Solubility

Solubility, concentration of dissolved substance in a saturated solution, i.e., in a solution that is in equilibrium with the same substance in the form of a separate phase. - Solubility of gases. In the case of contact between a liquid and a gas phase, the solubility of a gas shows a very strong dependence on its pressure. According to Henry's law, the amount of gas dissolving in a given liquid is proportional to the latter's pressure. Since, according to Boyle-Mariotte's law, pressure has the same effect on the gas content per unit volume of the gas phase itself, it can be said that the volume of gas dissolving in a liquid does not depend on its pressure. If we have not a pure gas but a gas mixture, then according to Dalton's law, the solubility of each gas is proportional to its partial pressure, regardless of other gases present in the mixture. The solubility of a gas is usually characterized by its "absorption coefficient." The latter equals the amount of gas that dissolves at a given temperature in 1 cm³ of liquid if its partial pressure equals 1 atm.; in this case, the amount of gas is expressed by the volume it occupies at 0° and atmospheric pressure. Real gases, not following the laws of ideal gases, also do not obey Henry-Dalton's law. This especially applies to gaseous substances having very high solubility, such as ammonia (NH₃) or hydrochloric acid (HCl). Their solubility is not proportional to partial pressure and therefore cannot be characterized by a constant absorption coefficient. With increasing temperature, the solubility of all gases decreases. Thus, for oxygen, the absorption coefficient at 20° is 3.10 × 10^-2, at 50° - 2.09 × 10^-2; for nitrogen, the corresponding values are 1.45 × 10^-2 and 1.02 × 10^-2. However small the amount of oxygen dissolving in water, it proves sufficient for the respiration of aquatic animals. The indicated dependence of gas solubility on temperature also has very great biological significance. With increasing temperature, the respiration of animals and their need for oxygen rapidly increase, while its content in water falls. Between these two quantities, a discrepancy arises, as a result of which the death of animals upon heating may actually depend on their suffocation due to lack of oxygen. In this case, it is possible to experimentally prove that saturation of water with oxygen makes animals less sensitive to temperature increase. For transporting absorbed oxygen from water or directly from air to their tissue elements, more highly organized animals do not limit themselves to dissolving oxygen in biological fluids, but resort to its labile binding with the help of hemoglobin or other substances acting in a similar manner. Solubility of liquids. Some liquids can dissolve or mix with each other in any proportions (e.g., sulfuric acid and water or alcohol and water); their solubility can be considered infinitely large. In other cases, it has for each temperature a definite, finite value. In contrast to gases, the solubility of liquids increases with increasing temperature. Liquids that partially dissolve each other at certain temperatures can, upon heating, acquire the ability to unlimited mixing. Thus, for example, phenol (carbolic acid) and water above 68.8° mix in any proportions. Upon cooling below this "critical" temperature, a mixture containing 36% phenol and 64% water by weight suddenly becomes cloudy and then separates into two phases: a saturated solution of water in phenol and a saturated solution of phenol in water. The lower the temperature, the more significantly do these two phases differ in composition, the wider the interval of concentrations at which a mixture of phenol with water gives a two-phase system. With increasing temperature, the solubility of water in phenol and of phenol in water increases, while the difference in composition of both solutions correspondingly decreases; at the critical point it becomes equal to zero. Solubility of solids is usually characterized by the number of grams of dry solid substance contained in 100 cm³ of saturated solution (or by the number of grams of solid substance that dissolves in 100 g of pure solvent). The solubility of solids, as well as that of liquids, usually increases with increasing temperature. The magnitude of solubility is determined by the equilibrium existing between the dissolved molecules of the solid and the molecules of the same substance present in the precipitate. The osmotic pressure of the dissolved molecules characterizes the tendency of the solid to pass into solution, its so-called "dissolving elasticity." More complex relationships arise in the case of dissolution of electrolytes, which in solution give molecules or particles of at least three kinds: a neutral molecule, an anion, and a cation. The solubility of an electrolyte equals the total amount passing into solution, regardless of whether it dissociates into ions or not. Meanwhile, only undissociated molecules (i.e., molecules of the same kind as those in the precipitate) are in direct equilibrium with the precipitate. Only their concentration (often negligibly small compared to the concentration of dissociation products) should represent a constant value for a given temperature. The concentration of dissolved molecules that have dissociated into ions is determined by dissociation equilibrium. According to the law of mass action, the product of dissociation products is proportional to the concentration of undissociated molecules. Since the latter remains constant in a saturated solution, the product of the concentrations of ions formed during dissociation is also a constant value at a given temperature. It has received the name "solubility product of ions." In the case of a binary electrolyte, if c₁ is the concentration of one ion, c₂ of the other, then their product c₁c₂=K is a constant value. Thus, the total solubility of the electrolyte consists of two quantities: the constant (for a given temperature) concentration of undissociated molecules and the concentration of ions determined by the "solubility product." Therefore, by decreasing the degree of dissociation (and along with it the concentration of ions) by adding a foreign electrolyte having a common ion with this electrolyte, one can correspondingly decrease the latter's solubility. Thus, for example, in a saturated solution of silver acetate, one can cause precipitation by adding silver nitrate. The increase in concentration of silver ions caused by this addition decreases the degree of dissociation of its acetate salt and causes the formation of an excess amount of undissociated molecules, which precipitate out. If the addition of an electrolyte with a common ion decreases dissociation and solubility, then the opposite effect is produced by binding one of the ions, forcing some amount of undissociated molecules to dissociate into ions, and an equal number of new molecules to pass from the precipitate into solution. This explains, for example, the easy solubility in acids of such a little water-soluble salt as acid potassium tartrate (KC₄H₄O₆), which dissociates into ions according to the equation: KC₄H₄O₆ ⇌ K⁺ + C₄H₄O₆^-. In an acid solution in the presence of an excess of H⁺ ions, the anions combine with them into molecules of weak, little-dissociated tartaric acid, and for the restoration of constancy of the ionic product [K⁺][C₄H₄O₆^-], an increasing amount of precipitate passes into solution and dissociates. Similarly behaves calcium oxalate, usually present in plant cells: CaC₂O₄ ⇌ Ca⁺⁺ + C₂O₄^--. In a neutral or alkaline solution it is very little soluble, but in an acidic environment the present oxalate ions are converted into relatively little-dissociated oxalic acid, thereby causing further dissolution of the precipitate. Such changes in solubility play an important role in the processes of normal and pathological precipitation and dissolution of salts in the organism. If an electrolyte is very little soluble, its saturated solution has such a low concentration that dissociation is practically complete. In such a case, the concentration of undissociated molecules is negligibly small, the concentration of each ion practically equals the molar concentration of the electrolyte, and the ionic product is calculated especially simply. It is given in the following table for several slightly soluble substances (at 18°): Electrolyte Solubility Ionic product K (in g·l^-1 solution) AgCl........ 1.42×10^-3 1.0×10^-10 BaSO₄ ......... 8.4×10^-6 1.0×10^-10 CaSO₄......... 2.3×10^-1 5.5×10^-5 AgBr........ 1.0×10^-5 4.5×10^-13 PbCl₂ ......... 1.5×10^-2 1.0×10^-4 CaCO₃......... 1.0×10^-3 4.3×10^-9 AgI........ 1.0×10^-8 2.0×10^-16 CaF₂ ......... 2.2×10^-4 1.8×10^-10 The ionic product of solubility retains a constant value only as long as the law of mass action, on the basis of which it is derived, is observed. As modern theory of activity shows (see Electrolytic dissociation), at sufficiently high total ionic concentration, the dissociation equilibrium of any electrolytes ceases to obey the law of mass action.

Therefore, the addition of a neutral salt or generally any strong electrolyte in high concentration contrary to the classical dissociation theory of Arrhenius exerts a noticeable influence on the solubility of electrolytes, somewhat increasing the ionic product (the so-called solubility anomaly). A very strong influence on solubility could be attributed to colloids: in many cases they apparently possess the ability to significantly increase solubility. For example, slightly soluble lime salts (calcium carbonate and calcium phosphate), which in a pure aqueous solution would be in a supersaturated state and would precipitate out, in the blood, in the presence of serum colloids, remain in solution. However, such a phenomenon by no means can be considered a true change in solubility. In reality, protective colloids (see Protective action) do not shift the equilibrium between the dissolved substance and the precipitate, but only more or less firmly stabilize the supersaturated solution and prevent the precipitation of the insoluble excess of the substance present in solution. Solubility depends not only on the properties of the substance being dissolved but to no lesser extent also on the nature of the solvent. Substances that are highly soluble in water may be poorly soluble in organic liquids (alcohol, ether, benzene, etc.) and vice versa. The modern theory of the structure of matter makes it possible to establish certain relationships between the chemical nature of a substance and its solubility in one or another medium. Substances having a polar molecular structure and approaching in this respect water itself (which is characterized by a large dipole moment and correspondingly high dielectric constant) dissolve well in it. Conversely, substances devoid of polarity or possessing it only to a slight degree show a strong affinity for non-aqueous solvents. This also applies to such non-aqueous organic media as lipoids, the solubility in which forms the basis of certain theories of cell permeability.