Different Bond Angles and Molecular Structures by MOT
Understanding how Molecular Orbital Theory (MOT) explains bond angles and molecular geometry
Introduction
Molecular Orbital Theory (MOT) is a powerful concept that explains how atomic orbitals combine to form molecular orbitals, determining the stability, shape, and bond angles of molecules. While VSEPR theory predicts geometry based on electron repulsion, MOT gives a deeper quantum-level insight into bond formation, hybridization, and bond order. Let us explore how MOT helps us understand different bond angles and structures of various compounds.
1. Linear Molecular Structure (180° Bond Angle)
In molecules such as BeCl₂ or CO₂, the atoms align in a straight line. According to MOT, the bonding involves the overlap of sp hybrid orbitals (for Be or C) with p orbitals (for O or Cl). Because the orbitals are oriented directly opposite each other, the resulting molecular geometry is linear with a bond angle of 180°.
In CO₂, the bonding consists of two sigma (σ) and two pi (π) bonds between carbon and oxygen. The central carbon atom is sp hybridized, producing a linear shape that minimizes electron repulsion.
2. Trigonal Planar Structure (120° Bond Angle)
In BF₃, boron undergoes sp² hybridization where one 2s and two 2p orbitals mix to form three equivalent sp² orbitals lying in one plane. These orbitals overlap with the p orbitals of fluorine to form σ bonds. The result is a trigonal planar geometry with bond angles of 120°. MOT explains that delocalized π molecular orbitals in such systems further stabilize the planar structure, as seen in SO₃.
3. Tetrahedral Structure (109.5° Bond Angle)
In methane, carbon undergoes sp³ hybridization forming four equivalent orbitals that point toward the corners of a tetrahedron. Each hydrogen 1s orbital overlaps with one sp³ orbital to form a σ bond. According to MOT, the molecular orbitals result in uniform electron density distribution, giving a perfect 109.5° angle between bonds.
This tetrahedral structure minimizes repulsion between bonding electron pairs, leading to high stability.
4. Trigonal Pyramidal Structure (107° Bond Angle)
Ammonia has three bonding pairs and one lone pair on nitrogen. Due to sp³ hybridization, nitrogen forms four orbitals — one containing a lone pair. The lone pair-bond pair repulsion is greater than bond pair-bond pair repulsion, which slightly compresses the bond angle to 107°.
MOT explains that the nitrogen lone pair occupies a higher-energy nonbonding molecular orbital, altering the effective geometry from tetrahedral to trigonal pyramidal.
5. Bent or V-Shaped Structure (104.5° Bond Angle)
Water is a classic example of a bent molecule. The oxygen atom has two lone pairs and two bond pairs, forming an sp³ hybridized system. However, because lone pair-lone pair repulsion is maximum, the H–O–H angle is reduced to 104.5°.
MOT interprets this by showing that the molecular orbitals of O–H bonds are formed from overlap between O(2p) and H(1s), with lone pairs occupying higher-energy nonbonding orbitals that distort the ideal tetrahedral angle.
6. Trigonal Bipyramidal Structure (120° and 90° Bond Angles)
Phosphorus in PCl₅ uses sp³d hybridization where one 3d, one 3s, and three 3p orbitals combine to form five sp³d hybrid orbitals. Three of them form an equatorial plane (120° apart), and two are placed axially (90° to the plane). Thus, MOT helps visualize two types of bonding interactions — axial and equatorial — giving rise to distinct bond lengths and angles.
7. Octahedral Structure (90° Bond Angle)
In SF₆, the sulfur atom undergoes sp³d² hybridization, forming six equivalent orbitals oriented towards the corners of an octahedron. All S–F bonds are equivalent with angles of 90°.
MOT further clarifies that bonding occurs through σ-type interactions, and π interactions are negligible due to the high electronegativity of fluorine.
Comparison Table of Bond Angles and Hybridization
| Compound | Shape | Hybridization | Bond Angle (°) | Explained by MOT |
|---|---|---|---|---|
| CO₂ | Linear | sp | 180 | σ and π bonding via sp–p overlap |
| BF₃ | Trigonal Planar | sp² | 120 | Delocalized π orbitals stabilize planar shape |
| CH₄ | Tetrahedral | sp³ | 109.5 | Equal overlap of sp³ and H 1s orbitals |
| NH₃ | Trigonal Pyramidal | sp³ | 107 | Nonbonding MO distorts ideal angle |
| H₂O | Bent | sp³ | 104.5 | Lone pairs occupy higher MOs |
| PCl₅ | Trigonal Bipyramidal | sp³d | 90,120 | Axial and equatorial bond differences |
| SF₆ | Octahedral | sp³d² | 90 | 6 σ-bonds equally distributed |
Conclusion
Molecular Orbital Theory gives a quantum-level picture of how atomic orbitals combine to form molecular orbitals, helping us interpret not only bond order and magnetic properties but also bond angles and geometries. While hybridization and VSEPR are quick prediction tools, MOT explains the why — the true reason behind molecular shapes and angle variations.
Hence, by analyzing bonding and antibonding orbitals, MOT provides a deeper understanding of how electron distributions govern molecular structures in chemistry.
