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 (𝝈) Pi Bonds (𝝅)

• Two same-axis overlap of s/p orbitals

• Variation in length depends on size of orbitals

• Two parallel overlap of p orbitals

• Weaker than 𝜎 bonds

Single Bond: One 𝜎 bond

Double Bond: One 𝜎 & one 𝜋 bond

Triple Bond: One 𝜎 & 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)

¼ of s character ¾ of p character

⅓ of s character ⅔ of p character

½ of s character ½ of p character

sp3 orbital (4 charge centres)

• Tetrahedral (109.5o)

• 4 𝜎 bonds

• 1 s orbital, 3 p orbitals

CH4 Each carbon creates 4 sp3 orbitals = 4 𝜎 bonds

sp2 orbital (3 charge centres)

• Trigonal Planar (120o)

• 3 𝜎 bonds + 1 𝜋 bond

• 1 s orbital + 2 p orbital

C2H4 Each carbon creates 3 sp2 orbitals = 3 𝜎 bonds

sp orbital (2 charge centres)

• Linear (180o)

• 2 𝜎 bonds + 2 𝜋 bonds

• 1 s orbital + 1 p orbital

C2H2

Each carbon creates 2 sp orbitals = 2 𝜎 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 = 𝑁𝑜.𝑜𝑓 𝐵𝑜𝑛𝑑𝑖𝑛𝑔 𝐸𝑙𝑒𝑐𝑡𝑟𝑜𝑛𝑠 −𝑁𝑜.𝑜𝑓𝐴𝑛𝑡𝑖−𝑏𝑜𝑛𝑑𝑖𝑛𝑔 𝐸𝑙𝑒𝑐𝑡𝑟𝑜𝑛𝑠

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 – 4)/2 = 2

F2 NO

Bond Order = (8 – 6)/2 = 1

Bond Order = (8 – 3)/2 = 2.5

Paramagnetic e-: Unpaired e-

Diamagnetic e-: Paired e-