COVALENT BONDING

Imagine two students who both need the same textbook. Nobody wants to hand it over permanently — they both need it. So they agree to share: the book stays between them, and both can pull it toward themselves whenever they need it. The book now “belongs” to both at once.

Covalent bonds work exactly this way. Instead of one atom stealing an electron from another, both atoms contribute electrons to a shared pair. That pair is simultaneously attracted to both nuclei — and the mutual tug-of-war holds the two atoms together.

Sometimes one shared pair isn't enough to satisfy the octet rule for both atoms. When that happens, atoms share two or even three pairs — forming double or triple bonds. More shared pairs means a stronger, shorter bond.

A Lewis structure is a diagram that accounts for every valence electron in a molecule — both the electrons involved in bonding and the ones sitting alone as lone pairs. The rules are simple: draw each atom, count all available valence electrons, and arrange them so every atom (except hydrogen) ends up with a full octet.

Electrons shared between two atoms are called bond pairs and are drawn as two dots between the atoms (·· ). Electrons that belong to one atom and aren't shared are lone pairs and sit beside their atom (: or ·· on the side).

Dot notation is precise, but in practice chemists switch to line notation — each bond pair of dots becomes a single line (—). A double bond is two lines (═), a triple bond is three (≡). Lone pairs stay as dots. This shorthand is far faster to draw and is the standard you'll see in every textbook and research paper.

Toggle between DOT NOTATION and LINE NOTATION to see how each pair of bonding dots collapses into a single line. Lone pair dots (gray) remain in both views — they are not part of the bond, but they affect the molecule's shape and reactivity.

When the two students sharing a textbook are equally determined, the book stays exactly in the middle. But if one student is more aggressive about pulling it toward themselves, the book drifts to their side. Covalent bonds work the same way — the shared electrons don't always sit perfectly in the center.

How uneven that sharing is determines a bond's polarity — the degree to which electrons are distributed unevenly between the two atoms. We can measure polarity precisely using the dipole moment (symbol μ, pronounced “mu”): the product of the partial charge and the distance between the two charge centers, measured in debyes (D). A bond with perfectly even sharing has μ = 0 D. A bond with uneven sharing has μ > 0 D — the larger the number, the greater the charge separation.

The atomic property that controls polarity is electronegativity — how strongly an atom pulls electrons toward itself. If both atoms have the same electronegativity (like H–H), the electrons split evenly and μ = 0 D: a nonpolar covalent bond. If one atom is more electronegative (like H–Cl, where chlorine pulls harder), the electrons drift toward the stronger atom: a polar covalent bond with μ > 0 D. For reference, H–Cl has μ = 1.08 D and H–F has μ = 1.86 D — fluorine's higher electronegativity creates a larger dipole even though the bond lengths are similar.

In a polar bond, the more electronegative atom picks up a partial negative charge (written δ⁻, “delta minus”) and the other atom gets a partial positive charge (δ⁺). These aren't full charges like Na⁺ or Cl⁻ — they're fractions of a charge, a subtle imbalance in where the electrons spend their time.

Notice how the spectrum is continuous — there's no sharp line between polar covalent and ionic. The cutoff (ΔEN ≈ 1.7) is a useful rule of thumb, not a law of nature. Bonds near the border have both covalent and ionic character.

Chemists use electrostatic potential (ESP) maps to visualize how charge is distributed across a molecule's surface. The surface itself is drawn at a constant electron density — like a skin over the molecule — and then colored by the electric potential at each point:

Compounds held together by covalent bonds are called molecular compounds. Because they don't form charged ions, they behave very differently from ionic compounds like NaCl: