Aqueous equilibria solubility-product constant

The solubility product constant (Ksp) of EDTA was determined by adjusting the pH of an aqueous solution to a low value using nitric acid, and leaving the system to reach equilibrium overnight at room temperature. The precipitate was filtered off, dried at 105°C, and weighed to determine the amount of solubilized material. Alternatively, the precipitate was analyzed by complexometric titration, using standardized 0.05 M Zn(II) solution and xylenol orange as indicator [12]. The estimated value of the solubility product is 10 24 66 (pKsp = 24.66). 

The solubility-product constant, Ksp, is the equilibrium constant for an ionic solid in contact with a saturated aqueous solution. The two processes with equal rates in this case are dissolution and crystallization. 

For aqueous systems, a unit activity is expected for the solid species (i.e., we assume that the chemical reactivity of a solid in water is unchanging as long as there is solid in equilibrium with the solution). Also, for dilute concentrations, we assume that the activities are equal to the concentrations of the species. With these assumptions, we can reduce the solubility product constant equation to... 

The solubility product constant is a useful parameter for calculating the aqueous solubility of sparingly soluble compounds under various conditions. It may be determined by direct measurement or calculated from the standard Gibbs energies of formation AjG° of the species involved at their standard states. Thus if = [M ] ", [A ]" is the equilibrium constant for the reaction... 

Notice that the value of this solubility product constant is small there are relatively few silver or chromate ions in aqueous solution at 25 °C. If changing the solution composition shifts the equilibrium in Equation 1 to the right, more solid will dissociate and the solubility will increase relative to the solubility of silver chromate in pure water at 25 °C. If the equilibrium shifts to the left as different solution compositions are used, then the solubility relative to that in pure water will decrease. 

The equilibrium equations that normally have to be considered in the EKR modeling of a soil contaminated by heavy metals can be classified into one of the following categories complex formation reactions, precipitation of the metal hydroxides or of other species, ion exchange reactions, surface complexation reactions, etc. Anyway, the autoionization of water always has to be considered and the precipitation of carbonates, together with the carbonate-bicarbonate equilibrium, should normally also be considered. However, the above equations have only considered the species in aqueous phase, so if a species precipitates, a new master species has to be included in this equilibrium system, whose concentration would be the amount of the precipitated species per unit volume of water. This additional degree of freedom is constrained by the solubility product constant of the precipitate (KO, because the new solid phase is in equilibrium with the aqueous phase. If there exists Np precipitated species, the pure-phase equilibria can be represented with the following equation ... 

As with any other equilibrium, the extent to which this dissolution reaction occurs is expressed by the magnitude of the equilibrium constant. Because this equilibrium equation describes the dissolution of a solid, the equilibrium constant indicates how soluble the solid is in water and is referred to as the solubility-product constant (or simply the solubility product). It is denoted K p, where sp stands for solubility product The equilibrium-constant expression for the equilibrium between a solid and an aqueous solution of its component ions (Ksp) is written according to the rules that apply to any other equilibrium-constant expression. Remember, however, that soUds do not appear in the equihbrium-constant expressions for heterogeneous equihbrium. ooo (Section 15.4) Thus, the solubihty-product expression for BaS04, which is based on Equation 17.15, is... 

This equilibrium constant is called a solubility-product constant. In general, the solubility-product constant (K p) is the equilibrium constant for the equilibrium that exists between a solid ionic solute and its ions in a saturated aqueous solution. 

Ions of salts that are slightly soluble form saturated aqueous solutions at low concentrations. The solubility equilibrium expression for such salts yields a constant—the solubility product constant,... 

Is there a relationship between the solubility product constant, K p, of a solute and the solute s molar solubility—its molarity in a saturated aqueous solution As shown in Examples 18-2 and 18-3, there is a definite relationship between them. As discussed in Section 18-4, calculations involving Kgp are generally more subject to error than are those involving other equilibrium constants, but the results are suitable for many purposes. In Example 18-2, we start with an experimentally determined solubility and obtain a value of Xgp. 

By convention, a solubility equilibrium is written in the direction of a solid dissolving to give aqueous ions, and the equilibrium constant for this reaction is called the solubility product ( sp). Here, for example, is the reaction... 

Sol id Sol utions. The aqueous concentrations of trace elements in natural waters are frequently much lower than would be expected on the basis of equilibrium solubility calculations or of supply to the water from various sources. It is often assumed that adsorption of the element on mineral surfaces is the cause for the depleted aqueous concentration of the trace element (97). However, Sposito (Chapter 11) shows that the methods commonly used to distinguish between solubility or adsorption controls are conceptually flawed. One of the important problems illustrated in Chapter 11 is the evaluation of the state of saturation of natural waters with respect to solid phases. Generally, the conclusion that a trace element is undersaturated is based on a comparison of ion activity products with known pure solid phases that contain the trace element. If a solid phase is pure, then its activity is equal to one by thermodynamic convention. However, when a trace cation is coprecipitated with another cation, the activity of the solid phase end member containing the trace cation in the coprecipitate wil 1 be less than one. If the aqueous phase is at equil ibrium with the coprecipitate, then the ion activity product wi 1 1 be 1 ess than the sol ubi 1 ity constant of the pure sol id phase containing the trace element. This condition could then lead to the conclusion that a natural water was undersaturated with respect to the pure solid phase and that the aqueous concentration of the trace cation was controlled by adsorption on mineral surfaces. While this might be true, Sposito points out that the ion activity product comparison with the solubility product does not provide any conclusive evidence as to whether an adsorption or coprecipitation process controls the aqueous concentration. 

The solubility product is the equilibrium constant for the dissolution of a solid salt into its constituent ions in aqueous solution. The common ion effect is the observation that, if one of the ions of that salt is already present in the solution, the solubility of a salt is decreased. Sometimes, we can selectively precipitate one ion from a solution containing other ions by adding a suitable counterion. At high concentration of ligand, a precipitated metal ion may redissolve by forming soluble complex ions. In a metal-ion complex, the metal is a Lewis acid (electron pair acceptor) and the ligand is a Lewis base (electron pair donor). 

In dilute aqueous solutions, it has been demonstrated experimentally for poorly soluble ionic salts (solubilities less than 0.01 molL ) that the mathematical product of the total molar concentrations of the component ions is a constant at constant temperature. This product, is called the solubility product. Thus for a saturated solution of a simple ionic compound AB in water, we have the dynamic equilibrium ... 

K,. is the conventional solubility product, and the subscript 0 indicates that the equilibrium reaction involves only uncomplexed aqueous species. If the solubility constant includes the formation of aqueous complexes, a notation analogous to that of Eq.(11.12) is used ... 

The strength of association between the ions in solution is expressed by various equilibrium constants. Stability (formation) constants refer to complex ions and ion pairs hydrolysis (deprotonation) constants refer to the loss of H+ from the water ligands surrounding central cations. Solubility products refer to the aqueous ion activities in equilibrium with solid phases. Some constants are reported in the literature in terms of concentrations rather than activities. Such constants are misnamed, since they depend both on the concentration and on the nature of other ions in solution. Converting concentrations to activities gives a much more useful value. 

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