Lewis Acids & Bases (HL) (DP IB Chemistry)

Lewis Acids & Bases

A more general definition of acids and bases was given by G.N. Lewis who defined them as:

• A Lewis acid is an lone pair acceptor

• A Lewis base is an lone pair donor

• The general mechanism for Lewis acids and bases can be represented as:

• This enabled a wider range of substances to be classed as acids or bases

• This can be shown in the following examples in which a hydroxide ion, OH^–, and ammonia, NH₃, donate a lone pair to a hydrogen ion

Diagram to show how OH^– and ammonia act as Lewis bases

The OH^– ion and ammonia act as Lewis bases in both examples by donating a lone pair of electrons

How are Brønsted-Lowry Acids and Bases Different from Lewis Acids and Bases?

• A Brønsted-Lowry acid is a species that can donate H⁺

• For example, hydrogen chloride (HCl) is a Brønsted-Lowry acid as it can donate a H⁺ ion

HCl (aq) → H⁺ (aq) + Cl^– (aq)

• Lewis acids, by definition, covers a boarder spectrum than Brønsted-Lowry acids

• Lewis acids are any compounds that are able to accept a lone pair of electrons which includes H⁺ itself

• Brønsted-Lowry acid and base theory considers acids as H⁺ donors only

• This does not of course occur in every reaction

• A Brønsted-Lowry base is a species that can accept H⁺

  • For example, a hydroxide (OH^–) ion is a Brønsted-Lowry base as it can accept H⁺ to form water

• Lewis bases and Brønsted-Lowry bases are in the same group of compounds as both of these must have a lone pair of electrons 

• The following molecules can behave as either Lewis bases and Brønsted-Lowry bases

  • Lewis bases as they can donate an electron pair

  • Brønsted-Lowry base as they can accept a proton

Hydroxide, cyanide and methylamine 

Examples of molecules that can behave both as Lewis bases and Brønsted-Lowry base