Covalent Bonds: Sharing Pairs
When neither atom will surrender an electron, they compromise: shared pairs that count for both. Welcome to molecules.
Builds on: 1.1 Ionic Bonds: Give and Take
The non-metal compromise
Two chlorine atoms meet. Both are one electron short; neither will give one up. The way out: each contributes one electron to a shared pair that sits between the two nuclei and — here’s the trick — counts toward both atoms’ shells at once. Each Cl now “sees” 8 outer electrons. This shared pair is a covalent bond, and the bonded unit is a molecule.
Single, double, triple
One shared pair makes a single bond (H–H). If one pair isn’t enough, atoms share two (a double bond, O=O) or even three (a triple bond, N≡N). More shared pairs bind tighter: nitrogen’s triple bond is so strong that N₂ — 78% of the air you breathe — behaves almost like a noble gas. Breaking it to make fertilizer (the Haber process) consumes about 1% of humanity’s entire energy supply.
Reading a molecule
Carbon has 4 valence electrons and needs 4 more, so it makes four bonds — the four hands that let it build chains, rings and eventually DNA. Oxygen makes two bonds, nitrogen three, hydrogen exactly one. With just that counting rule you can predict the shape of most small molecules: H₂O is oxygen with two hands holding two hydrogens; CO₂ is carbon double-handing two oxygens; CH₄ is carbon holding four hydrogens.
Electrons that stay home (not shared) are lone pairs. They take up space and shove the bonds around — that’s why water is bent at ~104.5° rather than straight. And a bent molecule with unevenly shared electrons becomes a little magnet… which is the next lesson, and ultimately the reason ice floats and life works.
⚗️ Lab — The Molecule Workshop
Six molecules, from a single shared pair to a triple bond.
- Check the octet counter under every atom: own electrons + shared pairs.
- Compare O₂ and N₂ — count the cyan pairs in the bond.
- Look at H₂O’s violet lone pairs pushing the molecule into its bent shape.