How is the Brønsted and Lowry Acid-Base theory defined?

A substance that is capable of donating a proton is an acid, while the one that receives that proton is a base. This very general definition of acids and bases was brought by the chemists J.N. Brønsted and T.M. Lowry in 1923, based on the transfer concept of H⁺ in an acid–base reaction.

Arrhenius defined protons H⁺ as isolated species, although today it is known that in solution they have a high attraction with water molecules and are forming hydronium ions (\({H_3}{O^ + }\)). Based on these two concepts we explore a known acid-base reaction:

\(H{C_2}{H_3}{O_2}_{\left( {ac} \right)} + {H_2}{O_{\left( l \right)}} \leftrightarrow {C_2}{H_3}{O_2 }{^ –{\left( {ac} \right)}} + \;{H_3}{O^ + }_{\left( {ac} \right)}\)

In this case, acetic acid is the one that donates an acidic hydrogen while water acts as a base, taking the donated proton. In turn, two new ionic species are formed, which are the acids and the conjugate bases of the acids and bases from which they came. In this case, the species \({C_2}{H_3}{O_2}^ – \) is the conjugate base of acetic acid while \({H_3}{O^ + }\) is the conjugate acid of water. Therefore, the conjugate acid-base pair only differs in the presence of an acidic hydrogen and, furthermore, the premise that every acid has its conjugate base and vice versa is fulfilled.

Now let's review the following reaction:

\(N{H_3}_{\left( {ac} \right)} + {H_2}{O_{\left( l \right)}} \leftarrow N{H_4}{^ + {\left( {ac } \right)}} + \;O{H^ – }_{\left( {ac} \right)}\)

In this case, we have a conjugate acid-base pair that is water and hydroxyl ion respectively, and a base, ammonia, with its conjugate pair, the species of acid character \(N{H_4}^ + \).

Now, you may wonder, how is it that water acts as both an acid and a base? That ability is known as amphotericism. That is, a substance that can act in both ways depending on who it is combined with is an amphoteric substance.

Just as we define conjugate pairs, they have a peculiar characteristic: the more acidic strength the acid in the pair has, the lower the basic strength. will have its conjugate base, and it is analogous to the case of the bases, the greater the strength of basicity the base has, its conjugate pair will decrease the strength of the acid. They will wonder what force are we talking about?

Well then, when an acid is strong we are talking about a species that is capable of completely donating acidic hydrogen, transferring all its protons to water and dissociating completely. Otherwise, weak acids are partially ionized in aqueous solution, this implies that part of the acid will be found as dissociated species and part will retain its structure. Let's look at the following typical examples:

\(HC{l_{\left( g \right)}} + {H_2}{O_{\left( l \right)}} \to C{l^ – }_{\left( {ac} \right) } + \;{H_3}{O^ + }_{\left( {ac} \right)}\)

This is a strong acid, since it dissociates completely, and similarly occurs with sodium hydroxide, which is a strong base:

\(NaO{H_{\left( s \right)}} \to N{a^ + }_{\left( {ac} \right)} + \;O{H^ – }_{\left( { ac} \right)}\)

If we explore the reaction of acetic acid in aqueous solution, we note that there is an equilibrium between the species, since the dissociation is not complete and, therefore, there is a thermodynamic acidity constant that governs the process, which is expressed as a function of the activities of the species; however, in dilute solutions, it can be estimated through the molar concentrations:

\(Ka = \frac{{\left[ {{C_2}{H_3}{O_2}^ – } \right]\left[ {{H_3}{O^ + }} \right]}}{{\left[ {H{C_2}{H_3}{O_2}} \right]}}\)

While for the case of weak bases we can describe the degree to which said base ionizes if we talk about its thermodynamic constant of basicity, such is the case of ammonia:

\(Kb = \frac{{\left[ {N{H_4}^ + } \right]\left[ {O{H^ – }} \right]}}{{\left[ {N{H_3}} \ right]}}\)

These constants are tabulated at reference temperatures while there is also a bibliography that indicates the level of acidity or basicity of certain compounds.

Finally, we will refer to the autoionization of water, as we have already seen, water has both a base and a conjugate acid, being able to describe this phenomenon in its ionization reaction:

\(2{H_2}{O_{\left( l \right)}} \leftrightarrow \) \(O{H^ – }_{\left( {ac} \right)} + {H_3}{O^ + }_{\left( {ac} \right)}\)

We could define this process as we did previously through the constant involved, which would be:

\(Kc = \frac{{\left[ {{H_3}{O^ + }} \right]\left[ {O{H^ – }} \right]}}{{{{\left[ {{H_2 }O} \right]}^2}}}\)

Resorting to a mathematical arrangement we could express the ionic product of water as the following constant:

\(Kw = \left[ {{H_3}{O^ + }} \right]\left[ {O{H^ – }} \right]\)

Whose value at 25ºC is constant and is: 1×10-14, which implies that, if the solution is neutral, that is, equal quantity of acid than base, each of the concentrations of the ionic species will be: 1×10-7 mol/L.