Salts solubility product

Solubility product — is the equilibrium constant Ksp of dissolution of a salt. Solubility products can be determined by direct determination of the -> concentrations of the dissolved salt, provided the activity constants are practically 1.0, or otherwise known. Solubility products can also be calculated from the standard -> Gibbs energies of formation AfG" of the species. 

Acid dissociation Aquated metal ions Amphoteric behaviour Solubilities of salts Solubility product constants... 

These are practically insoluble in water, are not hydrolysed and so may be prepared by addition of a sufficient concentration of sulphide ion to exceed the solubility product of the particular sulphide. Some sulphides, for example those of lead(II), copper(II) and silver(I), have low solubility products and are precipitated by the small concentration of sulphide ions produced by passing hydrogen sulphide through an acid solution of the metal salts others for example those of zincfll), iron(II), nickel(II) and cobalt(II) are only precipitated when sulphide ions are available in reasonable concentrations, as they are when hydrogen sulphide is passed into an alkaline solution. 

The sodium salt of methyl red may be prepared by dissolving the crude product in an equal weight of 35 per cent, sodium hydroxide which has been diluted to 350 ml., hitoring, and evaporating under diminished pressure (Fig. II, 37, I). The resulting sodium salt forms orange leaflets. This water-soluble product is very convenient for use as an indicator. Incidentally, the toluene extraction is avoided. 

Eor a sparingly soluble salt, the product of the total molecular concentrations of the ions is a constant for a given temperature. Thus for the dissolution of the electrolyte. 

The sodium salt of CS [9005-22-5] is prepared by reaction of cellulose with sulfuric acid in alcohol followed by sodium hydroxide neutrali2ation (20). This water-soluble product yields relatively stable, clear, and highly viscous solutions. Introduced as a thickener for aqueous systems and an emulsion stabilizer, it is now of no economic significance. 

The colloidal palladium solution is prepared as follows A solution of a palladium salt is added to a solution of an alkali salt of an acid of high molecular weight, the sodium salt of protalbinic acid being suitable. An excess of alkali dissolves the precipitate formed, and the solution contains tine palladium in the form of a hydrosol of its hydroxide. The solution is purified by dialysis, and the hydroxide reduced with hydrazine hydrate. On further dialysis and evaporation to dryness a water-soluble product is obtained, consisting of colloidal palladium and sodium protalbinate, the latter acting as a protective colloid. 

The Institut Fran ais du Petrole has developed and commercialized a process, named Dimersol X, based on a homogeneous catalyst, which selectively produces dimers from butenes. The low-branching octenes produced are good starting materials for isononanol production. This process is catalyzed by a system based on a nickel(II) salt, soluble in a paraffinic hydrocarbon, activated with an alkylalumini-um chloride derivative directly inside the dimerization reactor. The reaction is sec-... 

For sparingly soluble salts (i.e. those of which the solubility is less than 0.01 mol per L) it is an experimental fact that the mass action product of the concentrations of the ions is a constant at constant temperature. This product Ks is termed the solubility product . For a binary electrolyte ... 

It is important to note that the solubility product relation applies with sufficient accuracy for purposes of quantitative analysis only to saturated solutions of slightly soluble electrolytes and with small additions of other salts. In the presence of moderate concentrations of salts, the ionic concentration, and therefore the ionic strength of the solution, will increase. This will, in general, lower the activity coefficients of both ions, and consequently the ionic concentrations (and therefore the solubility) must increase in order to maintain the solubility product constant. This effect, which is most marked when the added electrolyte does not possess an ion in common with the sparingly soluble salt, is termed the salt effect. 

The great importance of the solubility product concept lies in its bearing upon precipitation from solution, which is, of course, one of the important operations of quantitative analysis. The solubility product is the ultimate value which is attained by the ionic concentration product when equilibrium has been established between the solid phase of a difficultly soluble salt and the solution. If the experimental conditions are such that the ionic concentration product is different from the solubility product, then the system will attempt to adjust itself in such a manner that the ionic and solubility products are equal in value. Thus if, for a given electrolyte, the product of the concentrations of the ions in solution is arbitrarily made to exceed the solubility product, as for example by the addition of a salt with a common ion, the adjustment of the system to equilibrium results in precipitation of the solid salt, provided supersaturation conditions are excluded. If the ionic concentration product is less than the solubility product or can arbitrarily be made so, as (for example) by complex salt formation or by the formation of weak electrolytes, then a further quantity of solute can pass into solution until the solubility product is attained, or, if this is not possible, until all the solute has dissolved. 

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... 

If the dissociation constant of the acid HA is very small, the anion A- will be removed from the solution to form the undissociated acid HA. Consequently more of the salt will pass into solution to replace the anions removed in this way, and this process will continue until equilibrium is established (i.e. until [M + ] x [A-] has become equal to the solubility product of MA) or, if sufficient hydrochloric acid is present, until the sparingly soluble salt has dissolved completely. Similar reasoning may be applied to salts of acids, such as phosphoric(V) acid (K1 = 7.5 x 10-3 mol L-1 K2 = 6.2 x 10-8 mol L-1 K3 = 5 x 10 13 mol L-1), oxalic acid (Kx = 5.9 x 10-2 mol L-K2 = 6.4 x 10-5molL-1), and arsenic)V) acid. Thus the solubility of, say, silver phosphate)V) in dilute nitric acid is due to the removal of the PO ion as... 

Factor 1, which is concerned with the completeness of precipitation, has already been dealt with in connection with the solubility-product principle, and the influence upon the solubility of the precipitate of (i) a salt with a common ion, (ii) salts with no common ion, (iii) acids and bases, and (iv) temperature (Sections 2.6-2.11). 

Solutions which prevent the hydrolysis of salts of weak acids and bases. If the precipitate is a salt of weak acid and is slightly soluble it may exhibit a tendency to hydrolyse, and the soluble product of hydrolysis will be a base the wash liquid must therefore be basic. Thus Mg(NH4)P04 may hydrolyse appreciably to give the hydrogenphosphate ion HPO and hydroxide ion, and should accordingly be washed with dilute aqueous ammonia. If salts of weak bases, such as hydrated iron(III), chromium(III), or aluminium ion, are to be separated from a precipitate, e.g. silica, by washing with water, the salts may be hydrolysed and their insoluble basic salts or hydroxides may be produced together with an acid ... 

I. Sodium tetraphenylborate Na+ [B(C6H5)4] . This is a useful reagent for potassium the solubility product of the potassium salt is 2.25 x 10 8. Precipitation is usually effected at pH 2 or at pH 6.5 in the presence of EDTA. Rubidium and caesium interfere ammonium ion forms a slightly soluble salt and can be removed by ignition mercury(II) interferes in acid solution but does not do so at pH 6.5 in the presence of EDTA. 

Two procedures have been developed for the aminohydroxylation of a, 3-unsat-urated amides Procedure A for products that are insoluble in the reaction mixture and Procedure B for soluble products (Scheme 12.17) [48]. These differ only in that the former requires a 10-25% excess of chloramine-T and t-BuOH as the cosolvent, while the latter uses only one equivalent of the chloramine salt and MeCN as the cosolvent. The excess of chloramine-T in Procedure A allows better turnover near the end of the reaction, and the trace amount of p-toluenesulfonamide byproduct can be removed by recrystallization. However, elimination of the necessity to remove p-toluenesulfonamide far outweighed the inconvenience of slightly longer reaction times needed in procedure B without the use of excess chloramine salt. 

Other useful solid-state electrodes are based on silver compounds (particularly silver sulfide). Silver sulfide is an ionic conductor, in which silver ions are the mobile ions. Mixed pellets containing Ag2S-AgX (where X = Cl, Br, I, SCN) have been successfiilly used for the determination of one of these particular anions. The behavior of these electrodes is determined primarily by the solubility products involved. The relative solubility products of various ions with Ag+ thus dictate the selectivity (i.e., kt] = KSp(Agf)/KSP(Aw)). Consequently, the iodide electrode (membrane of Ag2S/AgI) displays high selectivity over Br- and Cl-. In contrast, die chloride electrode suffers from severe interference from Br- and I-. Similarly, mixtures of silver sulfide with CdS, CuS, or PbS provide membranes that are responsive to Cd2+, Cu2+, or Pb2+, respectively. A limitation of these mixed-salt electrodes is tiiat the solubility of die second salt must be much larger than that of silver sulfide. A silver sulfide membrane by itself responds to either S2- or Ag+ ions, down to die 10-8M level. 

Solid Bi2S3 does not appear in the expression for K,p, because it is a pure solid and its activity is 1 (Section 9.2). A solubility product is used in the same way as any other equilibrium constant. However, because ion-ion interactions in even dilute electrolyte solutions can complicate its interpretation, a solubility product is generally meaningful only for sparingly soluble salts. Another complication that arises when dealing with nearly insoluble compounds is that dissociation of the ions is rarely complete, and a saturated solution of Pbl2, for instance, contains substantial... 

STRATEGY First, we write the chemical equation for the equilibrium and the expression for the solubility product. To evaluate Ksp, we need to know the molarity of each type of ion formed by the salt. We determine the molarities from the molar solubility, the chemical equation for the equilibrium, and the stoichiometric relations between the species. We assume complete dissociation. 

Sei.f-Tfst 11.11A The solubility product of silver sulfate, Ag2S04, is 1.4 X 10 Estimate the molar solubility of the salt. 

The solubility product is the equilibrium constant for the equilibrium between an undissolved salt and its ions in a saturated solution. 

FIGURE 11.17 The relative magnitudes of the solubility quotient, and the solubility product, K y are used to decide whether a salt will dissolve (left) or precipitate (right). When the concentrations of the ions are low (left) Qsp is smaller than (Csp when the ion concentrations are high (right), Qsp is larger than K. ... 

The values for the solubility products of some sparingly soluble salts are listed in Table 11.4. 

In arid environments, where the soluble products of weathering are not completely removed from the soil, saline solutions may circulate in the soil as well as in rock fractures. If upon evaporation the salt concentration increases above its saturation point, salt crystals form and grow (Goudie et al, 1970). The growth of salt crystals in crevices can force open fractures. Salt weathering occurs in cold or hot deserts or areas where salts accumulate. Boulders, blocks. 

The solubility product (. sp) describes the equilibrium of a salt dissolving in water. In the laboratory and in industry, solubility equilibria are often exploited in the opposite direction. Two solutions are mixed to form a new solution in which the solubility product of one substance is exceeded. This salt precipitates and can be collected by filtration. Example illustrates how precipitation techniques can be used to remove toxic heavy metals from aqueous solutions. 

Almost all the Earth s carbon is found in the lithosphere as carbonate sediments that have precipitated from the oceans. Shells of aquatic animals also contribute CaC03 to the lithosphere. Carbon returns to the hydrosphere as carbonate minerals dissolve in water percolating through the Earth s crust. This process is limited by the solubility products for carbonate salts, so lithospheric carbonates represent a relatively inaccessible storehouse of carbon. 

C18-0073. For the following salts, write a balanced equation showing the solubility equilibrium and write the solubility product expression for each (a) silver chloride (b) barium sulfate (c) iron(H) hydroxide and (d) calcium phosphate. 

Depending on electrolyte composition, the metal will either dissolve in the anodic reaction, that is, form solution ions [reaction (1.24)], or will form insoluble or poorly soluble salts or oxides precipitating as a new solid phase next to the electrode surface [reaction (1.28)]. Reacting metal electrodes forming soluble products are also known as electrodes of the first kind, and those forming solid products are known as electrodes of the second kind. 

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