Molecular Structure and Orbitals: VSEPR, Hybridization, and Molecular Orbital Theory

Topic 4: Molecular Structure and Orbitals

Molecular Structure

Molecular structure refers to the three-dimensional arrangement of atoms in a molecule. Understanding molecular structure is essential for predicting chemical reactivity and physical properties.

• Bond Distance: The distance between the nuclei of two bonded atoms. $1~ ext{Å} = 10^{-10}~ ext{m}$ (angstrom), $1~ ext{pm} = 10^{-12}~ ext{m}$ (picometer)

• Bond Angle: The angle between any two bonds that include a common atom. Bond angles are influenced by:

  • Number of bonds

  • Number of unshared electron pairs (lone pairs) on the central atom

  Example: In methane (CH4), the bond angle is 109.5°.

Valence Shell Electron-Pair Repulsion (VSEPR) Theory

VSEPR theory predicts the geometry of molecules based on the repulsion between electron pairs in the valence shell of the central atom. The structure is determined by minimizing electron pair repulsions.

• Regions of Electron Density:

  • Covalent bonds (bonding pairs)

  • Unshared pairs of electrons (lone pairs)

• Rules on Repulsion:

  • Bond pair–bond pair < lone pair–bond pair < lone pair–lone pair

  • Bond angle between bonding pairs decreases as the number of lone pairs increases on the central atom.

Number of Lone Pairs	Bond Angle
CH4 (0)	109.5°
NH3 (1)	107°
H2O (2)	104.5°

Effect of Electronegativity

• Increasing size and lower electronegativity of the central atom permit lone pairs to be drawn out further, decreasing repulsion between bonding pairs.

• Repulsion exerted by bonding pairs decreases as the electronegativity of the bonded atoms increases.

Molecule	Bond Angle	bp-lp Repulsion
H2O	104.5°	Strongest
H2S	92.1°	Intermediate
H2Se	90.5°	Weakest

Multiple Bonds and VSEPR

• Repulsion exerted by triple bonds > double bonds > single bonds.

• Double bonds are considered as one effective pair in VSEPR.

• Bond angles decrease as the number of electron pairs increases.

Special Geometries

• Trigonal Bipyramidal: Five electron pairs; minimal repulsion achieved by two trigonal-based pyramids sharing a common base. Lone pairs occupy equatorial positions (120° away).

• Octahedral: Six electron pairs; 90° bond angles. Lone pairs can occupy any position, but if two, they are opposite each other.

Molecules Containing No Single Central Atom

VSEPR can be applied to molecules with no single central atom, such as methanol. The arrangement of electron pairs and lone pairs around each atom determines the overall molecular structure.

Accuracy of the VSEPR Model

• VSEPR accurately predicts structures for most molecules formed from non-metallic elements.

• Can be used for molecules with hundreds of atoms.

• Fails in certain instances, e.g., phosphine (PH3) and ammonia (NH3) have similar Lewis structures but different bond angles.

Dipole Moment

A molecule with a center of positive charge and a center of negative charge is said to be dipolar or to possess a dipole moment.

• Represented by an arrow pointing to the negative charge center.

• Electrostatic potential diagrams show charge distribution.

• Red: Most electron-rich region; Blue: Most electron-poor region.

• Example: HF molecule has a significant dipole moment.

Bond Polarity

Bond polarity arises when atoms with different electronegativities form a bond, resulting in unequal sharing of electrons.

• Diatomic molecules with polar bonds exhibit dipole moments.

• Polyatomic molecules may possess polar bonds but lack a net dipole moment if bond polarities cancel out.

Shape	General Example	Classification	Bond Polarity
Linear	CO2	Nonpolar	Polar bonds, no net dipole
Tetrahedral	CCl4	Nonpolar	Polar bonds, no net dipole
Trigonal Pyramidal	NH3	Polar	Net dipole moment

Valence Bond Theory

Valence Bond Theory describes covalent bond formation as the overlap of atomic orbitals, allowing electrons of opposite spin to share space and form a bond.

• Increased overlap brings nuclei closer until a balance is reached between repulsion and attraction.

• Atoms cannot get too close due to internuclear repulsions.

Hybrid Orbitals

Hybrid orbitals are formed by mixing atomic orbitals to create new orbitals of equal energy (degenerate orbitals).

• Mixing two orbitals creates two hybrid orbitals; mixing three creates three, etc.

sp Hybridization

• Mixing one s and one p orbital yields two sp hybrid orbitals, oriented 180° apart (linear geometry).

• Example: BeF2 molecule.

sp2 Hybridization

• Mixing one s and two p orbitals yields three sp2 hybrid orbitals, oriented 120° apart (trigonal planar geometry).

• Example: BF3 molecule.

sp3 Hybridization

• Mixing one s and three p orbitals yields four sp3 hybrid orbitals, oriented 109.5° apart (tetrahedral geometry).

• Example: CH4 molecule.

Hypervalent Molecules

• Elements with more than an octet can use d orbitals for bonding (e.g., period 3 and below).

• dsp3 hybridization: Combination of one d, one s, and three p orbitals (trigonal bipyramidal).

• d2sp3 hybridization: Combination of two d, one s, and three p orbitals (octahedral).

Summary Table: Hybrid Orbitals

Number of Effective Pairs	Arrangement of Pairs	Hybridization Required
2	Linear	sp
3	Trigonal Planar	sp2
4	Tetrahedral	sp3
5	Trigonal Bipyramidal	dsp3
6	Octahedral	d2sp3

Types of Bonds: Sigma and Pi Bonds

• Sigma (σ) bonds: Formed by head-to-head overlap; electron density is along the internuclear axis.

• Pi (π) bonds: Formed by side-to-side overlap; electron density is above and below the internuclear axis.

• Single bonds are always σ bonds; double and triple bonds have one σ and one or two π bonds, respectively.

Molecular Orbital Theory

Molecular Orbital (MO) Theory describes the energy of electrons in a molecule using wave properties. Molecular orbitals are formed by the combination of atomic orbitals.

• Bonding Orbitals: Constructive combinations of atomic orbitals.

• Antibonding Orbitals: Destructive combinations; have a nodal plane where electron density is zero.

• Number of MOs equals the number of atomic orbitals used.

Molecular Orbital Diagram

• Shows how atomic orbitals combine to give molecular orbitals.

• Bond Order: $\text{Bond order} = \frac{\text{number of bonding electrons} - \text{number of antibonding electrons}}{2}$

• Larger bond order indicates stronger bonds.

Can H2 and He2 Form?

• H2: Bond order = 1 (stable molecule)

• He2: Bond order = 0 (unstable, does not exist)

Homonuclear Diatomic Molecules

• Composed of two identical atoms; valence orbitals contribute to MO.

• Both bonding and antibonding π molecular orbitals are formed from p orbitals.

• MO diagrams show the relative energies and electron configurations.

s and p Orbital Interactions

• s orbitals can interact with p orbitals, affecting the energy levels of MOs.

• Degenerate orbitals have equal energy.

MO Diagram and Magnetism

• Diamagnetism: All electrons are paired; substance is weakly repelled by a magnetic field.

• Paramagnetism: One or more unpaired electrons; substance is attracted to a magnetic field.

• Example: O2 is paramagnetic.

MO Diagram of Homonuclear Diatomic Molecules

Molecule	Bond Order	Magnetism
B2	1	Paramagnetic
C2	2	Diamagnetic
N2	3	Diamagnetic
O2	2	Paramagnetic
F2	1	Diamagnetic
Ne2	0	Diamagnetic

Heteronuclear Diatomic Molecules

• Consist of atoms from different elements; atomic orbitals have different energies.

• More electronegative atom has orbitals lower in energy; bonding orbitals resemble the more electronegative atom.

• Example: HF molecule; electron probability closer to fluorine.

Additional info:

• Examples and practice problems are provided throughout the notes to reinforce concepts, such as predicting molecular shapes, bond angles, and hybridization for various molecules.

• Tables and diagrams are used to summarize key points and visualize molecular structures and orbital interactions.