Theory, Arrhenius Lewis

The pH of a solution with a hydrogen ion concentration of 0.003 M is between what two whole numbers (Table 15.4) Which is more acidic, tomato juice or blood (Table 15.5) Use the three acid-base theories (Arrhenius, Brpnsted-Lowiy, and Lewis) to define an acid and a base. 

In the previous sections, you learned about three acid-base theories Arrhenius, Bronsted-Lowry, and Lewis. The Bronsted-Lowry theory is especially useful for describing acid-base reactions that take place in aqueous solutions. This section will use the Bronsted-Lowry description to explore reactions between acids and bases. 

A note on good practice The entities that are regarded as acids and bases are different in each theory. In the Lewis theory, the proton is an acid in the Bronsted theory, the species that supplies the proton is the acid. In both the Lewis and Bronsted theories, the species that accepts a proton is a base in the Arrhenius theory, the species that supplies the proton acceptor is the base (Fig. 10.61. 

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. 

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. 

A) In addition to the more modem Bronsted and Lewis theories, it is important not to forget the classic Arrhenius theory in its modern form, the so-called solvents theory, where it can be applied, i.e., with solvents that undergo self-dissociation in this form it was originally formulated in 1949 by Jander3 in Germany and is illustrated by the following reaction equations ... 

From our previous treatment of the Arrhenius, Bransted and Lewis acid-base theories, the importance of the choice between the divergent solvent types clearly appeared if we now confine ourselves to solvents to which the proton theory in general is applicable, this leads to a classification of eight classes as already proposed by Bronsted35,36 (Table 4.3). 

Many organic and biological reactions are acid-base reactions that do not lit within the Arrhenius or Bronsted-Lowry theories. Experienced chemists find the Lewis theory to be very usefol because so many other chemical reactions are covered by it. The less experienced sometimes find the theory less useful, but as their knowledge expands so does its utility. 

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. 

Skill 10.1 Analyzing acids and bases according to acid-base theories (i.e., Arrhenius, Bronsted-Lowry, Lewis)... 

Carbon dioxide reacts with the hydroxide ion to produce the bicarbonate anion. Write the Lewis dot structures for each reactant and product. Label each as a Bronsted acid or base. Explain the reaction using the Bronsted theory. Why would the Arrhenius theory provide an inadequate description of this reaction ... 

The theory of acids and bases postulated by Arrhenius, Br0nsted and Lewis is presently generalized in the Pearson theory of hard and soft acids and bases [26-31]. 

The original Arrhenius definition of a base has been extended by the Lowry-Bronsted theory and by the Lewis theory. See acid. 

A treatment on the basis of complete ionisation and electrical interionic action given by S. R. Milner was much simplified by P. Debye and E. Hiickel. G. N. Lewis proposed an arbitrary function of the concentration called the fugacity /, and an activity <2, such that the laws of ideal solutions are obeyed if activities are substituted for concentrations,/c = a/c being called the activity coefficient. The theory of Debye and Hiickel shows that the activity coefficient is a function of the square root of the concentration of the completely ionised electrolyte multiplied by a factor w depending on the valencies of the ions (an effect which has no place in Arrhenius s theory) /c = i - Awc >, where is a function of temperature and the properties of the pure solvent. In place of Arrhenius s a, the degree of dissociation, a conductivity coefficient / =A/Ao (or A/Aoo), which is numerically different from/c is found from the equation where a is a constant depending on the temperature... 

Arrhenius himself was recalcitrant on one point he throughout insisted that his interpretation of conductivity in terms of a dissociation equilibrium applies to strong electrolytes as well as to weak ones for an account of this controversy see ref. 13. Suggestions that strong electrolytes are completely dissociated and that ionic interactions must be invoked to explain the conductivities of their solutions were made by G.N. Lewis (14-16), Niels Bjerrum (17, 18), W. Sutherland (14) and S.R. Milner (20-24). Bjerrum provided spectroscopic evidence for this point of view, and Sutherland and Milner did important work on the theory of ionic interactions. Finally, in 1923, Debye and Hllckel (25, 26) developed their comprehensive treatment of strong electrolytes. Arrhenius, however, rejected these ideas and for the most part refused to discuss them. At a meeting of the Faraday Society held in January, 1919, Arrhenius commented that Bjerrum s idea "seems not to agree very well with experiment" (27), and he maintained this position until his death in 1927. 

Our emphasis throughout this chapter has been on water as the solvent and on the proton as the source of addic properties. In such cases we find the Bronsted-Lowry definition of adds and bases to be the most useful. In fact, when we speak of a substance as bdng addic or basic, we are usually thinking of aqueous solutions and using these terms in the Arrhenius or Bronsted-Lowry sense. The advantage of the Lewis theory is that it allows us to treat a wider variety of... 

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. 

Lewis proposed his stiU broader and more useful definition of acids and bases in the late 1920s and early 1930s. Classifying acids as electron-pair acceptors and bases as electron-pair donors, he thereby liberated acid—base theory entirely from its former dependence on the presence of hydrogen. The advantage of the Lewis definition is that a larger number of reactions can be classified as acid-base than under either the Arrhenius or Bronsted-Lowry definitions. The classic example used to demonstrate the more general nature of the Lewis definition is the gas-phase reaction between boron trifluoride and ammonia, as represented in Equation (4.1) ... 

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