Arrhenius, Bronsted, and Lewis

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Sean Sugai 4E
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Joined: Wed Sep 11, 2019 12:17 am

Arrhenius, Bronsted, and Lewis

Postby Sean Sugai 4E » Mon Nov 25, 2019 2:37 pm

What are the differences between Arrhenius, Bronsted, and Lewis acids and bases?

Sebastian Lee 1L
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Re: Arrhenius, Bronsted, and Lewis

Postby Sebastian Lee 1L » Mon Nov 25, 2019 3:12 pm

The definition of an Arrhenius acid is a molecule that forms a hydronium ion (H3O+) in the presence of water. An Arrhenius base forms a hydroxide ion (OH-) in water. This definition applies to most Bronsted acids/bases.

A Bronsted acid is a molecule that is a proton donor, meaning that it will give an H+ to another molecule, like water. HCl (aq) + H20 (l) --> H3O+ (aq) + Cl- (aq). In this example, hydrochloric acid gives a proton to water to form a hydronium ion. A Bronsted base is the molecule that receives a proton. In the above example, you would consider water to be the base. Another example: NH3 (aq) + H2O (l) --> NH4+ (aq) + OH- (aq). The ammonia (NH3) is a bronsted base because it accepts the H+ proton from water to form the ammonium ion. In this case water is the acid since it gives the proton.

A Lewis acid is a molecule that accepts a lone pair of electrons while a Lewis base donates a lone pair of electrons. The previous example with ammonia can also be interpreted with the Lewis definition since the NH3 is donating a lone pair (on the N) to the Hydrogen from the water. Similarly, HCl is a lewis acid because it accepts a lone pair from water to form a chloride anion. However, there are a few Lewis acid/bases that are NOT Bronsted acid/bases. For example, BF3 is a Lewis acid because it will accept a lone pair from F- in the reaction: BF3 + F- --> BF4-. However, BF3 doesn't donate any proton so you wouldn't consider it a Bronsted acid.

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