Lewis acids and bases

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 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, ←. Some sources indicate the Lewis base with a pair of dots, which allows consistent representation of the transition from the base itself to the complex with the acid:
A center dot may also be used to represent a Lewis adduct, such as. Another example is boron trifluoride diethyl etherate,. In a slightly different usage, the center dot is also used to represent hydrate coordination in various crystals, as in 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, 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, 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. Textbooks disagree on this point: some asserting that alkyl halides are electrophiles but not Lewis acids, while others describe alkyl halides 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:
In this adduct, all four fluoride centres are equivalent.
Both BF₄⁻ and BF₃OMe₂ are Lewis base adducts of boron trifluoride.
Many adducts violate the octet rule, such as the triiodide anion:
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:

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:
In this case, an intermediate can be isolated.
Many metal complexes serve as Lewis acids, but usually only after dissociating a more weakly bound Lewis base, often water.

H⁺ as Lewis acid

The proton 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. With this simplification in mind, acid-base reactions can be viewed as the formation of adducts:

H⁺ + NH₃ →
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 and creating the strongly acidic, that is, electrophilic, carbonium ion.

Lewis bases

A Lewis base is an atomic or molecular species where the highest occupied molecular orbital 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 NH_3−xR_x where R = alkyl or aryl. Related to these are pyridine and its derivatives.
phosphines of the formula PR_3−xAr_x.
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 of the parent acid: acids with high 's give good Lewis bases. As usual, a weaker acid has a stronger conjugate base.

Examples of Lewis bases based on the general definition of electron pair donor include:
*simple anions, such as H⁻ and F⁻
*other lone-pair-containing species, such as H₂O, NH₃, HO⁻, and CH₃⁻
*complex anions, such as sulfate
*electron-rich -system Lewis bases, such as ethyne, ethene, and benzene

The strength of Lewis bases have been evaluated for various Lewis acids, such as I₂, SbCl₅, and BF₃.

Lewis base	Donor atom	Enthalpy of complexation
Quinuclidine	N	150
Et₃N	N	135
Pyridine	N	128
Acetonitrile	N	60
DMA	O	112
DMSO	O	105
THF	O	90.4
Et₂O	O	78.8
Acetone	O	76.0
EtOAc	O	75.5
Trimethylphosphine	P	97.3
Tetrahydrothiophene	S	51.6

Applications of Lewis bases

Nearly all electron pair donors that form compounds by binding transition elements can be viewed ligands. Thus, a large application of Lewis bases is to modify the activity and selectivity of metal catalysts. Chiral Lewis bases, generally multidentate, confer chirality on a catalyst, enabling asymmetric catalysis, which is useful for the production of pharmaceuticals. The industrial synthesis of the anti-hypertension drug mibefradil uses a chiral Lewis base, for example.

Hard and soft classification

Lewis acids and bases are commonly classified according to their hardness or softness. In this context hard implies small and nonpolarizable and soft indicates larger atoms that are more polarizable.

typical hard acids: H⁺, alkali/alkaline earth metal cations, boranes, Zn²⁺
typical soft acids: Ag⁺, Mo, Ni, Pt²⁺
typical hard bases: ammonia and amines, water, carboxylates, fluoride and chloride
typical soft bases: organophosphines, thioethers, carbon monoxide, iodide

For example, an amine will displace phosphine from the adduct with the acid BF₃. In the same way, bases could be classified. For example, bases donating a lone pair from an oxygen atom are harder than bases donating through a nitrogen atom. Although the classification was never quantified it proved to be very useful in predicting the strength of adduct formation, using the key concepts that hard acid—hard base and soft acid—soft base interactions are stronger than hard acid—soft base or soft acid—hard base interactions. Later investigation of the thermodynamics of the interaction suggested that hard—hard interactions are enthalpy favored, whereas soft—soft are entropy favored.

Quantifying Lewis acidity

Many methods have been devised to evaluate and predict Lewis acidity. Many are based on spectroscopic signatures such as shifts NMR signals or IR bands e.g. the Gutmann-Beckett method and the Childs method.
The ECW model is a quantitative model that describes and predicts the strength of Lewis acid base interactions, −ΔH. The model assigned E and C parameters to many Lewis acids and bases. Each acid is characterized by an E_A and a C_A. Each base is likewise characterized by its own E_B and C_B. The E and C parameters refer, respectively, to the electrostatic and covalent contributions to the strength of the bonds that the acid and base will form. The equation is
The W term represents a constant energy contribution for acid–base reaction such as the cleavage of a dimeric acid or base. The equation predicts reversal of acids and base strengths. The graphical presentations of the equation show that there is no single order of Lewis base strengths or Lewis acid strengths. and that single property scales are limited to a smaller range of acids or bases.