Lewis acid-base theory

Diagram of some bases and acids

A Lewis acid is a chemical species containing an empty orbital that 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 that is not involved in a bond, but can 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 only lone pair of electrons. Trimethylborane (Me₃B) is a Lewis acid, since it is capable of accepting a lone pair. In a Lewis adduct, the Lewis acid and base share an electron pair provided by the Lewis base, forming a dative bond. In the context of a specific chemical reaction between NH₃ and Me₃B, the lone pair of NH₃ will form a dative bond with the empty orbital of Me ₃ B to form an NH₃ • BMe₃. The terminology refers to the contributions 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 noun forms, nucleophilicity and electrophilicity, emphasize the kinetic aspect of reactivity, while Lewis basicity and Lewis acidity emphasize the thermodynamic aspect of the formation of Lewis adducts.

Representing 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 to the Lewis acid using dative bond notation, for Example, Me₃B ← NH₃. Some sources indicate the Lewis base with a pair of dots (the explicit electrons being donated), allowing a consistent representation of the transition from the base itself to the complex with the acid:

Me.₃B +:NH₃ → Me.₃B:NH₃

A center point can 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 point is also used to represent coordination of hydrates in various crystals, as in MgSO₄ • 7H₂O for hydrated magnesium sulfate, regardless 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 nondative covalent bonds, for the most part, the distinction simply takes note of the source of the electron pair and the dative bonds, a Once formed, they behave simply as other covalent bonds do, although they usually have considerable polar character. Also, in some cases (for example, sulfoxides and amine oxides like R₂S → O and R₃N → O), the use of the dative bond arrow it is just a notational convenience to avoid drawing formal charges. In general, however, the donor-acceptor bond is simply considered to be somewhere on a continuum between the idealized covalent bond and the ionic bond.

Lewis acids

The main structural changes accompany the bonding of the Lewis base to the co-ordinatively unsaturated Lewis acid BF₃

Classically, the term "Lewis acid" is restricted to trigonal planar species with an empty p orbital, such as BR_3, where R can be an organic substituent or a halide. For the purposes of the discussion, even complex compounds such as Et3Al2Cl3 and AlCl3 are treated as trigonal planar Lewis acids. Metal ions such as Na⁺, Mg²⁺, and Ce³⁺, which invariably form complexes with additional ligands, are often sources of derivatives. coordinately unsaturated that form Lewis adducts upon reaction with a Lewis base. Other reactions could simply be called "acid-catalyzed" reactions. Some compounds, such as H₂O, are both Lewis acids and Lewis bases, because they can accept an electron pair or donate an electron pair, depending on the reaction.

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

• proton (H⁺) and onio ions composed of acids, such as NH4+ and H3O+
• cationes of transition metals of high oxidation state, for example, Fe³⁺;
• other metal cations, such as Li+ and Mg2+, often as their complexes here or ether,
• wheat flat species, such as BF3 and carbocationes H₃C⁺
• phosphorus, arsenic and antimony pentahalides
• π poor in electrons, such as enones and tetracianoethylenes.

Once again, the description of a Lewis acid is often used loosely. For example, in solution, naked protons do not exist.

Simple Lewis Acids

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

BF₃ + F^- → BF₄^-

In this adduct, the four fluoride centers (or more accurately, ligands) are equivalent.

BF₃ + OMe₂ → BF₃OMe₂

Both BF₄^- and BF₃OMe₂ are Lewis base adducts of boron trifluoride.

In many cases, adducts violate the octet rule, like the triiodide anion:

I₂ + I⁻ → I₃⁻

The variability of the colors of iodine solutions reflects the varying abilities of the solvent to form adducts with the Lewis acid I₂.

In some cases, Lewis acid is capable of joining two Lewis bases, a famous example being the formation of hexafluorosilicate:

SiF₄ + 2 F⁻ → SiF₆²⁻

Complex Lewis Acids

Most of the compounds considered Lewis acids require an activation step before the formation of the adduct with the Lewis base. Well-known cases are aluminum trihalides, which are generally considered Lewis acids. Aluminum trihalides, unlike boron trihalides, do not exist in the AlX₃ form, but rather as aggregates and polymers that must be broken down by the Lewis base. A simpler case is the formation of borane adducts. The monomeric BH₃ does not exist in any appreciable form, so the borane adducts are generated by diborane degradation:

B₂H₆ + 2 H⁻ → 2 BH₄⁻

In this case, an intermediate B₂H₇^- can be isolated.

Many metal complexes serve as Lewis acids, but usually only after dissociating a more loosely bound Lewis base, often water.

[Mg(H)₂O)₆]²⁺ + 6 NH₃ → [Mg(NH₃)₆]²⁺ + 6 H₂O

H+ as Lewis acid

The proton (H⁺) is one of the strongest, but also one of the most complicated Lewis acids. It is a convention to ignore the fact that a proton is strongly solvated (bound to the solvent). With this simplification in mind, acid-base reactions can be viewed as the formation of adducts:

• H⁺ + NH₃ → NH₄⁺
• H⁺ + OH⁻ → H₂O

Applications of Lewis acids

A typical example of a Lewis acid in action is the Friedel-Crafts alkylation reaction. The key step is the acceptance by AlCl₃ of a lone pair of chloride ions, forming AlCl ₄^- and the creation of the strongly acidic, electrophilic, carbonium ion.

RCl + AlCl₃ → R⁺ + AlCl₄⁻

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 alkylamines. Other common Lewis bases include pyridine and its derivatives. Some of the major classes of Lewis bases are

• NH formula amines_3−xR_x where R = alkylo or arilo. Related to these are the piridine and its derivatives.
• PR formula phosphine_3−xA_xWhere R = alkylo, A = arilo.
• 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 pK_a of the source acid: acids with high pK_a 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, like H- and F-
  • other species containing lone pairs, such as H₂O, NH₃, HO- and CH₃^-
  • complex anions, such as sulfate
  • Lewis bases of the π system rich in electrons, such as etino, eteno and benceno

The strength of Lewis bases has been evaluated for various Lewis acids, such as I₂, SbCl ₅, and BF₃.

Applications of Lewis Bases

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

Many Lewis bases are "multidentate," that is, they can form multiple bonds with the Lewis acid. These multidentate Lewis bases are called chelating agents.

Hard and soft classification

Lewis acids and bases are commonly classified according to their hardness or softness. In this context, strong implies small and non-polarizable and weak indicates larger atoms that are more polarizable.

• Typical hard acids: H⁺, cationes de metales alcalinos/alcalinotérreos, boranos, Zn²⁺
• Typical soft acids: Ag⁺, Mo(0), Ni(0), Pt²⁺
• Typical hard bases: ammonia and amines, water, carboxylates, fluoride and chloride
• Typical soft bases: organophosphine, thyoeter, carbon monoxide, iodide

For example, an amine will displace the phosphine from the adduct with BF₃ acid. Bases could be classified in the same way. For example, bases that donate a lone pair from an oxygen atom are harder than those that donate through a nitrogen atom. Although the classification was never quantified, it proved 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—hard acid-base interactions. soft base or soft acid—hard base. Later investigation of the thermodynamics of the interaction suggested that hard–hard interactions are favored by enthalpy, while soft–soft interactions are favored by entropy.

Quantifying Lewis Acidity

Many methods have been devised to assess and predict Lewis acidity. Many are based on spectroscopic signatures, such as signal changes in NMR or infrared bands, for example 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 parameters E and C to many Lewis acids and bases. Each acid is characterized by an E_A and a C_A. Each base is equally characterized by its own E_B and C_B. The parameters E and C 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 _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 dimer acid or base. The equation predicts the reversal of acid and base concentrations. Graphical presentations of the equation show that there is no single order of Lewis base concentrations or Lewis acid concentrations, and that individual property scales are limited to a smaller range of acids or bases.

History

MO diagram depicting the formation of a covalent link between two atoms

The concept originated with Gilbert N. Lewis, who studied chemical bonding. In 1923, Lewis wrote An acidic substance is one that can use a lone pair of electrons from another molecule to complete the stable group of one of its own atoms. The Brønsted-acid-base theory Lowry was 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 need not 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 standard enthalpy of formation of an adduct, can be predicted by the two-parameter Drago- Wayland.

Restatement of Lewis's theory

Lewis had suggested in 1916 that two atoms are held together in a chemical bond by sharing a pair of electrons. When each atom contributed one electron to the bond, it was called a covalent bond. When both electrons come from one of the atoms, it is called a dative covalent bond or a coordinate bond. The distinction is not very clear. For example, in the formation of an ammonium ion from ammonia and hydrogen, the ammonia molecule donates an electron pair to the proton; the identity of the electrons is lost in the ammonium ion that is formed. However, 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 a Lewis acid is an atomic or molecular species with a low-energy localized empty atomic or molecular orbital. This lowest energy molecular orbital (LUMO) can accommodate one pair of electrons.

Comparison with Brønsted-Lowry theory

A Lewis base is often a Brønsted-Lowry base, since it can donate an electron pair to H⁺; the proton is a Lewis acid, since it can accept a pair of electrons. The conjugate base of a Brønsted-Lowry acid is also a Lewis base, since the loss of H+ from the acid leaves the electrons that were used for the A-H bond as a lone pair in the conjugate base. However, a Lewis base can be very difficult to protonate and still react with a Lewis acid. For example, carbon monoxide is 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 not with BF ₃. This example demonstrates that steric factors, in addition to electronic configuration factors, play a role in determining the strength of the interaction between the bulky di-t-butylpyridine and the minute proton.

Further reading

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