Lewis acids and bases
A Lewis acid (named for the American physical chemist
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.[3]
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
- Me3B + :NH3 → Me3B:NH3
A center dot may also be used to represent a Lewis adduct, such as Me3B·NH3. Another example is
Although there have been attempts to use computational and experimental energetic criteria to distinguish dative bonding from non-dative covalent bonds,
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, CH+3 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 (SN2 reaction). Textbooks disagree on this point: some asserting that alkyl halides are electrophiles but not Lewis acids,
Simple Lewis acids
Some of the most studied examples of such Lewis acids are the boron trihalides and
- BF3 + F− → BF−4
In this adduct, all four fluoride centres (or more accurately, ligands) are equivalent.
- BF3 + OMe2 → BF3OMe2
Both BF4− and BF3OMe2 are Lewis base adducts of boron trifluoride.
Many adducts violate the octet rule, such as the triiodide anion:
- 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 of
- 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 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.[10] 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:
- B2H6 + 2 H− → 2 BH−4
In this case, an intermediate B2H−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(H2O)6]2+ + 6 NH3 → [Mg(NH3)6]2+ + 6 H2O
H+ as Lewis acid
The proton (H+) [11] 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 the
- RCl +AlCl3 → R+ + AlCl−4
Lewis bases
A Lewis base is an atomic or molecular species where the
- amines of the formula NH3−xRx where R = alkyl or aryl. Related to these are pyridine and its derivatives.
- phosphines of the formula PR3−xArx.
- 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 pKa of the parent acid: acids with high pKa's give good Lewis bases. As usual, a
- Examples of Lewis bases based on the general definition of electron pair donor include:
The strength of Lewis bases have been evaluated for various Lewis acids, such as I2, SbCl5, and BF3.[12]
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
Nearly all electron pair donors that form compounds by binding transition elements can be viewed
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 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.[citation needed]
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
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 EA and a CA. Each base is likewise characterized by its own EB and CB. 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 = EAEB + CACB + 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.[15][16] and that single property scales are limited to a smaller range of acids or bases.
History
The concept originated with
Reformulation of Lewis theory
Lewis had suggested in 1916 that two
A more modern definition of a Lewis acid is an atomic or molecular species with a localized empty
Comparison with Brønsted–Lowry theory
A Lewis base is often a Brønsted–Lowry base as it can donate a pair of electrons to H+;[11] 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 as loss 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 to protonate, yet 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,[19] 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
References
- ^
- ^ ISBN 9780598985408. From p. 142: "We are inclined to think of substances as possessing acid or basic properties, without having a particular solvent in mind. It seems to me that with complete generality we may say that a basic substance is one which has a lone pair of electrons which may be used to complete the stable group of another atom, and that an acid substance is one which can employ a lone pair from another molecule in completing the stable group of one of its own atoms. In other words, the basic substance furnishes a pair of electrons for a chemical bond, the acid substance accepts such a pair."
- OCLC 55600610.[page needed]
- .
- ^ ISBN 0-471-60180-2.[page needed]
- OCLC 1007924903.)
{{cite book}}
: CS1 maint: location missing publisher (link - OCLC 48850987.
- PMID 11671188.
- ISBN 0-7506-3365-4.[page needed]
- ^
- ^ Christian Laurence and Jean-François Gal "Lewis Basicity and Affinity Scales : Data and Measurement"
Wiley, 2009. ISBN 978-0-470-74957-9.[page needed]
- ISBN 978-3-540-64336-4.
- doi:10.1139/v82-117.
- .
- .
- ^ Miessler, L. M., Tar, D. A., (1991) p. 166 – Table of discoveries attributes the date of publication/release for the Lewis theory as 1923.
- S2CID 95865413.
- .
Further reading
- ISBN 0-471-03902-0.
- Yamamoto, Hisashi (1999). Lewis acid reagents : a practical approach. New York: Oxford University Press. ISBN 0-19-850099-8.