Acid-base solvent theory definition

Since Arrhenius, definitions have extended the scope of what we mean by acids and bases. These theories include the proton transfer definition of Bronsted-Lowry (Bronsted, 1923 Lowry, 1923a,b), the solvent system concept (Day Selbin, 1969), the Lux-Flood theory for oxide melts, the electron pair donor and acceptor definition of Lewis (1923, 1938) and the broad theory of Usanovich (1939). These theories are described in more detail below. 

Almost all of the reactions that the practicing inotganic chemist observes in the laboratory take place in solution. Although water is the best-known solvent, it is not the only one of importance to the chemist. The organic chemist often uses nonpolar solvents sud) as carbon tetrachloride and benzene to dissolve nonpolar compounds. These are also of interest to Ihe inoiganic chemist and, in addition, polar solvents such as liquid ammonia, sulfuric acid, glacial acetic acid, sulfur dioxide, and various nonmctal halides have been studied extensively. The study of solution chemistry is intimately connected with acid-base theory, and the separation of this material into a separate chapter is merely a matter of convenience. For example, nonaqueous solvents are often interpreted in terms of the solvent system concept, the formation of solvates involve acid-base interactions, and even redox reactions may be included within the (Jsanovich definition of acid-base reactions. 

These ideas were rather limiting since they only applied to aqueous solutions. There were situations where acid-base reactions were taking place in solvents other than water, or even in no solvent at all. This problem was addressed in 1923 by the Danish chemist Johannes Bronsted (1879-1947) and the English chemist Thomas Lowry (1874-1936) when they independentiy proposed a more general definition of acids and bases, and the study of acids and bases took a great step forward. This theory became known as the Bronsted-Lowry theory of acids and bases. 

Arrhenius acid-base theory - Arrhenius developed the theory of the electrolytic dissociation (1883-1887). According to him, an acid is a substance which delivers hydrogen ions to the solution. A base is a substance which delivers hydroxide ions to the solution. Accordingly, the neutralization reaction of an acid with a base is the formation of water and a salt. It is a so-called symmetrical definition because both, acids and bases must fulfill a constitutional criterion (presence of hydrogen or hydroxide) and a functional criterion (to deliver hydrogen ions or hydroxide ions). The theory could explain all of the known acids at that time and most of the bases, however, it could not explain the alkaline properties of substances like ammonia and it did not include the role of the solvent. -> Sorensen (1909) introduced the -> pH concept. 

The definition of acids can extend according to the solvent theory as follows (this definition is symmetric concerning the acid-base relationship in contrast to the Br0tisted theory) acids/bases increase/decrease or increase/decrease OH . 

The Lewis definitions of acid-base interactions are now over a half a century old. Nevertheless they are always useful and have broadened their meaning and applications, covering concepts such as bond-formation, central atom-ligand interactions, electrophilic-nucleophilic reagents, cationic-anionic reagents, charge transfer complex formation, donor-acceptor reactions, etc. In 1923 Lewis reviewed and extensively elaborated the theory of the electron-pair bond, which he had first proposed in 1916. In this small volume which had since become a classic, Lewis independently proposed both the proton and generalized solvent-system definitions of acids and bases. He wrote ... 

As G. N. Lewis said, We frequently define an acid or a base as a substance whose aqueous solution gives, respectively, a higher concentration of hydrogen ion or hydroxide ion than that furnished by pure water. This is a very one-sided definition. In 1923, Bronsted and Lowry expanded the definitions of acids and bases to include species that do not involve solvent participation. According to the Bronsted-Lowry definition, an acid is any proton donor, whereas a base is any proton acceptor. This broader definition expanded acid-base theory to include gaseous species, such as HCI (g) and NH3 (g). It also allowed for the inclusion of acid-base reactions occurring in nonionizing solvents, such as benzene, as shown by Equation (14.6) ... 

According to the Arrhenius theory of acids and bases, the acidic species in water is the solvated proton (which we write as H30+). This shows that the acidic species is the cation characteristic of the solvent. In water, the basic species is the anion characteristic of the solvent, OH-. By extending the Arrhenius definitions of acid and base to liquid ammonia, it becomes apparent from Eq. (10.3) that the acidic species is NH4+ and the basic species is Nl I,. It is apparent that any substance that leads to an increase in the concentration of NH4+ is an acid in liquid ammonia. A substance that leads to an increase in concentration of NH2- is a base in liquid ammonia. For other solvents, autoionization (if it occurs) leads to different ions, but in each case presumed ionization leads to a cation and an anion. Generalization of the nature of the acidic and basic species leads to the idea that in a solvent, the cation characteristic of the solvent is the acidic species and the anion characteristic of the solvent is the basic species. This is known as the solvent concept. Neutralization can be considered as the reaction of the cation and anion from the solvent. For example, the cation and anion react to produce unionized solvent ... 

The first clear definition of acidity can be attributed to Arrhenius, who between 1880 and 1890 elaborated the theory of ionic dissociation in water to explain the variation in strength of different acids.3 Based on electrolytic experiments such as conductance measurements, he defined acids as substances that dissociate in water and yield the hydrogen ion whereas bases dissociate to yield hydroxide ions. In 1923, J. N. Brpnsted generalized this concept to other solvents.4 He defined an acid as a species that can donate a proton and defined a base as a species that can accept it. This... 

Seeing that the Brpnsted definitions of acids and bases are not related to a specific solvent, this theory can readily explain the reaction shown in Eq. (5.4). In that case, HCI donates a proton to NH3, resulting in the formation of the ionic salt NH4C1. Therefore, HCI is an acid. Because NH3 accepts a proton, it is acting as a base. Likewise, the Brpnsted theory is applicable to many reactions in which there is a solvent other than water, which makes the Brpnsted theory much more generally applicable than the Arrhenius theory. 

See pH. When dealing with chemical reactions in solvents other than water, it is sometimes convenient to define an acid as a substance that ionizes to give the positive ion of the solvent. The common definitions of acid have been extended as more detailed studies of chemical reactions have been made. The Lowry-Brpnsted definition of an acid as a substance that can give up a proton is more useful in connection with an understanding of bases (see base). Perhaps the most significant contribution to the theory of acids was the electron-pair concept introduced by G. N. Lewis around 1915. 

The more general working definition of acids and bases we have been using is due to Franklin, who in 1905 developed a theory in which the solvent plays a central role. According to this view, an acid is a solute that gives rise to a cation characteristic of the solvent, and a base is a solute that jdelds a dissolved ion which is also characteristic of the solvent. 

A more quantitative prediction of activity coefficients can be done for the simplest cases [18]. However, for most electrolytes, beyond salt concentrations of 0.1 M, predictions are a tedious task and often still impossible, although numerous attempts have been made over the past decades [19-21]. This is true all the more when more than one salt is involved, as it is usually the case for practical applications. Ternary salt systems or even multicomponent systems with several salts, other solutes, and solvents are still out of the scope of present theory, at least, when true predictions without adjusted parameters are required. Only data fittings are possible with plausible models and with a certain number of adjustable parameters that do not always have a real physical sense [1, 5, 22-27]. It is also very difficult to calculate the activity coefficients of an electrolyte in the presence of other electrolytes and solutes. Even the definition is difficult, because electrolyte usually dissociate, so that extrathermodynamical ion activity coefficients must be defined. The problem is even more complex when salts are only partially dissociated or when complex equilibriums of gases, solutes, and salts are involved, for example, in the case of CO2 with acids and bases [28, 29]. 

The theory of solvent systems conforms to the experimental fact that there are many other substances besides those containing hydrogen which exhibit typical acid properties. But it makes the definitions of acid and base as rigidly dependent upon the solvent as does the water theory. 

However, when there seems to be no doubt that a primary acid (e.g., boron trichloride in the absence of a solvent) is involved, there is no difficulty in using the term neutralization. At any rate, the electronic theory of acids and bases gives definite meaning to the word. The type equation of the Br0nsted theory... 

---