Solutions of weak electrolytes

The methods for obtaining expressions for the chemical potential of a component that is a weak electrolyte in solution are the same as those used for strong electrolutes. For illustration we choose a binary system whose components are a weak electrolyte represented by the formula M2A and the solvent. We assume that the species are M +, MA , A2-, and M2A. We further assume that the species are in equilibrium with each other according to 

The first problem is to obtain relations between the chemical potential of the component and those of the species. The Gibbs energy of a solution containing n, moles of solvent and n2 moles of solute is given by 

Neither the mole numbers of the species nor their chemical potentials are all independent, but are subject to the condition of mass balance 

When these four condition equations are used in Equation (8.193) to make all of the remaining mole numbers and chemical potentials independent, we 

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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. 

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. 

Any one of these expressions for fCa represents what is known as Ostwald s dilution law, which has essentially been obtained by applying the law of mass action to solutions of weak electrolytes. It deals with the variation in the degree of dissociation with concentration or dilution of solutions of weak electrolytes. It is not applicable to solutions of strong electrolytes. The failure of strong electrolytes to obey Ostwald s dilution law is known as the anomaly of strong electrolytes. 

The first theory of solutions of weak electrolytes was formulated in 1887 by S. Arrhenius (see Section 1.1.4). If the molar conductivity is introduced into the equations following from Arrhenius concepts of weak electrolytes, Eq. (2.4.17) is obtained, known as the Ostwald dilution law this equation... 

The conductivity also increases in solutions of weak electrolytes. This second Wien effect (or field dissociation effect) is a result of the effect of the electric field on the dissociation equilibria in weak electrolytes. For example, from a kinetic point of view, the equilibrium between a weak acid HA, its anion A" and the oxonium ion H30+ has a dynamic character ... 

Many of the undesirable substances present in gaseous or liquid streams form volatile weak electrolytes in aqueous solution. These compounds include ammonia, hydrogen sulfide, carbon dioxide and sulfur dioxide. The design and analysis of separation processes involving aqueous solutions of these materials require accurate representation of the phase equilibria between the solution and the vapor phase. Relatively few studies of these types of systems have been published concerning solutions of weak electrolytes. This paper will review the methods that have been used for such solutions and, as an example, consider the alkanolamine solutions used for the removal of the acid gases (H2S and C02) from gas streams. 

Solutions of Weak Electrolytes Van Krevelen et al. (2) measured the vapor pressures of aqueous... 

About the same time Beutier and Renon (11) also proposed a similar model for the representation of the equilibria in aqueous solutions of weak electrolytes. The vapor was assumed to be an ideal gas and < >a was set equal to unity. Pitzer s method was used for the estimation of the activity coefficients, but, in contrast to Edwards et al. (j)), two ternary parameters in the activity coefficient expression were employed. These were obtained from data on the two-solute systems It was found that the equilibria in the systems NH3+ H2S+H20, NH3+C02+H20 and NH3+S02+H20 could be represented very well up to high concentrations of the ionic species. However, the model was unreliable at high concentrations of undissociated ammonia. Edwards et al. (1 2) have recently proposed a new expression for the representation of the activity coefficients in the NH3+H20 system, over the complete concentration range from pure water to pure NH3. it appears that this area will assume increasing importance and that one must be able to represent activity coefficients in the region of high concentrations of molecular species as well as in dilute solutions. Cruz and Renon (13) have proposed an expression which combines the equations for electrolytes with the non-random two-liquid (NRTL) model for non-electrolytes in order to represent the complete composition range. In a later publication, Cruz and Renon (J4J, this model was applied to the acetic acid-water system. 

Conclusions The correlation of vapor-liquid equilibria in aqueous solutions of weak electrolytes is important for the separation of undesirable components from gases and liquids. The major problem in such correlations is the estimation of the activity... 

In principle, this system of 20 equations can be solved provided the equilibrium constants, activities, Henry-constants and fugacities are available. While some results for most of these properties are available, there exists no approved method for calculating activities in concentrated aqueous solutions of weak electrolytes therefore, several approximations were developed. ... 

At high field strengths a conductance Increase Is observed both In solution of strong and weak electrolytes. The phenomena were discovered by M. Wien (6- ) and are known as the first and the second Wien effect, respectively. The first Wien effect Is completely explained as an Increase In Ionic mobility which Is a consequency of the Inability of the fast moving Ions to build up an Ionic atmosphere (8). This mobility Increase may also be observed In solution of weak electrolytes but since the second Wien effect Is a much more pronounced effect we must Invoke another explanation, l.e. an Increase In free charge-carriers. The second Wien effect Is therefore a shift in Ionic equilibrium towards free ions upon the application of an electric field and is therefore also known as the Field Dissociation Effect (FDE). Only the smallness of the field dissociation effect safeguards the use of conductance techniques for the study of Ionization equilibria. 

One of the many ways to classify inorganic compounds is into electrolytes, nonelectrolytes, and weak electrolytes. When electrolytes are dissolved in water, the resulting solution is a good conductor of electricity the water solutions of nonelectrolytes do not conduct electricity the solutions of weak electrolytes are very poor conductors. Water itself is an extremely poor conductor of electricity. A flow of current is a movement of electrical charges caused by a difference in potential (voltage) between the two ends of the conductor. 

Electrodes B consist of fine platinum wires supported upon glass rods, and are to be used with a lamp of about 15 watts. They are to be used in testing the conductivity of solutions of weak electrolytes in a 3-inch vial. This vial may be raised until the electrodes are immersed in the liquid. Before testing the conductivity of any given solution rinse the platinum electrodes with... 

Activity coefficients of non-ionized molecules do not differ appreciably from unity. In dilute solutions of weak electrolytes the differences between activities and concentrations (calculated from the degree of dissociation) is negligible. 

Strong electrolytes completely dissociate into ions and conduct electricity well. Weak electrolytes provide few ions in solution. Therefore, even in high concentrations, solutions of weak electrolytes conduct electricity weakly. Ionic compounds are usually strong electrolytes. Covalent compounds may be strong electrolytes, weak electrolytes, or nonconductors. 

Electroacoustic method, 80 Electrokinetic potential effect on interaction energy between colloidal particles, 248 effect of ionic strength on, 66 of macroscopic samples, 79 of oxides, effect of hydration time on, 76 relationship to thd. 649, 656 in solution of weak electrolyte, 242 at very high ionic strengths, 266, 267 Electronegativity, correlation with PZC, 213... 

It is very likely that in solutions of weak electrolytes a reversible equilibrium exists between the two molecular types, one of which behaves as a strong electrolyte while the other has practically no electrolyte properties. These forms may be called respectively the... 

Laskowski, J.S., Electrokinetic measurements in aqueous solutions of weak electrolyte type surfactants, 7. Colloid Interf. Sci., 159, 349, 1993. 

In solutions of weak electrolytes, the concentration of ions is often sufficiently low for the solution to approximate to ideality. However, with moderately weak electrolytes, it becomes possible to have higher concentrations of ions present and deviations from the behaviour described above become apparent. For such electrolytes corrections for non-ideality must be made (see Chapter 12). 

Solutions of weak electrolytes are not confined to acids and bases, e.g. solutions of the 1-1 electrolyte, tetrabutylammonium iodide, N(C4H9)4+n(aq), the 2-2 electrolyte, MgS04(aq), and the 3-3 lanthanum hexacyanoferrate(III), La Fe(CN)6 (aq), all form ion pairs in aqueous solution, and so are weak electrolytes. 

For electrolytes which are not fully dissociated highly accurate data should be able to detect the onset of ion pair formation and, as the concentration increases, significant association is expected. The higher the charge type the lower the concentration at which association will be observed. Ion association results in Aobsvd values being lower than expected. For solutions of weak electrolytes where undissociated molecules are present similar behaviour is observed, viz. Aobsvd values will approach the limiting law slope from below. 

This last condition is fulfilled when the ionic concentrations are very low, as they are in fact in dilute solutions of weak electrolytes. The dissociation constants of substances such as weak organic acids can be determined by a combination of the formulae of Ostwald and Arrhenius, but the procedure is quite inadmissible for salts. Here the degree of dissociation is large. In fact the value of a is often indistinguishable from unity, and the mutual influences of the ions are considerable. They are calculable in principle by methods due to Debye and Hiickel, and operate differently on different properties. The procedure outlined on p. 276 allows the calculation of the activity coefiicients. In general the thermodynamic properties of the salt in solution correspond to those of a system with apparently incomplete dissociation, not because the concentrations of the ions are reduced by molecule formation but because the activity coefficients are lowered by mutual ionic influences. 

Reactive separations are tvidespread operations. In typical reactive separation processes such as reactive absorption or distillation the superposition of reaction and separation is deliberately used. In other cases, simultaneous reaction and separation simply cannot be avoided. This is, for instance, the case when side reactions occur in separation equipment or when intrinsically chemically reactive mixtures, such as solutions of weak electrolytes or formaldehyde solutions, have to be separated. Furthermore, in many reactors products are directly removed, which is basically a reactive separation. 

The formaldehyde + water + methanol system is just one example for many technically important systems, which are intrinsically chemically reactive. Other examples include aqueous solutions of weak electrolytes, such as amine solutions used to scrub carbon dioxide and other sour gases from gaseous streams, or the solutions containing bases used for chemical extraction of acids from aqueous streams. In many of these cases, the key to the development of predictive thermodynamic models is a quantitative model of the often complex chemical reactions in those mixtures. The necessary information often only can be obtained using spectroscopic methods [25]. 

For equivalent concentrations, solutions of strong electrolytes contain many more ions than do solutions of weak electrolytes. As a result, solutions of strong electrolytes are better conductors of electricity. Consider two solutions, 1M HCl and 1M HC2H3O2. Hydrochloric acid is almost 100% ionized acetic acid is about 1% ionized. (See Figure 15.3.) Thus, HCl is a strong acid and HC2H3O2 is a weak acid. Hydrochloric acid has about 100 times as many hydronium ions in solution as acetic acid, making the HCl solution much more acidic. 

The conductivity of a strong electrolyte is also increased at high field strengths because the ions are moving so fast that the ionic atmospheres are unable to form completely. This is known as the Wien effect it will also operate in solutions of weak electrolytes, but it is considerably smaller than the dissociation field effect and can be eliminated at sufficiently high field strengths. 

This was fine for dilute solutions of weak electrolytes but the contradiction with experimental results for strong electrolytes led to the realization that the electrostatic force between ions must be taken into account. 

Although Pitzer s method of calculating activity coefficients has been applied with some success to solutions containing weak electrolytes (OPl, OP14, it s form is best suited to strong electrolyte solutions. This is becausei in solutions of weak electrolytes, significant concentrations of molecular solutes are present. While... 

Pitzer considered the ion-ion interactions, ion-molecule and molecule-molecule interactions, which are important in solutions of weak electrolytes, are ignored. In 1979, Chen et al. (Cl) presented an extension to Pitzer s method to allow for both molecular and ionic solutes. More recently they proposed the addition of a local composition expression to account for the short range interactions between all species (C2, C3, C4),... 

In practice, this condition is satisfied whenever very dilute solutions or solutions of weak electrolytes are measured with low-frequency alternating currents (/=40 - 80 Hz) or when solutions having conductivities of about 100 pS/cm are measured with frequencies in the kilohertz range. In this way, y s can be determined fairly accurately with a Wheatstone bridge circuit, shown in simplified form in Figure 14. 

At large applied voltages (i.e., in strong electric field of about lO -lO V/cm), the velocity of ions is too large to allow perfect establishment of the cloud of solvent molecules around the ions. The actually observed conductance is larger than expected the effect is called electrophoretic or (first) Wien effect. In solutions of weak electrolytes (i.e., incompletely dissociated ones), the large strength of the electric field may cause artificially enhanced dissociation this also results in increased conductance. It is called dissociation... 

The approximate constancy of the values may be attributed partly to the fact that the concentration of the free ions in the various solutions is very low, so that is close to unity, and partly to the fact that the Ostwald assumption a=A/Aop is probably fairly near to the truth under the same conditions of high dilution of the ions. In general, the Ostwald method gives reasonably satisfactory results for dilute solutions of weak electrolytes for which the dissociation constant is of the order 10 or less. 

Solutions of weak electrolytes conduct electricity poorly because only a small quantity of ions forms when the electrolyte dissolves. Solutions of nonelectrolytes do not conduct electricity because essentially no ions exist in the solution. 3. The movement of ions makes up an electric current in solution. Soluble molecular compounds are generally neutral, and are usually nonelectrolytes. Some molecular compounds can react with water, forming ions as a... 

Both of these examples of equilibrium are very important in real-world analytical laboratories. Solutions of weak electrolytes are used to control the acidity of solutions in laboratory analyses. Such solutions are known as buffer solutions. 

A weak electrolyte is a compound that dissolves in water mostly as undissociated molecules. Only a few of the dissolved solute molecules separate, producing a small number of ions in solution. Thus, solutions of weak electrolytes do not conduct electrical current as well as solutions of strong electrolytes. When the electrodes are placed in a solution of a weak electrolyte, the glow of the light bulb is very dim. For example, an aqueous solution of the weak electrolyte HF contains mostly HF molecules and only a few and F ions. As more H+ and F ions form, some recombine to give HF molecules. These forward and reverse reactions of molecules to ions and back again are indicated by two arrows between the reactants and products that point in opposite directions ... 

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