Post on 22-Jan-2022
L7: VSEPR No. of Charge Centres
No. of Bonding
Pairs
No. of Lone Pairs
Electron Pair Geometry/
Shape
Molecular Geometry/
Shape Dipole? Angle Example
2 2 0 Linear Linear No 180o
3
3 0 Trigonal Planar
Trigonal Planar
No 120o
2 1 Bent/V-shaped
Yes 117o
4
4 0
Tetrahedral
Tetrahedral No 109.5o
3 1 Pyramidal Yes 107o
2 2 Bent/V-shaped
Yes 104.5o
5
5 0
Trigonal Bipyramidal
Trigonal Bipyramidal
No 90o
120o 180o
4 1 Distorted
Tetrahedron/ See Saw
Yes
3 2 T-shaped Yes 90o
180o
2 3 Linear No 180o
6
6 0
Octahedral
Octahedral No
90o 180o 5 1
Square Pyramidal
Yes
4 2 Square Planar No
*Dipole property in the table applies only to molecules with identical outer atoms.
L8: Hybridisation Covalent bonds consists of shared pair of e-, creating an area of e- density between atoms.
Sigma Bonds (π) Pi Bonds (π )
β’ Two same-axis overlap of s/p orbitals
β’ Variation in length depends on size of orbitals
β’ Two parallel overlap of p orbitals
β’ Weaker than π bonds
Single Bond: One π bond
Double Bond: One π & one π bond
Triple Bond: One π & two π bond
Formation of Covalent Bonds
Excitation: Occurs when e- is promoted within atom (e.g. from 2s to empty 2p orbital)
E.g. Carbon e- configuration: 1s2 2s2 2px1 2py
1
How Hybridisation Works
Process by which atomic orbitals within an atom to produce hybrid orbitals of intermediate energy.
Formation of stronger covalent bonds is possible from hybrid orbitals.
sp3 hybrid orbital (4) sp2 hybrid orbital (3) sp hybrid orbital (2)
ΒΌ of s character ΒΎ of p character
β of s character β of p character
Β½ of s character Β½ of p character
sp3 orbital (4 charge centres)
β’ Tetrahedral (109.5o)
β’ 4 π bonds
β’ 1 s orbital, 3 p orbitals
CH4 Each carbon creates 4 sp3 orbitals = 4 π bonds
sp2 orbital (3 charge centres)
β’ Trigonal Planar (120o)
β’ 3 π bonds + 1 π bond
β’ 1 s orbital + 2 p orbital
C2H4 Each carbon creates 3 sp2 orbitals = 3 π bonds
sp orbital (2 charge centres)
β’ Linear (180o)
β’ 2 π bonds + 2 π bonds
β’ 1 s orbital + 1 p orbital
C2H2
Each carbon creates 2 sp orbitals = 2 π bonds
Commented [J1]: Images sourced from: https://chemistry.boisestate.edu/richardbanks/inorganic/bonding%20and%20hybridization/bonding_hybridization.htm
L9: Molecular Orbital (MO) Theory
Bond Order = ππ.ππ π΅ππππππ πΈππππ‘ππππ βππ.πππ΄ππ‘πβπππππππ πΈππππ‘ππππ
2
Molecular Orbitals Formation
S-Orbital
P-Orbital
H2 H2- H2
2-
Bond Order = 2/2 = 1
Bond Order = (2-1)/2 = 0.5
Bond Order = (2-2)/2 = 0
He2 He2+
Bond Order = (2-2)/2 = 0
Bond Order = (2-1)/2 = 0.5
O2
Bond Length = (8 β 4)/2 = 2
F2 NO
Bond Order = (8 β 6)/2 = 1
Bond Order = (8 β 3)/2 = 2.5
Paramagnetic e-: Unpaired e-
Diamagnetic e-: Paired e-