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:
| A+ | + | :B– | → | A←:B |
|
Lewis acid (lone pair acceptor) |
Lewis base (lone pair donor) |
Coordinate bond |
- 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