Precipitation and the Solubility Product

So far, we have considered only cases in which a single slightly soluble salt attains equilibrium with its component ions in water. The relative concentrations of the cations and anions in such solutions echo their relative numbers of moles in the original salt. Thus, when AgCl is dissolved, equal numbers of moles of Ag (aq) and Cl (aq) ions result, and when Ag2S04 is dissolved, twice as many moles of Ag (aq) ions as SO aq) are produced. A solubility product relationship such as 

FIGURE 16.6 Some solid silver chloride is in contact with a solution containing Ag (aq) and Cl (aq) ions. If a solubility equilibrium exists, then the product Q of the concentrations of the ions [Ag ] X [Cl ] is a constant, /Cjp (curved line). When Q exceeds K p, solid silver chloride tends to precipitate until equilibrium is attained. When Q is less than K p, additional solid tends to dissolve. If no solid is present, Q remains less than K p. 

Suppose a solution is prepared by mixing one soluble salt, such as AgN03, with a solution of a second, such as NaCl. Will a precipitate of very slightly soluble silver chloride form To reach an answer, the reaction quotient Q that was defined in connection with gaseous equilibria (see Section 14.6) is used. The initial reaction quotient Qo, when the mixing of the solutions is complete but before any reaction occurs, is 

IfQo JCgpj no solid silver chloride c n 3ppe3.r On the other h3.nd if 

An emulsion of silver chloride for photographic film is prepared by adding a soluble chloride salt to a solution of silver nitrate. Suppose 500 mL of a solution of CaCb with a chloride ion concentration of 8.0 X 10 M is added to 300 mL of a 0.0040 M solution of AgN03. Will a precipitate of AgCl(s) form when equilibrium is reached  

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The supersaturation, S, is typically defined as the ratio of the local instantaneous solute concentration to the equilibrium concentration or solubility in the fluid phase. In the case of precipitation supersaturation is instead defined as the ratio of the product of the concentration of ions constituting the precipitate and the solubility product in the fluid phase. 

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 ... 

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. ... 

Fluorides of the lantlmnides, yttrium, and scsmdium normally precipitate from aqueous media as the hemihydrates, LnF3-MsH20. It is these hemihydrates which are in equilibrium with saturated solution, and the solubility-product measurements refer to these hemihydrates. 

You can use the relationship between the ion product expression and the solubility product expression to predict whether a precipitate will form in a given system. One common system involves mixing solutions of two soluble ionic compounds, which react to form an ionic compound with a very low solubility. If Qsp > Kp. based on the initial concentrations of the ions in solution, the sparingly soluble compound will form a precipitate. 

Dissolution of Sparingly Soluble Salts. Obtain precipitates of calcium carbonate and calcium oxalate in test tubes by reacting the relevant salts. Decant the solutions and pour an acetic acid solution onto the moist precipitates. What happens Repeat the experiment, but use hydrochloric acid instead of the acetic acid. Write the equations of the chemical reactions in the molecular and net ionic forms. Explain the results obtained on the basis of the dissociation constants of the acids and the solubility product. 

By running a potentiometric precipitation titration, we can determine both the compositions of the precipitate and its solubility product. Various cation- and anion-selective electrodes as well as metal (or metal amalgam) electrodes work as indicator electrodes. For example, Coetzee and Martin [23] determined the solubility products of metal fluorides in AN, using a fluoride ion-selective LaF3 single-crystal membrane electrode. Nakamura et al. [2] also determined the solubility product of sodium fluoride in AN and PC, using a fluoride ion-sensitive polymer membrane electrode, which was prepared by chemically bonding the phthalocyanin cobalt complex to polyacrylamide (PAA). The polymer membrane electrode was durable and responded in Nernstian ways to F and CN in solvents like AN and PC. 

It has been observed that when an aqueous solution of pyrazine was added to an excess of silver nitrate solution, a precipitate of shiny white platelets formed immediately. If the order was reversed, however, and silver nitrate was added to an excess of pyrazine, the precipitate was formed only very slowly and after cooling. In each case, the precipitate was found to be Ag(pyrazine)N03 and the solubility product was determined as 2.3 x 10-4.103... 

Zinc sulphide does not precipitate from an acidified solution because the S ion concentration is repressed by the H+ ions of the strong acid, H2S 2H+ + S, and the solubility product of zinc sulphide cannot be reached. Acetate ions, however, remove H+ ions, and, the hydrogen sulphide thus being allowed to ionize to a greater extent, the solubility product of zinc sulphide is exceeded and the white precipitate appears. (See Solubility Product, page 131, and Experiment 22, page 175.)... 

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. 

From a knowledge of the solubility rules (see Section 4.2) and the solubility products listed in Table 16.2, we can predict whether a precipitate will form when we mix two solutions or add a soluble compound to a solution. This ability often has practical value. In industrial and laboratory preparations, we can adjust the concentrations of ions until the ion product exceeds K p in order to obtain a given compound (in the form of a precipitate). The ability to predict precipitation reactions is also useful in medicine. For example, kidney stones, which can be extremely painful, consist largely of calcium oxalate, CaC204 (K p = 2.3 X 10 ). The normal physiological concentration of calcium ions in blood plasma is about 5 mM (1 mM = 1 X 10 M). Oxalate ions ( 204 ), derived from oxalic acid present in many vegetables such as rhubarb and spinach, react with the calcium ions to form insoluble calcium oxalate, which can gradually build up in the kidneys. Proper adjustment of a patient s diet can help to reduce precipitate formation. Example 16.10 illustrates the steps involved in precipitation reactions. 

This test was performed in absence of any support. Dissolution of PdCl2 was obtained using 10" M HCl. According to literature [14], H2PdCl4 is formed in this conditions. As given by Chariot [15], the dissociation constant (pIQ) of PdCU is 13.2 and the solubility product (pKg) of Pd(OH)2 is 24. Thus precipitation of Pd-hydroxide should occur when pH becomes superior to 6.1. A rapid precipitation is indeed observed at pH 10 and room temperature. 

The pA ad and pAiy values (negative logarithms of the adsorption constant and the solubility product) are given in Table 4. In the case of Co and Ni it is difficult to separate the effects of precipitation and adsorption. 

From knowledge of the solubility rules (see Section 12.5) and the solubility products listed in Table 12.3, we can predict whether a precipitate will form when we mix two solutions or add a soluble compound to a solution. This ability often has practical value. 

Prom the point of view of chemistry, this kind of substances occur precipitation when its concentration in solution is larger than the solubility and the solubility product (Ksp). When the ion concentration of the system reaches the solubility product, the concentration of the OH and Ru + can be calculated according to the pH of the system. According to Eq. (6.21), the calculated concentration of the OH (i.e., pH value) and the lowest concentration of the Ru + which can generate Ru(OH)s precipitation are shown in Table 6.31. It is commonly considered that if the concentration of the Ru + is lower than 10 mol-L , it has been completely precipitated. It is seen from Table 6.31 that when the pH is fom, the concentration of Ru + ion in the solution has reached that of the completely precipitated. In fact, the pH of NHs H2O solution is higher than 11, the concentration of the Ru + ions is below 10 mol L . So the Ru + ions can be completely precipitated even in the acid environment and must be completely precipitated in the alkali environment. Therefore, it is a feasible method of precipitation which can load Ru + and remove the Cl simultaneity. 

Krishnamurthy et al. (1988b) observed, that R element type and concentration both affect the potential at which pseudopassivation occurs, and also the value of the current density in the pseudopassivation region. The magnitude of both these parameters will be determined by the nature of the R oxide formed, its ionic or electronic conductivity and the solubility product of the oxide which will govern the kinetics of precipitation. 

The precipitation of evaporites from marine and brine sources depends upon a number of faetors. Prominent among these are the concentrations of the evaporite ions in the water and the solubility products of the evaporite salts. The presence of a common ion decreases solubility for example, CaS04 precipitates more readily from a brine that eontains Na2S04 than it does from a solution that contains no other source of sulfate. The presence of other salts that do not have a common ion increases solubility because it decreases activity coefficients. Differences in temperature result in significant differences in solubility. 

In most precipitation titrations, silver nitrate solution is used and a silver salt, AgA, precipitates from solutions of an anion A. At the end-point, [Ag ]=[A ] and the solubility product S of AgA is given by ... 

Aqueous ammonia can also behave as a weak base giving hydroxide ions in solution. However, addition of aqueous ammonia to a solution of a cation which normally forms an insoluble hydroxide may not always precipitate the latter, because (a) the ammonia may form a complex ammine with the cation and (b) because the concentration of hydroxide ions available in aqueous ammonia may be insufficient to exceed the solubility product of the cation hydroxide. Effects (a) and (b) may operate simultaneously. The hydroxyl ion concentration of aqueous ammonia can be further reduced by the addition of ammonium chloride hence this mixture can be used to precipitate the hydroxides of, for example, aluminium and chrom-ium(III) but not nickel(II) or cobalt(II). 

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. 

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