The Bronsted-Lowry Theory
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The Bronsted-Lowry Theory
In 1923 Bronsted and Lowry proposed the proton transfer theory of acids and bases.
Their theory states that an acid is a substance which donates protons, and a base is a substance which accepts protons.
If we observe the behaviour of a strong acid such as hydrochloric acid, we note that the hydrogen chloride fully dissociates, forming hydrogen ion, H+, and chloride ions, Cl.
The hydrogen ions are also known as protons since a hydrogen ion consists of a single proton only. These protons are donated to the water molecules forming oxonium ions, H3O+.
If we now observe a weak acid such as ethanoic acid, the acid molecules are only partially dissociated in water, and so the equilibrium is established:
Even so, the acid molecules donate protons to water molecules.
Strong bases such as an alkali metal hydroxides are fully ionised in water:
They are bases because the hydroxide ions can accept hydrogen ions:
This reaction is common to all neutralisation reactions between acids and bases in aqueous solution.
Note: In aqueous solutions the hydrogen ions exist in their hydrated forms, that is, as oxonium ions.
If we examine an example of equilibria a little more closely, we can begin to label the species in the equilibrium mixture.
The equilibrium mixture is said to consist of two conjugate pairs of acids and bases:
NH3 and NH4+ are a conjugate pair. HCl and Cl- are a conjugate pair also.
Note: In each conjugate pair, the acid and base only differ from one another by a proton. Each acid has its own conjugate base.
A strong acid always has a weak conjugate base.
A weak acid always has a strong conjugate base.
The equilibrium law can be applied to aqueous solutions of acids. For example, the following equilibrium is established in an aqueous solution of ethanoic acid:
The equilibrium constant is given by Ka =
 represents concentration.
Ka is called the Acid Dissociation Constant.
The acid dissociation constant is a measure of the strength of an acid. For an acid such as hydrochloric acid which is virtually fully dissociated in aqueous solution, its value is extremely large.
On the other hand, for weak acids, values of Ka can be extremely small. It is often more convenient to compare the strengths of acids using pKa values, where pKa is given by pKa = -lgKa.
For most acids this gives the range of values from 0-14.
Strong acids have low pKa values and weak acids have large values.
The electronic conductivity of even the purest water never falls to exactly zero. This is due to the SELF-IONISATION of water. This can be represented by:
Applying the equiibrium law to this equilibrium, we obtain:
- Kw = [H3O+][OH-] eq
- Kw = [H+][OH-] eq
Kw is called the ionic product of water. It has units of (concentration)2; that is mol2dm-2. The exact value depends on the temperature.At 25oC its value is 1.0 x 10-14.
This gives a pKw value of 14, where: pKw = -lgKw
The ionic product of water is related to the dissociation constants pKa and pKb of an acid and its conjugate base respectively.
pKa + pKb = pKw = 14 (at 25o C)
Thus, if the pKa value of an acid is known, the pKb value of its conjugate base can be found.