Why do phosphorus, arsenic, antimony, and bismuth favor single bonds over multiple bonds?

In the nitrogen group, all elements possess a lone pair of electrons in their outermost shells. While this lone pair significantly weakens single bonds in nitrogen—leading it to favor stable, triply bonded diatomic molecules—the bond-weakening effect of the lone pair is far less pronounced for phosphorus, arsenic, antimony, and bismuth.1

Because this destabilizing effect is minimal in these heavier elements, they do not gain the same energetic advantage from forming multiple bonds that nitrogen does. Consequently, these elements favor the formation of single bonds and typically exist as condensed, singly bonded structures rather than as diatomic molecules.1