Strong electrolytes chemical equilibrium

The form in which chemical analyses of sea water are given records the history of our thought concerning the nature of salt solutions. Early analytical data were reported in terms of individual salts NaCl, CaSO/i, and so forth. After development of the concept of complete dissociation of strong electrolytes, chemical analyses of sea water were given in terms of individual ions Na+, Ca++, Cl-, and so forth, or in terms of known undissociated and partly dissociated species, e.g., HC03 , In recent years there has been an attempt to determine the thermodynamically stable dissolved species in sea water and to evaluate the relative distribution of these species at specified conditions. Table 1 lists the principal dissolved species in sea water deduced from a model of sea water that assumes the dissolved constituents are in homogeneous equilibrium, and (or) in equilibrium, or nearly so, with solid phases. 

To test the validity of the extended Pitzer equation, correlations of vapor-liquid equilibrium data were carried out for three systems. Since the extended Pitzer equation reduces to the Pitzer equation for aqueous strong electrolyte systems, and is consistent with the Setschenow equation for molecular non-electrolytes in aqueous electrolyte systems, the main interest here is aqueous systems with weak electrolytes or partially dissociated electrolytes. The three systems considered are the hydrochloric acid aqueous solution at 298.15°K and concentrations up to 18 molal the NH3-CO2 aqueous solution at 293.15°K and the K2CO3-CO2 aqueous solution of the Hot Carbonate Process. In each case, the chemical equilibrium between all species has been taken into account directly as liquid phase constraints. Significant parameters in the model for each system were identified by a preliminary order of magnitude analysis and adjusted in the vapor-liquid equilibrium data correlation. Detailed discusions and values of physical constants, such as Henry s constants and chemical equilibrium constants, are given in Chen et al. (11). 

A wide variety of data for mean ionic activity coefficients, osmotic coefficients, vapor pressure depression, and vapor-liquid equilibrium of binary and ternary electrolyte systems have been correlated successfully by the local composition model. Some results are shown in Table 1 to Table 10 and Figure 3 to Figure 7. In each case, the chemical equilibrium between the species has been ignored. That is, complete dissociation of strong electrolytes has been assumed. This assumption is not required by the local composition model but has been made here in order to simplify the systems treated. 

Nitric acid is a strong electrolyte. Therefore, the solubilities of nitrogen oxides in water given in Ref. 191 and based on Henry s law are utilized and further corrected by using the method of van Krevelen and Hofhjzer (77) for electrolyte solutions. The chemical equilibrium is calculated in terms of liquid-phase activities. The local composition model of Engels (192), based on the UNIQUAC model, is used for the calculation of vapor pressures and activity coefficients of water and nitric acid. Multicomponent diffusion coefficients in the liquid phase are corrected for the nonideality, as suggested in Ref. 57. 

For a 1 1 strong electrolyte with ion pairing, the concentration of both ions is aMj, where a is the fraction of ion pairs that are dissociated, (putij = ,e, the extent of the dissociation reaction). The chemical potential and activity of such electrolytes are easily calculated, because, at equilibrium,... 

From Eqn. (14) it follows that with an exothermic reaction - and this is the case for most reactions in reactive absorption processes - decreases with increasing temperature. The electrolyte solution chemistry involves a variety of chemical reactions in the liquid phase, for example, complete dissociation of strong electrolytes, partial dissociation of weak electrolytes, reactions among ionic species, and complex ion formation. These reactions occur very rapidly, and hence, chemical equilibrium conditions are often assumed. Therefore, for electrolyte systems, chemical equilibrium calculations are of special importance. Concentration or activity-based reaction equilibrium constants as functions of temperature can be found in the literature [50]. 

Also the so called degree of dissociation, determined from the colligative properties, does not agree with the result obtained from the measurement of the electrical conductance. Finally the law of chemical equilibrium, applicable to the dissociation of weak electrolytes, cannot be applied to the strong ones. 

Partial Molar Volume 1 0. Cryoscopic Determination of Molar Mass 1 1. Freezing-Point Depression of Strong and Weak Electrolytes 12. Chemical Equilibrium in Solution... 

The factor by which the. ion concentration is to be multiplied to obtain the ion activity is called the activity coefficient. For all strong electrolytes containing only univalent ions (HCl, NaCl, KNOg, etc.) its values are approximately 0.80 at 0.1 F, 0.90 at 0.01 F, and 0.96 at 0.001 F, approaching 1 only in very dilute solutions. These activity coefficients are of significance in connection with chemical equilibrium which is to be discussed later. 

There exist solutions whose properties, such as their electrical conductivity, indicate that the solute molecules are at least partially dissociated into ions. These solutions are termed electrolyte solutions. In general, solute molecules in which the chemical bonds have a large degree of ionic, rather than covalent, character will dissolve in polar solvents to yield electrolyte solutions. A solute is called a strong electrolyte if it completely dissociates into ions in solutions. Weak electrolytes are those for which an equilibrium is set up between undissociated molecules and constituent ions in solutions. 

Here, in contrast to Eq. (12-3) in the strong-electrolyte case, the equality of chemical potentials is assumed to imply the existence of a chemical equilibrium. Since it is impossible to freeze the reaction [Eq. (12-38)] in order to study the components independently, the chemical-equilibrium assumption is nonoperational. 

Any substance whose aqueous solution contains ions is called an electrolyte. Any substance that forms a solution containing no ions is a nonelectrolyte. Electrolytes that are present in solution entirely as ions are strong electrolytes, whereas those that are present partly as ions and partly as molecules are weak electrolytes. Ionic compounds dissociate into ions when they dissolve, and they are strong electrolytes. The solubility of ionic substances is made possible by solvation, the interaction of ions with polar solvent molecules. Most molecular compounds are nonelectrolytes, although some are weak electrolytes, and a few are strong electrolytes. When representing the ionization of a weak electrolyte in solution, half-arrows in both directions are used, indicating that the forward and reverse reactions can achieve a chemical balance called a chemical equilibrium. 

The chemical potential of the solid crystal salt B is in phase equilibrium with the dissolved salt B in the liquid or aqueous phase. In aqueous systems we are primarily dealing with salts of strong electrolytes, which in water dissociate completely to the constituent cations and anions of the salt. The chemical potential of the dissolved salt is then given by... 

Guggenheim, E. A. Stokes, R. H. Equilibrium Properties of Aqueous Solutions of Single Strong Electrolytes in The International Encyclopedia of Physical Chemistry and Chemical Physics Robinson, R. A., Ed. Pergamon Press Glasgow, 1969 Vol. 1, Topic 15. 

A prime requirement for all our calculations must be stressed at the start Solutions must be at equilibrium in order for calculations with equilibrium constants to be valid. This is obvious, but can cause great difficulty because chemical knowledge must be brought into play to decide whether a chemically stable solution is present. For example, how does one treat the acidity of a mixture which is made so that it will be 0.1 M in hydrochloric acid and 0.2 M in sodium hydroxide No such solution exists. One must know that these are strong electrolytes and that such high concentrations of H and OH react to form water. Thus, the mixture given turns into 0.1 M Na, Cl and 0.1 M Na, OH", that is, half the base is neutralized by the acid present. The approximate pH is 13. We do not follow the procedure of one canny student who replied that the pH of the 0.1 M acid is 1.0 and that of the 0.2 M base is 13.3, and thus the average is 7.15. 

In contrast to the situation with strong electrolytes, the degree of ionization of weak electrolytes is in good agreement with the law of chemical equilibrium. For the reaction... 

First, because of the electrostatic forces exerted over a long distance, it is certain that the exchange energy Wab between an ion Na and an ion Ca is far from negligible. The solution obtained will be far from perfect and solubility will be limited. In this case, it is extremely difficult to model the imperfection of the solution. The easiest way to do so, when the solution is involved in a state of equilibrium, is to use quasi-chemical modeling with structure elements, thus enabling us to consider the solution to be perfect. Certain authors advocate the use of the quasi-chemical approximation, but with the supposition that the solution obeys the Debye and Hiickel theory of strong electrolytes. ... 

Alkali-metal halides are textbook examples of strong electrolytes. However, conductance and potentiometric measurements reveal that there are some salts which behave differently and do form ion pairs by the strong attraction of the unlike ions. For these systems, a chemical model of ion pairing as proposed in Refs. 17 to 20 can be applied to consider the equilibrium between the completely dissociated electrolyte and the ion pair... 

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. 

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... 

In general, the chemical potential of the solution in the micellar phase must equal that in the surrounding aqueous medium when thermodynamic equilibrium is established. Nonpolar solutes, such as the permanent gases, which do not interact strongly with either phase may be distributed rather evenly over the whole microheterogeneous system (39). On the other hand, typical electrolytes are practically restricted to the aqueous medium, while molecules of hydrophobic substances, e.g. hydrocarbons, are almost totally sequestered in the micelles. 

A mixed-bed resin is a mixture of chemically equivalent amounts of a strongly acidic cation exchanger in the free acid form and a strongly basic anion exchanger in the hydroxide form. If such a mixture is stirred with a solution of any electrolyte Caf An until equilibrium is reached, the following processes occur ... 

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