3.4.3 Acids and Bases - Bronsted.Lowry acid-base equilibria in aqueous solution
Students should:
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Bronsted-Lowry acids and bases
Bronsted-Lowry theory defines acids as proton (H+ ion) donors (and bases as proton acceptors).
For a compound to act as a Bronsted-Lowry acid, it must have a hydrogen atom
in it, which it is capable or losing while remaining fairly stable. A Bronsted-Lowry
base must be capable of accepting a hydrogen ion while remaining relatively
stable (or reacting to form a stable compound...eg water and a salt). Some compounds
(such as water) may act as both ie (H2O
OH- or H3O+)
Acid - base reactions always involve an acid-base conjugate pair...one is an acid, one is its conjugate base
Examples: | |||
Compound | hydrochloric acid | ethanoic acid | water |
Acid form | HCl | CH3COOH | H2O |
Conjugate base | Cl- | CH3COO- | OH- |
The conjugate base always has one less H atom that the acid (or the acid one more than the base). In compounds where there are many hydrogen atoms, the one which is held the weakest is generally the one which is lost, and this must be reflected when writing the formula...as in CH3COOH