Molecular Polarity Quiz

Study Notes

Molecules with polar bonds can be non-polar if their geometry is symmetrical, e.g., carbon dioxide (CO₂).
In CO₂, the linear shape causes the dipoles of the C=O bonds to cancel out.

Polar Molecule: Water (H₂O), has a bent shape leading to an uneven charge distribution.
Non-Polar Molecule: Oxygen (O₂), consists of two identical oxygen atoms resulting in an even charge distribution.

Polarity affects solubility significantly; polar molecules dissolve well in polar solvents like water through hydrogen bonding.
Non-polar molecules do not dissolve well in water, forming immiscible layers due to lack of interaction.

Water is polar due to its bent structure and electronegativity difference between hydrogen and oxygen.
This creates a dipole moment with partial negative charge on oxygen and partial positive charges on hydrogen.

Polar Molecules: Experience dipole-dipole interactions and hydrogen bonding (e.g., H₂O).
Non-Polar Molecules: Experience only London dispersion forces, which are weaker, leading to lower boiling/melting points.

Dipole-Dipole Interaction: Hydrogen chloride (HCl) demonstrates interactions between polar molecules due to permanent dipoles.
Hydrogen Bonding: Water (H₂O) exhibits hydrogen bonding, which contributes to its unique properties.
Dispersion Forces: Helium (He) shows weak dispersion forces, the only type present in non-polar molecules.
Ion-Dipole Interaction: Sodium chloride (NaCl) in water involves interactions between ions and polar water molecules.

Symmetrical shapes (linear, tetrahedral) often lead to non-polar molecules as dipoles cancel.
Asymmetrical shapes (bent, trigonal pyramidal) usually result in polar molecules due to unbalanced dipoles.

Have a measurable dipole moment due to uneven electron sharing.
Dissolve well in polar solvents (e.g., water).
Typically possess higher boiling/melting points due to stronger intermolecular forces, e.g., hydrogen bonding.
Examples include water (H₂O), ammonia (NH₃), and hydrogen chloride (HCl).

Lack a permanent dipole moment due to even electron distribution.
Dissolve well in non-polar solvents (e.g., oil).
Generally have lower boiling and melting points compared to polar molecules, influenced by weak London dispersion forces.
Examples include methane (CH₄), carbon dioxide (CO₂), nitrogen (N₂), and oxygen (O₂).

Polar molecules demonstrate higher boiling points due to strong intermolecular forces requiring more energy to overcome.
Non-polar molecules have lower boiling points, as they are held together by weaker van der Waals forces requiring less energy for phase transition.