Molecular Geometry and Bonding Theories: VSEPR, Hybridization, and Molecular Polarity

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Chapter 9: Molecular Geometry and Bonding Theories

9.1 Molecular Shapes

The three-dimensional shape of a molecule is crucial in determining its physical and chemical properties. While Lewis structures show the connectivity of atoms and the types of bonds, they do not represent the actual spatial arrangement of atoms.

Lewis Structures: Illustrate the number and types of bonds but depict all atoms in a single plane, which is not always accurate for real molecules.
Actual Molecular Shape: Atoms are arranged in three dimensions to minimize electron pair repulsions, often resulting in geometries such as tetrahedral, trigonal planar, or linear.
Example: The Lewis structure of carbon tetrachloride (CCl4) shows four Cl atoms around a central C atom, but the actual shape is tetrahedral.

Lewis structure of CCl4 Tetrahedral, ball-and-stick, and space-filling models of CCl4

9.2 Valence-Shell Electron-Pair Repulsion (VSEPR) Model

The VSEPR model is used to predict the shapes of molecules based on the repulsion between electron domains (regions where electrons are likely to be found) around a central atom.

Electron Domain: A region where electrons are most likely found, including bonding pairs (shared between atoms) and non-bonding pairs (lone pairs on a single atom).
Bonding Pair: Electrons shared between two atoms.
Non-bonding Pair (Lone Pair): Electrons localized on one atom.
Electron Domain Geometry: The arrangement of electron domains around the central atom to minimize repulsions.

Two balloons linear orientation Three balloons trigonal-planar orientation Four balloons tetrahedral orientation

Counting Electron Domains

When using VSEPR, only the electron domains around the central atom are counted. Double and triple bonds count as one domain each.

Common Electron Domain Geometries

Number of Electron Domains	Arrangement	Electron-Domain Geometry	Predicted Bond Angles
2	Linear	Linear	180°
3	Trigonal planar	Trigonal planar	120°
4	Tetrahedral	Tetrahedral	109.5°

Table of electron-domain geometries for 2, 3, and 4 domains

Examples of Electron Domain Geometries

CO2: 2 domains, linear geometry, 180° bond angle.
SO3: 3 domains, trigonal planar geometry, 120° bond angle.
CCl4: 4 domains, tetrahedral geometry, 109.5° bond angle.

Table 9.1: Electron-domain geometries as a function of number of electron domains Trigonal bipyramidal and octahedral electron domain geometries

Expanded Electron Domain Geometries

Number of Electron Domains	Electron-Domain Geometry	Bond Angles
5	Trigonal bipyramidal	120°, 90°
6	Octahedral	90°

Trigonal bipyramidal and octahedral electron domain geometries

9.2 Molecular Geometry

Molecular geometry describes the arrangement of only the atoms (not lone pairs) in a molecule. The presence of lone pairs can alter the observed geometry from the electron domain geometry.

Example: Ammonia (NH3) has a tetrahedral electron domain geometry but a trigonal pyramidal molecular geometry due to one lone pair on nitrogen.
Example: Ozone (O3) and SnCl3- have bent and trigonal pyramidal geometries, respectively.

Lewis structure of NH3 Lewis structure of O3 Lewis structure of SnCl3-

Summary Table: Electron-Domain and Molecular Geometries

Number of Electron Domains	Electron-Domain Geometry	Molecular Geometry	Example
2	Linear	Linear	CO2
3	Trigonal planar	Trigonal planar / Bent	SO3 / O3
4	Tetrahedral	Tetrahedral / Trigonal pyramidal / Bent	CH4 / NH3 / H2O

Table 9.2: Electron-domain and molecular geometries for two, three, and four electron domains

9.3 Molecular Shape and Molecular Polarity

The polarity of a molecule depends on both the polarity of individual bonds and the overall molecular geometry. Bond dipoles are vector quantities; their sum determines the molecular dipole moment.

Non-polar Molecules: If all outer atoms are the same and the molecular geometry is symmetrical, bond dipoles cancel, resulting in a non-polar molecule (e.g., CCl4, CO2).
Polar Molecules: If the geometry is asymmetrical or the outer atoms are different, bond dipoles do not cancel, resulting in a polar molecule (e.g., CH3Cl, H2O).

Molecular polarity examples

9.4 Covalent Bonding and Orbital Overlap

Valence bond theory explains covalent bonding as the overlap of atomic orbitals from different atoms. The strength and directionality of bonds depend on the extent and orientation of this overlap.

Sigma (σ) Bonds: Formed by head-on overlap of orbitals along the internuclear axis (e.g., s-s, s-p, or p-p overlap).
Pi (π) Bonds: Formed by side-to-side overlap of p orbitals above and below the internuclear axis; present in double and triple bonds.

Sigma bond formation in H2 Sigma bond formation in HCl Sigma bond formation in Cl2 Sigma and pi bond comparison

9.5 Hybrid Orbitals

Hybridization is the mixing of atomic orbitals on a central atom to form new, equivalent hybrid orbitals that explain observed molecular geometries.

sp Hybridization: Linear geometry, 180° bond angles (e.g., BeF2).
sp2 Hybridization: Trigonal planar geometry, 120° bond angles (e.g., BF3).
sp3 Hybridization: Tetrahedral geometry, 109.5° bond angles (e.g., CH4).

Hybridization of atomic orbitals sp hybrid orbitals Electron configuration of F and Be Be ground state electron configuration Be excited state electron configuration sp hybridization diagram sp hybridization overlap

Summary Table: Hybridization and Geometry

Atomic Orbital Set	Hybrid Orbital Set	Geometry	Examples
s, p	sp	Linear (180°)	BeF2, HgCl2
s, p, p	sp2	Trigonal planar (120°)	BF3, SO3
s, p, p, p	sp3	Tetrahedral (109.5°)	CH4, NH3, H2O

Table 9.4: Geometric arrangements characteristic of hybrid orbital sets

9.6 Multiple Bonds

Multiple bonds consist of one sigma (σ) bond and one or more pi (π) bonds. Sigma bonds are stronger due to greater orbital overlap, while pi bonds are weaker and result from side-to-side overlap of unhybridized p orbitals.

Single Bond: One σ bond (e.g., H–Cl, H–H, Cl–Cl).
Double Bond: One σ bond and one π bond (e.g., O2).
Triple Bond: One σ bond and two π bonds (e.g., N2).

Sigma and pi bond comparison

Example: Formaldehyde (CH2O)

Electron Domain Geometry: Trigonal planar around carbon.
Hybridization: Carbon is sp2 hybridized.
Bonding: C–H bonds are σ bonds (sp2–1s overlap); C=O bond consists of one σ (sp2–sp2) and one π (p–p) bond.

Example: Acetonitrile (CH3CN)

Hybridization: Central carbon is sp hybridized; terminal carbon is sp3 hybridized; nitrogen is sp hybridized.
Bonds: C–H and C–C are σ bonds; C≡N is one σ and two π bonds.

Example: Ethanol (C2H5OH)

Hybridization: All central atoms are sp3 hybridized.
Bonds: All single bonds (σ bonds).

Additional info: The number of electrons in the π-system is equal to the number of electrons in unhybridized p orbitals involved in π bonding. Resonance structures can delocalize π electrons over multiple atoms.