A
Lewis acid is defined by
IUPAC as a "molecular entity (and the corresponding chemical species) that is an electron-pair acceptor and therefore able to react with a Lewis base to form a Lewis
adduct, by sharing the electron pair furnished by the Lewis base." An illustrative example is given by the reaction of
trimethylboron and
ammonia to give the adduct Me
3BNH
3.
A Lewis acid is defined to be any species that accepts lone pair electrons. A Lewis base is any species that donates lone pair electrons. Thus, H+ is a Lewis acid, since it can accept a lone pair, while OH− and NH3 are Lewis bases, both of which donate a lone pair.
Depicting adducts
In many cases, the interaction between the boron and nitrogen is indicated by an arrow, Me
3B←-NH
3, the direction of which points from the Lewis base toward the Lewis acid. Some sources indicate the Lewis base with a pair of dots both in the precursor Lewis base as well as the adduct, as shown in this equation:
Me3B + :NH3 → Me3B:NH3
In general, however, the donor acceptor bond is viewed as a step in a continuum between idealized covalent bonding and ionic bonding. The
Brønsted–Lowry acid–base theory 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 doesn't 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 standard enthalpy of formation of an adduct can be predicted by the Drago–Wayland two-parameter equation.
Reformulation of Lewis 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 was called a dative covalent bond or
coordinate bond. The distinction is not very clear-cut. For example, in the formation of an ammonium ion from ammonia and hydrogen the
ammonia molecule donates a pair of electrons to the
proton; the identity of the electrons is lost in the
ammonium ion 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 that with a localized empty atomic or molecular orbital 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
+; 2,6-di-
t-butylpyridine reacts to form the hydrochloride salt with HCl but does not react with BF
3. This example demonstrates that steric factors, in addition to electronic factors, play a role in determining the strength of the interaction between the bulky di-
t-butylpyridine and tiny proton.
A Brønsted–Lowry acid is a proton donor, not an electron-pair acceptor.
Lewis acids
ion (diagram should show
adduct having a positive charge)]]
Lewis acids are diverse. Simplest are those that react directly with the Lewis base. But more common are those that undergo a reaction prior to forming the adduct.
Simple Lewis acids
The most studied examples of such Lewis acids are the boron trihalides and organoboranes, but other compounds exhibit this behavior:
:BF
3 + F
− → BF
4−
In this adduct, all four fluoride groups (or more accurately,
ligands) are equivalent.
:BF
3 + OMe
2 → BF
3OMe
2
Both BF
4− and BF
3OMe
2 are Lewis base adducts of boron trifluoride.
In many cases, the adducts violate the octet rule, such as the triiodide anion:
:I2 + I− → I3−
The variability of the colors of iodine solutions reflects the variable abilities of solvent to form adducts with the Lewis acid I2.
In some cases, the Lewis acids are capable of binding two Lewis bases, a famous example being the formation of hexafluorosilicate:
:SiF4 + 2 F− → SiF62−
Complex Lewis acids
Most compounds considered to be Lewis acids require an activation step prior to formation of the adduct with the Lewis base. Well known cases are the aluminium trihalides, which are widely viewed as Lewis acids. Aluminium trihalides, unlike the boron trihalides, do not exist in the form AlX
3, but as aggregates and polymers that must be degraded by the Lewis base. A simpler case is the formation of adducts of borane. Monomeric BH
3 does not exist appreciably, so the adducts of borane are generated by degradation of diborane:
:B
2H
6 + 2 H
− → 2 BH
4−
In this case, an intermediate B
2H
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(H
2O)
6]
2+ + 6 NH
3 → [Mg(NH
3)
6]
2+ + 6 H
2O
H+ as Lewis acid
The proton (H
+)]] The key step is the acceptance by AlCl
3 of a chloride ion lone-pair, forming AlCl
4− and creating the strongly acidic, that is,
electrophilic, carbonium ion.
:RCl +AlCl
3 → R
+ + AlCl
4−
Lewis bases
A Lewis base is an atomic or molecular species where the
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 NH3−xRx where R = alkyl or aryl. Related to these are pyridine and its derivatives.
phosphines of the formula PR
3−xA
x, where R = alkyl, A = aryl.
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 parent acid: acids with high pK
a's give good Lewis bases.
The strength of Lewis bases have been evaluated for various Lewis acids, such as I2, SbCl5, and BF3.
{| class="wikitable collapsible" align="center" border="1" cellspacing="0" cellpadding="0" style="margin: 0 0 0 0em; background: #FFFFFF; border-collapse: collapse; border-color: #C0C030;"
!+ colspan="3" align="center" style="background:#ffdead;"| Heats of binding of various bases to BF3
|-
! colspan=1 align="center" style="background:#ffdead;"| Lewis base
! colspan=1 align="center" style="background:#ffdead;"| donor atom
! colspan=1 align="center" style="background:#ffdead;"| Enthalphy of Complexation (kJ/mol)
|-
|Et3N
|N
|135
|-
|quinuclidine
|N
|150
|-
|pyridine
|N
|128
|-
|Acetonitrile
|N
|60
|-
|Et2O
|O
|78.8
|-
|THF
|O
|90.4
|-
|acetone
|O
|76.0
|-
|EtOAc
|O
|75.5
|-
|DMA
|O
|112
|-
|DMSO
|O
|105
|-
|Tetrahydrothiophene
|S
|51.6
|-
|PMe3
|P
|97.3
|-
|}
Applications of Lewis bases
Nearly all of the compounds formed by the transition elements can be viewed as collections of the Lewis bases – or
ligands – bound to a metal. Thus a large application of Lewis bases is to modify the activity and selectivity of metal catalysts. Chiral Lewis bases thus confer
chirality on a catalyst, enabling
asymmetric catalysis, which is useful for the production of
pharmaceuticals.
Many Lewis bases are "multidentate," that is they can form several bonds to the Lewis acid. These multidentate Lewis acids are called chelating agents.
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(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, 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
hard acid — hard base interactions are stronger than hard acid — soft base or soft acid — hard base interactions.
soft acid — soft base interactions are stronger than soft acid — hard base or hard acid — soft base interactions.
Later investigation of the thermodynamics of the interaction suggested that hard—hard interactions are
enthalpy favored, whereas soft—soft are
entropy favored.
References
Further reading
See also
2-Methyltetrahydrofuran
acid
base
acid–base reaction
Brønsted–Lowry acid–base theory
chiral Lewis acid
Category:Acid-base chemistry
Category:Acids
Category:Bases
![Jenny Lewis - Acid Tongue [QMU]](http://web.archive.org./web/20110830012427im_/http://i.ytimg.com/vi/2Z-e7XYW1Qk/0.jpg "Jenny Lewis - Acid Tongue [QMU]")
