dipole moment
/ DY-pole MOH-ment /
Imagine a tiny molecule as a microscopic tug-of-war over its shared electrons. If one atom pulls harder, the electrons drift toward it, leaving that end slightly negative and the other end slightly positive — like a battery with a plus and a minus end. The dipole moment is the number that measures how strong and which-way that built-in lopsidedness is.
The dipole moment is a vector quantity equal to the size of the separated charge multiplied by the distance between the positive and negative centers. It points from the positive end toward the negative end and is measured in debye units. A molecule has a permanent dipole moment when its bonds are polar and its shape does not let those bond pulls cancel out: water (bent) has a large one, while carbon dioxide (straight, symmetric) has none even though each bond is polar.
Why it matters: the dipole moment tells you at a glance whether a molecule is polar, and polarity governs boiling point, solubility, how molecules line up, and how strongly they feel dipole–dipole and ion–dipole forces. The honest caveat is that a zero overall dipole moment does not mean the bonds are nonpolar — it can simply mean a symmetric shape made the individual bond dipoles cancel.
Carbon dioxide and water both have polar bonds, but CO₂ is straight so its two bond dipoles cancel to give zero, while bent water leaves a strong net dipole moment of about 1.85 debye.
Shape, not just bonds, decides a molecule's net dipole moment.
Bond polarity and molecular dipole moment are not the same. Each bond can have its own dipole, but the molecule's net dipole moment is the vector sum of them all — a symmetric molecule can have very polar bonds yet a zero overall dipole moment.