling points due to strong
electrostatic force of attraction between the ions
• Good conductor of electricity in molten or dissolved state
• Does not conduct electricity in solid state as ions are not free
to move
• Are soluble in polar solvent like water as solvent interacts
with the ions of ionic solid
•The chemical reactions between ionic compounds in aqueous
solution involves the combination between their ions, such
reactions are called ionic reactions.

This thermochemical cycle was devised by Born and Haber in 1919. It
relates the lattice energy of a crystalline substance to other
thermochemical data. The Born-Haber cycle is the application of
Hess's law to the enthalpy of formation of an ionic solid at 298 K.

Formation of crystalline sodium chloride form sodium metal and
chlorine gas can be described by the reaction.
Na(s) + ½ Cl2
(g) → NaCl (crystal) ∆r
H = ∆f
H° = − 411 kJ mol−1
(energy evolved)
This overall reaction can be considered to proceed in a stepwise manner as
follows
The signs of the energy involved in each step follow the rule that energy
evolved is negative and energy absorbed is positive. These steps are
summarized in Fig. 6.3.
From the Born-Haber cycle the value for any one of the steps can be
calculated if data for all the other steps are known.

• Pairs of e- are shared between non-
metal atoms
• electronegativity difference < 2.0
• forms polyatomic ions

• Between nonmetallic elements of similar
electronegativity.
• Formed by sharing electron pairs
• Stable non-ionizing particles, they are not
conductors at any state
• Examples; O2, CO2, C2H6, H2O, SiC

• Bonds in all polyatomic ions
and diatomic are all covalent
bonds
• Two types of covalent bond:
• Non-Polar Covalent Bond
• Polar Covalent Bond

when electrons are shared
equally
H2 or Cl2

Oxygen AtomOxygen Atom Oxygen AtomOxygen Atom
Oxygen Molecule (OOxygen Molecule (O2)2)

when electrons are shared but
shared unequally
H2O

- water is a polar molecule because oxygen is more
electronegative than hydrogen, and therefore electrons
are pulled closer to oxygen.

•Covalent bond can be represented by line drawn
between symbols of element whose atoms are
involved in sharing of electron
•Single Bond : Sharing single electron pairs
between two atoms
•Double Bond : Sharing two electron pairs
between two atoms
•Triple Bond : Sharing three electron pairs
between two atoms

Hydrogen (H2) has a single bond between atoms. Oxygen (O2)
has a double bond between atoms, indicated by two lines (=).
Nitrogen (N2) has a triple bond between atoms, indicated by
three lines ( ). Each bond represents an electron pair.≡

• Compounds formed exist as discrete molecules
•Weak intermolecular force due to small molecular size
•Mainly exist in liquid or gaseous state
•Sugar, urea, starch etc. exist in solid state
•Low melting and Boiling points due to weak attractive forces
•Poor conductor of electricity in fused or dissolved state
•Less soluble in water
•Gives molecular reaction

VSEPR (Valence Shell Electron Pair Repulsion) Theory is
a model for understanding and predicting the shape of
molecules.
This theory is based on the principle that electrons will
spread out as far as possible from each other as a result
of their mutual repulsion.

We refer to the electron
pairs as electron domains.
A double or triple bond
counts as one electron
domain.
This molecule has four
electron domains.

These are the electron-
domain geometries for two
through six electron domains
around a central atom.

The molecular geometry is defined by the positions
of only the atoms in the molecules, not the
nonbonding pairs (lone pairs).

Within each electron domain,
there may be several
molecular geometries.

In this domain, there is only one molecular geometry:
linear.
NOTE: If there are only two atoms in the molecule, the
molecule will be linear no matter what the electron
domain is.

There are two molecular geometries:
• Trigonal planar, if all the electron domains are bonding
• Bent, if one of the domains is a nonbonding pair

There are three molecular geometries:
• Tetrahedral, if all are bonding pairs
• Trigonal pyramidal if one is a nonbonding pair
• Bent if there are two nonbonding pairs

There are four distinct
molecular geometries:
• Trigonal bipyramidal
• Seesaw
• T-shaped
• Linear

In the trigonal bipyramidal
electron domain geometry,
the electron domains have
two distinct positions:
• Axial
• Equatorial

Nonbonding electron pairs occupy the equatorial,
rather than axial, positions in this geometry.

All positions are
equivalent in the
octahedral domain.
There are three
molecular
geometries:
• Octahedral
• Square pyramidal
• Square planar

Predicting the Molecular Shape of a Molecule:
Examples
What are the molecular shapes of CO2 and NF3?

Step 1: Draw the structural formula (using the outside
atoms to determine the bonds).
Oxygen needs
two bonds Fluorine needs
one bond

Step 2: Add in lone pairs to the central atom (first
determine how many valence electrons it has).
Carbon has four valence
electrons. Each bond includes
one of these valence electrons,
so there are none left over.
Nitrogen has five valence
electrons. Each bond includes
one of these valence electrons,
so there are two left over (one
lone pair).

Step 3: Count the total number of electron domains,
and the number of lone pairs. Use this to determine
the shape.
2 electron domains,
0 lone pairs
4 electron domains,
1 lone pair

Nonbonding electron pairs (lone pairs) have a
stronger repulsion than bonding electrons. This
causes the actual bond angles to be slightly less
than otherwise predicted.
CH4 NH3 H2O

In larger molecules,
it makes more sense
to talk about the
geometry about a
particular atom
rather than the
geometry of the
molecule as a whole.

This approach makes sense,
especially because larger
molecules tend to react at a
particular site in the molecule.

To tell if bonds are polar or nonpolar, we simply look at
the difference in electronegativity for the two atoms.
To tell if a molecule is polar or nonpolar,
we need to know its shape.

Nonpolar molecules have a symmetrical distribution of
charge (even if the bonds themselves are polar).
Polar molecules have an asymmetrical distribution of
charge (even if the bonds themselves are nonpolar).
H2O
CO2

• Cannot explain shape of very much polar
compound
• Is also unable to explain the shape of very
much polar in which delocalized p electrons
are very much
• VSEPR theory does not explain the shapes of
molecules having inert electron pair

Electronegativity is a measure of an atoms
ability to attract electrons from an atom to
which it is bonded

Mulliken defined electronegativity as:
• ½ (electron affinity + ionization energy)
• This approach yields very high electronegativity
values for He and Ne, even though they do not
form compounds

Electronegative atoms bonded to a less
electronegative central atom tend to draw electron
density away from the central atom, thus lowering
repulsion.
PFPF33 PClPCl33 PBrPBr33
97.897.8oo
100.3100.3oo
101101oo

• The dipole moment (µ) of a molecule is the product of the
magnitude of the charge (ε) and the distance (d) that separates
the centers of positive and negative charge
µ= ε∗d
• A unit of dipole moment is the debye (D)
• One debye (D) is equal to 1 x 10–10
esu Α

• A polar covalent bond has a bond dipole; a separation of positive
and negative charge centers in an individual bond
• Bond dipoles have both a magnitude and a direction (they are
vector quantities)
• Ordinarily, a polar molecule must have polar bonds, BUT … polar
bonds are not sufficient
• A molecule may have polar bonds and be a nonpolar molecule – IF
the bond dipoles cancel.

• CO2 has polar bonds, but is a
linear molecule; the bond
dipoles cancel and it has no
net dipole moment (µ = 0 D).
• The water molecule has
polar bonds also, but is an
angular molecule.
• The bond dipoles do not
cancel (µ = 1.84 D), so
water is a polar molecule. Net dipole
No net
dipole

The first quantum mechanical model to explain the nature and stability of a
covalent bond was formulated by Heitler and London 1927. This theory was
then modified by Pauling and Slater in 1931. This theory is commonly known
as the Valence-bond theory.
The main postulates of the valence bond theory are:
i. A covalent bond is formed due to the overlap of the outermost half-filled
orbitals of the combining atoms. The strength of the bond is determined by
the extent of overlap.
ii. The two half-filled orbitals involved in the covalent bond formation should
contain electrons with opposite spins. The two electrons then move under the
influence of both the nuclei.
iii. The completely-filled orbitals (orbitals containing two paired electrons) do not
take part in the bond formation.

iv. An s-orbital does not show any preference for direction. The non-
spherical orbitals such as, p- and d-orbitals tend to form bonds in
the direction of the maximum overlap, i.e., along the orbital axis.
v. Between the two orbitals of the same energy, the orbital which is
non-spherical (e.g., p- and d- orbitals forms stronger bonds than
the orbital which is spherically symmetrical, e.g., s-orbital.
vi. The valence of an element is equal to the number of half-filled
orbitals present in it.
In the valence bond model, the stability of a molecule is explained in
terms of the following types of interactions.
a. electron - nuclei attractive interactions, i.e., the electrons of one
atom are attracted by the nucleus of the other atom also.
b. electron - electron repulsive interactions, i.e., electrons of one
atom are repelled by the electrons of the other atom.
c. nucleus - nucleus repulsive interactions, i.e., nucleus of one atom
is repelled by the nucleus of the other atom.

The attractive and the repulsive interactions oppose each other.
When the attractive interactions are stronger than the repulsive
interactions, certain amount of energy is released. Due to the
lowering of energy the molecule becomes stable.
Various interactions which act between the two atoms are shown in Fig.

Valence Bond Description of Hydrogen Molecule

TYPES OF OVERLAPPING
Various types of atomic orbital overlap leading
to the formation of covalent bond are:
1. s − s overlap .
In this type of overlap, half-filled
s -orbitals of the two combining atoms overlap
each other. This is shown in Fig.

2. s − p overlap .
Here a half-filled s -orbital of one atom overlaps with one of the p
-orbitals having only one electron in it. This is shown in Fig.
3. p−p overlap along the orbital axis. This is called head on, end-on or end-
to-end linear overlap. Here, the overlap of the two half-filled p-orbitals
takes place along the line joining the two nuclei. This is shown in Fig.

4. p−p sideways overlap. This is also called lateral overlap. In this types
of overlap, two p-orbitals overlap each other along a line
perpendicular to the internuclear axis, i.e., the two overlapping p-
orbitals are parallel to each other. This is shown in Fig.
TYPES OF COVALENT BONDS: SIGMA (σ) AND PI (π) BONDS
The overlapping of orbitals is possible in two ways.
(i) along their orbital axis so that the electron density along the axis is
maximum.
(ii) along a direction perpendicular to the bond axis due to sideways
overlap of the orbitals.
Depending upon the manner in which the two atomic orbitals overlap
with each other, two types of bonds are formed. These are called, sigma
(σ) bond, and pi (π) bond.

A covalent bond formed due to the overlap of orbitals of the two atoms
along the line joining the two nuclei (orbital axis) is called sigma ( σ )
bond. For example, the bond formed due to s-s and s-p, and p-p
overlap along the orbital axis are sigma bonds, (by convention Z-axis
is taken as inter-nuclear axis.
A covalent bond formed between the two atoms due to the sideways
overlap of their p -orbitals is called a pi ( π ) bond

Sigma (σ) bond Pi (π) bond
1. It is formed due to axial overlap of the
twoorbitals. The overlap may be of s-s, s-p, p-
p orbitals.
1. This bond is formed by the lateral (sideways)
overlap of two p-orbitals.
2. There can be only one sigma bond
between atoms.
2. There can be more than one π-bond between
the two atoms.
3. The electron density is maximum and
cylindrically symmetrical about the bond
axis.
3. The electron density is high along a direcion
tion at right angle to the bond axis.
4. The bonding is relatively strong. 4. The bonding due to a π-bond is weak.
5. Free rotation of atoms about sigma (σ)
bond is possible.
5. Free rotation about a π bond is not possible.
6. It can be formed independently, i.e., there
can be a sigma (σ) bond without having a π
bond in any molecule.
6. The π bond is formed only after σ bond has
been formed.

The concept of hybridization is used to explain the nature of bonds, and
shape of the polyatomic molecules. For an isolated atom hybridization has
no meaning.
According to the concept of hybridization, certain atomic orbitals of nearly
the same energy undergo mixing to produce equal number of new orbitals.
The new orbitals so obtained are called hybrid orbitals.
The process of mixing of the atomic orbitals to form new hybrid orbitals is
called hybridization.
All hybrid orbitals of a particular kind have equal energy, identical shapes
and are symmetrically oriented in space.
The types of atomic orbitals involved in hybridization, and the nature of
hybridization depends upon the requirements of the reaction. For example,
carbon in methane (CH4
) shows sp3
hybridization. In ethene (ethylene, C2
H4
),
it exhibits sp2
hybridization. In ethyne (acetylene, C2
H2
) it shows sp
hybridization.

The hybrid orbitals are designated according to the type and the number of
atomic orbitals merging together. For example,
Mixing orbitals Hybrid orbital Hybridisation
one s and three p- four sp3
orbitals sp3
hybridisation
one s and two p- three sp2
orbitals sp2
hybridisation
one s and one p- two sp orbitals sp hybridisation
Conditions Necessary For Hybridisation
Atomic orbitals undergo hybridisation only if the following conditions are satisfied.
(i) Atomic orbitals of the same atom participate in hybridisation. Electrons present in
these atomic orbitals do not participate in the hybridisation, and occupy the hybrid
orbitals as usual.
(ii) The atomic orbitals participating in hybridisation should have nearly equal
energy.
(ii) Characteristics of Hybrid Orbitals
The characteristics of hybrid orbitals are:
i. The number of hybrid orbitals formed is equal to the number of the atomic
orbitals participating in hybridisation.

ii. All hybrid orbitals are equivalent in shape, and energy, but different
from the participating atomic orbitals.
iii. A hybrid orbital which takes part in the bond formation must contain
only one electron in it.
iv. A hybrid orbital, like atomic orbitals, cannot have more than two
electrons. The two electrons should have their spins paired.
v. Due to the electronic repulsions between the hybrid orbitals, they
tend to remain at the maximum distance from each other
Types of Hybridisation
Depending upon the nature of the orbitals involved in hybridisation,
different types of hybridisation become possible. The type of
hybridisation shown by an atom depends upon the requirements of
the reaction.
• sp3
hybridisation: In any atom, corresponding to energy levels (or
shells) for which n ≥ 2, there is one s orbital and three p orbitals. For
example, for n = 2, we have one 2s and three 2p orbitals; for n = 3, we
have one 3s, and three 3p orbitals. These four orbitals undergo mixing
to provide four new hybrid orbitals.
s + (px
+ py
+ pz
) → sp3
one s-orbital three p-orbitals four hybrid
orbitals

Hybridisation sp3
, geometry tetrahedral, bond angle 109.5’

• sp2
hybridisation: In certain reactions, one s and two p (say px
and py
)
orbitals of an atom undergo mixing to produce three equivalent sp2
hybridised orbitals. The three sp2
hybrid orbitals are oriented in a
plane along the three corners of an equilateral triangle, i.e., they are
inclined to each other at an angle of 120°. The third p-orbital (say pz
here) remains unchanged. Each hybrid orbital has 33.3% s-character
and 66.7% p-character. Formation of sp2
-hybrid orbitals from one s
and two p-orbitals is shown in Fig. 6.43.
s + (px
+ py
) → sp2
one s-orbital two p-orbitals three hybrid orbitals

Boron trifluoride has a plane trigonal
shape in which all three bonds are
identical.
Hybridisation sp2
, geometry trigonal planar, bond angle 120’

• sp hybridisation. In this type of hybridisation, one s and one p (say pz
)
orbitals belonging to the same main energy level hybridise to give two sp
hybrid orbitals. These sp hybrid orbitals are oriented at an angle of 180°
to each other. Each hybrid orbital has 50% s- and 50% p- character. The
other two p-orbitals (say 2px
and 2py
) remain unhybridised and are
oriented at right angles to each other and to the internuclear axis.
Hybridisation sp, geometry Linear, bond angle 180’

Formation of ethyne (acetylene) molecule.

•When some compounds cannot be represented by a single
definite structure rather more than one structure
•Thus, the various structure written for a compound to explain
the known properties of the compound are called as resonating
or contributing or canonical structure
•This phenomenon is called resonance
•The real structure is a resonance hybrid of all the resonating
structure
•The actual energy of the resonance hybrid and the most stable
one of the resonating structures is called resonance energy

Conditions of Resonance :
1.Different contributing structures should have the same position of
the constituent, though may have different electronic arrangement
2.The number of unpaired electrons should be the same in all
resonating structures
3.The contributing structure should have nearly the same energy
4.The bond length and bond angles should closer to the real
structure

Characteristics of Resonance :
1.Resonating structure are imaginary and do not have real existence
2.Resonance hybrid is more stable i.e. its energy is least among
different resonating forms
3.The difference between the energy of resonance hybrid and most
stable resonating form is called resonance energy.
4.Larger the value of resonance energy greater the stability of
hybrid resonance
5.Bond length in hybrid structures are intermediate of the bond
lengths in various resonating forms

This is the Lewis
structure we would
draw for ozone, O3.
-
+

• But this is at odds with
the true, observed
structure of ozone, in
which…
– …both O—O bonds are
the same length.
– …both outer oxygens
have a charge of −1/2.

• One Lewis structure
cannot accurately depict
a molecule such as
ozone.
• We use multiple
structures, resonance
structures, to describe
the molecule.

Just as green is a synthesis of
blue and yellow…
…ozone is a synthesis of
these two resonance
structures.

• In truth, the electrons that form the second C—O bond
in the double bonds below do not always sit between
that C and that O, but rather can move among the two
oxygens and the carbon.
• They are not localized, but rather are delocalized.

• The organic compound
benzene, C6H6, has two
resonance structures.
• It is commonly depicted
as a hexagon with a circle
inside to signify the
delocalized electrons in
the ring.

• A covalent bond is formed by two atoms sharing a
pair of electrons. The atoms are held together
because the electron pair is attracted by both of the
nuclei.
• In a simple covalent bond, each atom supplies one
electron to the bond - but that doesn't have to be
the case.
• A co-ordinate bond is a covalent bond (a shared pair
of electrons) in which both electrons come from the
same atom.

Properties of Coordinate bond :
1.Are generally soluble in water and organic solvents
2.Boiling and melting points of these compounds are less than electrovalent
compounds but are higher than covalent compounds
3.Compounds ionize in aqueous solution giving simple and complex ions
4.These bonds are also directional and stereoisomerism is also found
5.Molecules possess definite shape and definite bond angles, thus have definite
geometry

A B
ψA ψB
ψAB = N(cA ψA + cBψB)
ψ2
AB = (cA2
ψA2
+ 2cAcB ψA ψB + cB2
ψB 2
)
Overlap integral
The wave function for the molecular orbitals can be approximated by taking linearThe wave function for the molecular orbitals can be approximated by taking linear
combinations of atomic orbitals.combinations of atomic orbitals.
Probability density
c – extent to which each AO
contributes to the MO

cA = cB = 1
+. +. . .+
bondingψg
Amplitudes of wave
functions added
ψg = N [ψA + ψB]
Constructive interferenceConstructive interference

ψ2
AB = (cA2
ψA2
+ 2cAcB ψA ψB + cB2
ψB 2
)
electron density on original atoms,electron density on original atoms,
density between atomsdensity between atoms

The accumulation of electron density between the nuclei put the electron in a position
where it interacts strongly with both nuclei.
The energy of the molecule is lower
Nuclei are shielded from each other

Amplitudes of wave
functions
subtracted.
Destructive interferenceDestructive interference
Nodal plane perpendicular to the H-H bondNodal plane perpendicular to the H-H bond
axis (en density = 0)axis (en density = 0)
Energy of the en in this orbital is higher.Energy of the en in this orbital is higher.
+. -. ..
node
antibonding
ψu = N [ψA - ψB]
cA = +1, cB = -1 ψu
+ -
ΨA-ΨB

The electron is excluded from internuclear regionThe electron is excluded from internuclear region  destabilizingdestabilizing
AntibondingAntibonding

When 2 atomicWhen 2 atomic orbitalsorbitals combine there are 2combine there are 2
resultantresultant orbitalsorbitals..
low energy bonding orbitallow energy bonding orbital
high energyhigh energy antibondingantibonding orbitalorbital
1sb 1sa
σ1s
σ*
E
1s
MolecularMolecular
orbitalsorbitals
EgEg. s. s orbitalsorbitals

Molecular potential energy curve shows the variation of
the molecular energy with internuclear separation.

Looking at the Energy Profile
• Bonding orbital
• called 1s orbital
• s electron
• The energy of 1s orbital
decreases as R decreases
• However at small separation,
repulsion becomes large
• There is a minimum in potential
energy curve

11.4 eV
109 nm
H2H2
Location of
Bonding orbital
4.5 eV
LCAO of n A.O ⇒ n M.O.

The overlap integral
∫= τψψ dS BA
*
The extent to which two atomic orbitals on different atom overlaps : theThe extent to which two atomic orbitals on different atom overlaps : the
overlap integraloverlap integral

S > 0 Bonding S < 0 anti
S = 0 nonbonding
Bond strength depends on theBond strength depends on the
degree of overlapdegree of overlap

Homonuclear Diatomics
• MOs may be classified according to:
(i) Their symmetry around the molecular axis.
(ii) Their bonding and antibonding character.
∀σ1s< σ1s*< σ2s< σ2s*< σ2p< πy(2p) = πz(2p)
<πy*(2p) =πz*(2p)<σ2p*.

dx2
-dy2
and dxy
Cl4Re ReCl4
2-

First period diatomic moleculesFirst period diatomic molecules
σ1s2
H
Energy
HH2
1s 1s
σg
σu*
Bond order =
½ (bonding electrons – antibonding electrons)
Bond order: 1

σ1s2
, σ*
1s2He
Energy
HeHe2
1s 1s
σg
σu*
Molecular Orbital theory is powerful because it allows us to predict whether
molecules should exist or not and it gives us a clear picture of the of the
electronic structure of any hypothetical molecule that we can imagine.
Diatomic molecules: The bonding in He2
Bond order: 0

Second period diatomic moleculesSecond period diatomic molecules
σ1s2
, σ*
1s2
, σ2s2
Bond order: 1
Li
Energy
LiLi2
1s 1s
1σg
1σu*
2s 2s
2σg
2σu*

σ1s2
, σ*
1s2
, σ2s2
,
σ*
2s2
Bond order: 0
Be
Energy
BeBe2
1s 1s
1σg
1σu*
2s 2s
2σg
2σu*
Diatomic molecules: Homonuclear Molecules of the Second Period

Molecular orbital theory was put forward by R.S. Mulliken to explain the nature of
bonding in
the molecules of covalent compounds. Mulliken was awarded Nobel Prize for Chemistry
in 1966.
Major postulates of the theory are:
(i) The wavefunction of an electron in a molecule is called molecular orbital (MO). The
molecular orbital surrounds all the nuclei in the molecule, i.e., MO’s are polycentric.
(ii) The atomic orbitals (AO’s) of nearly equal energy, and appropriate symmetry
combine to
give equal number of MO’s. The MO’s are constructed by the linear combination of the
atomic
orbitals (LCAO method).
(iii) MO of lower energy is called bonding molecular orbital ( ψb ), while that of higher
energy as antibonding molecular orbital (ψ a ),

(v) The electrons of the constituent atoms of a molecule are distributed over all
the available
MO’s in accordance with the Aufbau principle, the Pauli's exclusion principle and
Hund’s rule.
(vi) Like atomic orbitals (AO’s), the molecular orbitals can also be arranged
according to their energies. The internuclear axis is taken to be in the z-direction.
For the molecule or molecular ions formed from Li, Be, B, C, and N, the energies
of 2s and 2p orbitals are quite close to each other. Because of the repulsion
between the electrons that occupy 2s and 2p orbitals, the energy of the σ2p
molecular orbital gets raised. Relative to π 2p orbitals.

Splitting patterns for the second row Diatomic
If we combine the splitting schemes for the 2s and 2p orbitals, we can predict
bond order in all of the diatomic molecules and ions composed of elements in the first
complete row of the periodic table. Remember that only the valence orbitals of the
atoms need be considered.

One minor complication that you should be aware of is that the
relative energies of the σ and π bonding molecular orbitals are
reversed in some of the second-row diatomics.

The presence of one or more unpaired electrons accounts for the paramagnetic
nature of the molecule. The electronic configuration in which all the
electrons are paired indicate the diamagnetic nature of the species.
The strength of a chemical bond is described in terms of a parameter called bond
order.
As per definition, the bond order is expressed as,
Bond order = (No. of electrons in BMO-No. of electrons in ABMO)/2=(Nb-Na)/2
where, N b is the total number of electrons in bonding MOs.
N ais the total number of electrons in antibonding MOs.
(a) When, N b > N a : Bond order > 0 (+ ve). Then, a stable bond formation is
indicated.
(b) When, N b =N a : Bond order =0. Then, the bond is unstable. In fact. such a
bond is not formed.
Conditions For the Formation of MOs From the Atomic Orbitals
Formation of MOs by the combination of atomic orbitals takes place only if the
following conditions are satisfied:
(i) The combining atomic orbitals should have nearly equal energies. Only the
atomic orbitals of nearly the same energy combine to form MOs. For
example, 1s atomic orbitals of two atoms can combine to form one
bonding (σ 1s ) and one antibonding (σ * 1s ) orbitals. The 1s atomic
orbital of one atom cannot combine with 2s or 2p atomic orbital of the
other atom.

(ii) The combining atomic orbitals should have the same symmetry. The
atomic orbitals are oriented in space. Only those atomic orbitals can
combine to form molecular orbitals which have the same symmetry
about the molecular axis. For example, a pxorbital of an atom can
combine with a p xorbital of another atom. A p xorbital cannot
combine with a pzorbital.
(iii) The combining atomic orbitals should overlap effectively. MOs are
formed only if the combining atomic orbitals overlap to a reasonable
extent.
In-phase and out-of-phase wave combinations
“Matter waves” corresponding to the two separate hydrogen 1s orbitals
interact; both in-phase and out-of-phase combinations are possible,
and both occur. One of the resultants is the bonding orbital that we
just considered. The other, corresponding to out-of-phase
combination of the two orbitals, gives rise to a molecular orbital that
has its greatest electron probability in what is clearly the antibonding
region of space. This second orbital is therefore called an
antibonding orbital.

If a hydrogen atom is bonded to a highly electronegative element such as
fluorine, oxygen, nitrogen, then the shared pair of electrons lies more
towards the electronegative element. This leads to a polarity in the bond
in such a way that a slight positive charge gets developed on H-atom, viz.,
H+δ
: O−δ
H+δ
: F−δ
H+δ
: N−δ
This positive charge on hydrogen can exert electrostatic attraction on the
negatively charged electronegative atom of the same or the other
molecule forming a bridge-like structure such as
Xδ−
− Hδ+
× × × × × × Yδ−
− Hδ+
where X and Y are the atoms of strongly electronegative elements.
The bond between the hydrogen atom of one molecule and a more
electronegative atom of the same or another molecule is called
hydrogen bond.

A hydrogen bond is shown by a dotted line (.....).
Hydrogen bond is not a covalent bond as the 1s orbital of hydrogen is
already completed, and the 2s level is high up in its energy.
Conditions Necessary For the Formation of Hydrogen Bond
Hydrogen bond is formed only when the following conditions are
satisfied. (i) Only the molecules in which hydrogen atoms is linked to
an atom of highly electro-negative element, are capable of forming
hydrogen bonds.
(ii) The atom of the highly electronegative element should be small.
These conditions are met by fluorine, oxygen and nitrogen atoms. As a
results, all compounds containing hydrogen atom linked to an atom
of either N, O, or F exhibit hydrogen bonding.
Some Typical Compounds Showing Hydrogen Bonding
Hydrogen fluoride (HF).

Water (H2
O).
Ice (H2
O(s)).
Each Oxygen is "linked" in by a combination of a covalent bond and a
hydrogen bond to 4 other Oxygens.

Notice that each Oxygen can be linked to Hydrogen in one of two ways.
Or
Types of Hydrogen Bonding
There are two types of hydrogen bonding, viz.,
(a) Intermolecular hydrogen bonding (b) Intermolecular hydrogen
bonding
When the hydrogen bonding is between the H-atom of one molecule
and an atom of the electronegative element of another molecule, it is
termed as intermolecular hydrogen bonding. For example, hydrogen
bonding in water, ammonia etc., is intermolecular hydrogen bonding.

The intramolecular hydrogen bonding is between the hydrogen of one
functional group, and the electronegative atom of the adjacent
functional group in the same molecule. For example, the molecule of
o-nitrophenol, shows intramolecular hydrogen bonding. The p-
nitrophenol shows intermolecular hydrogen bonding.

Chemical bonding and molecular structure

More Related Content

What's hot

Viewers also liked

Similar to Chemical bonding and molecular structure

Recently uploaded

Chemical bonding and molecular structure