Lewis base

(noun)

Any compound that can donate a pair of electrons and form a coordinate covalent bond.

Related Terms

Examples of Lewis base in the following topics:

  • Lewis Acid and Base Molecules

    • Lewis bases are electron-pair donors, whereas Lewis acids are electron-pair acceptors.
    • A Lewis acid is defined as an electron-pair acceptor, whereas a Lewis base is an electron-pair donor.
    • A Lewis base, therefore, is any species that donates a pair of electrons to a Lewis acid.
    • Under the Lewis definition, hydroxide acts as the Lewis base, donating its electron pair to H+.
    • We first look at the Bronsted-Lowry theory, and then describe Lewis acids and bases according to the Lewis Theory.

  • Metal Cations that Act as Lewis Acids

    • Transition metals can act as Lewis acids by accepting electron pairs from donor Lewis bases to form complex ions.
    • The modern-day definition of a Lewis acid, as given by IUPAC, is a molecular entity—and corresponding chemical species—that is an electron-pair acceptor and therefore able to react with a Lewis base to form a Lewis adduct; this is accomplished by sharing the electron pair furnished by the Lewis base.
    • However, metal ions such as Na+, Mg2+, and Ce3+ often form Lewis adducts upon reacting with a Lewis base.
    • Ligands create a complex when forming coordinate bonds with transition metals ions; the transition metal ion acts as a Lewis acid, and the ligand acts as a Lewis base.
    • Usually, metal complexes can only serve as Lewis acids after dissociating from a more weakly bound Lewis base, often water.

  • Acid-Base Reactions

    • According to the Lewis theory, an acid is an electron pair acceptor, and a base is an electron pair donor.
    • Lewis bases are also Brønsted bases; however, many Lewis acids, such as BF3, AlCl3 and Mg2+, are not Brønsted acids.
    • Two examples of Lewis acid-base equilibria that play a role in chemical reactions are shown in equations 1 & 2 below.
    • A terminology related to the Lewis acid-base nomenclature is often used by organic chemists.
    • Here the term electrophile corresponds to a Lewis acid, and nucleophile corresponds to a Lewis base.

  • Nature of Acids and Bases

    • Lewis acid: any substance that can accept a pair of electrons.
    • A list of various Lewis bases (right) and Lewis acids (left).
    • Acids + Bases Made Easy!
    • Part 1 - What the Heck is an Acid or Base?
    • In this video I introduce to you guys what the heck an Acid and Base really is forgetting the Lewis or Bronstead/Lowry definitions and then we'll go more in depth in parts 2,3, and 4.

  • Trihalides: Boron-Halogen Compounds

    • Trihalides adopt a planar trigonal structure and are Lewis acids.
    • The trihalides form planar trigonal structures and are Lewis acids because they readily form adducts with electron-pair donors, which are called Lewis bases.
    • All three lighter boron trihalides, BX3 (X = F, Cl, Br), form stable adducts with common Lewis bases.
    • Such measurements have revealed the following sequence for the Lewis acidity: BF3 < BCl3 < BBr3 (in other words, BBr3 is the strongest Lewis acid).
    • BF3 is commonly referred to as "electron deficient" because of its exothermic reactivity toward Lewis bases.

  • Strong Bases

    • As discussed in the previous concepts on bases, a base is a substance that can: donate hydroxide ions in solution (Arrhenius definition); accept H+ ions (protons) (Bronsted-Lowry definition); or donate a pair of valence electrons (Lewis definition).
    • Strong bases are capable of deprotonating weak acids; very strong bases can deprotonate very weakly acidic C–H groups in the absence of water.
    • Generally, the alkali metal bases are stronger than the alkaline earth metal bases, which are less soluble.
    • When writing out the dissociation equation of a strong base, assume that the reverse reaction does not occur, because the conjugate acid of a strong base is very weak.
    • Unlike weak bases, which exist in equilibrium with their conjugate acids, the strong base reacts completely with water, and none of the original anion remains after the base is added to solution.

  • The Brønsted-Lowry Definition of Acids and Bases

    • Originally, acids and bases were defined by Svante Arrhenius.
    • Keep in mind that acids and bases must always react in pairs.
    • Here, ammonia is the Brønsted-Lowry base.
    • Chemistry 12.1 What are Acids and Bases?
    • We first look at the Brønsted-Lowry theory, and then describe Lewis acids and bases according to the Lewis Theory

  • Bonding in Coordination Compounds: Valence Bond Theory

    • Lewis proposed that the basis of chemical bonding is in the ability of atoms to share two bonding electrons.
    • In 1927, physicist Walter Heitler and collaborator Fritz London developed the Heitler-London theory, which enabled the calculation of bonding properties of the covalently bonded diatomic hydrogen molecule (H2) based on quantum mechanical considerations.
    • Finally, Linus Pauling integrated Lewis' proposal and the Heitler-London theory to give rise to two additional key concepts in valence bond theory: resonance and orbital hybridization.
    • Valence bond structures are similar to Lewis structures, except where a single Lewis structure is insufficient, several valence bond structures can be used.
    • Calculate the theoretical hybridization of a metal in a coordination complex based on valence bond theory

  • Regioselectivity and Lewis Acid Catalysis

    • Unfortunately, neither molecular orbital symmetry analysis nor the simple mnemonic rules based on electron counts explain these regioselectivities.
    • Both Diels-Alder and ene reactions are catalyzed by Lewis acids.
    • In some cases Lewis acid catalysis may change the regioselectivity of a Diels-Alder reaction.
    • In many cases, this analysis of HOMO and LUMO orbital coefficients also provides a good explanation for the beneficial influence of Lewis acid catalysis.
    • Lewis acids complex with the basic oxygen atom of such functions, rendering them more electrophilic.

  • Lewis Structures for Polyatomic Ions

    • The Lewis structure of an ion is placed in brackets and its charge is written as a superscript outside of the brackets, on the upper right.
    • The total number of electrons represented in a Lewis structure is equal to the sum of the numbers of valence electrons in each individual atom.
    • Non-valence electrons are not represented in Lewis structures.
    • Lewis structures for polyatomic ions are drawn by the same methods that we have already learned.
    • When counting electrons, negative ions should have extra electrons placed in their Lewis structures; positive ions should have fewer electrons than an uncharged molecule.