Lewis acid

A Lewis acid (named for American physical chemist Gilbert N. Lewis ) is a chemical species that contains an empty orbital which is able to accept 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 that is not involved in the bond but can form a dative bond with a Lewis acid to form a Lewis adduct . For example, NH 3 is a Lewis base because it stores its own electrons. Can donate to a single couple . Trimethylborane (Me3b ) is a Lewis acid because it is able to accept a lone pair. In a Lewis adduct, a Lewis acid and a base share an electron pair furnished by a Lewis base, forming a dative bond. [1] In the context of a typical chemical reaction between NH3 and Me3B, the lone pair from NH 3 will form a basic bond with the vacant orbital of Me3B to form an adduct NH3 •BMe3 . The terminology refers to the contribution of Gilbert N. Lewis .

The terms nucleophile and electrophile are more or less interchangeable with Lewis base and Lewis acid respectively. However, these terms, especially their abstract nouns, nucleophilicity and electrophilicity , emphasize the kinetic aspect of reactivity, while Lewis basicity and Lewis acidity emphasize the thermodynamic aspect of Lewis addition formation.

Illustration of addictions

In many cases, the interaction between a Lewis base and a Lewis acid in a compound is indicated by an arrow indicating the Lewis acid donating electrons using the notation of a dative bond to the Lewis acid—for example, 3 b←nh 3 . Some sources indicate a Lewis base with a pair of dots (the apparent electrons being donated), which allows a consistent representation of the transition from base to complex with acids:

Me 3b + :NH 3 → Me 3 B: NH 3

A focal point can also be used to denote a Lewis adduct, such as Me 3 B•NH 3 . Another example is boron trifluoride diethyl ether , BF 3 • Et 2 O. (In a slightly different usage, the center point is also used to represent the hydrate coordination in various crystals, such as MgSO 4 • 7H 2 O for hydrated magnesium sulfate , regardless of whether the water forms a basic bond with the metal. is.)

Although attempts have been made to use computational and experimental energetic criteria to distinguish basic bonds from non-dative covalent bonds, for the most part, the difference is only between the source of the electron pair and the origin once formed. Pays attention to ties. , behave in the same way as other covalent bonds, although they usually have a fairly polar character. Furthermore, in some cases (for example, R 2 s → O and R 3 n → O as in sulfoxides and amine oxides), the use of the basic bond arrow is just an indicative feature to avoid the depiction of formal charges. In general, however, the donor–acceptor bond is seen to be somewhere along a continuum between the ideal covalent bond and the ionic bond.

Lewis acid

Classically, the term “Lewis acid” is limited to tri-planar species with an empty p orbital , such as Br where R can be an organic material or a halide. [ citation needed ] For the purposes of discussion, even complex compounds such as Et 3 Al 2 Cl 3 and AlCl 3 are treated as trigonal planar Lewis acids. Metal ions such as Na+ , Mg2+ , and Ce3+ , which are always complexed with additional ligands , often synergisticallyThere are sources of unsaturated derivatives that form a Lewis addition upon reaction with a Lewis base . [ citation needed ] Other reactions may simply be referred to as “acid-catalyzed” reactions. Some compounds, such as H2O , are both Lewis acids and Lewis bases, as they can either accept a pair of electrons or donate a pair of electrons, depending on the reaction.

Lewis acids are diverse. The simplest are those that react directly with a Lewis base. But more common are those that undergo a reaction before forming an addiction. Examples of Lewis acids based on the general definition of electron pair acceptor include:

  • protons (H + ) and acidic compounds such as onium ions , such as NH 4 + and H 3 O +
  • Higher oxidation state transition metal cations, for example, Fe 3+ ;
  • Other metal cations, such as Li + and Mg2 + , often form their aqua or ether complexes,
  • Tri-planar species, such as BF3 and the carbocation H3C +
  • pentahalides of phosphorus, arsenic and antimony
  • Poor electron -systems , such as enones and tetracyanoethylenes .

Again, the description of Lewis acids is often used loosely. For example, bare protons do not exist in solution.

Simple lewis acid

Some of the most studied examples of such Lewis acids are boron trihalides and organoboranes , but other compounds exhibit this behavior:bf 3 + f  → bf 

In this addition, all four fluoride centers (or more precisely, ligands ) are equal.bf 3 + ome 2 → bf 3 ome 2

Both BF  and BF 3 OMe 2 are Lewis base adducts of boron trifluoride.

In many cases, adducts violate the octet rule , such as the triiodide anion: I2 + I → I

The variability of the colors of iodine solutions reflects the changing abilities of the solvent to form additives with Lewis acid I 2 .

In some cases, a Lewis acid is able to bind to two Lewis bases, a well-known example being the formation of hexafluorosilicate:SiF 4 + 2 F  → SiF 2−

Complex Lewis Acid

Most compounds considered to be Lewis acids require an activation step before the addition can form with a Lewis base. Well-known cases are aluminum trihalides, widely viewed as Lewis acids. Aluminum trihalides, unlike boron trihalides, do not exist as AlX3, but as aggregates and polymers that must be degraded by Lewis bases [6] A simpler case is the creation of borane adducts. Monomeric BH is not sufficiently present, so borane adducts are produced by degradation of diborane:b 2h 6 + 2h  → 2bh 

In this case, an intermediate B 2 H  can be isolated.

Many metal complexes act as Lewis acids, but usually only the more weakly bonded Lewis bases, often followed by dissociation with water.[Mg (H 2 O) 6 ] 2+ + 6 NH 3 → [Mg (NH 3 ) 6 ] 2+ + 6 H 2 O

H+ as Lewis acid

The proton (H + )  [7] is one of the strongest, but it is also one of the most complex Lewis acids. It is customary to ignore the fact that a proton is heavily soluble (bound to a solvent). With this simplification in mind, acid–base reactions can be viewed as the formation of adducts:

  • + + NH 3 → NH +
  • + + OH  → H 2 O

Applications of Lewis Acids

A typical example of a Lewis acid in action is in the Friedel–Crafts alkylation reaction. [5] Acceptance of a chloride ion lone-pair by AlCl 3 , forming AlCl  and forming the strongly acidic, i.e. electrophilic, carbonium ion.RCl +AlCl 3 → R + + AlCl 

Lewis Bess

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. Some of the main classes of Lewis bases are

  • Amines of the formula NH 3− x R x where R = alkyl or aryl. Related to these are pyridine and its derivatives.
  • The formula for phosphines is Pr x a x , where r = alkyl, a = aryl.
  • Compounds of O, S, Se and Te in oxidation state -2 including water, ether, ketone

The most common are Lewis base ions. strength of Lewis basicityThe pK of the basic acid is correlated with a: Acids with high pK a give good Lewis bases. As always, a weak acid has a strong 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 HO, NH3 , HO- and CH3-
    • complex ions such as sulfate
    • The electron-rich -system of Lewis bases, such as ethyne , ethene, and benzene

The strength of Lewis bases has been evaluated for various Lewis acids, such as I 2 , SbCl 5 , and BF 3 . [8]

Lewis Bessdonor atomEnthalpy of Complexity (kJ/mol)
at 3 nNo135
at 2 oO78.8

Applications of Lewis bases

Almost all electron pair donors that form compounds by binding to transition elements can be viewed as Lewis bases—or collections of ligands. Thus a major application of Lewis bases is to modify the activity and selectivity of metal catalysts. Chiral Lewis bases thus confer chirality over catalysts, enabling asymmetric catalysis, which is useful for the production of pharmaceuticals.

Many Lewis bases are “multidentate”, i.e. they can form multiple bonds to Lewis acids. These multifunctional Lewis bases are called chelating agents.

Hard and soft classification

Lewis acids and bases are usually classified according to their hardness or softness. Hard in this context means smaller and non-polar and soft indicates larger atoms that are more polarizable.

  • Typical hard acids: H + , alkali/alkaline earth metal cations, borane, Zn 2+
  • Typical soft acids: Ag + , Mo(0), Ni(0), Pt 2+
  • Typical hard bases: ammonia and amines, water, carboxylates, fluorides and chlorides
  • Typical soft bases: organophosphine, thioether, carbon monoxide, iodide

For example, an amine acid will displace phosphine by addition with BF3 In the same way, bases can 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 additive formation using the key concepts that hard acid–hard bases and soft acid–soft base interactions are hard acid–soft bases or hard acid–soft bases. Soft acids are stronger than hard. base interaction. Subsequent investigations into the thermodynamics of the interaction suggested that the harder–harder interactions favor enthalpy, while the softer one favors the softer entropy.

Quantification of Lewis Acidity

Several methods have been devised to evaluate and predict Lewis acidity. Many are based on spectroscopic signatures such as shift NMR signals or IR bands such as the Gutmann–Beckett method and the Childs [9] method.

The ECW model -ΔH is a quantitative model that describes and predicts the strength of Lewis acid-base interactions. The model gave the E and C parameters to several Lewis acids and bases. Each acid is characterized by E A and C A. Similarly each base is characterized by its own E B and C B. The E and C parameters refer to the electrostatic and covalent contributions, respectively, to the strength of the bonds that form acids and bases. the equation is

−ΔH = E A E B + C A C B + W

The W term represents a constant energy contribution to an acid–base reaction such as the cleavage of a dimeric acid or base. The equation predicts the inversion of acid and base strengths. The graphical representation of the equation shows that there is no single sequence of Lewis base strengths or Lewis acid strengths. [10] [11] And that single property scale is limited to a small range of acids or bases.


The concept originated from Gilbert N. Lewis who studied chemical bonding. In 1923, Lewis wrote that an acid substance is one that can employ an electron lone pair from another molecule to complete the stable group of one of its own atoms. [2] [12] The Bronsted–Lowry acid–base theory was published in the same year. The two principles are different but complementary. A Lewis base is also a Brnsted–Lowry base, but a Lewis acid is not required to be a Brnsted–Lowry acid. The classification followed in 1963 into hard and soft acids and bases (HSAB principle). The strength of the Lewis acid–base interaction, as measured by the standard enthalpy of formation of an adduct, can be estimated by the Drago–Weland two-parameter equation.

Improvement of lewis theory

Lewis suggested in 1916 that two atoms are held together by sharing a pair of electrons in a chemical bond. [13] When each atom contributed an electron to the bond, it was called a covalent bond. When both electrons come from one atom, it was called a basic covalent bond or coordination bond. The distinction is not very clear. For example, in the formation of the ammonium ion from ammonia and hydrogen, the ammonia molecule donates a pair of electrons to a proton; [7] The identity of the electrons in the formed ammonium ion is lost. Nevertheless, Lewis suggested that an electron-pair donor be classified as a base and an electron-pair acceptor be classified as an acid.

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

Comparison with Bronsted-Lowry theory

A Lewis base is often a Brnsted–Lowry base because it can donate a pair of electrons to H + ; [7] The proton is a Lewis acid because it can accept a pair of electrons. The conjugate base of a Brnsted–Lowry acid is also a Lewis base because the loss of H+ from the acid leaves electrons that were used for the A–H bond to form a lone pair on the conjugate base. Although it can be very difficult to protonate a Lewis base, it still reacts with Lewis acids. For example, carbon monoxide is a very weak Brnsted–Lowry base but it forms a strong adduct with BF3.

In a further comparison of Lewis and Bronsted-Lowry acidity by Brown and Kanner, [14] 2,6-di- t – butylpyridine reacts with HCl to form the hydrochloride salt but does not react with BF 3 . This example demonstrates that steric factors, in addition to electron configuration factors, play a role in determining the strength of the interaction between the heavy di- t- butylpyridine and the small proton.

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