Equilibrium constant weak electrolyte dissociation

Compounds of this type may be classified as strong electrolytes, which dissociate almost completely into ions in solution, or as weak electrolytes, which only dissociate to a small extent in solution. Since strong electrolytes are almost completely dissociated in solution, measurement of the equilibrium constant for their dissociation is very difficult. For weak electrolytes, however, the dissociation can be expressed by the law of mass action in terms of the equilibrium constant. 

In the pure state, water is dissociated to a very small extent and behaves as a weak electrolyte. The equilibrium constant of the dissociation, H2O —- H -1- OH , is given by. 

As the titration begins, mostly HAc is present, plus some H and Ac in amounts that can be calculated (see the Example on page 45). Addition of a solution of NaOH allows hydroxide ions to neutralize any H present. Note that reaction (2) as written is strongly favored its apparent equilibrium constant is greater than lO As H is neutralized, more HAc dissociates to H and Ac. As further NaOH is added, the pH gradually increases as Ac accumulates at the expense of diminishing HAc and the neutralization of H. At the point where half of the HAc has been neutralized, that is, where 0.5 equivalent of OH has been added, the concentrations of HAc and Ac are equal and pH = pV, for HAc. Thus, we have an experimental method for determining the pV, values of weak electrolytes. These p V, values lie at the midpoint of their respective titration curves. After all of the acid has been neutralized (that is, when one equivalent of base has been added), the pH rises exponentially. 

The conductivity of a solution containing such molecular ions may be small compared with the value that would result from complete dissociation into atomic ions. In this way, in the absence of neutral molecules, we can have a weak electrolyte. The association constant for (29) has a value that is, of course, the reciprocal of the dissociation constant for the molecular ion (PbCl)+ the logarithms of the two equilibrium constants have the same numerical value, but opposite sign. 

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). 

DETERMINATION OF EQUILIBRIUM CONSTANTS FOR DISSOCIATION OF WEAK ELECTROLYTES... 

It is instructive to consider the effect of dissociation on the adsorption of amphipathic substances since many of the compounds that behave according to curve 3 are electrolytes. We consider only the case of strong 1 1 electrolytes for weak electrolytes the equilibrium constant for dissociation must be considered. 

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. 

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]. 

The equilibrium constant given by Eq. (9) using a values obtained from Eq. (8) differs from K, the true equilibrium constant in terms of activities, owing to the omission of activity coefficients (y ) from the numerator of Eq. (9) and the approximations inherent in Eq. (8). At the very low ionic concentrations encountered in the dissociation of a weak electrolyte, a simple extrapolation procedure can be developed to obtain from the values of Since y is an excellent approximation, it follows that... 

Weak acids are weak electrolytes and do not dissociate completely. An equilibrium exists between the reactants and the products, and the equilibrium constant must be taken into account to solve for the pH value. When a weak acid (HA) is dissolved in water, the conjugate base (A ) and conjugate acid (H+) are... 

When the source of the catalytically active hydrogen ion is a weak acid, one has to consider the weak electrolyte equilibrium involved and the change of the dissociation constant with electrolyte concentration, medium, and temperature. Br0nsted (7) termed this phenomenon secondary kinetic salt effect, but the writer would prefer to omit the word kinetic and substitute electrolyte for salt. The understanding of these... 

The dissociation of weak electrolytes or the solubility of slightly soluble substances can be quantitatively described by equilibrium constants. Equilibrium constants for completely dissolved and dissociated electrolytes are effectively infinite. Consider the dissociating species AB ... 

The weak electrolyte AB dissociates to A" and B, with a thermodynamic equilibrium constant K°q of 2 X 10 (a) Calculate the molar equilibrium constant i eq-... 

In Chapter 6, we discussed the thermodynamic equilibrium constant based on activities rather than on concentrations. Diverse salts affect the activities and therefore the extent of dissociation of weak electrolytes such as weak acids or bases. 

The equivalent conductivity of a weak electrolyte varies approximately with (Fig. 31.3). Explain this in terms of the equilibrium constant for small degree of dissociation. 

Conversely, Eq. (21.42) can be used to determine the degree of dissociation a of a weak electrolyte at a given concentration c by measuring the molar conductivity. Moreover, with the help of Eq. (21.41), the equilibrium constant of the substance becomes accessible. However, for these calculations we need the limiting molar conductivity A . This quantity is very difficult to find experimentally because the steep rise of the A at low concentrations makes an extrapolation to infinite dilution very uncertain. The law of independent migration of ions [Eq. (21.35)] offers a way out. hi the case of infinite dilution, the limiting molar conductivity of acetic acid is the sum of the contributions of cation and anion ... 

SECTION 16.6 Weak acids are weak electrolytes only some of the molecules exist in solution in ionized form. The extent of ionization is expressed by the acid-dissociation constant, K , which is the equilibrium constant for the reaction HA(aq), H (flq) + A (oq), which... 

Care must be taken in evaluating the ionic strength contribution of weak electrolytes. For example, if any of the solutions above contained phosphoric or acetic acid also, the ionic strength would be essentially the same, because only the dissociated phosphoric or acetic acids contributes, and this is generally very small. If, on the other hand, H3PO4 is the only solute present, then an approximate equilibrium calculation must be carried out (see Example 3.3), an ionic strength calculated, and the process repeated until values of I remain constant. This may take one or two successive approximations. 

Note the strong dependence on the charge number. As shown elsewhere, the equilibrium constant of a 1 1 weak electrolyte like acetic acid is increased by an electric field of 100 kV cm to about 14%, that for a 2 2 electrolyte like MgS04 to about 110%. ° Compared to simple dipolar equilibria of small molecules where electric-field-induced changes in K are very small, we see that the dissociation step of simple ion pairs is associated... 

We use the hydrolysis of A into P and Q as an illustration. Examples are the hydrolysis of benzylpenicillin (pen G) or the enantioselective hydrolysis of L-acetyl amino acids in a DL-mixture, which yields an enantiomerically pure L-amino acid as well as the unhydrolysed D-acetyl amino acid. In concentrated solutions these hydrolysis reactions are incomplete due to the reaction equilibrium. It is evident that for an accurate analysis of weak electrolyte systems, the association-dissociation reactions and the related phase behaviour of the reacting species must be accounted for precisely in the model [42,43]. We have simplified this example to neutral species A, P and Q. The distribution coefficients are Kq = 0.5 and Kp = K = 2. The equilibrium constant for the reaction K =XpXQ/Xj = 0.01, where X is a measure for concentration (mass or mole fractions) compatible with the partition coefficients. The mole fraction of A in the feed (z ) was 0.1, which corresponds to a very high aqueous feed concentration of approximately 5 M. We have simulated the hydrolysis conversion in the fractionating reactor with 50-100 equilibrium stages. A further increase in the number of stages did not improve the conversion or selectivity to a significant extent. Depending on the initial estimate, the calculation requires typically less than five iterations. 

As we have seen in several examples in this chapter, HCN acts as an acid in aqueous solutions. We introduced a few fundamental concepts of acids and bases in Chapter 3, but the context of equilibrium allows us to explore them further. Recall that we distinguished between strong acids (or bases), which dissociate completely in solution, and weak acids (or bases), which dissociate only partially. At this point in our study of chemistry, we should realize that this partial dissociation of weak electrolytes was an example of a system reaching equilibrium. So we can use equilibrium constants to characterize the relative strengths of weak acids or bases. One common way to do this is to use the pH scale, which we will define in this section. 

Conductance reflects the concentration of ions in the solution, not of neutral molecules. It is therefore a suitable method for determining the equilibrium constants of both strong and weak electrolytes. These determinations are based on Eqs. (44) and (46) and on other relationships not given here. Equation (46) is especially suitable for strongly associated electrolytes where the interactions between the ions are small and the resulting values of and Ka are reasonably accurate. The dissociation of water, alcohols, etc., as well as the solubility product can be estimated on the basis of conductance. A more complicated approach must be taken for cases where the solvent itself is strongly dissociated. 

Ostwald (2) combined the Arrhenius theory of electrolytic dissociation with the law of mass action and calculated the dissociation constants of various weak acids from the results of conductivity measurements. The existence of complex ions could be deduced from distribution experiments (3) and solubility behavior (4) as well as from rate studies, and several equilibrium constants were determined. 

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