L7: VSEPR
L7: VSEPR No. of Charge Centres 2
No. of Bonding Pairs 2
No. of Lone Pairs 0
3
0
2
Angle
No
180o
Trigonal Planar
No
120o
1
Bent/Vshaped
Yes
117o
4
0
Tetrahedral
No
109.5o
3
1
Pyramidal
Yes
107o
2
2
Bent/Vshaped
Yes
104.5o
5
0
Trigonal Bipyramidal
No
4
1
Distorted Tetrahedron/ See Saw
Yes
T-shaped
Yes
90o 180o
180o
5
6
Molecular Geometry/ Shape Linear
Dipole?
3
4
Electron Pair Geometry/ Shape Linear Trigonal Planar
Tetrahedral
Trigonal Bipyramidal
3
2
2
3
Linear
No
6
0
Octahedral
No
5
1
Square Pyramidal
Yes
4
2
Square Planar
No
Octahedral
90o 120o 180o
90o 180o
*Dipole property in the table applies only to molecules with identical outer atoms.
Example
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’ Two parallel overlap of p orbitals β’ Variation in length depends on size of β’ Weaker than π bonds orbitals
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 2py1
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 β of s character Β½ of s character ΒΎ of p character β of p 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 Commented [J1]: Images sourced from: https://chemistry.boisestate.edu/richardbanks/inorganic/b onding%20and%20hybridization/bonding_hybridization.htm
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
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-
Bond Order = 2/2 = 1 He2
Bond Order = (2-1)/2 = 0.5 He2+
Bond Order = (2-2)/2 = 0
Bond Order = (2-1)/2 = 0.5 O2
Bond Length = (8 β 4)/2 = 2
H22-
Bond Order = (2-2)/2 = 0
F2
NO
Bond Order = (8 β 6)/2 = 1
Bond Order = (8 β 3)/2 = 2.5
Paramagnetic e-: Unpaired eDiamagnetic e-: Paired e-
No. of Bonding Pairs 2
No. of Lone Pairs 0
3
0
2
Angle
No
180o
Trigonal Planar
No
120o
1
Bent/Vshaped
Yes
117o
4
0
Tetrahedral
No
109.5o
3
1
Pyramidal
Yes
107o
2
2
Bent/Vshaped
Yes
104.5o
5
0
Trigonal Bipyramidal
No
4
1
Distorted Tetrahedron/ See Saw
Yes
T-shaped
Yes
90o 180o
180o
5
6
Molecular Geometry/ Shape Linear
Dipole?
3
4
Electron Pair Geometry/ Shape Linear Trigonal Planar
Tetrahedral
Trigonal Bipyramidal
3
2
2
3
Linear
No
6
0
Octahedral
No
5
1
Square Pyramidal
Yes
4
2
Square Planar
No
Octahedral
90o 120o 180o
90o 180o
*Dipole property in the table applies only to molecules with identical outer atoms.
Example
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’ Two parallel overlap of p orbitals β’ Variation in length depends on size of β’ Weaker than π bonds orbitals
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 2py1
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 β of s character Β½ of s character ΒΎ of p character β of p 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 Commented [J1]: Images sourced from: https://chemistry.boisestate.edu/richardbanks/inorganic/b onding%20and%20hybridization/bonding_hybridization.htm
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
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-
Bond Order = 2/2 = 1 He2
Bond Order = (2-1)/2 = 0.5 He2+
Bond Order = (2-2)/2 = 0
Bond Order = (2-1)/2 = 0.5 O2
Bond Length = (8 β 4)/2 = 2
H22-
Bond Order = (2-2)/2 = 0
F2
NO
Bond Order = (8 β 6)/2 = 1
Bond Order = (8 β 3)/2 = 2.5
Paramagnetic e-: Unpaired eDiamagnetic e-: Paired e-
