Lecture 7
CHEM101
Hybridisation of
Atomic Orbitals
Dr. Noha Osman
Learning Outcomes
•
•
•
•
Understand the valence bond theory
Understand the concept of hybridization.
Understand the different types of orbital hybridization.
Understand the relationship between hybridization and
molecular geometry
• Determine the effect of multiple bonds.
• Be able to predict the type of hybridization in given
covalent compounds.
• Know the different types of intermolecular forces.
2
Covalent Bonding
• Lewis theory: covalent bond is merely formed by the
sharing of electrons.
H-H ≠ F-F
(different bond enthalpies and bond lengths)
• Quantum mechanics: more complete explanation of
covalent bonding; two theories were postulated:
1- Valence Bond Theory:
Uses atomic orbitals of individual atoms to explain covalent
bonding.
2- Molecular Orbital Theory:
Assumes the formation of molecular orbitals to explain
covalent bonding.
3
Valence bond theory
This theory assumes that a covalent
bond is formed when:
1- atomic orbital of one atom merges
with that of another atom. The orbitals
are then said to share a common region
of space, or to overlap.
2- Overlap results in energy release,
therefore the energy of the system
reaches a minimum and is most stable
(bond formation: exothermic reaction!)
4
Hybridisation
• VB theory uses hypothetical hybrid orbitals, which are
atomic orbitals obtained when two or more
nonequivalent orbitals of the same atom combine in
preparation for covalent bond formation.
• Hybridization is the term applied to the mixing of atomic
orbitals in an atom (usually a central atom) to generate a
set of equivalent hybrid orbitals.
• The valence orbitals, associated with the highest
principal quantum level that contain electrons on a given
atom are responsible for hybridization.
5
Real Life Hybridisation
A ZONKEY is a HYBRID of
a DONKEY and a ZEBRA
!!!
6
Types of Hybridisation
sp3 Hybridisation
sp2 Hybridisation
sp Hybridisation
sp3d Hybridisation
sp3d2 Hybridisation
7
sp3 Hybridisation, CH4 molecule
The electronic configuration of C is 1s2
2s2
↑↓
2p2
↑ ↑
• It might be expected that C would form only two bonds with 2 H
atoms, since it has two unpaired electrons.
• However, it actually forms four C-H bonds in methane!
• This can be explained by assuming that an electron from the 2s is
energetically excited to the empty 2p orbital:
2p3
2s1
↑
↑ ↑ ↑
Thus, four bonds can be formed:
1 bond: overlap of 2s orbital of C
and 1s orbital of H
3 bonds: overlap of 2p orbitals of C
with 1s orbitals of three H atoms.
8
sp3 Hybridisation, CH4 molecule (continue)
• Since all four C-H bonds are identical, then this means that all four C
orbitals involved in bonding are identical, too!
• This can be explained by:
Hybridisation
2p
sp3
sp3
C-H bonds
2s
1s
Electron Promotion
1s
Hybrid Orbitals
1s
Forming Bonds 9
sp3 Hybridisation
Four hybrid orbitals are directed
toward the four corners of a
regular tetrahedral.
CH4 has a tetrahedral shape,
and all the HCH angles are
109.5°.
10
sp3 Hybridisation, NH3 molecule
2p
sp3
lone
pair
2s
1s
N atom
(ground state)
sp3
1s
N atom
(hybridized state)
N-H
bonds
1s
N atom
(in NH3)
11
sp3 Hybridisation, NH3 molecule (continue)
Three of the four hybrid orbitals
form covalent N-H bonds
The fourth hybrid orbital
accommodates the lone pair on
nitrogen
Repulsion between the lone-pair
electrons and bonding-electron
pairs decreases the HNH bond
angles from 109.5° to 107.3°
Thus, Hybridisation and VSEPR model are
RELATED!
12
How to determine the type of hybridization?
Count the number of effective electron pairs around the
atom.
What is considered as an effective electron pair?
Lone pair.
Single bond.
Double bond.
Triple bond.
Each of the above is counted as one
effective electron pair.
13
Relationship between the number of effective pairs
and their hybrid orbitals
No of effective
electron pairs
Type of
hybridization
Arrangement
Of hybrid
orbitals
2
sp
Linear
3
sp2
Triangular planar
4
sp3
Tetrahedral
5
sp3d
Triangular
bipyramidal
6
sp3d2
Octahedral
14
sp2 Hybridisation, BF3 molecule
B: ground state electron
configuration, can form only
one bond!
B: energy promoted an
electron from the 2s to an
empty 2p orbital, thus can
form three bonds!
Since all B-F bonds are
identical, therefore sp2
hybridization must have
occurred!
sp2 orbitals
Empty p
orbital 15
sp2 Hybridisation, BF3 molecule (continue)
Three hybrid orbitals are
directed toward the 3 corners of
an equatorial, planar triangle.
BF3 has a trigonal planar shape,
and all the FBF angles are 120°.
16
Types of Covalent Bonds
VB Theory: Two ways orbitals can overlap to form covalent bonds
between atoms
Sigma Bonds ()
- end-to-end orbital overlap.
- electron density concentrated
between the nuclei of the
bonding atoms.
Pi Bonds ()
- Sideways orbital overlap.
- electron density concentrated
above and below the plane of
the nuclei of the bonding
atoms.
17
sp2 Hybridisation, C2H4 molecule
C: ground state electron
configuration.
C: energy promoted an
electron from the 2s to an
empty 2p orbital.
sp2 hybridised state
sp2 orbitals
2p orbital
18
sp2 Hybridisation, C2H4 molecule
(continue)
The three sp2 orbitals on each carbon atom
form two  bonds with the two hydrogen
1s orbitals and a third  bond with the sp2
hybrid orbital of the adjacent C atom.
The two unhybridised 2p orbitals of
each of the C atoms form a  bond
by overlapping sideways.
19
sp Hybridisation, BeH2 molecule
Be: ground state electron
configuration, cannot form
any bonds!
Be: energy promoted an
electron from the 2s to the
2p orbital, thus can form
two bonds!
Since both Be-H bonds are
identical, therefore sp
hybridization must have
occurred!
sp Hybridisation, BeH2 molecule (continue)
Be atom:
1s atomic
orbital of H
..
..
1s atomic
orbital of H
Overlap region
21
sp Hybridisation, CO2 molecule
Carbon atom
sp orbitals
Two 2p orbitals
Two effective pairs;
sp hybridization
..
..
O=C=O
..
..
Oxygen atom
sp2 orbitals
2p orbital
Three effective pairs;
sp2 hybridization
22
sp Hybridisation, CO2molecule (continue)
Note that two  bonds are formed between the two sp hybridized
orbitals of C and sp2 hybridized orbitals of each of the O atoms.
Note that the two 2p orbitals
remain unchanged on the sp
hybridized carbon. These are
used to form the two  bonds
with the 2p orbitals on each
of the oxygen atoms.
23
sp Hybridisation, N3 molecule
:NN:
sp orbitals Two 2p orbitals
Two effective pairs;
sp hybridization
24
sp3d Hybridization
•
E.g. SF4 & PCl5 : a set of five effective pairs around the S P
P atom is available.
• In general, a set of five effective pairs around a given atom
always requires a trigonal bipyramidal arrangement, which
in turn requires sp3d hybridization of that atom.
sp3d2 Hybridization
•
•
E.g. SF6 (6 bonding pairs) & XeF4 (4 bondig pairs + 2 lone
pairs around Xe).
In general, six electron pairs around an atom are always
arranged octahedrally and require sp3d2 hybridization of
the atom.
25
Hybridization with d-Orbitals, 3rd Row:
Extension of Octet Rule
s
p
d
Hybridization
F
F
sp3d
S: 3s23p4
F F F
Bonding
F
lone
pair
trigonal-bipyramidal
SF4
S
F
F
s
p
d
Hybridization
Bonding
F
F F F F
F
S: 3s23p4
sp3d2 F
octahedral
F
SF6
S
F
F
F
26
F
Practice Excercise
For each of the following molecules or ions, predict the
hybridization of each atom and describe the molecular
structure.
a. CO
The Lewis structure is :C O:
Each atom has two effective pairs, which means that
both are sp hybridized.
27
b. BF4The Lewis structure for such ion is as above.
Four pairs of electrons around B, thus, sp3 hybridized;
tetrahedral structure.
28
Intermolecular Forces
Intramolecular or bonding forces are found within a molecule.
Intermolecular or nonbonding forces are attractive forces found
between molecules.
Intermolecular forces are relatively weak compared to bonding
forces and thus require less energy to overcome. They are
responsible for physical properties of matter.
29
Types of Intermolecular Forces
van der Waals
Ion-Dipole
Forces
Hydrogen
Bonding
30
DipoleDipole
Forces
Ioninduced
dipole
forces
van der
Waals
Dipoleinduced
dipole
Forces
London
Dispersion
forces
31
van der Waals
1- Dipole-Dipole Forces
• Are attractive forces found between polar molecules,
that is, molecules that possess net dipole moments.
• Molecules align themselves to maximize attractive
interactions while minimize repulsive interactions.
• Dipole forces are 1% as strong as covalent bonds and
become weaker especially in the gas phase where the
distance between dipole increases.
The positive pole of one
polar molecule attracts the
negative pole of another.
solid
32
2- Ion-induced dipole interaction
Are attractive forces that arise between ions and temporary
dipoles induced in atoms or molecules.
3- Dipole-induced dipole interaction
Are attractive forces that arise between dipoles and temporary
dipoles induced in atoms or molecules.
33
4- London Dispersion forces
• Non-polar molecules and noble gases exhibit
temporary dipoles, where it happens that for a very
short time period the electrons are not distributed
symmetrically around the molecule (the electron
cloud is more dense on one side of the molecule in
comparison to the other) leading to temporary δ+
and δ- charges.
(Instantaneous dipole)
• This instantaneous dipole distorts the electron
distribution in neighboring molecules thus inducing
dipoles in these neighboring molecules.
(Induced dipole)
• This leads to interatomic or intermolecular attraction
that is relatively weak and short lived.
34
Ion-Dipole forces
• Are forces which attract an ion and a polar molecule to
each other.
• Strength of interaction depends on the charge and size
of the ion and on the magnitude of the dipole moment
and size of the molecule.
Interaction of a water molecule with a Na+ ion and a Mg2+ ion.
35
Hydrogen Bonding
• Hydrogen bonding is possible for molecules that have a
hydrogen atom covalently bonded to a small, highly
electronegative atom with lone electron pairs, specifically N,
O, or F.
• An intermolecular hydrogen bond is the attraction between
the H atom of one molecule and a lone pair of the N, O, or F
atom of another molecule.
• ↑ ∆ EN → ↑ bond polarity → ↑ strength of H bond.
36
Comparison of Strengths of Intermolecular Forces
Thank
You