Molecular geometry describes how atoms arrange themselves in 3D space, while polarity describes how electrical charge is distributed within a molecule. These two ideas are tightly connected: the shape of a molecule determines whether it's polar or nonpolar, which in turn controls physical properties like boiling point and solubility.

This section covers VSEPR theory (the main tool for predicting molecular shapes), common geometries you need to know, and how to determine whether a molecule is polar or nonpolar.

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Molecular Geometry

VSEPR Theory and Molecular Shapes

VSEPR stands for Valence Shell Electron Pair Repulsion. The core idea is simple: electron pairs around a central atom repel each other, so they spread out as far apart as possible. The resulting arrangement determines the molecule's shape.

Two types of electron pairs matter here:

Bonding pairs form covalent bonds between atoms.
Lone pairs sit on the central atom without bonding to anything. They still take up space and push other electron pairs away.

To use VSEPR, count the total number of electron domains around the central atom. An electron domain is any region of electron density: a single bond, a double bond, a triple bond, or a lone pair. A double bond counts as one domain, not two. The total number of domains tells you the electron geometry, and the arrangement of just the atoms (ignoring lone pairs) gives you the molecular geometry.

VSEPR Theory and Molecular Shapes, Molecular Structure and Polarity (4.6) – Chemistry 110

Common Molecular Geometries

Linear — 2 electron domains, 0 lone pairs

Bond angle: 180°
The atoms sit in a straight line.
Examples: $CO_2$ , $HCN$

Trigonal Planar — 3 electron domains, 0 lone pairs

Bond angle: 120°
All atoms lie in a flat plane, evenly spaced.
Examples: $BF_3$ , $CO_3^{2-}$

Bent (from 3 domains) — 3 electron domains, 1 lone pair

Bond angle: less than 120° (the lone pair squeezes the bonding pairs closer together)
Example: $SO_2$

Tetrahedral — 4 electron domains, 0 lone pairs

Bond angle: 109.5°
Electron domains point toward the four corners of a tetrahedron.
Examples: $CH_4$ , $NH_4^+$

Trigonal Pyramidal — 4 electron domains, 1 lone pair

Bond angle: about 107°
Looks like a tetrahedron with one corner "missing" where the lone pair sits.
Example: $NH_3$

Bent (from 4 domains) — 4 electron domains, 2 lone pairs

Bond angle: about 104.5°
Example: $H_2O$

Water ( $H_2O$ ) is a common source of confusion. It has 4 electron domains (2 bonding pairs + 2 lone pairs), so its electron geometry is tetrahedral, but its molecular geometry is bent because you only describe the positions of the atoms, not the lone pairs. The 2 lone pairs compress the bond angle from 109.5° down to about 104.5°.

Factors Influencing Molecular Shape

Lone pairs repel more strongly than bonding pairs. Lone pair electrons are held closer to the central atom, so they push harder on neighboring pairs. This is why lone pairs reduce bond angles below the "ideal" values.
Multiple bonds count as a single electron domain in VSEPR. A double or triple bond occupies more space than a single bond, but you still count it as one domain when predicting geometry.
Larger central atoms can accommodate more electron domains, which is why elements in lower periods can sometimes have 5 or 6 domains (expanded octets).

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Molecular Polarity

Understanding Polarity

Polarity refers to an uneven distribution of electrical charge within a molecule. It arises from electronegativity differences between bonded atoms. When one atom pulls electrons more strongly than the other, the bond becomes polar, with a partial negative charge ( $\delta^-$ ) on the more electronegative atom and a partial positive charge ( $\delta^+$ ) on the less electronegative one.

A few key points about electronegativity:

Electronegativity generally increases across a period (left to right) and decreases down a group on the periodic table.
A bond is considered polar covalent when the electronegativity difference between the two atoms is roughly between 0.5 and 1.7 on the Pauling scale.
Below about 0.5, the bond is essentially nonpolar covalent. Above about 1.7, it's typically ionic.

But here's the critical part: having polar bonds does not automatically make a polar molecule. You also need to consider the geometry.

Determining Whether a Molecule Is Polar or Nonpolar

To figure out if a molecule is polar, follow these steps:

Draw the Lewis structure and identify polar bonds (look for electronegativity differences between bonded atoms).
Determine the molecular geometry using VSEPR.
Check for symmetry. If the polar bonds are arranged symmetrically, their dipoles cancel out and the molecule is nonpolar. If the arrangement is asymmetrical, the dipoles don't cancel, and the molecule is polar.

Polar molecules have a net dipole moment due to uneven charge distribution:

$H_2O$ is polar. It has polar O–H bonds, and its bent shape means the bond dipoles point in roughly the same direction instead of canceling.
$NH_3$ is polar. Its trigonal pyramidal shape (3 bonding pairs + 1 lone pair) creates an asymmetric charge distribution.

Nonpolar molecules have symmetrical charge distribution:

$CO_2$ is nonpolar. Even though each C=O bond is polar, the linear geometry means the two bond dipoles point in exactly opposite directions and cancel out.
$CH_4$ is nonpolar. The four C–H bonds are arranged in a perfect tetrahedron, so the dipoles cancel. (The C–H electronegativity difference is small to begin with, and the symmetry eliminates whatever slight polarity exists.)

Quick rule of thumb: if a molecule has polar bonds and lone pairs on the central atom, it's almost always polar. The lone pairs break the symmetry.

Dipole Moments and Their Significance

The dipole moment is a measure of how polar a molecule is. It's a vector quantity, meaning it has both magnitude and direction. Dipole moments are measured in Debye units (D).

Each polar bond has its own bond dipole. The net dipole moment is the vector sum of all bond dipoles in the molecule.
If the vectors cancel (symmetric geometry), the net dipole moment is zero and the molecule is nonpolar.
If they don't cancel, the molecule has a nonzero dipole moment and is polar.

Why does this matter? Polarity directly affects physical properties:

Boiling point: Polar molecules experience dipole-dipole attractions (a type of intermolecular force), so they tend to have higher boiling points than nonpolar molecules of similar size.
Solubility: "Like dissolves like." Polar molecules dissolve well in polar solvents (like water). Nonpolar molecules dissolve well in nonpolar solvents (like hexane or oil).
Intermolecular forces: Dipole-dipole interactions between polar molecules are stronger than the London dispersion forces that act between nonpolar molecules of comparable size. This difference in intermolecular force strength is what drives the boiling point and solubility trends above.

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