A polar covalent bond forms when two atoms share electrons unequally: the more electronegative atom pulls electron density toward itself and becomes δ−, leaving the other atom δ+. That separation of charge is a bond dipole, a vector pointing from δ+ to δ−.
A molecule's net dipole moment is the vector sum of all its bond dipoles (plus lone-pair contributions), so shape decides polarity. Symmetric molecules like CO₂ and CCl₄ have bond dipoles that cancel and are nonpolar, while bent water and asymmetric molecules keep a net dipole and are polar.
1. Electronegativity makes a bond polar
Every covalent bond is a tug-of-war for the shared electron pair. Electronegativity measures how strongly an atom pulls that pair toward itself. When two bonded atoms have different electronegativities, the electrons spend more time near the greedier atom. That atom takes on a partial negative charge, written δ−, and its partner is left electron-poor, written δ+.
The bigger the electronegativity gap, the more polar the bond. In chloromethane, chlorine (3.2) outpulls carbon (2.6), so the C–Cl bond is clearly polar with Cl as δ−. In hydrogen fluoride, fluorine is the most electronegative element of all, giving one of the most polar single bonds.
2. A bond dipole is a vector from δ+ to δ−
Chemists picture the charge separation as an arrow: the bond dipole points from the δ+ atom toward the δ− atom (the crossed end sits on the positive side). Its length represents the dipole's magnitude, which grows with both the size of the charge separation and the bond length. Because it has both a direction and a magnitude, a dipole is a vector — and vectors add head-to-tail.
3. The molecular dipole is the vector sum of bond dipoles
Most molecules have several polar bonds. To find the net (molecular) dipole moment, you add every bond dipole as a vector and include any push from lone pairs. If the arrows partly oppose one another, they subtract; if they reinforce, they add. This is why you can never judge a whole molecule from a single bond in isolation — geometry sets the directions of the arrows.
Water is the classic example. Its two O–H bond dipoles both point toward the δ− oxygen, and because the molecule is bent (about 104.5°), those two arrows do not cancel. They add to a strong net dipole pointing from the H's up through the oxygen, reinforced by oxygen's two lone pairs. Ammonia is similar — three N–H dipoles plus a lone pair give a pyramidal molecule a clear net dipole.
4. Symmetry makes bond dipoles cancel
When polar bonds are arranged symmetrically, their dipole vectors point in balanced, opposite directions and sum to zero. The molecule then has polar bonds but no net dipole — it is nonpolar overall. Carbon dioxide is linear, so its two C=O dipoles point in exactly opposite directions and cancel. Carbon tetrachloride and tetrafluoromethane are tetrahedral; their four identical bond dipoles pull outward symmetrically and cancel. Methane is nonpolar for the same reason (and its C–H bonds are barely polar to begin with).
Break the symmetry and the cancellation fails. Dichloromethane has two C–Cl and two C–H bonds on the same tetrahedral carbon; the Cl dipoles no longer have partners pointing the opposite way, so a net dipole survives and the molecule is polar. Chloroform is polar for the same reason.
5. Polarity controls boiling point and solubility
Polar molecules attract one another through their opposite ends (dipole–dipole forces, and hydrogen bonding when H sits on N, O, or F). These extra attractions take energy to overcome, so polar substances tend to have higher boiling points than nonpolar molecules of similar size. They also follow the rule "like dissolves like": polar solvents such as water, methanol, and ethanol dissolve polar and ionic solutes, while nonpolar solvents like carbon tetrachloride dissolve fats and oils.
6. The δ+ carbon is where organic reactions begin
Polarity is not just about physical properties — it is the roadmap for reactivity. A polar bond leaves the δ+ atom electron-poor and electrophilic, so electron-rich nucleophiles attack there. In a carbonyl group the oxygen is δ− and the carbon is a strongly δ+ electrophilic site; nucleophiles add straight to that carbon. In an alkyl halide the halogen is δ− and the carbon it is bonded to is δ+, the target of substitution and elimination reactions.
7. Summary
Bond polarity comes from an electronegativity difference: the more electronegative atom is δ−, its partner δ+, and the bond carries a dipole vector pointing from δ+ to δ−. A molecule's net dipole is the vector sum of those bond dipoles plus lone-pair effects, so shape is decisive: symmetric molecules (CO₂, CCl₄, CF₄, CH₄) cancel to nonpolar, while bent, pyramidal, or otherwise asymmetric molecules (water, ammonia, CH₂Cl₂, chloroform) keep a net dipole and are polar. Polarity raises boiling points, governs solubility through "like dissolves like," and — most importantly for organic chemistry — marks the δ+ carbon as the electrophilic site where nucleophiles attack.
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Shape. CO₂ is linear, so its two C=O bond dipoles point in exactly opposite directions and cancel to zero. Water is bent (~104.5°), so its two O–H dipoles do not cancel — they add to a net dipole, making water polar.
The arrow points from the δ+ (less electronegative) atom toward the δ− (more electronegative) atom. Its length represents the dipole's magnitude, which increases with a larger charge separation and a longer bond.
CCl₄ is tetrahedrally symmetric: four identical C–Cl dipoles pull outward evenly and sum to zero. In CH₂Cl₂ the two C–Cl dipoles are no longer balanced by opposing C–Cl bonds (the other two positions hold weakly polar C–H bonds), so a net dipole survives.
The polar bond leaves that carbon electron-poor and electrophilic, so electron-rich nucleophiles attack there. This is the starting point for nucleophilic addition to C=O carbonyls and for substitution/elimination at the C bonded to a halogen.
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