Electrolytes, strong-weak rules

Regardless of classification, polysaccharides express varying degrees of ethanol tolerance neutral species are mostly unresponsive to electrolytes, including weak acid, and protonated species precipitate with low solvent retention from strong acid and Ca2+ solutions. Exceptions to many general rules abound for example, polyacids are generally the most stable species,... 

Let ns snmmarize the main points in this section. Compounds that dissolve in water are soluble those that dissolve little, or not at all, are insoluble. Soluble substances are either electrolytes or nonelectrolytes. Nonelectrolytes form noncon-dncting aqneons solutions because they dissolve completely as molecules. Electrolytes form electrically conducting solutions in water because they dissolve to give ions in solntion. Electrolytes can be strong or weak. Almost all soluble ionic substances are strong electrolytes. Soluble molecular substances usually are nonelectrolytes or weak electrolytes the latter solution consists primarily of molecules, but has a small percentage of ions. Ammonia, NH3, is an example of a molecular substance that is a weak electrolyte. A few molecular substances (such as HCl) dissolve almost entirely as ions in the solution and are therefore strong electrolytes. The solubility rules can be used to predict the solubility of ionic compounds in water. 

Reactions often involve ions in aqueous solution. Many of the compounds in such reactions are electrolytes, which are substances that dissolve in water to give ions. Electrolytes that exist in solution almost entirely as ions are called strong electrolytes. Electrolytes that dissolve in water to give a relatively small percentage of ions are called weak electrolytes. The solubility rules can be used to predict the extent to which an ionic compound will dissolve in water. Most soluble ionic compounds are strong electrolytes. 

In aqueous electrolyte solutions the molar conductivities of the electrolyte. A, and of individual ions, Xj, always increase with decreasing solute concentration [cf. Eq. (7.11) for solutions of weak electrolytes, and Eq. (7.14) for solutions of strong electrolytes]. In nonaqueous solutions even this rule fails, and in some cases maxima and minima appear in the plots of A vs. c (Eig. 8.1). This tendency becomes stronger in solvents with low permittivity. This anomalons behavior of the nonaqueous solutions can be explained in terms of the various equilibria for ionic association (ion pairs or triplets) and complex formation. It is for the same reason that concentration changes often cause a drastic change in transport numbers of individual ions, which in some cases even assume values less than zero or more than unity. 

All biological systems contain aqueous electrolyte solutions. These solutions consist of strong electrolytes (inorganic salts) as well as various organic substances with acidic or basic functional groups which usually behave as weak electrolytes. The solutions are often gel-like in their consistency because of the polyelectrolytes, proteins, and other macromolecules contained in them. The pH values of biological solutions as a rule are between 6.7 and 7.6. 

The law is in fact of very wide application it holds for nonionic as well as ionic reactions. The degree of ionization of weakly ionized substances can be calculated with high precision according to the law. But the behavior of strong electrolytes does not conform as closely to this law, and the law is of value only in a qualitative fashion to predict the extent of the ionization of these substances. In this connection we may recall Rule 4 for writing ionized equations, which directed to treat all strong electrolytes as if they were completely ionized. 

For low-molecular weak electrolytes the concentration dependence of conductance is more complex, as in addition to the interionic friction effect it is strongly influenced by the association-dissociation reactions taking place in the solutions. However, as these in general follow the mass action law and thus, in simpler cases, the van t Hoff dilution law, their conductivity behavior is predictable. As a rule their equivalent conductivity steeply increases on dilution due to the increased dissociation of the electrolyte. 

In this chapter we discuss some of the properties of electrolyte solutions. In Sec. 12-1, the chemical potential and activity coefficient of an electrolyte are expressed in terms of the chemical potentials and activity coefficients of its constituent ions. In addition, the zeroth-order approximation to the form of the chemical potential is discussed and the solubility product rule is derived. In Sec. 12-2, deviations from ideality in strong-electrolyte solutions are discussed and the results of the Debye-Hiickel theory are presented. In Sec. 12-3, the thermodynamic treatment of weak-electrolyte solutions is given and use of strong-electrolyte and nonelectrolyte conventions is discussed. 

In the following sections we will look at how a compound s composition lets us predict whether it is a strong electrolyte, weak electrolyte, or nonelectrolyte. For the moment, you need only to remember that water-soluble ionic compounds are strong electrolytes. Ionic compounds can usually be identified by the presence of both metals and nonmetals [for example, NaCl, FeS04, and A1(N03)3]. Ionic compounds containing the ammonium ion, NH [for example, NH4Br and (NH4)2C03], are exceptions to this rule of thumb. 

Anions of common strong acids, such as C104, S04, CF, NOa , etc. exhibit as a rule only weak complexing interactions, if any. Nevertheless, even weak complexation may be of importance in electrode kinetics if the complex ion undergoes electrode reaction more easily than the free metal ion, as is often the case, especially with chlorides. In such cases, the complex takes the role of an electroactive species, as already discussed for the hydroxo complexes. Thus, e.g., nickel can hardly be anodically dissolved at all if chloride ions are not present in the solution. In sulfate electrolytes, the oxidation product (some oxygen-containing species) forms a passive film and further dissolution is blocked soon after an anodic overpotential is imposed upon the electrode. The phenomenon of passivity is discussed elsewhere (cf. Volume 4). At this point, one should note that passivity is absent in the presence of chlorides. 

The roots of VB theory in chemistry can be traced to the famous paper of Lewis The Atom and The Molecule / which introduces the notions of electron-pair bonding and the octet rule. Lewis was seeking an understanding of weak and strong electrolytes in solution, and this interest led him to formulate the concept of the chemical bond as an intrinsic property of the molecule that varies between the covalent (shared-pair) and ionic situations. Lewis paper predated the introduction of quantum mechanics by 11 years, and constitutes... 

For weakly and moderately associated electrolytes, 10 < Ka < lO, generally no problem occurs in obtaining reliable values for Ka and Aq from conductivity measurements. However, strong association, as known for many salts in the solvent class 6, often entails unrealistic Ao-values and Ka-values. This is a problem of data analysis caused by the large extrapolation of very small Ai-values at the lowest concentration of a ran when compared with the expected Ao-value. Improved estimates for Ao are obtained with the help of the Walden rule at constant temperature T, if Ao is known for the electrolyte in another solvent (2) where a small association constant does not prohibit its determination, and rj is known for both solvents (1, 2). 

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