Slightly soluble salts

Solubilization. The solubiUty product of a slightly soluble salt determines the concentration of metal ion that can be present in solution with the anion of that salt. For the salt MX the solubiUty product is... 

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. 

FIGURE 11.15 If the concentration of one of the ions of a slightly soluble salt is increased, the concentration of the other decreases to maintain a constant value of Ksp. (a) The cations (pink) and anions (green) in solution, (b) When more anions are added (together with their accompanying spectator ions, which are not shown), the concentration of cations decreases. In other words, the solubility of the original compound is reduced by the presence of a common ion. In the insets, the blue background represents the solvent (water). 

The equilibrium constant is small, indicating a slightly soluble salt, as we would expect for a naturally occurring mineral like gypsum. 

Metal insoluble-salt These consist of a metal in contact with one of its slightly soluble salts this salt in turn is in contact with a solution containing the anion of the salt. An example is represented as Ag AgCl Or (c). The electrode process at such an electrode as AgCl (s) Ag + Cl" Ag + e- —> Ag (s) or overall, AgCl (s) + e- Ag (s) + Cl". The electrode reaction involves only the concentration of Cl" as a variable, in contrast with the Ag Ag electrode, which has the Ag concentration as a variable. The most frequently electrode of this type is the calomel electrode (see text for description). 

Plan (1) Calculate the molar solubility of the slightly soluble salt, which is the number of moles of the salt... 

If a slightly soluble salt solution is at equilibrium and we add a solution containing one of the ions involved in the equilibrium, the solubility of the slightly soluble salt decreases. For example, consider the PbS04 equilibrium ... 

Suppose we add a solution of Na2S04 to this equilibrium system. The additional sulfate ion will disrupt the equilibrium by Le Chatclier s principle and shift it to the left. This decreases the solubility. The same would be true if you tried to dissolve PbS04 in a solution of Na2S04 instead of pure water—the solubility would be less. This application of Le Chatelier s principle to equilibrium systems of a slightly soluble salt is the common-ion effect. 

The equilibrium constant expression associated with systems of slightly soluble salts is the solubility product constant, Ksp. It is the product of the ionic concentrations, each one raised to the power of the coefficient in the balanced chemical equation. It contains no denominator since the concentration of a solid is, by convention, 1, and for this reason it does not appear in the equilibrium constant expression. The Ksp expression for the PbS04 system is ... 

Knowing the value of the solubility product constant can also allow us to predict whether or not a precipitate will form if we mix two solutions, each containing an ion component of a slightly soluble salt. We calculate the reaction quotient (many times called the ion product), which has the same form as the solubility product constant. We take into consideration the mixing of the volumes of the two solutions, and then compare this reaction quotient to the K.p. If it is greater than the Ksp then precipitation will occur until the ion concentrations reduce to the solubility level. 

The common-ion effect is an application of Le Chatelicr s principle to equilibrium systems of slightly soluble salts. A buffer is a solution that resists a change in pH if we add an acid or base. We can calculate the pH of a buffer using the Henderson-Hasselbalch equation. We use titrations to determine the concentration of an acid or base solution. We can represent solubility equilibria by the solubility product constant expression, Ksp. We can use the concepts associated with weak acids and bases to calculate the pH at any point during a titration. 

Reference Electrodes By definition, the normal hydrogen electrode (N H E) is the reference for electrode potentials (see Sect. 2.3.2.1), but practically it is scarcely usable. A reference electrode (RE) has to provide a well-defined potential between the electrolyte and its electric connector, joined with the input of the measuring instrument. Usually, a metal and a slightly soluble salt of this metal is applied (secondary electrode) [76, 77]. The electrolyte in the RE is connected to the electrolyte in the electrochemical cell via a diaphragm, which has to separate both electrolytes, as far as possible without a potential difference (see below). 

A) The reaction that produced the precipitate is Pb (a,) + 2C1 (aq) PbCl2(,). Lead chloride is a slightly soluble salt, with a solubility of 10 g/L at 20°C. The solubility of PbCF increases very rapidly as the temperature rises. At 100°C it has a solubility of 33.5 g/L. However, PbCl2 precipitates very slowly, particularly when other ions that form insoluble chlorides are not present. PbCb dissolves in excess chloride ion as a result of the formation of a complex ion, tetrachloroplumbate(II) ion ... 

In brief, dissolution of slightly soluble salts in water is an equilibrium process. 

A precipitate forms when a cation and an anion bind to form a slightly soluble salt. Precipitation reactions are equilibrium reactions. For example when... 

Since the dissolution of a slightly soluble salt in water is an equilibrium, an equilibrium expression can be written. This expression is known as the solubility product. The constant for the expression is named the solubility product constant and denoted by K p. For example, the solubility product for the reaction below ... 

Write the mass balance for a saturated solution of the slightly soluble salt Ag3P04, which produces PO and 3Ag+ when it dissolves. 

Consider saturated solutions of the slightly soluble salts AgBr and BaC03. 

To prevent the precipitation of a slightly soluble salt, some substance must be added which will keep the concentration of one of the ions so low that the solubility product of the salt is not reached. 

Electrodes classified as second-class electrode systems are those in which the electrode is in direct contact with a slightly soluble salt of the electroactive species such that the potentiometric response is indicative of the concentration of the inactive anion species. Thus the silver/silver-chloride electrode system, which is representative of this class of electrodes, gives a potential response that is directly related to the logarithm of the chloride ion activity... 

Third-class electrodes are really a specialized case of second-class electrodes. They consist of the metal being in direct contact with a slightly soluble salt of the metal, which is then used to monitor the activity of an electroinactive metal ion in equilibrium with a more soluble salt that includes the same anion as the electrode-salt system. For example, the concentration of calcium ions in equilibrium with solid calcium oxalate may be monitored using a silver/ silver oxalate electrode system. The concentration of calcium ion affects the concentration of oxalate ion, which in turn controls the concentration of silver... 

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 ... 

Equilibrium constant expressions for gases and slightly soluble salts... 

Consider the chemical equation for AgCl dissolved in water to make a saturated solution AgCl(s) <—> Ag1+(aq) + Cl1 (aq). At 298 K the solubility product constant is 1.8 x 10-10, which indicates that is a slightly soluble salt. There is a way of making AgCl even less soluble, via the common ion effect. Consider the following, when an ion that is already present is added to the solution, the equilibrium will shift to consume the increase in concentration of the ion. 

The mass action expression for this slightly soluble salt is... 

Solubility Solubility Product Constant The ability of a substance to dissolve in another substance. The equilibrium constant of a slightly soluble salt. 

In spite of its limitations (as outlined in the previous section) the solubility product relation is of great value in qualitative analysis, since with its aid it is possible not only to explain but also to predict precipitation reactions. The solubility product is in reality an ultimate value which is attained by the ionic product when equilibrium has been established between the solid phase of the slightly soluble salt and the solution. If conditions are such that the ionic product is different from the solubility product, the system will seek to adjust itself in such a manner that the ionic product attains the value of the solubility product. Thus, if the ionic product is arbitrarily made greater than the solubility product, for example by the addition of another salt with a common ion, the adjustment of the system results in the precipitation of the solid salt. Conversely, if the ionic product is made smaller than the solubility product, as, for instance, by diminishing the concentration of one of the ions, equilibrium in the system is attained by some of the solid salt passing into solution. 

Chantooni and Kolthoff " derived equations which permit the calculation of hydration constants of cations and anions from the solubility products of slightly soluble salts in solutions of acetonitrile with various concentrations of water. The ionic solubility of a salt was determined by measuring the conductance. The water concentration of the acetonitrile solution was always less than 1 M. The total ionic solubility product was expanded in powers of the water concentration. The coefficients are related to the individual ionic hydration constants and were evaluated by... 

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