The theories of acids and bases are based on the concept given by Antoine Lavoisier in 1776, who had limited knowledge of strong acids, including nitric and sulfuric. Lavoisier claimed that the acidity of a substance depended on how much oxygen it contained, since he did not know the actual compositions of hydrogen halides and other strong acids.
This theory was taken as the true definition of acid for several decades, even when scientists like Berzelius and von Liebig made modifications and proposed other visions, but it was not until Arrhenius that it began to be seen more clearly how acids and bases worked.
Thomas Martin Lowry, one of the acid and base theorists
Following Arrhenius, the physicochemists Brönsted and Lowry independently developed their own theory, until Lewis came along to propose an improved and more accurate version of it.
This set of theories are used to this day and are said to be the ones that helped form modern chemical thermodynamics.
Arrhenius theory
The Arrhenius theory is the first modern definition of acids and bases, and it was proposed by the physicochemist of the same name in 1884. It states that a substance is identified as acid when it forms hydrogen ions by dissolving in water.
That is, the acid increases the concentration of H + ions in aqueous solutions. This can be demonstrated with an example of the dissociation of hydrochloric acid (HCl) in water:
HCl (aq) → H + (aq) + Cl - (aq)
According to Arrhenius, bases are those substances that release hydroxide ions when they dissociate in water; that is, it increases the concentration of OH - ions in aqueous solutions. An example of an Arrhenius base is the dissolution of sodium hydroxide in water:
NaOH (aq) → Na + (aq) + OH - (aq)
The theory also states that, as such, there are no H + ions, but that this nomenclature is used to denote a hydronium ion (H 3 O +) and that this was referred to as a hydrogen ion.
The concepts of alkalinity and acidity were only explained as the concentrations of hydroxide and hydrogen ions, respectively, and the other types of acid and base (their weak versions) were not explained.
Brönsted and Lowry theory
Johannes Nicolaus Bronsted
This theory was developed independently by two physicochemicals in 1923, the first in Denmark and the second in England. They both had the same vision: Arrhenius's theory was limited (since it completely depended on the existence of an aqueous solution) and did not correctly define what an acid and a base were.
For this reason, chemists worked around the hydrogen ion and made their claim: acids are the substances that release or donate protons, while bases are those that accept those protons.
They used an example to demonstrate their theory, which involved an equilibrium reaction. He claimed that each acid had its conjugate base, and that each base also had its conjugate acid, like this:
HA + B ↔ A - + HB +
Like, for example, in the reaction:
CH 3 COOH + H 2 O ↔ CH 3 COO - + H 3 O +
In the previous reaction, acetic acid (CH 3 COOH) is an acid because it donates a proton to water (H 2 O), thus becoming its conjugate base, the acetate ion (CH 3 COO -). In turn, water is a base because it accepts a proton from acetic acid and becomes its conjugated acid, the hydronium ion (H 3 O +).
This reverse reaction is also an acid-base reaction, as the conjugated acid is converted to acid and the conjugated base is converted to base, through the donation and acceptance of protons in the same way.
The advantage of this theory over Arrhenius is that it does not require an acid to dissociate to account for acids and bases.
Lewis theory
Physicochemist Gilbert Lewis began studying a new definition of acids and bases in 1923, the same year that Brönsted and Lowry offered their own theory on these substances.
This proposal, which was published in 1938, had the advantage of removing the hydrogen (or proton) requirement from the definition.
He himself had said, in relation to the theory of his predecessors, that "restricting the definition of acids to substances containing hydrogen was as limiting as restricting oxidizing agents to those containing oxygen."
Broadly speaking, this theory defines bases as the substances that can donate a pair of electrons, and acids as those that can receive this pair.
More precisely, it states that a Lewis base is one that has a pair of electrons, that is not bound to its nucleus and can be donated, and that Lewis acid is one that can accept a free pair of electrons. However, the definition of Lewis acids is loose and depends on other characteristics.
An example is the reaction between trimethylborane (Me 3 B) - which acts as a Lewis acid because it has the ability to accept a pair of electrons - and ammonia (NH 3), which can donate its free electron pair.
Me 3 B +: NH 3 → Me 3 B: NH 3
A great advantage of the Lewis theory is the way in which it complements the model of redox reactions: the theory suggests that acids react with bases to share an electron pair, without changing the oxidation numbers of any of their atoms.
Another advantage of this theory is that it allows us to explain the behavior of molecules such as boron trifluoride (BF 3) and silicon tetrafluoride (SiF 4), which do not have the presence of H + or OH - ions, as required by the previous theories.
References
- Britannica, E. d. (sf). Encyclopedia Britannica. Retrieved from britannica.com
- Brønsted – Lowry acid – base theory. (sf). Wikipedia. Retrieved from en.wikipedia.org
- Clark, J. (2002). Theories of acids and bases. Retrieved from chemguide.co.uk