Molecular Shapes, Orbitals & Orbital
Hybridisation
Session 8
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Review of General Chemistry, UNAM School of Medicine
Dr L.H.A. Prins (Ph.D.)
Dept. of Pharmacy
UNAM
Learning Outcomes
By the end of this session, the student should:
Understand molecular shapes/geometry. (Valence Shell
Electron- Pair Repulsion)
Understand electron-domain shapes/geometry.
Understand orbital shapes.
Understand orbital overlap ( & π bonds).
Understand orbital hybridisation.
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Molecular shapes & VSEPR model
How do we predict 3D geometry of molecules & ions?
Valence Shell Electron-Pair Repulsion (VSEPR) model = reliable
method for predicting shapes of covalent molecules & polyatomic ions.
Idea based on: Bond & lone-pair e- in valence shell of atoms, repel
each other & seek to be as far apart as possible.
Positions assumed by valence e- of an atom thus define angles between
bonds to surrounding atoms.
VSEPR – used for structures of main group elements (Group IA –
VIIIA).
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Molecular shapes & VSEPR model
Formula examples represented as AXn
A = central atom.
X = attached atom.
n = number of X bonded to A.
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Molecular shapes & VSEPR model
Linear
Trigonal-planar
Bent/V-shaped
Tetrahedral
Trigonal-pyramidal
Trigonal-bipyramidal
T-shaped
See saw
Octahedral
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Formula =
Formula =
Formula =
Formula =
Formula =
Formula =
Formula =
Formula =
Formula =
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AX2
AX3
..
: AX2
AX4
: AX3
AX5
..
: AX3
: AX4
AX6
Self-study!!
Linear (AX2)
Bonding pairs: 2
Linear geometry for 2 bond pairs
involve a central atom with less
than 8 e- in valence (bond) shell.
Lone pairs: 0
Example: BeCl2
Cl Be Cl
Bond angle = 180
(other examples: CO2, Ethyne, etc.)
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Trigonal-planar (AX3)
Bonding pairs: 3
Lone pairs: 0
Trigonal-planar geometry for 3 bond
pairs involve a central atom with less
than 8 e- in valence (bond) shell.
Example: BCl3
Cl
Bond angle = 120
B
Cl
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Cl
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..
Bent / Angular / V-shaped (:AX2)
Bonding pairs: 2
Lone pairs: 2
Example: H2O
H
O
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Lone pair e- - considered “fat”.
Bond pair e- - considered “skinny”.
H
O
H
Lone pair e- “push” the bond pairs closer
together.
H
Bond angle = 105
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Tetrahedral (AX4)
Bonding pairs: 4
Lone pairs: 0
Tetrahedral geometry for 4 bond pairs
involve a central atom with 8 e- in
valence (bond) shell.
Example: CH4
Bond angle = 109.5
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Trigonal-pyramidal (:AX3)
Bonding pairs: 3
Lone pairs: 1
Example: NH3
Trigonal-pyramidal geometry for 3
bond pairs & 1 lone pair involve a
central atom with 6 e- in valence (bond)
shell.
Bond angle = 107.5
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Trigonal-bipyramidal (AX5)
Bonding pairs: 5
Lone pairs: 0
Trigonal-bipyramidal geometry for 5
bond pairs involve a central atom with
10 e- in valence (bond) shell.
Example: PF5 Axial atom.
*
#
Bond angles:
*Between equatorial
atoms= 120
#Between equatorial &
axial atoms= 90
Equatorial atom.
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e- Domain geometry:
..
Tetrahedral (:AX2)
Bonding pairs: 2
Lone pairs: 2
Example: H2O
H
O
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H
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e- domain geometry =
Tetrahedral.
e- Domain geometry:
Tetrahedral (:AX3)
Bonding pairs: 3
Lone pairs: 1
Example: NH3
..
H
H
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N
H
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e- domain geometry =
Tetrahedral.
Multiple bonds & molecular
shapes
Influence of multiple bonds (unsaturated bonds) on
molecular geometry?
e- pairs involved in multiple bonds are all shared between
the same 2 atoms & therefore occupy the same region of
space.
Multiple bonds contribute to molecular geometry the
same as single bonds.
Example: CO2
O C O
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Linear molecular
geometry.
Bond angle = 180
Review of General Chemistry, UNAM School of Medicine
Orbital shapes
Atomic orbital: Predicts the volume of space around a nucleus
where an e- can most likely be found.
Pauli exclusion principle: No more than 2 e- of opposite spin per
orbital.
Orbital shapes: s, p, d (& f).
e- amount per orbital:
s
2ep 6ed 10e(f 14e-)
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s & p = most important in
Organic Chemistry.
Orbital shapes
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Orbital shapes
• Order of filling of subshells/sublevels:
Lowest energy
1s
2s
2p
3s
3p
3d
4s
4p
4d
4f
5s
5p
5d
5f
6s
6p
6d
6f
Highest energy
Remember Hund’s rule!!
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Orbital shapes
Sublevels of the First 4 Energy Levels.
Principal Energy
Level (Shells)
Sublevel
(Subshells)
# of Orbitals in
Sublevel
Total Possible
Occupying e-
1
s
1
2
2
s
1
2
p
3
6
s
1
2
p
3
6
d
5
10
s
1
2
p
3
6
d
5
10
f
7
14
3
4
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8 e-
18 e-
32 e-
Orbital overlap – covalent bonding
2 methods of orbital overlap during covalent bonding:
1. Linear orbital overlap ( covalent bonding).
Examples (H2, HCl, Cl2):
bond = along the orbital axis.
In general: Where there is a single bond between two atoms, there is a bond.
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Orbital overlap – covalent bonding
Parallel orbital overlap ( covalent bonding).
Example (CH2CH2):
2.
bond = sideways/parallel overlap (not along the orbital axis.)
The e- in a bond stay between the atoms & the e- in a bond are
delocalised on either side of the bonding plane. (Enhanced reactivity)
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Orbital hybridisation (for Carbon)
Showing orbital overlap for more complex molecules not
possible with simple s, p & d orbitals.
By mixing s, p & d atomic orbitals, hybrid orbitals are created.
Hybrid orbitals have a 3D orientation matching compound
geometry/shape.
Hybridised orbitals relevant to Organic Chemistry:
sp3 – Combination of 1 x s & 3 x p orbitals = 4 x sp3 orbitals.
Used for single bonds between atoms. Tetrahedral.
sp2 - Combination of 1 x s & 2 x p orbitals = 3 x sp2 orbitals.
Used for double bonds between atoms. Trigonal Planar.
sp - Combination of 1 x s & 1 x p orbital = 2 x sp orbitals.
Used for triple bonds between atoms. Linear.
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Orbital hybridisation (for Carbon)
Ground electronic state of C suggests it should form 2 bonds
as there are two unpaired e-.
How & why does C form 4 bonds?
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Orbital hybridisation (for Carbon)
Promotion of one of the two 2s e- increases energy, but
formation of four bonds causes a four-fold decrease.
Need to modify the orbital model to explain the tetravalent
character of C.
This modification is called orbital hybridisation.
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Review of General Chemistry, UNAM School of Medicine
sp3 orbitals & the structure of Methane
Carbon has 4 valence electrons (Excited state: 2s1 2px1 2p1y 2p1z).
In CH4, all C–H bonds are identical (tetrahedral). Why?
sp3 hybrid orbitals: s orbital & three p orbitals combine to form
four equivalent, unsymmetrical, tetrahedral orbitals (sppp = sp3).
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sp3 orbitals & the structure of Methane
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Structure of Methane (CH4)
sp3 orbitals on C overlap with 1s orbitals on 4 x H atoms to
form four identical C-H bonds.
(Each C–H bond has a length of 109 pm.)
Bond angle: Each H–C–H is 109.5°, the tetrahedral
angle.
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sp3 orbitals & the structure of Ethane
Two C’s bond to each other by overlap of an sp3 orbital from
each.
Three sp3 orbitals on each C overlap with 3 x H 1s orbitals to
form six C–H bonds.
(C–C bond is 154 pm long.)
All bond angles of ethane are tetrahedral.
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sp2 orbitals & the structure of Ethene
sp2 hybrid orbitals: s orbital combines with two p orbitals,
giving 3 orbitals (spp = sp2). This results in a double bond.
sp2 orbitals are in a plane with 120° angles (trigonal planar).
Remaining p orbital is perpendicular to the plane.
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sp2 orbitals & the structure of Ethene
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Bonds from sp2 hybrid orbitals
(Ethene)
Two sp2-hybridised orbitals overlap to form a bond.
Two p orbitals (on sp2 orbitals) overlap side-to-side to
form a pi () bond.
The & bonds result in the sharing of 4 e- and the
formation of a C-C double bond.
Ethene:
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Structure of Ethene
H atoms form bonds with 4 sp2 orbitals.
Bond angle: H–C–H & H–C–C bond angles = ±
120°; the trigonal-planar angle.
C–C double bond in ethene (ethylene) shorter & stronger
than single bond in ethane.
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sp orbitals & the structure of Ethyne
One s orbital & one p orbital combine to form two sp hybrid
orbitals.
Two p orbitals remain unchanged.
sp orbitals are 180° apart on x-axis (linear).
The two unchanged p orbitals are perpendicular on y-axis &
z-axis.
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sp orbitals & the structure of Ethyne
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Orbitals of Ethyne
Two sp hybrid orbitals from each C form bond.
pz orbitals from each C form a bond by sideways
overlap & py orbitals overlap similarly.
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Orbitals of Ethyne
6 e- shared in C C bond.
Two sp orbitals (those not involved in C C bond) form
bonds with 2 hydrogen atoms (1s).
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Summary: Orbital hybridisation
Carbon (C) uses hybrid orbitals to form bonds in organic molecules.
In single bonds with tetrahedral geometry, carbon has four sp3 hybrid
orbitals.
In double bonds with trigonal planar geometry, carbon uses three
equivalent sp2 hybrid orbitals & one unhybridised p orbital.
In triple bonds with linear geometry, carbon uses two equivalent sp
hybrid orbitals & two unhybridised p orbitals.
Atoms such as nitrogen and oxygen hybridise to form strong, oriented
bonds.
The N atom in ammonia & the O atom in water are sp3-hybridised.
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Quiz: Molecular shapes
Predict the molecular geometry of the following using
VSEPR model:
1.
2.
3.
4.
5.
6.
7.
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OF2
PF3
CS2
NSF
SO42PH3
SO3
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Quiz: Orbital hybridisation
If a central C atom is sp2 hybridised, what is the molecular
geometry?
9. If a C atom is sp hybridised, how many p-orbitals are left over to
make bonds?
10. If two C atoms are bound via overlap of sp hybrids, how many
C-C bonds & how many C-C bonds result?
11. What kind of hybrid orbitals do the C atom have in
formaldehyde, H2C=O?
12. What kind of hybridisation would you expect for the C atoms in
ethene/ethylene (C2H4)?
13. Where are the electrons in a bond & those in a bond?
8.
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END
Thank you
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Conceptualisation
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