Solubility equilibria cations

When a precipitate has been formed during the qualitative analysis of the ions present in a solution, it may be necessary to dissolve the precipitate again to identify the cation or anion. One strategy is to remove one of the ions from the solubility equilibrium so that the precipitate will continue to dissolve in a fruitless chase for equilibrium. Suppose, for example, that a solid hydroxide such as iron(IIl) hydroxide is in equilibrium with its ions in solution ... 

All sodium salts are soluble, and so are all nitrate salts, so It makes sense that neither of these ions participates in a solubility equilibrium. Furthermore, nitrate and sodium cations are neither acidic nor basic, so it makes sense that neither participates in an acid-base equilibrium. 

An ion pair is a close association of a cation and an anion in solution, whereas the ion product is the value obtained when the initial concentrations for the dissolved ions involved in the solubility equilibrium are inserted into the equilibrium constant expression. 

STRATEGY As usual, we begin by writing the chemical equation for the solubility equilibrium and the expression for Ksp. The molar solubility is the molarity of formula units in the saturated solution. Because each formula unit produces a known number of cations and anions in solution, we can express the molarities of the cations and anions in terms of s. Then we express Ksp in terms of s and solve for s. Assume complete dissociation. 

The formation of a complex can also remove an ion and disturb the solubility equilibrium until more solid dissolves. We first met complexes in Section 2.13, where we saw that they were species formed by the reaction of a Lewis acid and a Lewis base. In this section, we consider complexes in which the Lewis acid is a metal cation, such as Ag+. An example is the formation of Ag(NH3)2+ when an aqueous solution of the Lewis base ammonia is added to a solution of silver bromide ... 

The solubility of an ionic compound increases dramatically if the solution contains a Lewis base that can form a coordinate covalent bond (Section 7.5) to the metal cation. Silver chloride, for example, is insoluble in water and in acid, but it dissolves in an excess of aqueous ammonia, forming the complex ion Ag(NH3)2 + (Figure 16.13). A complex ion is an ion that contains a metal cation bonded to one or more small molecules or ions, such as NH3, CN-, or OH-. In accord with Le Chatelier s principle, ammonia shifts the solubility equilibrium to the right by tying up the Ag+ ion in the form of the complex ion ... 

In Section 16.15, we ll see that mixtures of metal cations, M2+, can be separated into two groups by selective precipitation of metal sulfides, MS. For example, Pb2+, Cu2+, and Hg2+, which form very insoluble sulfides, can be separated from Mn2+, Fe2+, Co2+, Ni2+, and Zn2+, which form more soluble sulfides. The separation is carried out in an acidic solution and makes use of the following solubility equilibrium ... 

These electrodes are based on two equilibria the electrochemical equilibrium involving formation of the interfacial potential and the solubility equilibrium between the cation and its sparsely soluble salt. The most popular electrode of this type is the silver/silver chloride electrode. The electrochemical equilibrium is the same as for the Ag/Ag+ electrode described above (6.27) and the solubility equilibrium is... 

When the metal ion of an insoluble salt forms a complex ion, the aquo cation is removed from solution, shifting the solubility equilibrium toward solution species. 

Colloidal dispersions owe their stability to a surface charge and the resultant electrical repulsion of charged particles. This charge is acquired by adsorption of cations or anions on the surface. For example, an ionic precipitate placed in pure water will reach solubility equilibrium as determined by its solubility product, but the solid may not have the same attraction for both its ions. Solid silver iodide has greater attraction for iodide than for silver ions, so that the zero point of charge (the isoelectric point) corresponds to a silver ion concentration much greater than iodide, rather than to equal concentrations of the two ions. The isoelectric points of the three silver halides are ° silver chloride, pAg = 4, pCl = 5.7 silver bromide, pAg = 5.4, pBr = 6.9 silver iodide, pAg = 5.5, pi = 10.6. For barium sulfate the isoelectric point seems to be dependent on the source of the product and its de ee of perfection. ... 

Figure 7.1. Solubility of simple salts as a function of the common anion concentration (Example 7.2). The cations and anions of these salts do not protolyze in the neutral pH range. The equilibrium solubility is given by the metal-ion concentration. At high anion or cation concentration, complex formation or ion-pair binding becomes possible (dashed lines). If the salt is dissolved in pure water (or in an inert electrolyte), the solubility is defined by the electroneutrality z[Me" J = /i[anion ]. If z = n (e.g., BaS04), the solubility is given by the intersection (-I-). If z the electroneutrality condition is fulfilled at a point slightly displaced from the intersection (t). The insert exemplifies the solubility equilibrium for Cap2 ( o = 10" ) and lists the domains of over- and undersaturation.

However, the case in which the solubility of a solid can be calculated from the known analytical concentration of added components and from the solubility product alone is very seldom encountered. Ions that have dissolved from a crystalline lattice frequently undergo chemical reactions in solution, and therefore other equilibria in addition to the solubility product have to be considered. The reaction of the salt cation or anion with water to undergo acid-base reactions is very common. Furthermore, complex formation of salt cation and salt anion with each other and with one of the constituents of the solution has to be considered. For example, the solubility of FeS(s) in a sulfide-containing aqueous solution depends on, in addition to the solubility equilibrium, acid-base equilibria of the cation (e.g., Fe + H2O = FeOH + H ) and of the anion (e.g., S + HjO = HS + OH, and HS" + H2O = H2S + OH ), as well as on equilibria describing complex formation (e.g., formation of FeHS" or FeSi ). 

Formed by the reaction of a metal cation, acting as a Lewis acid, with a Lewis base May disturb a solubility equilibrium by reducing the concentration of metal ions -> Net result is an increase in the solubility of salts. 

Analytical 13 marks Solubility equilibrium Equilibria of sparingly soluble salts, precipitation titrations and calculations of cation concentration during the course of titrations and indicators used in precipitation titrations... 

Due to the dynamic equilibrium between the atmospheric carbon dioxide and the oceanic bicarbonate and carbonate anions, the greatest amount of soluble calcium cations is contained in the ocean. This mass is four orders of magnitude higher than the total mass of bound calcium in living and dead matter of both terrestrial and aquatic organisms. The average calcium content in the seawater is 408 mg/L, and the overall pool is 559 x lO tons. 

The cations in many slighdy soluble compounds can form complex war.This often results in dissolution of the slightly soluble compound. Some metal ions share electron pairs donated by molecules and ions such as NH3, CN, OH, F, Cl, Br, and T to form coordinate covalent bonds to metal ions. Coordinate covalent bonds are formed as these electron-donating groups (ligands) replace H2O molecules from hydrated metal ions. The decrease in the concentration of the hydrated metal ion shifts the solubility equilibrium to the right. 

Briefly describe each of the following ideas, methods, or phenomena (a) common-ion effect in solubility equilibrium (b) fractional precipitation (c) ion-pair formation (d) qualitative cation analysis. 

Potentiometric electrodes are divided into two classes metallic electrodes and membrane electrodes. The smaller of these classes are the metallic electrodes. Electrodes of the first kind respond to the concentration of their cation in solution thus the potential of an Ag wire is determined by the concentration of Ag+ in solution. When another species is present in solution and in equilibrium with the metal ion, then the electrode s potential will respond to the concentration of that ion. Eor example, an Ag wire in contact with a solution of Ck will respond to the concentration of Ck since the relative concentrations of Ag+ and Ck are fixed by the solubility product for AgCl. Such electrodes are called electrodes of the second kind. 

The data given in Tables 1.9 and 1.10 have been based on the assumption that metal cations are the sole species formed, but at higher pH values oxides, hydrated oxides or hydroxides may be formed, and the relevant half reactions will be of the form shown in equations 2(a) and 2(b) (Table 1.7). In these circumstances the a + will be governed by the solubility product of the solid compound and the pH of the solution. At higher pH values the solid compound may become unstable with respect to metal anions (equations 3(a) and 3(b), Table 1.7), and metals like aluminium, zinc, tin and lead, which form amphoteric oxides, corrode in alkaline solutions. It is evident, therefore, that the equilibrium between a metal and an aqueous solution is far more complex than that illustrated in Tables 1.9 and 1.10. Nevertheless, as will be discussed subsequently, a similar thermodynamic approach is possible. 

An aqueous solution of a soluble salt contains cations and anions. These ions often have acid-base properties. Anions that are conjugate bases of weak acids make a solution basic. For example, sodium fluoride dissolves in water to give Na, F, and H2 O as major species. The fluoride anion is the conjugate base of the weak acid HF. This anion establishes a proton transfer equilibrium with water ... 

A great many electrolytes have only limited solubility, which can be very low. If a solid electrolyte is added to a pure solvent in an amount greater than corresponds to its solubility, a heterogeneous system is formed in which equilibrium is established between the electrolyte ions in solution and in the solid phase. At constant temperature, this equilibrium can be described by the thermodynamic condition for equality of the chemical potentials of ions in the liquid and solid phases (under these conditions, cations and anions enter and leave the solid phase simultaneously, fulfilling the electroneutrality condition). In the liquid phase, the chemical potential of the ion is a function of its activity, while it is constant in the solid phase. If the formula unit of the electrolyte considered consists of v+ cations and v anions, then... 

Electrodes of the second kind. These electrodes consist of three phases. The metal is covered by a layer of its sparingly soluble salt, usually with the character of a solid electrolyte, and is immersed in a solution containing the anions of this salt. The solution contains a soluble salt of this anion. Because of the two interfaces, equilibrium is established between the metal atoms and the anions in solution through two partial equilibria between the metal and its cation in the sparingly soluble salt and between the anion in the solid phase of the sparingly soluble salt and the anion in solution (see Eqs (3.1.24), (3.1.26) and (3.1.64)). 

A specific feature of reactions occurring in the autoclave is that the least soluble compounds are always precipitated from the homogeneous phase of the reaction. As a result, the equilibrium of the reaction is always shifted to the formation of these very insoluble compounds. Thus, it becomes clear that by varying the composition of the reaction mixture (mainly due to the introduction of new cations and anions) practically all types of the cluster forms being generated in the given system can be obtained in the solution. This is a clear advantage of the hydrothermal technique for cluster synthesis in the autoclave. 

Several chemical geothermometers are in widespread use. The silica geothermometer (Fournier and Rowe, 1966) works because the solubilities of the various silica minerals (e.g., quartz and chalcedony, Si02) increase monotonically with temperature. The concentration of dissolved silica, therefore, defines a unique equilibrium temperature for each silica mineral. The Na-K (White, 1970) and Na-K-Ca (Fournier and Truesdell, 1973) geothermometers take advantage of the fact that the equilibrium points of cation exchange reactions among various minerals (principally, the feldspars) vary with temperature. 

Sol id Sol utions. The aqueous concentrations of trace elements in natural waters are frequently much lower than would be expected on the basis of equilibrium solubility calculations or of supply to the water from various sources. It is often assumed that adsorption of the element on mineral surfaces is the cause for the depleted aqueous concentration of the trace element (97). However, Sposito (Chapter 11) shows that the methods commonly used to distinguish between solubility or adsorption controls are conceptually flawed. One of the important problems illustrated in Chapter 11 is the evaluation of the state of saturation of natural waters with respect to solid phases. Generally, the conclusion that a trace element is undersaturated is based on a comparison of ion activity products with known pure solid phases that contain the trace element. If a solid phase is pure, then its activity is equal to one by thermodynamic convention. However, when a trace cation is coprecipitated with another cation, the activity of the solid phase end member containing the trace cation in the coprecipitate wil 1 be less than one. If the aqueous phase is at equil ibrium with the coprecipitate, then the ion activity product wi 1 1 be 1 ess than the sol ubi 1 ity constant of the pure sol id phase containing the trace element. This condition could then lead to the conclusion that a natural water was undersaturated with respect to the pure solid phase and that the aqueous concentration of the trace cation was controlled by adsorption on mineral surfaces. While this might be true, Sposito points out that the ion activity product comparison with the solubility product does not provide any conclusive evidence as to whether an adsorption or coprecipitation process controls the aqueous concentration. 

Sometimes a metal electrode may be directly responsible to the concentration of an anion which either gives rise to a complex or a precipitate with the respective cations of the metal. Therefore, they are termed as second-order electrodes as they respond to an ion not directly involved in the electron transfer process. The silver-silver chloride electrode, as already described in Section 16.3.1.1.3, is a typical example of a second-order electrode. In this particular instance, the coated Ag wire when dipped in a solution, sufficient AgCl dissolves to saturate the layer of solution just in contact with the respective electrode surface. Thus, the Ag+ ion concentration in the said layer of solution may be determined by the status of the solubility product (Kvfa equilibrium ... 

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