Lewis acids and bases

Diagram of some Lewisbasesandacids

ALewis acid(named for the American physical chemistGilbert N. Lewis) is a chemical species that contains an emptyorbitalwhich is capable of accepting anelectron pairfrom a Lewisbaseto form a Lewisadduct.ALewis base,then, is any species that has a filled orbital containing an electron pair which is not involved inbondingbut may form adative bondwith a Lewis acid to form a Lewis adduct. For example,NH₃is a Lewis base, because it can donate itslone pairof electrons.Trimethylborane(${\ce {(CH3)3 B}}$) 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 specificchemical reactionbetween NH₃and Me₃B, a lone pair from NH₃will form a dative bond with the empty orbital of Me₃B to form an adduct NH₃•BMe₃.The terminology refers to the contributions ofGilbert N. Lewis.

The termsnucleophileandelectrophileare sometimes interchangeable with Lewis base and Lewis acid, respectively. These terms, especially their abstract noun formsnucleophilicityandelectrophilicity,emphasize thekineticaspect of reactivity, while the Lewis basicity and Lewis acidity emphasize thethermodynamicaspect 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 adative 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 asMe₃B·NH₃.Another example isboron trifluoride diethyl etherate,BF₃·Et₂O.In a slightly different usage, the center dot is also used to representhydrate coordinationin various crystals, as inMgSO₄·7H₂Ofor hydratedmagnesium 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 asR₂S → OandR₃N → 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 idealizedcovalent bondingandionic bonding.

Lewis acids

Major structural changes accompany binding of the Lewis base to the coordinatively unsaturated, planar Lewis acid BF₃

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+3can 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 (S_N2 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.TheIUPACstates 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 andorganoboranes:

BF₃+ F⁻→BF−4

In this adduct, all four fluoride centres (or more accurately,ligands) are equivalent.

BF₃+ OMe₂→ BF₃OMe₂

Both BF₄⁻and BF₃OMe₂are Lewis base adducts of boron trifluoride.

Many adducts violate theoctet rule,such as thetriiodideanion:

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 ofhexafluorosilicate:

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 asEt₃Al₂Cl₃andAlCl₃are 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 BH₃does not exist appreciably, so the adducts of borane are generated by degradation of diborane:

B₂H₆+ 2 H⁻→ 2BH−4

In this case, an intermediateB₂H−7can 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

Theproton(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 theFriedel–Crafts alkylationreaction.The key step is the acceptance by AlCl₃of a chloride ion lone-pair, formingAlCl−4and 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 thehighest occupied molecular orbital(HOMO) is highly localized. Typical Lewis bases are conventionalaminessuch as ammonia andalkylamines. Other common Lewis bases includepyridineand its derivatives. Some of the main classes of Lewis bases are

• amines of the formula NH_3−xR_xwhere R = alkyl oraryl.Related to these are pyridine and its derivatives.
• phosphinesof 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 thepK_aof the parent acid: acids with highpK_a's give good Lewis bases. As usual, aweaker acidhas a strongerconjugate base.

• Examples of Lewis bases based on the general definition of electron pair donor include:
  • simple anions, such asH⁻andF⁻
  • other lone-pair-containing species, such as H₂O, NH₃,HO⁻,and CH₃⁻
  • complex anions, such assulfate
  • electron-richπ-system Lewis bases, such asethyne,ethene,andbenzene

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

Heats of binding of various bases toBF₃

Lewis base	Donor atom	Enthalpy of complexation (kJ/mol)
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

Main article:Homogeneous catalysis

Nearly all electron pair donors that form compounds by binding transition elements can be viewedligands.Thus, a large application of Lewis bases is to modify the activity and selectivity ofmetal catalysts.Chiral Lewis bases, generallymultidentate,conferchiralityon a catalyst, enablingasymmetric catalysis,which is useful for the production ofpharmaceuticals.The industrial synthesis of the anti-hypertension drugmibefradiluses a chiral Lewis base (R-MeOBIPHEP), for example.

Hard and soft classification

Main article:HSAB theory

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(0), Ni(0), Pt²⁺
• typical hard bases: ammonia and amines, water, carboxylates, fluoride and chloride
• typical soft bases: organophosphines, thioethers, carbon monoxide, iodide

For example, anaminewill displacephosphinefrom 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 areenthalpyfavored, whereas soft—soft areentropyfavored.

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. theGutmann-Beckett methodand the Childsmethod.

TheECW modelis 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_Aand a C_A.Each base is likewise characterized by its own E_Band 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

−ΔH = E_AE_B+ C_AC_B+ W

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.

History

MO diagramdepicting the formation of a dative covalent bond between two atoms

The concept originated withGilbert N. Lewiswho studiedchemical bonding.In 1923, Lewis wroteAn acid substance is one which can employ an electron lone pair from another molecule in completing the stable group of one of its own atoms.TheBrønsted–Lowry acid–base theorywas published in the same year. The two theories are distinct but complementary. A Lewis base is also a Brønsted–Lowry base, but a Lewis acid does not need to be a Brønsted–Lowry acid. The classification into hard and soft acids and bases (HSAB theory) followed in 1963. The strength of Lewis acid-base interactions, as measured by the standardenthalpyof formation of an adduct can be predicted by the Drago–Wayland two-parameter equation.

Reformulation of Lewis theory

Lewis had suggested in 1916 that twoatomsare held together in a chemical bond by sharing a pair of electrons.When each atom contributed one electron to the bond, it was called acovalent bond.When both electrons come from one of the atoms, it was called a dative covalent bond orcoordinate bond.The distinction is not very clear-cut. For example, in the formation of an ammonium ion from ammonia and hydrogen theammoniamolecule donates a pair of electrons to theproton;the identity of the electrons is lost in theammoniumion that is formed. Nevertheless, Lewis suggested that an electron-pair donor be classified as a base and an electron-pair acceptor be classified as acid.

A more modern definition of a Lewis acid is an atomic or molecular species with a localized emptyatomicormolecularorbital of low energy. This lowest-energy molecular orbital (LUMO) can accommodate a pair of electrons.

Comparison with Brønsted–Lowry theory

A Lewis base is often a Brønsted–Lowry base as it can donate a pair of electrons to H⁺;the proton is a Lewis acid as it can accept a pair of electrons. The conjugate base of a Brønsted–Lowry acid is also a Lewis base asloss of H⁺from the acid leaves those electrons which were used for the A—H bond as a lone pair on the conjugate base. However, a Lewis base can be very difficult toprotonate,yet still react with a Lewis acid. For example,carbon monoxideis a very weak Brønsted–Lowry base but it forms a strong adduct with BF₃.

In another comparison of Lewis and Brønsted–Lowry acidity by Brown and Kanner,2,6-di-t-butylpyridine reacts to form the hydrochloride salt with HCl but does not react with BF₃.This example demonstrates that steric factors, in addition to electron configuration factors, play a role in determining the strength of the interaction between the bulky di-t-butylpyridine and tiny proton.

See also

• Acid
• Base (chemistry)
• Acid–base reaction
• Brønsted–Lowry acid–base theory
• Chiral Lewis acid
• Frustrated Lewis pair
• Gutmann–Beckett method
• ECW model
• Philosophy of chemistry

Further reading

• Jensen, W.B.(1980).The Lewis acid-base concepts: an overview.New York: Wiley.ISBN0-471-03902-0.
• Yamamoto, Hisashi (1999).Lewis acid reagents: a practical approach.New York: Oxford University Press.ISBN0-19-850099-8.