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Brønsted-Lowry's acids and bases

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pH calculation lectures » Brønsted-Lowry's acids and bases

As all reactions we are interested in take place in water (or other solvent, but we won't touch the subject here) and water dissociates itself into H+ and OH- ions, classic definition of acid as a substance that dissociates producing H+ ions (the one we used in equation 1.1) becomes a little bit problematic. Consider the solution of salt of weak base BOH. Such solution contains B+ ions, that are between products of BOH dissociation:

base dissociation

with equilibrium described by the already mentioned in the previous section base dissociation constant:

base dissociation constant, eq. 2.12.1

For the equilibrium BOH molecules are needed. As there are already OH- ions from the water dissociation present in the solution, they will react with B+. This will lower OH- concentration, forcing water to dissociate further. Final solution in equilibrium will contain some BOH molecules and some excess of H+ - so it will be acidic, even if we haven't add any acid! Seems that B+ is an acid - even if it doesn't dissociate to give H+ ions.

To overcome this inconsistency Brønsted and Lowry proposed independently in 1923 new definitions of acid and base: acid is a substance that can donate the proton and base is a substance than can accept the proton. The most important outcome of this definition is the fact that every acid loosing its proton becomes a Brønsted-Lowry base (as it has a free "slot" for the proton) and that every base when protonated becomes a Brønsted-Lowry acid (it has a proton that is can release). These pairs of acid and base are called conjugate. In other words every acid loosing proton becomes its conjugate base, and every protonated base becomes its conjugate acid.

This approach has some interesting implications. Let's take a reaction of conjugate base A- with water:

reaction between conjugate base and water

Its equilibrium constant is

eq. 2.22.2

(water concentration assumed constant). Multiplying this equation by the equation for acid dissociation constant we get

eq. 2.32.3

[A-] and [HA] cancel out leaving

base dissociation constant, eq. 2.42.4


base dissociation constant, eq. 2.52.5

The most important lesson is that to describe acid/base properties of substance in the solution (not necessarily water solution) we can use either Ka or Kb value. Sometimes it is more convenient to use Kb for calculations, but whenever selection of constant doesn't matter we will concentrate our efforts around Ka according to widely accepted convention.

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