TYPES OF HYBRIDIZATION AND
GEOMETRY OF MOLECULES
Children’s club lecture
20-04-2010
K. Suthagar
1
CHEMICAL
BONDING
Chemical Bonding
Problems and questions —
How is a molecule or
polyatomic ion held
together?
Why are atoms distributed at
strange angles?
Why are molecules not flat?
Can we predict the structure?
How is structure related to
chemical and physical
properties?
Chemical Bonds
 Chemical bonds are the attractive forces that hold atoms together in the
form of compounds. They are formed when electrons are shared between
two atoms.
There are 3 types of bonds...covalent bonds, polar covalent bonds and ionic
bonds
4
Important concepts in chemical bonding and molecular structure
 The attractive force which holds together the constituent particles (atoms, ions or
molecules) in chemical species is known as chemical bond.
 Tendency of atoms of various elements to attain stable configuration of eight electrons
in their valence shell is the cause of chemical combination.
 The principle of attaining a maximum of eight electrons in the valence shell or
outermost shell of atoms is known as octet rule.
The tendency of an atom to take part in chemical combination is determined by the
number of valence electrons (electrons in the outermost shell of an atom).
The atoms acquire the stable noble gas configuration of having eight electrons in the
outermost shell (called octect rule) by mutual sharing or by transfer of one or more
electrons.
The valency (number of electrons an atom loses, gains or mutually shares to attain
noble gas configuration) of an element is either equal to the number of valence electrons
or equal to 8 minus the number of valence electrons.
5
Review of Chemical Bonds
 There are 3 forms of bonding:
 _________—complete transfer of
Most bonds are
somewhere in between
ionic and covalent.
1 or more electrons from one
atom to another (one loses, the
other gains) forming oppositely
charged ions that attract one
another
 _________—some valence
electrons shared between atoms
 _________ – holds atoms of a
metal together
General Properties of Covalent Compound
The covalent compounds do not exist as ions but they exist as molecules.
The melting and boiling points of covalent compounds are generally low.
 Covalent compounds are generally insoluble or less soluble in water and other polar
solvents.. However, these are soluble in non- polar solvents.
Covalent compounds do not give ions in solution, these are poor conductors of
electricity in the fused or dissolved state.
7
Characteristics of Covalent Compounds
 Covalent compounds are formed by the mutual sharing of electrons.
 There is no transfer of electrons from one atom to another and therefore no charges
are created on the atom. No ions are formed. These compounds exist as neutral
molecules and not as ions. Although some of the covalent molecules exist as solids,
they do not conduct electricity in fused or molten or dissolved state.
 ii) They possess low melting and boiling points. This is because of the weak
intermolecular forces existing between the covalent molecules. Since, no strong
coulombic forces are seen; some of covalent molecules are volatile in nature. Mostly
covalent compounds possess low melting and boiling points.
 iii) Covalent bonds are rigid and directional therefore different shapes of covalent
molecules are seen.
 iv) Most of the covalent molecules are non polar and are soluble in nonpolar (low
dielectric constant) solvents like benzene, ether etc and insoluble in polar solvents
like water. Carbon tetrachloride (CCl4) is a covalent nonpolar molecule and is
soluble in benzene.
8
Co-Ordinate Covalent Bond
 Covalent type bond in which both the electrons in the shared pair come
from one atom is called a coordinate covalent bond. Co-ordinate covalent
bond is usually represented by an arrow () pointing from donor to the
acceptor atom.
 Co-ordinate Covalent bond is also called as dative bond, donor – acceptor
bond, semi- polar bond or co-ionic bond. The electrostatic force of
attraction which holds the oppositely charged ions together is known as
ionic bond or electrovalent bond.
 Ionic compounds will be formed more easily between the elements with
comparatively low ionization enthalpy and elements with comparatively
high negative value of electron gain enthalpy.
9
General Properties of Ionic Compounds
 Ionic compounds usually exist in the form of crystalline solids. They have
high melting and boiling points and this ionic compounds are generally
soluble in water and other polar solvents having high dielectric constants.
Ionic compounds are good conductors of electricity in the solutions or in
their molten states.
 The chemical reactions of ionic compounds are characteristic of the
constituent ions and are known as ionic reactions. In ionic - compounds,
each ion is surrounded by oppositely charged ions uniformly distributed all
around the ion and therefore, electrical field is nondirectional.
10
The type of bond can usually be calculated by
finding the difference in electronegativity of the two
atoms that are going together.
Electronegativity
Difference
 If the difference in electronegativities
is between:
 1.7 to 4.0: Ionic
 0.3 to 1.7: Polar Covalent
 0.0 to 0.3: Non-Polar Covalent
Example: NaCl
Na = 0.8, Cl = 3.0
Difference is 2.2, so
this is an ionic bond!
Ionic Bonds
All those ionic compounds were made
from ionic bonds. We’ve been through
this in great detail already. Positive
cations and the negative anions are
attracted to one another!)
Therefore, ionic compounds are usually
between metals and nonmetals (opposite
ends of the periodic table).
Electron
Distribution
in Molecules
G. N. Lewis
1875 - 1946
 Electron distribution is
depicted with Lewis
(electron dot)
structures
 This is how we can
decide how many
atoms will bond
covalently!
(In ionic bonds, it
was decided with
charges)
Bond and Lone Pairs
Valence electrons are distributed
as shared or BOND PAIRS and
unshared or LONE PAIRS.
••
H
Cl
•
•
••
shared or
bond pair
lone pair (LP)
This is called a LEWIS
structure.
Bond Formation
A bond can result from an overlap of
atomic orbitals on neighboring atoms.
••
H
+
Cl
••
••
•
•
H
Cl
•
•
••
Overlap of H (1s) and Cl (2p)
Note that each atom has a single, unpaired electron.
Review of Valence Electrons
Remember from the electron chapter
that valence electrons are the
electrons in the OUTERMOST energy
level… that’s why we did all those
electron configurations!
B is 1s2 2s2 2p1; so the outer energy
level is 2, and there are 2+1 = 3
electrons in level 2. These are the
valence electrons!
Br is [Ar] 4s2 3d10 4p5
How many valence electrons are
present?
Review of Valence Electrons
Number of valence electrons of a main (A)
group atom = Group number
Lewis electron dot symbols
The symbols of the element is written first. This represents the
nucleus of the element with all the inner electrons that do not take
part in the bond formation.
The valence electrons are then written as dots or (small cross
marks) around the symbol. They are spread in a pair on four sides of
the symbol.
In case of ions the charge is shown with the symbol
Electron dot symbol representation of F2
19
Valence Shell Electron Pair Repulsion (VSEPR) theory
The shape of a molecule can be determined from the arrangement and repulsions
between the electron pairs present in the valence shell of central atom of that
molecule.
ii) There are two types of valence shell electron pairs viz., i) Bond pair and ii) Lone
pair
iii) The electron pairs in the valence shell the repel each other and determines the
shape of the molecule. The magnitude of the repulsion depends upon the type of
electron pair.
iv) The bond pair is attracted by nuclei they occupy less space and hence it causes
less repulsion. Whereas, the lone pairs are only attracted by one nucleus and
hence occupy more space. As a result, the repulsion caused by them is greater.
v) The order of repulsion between different types of electron pairs is as follows :
Lone pair - Lone pair > Lone Pair -Bond pair > Bond pair- Bond pair
vi) When the valence shell of central atom contains only bond pairs, the molecule
gets sym-metrical structure, whereas the symmetry is distorted when there are
lone pairs along with bond pairs.
vii) The bond angle decreases due to the presence of lone pairs.
viii) The repulsion increases with increase in the number of bonds between
two atoms
ix) Triple bond causes more repulsion then double bond which in turn
cause more repulsion than single bonds
x) The repulsion between electron pairs increases with increase
electronegativity of central atom and hence the bond angle increases.
xi) Shapes of molecules can be predicted from the number of electron pairs
in the valence shell of central atom
Lewis dot structures of Cl2 , O 2 and ethane molecules
22
Steps for Building a Dot Structure
Ammonia, NH3
1. Decide on the central atom; never H. Why?
If there is a choice, the central atom is atom of
lowest affinity for electrons. (Most of the time, this is
the least electronegative
Therefore, N is central on this one
2. Add up the number of valence electrons
that can be used.
H = 1 and N = 5
Total = (3 x 1) + 5
= 8 electrons / 4 pairs
Building a Dot Structure
3.
Form a single bond
between the central atom and
each surrounding atom (each
bond takes 2 electrons!)
H N H
4.
Remaining electrons form LONE PAIRS to
complete the octet as needed (or duet in the case of H).
3 BOND PAIRS and 1 LONE PAIR.
H
••
H N H
Note that N has a share in 4 pairs (8 electrons),
while H shares 1 pair.
5. Check to make sure there are 8 electrons around each
atom except H. H should only have 2 electrons. This
includes SHARED pairs.
H
Carbon Dioxide, CO2
1. Central atom =
2. Valence electrons =
3. Form bonds.
C 4 eO 6 e- X 2 O’s = 12 eTotal: 16 valence electrons
This leaves 12 electrons (6 pair).
4. Place lone pairs on outer atoms.
5. Check to see that all atoms have 8 electrons around it except for H, which can
have 2.
Carbon Dioxide, CO2
C 4 eO 6 e- X 2 O’s = 12 eTotal: 16 valence electrons
How many are in the drawing?
6. There are too many electrons in our drawing. We
must form DOUBLE BONDS between C and O.
Instead of sharing only 1 pair, a double bond shares 2
pairs. So one pair is taken away from each atom and
replaced with another bond.
Double and even
triple bonds are
commonly
observed for C, N,
P, O, and S
H2CO
SO3
C2F4
Violations of the Octet Rule
(Honors only)
Usually occurs with B and elements of
higher periods. Common exceptions
are: Be, B, P, S, and Xe.
Be: 4
B: 6
P: 8 OR 10
S: 8, 10, OR 12
Xe: 8, 10, OR 12
SF4
BF3
Significance of Lewis Symbols
a) The Lewis symbols indicate the number of electrons in the outermost or valence shell
which helps to calculate common or group valency.
b) The common valency of an element is either equal to number of dots or
valence electrons in the Lewis symbol or it is equal to 8 minus the number of dots or
valence electrons.
c) The bond formed by mutual sharing of electrons between the combining atoms
of the same or different elements is called a covalent bond.
d) If two atoms share one electron pair, bond is known as single covalent bond and is
represented by one dash (–).
e) If two atoms share two electron pairs, bond is known as double covalent bond and is
represented by two dashes (=).
f) If two atoms share three electron pairs, bond is known as triple covalent bond and is
represented by three dashes ( º).
g) The formal charge of an atom in a polyatomic ion or molecule is defined as the
difference between the number of valence electrons in an isolated (or free) atom
and the number of electrons assigned to that atom in a Lewis structure.
29
MOLECULAR GEOMETRY
MOLECULAR GEOMETRY
VSEPR
 Valence Shell Electron Pair
Repulsion theory.
 Most important factor in
determining geometry is
relative repulsion between
electron pairs.
Molecule adopts the shape
that minimizes the
electron pair repulsions.
Some Common Geometries
Linear
Trigonal Planar
Tetrahedral
Structure Determination by VSEPR
Water, H2O
2 bond pairs
2 lone pairs
The molecular geometry is
BENT.
The electron pair geometry is
TETRAHEDRAL
Structure Determination by
VSEPR
Ammonia, NH3
The electron pair geometry is tetrahedral.
lone pair of electrons
in tetrahedral position
N
H
H
H
MOLECULAR GEOMETRY — the positions of
the atoms — is TRIGONAL PYRAMID.
The
Bond Polarity
HCl is POLAR because it
has a positive end and a
negative end. (difference
in electronegativity)
+d -d
••
••
H Cl
••
Cl has a greater share in
bonding electrons than
does H.
Cl has slight negative charge (-d) and H has
slight positive charge (+ d)
Bond Polarity
 This is why oil and water will not mix! Oil
is nonpolar, and water is polar.
 The two will repel each other, and so you
can not dissolve one in the other
Bond Polarity
 “Like Dissolves Like”
 Polar dissolves Polar
 Nonpolar dissolves
Nonpolar
Put on your 3-D glasses!
Hybridization
 Uses modifications of molecular models to account for
observed structures of molecules or ions
 Is a mixing of the native atomic orbitals to form special
hybrid orbitals for bonding
 The special orbitals will then strive to be as far away from
each other in space as they can be
Types of Hybridization
sp3
sp2
sp
dsp3
d2sp3
Types
of
Bonds
 Sigma (σ) bonds
 End-to-end bonding
 There is an overlap
 One in every type of bond
 Pi (π) bonds
 Side-to-side bonding
 There is no overlap
 1 in a double bond
 2 in a triple bond
Native s orbital
x
z
y
Native p orbitals
x
z
y
px
x
z
y
x
py
z
y
pz
Hybrid sp3 orbitals
x
z
x
z
y
y
x
x
z
z
y
y
Hybrid sp3 orbitals
x
z
y
Hybrid sp3 orbitals
x
z
y
sp3 orbitals
4 effective pairs
109.5°, 107.3°(w/1 lone pair), or 104.5° (w/2 lone pairs)
tetrahedral, pyramidal, or bent
Hybrid sp3 orbitals
overlapping orbitals
H
x
H
z
CH4
C
H
H
y
4 σ bonds in
the molecule
sp3 hybridization –Methane
48
Orbital picture and structure of methane
49
Native s orbital
x
z
y
Native p orbitals
x
z
y
px
x
z
y
x
Leave this one as is
py
z
y
pz
Hybrid sp2 orbitals
x
z
x
y
x
z
y
z
y
Hybrid sp2 orbitals
x
z
y
sp2 orbitals
3 effective pairs
120°
trigonal planar
Hybrid sp2 orbitals
Remember the unhybridized p orbital?
x
z
y
sp2 orbitals
3 effective pairs
120°
trigonal planar
Hybrid sp2 orbitals
overlapping orbitals
side-by-side orbitals
H
H
x
C
H
H
z
C2H4
x
C
z
y
y
5 σ bonds
1 π bond
in molecule
sp2 hybridization Ethene
56
Orbital picture of ethylene
57
Native s orbital
x
z
y
Native p orbitals
x
z
y
px
x
z
y
x
Leave this one as is
py
Leave this one as is
z
y
pz
Hybrid sp orbitals
x
x
z
z
y
y
Hybrid sp orbitals
x
z
sp orbitals
2 effective pairs
180°
linear
y
Hybrid sp orbitals
Remember the unhybridized p orbitals?
x
z
sp orbitals
2 effective pairs
180°
linear
y
Hybrid sp orbitals
overlapping orbitals
x
side-by-side orbitals
x
z
z
y
C2H2
y
Hybrid sp orbitals
H
x
z
C2H2
x
C
C
y
y
z
H
sp hybridization – Ethyne
65
Orbital picture of ethyne
66
Hybridization and molecular shapes of some molecules involving sporbitals
67
Hybridization involving d orbitals
 sp3d hybridization
 sp3d2 hybridization
 sp3d3 hybridization
 dsp2 hybridization
sp3d hybridization
Formation of five sp3d hybrid orbitals, which adopt trigonal bipyramidal
geometry Phosphorus penta fluoride involves sp3d-Hybridization
68
Geometry of PF5 molecule
Formation of PF5 molecule involving sp3d Hybridization
Structure of PF5
69
sp3d2 Hybridization - Geometry of SF6 molecule
Formation of SF6 molecule involving sp3d2 Hybridization
Octahedral geometry of SF6 molecule
70
Geometry of IF7 molecule - sp3d3 hybridization
Formation of IF7 molecule involving sp3d3 hybridization
Pentagonal bipyramidal geometry of IF7 molecule
71
dsp2 hybridization Geometry of [Ni(CN)4]2-
Formation of [Ni(CN)4]2- involving dsp2 hybridization
Structure of [Ni(CN)4]2-
72
Molecular Geometry
bond length,
angle
Lewis structures
bonding
geometry
VSEPR
Valence
determined experimentally
Shell
Electron
octahedron
Pair
Repulsion
90o bond angles
small groups
big groups
trigonal bipyramid
equatorial
120o
axial
180o
tetrahedron
109.5o
trigonal planar
120o
linear
geometry
apply to Chemistry
180o
180o
linear
BeCl2
valence e- =
2+
(2 x 7)
= 16e-
fewer than 8e-
..
two
Be
..
..
Cl
..
..
..
Cl
valence pairs on Be
bonding e-
linear molecule
180o
linear
CO2
valence e- =
two
C
..
O
..
valence pairs on C
..
..
..
O
..
4+
(2 x 6)
= 16e-
..
O
..
C
..
O
..
ignore double bonds
single and double bonds same
molecular geometry
linear
molecular shape
linear
120o
trigonal planar
SO2
..
..
O
:
S
..
O
..
..
..
O
..
= 18e-
:
..
O
..
(2 x 6)
..
three
S
6+
..
..
..
O
..
:
valence e- =
S
valence pairs on S
two bonding pairs
one lone pair
molecular geometry
molecular shape
< 120o
trigonal
bent
..
..
O
tetrahedral
109.5o
CH4
valence e- =
4+ (4 x 1)
= 8eH
four
valence pairs on C
H
109.5o
C
H
H
molecular geometry
molecular shape
tetrahedral
tetrahedral
tetrahedral
109.5o
NH3
valence e- =
5+
(3 x 1)
= 8e-
:
four
H
valence pairs on N
N
H
three bonding pairs
H
one lone pair
molecular geometry
molecular shape
< 109.5o
tetrahedral
trigonal pyramid
tetrahedral
109.5o
H2O
valence e- =
6+
(2 x 1)
= 8e-
:
four
valence pairs on O
two bonding pairs
H
O
H
two lone pair
:
molecular geometry
molecular shape
< 109.5o
tetrahedral
bent
bipyramidal
120o and 1800
PCl5
5+
(5 x 7)
Cl
= 40e-
five
valence pairs on P
P
90o
molecular geometry
120o
molecular shape
bipyramidal
bipyramidal
..
..
180o
..
..
..
Cl
..
Cl
..
..
..
Cl
..
..
Cl
..
..
..
valence e- =
..
bipyramidal
120o and 1800
SF4
valence e- =
6+
(4 x 7)
= 34e-
:
five
..
four bonding pairs
..
one lone pair
S
< 180o
molecular geometry
molecular shape
..
F
..
..
F
..
.. ..
..
F
..
..
F
..
valence pairs on S
bipyramidal
seesaw
bipyramidal
120o and 1800
ClF3
valence e- =
7+
(3 x 7)
= 28e-
:
five
..
three bonding pairs
two lone pair
molecular geometry
molecular shape
90o
180o
Cl
..
F
..
..
F
..
.. ..
..
F
..
valence pairs on Cl
bipyramidal
T
120o and 1800
bipyramidal
ICl2valence e- =
7+
(2 x 7)
+ e-
= 22e-
:
five
..
I
two bonding pairs
..
Cl
..
three lone pair on I
molecular geometry
molecular shape
bipyramidal
linear
..
..
Cl
..
valence pairs on I
octahedral
90o
BrF5
valence e- =
7+
(5 x 7)
= 42e-
:
..
six
valence pairs on Br
Br
..
..
five bonding pairs
molecular geometry
molecular shape
..
..
..
F
one lone pair
..
..
..
F
..
F
..
..
..
..
F
..
F
octahedral
square pyramidal
octahedral
90o
XeF4
valence e- =
8+
(4 x 7)
= 36e-
:
valence pairs on Xe
two lone pair
molecular geometry
molecular shape
..
..
four bonding pairs
Xe
:
six
..
F
..
..
F
..
..
..
..
F
..
..
F
..
octahedral
square planar
: :
..
S
:F
F
B
:
..
..
:
O
S
..
..
F
:
:
..
O
..
S
C
: :
..
..
O
..
Be
..
..
H
..
..
..
O
..
Cl
..
..
O
..
..
..
..
O
..
Cl
..
H
C
..
..
O
:
N
H
..
O
..
..
..
:
..
O
..
O
H
:
H
C
S
H
O
O
H
:
..
..
..
..
..
..
..
Cl
..
I
..
Cl
..
..
..
..
..
F
..
F
..
Cl
.. ..
..
..
.. ..
..
F
..
..
:
:
:
..
S
..
F
..
..
F
..
..
S
F
..
F
..
..
F
..
..
F
..
..
F
..
..
..
..
:
..
F
..
..
F
..
..
P
..
Cl
..
..
Cl
..
..
F
..
H
..
O
H
..
Cl
..
..
Cl
..
..
..
:
Cl
..
F
..
..
:
Br
..
..
..
F
..
..
F
..
..
..
..
..
..
Xe
..
F
..
..
F
..
..
F
..
..
F
..
:
:
..
..
F
..
..
F
..
90
Geometry of molecules containing one or more lone pairs in central atom
91
Questions?