The term Lewis acid is named after Gilbert N. Lewis. Early definitions of acids and bases were based on reactions with metals and alkalis. Arrhenius defined acids as substances that form H+ ions in water, while Bronsted-Lowry defined acids as proton donors. Lewis defined acids as electron pair acceptors and bases as electron pair donors. This definition includes all Arrhenius and Bronsted-Lowry acids and also many substances that do not meet their criteria. Boron trifluoride is an example of a Lewis acid that can form a coordinated bond with a Lewis base like ammonia.
The term Lewis acid is named after the American chemist Gilbert N. Lewis. Early chemists recognized an acid as a sour-tasting substance that reacts with some metals and neutralizes the bases, or alkalis, producing a salt. Since the late 19th century, however, attempts have been made to define acids and bases more rigorously and to explain what actually happens in an acid-base reaction. Lewis’s is the broadest definition.
In 1883, the Swedish chemist Svante Arrhenius defined an acid as a substance which forms hydrogen ions (H+) in aqueous solution and a base as a substance which forms hydroxide ions (OH-). H+ ions, which are simply protons, are too reactive to exist in an aqueous solution and associate with water molecules to form hydronium ions (H3O+). Arrhenius’ definition has proved very useful and covers most of the compounds commonly considered acids. For example, hydrochloric acid, a solution of hydrochloric acid gas in water, gives H+ ions which form hydronium ions in solution: HCl + H2O → H3O+ + Cl-. This definition remained the standard well into the 20th century and is still frequently used today.
A distinguishing feature of all acids is that they neutralize bases to produce salts. An example is the reaction of hydrochloric acid with sodium hydroxide (NaOH) to produce sodium chloride and water (H2O): H3O+Cl- + Na+OH- → Na+Cl- + H2O. Here, the H+ ions supplied by hydrochloric acid combined with the OH- ions supplied by sodium hydroxide to produce water, while the Na+ and Cl- ions combined to produce salt, according to Arrhenius’ theory; however, similar reactions can occur between compounds that do not fit Arrhenius’ definitions of acids and bases. For example, hydrochloric acid gas can react with ammonia gas to form the ammonium chloride salt: HCl + NH3 → NH4+Cl-. Two compounds have combined to form a salt, but since they are not in solution, there are no H+ or OH- ions present, so the reactants do not qualify as Arrhenius acid and base.
In 1923, two chemists, Johaness Bronsted and Thomas Lowry, independently came up with a new definition. They suggested that an acid was a proton donor and a base a proton acceptor. In an acid-base reaction, the acid contributes a proton, or H+ ion, to the base; however, neither reactant needs to be in solution, with the H+ or OH- ions actually present prior to the reaction. This definition includes all Arrhenius acids and bases, but also explains the combination of gaseous hydrochloric acid and ammonia as an acid-base reaction: the covalent hydrochloric acid gave a proton to ammonia to form an ammonium ion (NH4+), which forms an ionic compound with the Cl ion.
The American chemist Gilbert N. Lewis suggested, also in 1923, an extended concept of acids and bases as electron pair acceptors and donors, respectively. By this definition, an acid-base reaction involves the reactants forming a coordinate bond — a covalent bond in which both shared electrons come from the same atom — with the electrons from the base. In the HCl-NaOH reaction described above, the H+ ion supplied by HCl accepts an electron pair from the OH- ion supplied by NaOH to form water.
According to this theory, therefore, a Lewis base is a compound that has an unbonded pair of electrons available for bonding. The structure of Lewis acid is such that it can achieve a stable configuration by forming a coordinate bond with a Lewis base. Bases must not contain hydroxide ions or accept protons, and a Lewis acid must not contain hydrogen or donate protons. The definition of Lewis acid includes all Arrhenius and Bronsted-Lowry acids and also many substances that do not meet the Bronsted-Lowry or Arrhenius criteria.
A good example of such a substance is boron trifluoride (BF3). In this compound, the boron, which normally has three electrons in its outer shell, has formed covalent bonds, sharing one pair of electrons with each of the three fluorine atoms. While the compound is stable, it has room for two more electrons in its outer shell. It can then form a coordinated bond with an electron pair donor, in other words a base.
For example, it can combine with ammonia (NH3), which has a nitrogen atom with an unbonded electron pair, because three of the five electrons in the nitrogen’s outer shell are in covalent bonds with the three hydrogen atoms. The combination of boron trifluoride and ammonia is therefore as follows: BF3 + :NH3 → BF3:NH3 — the “:” represents the electron pair of the ammonia nitrogen atom. Boron trifluoride therefore behaves like a Lewis acid and ammonia like a base.
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