How do dipole moments affect molecular polarity and properties?

The net electric dipole moment (p) is the definitive measure of a molecule's polarity. This polarity is one of the most important factors determining the physical properties of a substance.

1. Defining Polarity:
   • Polar Molecules (p ≠ 0): Molecules with a non-zero net dipole moment (like water, H₂O, or ammonia, NH₃) are polar. They have a permanent "positive end" and "negative end."
   • Non-polar Molecules (p = 0): Molecules with a zero net dipole moment (like CO₂ or CH₄) are non-polar. This happens either because their bonds are non-polar (e.g., N₂) or their polar bonds are arranged symmetrically and cancel out (e.g., CO₂).
2. Effect on Physical Properties (Intermolecular Forces):
   • Boiling and Melting Points: Polar molecules attract each other via strong dipole-dipole interactions (the δ+ of one molecule attracts the δ- of another). These are much stronger than the weak London dispersion forces that hold non-polar molecules together.
   • Result: It takes more energy (a higher temperature) to separate polar molecules. This is why water (H₂O, p=1.85 D) has a boiling point of 100°C, while methane (CH₄, p=0 D), which has a similar mass, has a boiling point of -161.5°C.
   • Solubility (Rule: "Like Dissolves Like"):
     • Polar solvents (like water) are excellent at dissolving other polar substances (like sugar) and ionic compounds (like salt, NaCl).
     • Non-polar solvents (like oil or hexane) are used to dissolve other non-polar substances (like grease).
     • This is why oil (non-polar) and water (polar) do not mix.