Solubility product Ksp

LAS is moderately sensitive to water hardness such that at certain concentration levels governed by the solubility product (Ksp) and the CMC, Ca(LAS)2 will precipitate out of solution [26]. The precipitation boundary diagram provides a useful method to study the interaction of LAS with calcium ions. From this diagram, approximate values of the CMC and Ksp for LAS and Ca2+ can be approximated [27]. 

The equilibrium constant for the solubility equilibrium between an ionic solid and its dissolved ions is called the solubility product, Ksp, of the solute. For example, the solubility product for bismuth sulfide, Bi2S3, is defined as... 

Sometimes it is important to know under what conditions a precipitate will form. For example, if we are analyzing a mixture of ions, we may want to precipitate only one type of ion to separate it from the mixture. In Section 9.5, we saw how to predict the direction in which a reaction will take place by comparing the values of J, the reaction quotient, and K, the equilibrium constant. Exactly the same techniques can be used to decide whether a precipitate is likely to form when two electrolyte solutions are mixed. In this case, the equilibrium constant is the solubility product, Ksp, and the reaction quotient is denoted Qsp. Precipitation occurs when Qsp is greater than Ksp (Fig. 11.17). 

The equilibrium constant for this reaction is actually the solubility product, Ksp = [Ag+][C1 ], for silver chloride (Section 11.8). 

As a new rule of thumb [473], in 0.15 M NaCl (or KC1) solutions titrated with NaOH (or KOH), acids start to precipitate as salts above log (S/So) 4 and bases above log (S/.S o) 3. It is exactly analogous to the cliff 3-4 rule let us call the solubility equivalent the sdijf3-4 rule [473], Consider the case of the monopro-tic acid HA, which forms the sodium salt (in saline solutions) when the solubility product Ksp is exceeded. In additions to Eqs. (3.1) and (6.1), one needs to add the following reaction/equation to treat the case ... 

The insolubility of calcium carbonate is clearly evident from the value of the solubility product, Ksp, in water at 25°C Ksp = 8.7 x 10-9. The carbonate ions are produced in seawater by the dissociation of carbonic acid that forms from the... 

Silver chloride is fairly insoluble (see p. 332), with a solubility product Ksp of 1.74 x 1CT10 mol2 dm-6. Its concentration in pure distilled water will, therefore, be 1.3 x 10-5 mol dm-3, but adding magnesium sulphate to the solution increases it solubility appreciably see Figure 7.10. 

The lanthanide fluorides Eire all sparingly soluble in water, as are the fluorides of yttrium and scEmdiiun. Thus, solubility information is generally presented in the form of the solubility product Ksp. 

Solubility data are presented for practically all entries. Quantitative data are also given for some compounds at different temperatures. In general, ionic substances are soluble in water and other polar solvents while the non-polar, covalent compounds are more soluble in the non-polar solvents. In sparingly soluble, slightly soluble or practically insoluble salts, degree of solubility in water and occurrence of any precipitation process may be determined from the solubility product, Ksp, of the salt. The smaller the Ksp value, the less its solubility in water. 

A central concept necessary to understanding the mechanisms of CD is that of the solubility product (Ksp). The solubility product gives the solubility of a sparingly soluble ionic salt (this includes salts normally termed insoluble ). Consider a very sparingly soluble salt (say, CdS) in equilibrium with its saturated aqueous solution ... 

Thus, if is large enough, for a given solubility product Ksp, and the pH and free EDTA concentration are appropriate, the solid M(OH)m of reaction 13.13 will pass entirely into solution (see Exercise 13.6). Since the donor atoms in these ligands are hard oxygen, they are particularly effective against hard Mm+ ions such as Ca2+ or Fe3+. 

The Mn3+/Mn2+ couple is ostensibly pH independent, but we must bear in mind that manganese(III), in particular, will hydrolyze unless the acidity is very high, and both Mn(OH)3 and Mn(OH)2 will come out of solution in alkaline media. Small degrees of hydrolysis in solution have little impact on E°, but precipitation of hydroxides (and all metal hydroxides except those of the alkali metals and Ca, Sr, Ba, and Ra are poorly soluble in water) affects E° profoundly. Thus, in alkaline media, the electrode potentials (often called Eh by geologists, wherever [H+] is not the standard value) of Mn2+ and Mn3+ are controlled by the solubility products (Ksp) of Mn(OH)2 and Mn(OH)3, respectively. In practice, Mn(OH)3 tends to dehydrate to MnO(OH), so we consider the Mn2+(aq)/Mn(s) couple ... 

The equilibrium constant for this reaction is actually the solubility product, Ksp, for silver chloride (Section 11.10). It does not matter that overall the reaction is not a redox reaction so long as it can be expressed as the differ- ence of two reduction half-reactions. Because silver chloride is almost insol-i uble, we expect K to be very small (and E° to be negative). 

The numerical value of a solubility product Ksp is measured by experiment. For example, we could determine Ksp for CaF2 by adding an excess of solid CaF2 to water, stirring the mixture to give a saturated solution of CaF2, and then... 

The solubility product, Ksp, for an ionic compound is the equilibrium constant for dissolution of the compound in water. The solubility of the compound and Ksp are related by the equilibrium equation for the dissolution reaction. The solubility of an ionic compound is (1) suppressed by the presence of a common ion in the solution (2) increased by decreasing the pH if the compound contains a basic anion, such as OH-, S2-, or CO32- and (3) increased by the presence of a Lewis base, such as NH3, CN-, or OH-, that can bond to the metal cation to form a complex ion. The stability of a complex ion is measured by its formation constant, Kf. 

The intrinsic solubility of the charged form of an organic compound is governed by its solubility product, Ksp, described here (Anderson and Conradi, 1985) fora 1 1 salt of a monoprotic acid, HA ... 

The zero-point G0 depends on the solvent and conceivably on other variables. Thus, the distribution coefficient K between two phases correspond to the species being exp[(< oi — G02)/RT] times more soluble in the phase 2 than in the phase 1. The solubility product Ksp is an expression of G0 in the solid compared with G0 in the solution, if we accept to speak about free energies of ions and not only of neutral species. Since Arrhenius convinced his reluctant contemporaries about the predominant (or complete) dissociation of salts in aqueous solution, it became acceptable to write complex formation constants involving concentrations of ions and not only of neutral molecules. 

Quadratic equations arise frequently in the mathematical descriptions of common physical and chemical processes. For instance, silver chloride is only very slightly soluble in water. It has been determined experimentally that the solubility product Ksp of silver chloride at 25C is 1.56 x 10 12 M2, meaning that in a saturated solution the concentrations of silver ion and chloride ion satisfy the relationship... 

We could also calculate the solubility from the solubility product, Ksp. Suppose the solubility of solid AgCl in water is taken to be s mol dm-3, then on dissolution to form a saturated solution, then equal concentrations ofs mol dm 3 of Ag+ ions and s mol dm-3 of Cl- ions enter the solution ... 

As shown in Equation (2.164b), the solubility of slightly soluble substances is simply related to the solubility product (Ksp). However, the solubility is dependent on the ionic strength of a solution. A slightly soluble salt dissolves in water ... 

Solubility, which refers to the amount of a substance that can be dissolved in solution, differs from the solubility product. Ksp for a substance remains constant as long as temperature is constant, while the solubility of a substance is subject to change. 

Still other conventional equilibrium constants, such as the solubility product Ksp for electrolytes and the ion product Kw = [H+(a )][OH (a )] for water, can be defined by absorbing near-constant activities (such as those of pure solids or liquids) into the constant K. Although such modified eq-type expressions may have practical utility, their... 

The equilibrium solubilities are deflned by the respective solubility products, Ksp (values from ref. 30). 

I". Cadmium sulfide (CdS) is used in some semiconductor applications. Calculate the value of the solubility product (Ksp) for CdS given the following standard reduction potentials ... 

However, factors such as the solubility product (Ksp) of the salt, common-ion effects, and hygroscopi-city may disfavor the salts of first choice using the above criterion, e.g., hydrochloride or sodium salts. The formation of hydrochloride salts does not always enhance solubility above that of the free base. The lower solubility of a hydrochloride salt in dilute HCl, relative to that of the free base, is attributed to the common ion effect of the chloride ion on the solubility product equilibrium of the salts. The common-ion effect suppresses the solubility product equilibrium. This is particularly relevant to the HCl salts of drugs administered orally, resulting in contact with... 

The equilibrium of a solvation reaction has its own equilibrium constant called the solubility product Ksp. Use Ksp the same way you would use any other equilibrium constant. Remember that solids and pure liquids have an approximate mole fraction of one and can be excluded from the equilibrium expression. Thus, solids are left out of the solubility product expression as in the example of the Ksp for barium hydroxide shown below. 

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