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Deprotonation is the removal (transfer) of a proton (a hydrogen cation, H+) from a Bronsted–Lowry acid in an acid-base reaction. The species formed is the conjugate base of that acid. The complementary process, when a proton is added (transferred) to a Bronsted–Lowry base, is protonation. The species formed is the conjugate acid of that base.
A species that can either accept or donate a proton is referred to as amphiprotic. An example is the H2O (water) molecule, which can gain a proton to form the hydronium ion, H3O+, or lose a proton, leaving the hydroxide ion, OH−.
The relative ability of a molecule to give up a proton is measured by its pKa value. A low pKa value indicates that the compound is acidic and will easily give up its proton to a base. The pKa of a compound is determined by many things, but the most significant is the stability of the conjugate base. This is primarily determined by the ability (or inability) of the conjugated base to stabilize negative charge. One of the most important ways of assessing a conjugate base's ability to distribute negative charge is using resonance. Electron withdrawing groups (which can stabilize the molecule by increasing charge distribution) or electron donating groups (which destabilize by decreasing charge distribution) present on a molecule also determine its pKa. The solvent used can also assist in the stabilization of the negative charge on a conjugated base.
Bases used to deprotonate depend on the pKa of the compound. When the compound is not particularly acidic, and, as such, the molecule does not give up its proton easily, a base stronger than the commonly known hydroxides is required. Hydrides are one of the many types of powerful deprotonating agents. Common hydrides used are sodium hydride and potassium hydride. The hydride forms hydrogen gas with the liberated proton from the other molecule. The hydrogen is dangerous and could ignite with the oxygen in the air, so the chemical procedure should be done in an inert atmosphere (e.g., nitrogen).
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