Incomplete electrolytic dissociation

An electrolyte in solution dissociates into two (in the case of NaCl) or three (in the case of CaCh) particles, and therefore the colligative effects of such solutions are multiplied by the number of dissociated ions formed per molecule. However, because of incomplete electrolyte dissociation and associations between the solute and solvent molecules, many solutions do not behave in the ideal case, and a 1-molal solution may give an osmotic pressure lower than theoretically expected. The osmotic activity coefficient is a factor used to correct for the deviation from the "ideal behavior of the system ... 

The model of electrical conductivity of molten salt mixtures based on incomplete electrolytic dissociation of components was proposed by DanCk (1989). The dissociation degree of a component is affected by the presence of second component. Consequently, the dissociation degree of both components in the system is not constant, but changes with composition, affecting the concentration of the conducting particles in the electrolyte. This effect is caused by interactions of components, given by the nature of the repulsive forces between ions, determining their actual coordination sphere. 

Complete and Incomplete Ionic Dissociation. Brownian Motion in Liquids. The Mechanism of Electrical Conduction. Electrolytic Conduction. The Structure of Ice and Water. The Mutual Potential Energy of Dipoles. Substitutional and Interstitial Solutions. Diffusion in Liquids. 

Complete and Incomplete Ionic Dissociation. In the foregoing chapter mention has been made of electrolytes that are completely dissociated in solution, and of weak electrolytes where free ions are accompanied by a certain proportion of neutral molecules. In the nineteenth century it was thought that aqueous solutions of even the strongest electrolytes contained a small proportion of neutral molecules. Opinion as to the relation between strong and weak electrolytes has passed through certain vicissitudes and we shall describe later how this problem has been resolved. 

In a weak electrolyte (e.g. an aqueous solution of acetic acid) the solute molecules AB are incompletely dissociated into ions and according to the familiar chemical equation... 

Incomplete Dissociation into Free Ions. As is well known, there are many substances which behave as a strong electrolyte when dissolved in one solvent, but as a weak electrolyte when dissolved in another solvent. In any solvent the Debye-IIiickel-Onsager theory predicts how the ions of a solute should behave in an applied electric field, if the solute is completely dissociated into free ions. When we wish to survey the electrical conductivity of those solutes which (in certain solvents) behave as weak electrolytes, we have to ask, in each case, the question posed in Sec. 20 in this solution is it true that, at any moment, every ion responds to the applied electric field in the way predicted by the Debye-Hiickel theory, or does a certain fraction of the solute fail to respond to the field in this way In cases where it is true that, at any moment, a certain fraction of the solute fails to contribute to the conductivity, we have to ask the further question is this failure due to the presence of short-range forces of attraction, or can it be due merely to the presence of strong electrostatic forces ... 

A study of the concentration dependence of the molar conductivity, carried out by a number of authors, primarily F. W. G. Kohlrausch and W. Ostwald, revealed that these dependences are of two types (see Fig. 2.5) and thus, apparently, there are two types of electrolytes. Those that are fully dissociated so that their molecules are not present in the solution are called strong electrolytes, while those that dissociate incompletely are weak electrolytes. Ions as well as molecules are present in solution of a weak electrolyte at finite dilution. However, this distinction is not very accurate as, at higher concentration, the strong electrolytes associate forming ion pairs (see Section 1.2.4). 

For strong electrolytes, the activity of molecules cannot be considered, as no molecules are present, and thus the concept of the dissociation constant loses its meaning. However, the experimentally determined values of the dissociation constant are finite and the values of the degree of dissociation differ from unity. This is not the result of incomplete dissociation, but is rather connected with non-ideal behaviour (Section 1.3) and with ion association occurring in these solutions (see Section 1.2.4). 

Interionic forces are relatively less important for weak electrolytes because the concentrations of ions are relatively rather low as a result of incomplete dissociation. Thus, in agreement with the classical (Arrhenius) theory of weak electrolytes, the concentration dependence of the molar conductivity can be attributed approximately to the dependence of the degree of dissociation a on the concentration. If the degree of dissociation... 

If a solute of the general formula AX (A is the chiral ion and X is an achiral ion) dissociates completely into ions once dissolved, then the solubility of the racemic conglomerate, SR, is equal to n%V2-SA (where SA is concentration of A in a solution saturated with AX ). If the solute is of the type AX, then 5 = V2-5a. The subscript n refers to the achiral ion and may be fractional, and so A2X must be represented by AXi/. If dissociation of AX is incomplete, SA lies between n i/2-SA and 2SA. For weakly dissociated electrolytes (such as carboxylic acids), SR is approximately 2SA. 

Raji Heyrovska [18] has developed a model based on incomplete dissociation, Bjermm s theory of ion-pair formation, and hydration numbers that she has found fits the data for NaCl solutions from infinite dilution to saturation, as well as several other strong electrolytes. She describes the use of activity coefficients and extensions of the Debye-Hiickel theory as best-fitting parameters rather than as explaining the significance of the observed results. ... 

A common sitnation is that the electrolyte is completely dissociated in the aqueons phase and incompletely, or hardly at all, in the organic phase of a ternary solvent extraction system (cf. Chapter 3), since solvents that are practically immiscible with water tend to have low valnes for their relative permittivities e. At low solnte concentrations, at which nearly ideal mixing is to be expected for the completely dissociated ions in the aqneons phase and the undissociated electrolyte in the organic phase (i.e., the activity coefficients in each phase are approximately nnity), the distribntion constant is given by... 

PK. A measurement of the complete ness of an incomplete chemical reaction. It is defined as the negative logarithm ito the base 101 of the equilibrium constant K for the reaction in question. The pA is most frequently used to express the extent of dissociation or the strength of weak acids, particularly fatty adds, amino adds, and also complex ions, or similar substances. The weaker an electrolyte, the larger its pA. Thus, at 25°C for sulfuric add (strong acid), pK is about -3,0 acetic acid (weak acid), pK = 4.76 bone acid (very weak acid), pA = 9.24. In a solution of a weak acid, if the concentration of undissociated acid is equal to the concentration of the anion of the acid, the pAr will be equal to the pH. 

Ostwald s dilution law — A weak - electrolyte is dissociated incompletely upon dissolution in a solvent. The chemical equilibrium of dissociation of a weak acid HA into protons and acid anions is described by ... 

Strong Electrolytes. Solutes of this type, such as HCl, are completely dissociated in ordinary dilute solutions. However, their colligative properties when interpreted in terms of ideal solutions appear to indicate that the dissociation is a little less than complete. This fact led Arrhenius to postulate that the dissociation of strong electrolytes is indeed incomplete. Subsequently this deviation in colligative behavior has been demonstrated to be an expected consequence of interionic attractions. 

Where both the acid and the base are strong electrolytes, the neutralization point will be at pH = 7 and the end point break will be distinct unless the solutions are very dilute (< 10" 3 mol dm"3). The composition of the titrand at any point in the titration may W computed from the total amount of acid and base present. However, when one of the reactants is a weak acid or base the picture is less clear. The incomplete dissociation of the acid or base and the hydrolysis of the salt produced in the reaction must be taken into account when.calculations of end points and solution composition are made. These points have been considered in chapter 3 and are used in the indicator selection procedure outlined in the preceding section of this chapter. 

The difficulty in explaining the effects of inorganic solutes on the physical properties of solutions led in 1884 to Arrhenius theory of incomplete and complete dissociation of ionic solutes (electrolytes, ionophores) into cations and anions in solution, which was only very reluctantly accepted by his contemporaries. Arrhenius derived his dissociation theory from comparison of the results obtained by measurements of electroconductivity and osmotic pressure of dilute electrolyte solutions [6]. 

The value ofjc,- in the case of an electrolyte derives from the number of moles of ions of species i actually present in solution. This number need not be equal to the number of moles of i expected of dissolved electrolyte if, for instance, the electrolyte is a potential one, then only a fraction of the electrolyte may react with the solvent to form ions, i.e., the electrolyte may be incompletely dissociated. 

Thallium(I) halides are predominantly ionic, although there is a tendency toward increasing covalent character in the series of compounds TlCl (17%), TlBr (20%), and TII (28%). This increased degree of covalency results in decreased solubility for example, TIF is soluble in water whilst the other Tl halides are only sparingly soluble. The thallium(I) halides are classical examples of incompletely dissociated 1 1 electrolytes. The stability of halide complexes of Tl is low and follows the order TIF < TlCl < TlBr < TII, where for the series of halides, Kx = -, 0.8, 2.1, 5.0 and Ki = -, 0.2, 0.7, 1.5 respectively. The fluoride ion F is preferred to perchlorate as a noncomplexing counterion. Claims have been made for T1X species with n = 3 and 4 however, the formation of complexes in aqueous solution with n > 2 seems unlikely. 

X 10 equiv. per liter, but direct determination gives 2.22 X 10 equiv. per liter. The difference is partly due to incomplete dissociation and partly to the formation of complex ions. In other words, the lanthanum oxalate does not ionize to yield simple La+ "+ and C2O4 ions, as is assumed in the conductance method for determining the solubility in addition complex ions, containing both lanthanum and oxalate, are present to an appreciable extent in the saturated solution. It is necessary, therefore, to exercise caution in the interpretation of the results obtained from conductance measurements with saturated solutions of sparingly soluble electrolytes. 

In addition to the reason for incomplete dissociation just considered, there are some cases, e.g., weak acids and many salts of the transition and other metals, in which the electrolyte is not wholly ionized. These substances exist to some extent in the form of un-ionized molecules a weak acid, such as acetic acid, provides an excellent illustration of this type of behavior. The solution contains un-ionized, covalent molecules, quite apart from the possibility of ion-pairs. With sodium chlorides, and similar electrolytes, on the other hand, there are probably no actual covalent molecules of sodium chloride in solution, although there may be ion-pairs in which the ions are held together by forces of electrostatic attraction. 

For incompletely dissociated electrolytes this involves a knowledge of the degree of dissociation, which may not always be available with sufficient accuracy. It is for this reason that the Debye-Htickel equations are generally tested by means of data obtained with strong electrolytes, since they can be assumed to be completely dissociated. It is probable that some of the discrepancies observed with certain electrolytes of high valence types are due to incomplete dissociation for which adequate allowance has not been made. 

Fig. 59 this is particularly true if both the saturating salt and the added electrolyte are of high valence types. The deviations are often due to incomplete dissociation, and also to the approximations made in the derivation of the Debye-Hiickel equations as already seen, both these factors become of importance with ions of high valence. 

Incomplete dissociation would also reduce the screening power of electrolytes. Ions of higher valency could, as a result of Ion association, become entitles of lower vzdency this has substantial consequences because the valency occurs In the exponentials, see for instance sec. 3.5c. For otherwise strong electrolytes this complication starts when the relative dielectric permittivity Is low. say < 30. 

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