The hydride ion (H⁻) is a Lewis base. It has a lone pair of electrons it can donate to electron-poor atoms, which is the defining feature of a Lewis base. This makes it the chemical opposite of the proton (H⁺), which is one of the simplest Lewis acids because it has an empty orbital ready to accept electrons.
Why H⁻ Qualifies as a Lewis Base
A Lewis base donates an electron pair. A Lewis acid accepts one. The hydride ion is a hydrogen atom that has gained an extra electron, giving it two electrons in its 1s orbital and a negative charge. That filled orbital is what makes it a Lewis base: it has electrons available to share with any atom that needs them.
Compare this to the proton, H⁺. When hydrogen loses its one electron, it’s left with a completely empty 1s orbital. That empty orbital is eager to accept a pair of electrons, making H⁺ a textbook Lewis acid. So hydrogen in its two ionic forms sits on opposite ends of Lewis acid-base chemistry: H⁺ is the acid, H⁻ is the base.
How H⁻ Behaves as a Base in Reactions
The hydride ion shows up across organic and inorganic chemistry as an electron pair donor, and its behavior consistently matches the Lewis base classification.
In organic chemistry, reagents like lithium aluminum hydride (LiAlH₄) and sodium borohydride (NaBH₄) work by delivering a hydride ion to a carbon atom that carries a partial positive charge, such as the carbon in a carbonyl group (C=O). The hydride donates its electron pair directly to that carbon in what’s called a nucleophilic attack. The result is that an aldehyde or ketone gets converted into an alcohol. The hydride is acting as both a Lewis base and a nucleophile here, since nucleophiles are simply Lewis bases that donate electrons to atoms other than hydrogen.
Sodium hydride (NaH) is one of the most common bases in organic chemistry labs. It works by donating its electron pair to a proton on molecules like alcohols, phenols, and ketones, pulling that proton off in a deprotonation reaction. This is a classic Brønsted base behavior, but it’s also Lewis base behavior: the hydride’s electron pair forms a bond with H⁺, which is a Lewis acid.
H⁻ in Coordination Chemistry
When hydride ions bond to transition metals, they act as ligands, donating their electron pair to the metal center. Metal hydride complexes are widespread in catalysis and industrial chemistry. The hydride serves as a Lewis base by sharing its electrons with the metal, which acts as a Lewis acid. These metal-hydride bonds can then be influenced by additional Lewis acids in what chemists call secondary coordination sphere interactions, where nearby electron-accepting molecules affect how reactive the hydride is.
H⁻ in Biological Systems
Your cells use hydride transfer constantly. The molecule NADH, often called the “energy currency” of metabolism, carries a hydride ion on its nicotinamide ring. During cellular energy production, NADH donates that hydride to other molecules. In mitochondrial complex I, for example, the hydride transfers from a specific position on the nicotinamide ring to a flavin molecule, a step that’s partially rate-limiting for the entire reaction. This biological hydride transfer follows the same Lewis base logic: the hydride’s electron pair moves to an electron-accepting site on another molecule.
How Strong a Lewis Base Is H⁻?
The hydride ion is a powerful Lewis base. One way chemists measure this is through hydride affinity, which quantifies how strongly a given Lewis acid attracts a hydride ion. Gas-phase calculations show that for common carbon-based cations (carbenium ions), hydride affinities run 35 to 60 kcal/mol higher than fluoride affinities. Since fluoride is itself a strong Lewis base, this tells you that H⁻ is an even more aggressive electron pair donor in many contexts.
This strong basicity is also why sodium hydride can deprotonate relatively weak acids. Acetonitrile, for instance, has a pKa around 25, making it quite reluctant to give up a proton. Yet NaH can pull it off. That reactivity sometimes causes complications in lab settings, where NaH may deprotonate unintended targets or even act as a reducing agent alongside its role as a base.
The Dihydrogen Bond: H⁻ Meets H⁺
One of the more interesting consequences of hydride’s Lewis base character is the dihydrogen bond. This forms when a negatively charged hydrogen (acting as the Lewis base) interacts with a positively charged hydrogen on a nearby molecule (acting as the Lewis acid). The interaction looks like H+δ···−δH, where the two hydrogens of opposite polarity attract each other. In some cases this interaction is reversible, but in others it leads to the release of molecular hydrogen gas (H₂) as the two hydrogens combine and leave their parent molecules behind.