Molecular Polarity

Molecule Type

Polarity Result

POLAR

Dipole Moment (μ): 0.00 D

Geometry: Linear

Bond Angle: 180°

Bond Polarity: Polar

Vector Sum: ≠ 0

Parameters

Electronegativity Difference (ΔEN): 0.9

Bond Length: 1.0 Å

Rotation Speed: 1.0x

Molecule Scale: 1.0x

Show Bond Dipole Vectors Show Total Dipole Moment Show Atom Labels Show Bonds

Physical Equations

Dipole Moment: μ = q × d

Bond Polarity Rule: ΔEN > 0.4 (polar bond)

Molecular Polarity Rule: μ_total = Σμ_bond ≠ 0 (polar molecule)

Legend

Partial Positive (δ+)

Partial Negative (δ-)

Dipole Vector

What is Molecular Polarity?

Molecular polarity is a measure of the uneven distribution of electron density in a molecule. It arises from the difference in electronegativity between atoms and the molecular geometry. A polar molecule has a net dipole moment, while a nonpolar molecule has either no polar bonds or symmetrical polar bonds that cancel each other out.

Bond Polarity

A covalent bond becomes polar when there is a significant difference in electronegativity (ΔEN > 0.4) between the bonded atoms. The more electronegative atom attracts the shared electrons more strongly, creating a partial negative charge (δ-) on that atom and a partial positive charge (δ+) on the other. This creates a bond dipole with magnitude μ = q × d, where q is the partial charge and d is the bond length.

Molecular Geometry and Polarity

The molecular geometry determines whether individual bond dipoles cancel out or add together. Symmetrical molecules like CO₂ (linear) and CCl₄ (tetrahedral) have polar bonds but are nonpolar overall because the bond dipoles cancel in opposite directions. Asymmetrical molecules like H₂O (bent) and NH₃ (trigonal pyramidal) are polar because the bond dipoles don't completely cancel.

Dipole Moment

The dipole moment (μ) is a vector quantity that measures the separation of positive and negative charges in a molecule. It's measured in Debye units (D). The total dipole moment of a molecule is the vector sum of all individual bond dipoles: μ_total = Σμ_bond. A molecule is polar if μ_total ≠ 0. Water has a dipole moment of 1.85 D, making it a highly polar molecule that dissolves many ionic compounds.

Molecular Examples

HCl: Diatomic molecule with ΔEN = 0.9, highly polar (μ = 1.08 D)

CO₂: Linear molecule with polar bonds that cancel, nonpolar (μ = 0 D)

H₂O: Bent geometry with bond angle 104.5°, polar (μ = 1.85 D)

CH₄: Tetrahedral symmetry, nonpolar despite polar C-H bonds (μ = 0 D)

NH₃: Trigonal pyramidal, polar (μ = 1.47 D)

Applications

Molecular polarity is crucial in understanding: solubility ("like dissolves like"), intermolecular forces (dipole-dipole interactions), boiling and melting points, surface tension, and chemical reactivity. Polar molecules like water are excellent solvents for ionic and polar substances, while nonpolar molecules like hexane dissolve nonpolar compounds. Polarity also affects drug design, protein folding, and membrane permeability in biological systems.