The term acid is derived from the Latin word acidus which means sour. Early chemists had a list of properties that were common to the substances that they considered to be acids or bases [eg acids had a sour taste, turned litmus red, reacted with some metals to produce a flammable gas (hydrogen) ..etc.. ]. They would assess a new substance as an acid or as a base (or as neither) by comparing the properties of the new substance against the list of properties.
The first modern approach to acid-base chemistry was by Arrhenius in 1887. He defined an acid as a substance which was capable of dissociating in water solution to produce hydrogen ions. This definition identified most of the substances which were considered to be acids at that time. A base was defined as a substance which dissociated in water solution to produce hydroxide ions. The theory was not totally satisfactory for several reasons. For example, some substances which had acidic properties did not contain hydrogen and some bases did not contain hydroxide ions. The theory also applied only to aqueous solutions.
The next development was the Bronsted-Lowry Theory (1923) and this is the approach which is generally accepted in biological and medical fields. An acid is defined as a substance which donates a hydrogen ion to another substance. This does not require an aqueous solution or dissociation into ions as in the Arrhenius definition. The substance which accepts the H+ from the acid is called the ‘conjugate base’. This idea of conjugate acid-base pairs is an important part of the Bronsted-Lowry approach. Acid strength is defined in terms of the strength of the tendency to donate the hydrogen ion to the solvent (ie water in biological systems). A strong acid has a high tendency to donate a proton to water; so the [H3O+] is high.
A more general definition of acids and bases is the approach of Lewis in 1923. The impetus here was the problem of substances which exhibited acidic properties in solution (eg CO2) but did not contain a H+. Lewis defined an acid as any compound that was a potential electron pair acceptor and a base as any compound that was a potential electron pair donor. In the Lewis scheme, H+ itself is an acid.
From the medical and biological perspective, the Bronsted-Lowry theory is easy to understand and encompasses all the biological acids and bases encountered in aqueous solutions. It is the preferred approach.(CO2 is not strictly an acid in the Bronsted-Lowry system as it has no hydrogen ion but it can be accommodated by considering carbonic acid ( H2CO3 ) as the acid.)
In reality, most physicians have a basic knowledge of acids and bases which is somewhat of an combination of the Arrhenius approach (acid: H+ in solution), the Bronsted-Lowry approach (acid = proton donor) and even the Lewis approach (eg CO2 as an acid). This level of understanding is generally satisfactory for clinical purposes. The table below summarises the different approaches.
Basic Principles of the Various Theories of Acids and Bases
Acid: a substance that has certain properties
(eg sour taste, turns litmus red)
Acid : H+ in aqueous solution
Base : OH- in aqueous solution
At neutrality: [H+] = [OH-]
Acid : H+ donor
Base : H+ acceptor
Conjugate acid-base pairs
No concept of neutrality
Acid : a potential electron-pair acceptor
Base : a potential electron-pair donor
Acid: a substance that donates a cation, or accepts an anion or an electron
Base: a substance that donates an anion, or accepts a cation.