The Evolution of Acid–Base Theories: From Early Concepts to the Brønsted–Lowry Proton Model

PART 3.
TOPIC 4. Proton Theory of Acids and Bases.
Theories of acids and bases are a set of fundamental physicochemical concepts that describe the nature and properties of acids and bases. All of them provide definitions of acids and bases—two classes of substances that react with each other. The task of the theory is to predict the products of the reaction between an acid and a base and the possibility of its occurrence, for which quantitative characteristics of the strength of acids and bases are used. The differences between the theories lie in the definitions of acids and bases, their strength characteristics, and, consequently, in the rules for predicting the products of their reactions. Each theory has its scope of applicability, and these scopes partially overlap.

Evolution of Ideas about Acid-Base Interactions.
Ideas about acid-base interactions are among the fundamental chemical concepts. The terms "acid" and "base" were formed in the 17th century, but their content has been repeatedly reviewed and refined. For instance, R. Boyle believed that acids were substances whose atoms had sharp protrusions (and thus a sharp taste), while bases had pores (and astringent taste). According to his theory, the neutralization reaction was explained by the acid's protrusions entering the pores of the base.
An image-based theory of acids and bases was proposed by N. Lémery. In his "Course of Chemistry" (1675), he attempted to explain the physical and chemical properties of substances in terms of their shape and structure. According to Lémery, acids have sharp spines on their surface that cause prickling sensations on the skin. Bases, which he called alkalis, consist of porous bodies. The "spines" of acids penetrate into the "pores," breaking or dulling, and acids are transformed into neutral salts[1].
Scientific understanding of the nature of acids and bases began to take shape at the end of the 18th century. In the works of A. Lavoisier, the acidic properties were linked to the presence of oxygen atoms in a substance. The well-known mineral and organic acids of that time did indeed contain oxygen. This hypothesis quickly proved to be untenable, as works by H. Davy and J. Gay-Lussac showed that there were many acids that did not contain oxygen (for example, hydrogen halides, cyanide acids), while many oxygen-containing compounds did not exhibit acidic properties.
Since the beginning of the 19th century, substances capable of reacting with metals to release hydrogen were considered acids (J. Liebig, 1839). Around the same time, J. Berzelius introduced an idea that explained the acid-base properties of substances based on their electrical "dualistic" nature. Thus, he classified electronegative oxides of nonmetals and some metals (such as chromium, manganese, etc.) as acids, while electropositive oxides of metals were considered bases. Therefore, acidity or basicity, according to Berzelius, is seen as a functional, not an absolute property of a compound. Berzelius was the first to attempt to quantitatively assess and predict the strength of acids and bases[2].
With the emergence of the theory of electrolytic dissociation by S. Arrhenius (1887), it became possible to describe acid-base properties based on the products of electrolyte ionization. Thanks to the work of W. Ostwald, the theory was extended to weak electrolytes.
In the early 20th century, American chemists G. Cady, E. Franklin, and C. Kraus developed the theory of solvosystems, which extended the principles of the Arrhenius-Ostwald theory to all solvents capable of self-dissociation.
Modern acid-base theories are based on the concepts of J. Brønsted and G. Lewis. There have been successful attempts to create generalized theories (M. Usanovich, 1939), but they have not found wide application[3].

Liebig's Hydrogen Theory
Definitions. An acid is a substance capable of reacting with a metal to release hydrogen. The concept of "base" is absent in this theory.
Reaction Products. When an acid reacts with a metal, a salt and hydrogen are produced.
Examples. Acid — HCl:
Reaction: 2HCl + Zn = ZnCl2 + H2↑
Reaction Conditions. Metals that are positioned left of hydrogen in the activity series react with strong acids. The weaker the acid, the more active the metal required for the reaction.
Quantitative Characteristics. Since this theory is rarely used, quantitative characteristics of acid strength (and, consequently, predictions of reaction direction) are not developed within this framework.
Area of Applicability. Prediction of interactions between hydrogen-containing substances and metals in any solvent.
Specific Features. According to this theory, ethyl alcohol and ammonia are weak acids, as they are capable of reacting with alkali metals:
2C2H5OH + 2Na = 2C2H5ONa + H2↑
2NH3 + 2Na = 2NaNH2 + H2↑

Arrhenius-Ostwald Electrolytic Dissociation Theory.
Electrolytic dissociation scheme of acetic acid in aqueous solution.
Main article: Electrolytic Dissociation Theory
Definitions. Acids are substances that form hydrated hydrogen cations H+ (or hydronium ions) and anions of the acidic residue in an aqueous solution.
Bases are substances that dissociate in an aqueous solution to form metal or ammonium cations and hydroxide anions OH−.
Salts are substances that dissociate to form metal or ammonium cations and anions of the acidic residue.
Reaction Products. In the reaction between an acid and a base (neutralization reaction), a salt and water are formed.
Examples. Acid — HCl (acidic residue Cl-):
HCl + H2O ↔ H3O+ + Cl-
Base — NaOH:
NaOH ↔ Na+ + OH-
Neutralization reaction (salt — NaCl):
HCl + NaOH = NaCl + H2O
Reaction Conditions. Strong acids react with strong bases. The weaker the acid, the stronger the base required for the reaction.
Quantitative Characteristics. The strength of an acid and base is characterized by their dissociation constants K.
For acid HA, K = [H+]·[A-]/[HA]
For base MeOH, K = [Me+]·[OH-]/[MeOH]
For a reaction between an acid and a base to occur, the product of their dissociation constants must be greater than 10−14 (the ion product of water).
Area of Applicability. This theory satisfactorily describes reactions between relatively strong acids and bases and the properties of their aqueous solutions. The concepts of dissociation degree and constant were used to classify electrolytes as strong and weak, and the concept of the hydrogen ion concentration was introduced, though applying this to alkaline media requires additional assumptions (the introduction of the ion product of water).
The theory can also be used to describe the hydrolysis of salts and the reactions between acids, bases, and salts. However, this requires cumbersome apparatus — the proton theory is much more convenient in such cases.
The Arrhenius-Ostwald theory is limited to aqueous solutions. It also does not explain the basic properties of ammonia, phosphine, and other compounds that do not contain hydroxyl groups.

Brønsted-Lowry Proton Theory.
Comparison of Lewis and Brønsted Acid-Base Interaction Models.
The proton theory of acids and bases was proposed in 1923 independently by Danish scientist J. Brønsted and English scientist T. Lowry. In this theory, the concepts of acids and bases are united into a single whole manifested in acid-base interaction: A + B + H+ (A is the acid, B is the base). According to this theory, acids are molecules or ions that can act as proton donors in a reaction, while bases are molecules or ions that accept protons (proton acceptors). Acids and bases are collectively known as protolytics.

Solvosystem Theory.
The solvosystem theory is an extension of the Arrhenius-Ostwald theory to other ionic (particularly protonic) solvents. It was proposed by American chemists G. Cady, E. Franklin, and C. Kraus.
Definitions. An ionic solvent is a solvent that self-dissociates into a cation and an anion. The cation is called the lione ion, and the anion is called the liate ion. A protonic solvent is a solvent capable of autoprotolysis, i.e., the transfer of an H+ ion from one molecule to another:
2HL ↔ H2L+ + L−
These solvents contain a sufficiently polar bond involving hydrogen and an unshared electron pair on some other nonmetal (most often nitrogen, oxygen, or fluorine).
Note: This definition implicitly incorporates the proton theory, as autoprotolysis is an acid-base reaction according to Brønsted-Lowry. It also implicitly incorporates the Lewis theory since it explains the reasons for the formation of lione ions.
The ion H2L+ is called the lione ion, and L– is called the liate ion.
Acids are substances that form the lione ion in a given solvent.
Bases are substances that form the liate ion in a given solvent.
Salts are substances that dissociate in a given solvent to form a cation and an anion, neither of which are a lione or liate ion.
Reaction Products. In the reaction between an acid and a base (neutralization reaction), salt and solvent are formed.