Dissociation: electrolytic

A simple compound can decompose not only by dissociating into atoms or molecules, thus into the elements, but also dissociation into ions must be considered. This phenomenon is already more complicated to begin with, because this is practically only of interest for aqueous solutions. On the one hand there is the dissociation energy, and on the other hand there is the energy of interaction of the ions with the molecules of the solvent (hydration energy p. 98). 

From this consideration it follows that strictly speaking no direct relation need exist between the nature of the bonding in the solid state and the dissociation in aqueous solution. However, salts, considered as the group of the compounds of bases and acids, are usually built up from ions also in the solid state 

The cadmium and lead halides are also incompletely dissociated, that is to say, that here in the first place ions [CdCl]+ and [PbCl]+ are formed [dissociation constants are small, of the order of 0.01-0.03 for Cdl+ this is even only 0.004]. 

From the conductivity in solutions which are not extremely dilute, it appears that in numerous other cases also ion pairs are formed (Bjerrum), that is to say, combinations of ions, each still with its hydration sheath and which thus do not correspond with molecules. It is especially the higher valency ions from the nature of things which exhibit this phenomenon thus this pair formation occurs, for example, in the alkali sulphates, alkaline earth nitrates and barium hydroxide through the formation of [MS04], [MN03]+ and [Ba(OH)] ions, furthermore [Ce4+(OH)-]3+, [Fe3+(OH)-]2+. 

The zinc halides behave in dilute solutions entirely as strong electrolytes, but in more concentrated solutions this is otherwise probably it is here a question of the direct formation of higher complexes, for example, [ZnX4]2-, the dissociation of which depends very much more on the concentration than for [ZnX]+. 

Of great importance for the development of solution theory were the studies of col-ligative solution properties, detected in the 1870s and 1880s by F. M. Raoult, J. H. van t Hoff, and others. These are properties that depend not on the chemical nature of solutes but on their concentration. Three such colligative properties exist  

Osmotic pressure of the solvent. In dilute ideal solutions the osmotic pressure n of the solution obeys the equation 

Fundamentals of Electrochemistry, Second Edition, By V. S. Bagotsky Copyright 2006 John Wiley Sons, Inc. 

Relative lowering of the solvent s vapor pressure. The equilibrium vapor pressure of the solvent over a dilute ideal solution p obeys the equation 

Elevation of boiling point and depression of freezing point of the solution. Within certain limits, the change in temperature of these phase transitions obeys the eqnation 

In general, the activity coeflicient of an electrolyte will be larger in a solvent of lower dielectric constant than in water. Transfer activity coefficients are likely to be large if the dielectric constant is small or if the ions are of high charge or small radius. 

FIGURE 4-2 Left, relation between the observed formation constant for the salt tetia n-butylammonium picrate and the dielectric constant of the solvent. From Inami, Bodenseh, and RamseyJ ) Right, relation between fraction dissociated and dielectric constant calculated from data in the left part of the figure. 

From (4-26), for a 1 1 electrolyte in water the Bjerrum critical distance is 3.6 X 10 cm. When it is realized that ions in water are normally highly solvated and that the sum of ionic crystal radii for typical anions and cations often approaches or exceeds 3.6 x 10 cm, it is reasonable to find that dissociation constants for ion pairs in water are large thus for sodium hydroxide the dissociation constant is about 5. On the other hand, for a 2 2 electrolyte in water the critical distance is 14.3 x 10 cm, and for a 1 1 electrolyte in ethanol, 11.5 x 10 cm. In these cases, even highly solvated ions can readily approach to the distance necessary to form an ion pair. For magnesium sulfate the dissociation constant in water is 6 x 10 , and for sodium sulfate, 0.2. 

In addition to ion-pair formation, other types of association reactions are important in nonaqueous solvents. Evidence for triple and quadruple ion formation in nonaqueous solvents is obtained from conductimetric, solvent extraction, calorimetric, or cryoscopic measurements. Self-association reactions, that is, equilibria such as 2HA (HA)2, have been reviewed and have been studied by dififer-ential-vapor-pressure techniques. Frequently the anion A obtained from an acid HA is poorly solvated stabilization then may occur by interaction (homoconjugation) with a second molecule of acid to give HAj . Homoconjugation can be studied by techniques such as spectroscopy or conductimetry. Thus, the homoconjugation 

Water is highly unusual in the extent of its interactions with solutes, but even minimal solvent-solute interactions can play a major role in the nature of chemical reactions. To calculate pH during acid-base titrations in a nonaqueous solvent, we must consider not only the equilibria discussed in Chapter 3 but also reactions discussed in Sections 4-2, 4-3, and 4-4. 

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This system of nomenclature has withstood the impact of later experimental discoveries and theoretical developments that have since the time of Guyton de Morveau and Lavoisier greatiy altered the character of chemical thought, eg, atomic theory (Dalton, 1802), the hydrogen theory of acids (Davy, 1809), the duahstic theory (Berzehus, 1811), polybasic acids (Liebig, 1834), Periodic Table (Mendeleev and Meyer, 1869), electrolytic dissociation theory (Arrhenius, 1887), and electronic theory and modem knowledge of molecular stmcture. 

V total number of moles from electrolyte dissociation ... 

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... 

Strictly speaking the hydrogen ion H+ exists in water as the hydroxonium ion H30 + (Section 2.4). The electrolytic dissociation of water should therefore be written ... 

Electrolytic cells 504 Electrolytic dissociation 19 Electrolytically generated reagents ... 

Two components in one phase, e.g.f the electrolytic dissociation of a salt in aqueous solution ... 

Amplitude of a process, 114. Andrew s diagram, 173 Anisotropic bodies, 193 Aphorism of Clausius, 83, 92 Arrhenius theory of electrolytic dissociation, 301 Aschistic process, 31, 36, 51 Atmosphere, 39 Atomic energy, 26 Availability, 65, 66 Available energy, 66, 77, 80, 98, 101... 

The limiting law dependence of a2 on nr and a better understanding of the nature of 72 can be obtained by considering the equilibrium that is present when the electrolyte dissociates. For HC1, this equilibrium is... 

It is perhaps desirable to point out that the bond type has no direct connection with ease of electrolytic dissociation in aqueous solution. Thus the nearly normal covalent molecule HI ionizes completely in water, whereas the largely ionic HF is only partially ionized. 

The above examples assume that the strong base KOH is completely dissociated in solution and that the concentration of OH ions was thus equal to that of the KOH. This assumption is valid for dilute solutions of strong bases or acids but not for weak bases or acids. Since weak electrolytes dissociate only slightly in solution, we must use the dissociation constant to calculate the concentration of [H" ] (or [OH ]) produced by a given molarity of a weak acid (or base) before calculating total [H" ] (or total [OH ]) and subsequendy pH. 

Number of cations, anions, and total ions v = Vj,+vJ) into which the given electrolyte dissociates (Chap. XIII). 

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. 

Thus, quantitative criteria that could be tested experimentally had now been formulated for the first time in the theory of electrolytic dissociation, in contrast to earlier theories. The good agreement between degrees of dissociation calculated from independent measurements of two different properties with Eqs. (7.5) and... 

Soon after inception of the theory of electrolytic dissociation, it was shown that two types of componnds exist that can dissociate upon dissolution in water (or other solvents) ... 

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. 

The theory of electrolytic dissociation was not immediately recognized universally, despite the fact that it could qualitatively and quantitatively explain certain fundamental properties of electrolyte solutions. For many scientists the reasons for spontaneous dissociation of stable compounds were obscure. Thus, an energy of about 770kJ/mol is required to break up the bonds in the lattice of NaCl, and about 430kJ/mol is required to split H l bonds during the formation of hydrochloric acid solution. Yet the energy of thermal motions in these compounds is not above lOkJ/mol. It was the weak point of Arrhenius s theory that this mismatch could not be explained. 

Between 1865 and 1887, Dmitri 1. Mendeleev developed the chemical theory of solutions. According to this theory, the dissolution process is the chemical interaction between the solutes and the solvent. Upon dissolution of salts, dissolved hydrates are formed in the aqueous solution which are analogous to the solid crystal hydrates. In 1889, Mendeleev criticized Arrhenius s theory of electrolytic dissociation. Arrhenius, in turn, did not accept the idea that hydrates exist in solutions. 

It was found in later work that it is precisely the idea of ionic hydration that is able to explain the physical nature of electrolytic dissociation. The energy of interaction between the solvent molecules and the ions that are formed is high enough to break up the lattices of ionophors or the chemical bonds in ionogens (for more details, see Section 7.2). The significance of ionic hydration for the dissociation of electrolytes had first been pointed out by Ivan A. Kablukov in 1891. 

According to modem views, the basic points of the theory of electrolytic dissociation are correct and were of exceptional importance for the development of solution theory. However, there are a number of defects. The quantitative relations of the theory are applicable only to dilute solutions of weak electrolytes (up to 10 to 10 M). Deviations are observed at higher concentrations the values of a calculated with Eqs. (7.5) and (7.6) do not coincide the dissociation constant calculated with Eq. (7.9) varies with concentration and so on. For strong electrolytes the quantitative relations of the theory are altogether inapplicable, even in extremely dilute solutions. 

Numerous measurements of the conductivity of aqueous solutions performed by the school of Friedrich Kohhansch (1840-1910) and the investigations of Jacobns van t Hoff (1852-1911 Nobel prize, 1901) on the osmotic pressure of solutions led the young Swedish physicist Svante August Arrhenius (1859-1927 Nobel prize, 1903) to establish in 1884 in his thesis the main ideas of his famous theory of electrolytic dissociation of acids, alkalis, and salts in solutions. Despite the sceptitism of some chemists, this theory was generally accepted toward the end of the centnry. 

The acidic character of acids depends on the availability ofhydrogen ions in their solution. An acid X3 is said to be stronger than another acid X2 if, in equimolar solutions, X3 provides more hydrogen ions than does X2. This will be possible provided that the degree of dissociation of X3 is greater than that of X2. Based on the Arrhenius theory of electrolytic dissociation, solutions may be classified in the manner shown in Figure 6.1. If the ionization of an acid is almost complete in water, the acid is said to be a strong acid, but if the... 

The elucidation of the electrical behavior of electrolytes owes much to Arrhenius, who was the originator of the theory of electrolytic dissociation, generally, known as the ionic theory. 

A state of dynamic equilibrium exists between the ionized and the non-ionized molecules (XY (non-ionized molecule) X+ (cation) + Y (anion)). The process of electrolytic dissociation is reversible. 

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