covalent (molecular) hydride

Think of water, ammonia, methane, and hydrogen sulfide — the small molecules that fill chemistry. All of them are hydrides: compounds of hydrogen with another element. But here the metal-like generosity is gone. When hydrogen meets a nonmetal of similar electronegativity, neither side surrenders an electron outright. Instead they share, forming discrete covalent molecules. These are the covalent or molecular hydrides, and they are by far the most common kind.

Because the bonding is covalent, these hydrides exist as separate molecules held to one another only by relatively weak forces, so most are gases or low-boiling liquids — H2O, NH3, CH4, HCl, H2S, PH3. The hydrogen here carries a small positive oxidation state (typically +1) because the partner nonmetal pulls electron density toward itself. A striking subgroup are the electron-deficient hydrides such as the boron hydrides (boranes), where there are not enough electrons to draw a normal two-centre bond between every pair of atoms, forcing exotic three-centre two-electron bonds. When the partner atom is small and very electronegative — N, O, F — the molecules can also link to each other through hydrogen bonds, which is why water boils so much higher than its size would suggest.

Covalent hydrides matter because they are the working substances of life and industry: water itself, ammonia (made on a vast scale by the Haber process for fertilizer), the simple alkanes that fuel us, the acids HF, HCl, and H2S. The honest nuance is that the line between ionic and covalent hydrides is a gradient, not a wall — beryllium and magnesium hydrides sit in between, polymeric and partly covalent, reminding you that bonding type is a spectrum set by the electronegativity difference.