Hydrogen and the Binary Hydrides

Core Idea

Hydrogen forms binary compounds with most elements, and these hydrides fall into three broad classes set by the partner's electronegativity. Very electropositive metals form saline (ionic) hydrides that contain the hydride ion H⁻. The p-block elements form molecular hydrides held together by covalent bonds. Many d- and f-block metals take up hydrogen into their lattices to form metallic hydrides, often with nonstoichiometric formulas.

How It's Best Learned

Place sample hydrides on a periodic table, NaH and CaH₂ on the left, PdHₓ in the d-block, and CH₄, NH₃, H₂O and HF on the right. For each class, predict the oxidation state of hydrogen, the physical state, and the reaction with water. Then plot boiling points of the group 14 to 17 hydrides to see where hydrogen bonding takes over.

Common Misconceptions

• Hydrogen always has oxidation state +1. (In saline hydrides such as NaH and CaH₂, hydrogen is −1, and these compounds release H₂ when they react with water.)
• Hydrides are all gases like CH₄ and NH₃. (Saline hydrides are high-melting ionic solids, and metallic hydrides are solids that often conduct electricity.)
• Metallic hydrides are ordinary compounds with fixed formulas. (Many have variable composition, such as PdHₓ, because hydrogen occupies holes in the metal lattice to different extents.)

Explainer

Hydrogen is the simplest element and the most abundant in the universe. It has three isotopes: protium (¹H), deuterium (²H or D) and radioactive tritium (³H). Its electronegativity sits in the middle of the scale, so it can lose its electron, share it, or gain one. That flexibility is why it forms compounds with almost every element.

Binary hydrides, compounds of hydrogen with one other element, divide into three classes. The class depends mainly on how electronegative the partner is.

Saline, or ionic, hydrides form with the alkali metals and the heavier alkaline earth metals. Examples are LiH, NaH and CaH₂. These metals transfer an electron to hydrogen, giving the hydride ion H⁻. The compounds are white crystalline solids with high melting points. H⁻ is a strong base and a good reducing agent, so saline hydrides react with water to release hydrogen gas. For example, NaH + H₂O → NaOH + H₂. Chemists use this reaction to dry solvents.

Molecular, or covalent, hydrides form with the p-block elements. These include CH₄, NH₃, H₂O, HF, H₂S and the boranes. They can be sorted by electron count. Electron-deficient hydrides, such as diborane B₂H₆, have too few electrons for ordinary two-electron bonds and use bridging hydrogens. Electron-precise hydrides, such as CH₄ and SiH₄, have exactly enough electrons for their bonds. Electron-rich hydrides, such as NH₃, H₂O and HF, carry lone pairs on the central atom. Those lone pairs and the very polar bonds allow hydrogen bonding. Hydrogen bonding explains why NH₃, H₂O and HF boil far above the trend set by the heavier hydrides of their groups.

Metallic hydrides form with many d- and f-block metals. Hydrogen enters holes in the metal lattice, and the product often keeps a metallic luster and conducts electricity. Compositions can vary continuously, as in PdHₓ. Palladium can absorb many times its own volume of hydrogen, and metallic hydrides have been studied as a way to store hydrogen for fuel.

Industrial hydrogen comes mostly from fossil sources. In steam reforming, methane reacts with steam at high temperature over a nickel catalyst: CH₄ + H₂O → CO + 3H₂. The water-gas shift reaction, CO + H₂O → CO₂ + H₂, then raises the hydrogen yield. Electrolysis of water is a smaller but growing source.