lewis acids and bases


  • [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 dative bond — for example, Me3B NH3.

  • • Examples of Lewis bases based on the general definition of electron pair donor include: o simple anions, such as H− and F− o other lone-pair-containing species, such as
    H2O, NH3, HO−, and CH3− o complex anions, such as sulfate o electron-rich π-system Lewis bases, such as ethyne, ethene, and benzene The strength of Lewis bases have been evaluated for various Lewis acids, such as I2, SbCl5, and BF3.

  • Some Lewis acids bind with two Lewis bases, a famous example being the formation of hexafluorosilicate: SiF4 + 2 F− SiF62− Complex Lewis acids[edit] Most compounds considered
    to be Lewis acids require an activation step prior to formation of the adduct with the Lewis base.

  • A Lewis base, then, is any species that has a filled orbital containing an electron pair which is not involved in bonding but may form a dative bond with a Lewis acid to form
    a Lewis adduct.

  • 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 as Me3B·NH3.

  • In a slightly different usage, the center dot is also used to represent hydrate coordination in various crystals, as in MgSO4·7H2O for hydrated magnesium sulfate, irrespective
    of whether the water forms a dative bond with the metal.

  • H+ as Lewis acid[edit] The proton (H+) [11] is one of the strongest but is also one of the most complicated Lewis acids.

  • 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.

  • [12] Applications of Lewis bases[edit] Main article: Homogeneous catalysis Nearly all electron pair donors that form compounds by binding transition elements can be viewed

  • In 1923, Lewis wrote An 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.

  • With this simplification in mind, acid-base reactions can be viewed as the formation of adducts: Applications of Lewis acids[edit] A typical example of a Lewis acid in action
    is in the Friedel–Crafts alkylation reaction.

  • 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.

  • A Lewis acid (named for the American physical chemist Gilbert N. Lewis) is a chemical species that contains an empty orbital which is capable of accepting an electron pair
    from a Lewis base to form a Lewis adduct.

  • The classification into hard and soft acids and bases (HSAB theory) followed in 1963.

  • Comparison with Brønsted–Lowry theory[edit] 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.

  • [2]The terms nucleophile and electrophile are sometimes interchangeable with Lewis base and Lewis acid, respectively.

  • The model assigned E and C parameters to many Lewis acids and bases.

  • The equation is The W term represents a constant energy contribution for acid–base reaction such as the cleavage of a dimeric acid or base.

  • 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, 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).

  • Reformulation of Lewis theory[edit] Lewis had suggested in 1916 that two atoms are held together in a chemical bond by sharing a pair of electrons.

  • Although there have been attempts to use computational and experimental energetic criteria to distinguish dative bonding from non-dative covalent bonds,[4] 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.

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

  • [13] Hard and soft classification Lewis acids and bases are commonly classified according to their hardness or softness.

  • Thus, a large application of Lewis bases is to modify the activity and selectivity of metal catalysts.

  • The industrial synthesis of the anti-hypertension drug mibefradil uses a chiral Lewis base (R-MeOBIPHEP), for example.

  • In a Lewis adduct, the Lewis acid and base share an electron pair furnished by the Lewis base, forming a dative bond.


Works Cited

[‘1. IUPAC, Compendium of Chemical Terminology, 2nd ed. (the “Gold Book”) (1997). Online corrected version: (2006–) “Lewis acid”. doi:10.1351/goldbook.L03508
2. ^ Jump up to:a b Lewis, Gilbert Newton (1923). Valence and the Structure of Atoms and
Molecules. American chemical society. Monograph series. New York, New York, U.S.A.: Chemical Catalog Company. p. 142. 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.”
3. ^ Anslyn,
Eric V. (2006). Modern physical organic chemistry. Dougherty, Dennis A., 1952-. Sausalito, CA: University Science. ISBN 1891389319. OCLC 55600610.[page needed]
4. ^ Lepetit, Christine; Maraval, Valérie; Canac, Yves; Chauvin, Remi (2016). “On the
Nature of the Dative Bond: Coordination to Metals and Beyond. The Carbon Case”. Coordination Chemistry Reviews. 308: 59–75. doi:10.1016/j.ccr.2015.07.018.
5. ^ Jump up to:a b March, J. “Advanced Organic Chemistry” 4th Ed. J. Wiley and Sons, 1992:
New York. ISBN 0-471-60180-2.[page needed]
6. ^ Vollhardt, K. Peter C. (2018). Organic chemistry : structure and function. Neil Eric Schore (8th ed.). New York. p. 73. ISBN 978-1-319-07945-1. OCLC 1007924903.
7. ^ Carey, Francis A. (2003). Organic
chemistry (5th ed.). Boston: McGraw-Hill. p. 46. ISBN 0-07-242458-3. OCLC 48850987.
8. ^ IUPAC, Compendium of Chemical Terminology, 2nd ed. (the “Gold Book”) (1997). Online corrected version: (2006–) “Electrophile (Electrophilic)”. doi:10.1351/goldbook.E02020
9. ^
Rowsell, Bryan D.; Gillespie, Ronald J.; Heard, George L. (1999). “Ligand Close-Packing and the Lewis Acidity of BF3 and BCl3”. Inorganic Chemistry. 38 (21): 4659–4662. doi:10.1021/ic990713m. PMID 11671188.
10. ^ Greenwood, N. N.; & Earnshaw, A.
(1997). Chemistry of the Elements (2nd Edn.), Oxford:Butterworth-Heinemann. ISBN 0-7506-3365-4.[page needed]
11. ^ Jump up to:a b c Traditionally, but not precisely, H+ ions are referred as “protons”. See IUPAC, Compendium of Chemical Terminology,
2nd ed. (the “Gold Book”) (1997). Online corrected version: (2006–) “hydron”. doi:10.1351/goldbook.H02904
12. ^ Christian Laurence and Jean-François Gal “Lewis Basicity and Affinity Scales : Data and Measurement” Wiley, 2009. ISBN 978-0-470-74957-9.[page
13. ^ Jacobsen, E.N.; Pfaltz, Andreas; Yamamato, H., eds. (1999). Comprehensive Asymmetric Catalysis. Berlin; New York: Springer. pp. 1443–1445. ISBN 978-3-540-64336-4.
14. ^ Childs, R.F; Mulholland, D.L; Nixon, A. (1982). “Lewis acid
adducts of α,β-unsaturated carbonyl and nitrile compounds. A nuclear magnetic resonance study”. Can. J. Chem. 60 (6): 801–808. doi:10.1139/v82-117.
15. ^ Vogel, Glenn C.; Drago, Russell S. (1996). “The ECW Model”. Journal of Chemical Education.
73 (8): 701. Bibcode:1996JChEd..73..701V. doi:10.1021/ed073p701.
16. ^ Cramer, Roger E.; Bopp, Thomas T. (1977). “Great e and C plot. Graphical display of the enthalpies of adduct formation for Lewis acids and bases”. Journal of Chemical Education.
54 (10): 612. Bibcode:1977JChEd..54..612C. doi:10.1021/ed054p612.
17. ^ Miessler, L. M., Tar, D. A., (1991) p. 166 – Table of discoveries attributes the date of publication/release for the Lewis theory as 1923.
18. ^ Lewis, Gilbert N. (April 1916).
“The atom and the molecule”. Journal of the American Chemical Society. 38 (4): 762–785. doi:10.1021/ja02261a002. S2CID 95865413.
19. ^ Brown, Herbert C.; Kanner, Bernard (1966). “Preparation and Reactions of 2,6-Di-t-butylpyridine and Related Hindered
Bases. A Case of Steric Hindrance toward the Proton”. Journal of the American Chemical Society. 88 (5): 986–992. doi:10.1021/ja00957a023.
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