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
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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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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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