Solubilities of sparingly soluble salts

Solubility of Sparingly Soluble Saits.—Determinations of the conductivity may also be employed, very advantageously, for the determination of the solubility of sparingly soluble salts. 

The equivalent conductivity of a solution is given, as we have seen, by the expression Aj, = K. p, where kis the specific conductance and the volume (in c.c.) in which i gram-equivalent of solute is dissolved. That is to say, if the saturated solution of a salt contains i gram-equivalent in Jt c.c., the conductivity would be equal to or, the volume in cubic centimetres containing i gram-equivalent of solute will be /k. A is, however, equal to aA., where a is the 

Experiment.—Determine the Solubility of Lead Sulphate or of Silver Chloride in Water at 25°. 

The conductivity of the water employed should first be determined at 25°. A quantity of finely powdered lead sulphate or silver chloride is then shaken repeatedly with the conductivity water in order to remove any impurities of a comparatively soluble nature. The well-washed salt is then placed along with conductivity water in a hard-glass vessel, which is placed in a thermostat at 25°, and shaken from time to time. After intervals of about quarter of an hour, a quantity of the solution is transferred to a conductivity cell, and the conductivity determined. This is repeated with fresh samples of the solution, until constant values are obtained. 

The conductivity so determined is corrected by subtracting the conductivity of the water employed, and the solubility 

---

An important application of the solubility product principle is to the calculation of the solubility of sparingly soluble salts in solutions of salts with a common... 

Ksp, known as the solubility product, is widely used as a measure of the solubility of sparingly soluble salts. It should be noted that the dimensions of this constant will change according to the stoichiometry of the reaction. 

The increase in solubility of benzoic acid in concentrated aqueous solutions of alkali metal benzoate (salting-in) has been explained by a mechanism including homoconjugation. Generally, homoconjugation increases the solubility of sparingly soluble salts. If HA is... 

Our approach is similar to one by J. L. Guinon, J. Garcia-Anton, and V. Perez-Herranz, Spreadsheet Techniques for Evaluating the Solubility of Sparingly Soluble Salts of Weak Acids, J. Chem. Ed. 1999, 76, 1157. 

Solubility of Sparingly Soluble Salts, the "Common Ion" Effect and Le Chatelier s Principle... 

Solubility of Sparingly Soluble Salts of Weak Acids or Weak Bases... 

Solubility of sparingly soluble salts of weak acids in strong mineral acids The solubility product principle enables us to give a simple explanation of this phenomenon, which is of relatively frequent occurrence in quantitative analysis. Typical examples are the solubilities of calcium oxalate or barium carbonate in hydrochloric acid. When dilute hydrochloric acid is added to a suspension of calcium oxalate, the following equilibria will occur simultaneously ... 

A simple equation [(3-27) or (3-28)] may therefore be used in almost all practical situations involving titration end points. But, in the calculation of solubilities of sparingly soluble salts of polybasic weak acids, the concentration of anion may be so low that successive hydrolysis steps must be considered (Section 7-4). 

Relate the solubilities of sparingly soluble salts in water to their solubility product constants (Section 16.2, Problems 7-16). 

Determine the effect of complex-ion formation on the solubility of sparingly soluble salts involving a common cation (Section 16.6, Problems 47-48). 

This equation describes the relationship between the aqueous solubility of sparingly soluble salts (So) and the empirical Setschenow salting-out constant k = 0.217/Sq. This relationship and the Setschenow equation are valid only at low concentrations of added salt. As the concentration of added salt increases, the apparent k value is not constant, but is dependent on the solubility and the rate of change of solubility with added salt concentration. It was concluded that the Setschenow treatment is generally inappropriate for description and analysis of common ion equilibria. 

Another aspect of the effect of electrolytes on the solubility of a salt is the concept of the solubility product for poorly soluble substances. The experimental consequences of this phenomenon are that if the concentration of a common ion is high, then the other ion becomes low in a saturated solution of the substance, that is, precipitation occurs. Conversely, the effect of foreign ions on the solubility of sparingly soluble salts is just the opposite, and the solubility increases. This is called the salt effect. 

Heterogeneous reactions behave similarly. The solubility of sparingly soluble salts is an example of a heterogeneous reaction where equilibrium is set up rapidly. 

These arguments are standard for all reversible reactions, and discussion of what happens when reactants and products are mixed, and, in particular, the prediction of the direction of reaction, are individual for each reaction. The basic principles, however, are identical for all cases. Typical reversible reactions for electrolyte solutions involve acid-base equilibria, ion pair, complex and chelate formation, and heterogeneous reactions such as solubility of sparingly soluble salts. 

Data on the solubility of sparingly soluble salts can be listed either in terms of the solubility product, Ks, as defined above, or as the raw experimental quantity, the solubility, s. The solubility is expressed as a number of mol per dm. ... 

These can be illustrated by the use of emf measurements to find the equilibrium constants for weak acids and bases, for the self ionisation of water, for the formation of a complex or ion pair and for the solubility of sparingly soluble salts. This, taken with the situations described in the previous worked problems, illustrates the extreme versatility of emf studies. 

The self ionisation of water results in H30" "(aq) and OH (aq) always being present in any aqueous electrolyte solution. The measured resistance, or conductance, is therefore a composite quantity made up from the contribution made by the ions of the electrolyte and the H30 (aq) and OH (aq). The contribution from the HsO+Caq) and OH (aq) can be obtained by measuring the resistance, or conductance, of the H20(l) used in preparing the electrolyte solutions under study. The conductivity of the water can then be found from this measured conductance and the cell constant. This conductivity must then be subtracted from the conductivities found for each electrolyte solution. Values below 5 x 10 S cm are acceptable. This may seem to be a small contribution, but with accurate work it could represent a significant contribution especially if low concentration solutions are being studied, e.g. solubilities of sparingly soluble salts (see Section 11.15), or solutions such as very weak acids or bases (see Section 11.14). 

The use of radioisotopes in analysis (see also Section 16.3) includes determinations of solubilities of sparingly soluble salts and vapour pressures of rather involatile substances, and investigations of solid solution formation and adsorption of precipitates. 

Why is the method of isotope dilution analysis used to determine the solubility of sparingly soluble salts rather than a method depending upon mass determination ... 

For very dilute solutions at 298 K, the numerical value of a concentration in molkg is equal to that in moldm, and the solubilities of sparingly soluble salts (see below) are generally expressed in mol dm. ... 

A,p, known as the solubility product, is widely used as a measure of the solubility of sparingly soluble salts. It should be noted that the dimensions of this constant will change according to the stoichiometry of the reaction. Temperature effects on solubility products arc readily assessed as most solubility reactions are clearly seen as cndothcrinic and disorder increasing. Raising the temperature will thus Increase A, p together with the solubility of the solid. 

The solubility of sparingly soluble salts may also be affected by the presence of impurities. Kemmer [1988] gives as an example the solubility of CaCO in the presence of magnesium. If magnesium is precipitated along with CaCO the residual concentration of calcium in solution may be increased. The inclusion of other impurities, such as strontium, has also been shown to increase the solubility of CaCOj. The use of empirical data from process plant therefore, has to be treated with caution since the presence of impurities may have a significant effect on the fouling propensity under certain circumstances, and will depend on location. 

This agrees weU with the literature value for the for HCOOH. Conductance measurements may also be used to find the solubility of sparingly soluble salts and complexation equilibrium constants. 

The utility of the above relationships in estimating the solubility of sparingly soluble salts is discussed in section 3.9.3. 

Experimental determinations of the conducting properties of electrolyte solutions are important essentially in two respects. Firstly, it is possible to study quantitatively the effects of interionic forces, degrees of dissociation and the extent of ion-pairing. Secondly, conductance values may be used to determine quantities such as solubilities of sparingly soluble salts, ionic products of self-ionizing solvents, dissociation constants of weak acids and to form the basis for conductimetric titration methods. 

---