Theory, of Arrhenius

The test of the validity of the theory of Arrhenius is not therefore to be found in the agreement between the values of i obtained from measurements of any properties of solutions which are conditioned by the osmotic pressure it is in quite another field—that of electrochemistry—that a comparison of known relations with the deductions from the theory may be instituted. 

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 theory of Arrhenius has proved its value with some minor supplements in explaining the behaviour of weak electrolytes, it failed, however, completely when applied to strong electrolytes. Success in the case of weak electrolytes can be explained by the fact that because of the comparatively small number of ions and because of the considerable distances between them, there is no substantial difference between ions and undissociated molecules as far as their individual behaviour is concerned. We may, therefore, assume that all these particles regardless of their nature play an equal part in the determination of the thermodynamic properties of the solution. The degree of dissociation has in this case a concrete meaning and individual weak electrolytes can be differentiated by their characteristic dissociation constants. 

In the case of strong electrolytes with a = 1 the last, mentioned equation changes to the form of Onsager s equation, while for very weak electrolytes or substantially diluted solutions (o, si 0, /A 1) it changes to the classical formula of Arrhenius. If we consider the problem of electrolytes from the point of view of the equation (III-31) we can see that there is no sharp dividing line between strong and weak electrolytes. It can equally be seen that the Debye-Hiickel-Onsager theory does not replace the theory of Arrhenius but merely corrects and suitably supplements it. 

The history of the ion-pair concept begins in the 1880s [50] with the theory of Arrhenius that described the electrolytic dissociation in solution as a function of electrolyte concentration and nature. This theory was completely eclipsed by the success of the D-H treatment of electrolyte activity in the 1920s [33] that also paved... 

Hydrogen was recognized as the essential element in acid by Davy after his work on hydrohalic acid. Theories of acid and base have played an important role ever since. The electrolytic dissociation theory of Arrhenius and Ostwald, the introduction of the pH scale for hydrogen-ion concentrations by Sorensen, the theory of acid-base titration and the use of indicators, and Bronsted s concept of acid and conjugate base as proton donors and acceptors are other landmarks in the recognition of hydrogen as an acid. 

Subsequent research has entirely confirmed the theory of Arrhenius. The chemical and electrochemical behaviour of solutions is closely connected with their ionic dissociation, and would be quite inexplicable without this theory. Rarely in the history of science has an idea proved so fruitful and suggestive, and led to the discovery of so many hitherto unsuspected relationships as this hypothesis of Arrhenius. 

The further development of the theory of Arrhenius is not within the scope of this book. Let us conclude by pointing out, in emphatic contradiction of the views of certain authors to tho contrary, that this theory has also rendered excellent service in explaining the behaviour of non-aqueous solutions. This has-been proved by the recent work of AValden on this subject. 

Other phenomena, such as the influence of electrolytes on the solubility of slightly soluble salts, the so-called neutral salt effect in catalytic phenomena, and the light absorption by strong electrolytes, are not elucidated by the classical theory of Arrhenius. 

Most dissociation constants recorded in the literature have been calculated on the basis of the old dissociation theory of Arrhenius. It is therefore of great practical interest to evaluate these data critically and to see how the thermodynamic dissociation constants may be calculated from them. Only the three most important methods for determining dissociation constants will be discussed. Sufficient information has been given... 

Composition of Salts. A famous example of school-made misconceptions of our students arises from the Dissociation Theory of Arrhenius. In 1884, he postulated that salt molecules are found in solid salts as the smallest particles and decompose into ions by dissolving in water . Later, with the concept of electrons, the misconception that atoms of salt molecules form ions through electron exchange was born. Today, experts recognize that there are no salt molecules, that ions exist all the time - even in the solid salt. By dissolving the solid salt, water molecules surround the ions, and hydrated ions are not connected, they move freely in the salt solution. 

The "Classical Ionic Theory of Arrhenius. Arrhenius,14 in 1887, proposed a method for computing the degree of dissociation, a, which has had an enormous influence on the progress of physical chemistry and upon related sciences. According to his method the degree of dissociation is found by means of the simple formula,... 

According to the theory of Arrhenius, the variations of A with concentration are due to shifts in equilibrium between undissociated and dissociated species. This idea was expressed quantitatively by the Russian-German physical chemist Friedrich Wilhelm Ostwald (1853-1932) in terms of a dilution law. Consider an electrolyte AB which exists in solution partly as the undissociated species AB and partly as the ions A" and B""... 

The theory of ionization in solution showed that hydrogen ions are the essential constituents of acids. However, the part played by the solvent in the ionization, or in other words, the probable mechanism of reaction according to which the ions are produced, is not shown in the theory of Arrhenius. It is true that the relative degrees of ionization of substances in different solvents were shown to be paralleled by certain other properties of the solvent, such as the dielectric constant, and that combination of the solvent... 

He employed a range of adds and he correlated the affinity (reactivity) of an acid with its catalytic power. He was therefore in a good position to appreciate Arrhenius s concept of electrolytic dissociation when the latter sent him a copy of his doctoral thesis in 1884. In 1887 Ostwald moved to Leipzig as professor of physical chemistry. For the remainder of his career he championed the ionic theory of Arrhenius against much opposition. He provided additional evidence for the theory, and he developed the theory of add-base indicators. He resigned from Leipzig in 1905, and in his retirement he worked on the theory of colours, as well as espousing many humanistic, educational and cultural causes. 

While this relatively simple and accurate relationship bears the name of Arrhenius, it was the work of fellow Nobel Prize winner Jacobus van t Hoff, a Dutch physical and organic chemist, that gave physical justification for the theory and the mathematical relationship that we use today. During his work with the theories of Arrhenius, van t Hoff made an important observation for simple reactions, the reaction rate doubles with an approximate 10 °C increase in temperature [1]. This observation can be mathematically applied to the Arrhenius equation to give a relationship that describes change in the system as a function of tanperature and time ... 

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