Lewis Acids & Bases (HL) (DP IB Chemistry): Revision Note
Lewis Acids & Bases
A more general definition of acids and bases was given by G.N. Lewis:
A Lewis acid is an lone pair acceptor
A Lewis base is an lone pair donor
This enabled a wider range of substances to be classed as acids or bases, beyond those that donate or accept protons
Lewis acid and base interactions
The interaction between a Lewis acid and base can be shown as:
A+ + :B- → A←:B
A coordinate covalent bond forms when the Lewis base donates a lone pair of electrons to the Lewis acid
Lewis theory is broader than Brønsted–Lowry theory because it includes reactions that do not involve proton transfer
Some species, such as OH- and NH3, act as both types of base:
As Lewis bases they donate a lone pair of electrons
As Brønsted–Lowry bases they accept a proton
Diagram to show how OH– and ammonia act as Lewis bases

Comparing Lewis and Brønsted–Lowry acid-base theory
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 are any species that are able to accept a lone pair of electrons
This includes H⁺
This means Lewis theory covers a broader spectrum than Brønsted–Lowry acids
Brønsted-Lowry theory strictly defines acids as H+ donors only
Not all acid–base reactions involve proton transfer
This is where Lewis theory is more broadly applicable
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
H+ (aq) + OH- (aq) → H2O (l)
Lewis bases and Brønsted-Lowry bases are in the same group of chemicals because they both must have a lone pair of electrons
Lewis bases donate a lone pair of electrons
Brønsted-Lowry bases accept a proton vis their lone pair of electrons
Examples of chemicals that can act as Lewis bases and Brønsted-Lowry bases include:
Hydroxide, OH-
Cyanide, CN-
Methylamine, CH3NH2

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