Lewis Acids and Bases

Lewis Acid and Base Molecules

Lewis bases are electron-pair donors, whereas Lewis acids are electron-pair acceptors.

Learning Objectives

Recognize Lewis acids and bases in chemical reactions.

Key Takeaways

Key Points

  • A Lewis acid is an electron -pair acceptor; a Lewis base is an electron-pair donor.
  • Some molecules can act as either Lewis acids or Lewis bases; the difference is context-specific and varies based on the reaction.
  • Lewis acids and bases result in the formation of an adduct rather than a simple displacement reaction, as with classical acids and bases. An example is HCl vs H+: HCl is a classical acid, but not a Lewis acid; H+ is a Lewis acid when it forms an adduct with a Lewis base.

Key Terms

  • covalent bond: a chemical bond in which two atoms are connected to each other by sharing two or more electrons
  • nucleophile: literally “lover of nuclei,” Lewis bases are often referred to as this because they seek to donate their electron pairs to electron-poor species, such as H+

A Lewis acid is defined as an electron-pair acceptor, whereas a Lewis base is an electron-pair donor. Under this definition, we need not define an acid as a compound that is capable of donating a proton, because under the Lewis definition, H+ itself is the Lewis acid; this is because, with no electrons, H+ can accept an electron pair.

A Lewis base, therefore, is any species that donates a pair of electrons to a Lewis acid. The “neutralization” reaction is one in which a covalent bond forms between an electron-rich species (the Lewis base) and an electron-poor species (the Lewis acid). For this reason, Lewis bases are often referred to as nucleophiles (literally, “lovers of nuclei”), and Lewis acids are sometimes called electrophiles (“lovers of electrons”). This definition is useful because it not only covers all the acid-base chemistry with which we are already familiar, but it describes reactions that cannot be modeled by Arrhenius or Bronsted-Lowry acid-base chemistry. For now however, we will consider how the Lewis definition applies to classic acid-base neutralization.

Applying the Lewis Definition to Classical Acid-Base Chemistry

Consider the familiar reaction of NaOH and HCl:

[latex]\text{NaOH}(\text{aq})+\text{HCl}(\text{aq})\rightarrow \text{NaCl}(\text{aq})+\text{H}_2\text{O}(\text{l})[/latex]

We have previously described this as an acid-base neutralization reaction in which water and a salt are formed. This is still completely correct, but the Lewis definition describes the chemistry from a slightly different perspective. When considering Lewis acids and bases, the only real reaction of interest is the net ionic reaction:

[latex]\text{OH}^-(\text{aq})+\text{H}^+(\text{aq})\rightarrow \text{H}_2\text{O}(\text{l})[/latex]

Under the Lewis definition, hydroxide acts as the Lewis base, donating its electron pair to H+. Thus, in this version of the neutralization reaction, what interests us is not the salt that forms, but the covalent bond that forms between OH and H+ to form water. A significant hallmark for Lewis acid-base reactions is the formation of such a covalent bond between the two reacting species. The reaction’s final product is known as an adduct, because it forms from the addition of the Lewis base to the Lewis acid.

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Lewis acids and bases: Lewis acids (BF3, top, and H+, bottom) react with Lewis bases (F, top, NH3, bottom) to form products known as adducts. Note that the first reaction cannot be described by Arrhenius or Bronsted-Lowry acid-base chemistry.

Beyond Classical Acid-Base Chemistry

By treating acid-base reactions in terms of electron pairs instead of specific substances, the Lewis definition can apply to reactions that do not fall under other definitions of acid-base reactions. For example, a silver cation behaves as a Lewis acid with respect to ammonia, which behaves as a Lewis base, in the following reaction:

[latex]\text{Ag}^+(\text{aq}) + 2\;\text{NH}_3 \rightarrow [\text{Ag}(\text{NH}_3)_2]^+[/latex]

This reaction results in the formation of diamminesilver(I), a complex ion; it is perfectly described by Lewis acid-base chemistry, but is unclassifiable according to more traditional Arrhenius and Bronsted-Lowry definitions.

Application to Organic Chemistry

In organic chemistry, it is useful to understand that nucleophiles are Lewis bases and electrophiles are Lewis acids. Nearly all reactions in organic chemistry can be considered Lewis acid-base processes.

What are acids and bases?: This lesson continues to describe acids and bases according to their definition. 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.

Learning Objectives

Recognize metals that function as Lewis acids.

Key Takeaways

Key Points

  • A Lewis acid is an electron pair acceptor; because metal ions have one or more empty orbitals, they act as Lewis acids when coordinating ligands.
  • Examples of metals that can act as Lewis acids include Na+, Mg2+, and Ce3+.
  • Metal ions rarely exist uncoordinated; they often have to dissociate from weaker ligands, like water, before complexing with other Lewis bases.

Key Terms

  • coordinate bond: a type of covalent bond in which two shared electrons originate from the same atom; a dative bond
  • ligand: the species that coordinates with a metal cation to form a complex ion
  • Complex ion: a compound consisting of a metal ion coordinated to various ligands in solution

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. Classically, the term “Lewis acid” was restricted to trigonal planar species with an empty p orbital, such as BR3 where R can be an organic substituent or a halide. However, metal ions such as Na+, Mg2+, and Ce3+ often form Lewis adducts upon reacting with a Lewis base.

Complex Ion Formation

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. The number of coordinate bonds is known as the complex’s coordination number. Common ligands include H2O and NH3 ; examples of complexes include the tetrachlorocobaltate(II) ion, [CoCl4]2- and the hexaqua-iron(III) ion, [Fe(H2O)6]3+.

Usually, metal complexes can only serve as Lewis acids after dissociating from a more weakly bound Lewis base, often water. For instance, Mg2+ can coordinate with ammonia in solutions, as shown below:

[latex][\text{Mg}(\text{H}_2\text{O})_6]^{2+} + 6\text{NH}_3 \rightarrow [\text{Mg}(\text{NH}_3)_6]^{2+} + 6\text{H}_2\text{O}[/latex]

Nearly all compounds formed by the transition metals can be viewed as collections of the Lewis bases—or ligands—bound to the metal, which functions as the Lewis acid. The product is known as a complex ion, and the study of these ions is known as coordination chemistry. One coordination chemistry’s applications is using Lewis bases to modify the activity and selectivity of metal catalysts in order to create useful metal-ligand complexes in biochemistry and medicine.

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Examples of metal-ligand coordination complexes: Examples of several metals (V, Mn, Re, Fe, Ir) in coordination complexes with various ligands. All these metals act as Lewis acids, accepting electron pairs from their ligands.