A Lewis acid is a chemical species that contains an empty orbital which is capable of accepting an electron pair from a Lewis base to form a Lewis adduct. A Lewis base, then, is any species that has a filled orbital containing an electron pair which is not involved in bonding but may form a dative bond with a Lewis acid to form a Lewis adduct. For example, NH3 is a Lewis base, because it can donate its lone pair of electrons. Trimethylborane [(CH3)3B or Me3B] is a Lewis acid as it is capable of accepting a lone pair. In a Lewis adduct, the Lewis acid and base share an electron pair furnished by the Lewis base, forming a dative bond. In the context of a specific chemical reaction between NH3 and Me3B, a lone pair from NH3 will form a dative bond with the empty orbital of Me3B to form an adduct NH3•BMe3. Lewis acids and bases are named for the American physical chemist Gilbert N. Lewis. The terms nucleophile and electrophile are sometimes interchangeable with Lewis base and Lewis acid, respectively. These terms, especially their abstract noun forms nucleophilicity and electrophilicity, emphasize the kinetic aspect of reactivity, while the Lewis basicity and Lewis acidity emphasize the thermodynamic aspect of Lewis adduct formation.
Depicting adducts In many cases, the interaction between the Lewis base and Lewis acid in a complex is indicated by an arrow indicating the Lewis base donating electrons toward the Lewis acid using the notation of a dative bond—for example, Me3B←NH3. Some sources indicate the Lewis base with a pair of dots (the explicit electrons being donated), which allows consistent representation of the transition from the base itself to the complex with the acid:
Me3B + :NH3 → Me3B:NH3 A center dot may also be used to represent a Lewis adduct, such as Me3B·NH3. Another example is boron trifluoride diethyl etherate, BF3·Et2O. In a slightly different usage, the center dot is also used to represent hydrate coordination in various crystals, as in MgSO4·7H2O for hydrated magnesium sulfate, irrespective of whether the water forms a dative bond with the metal. Although there have been attempts to use computational and experimental energetic criteria to distinguish dative bonding from non-dative covalent bonds, for the most part, the distinction merely makes note of the source of the electron pair, and dative bonds, once formed, behave simply as other covalent bonds do, though they typically have considerable polar character. Moreover, in some cases (e.g., sulfoxides and amine oxides as R2S → O and R3N → O), the use of the dative bond arrow is just a notational convenience for avoiding the drawing of formal charges. In general, however, the donor–acceptor bond is viewed as simply somewhere along a continuum between idealized covalent bonding and ionic bonding.
Lewis acids
Lewis acids are diverse and the term is used loosely. Simplest are those that react directly with the Lewis base, such as boron trihalides and the pentahalides of phosphorus, arsenic, and antimony. In the same vein, CH+3 can be considered to be the Lewis acid in methylation reactions. However, the methyl cation never occurs as a free species in the condensed phase, and methylation reactions by reagents like CH3I take place through the simultaneous formation of a bond from the nucleophile to the carbon and cleavage of the bond between carbon and iodine (SN2 reaction). Textbooks disagree on this point: some asserting that alkyl halides are electrophiles but not Lewis acids, while others describe alkyl halides (e.g. CH3Br) as a type of Lewis acid. The IUPAC states that Lewis acids and Lewis bases react to form Lewis adducts, and defines electrophile as Lewis acids.
Simple Lewis acids Some of the most studied examples of such Lewis acids are the boron trihalides and organoboranes:
BF3 + F− → BF−4 In this adduct, all four fluoride centres (or more accurately, ligands) are equivalent.
BF3 + OMe2 → BF3OMe2 Both BF4− and BF3OMe2 are Lewis base adducts of boron trifluoride. Many adducts violate the octet rule, such as the triiodide anion:
I2 + I− → I−3 The variability of the colors of iodine solutions reflects the variable abilities of the solvent to form adducts with the Lewis acid I2. Some Lewis acids bind with two Lewis bases, a famous example being the formation of hexafluorosilicate:
SiF4 + 2 F− → SiF2−6
Complex Lewis acids Most compounds considered to be Lewis acids require an activation step prior to formation of the adduct with the Lewis base. Complex compounds such as Et3Al2Cl3 and AlCl3 are treated as trigonal planar Lewis acids but exist as aggregates and polymers that must be degraded by the Lewis base. A simpler case is the formation of adducts of borane. Monomeric BH3 does not exist appreciably, so the adducts of borane are generated by degradation of diborane:
B2H6 + 2 H− → 2 BH−4 In this case, an intermediate B2H−7 can be isolated. Many metal complexes serve as Lewis acids, but usually only after dissociating a more weakly bound Lewis base, often water.
[Mg(H2O)6]2+ + 6 NH3 → [Mg(NH3)6]2+ + 6 H2O
H+ as Lewis acid The proton (H+) is one of the strongest but is also one of the most complicated Lewis acids. It is convention to ignore the fact that a proton is heavily solvated (bound to solvent). With this simplification in mind, acid-base reactions can be viewed as the formation of adducts:
H+ + NH3 → NH+4 H+ + OH− → H2O
Applications of Lewis acids A typical example of a Lewis acid in action is in the Friedel–Crafts alkylation reaction. The key step is the acceptance by AlCl3 of a chloride ion lone-pair, forming AlCl−4 and creating the strongly acidic, that is, electrophilic, carbonium ion.
RCl + AlCl3 → R+ + AlCl−4
Lewis bases A Lewis base is an atomic or molecular species where the highest occupied molecular orbital (HOMO) is highly localized. Typical Lewis bases are conventional amines such as ammonia and alkyl amines. Other common Lewis bases include pyridine and its derivatives. They are nucleophilic in nature. Some of the main classes of Lewis bases are:
amines of the formula NH3−xRx where R = alkyl or aryl. Related to these are pyridine and its derivatives. phosphines of the formula PR3−xArx. compounds of O, S, Se and Te in oxidation state −2, including water, ethers, ketones The most common Lewis bases are anions. The strength of Lewis basicity correlates with the pKa of the parent acid: acids with high pKa's give good Lewis bases. As usual, a weaker acid has a stronger conjugate base.
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