Aqueous solutions solubility product constant

Once the composition of the aqueous solution phase has been determined, the activity of an electrolyte having the same chemical formula as the assumed precipitate can be calculated (11,12). This calculation may utilize either mean ionic activity coefficients and total concentrations of the ions in the electrolyte, or single-ion activity coefficients and free-species concentrations of the ions in the electrolyte (11). If the latter approach is used, the computed electrolyte activity is termed an ion-activity product (12). Regardless of which approach is adopted, the calculated electrolyte activity is compared to the solubility product constant of the assumed precipitate as a test for the existence of the solid phase. If the calculated ion-activity product is smaller than the candidate solubility product constant, the corresponding solid phase is concluded not to have formed in the time period of the solubility measurements. Ihis judgment must be tempered, of course, in light of the precision with which both electrolyte activities and solubility product constants can be determined (12). 

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. 

Precipitation and dissolution of metal hydroxides The solubility product principle can also be applied to the formation of metal hydroxide precipitates these are also made use of in qualitative inorganic analysis. Precipitates will be formed only if the concentrations of the metal and hydroxyl ions are momentarily higher than those permitted by the solubility product. As the metal-ion concentration in actual samples does not vary much (10—1 —10 3 mol -1 is the usual range), it is the hydroxyl-ion concentration which has the decisive role in the formation of such precipitates. Because of the fact that in aqueous solutions the product of hydrogen- and hydroxyl-ion concentrations is strictly constant (A = 10 14 at 25°C, cf. Section 1.18), the formation of a metal-hydroxide precipitate depends mainly on the pH of the solution. Using the solubility product principle, it is possible to calculate the (minimum) pH required for the precipitation of a metal hydroxide. 

In Table 22-1 there are given values of the solubility-product constants at room temperature for many substances. More complete tables of values of these constants may be found in the handbooks and reference books mentioned at the end of Chapter 1. An extensive table App. Ill) and a discussion of the experimental data on which the values depend are given by W. M. Latimer,. T/ie Oxidation States of the Elements and Their Potentials in Aqueous Solution, Prentice-Hall, Inc., New York, 1938. 

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... 

A silver rod and a SHE are dipped into a saturated aqueous solution of silver oxalate, Ag2C204, at 25°C. The measured potential difference between the rod and the SHE is 0.589 V, the rod being positive. Calculate the solubility product constant for silver oxalate. 

Describe the process by which adding potassium hydroxide to a saturated aluminum hydroxide solution reduces the concentration of aluminum ions. Write the solubility equihbrium equation and solubility product constant expression for a saturated aqueous solution of aluminum hydroxide. 

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 true or thermodynamic solubility product constant K p will, for most substances, increase with the temperature by an amount that depends on the heat of solution. The addition of any organic solvent such as alcohol to an aqueous solution will generally result in a lower Kjp. This may be easily understood in terms of the increased work of separation of ions in a medium orxlower dielectric constant. 

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... 

The solubility product constant (Kjp) is a dissociation constant for metal complexes in aqueous solution. A smaller K,p value means that the complex is less likely to dissociate in water and implies that the complex is more stable. More stable metal sulfides drive Equation 3 to form the acid which then dopes the poiyaniline resulting in a large resistance decrease in the film. Table I shows the Ksp values for different metal sulfides. According to the table, CuS is the most stable metal sulfide and ZnS is the least stable. This is consistent with CuS exhibiting the largest response, whereas ZnS gives the smallest response. 

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. 

By experiment, it is found that 1.2 X 10 mol of lead(II) iodide, Pbl2, dissolves in 1 L of aqueous solution at 25°C. What is the solubility product constant at this temperature ... 

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. 

Urea-formaldehyde resins are generally prepared by condensation in aqueous basic medium. Depending on the intended application, a 50-100% excess of formaldehyde is used. All bases are suitable as catalysts provided they are partially soluble in water. The most commonly used catalysts are the alkali hydroxides. The pH value of the alkaline solution should not exceed 8-9, on account of the possible Cannizzaro reaction of formaldehyde. Since the alkalinity of the solution drops in the course of the reaction, it is necessary either to use a buffer solution or to keep the pH constant by repeated additions of aqueous alkali hydroxide. Under these conditions the reaction time is about 10-20 min at 50-60 C. The course of the condensation can be monitored by titration of the unused formaldehyde with sodium hydrogen sulfite or hydroxylamine hydrochloride. These determinations must, however, be carried out quickly and at as low temperature as possible (10-15 °C), otherwise elimination of formaldehyde from the hydroxymethyl compounds already formed can falsify the analysis. The isolation of the soluble condensation products is not possible without special precautions, on account of the facile back-reaction it can be done by pumping off the water in vacuum below 60 °C imder weakly alkaline conditions, or better by careful freeze-drying. However, the further condensation to crosslinked products is nearly always performed with the original aqueous solution. 

The salts of the series are bright red crystalline bodies. They are soluble in water, neutral in reaction, and dilute mineral acids do not transform them into aquo-salts. The least soluble member of the series is the sulphate [(NH3)4Co (OH)2 Co(NH3)4](S04)a.2H30, which is prepared by heating hydroxo-aquo-tetrammino-cobaltic sulphate at 100° C. till it is constant in weight. The mass is extracted with water and the sparingly soluble sulphate collected and dried. The crude product so obtained is converted into the chloride and an aqueous solution of this then treated with a solution of sodium sulphate, when a crystalline precipitate of the diol-sulphate is obtained. It is collected, washed with water, alcohol, and finally with ether. It forms small red needle-shaped crystals which contain two molecules of water of hydration. 

Tetracycline antibiotics are closely related derivatives of the polycyclic naphtha-cenecarboxamide. They are amphoteric compounds with characteristic dissociation constants corresponding to the acidic hydroxyl group at position 3 (pK about 3.3), die dimethylamino group at position 4 (pK, about 7.5), and the hydroxyl group at position 12 (pK about 9.4). In aqueous solutions of pH 4-7, tetracyclines exist as dipolar ions, but as the pH increases to 8-9 marked dissociation of the dimethylamine cation occurs. They are soluble in acids, bases, and alcohols but are quite insoluble in organic solvents such as chloroform. Their ultraviolet spectra show strong absorption at around 270 and 360 nm in neutral and acidic solutions. Tetracyclines are readily transformed into fluorescent products in the presence of metal ions or under alkaline conditions. 

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