Lewis Acids and Bases

Introduction

Chemical reactions originate from the flow of electrons from an electron source to an electron sink.  This is analogous to the flow of electricity in wires and circuits.  That is the electrons (electricity) flows from a negative source (powerplant) to a positive sink (your house). An important skill that will aid you in predicating the outcome of chemical reactions is the ability to identify which species can act as an electron source an which can act as an electron sink. Our first example of electrophiles are Lewis acids.

By definition, a Lewis base is an electron pair donor, and thus acts as a source of electrons. A Lewis acid is an electron pair acceptor, and thus acts as an electron sink. When a Lewis base donates a pair of electrons to a carbon atom, it is further described as a "nucleophile", which means nucleus loving. The nucleus is the positive part of an atom, thus something that is electron rich would seek, or "love" the nucleus, hence the name nucleophile. Conversely, an electrophile is electron "loving" species because it is electron deficient.

Nucleophilicity and basicity are two different animals.  We will discuss the difference again many times, but for know remember that nucleophilicity is a kinetic property and nucleophiles are sensitive to sterics, while on the other hand basicity is more of a thermodynamic property and equilibrium is important or implied.

Lewis Acids

A. Neutral compounds containing B or Al: BF₃, BH₃, AICl₃

Neutral boron and aluminum compounds have an empty p-orbital that can accept a pair of electrons and are notorously good Lewis acids.  We will see them behave like electrophiles in various reactions.

B. Carbocations: CH₃⁺, (CH₃)₃C⁺, etc.

Isoelectric to BH₃ and BF₃ discussed previously, carbocations also have an empty p-orbital. Since they are in group IV, having just three bonds (and no lone pair electrons) results in a formal charge of + 1 on the carbon (in addition to the empty orbital). Thus, these are very strong electrophiles that often cannot be isolated, but instead are reactive intermediates in substitution and elimination reactions.

C. Strong acids: H₂SO₄, HX. HNO₃ etc.

Strong acids are source of H⁺ (protons). While they are typically solvated, and bound to solvents, protons are just empty s orbitals and as such are Lewis acids.

D. Electronegative atoms bound to other electronegative atoms: (X2, HOOH, ROOH, O=O-O, etc)

O₃, F₂ and Cl₂ are very strong electrophiles (Lewis acids), peroxides (HOOH and HOOR) and Br₂ are
strong to moderate, and while I₂ is a weak electrophile.  Soon you will see the electrophilic nature of Br₂, Cl₂ and O₃ in halogenation and ozonolysis of alkenes.

Lewis Bases

A. Strong to Moderate Lewis Bases: Same as Bronsted Bases

Generally, the same structural features that make a strong Bronsted base also lead to strong Lewis bases. We will discuss in detail the differences between a Bronsted base and a nucleophile in coming sections. In general, an atom will be more willing to donate a pair of electron if: a) the lone pair electrons are on an atom of low electronegativity, and/or b) the lone pair electrons are far away from the nuclei, i.e. a large atom.

Typical organic Lewis bases include the following.

B. Weak Lewis Bases: Compounds with π bonds.

Compounds that contain π bonds can act as a Lewis bases, but donation of a pair of π electrons results in the loss of an octet for one of the atoms involved in the double bond. Thus, the π-electrons are not donated as readily as lone pair electrons.  Weak Lewis bases would include alkenes and alkynes.