Acid-base concepts Arrhenius theory

This theory was a milestone in the development of acid-base concepts it was the first to define acids and bases in terms other than that of a reaction between them and the first to give quantitative descriptions. However, the theory of Arrhenius is far more narrow than both its predecessors and its successors and, indeed, it is the most restrictive of all acid-base theories. 

The first scientific definition in this new field was described by a Nobel laureate, Swedish chemist Svante Arrhenius (1859-1927). His theory, which is now called the Arrhenius acid-base concept, states that acids are substances that produce hydrogen ions in an aqueous solution. Hydrogen ions (H+) do not exist in solutions because they are always attached to at least one water molecule- often written as hydroxonium ion (H3O+). More detailed studies show that H3O+ ions are not very common, either, and hydrogen ions are often attached to more than one water molecules. But this is only a matter of notation and different chemists use H+, H3O+, Hj02 or H,03+ to mean essentially the same thing. 

Ideas about adds and bases (or alkalis) date back to ancient times. The word acid is derived from the Latin acidus (sour). Alkali (base) comes from the Arabic al-qali, referring to the ashes of certain plants from which alkaline substances can be extracted. The acid-base concept is a major theme in the history of chemistry. In this section, we emphasize the view proposed by Svante Arrhenius in 1884 but also introduce a more modern theory proposed in 1923 by Thomas Lowry and by Johannes Bronsted. 

The Arrhenius definition of an acid and a base attributed acidity to the presence of H" (aq), and alkalinity to OH (aq). Br0nsted-Lowry theory generalizes the acid-base concept by focusing on proton transfer, rather than on particular aqueous ions. Here, we discuss an attempt to generalize it further by focusing on the changes in electronic structure that occur when acid-base reactions take place, ideas introduced by G. N. Lewis. 

Hydrogen was recognized as the essential element in acids by H. Davy after his work on the hydrohalic acids, and theories of acids and bases have played an important role ever since. The electrolytic dissociation theory of S. A. Arrhenius and W. Ostwald in the 1880s, the introduction of the pH scale for hydrogen-ion concentrations by S. P. L. Sprensen in 1909, the theory of acid-base titrations and indicators, and J. N. Brdnsted s fruitful concept of acids and conjugate bases as proton donors and acceptors (1923) are other land marks (see p. 48). The di.scovery of ortho- and para-hydrogen in 1924, closely followed by the discovery of heavy hydrogen (deuterium) and... 

It was G. N. Lewis who extended the definitions of acids and bases still further, the underlying concept being derived from the electronic theory of valence. It provided a much broader definition of acids and bases than that provided by the Lowry-Bronsted concept, as it furnished explanations not in terms of ionic reactions but in terms of bond formation. According to this theory, an acid is any species that is capable of accepting a pair of electrons to establish a coordinate bond, whilst a base is any species capable of donating a pair of electrons to form such a coordinate bond. A Lewis acid is an electron pair acceptor, while a Lewis base is an electron pair donor. These definitions of acids and bases fit the Lowry-Bronsted and Arrhenius theories, and cover many other substances which could not be classified as acids or bases in terms of proton transfer. 

The Arrhenius concept was of basic importance because it permitted quantitative treatment of a number of acid-base processes in aqueous solutions, i.e. the behaviour of acids, bases, their salts and mixtures of these substances in aqueous solutions. Nonetheless, when more experimental material was collected, particularly on reaction rates of acid-base catalysed processes, an increasing number of facts was found that was not clearly interpretable on the basis of the Arrhenius theory (e.g. in anhydrous acetone NH3 reacts with acids in the absence of OH- and without the formation of water). It gradually became clear that a more general theory was needed. Such a theory was developed in 1923 by J. N. Br0nsted and, independently, by T. M. Lowry. 

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 ... 

Our goal in this chapter is to help you understand the equilibrium systems involving acids and bases. If you don t recall the Arrhenius acid-base theory, refer to Chapter 4 on Aqueous Solutions. You will learn a couple of other acid-base theories, the concept of pH, and will apply those basic equilibrium techniques we covered in Chapter 14 to acid-base systems. In addition, you will need to be familiar with the log and 10 functions of your calculator. And, as usual, in order to do well you must Practice, Practice, Practice. 

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. 

Arrhenius in 1887 was the first person to give a definition of an acid and a base. According to him, an acid is one that gives rise to excess of in aqueous solution, whereas a base gives rise to excess of OH in solution. This was modified by Bronsted-Lowry in 1923 such that a proton donor was defined as an acid and a proton acceptor as a base. They also introduced the familiar concept of the conjugate acid-base pair. The final refinement to the acid-base theory was completed by Lewis in 1923, who extended the concept that acid is an acceptor of electron pairs while base is a donor of electron pairs. 

Sumfleth [6] also states, that the idea of proton transfer may be learned by students but cannot be applied in a new context. Sumfleth and Geisler [7] show that students accept the Broensted definition, but bases are interpreted mostly based on the Arrhenius idea. Therefore, the knowledge about Broensteds concept cannot be transferred to new contexts. Sumfleth states that most students cannot really apply acid-base theories, especially at the advanced levels. This is also evident for students who have chosen chemistry as their major . 

The Arrhenius model is still widely used, and substances that are called acid are usually acids in the Arrhenius sense. Since citric acid, for example, produces hydrogen ions in solution, it is therefore an acid. Inddenlally, it also tastes sour and is edible at the same time. The same is true for tartaric add. The most commonly consumed form of tartaric add is wine, but it is mixed with alcohol there, which may have a significant health efled. In the Arrhenius concept, a substance can be called acid or base without further specifications (bases decrease the concentration of hydrogen ions). In the remaining two theories, strictly speaking, a substaiKe cannot be called acid or base. All that can be said is that a certain substance behaves as an add (or base) in a certain chemical reaction with another reactarrt. 

The Arrhenius concept is important in that it has provided us with the first mechanistic approach to acid - base behaviour and has been instrumental for the development of more sophisticated theories. TTiere are, however, two major shortcomings in the Arrhenius model. 

This theory defines an acid as any compound that yields protons (H Ions) and a base as any compound that combines with a proton. In other words, acids are proton donors and bases are proton acceptors. It should be noted that as far as acids are concerned, Arrhenius and Bronsted - Lowry theories are similar in both cases acids give off ions. However, the concept of a base is much broader in the Bronsted theory, hydrorqrl ion being just one of the possible bases. Cited below are a few examples which wlU illustrate the point much better. 

The Arrhenius concept was the first successful theory of acids and bases. Then in 1923, Br0nsted and Lowry characterized acid-base reactions as proton-transfer reactions. According to the Br0nsted—Lowry concept, an acid is a proton donor and a base is a proton acceptor. The Lewis concept is even more general than the Br0nsted-Lowry concept. A Lewis acid is an electron-pair acceptor and a Lewis base is an electron-pair donor. Reactions of acidic and basic oxides and the formation of complex ions, as well as proton-transfer reactions, can be described in terms of the Lewis concept. 

The first substantial constitutive concept of acid and bases came only in 1887 when Arrhenius applied the theory of electrolytic dissociation to acids and bases. An acid was defined as a substance that dissociated to hydrogen ions and anions in water (Day Selbin, 1969). For the first time, a base was defined in terms other than that of an antiacid and was regarded as a substance that dissociated in water into hydroxyl ions and cations. The reaction between an acid and a base was simply the combination of hydrogen and hydroxyl ions to form water. 

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. 

The theory of Bronsted (1923) and Lowry (1923a, b) is of more general applicability to AB cements. Their definition of an acid as a substance that gives up a proton differs little from that of Arrhenius. However, the same is not true of their definition of a base as a substance capable of accepting protons which is far wider than that of Arrhenius, which is limited to hydroxides yielding hydroxide ions in aqueous solution. These concepts of Bronsted and Lowry can be defined by the simple equation (Finston Rychtman, 1982) ... 

The theory of electrolytic dissociation also provided the possibility for a transparent definition of the concept of acids and bases. According to the concepts of Arrhenius, an acid is a substance which upon dissociation forms hydrogen ions, and a base is a substance that forms hydroxyl ions. Later, these concepts were extended. 

There are certain limitations of the Arrhenius concept of acids and bases. Acids and bases have been described in terms of their aqueous solutions and not in terms of the entities themselves. The theory is thus applicable exclusively to aqueous solutions. An entity such as HC1 is accounted as an acid only when it is dissolved in water if dissolved in... 

Any text on acids and bases would not be deemed complete if mention were not made of the extended definition of acids and bases that is embodied in the Lowry-Bronsted theory. The theory basically proposed a more general definition of acids and bases to overpower the limitations of the theory arising from the Arrhenius concept. 

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... 

An understanding of the chemical mechanisms that give rise to the properties of acids evolved from a number of different theories of the nature of acids. Arrhenius proposed the Arrhenius Concept of Bases that an acid is a... 

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