Q.All Lewis acids are not Bronsted acids. Why?

A Lewis acid only accepts an electron pair; many (like BF3) have no ionisable hydrogen, so they cannot give a proton and hence are not Bronsted acids.

Definitions:

• A Bronsted acid is a substance that can donate a proton (H+).
• A Lewis acid is a substance that can accept a lone pair of electrons.

Every Bronsted acid, when it donates a proton, is also acting as an electron-pair acceptor in a wider sense, but the reverse is not true.

Many Lewis acids do not contain any hydrogen that they can release as H+. For example:

• BF3: boron is electron-deficient and accepts a lone pair, so it is a Lewis acid; but it has no hydrogen to donate, so it cannot be a Bronsted acid.
• AlCl3, similarly, is a Lewis acid but not a Bronsted acid. …

The Arrhenius concept defines acids and bases purely in terms of what ion they release when dissolved in water: H+ for acids, OH- for bases.

Arrhenius acid: A substance which, when dissolved in water, increases the concentration of hydrogen ions, H+ (in practice these combine with water to form hydronium ions, H3O+).

Example: HCl dissolves in water as HCl(aq) -> H+(aq) + Cl-(aq), increasing [H+].

Arrhenius base: A substance which, when dissolved in water, increases the concentration of hydroxide ions, OH-.

Example: NaOH dissolves in water as NaOH(aq) -> Na+(aq) + OH-(aq), increasing [OH-].

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A Bronsted base is a proton (H+) acceptor; NH3 is a common example.

The Bronsted-Lowry theory defines acids and bases in terms of proton transfer:

• A Bronsted acid is a proton (H+) donor.
• A Bronsted base is a proton (H+) acceptor.

Example: When ammonia is dissolved in water,

NH3 + H2O -> NH4+ + OH-

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