Bond Order, Length, and Energy

A single bond is one sigma bond: one pair of electrons shared directly between the nuclei. A double bond adds a π bond, a second pair of electrons shared in a region above and below the sigma bond. A triple bond adds a third pair, another π bond perpendicular to the first. Each extra shared pair pulls more electron density into the space between the two nuclei. That extra negative charge pulls both positive nuclei inward, shortening the distance between them. It also takes more energy to separate the atoms, because you have to break every one of those shared pairs to pull them apart. So going from a single to a double to a triple bond between the same two atoms, the bond gets shorter and the bond energy goes up.

Carbon-carbon bonds show this cleanly. Ethane (CH₃-CH₃) has a C-C single bond, about 154 pm long, with a bond energy near 347 kJ/mol. Ethene (CH₂=CH₂) has a C=C double bond, shorter at about 134 pm and stronger at about 614 kJ/mol. Ethyne (CH≡CH, acetylene) has a C≡C triple bond, shorter still at about 120 pm and stronger at about 839 kJ/mol. Same two atoms, three different bond orders, a clean trend in both directions.

That trend only holds when you keep the elements fixed. Compare bonds between different elements and length stops predicting strength reliably, because things like lone-pair repulsion change the picture. This sets up why bond order also predicts bond angles and molecular shape later: more electron pairs crowded around a central atom push neighboring groups farther apart, which is where structure and geometry pick up the story.