Molecular Structure

VSEPR (Valence Shell Electron Pair Repulsion) theory predicts the three-dimensional shape of a molecule by assuming that electron groups around a central atom repel each other and arrange themselves as far apart as possible to minimize repulsion.

An electron group is any region of electron density bonded to the central atom: a single bond, a double bond, a triple bond, or a lone pair. All of these count equally when determining the arrangement. The number of electron groups around the central atom determines the electron-pair geometry — the spatial arrangement of all electron groups, including lone pairs.

To apply VSEPR: (1) draw the Lewis structure, (2) count electron groups on the central atom, (3) identify the electron-pair geometry, and (4) name the molecular geometry based on the positions of the atoms only (ignoring lone pairs in the name).

The five fundamental electron-pair geometries correspond to 2 through 6 electron groups:

When all electron groups are bonding pairs, the molecular geometry matches the electron-pair geometry. When lone pairs are present, the molecular geometry is a subset: for example, four electron groups with one lone pair give a trigonal pyramidal shape (NH₃), and four groups with two lone pairs give a bent shape (H₂O).

Lone pairs occupy more space than bonding pairs because they are held closer to the nucleus and spread out more broadly. This extra repulsion compresses bond angles below their ideal values. In methane (CH₄), all four groups are bonding and angles are 109.5°. In ammonia (NH₃), one lone pair pushes the three N–H bonds closer together, reducing the angle to about 107°. In water (H₂O), two lone pairs compress the H–O–H angle further to about 104.5°.

For five electron groups, removing bonding pairs in favor of lone pairs produces progressively distorted shapes: seesaw (4 bonds, 1 lone pair), T-shaped (3 bonds, 2 lone pairs), and linear (2 bonds, 3 lone pairs). For six groups: square pyramidal (5 bonds, 1 lone pair) and square planar (4 bonds, 2 lone pairs). Lone pairs preferentially occupy equatorial positions in trigonal bipyramidal systems because equatorial sites have more room.

A molecule is polar if it has a net dipole moment — an overall uneven distribution of electron density. Two conditions must both be met:

1. The molecule must contain polar bonds (bonds between atoms with different electronegativities, producing individual bond dipoles).
2. The molecular geometry must be asymmetric so that the individual bond dipoles do not cancel.

Symmetry is the key factor. CO₂ has two polar C=O bonds, but its linear geometry makes the dipoles point in opposite directions and cancel — the molecule is nonpolar. H₂O also has two polar bonds, but its bent shape leaves a net dipole pointing from H toward O — the molecule is polar.

Similarly, CCl₄ (tetrahedral, all identical bonds) is nonpolar because of perfect symmetry, while CHCl₃ (one H replaces a Cl) is polar because the symmetry is broken. Molecular polarity governs many physical properties including boiling point, solubility, and interactions with other molecules.

Hybridization is the mixing of standard atomic orbitals (s, p, d) on a central atom to form a new set of equivalent hybrid orbitals oriented to match the observed molecular geometry. The type of hybridization is determined by the number of electron groups:

The quick rule: count all electron groups (bonding and lone pairs) on the central atom. Two groups = sp, three = sp², four = sp³, and so on. Water has four electron groups (2 bonds + 2 lone pairs), so oxygen is sp³ hybridized even though its molecular shape is bent.

Hybrid orbitals form sigma (σ) bonds — bonds with electron density concentrated directly between the two nuclei along the internuclear axis. Every single bond is a sigma bond, and every double or triple bond contains exactly one sigma bond.

The additional bonds in a double or triple bond are pi (π) bonds, formed by the side-by-side overlap of unhybridized p orbitals. A double bond consists of one σ bond + one π bond. A triple bond consists of one σ bond + two π bonds.

In ethylene (C₂H₄), each carbon is sp² hybridized with three sigma bonds (two C–H and one C–C). The remaining unhybridized p orbital on each carbon overlaps to form the pi bond, locking the molecule into a planar geometry. Pi bonds prevent rotation around the bond axis, which is why geometric (cis/trans) isomerism occurs in molecules with C=C double bonds.

For any molecule, a systematic four-step approach connects Lewis structure to full three-dimensional properties:

1. Draw the Lewis structure to identify bonding pairs, lone pairs, and multiple bonds on the central atom.
2. Count electron groups to determine the electron-pair geometry and hybridization.
3. Identify the molecular geometry by considering only the positions of atoms (excluding lone pairs from the shape name).
4. Assess polarity by checking whether bond dipoles cancel given the molecular geometry.

For example, sulfur dioxide (SO₂): the Lewis structure shows two double bonds and one lone pair on S (3 electron groups). Electron-pair geometry: trigonal planar. Hybridization: sp². Molecular geometry: bent (2 atoms + 1 lone pair). Polarity: the S=O dipoles do not cancel in the bent arrangement, so SO₂ is polar. This integrated approach works for any molecule.

A systematic approach to molecular geometry and polarity:

1. Draw the Lewis structure first. Geometry predictions depend on knowing the number of electron groups.
2. Count electron groups (bonding pairs + lone pairs) around the central atom. Double and triple bonds each count as one group.
3. Determine electron-pair geometry from the total number of electron groups (2 = linear, 3 = trigonal planar, 4 = tetrahedral, 5 = trigonal bipyramidal, 6 = octahedral).
4. Determine molecular geometry by considering only the positions of atoms (not lone pairs).
5. Assess polarity: if all outer atoms are identical and there are no lone pairs on the central atom, the molecule is likely nonpolar. Otherwise, check whether dipole moments cancel.

Common mistakes: counting a double bond as two electron groups, confusing electron-pair geometry with molecular geometry, assuming all tetrahedral molecules are nonpolar (bent and trigonal pyramidal are subsets of tetrahedral electron-pair geometry but are polar), and misidentifying hybridization by not matching it to the electron-pair geometry.