Salts solubility equilibrium

Electrodes such as Cu VCu which are reversible with respect to the ions of the metal phase, are referred to as electrodes of the first kind, whereas electrodes such as Ag/AgCl, Cl" that are based on a sparingly soluble salt in equilibrium with its saturated solution are referred to as electrodes of the second kind. All reference electrodes must have reproducible potentials that are defined by the activity of the species involved in the equilibrium and the potential must remain constant during, and subsequent to, the passage of small quantities of charge during the measurement of another potential. 

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. 

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. 

The amount of lead that remains in solution in surface waters depends upon the pH of the water and the dissolved salt content. Equilibrium calculations show that at pH >5.4, the total solubility of lead is approximately 30 pg/L in hard water and approximately 500 pg/L in soft water. Sulfate ions, if present... 

Pure-component properties from which prediction of salt effect in vapor-liquid equilibrium might be sought, include vapor pressure lowering, salt solubility, degree of dissociation and ionic properties (charges and radii) of the salt, polarity, structural geometry, and perhaps others. 

Salts soluble in ethanol as well as water have been found to break the azeotrope, while salts which are very soluble in water and only slightly soluble in the alcohol move the azeotrope to richer ethanol regions without breaking it. Furthermore, the salts or compounds which dissolve more in one component are found to raise the volatility of the other component. This finding is in conformity with that of previous workers in the field (8,11,12,13,18,19,23,24,27). In this work the equilibrium diagrams were obtained at atmospheric pressure (690-700 mmHg), and under saturation conditions. 

The solubility of oxygen in rivers and streams depends on various factors, such as temperature, salt concentration, depth, and biological activity. Its solubility equilibrium is 02(g) 02(aq). Compare the dissolved oxygen concentration at 1.0 atm total pressure (0.21 atm partial pressure) with its concentration for pure oxygen at 1.0 atm. 

As more salt is added, excess salt is present in the solid phase and the solution composition is invariant. Therefore, the pH is constant and the product of the cation and anion activities equals the solubility product, as deLned in Equation 15.5, in the absence of cation or anion from other sources including molecular complex forms (Amis, 1983). At this point, more salt will not dissolve, and the salt concentration represents the solubility of the drug in the speciLc salt form. To conLrm that the salt solubility has been reached, it should be veriLed (Anderson and Conradi, 1985) that the solid salt phase in equilibrium with the solution has not been contaminated with the uncharged form precipitate. 

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. 

Electrode of the second kind — Electrodes of the second kind [i-ii] contain a metal, a poorly soluble compound of this metal (which is usually a salt but it may be oxide or hydroxide as well), and an electrolyte which can establish a solubility equilibrium with the precipitate. Typical examples are Ag(s) AgCl(s) KCl(aq), Hg(l) Hg2Cl2(s) KCl(aq), Hg(l) HgO(s) NaOH(aq), Hg(l) HgS04(s) H2S04(aq). The solubility equilibrium, e.g., for the silver chloride electrode, is as follows... 

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. 

In dilute aqueous solutions, it has been demonstrated experimentally for poorly soluble ionic salts (solubilities less than 0.01 molL ) that the mathematical product of the total molar concentrations of the component ions is a constant at constant temperature. This product, is called the solubility product. Thus for a saturated solution of a simple ionic compound AB in water, we have the dynamic equilibrium ... 

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 ). 

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... 

Metal hydroxides can be described as salts of a strong base, the hydroxide ion. The solubility of salts in which the anion is a different weak or strong base is also affected by pH. For example, consider a solution of a slightly soluble fluoride, such as calcium fluoride. The solubility equilibrium is... 

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. 

Zone-B Solubility of the Weak Acid audits Salt Under Equilibrium Conditions Inside of the pH region enclosed by the limits of Zone-B, one hnds both the free acid and its salt form. The mass balance relationship in this zone defines the total solubility (Sj) at any particular pH value as the sum of the concentration of the free acid plus the concentration of its salt form ... 

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... 

Solubility Equilibrium. The misconceptions regarding the amount of solid materials in equilibrium and the dynamic aspect are equally important in the discussion. If one observes a saturated sodium chloride solution together with solid sodium chloride, and adds an additional portion of solid sodium chloride to it, this portion sinks down without dissolving (see E6.2). If one measures the density of the saturated solution before and after the addition of salt portions, one gets the same measurements (see E6.2). The concentration of the saturated solution does not depend on how much solid residue is present equilibrium sets in between the saturated solution and arbitrary amounts of solid residue (see Fig. 6.2) ... 

Problem Solubility equilibrium of a salt is not limited to the concentrations of the ions that deliver the pure salt solution. If, for example, a large amount of chloride ions are added to the saturated sodium chloride solution, then the equilibrium deviates in such a way that solid sodium chloride is formed and precipitates (Le Chatelier s principle of getting rid of the stress ). This way, the position of equilibrium is altered however, the product of concentrations of sodium ions and chloride ions remain constant solubility product. 

The pH also influences the solubility of salts that contain a basic anion. For example, the solubility equilibrium for Bap2 is... 

In this cycle, AsgjG° is related by equation 6.48 to the equilibrium constant, K, for the dissolution process for a sparingly soluble salt, the equilibrium constant is p. 

For a slightly soluble salt, an equilibrium is set up between the excess solid (MX) and the ions in solution ... 

Several different subscripts are used with the equilibrium constant When all reactants or products that appear in the equilibrium expression are gases, and pressures are used instead of molarity, the symbol for the equilibrium constant is Kp. If weak acids and bases are under investigation, the equilibrium constant symbol is written as K or K,. When the equilibrium constant refers to the autoionization of water, the symbol K, is used. Solubility product equilibria express the ion concentrations of partially soluble salts. The equilibrium constant symbol K p is used for solubility product equilibria. ... 

The solubilities of the scale-forming salts barium and strontium sulphates in aqueous solutions of sodium chloride have been reviewed by Raju and Atkinson (1988, 1989). Equations were proposed for the prediction of specific heat capacity, enthalpy and entropy of dissolution, etc., for all the species in the solubility equilibrium, and the major thermodynamic quantities and equilibrium constraints expressed as a function of temperature. Activity coefficients were calculated for given temperatures and NaCl concentrations and a computer program was used to predict the solubility of BaS04 up to 300 °C and SrS04 up to 125 °C. 

These expressions correctly describe the relationship between solubility product and solubility either in pure water or in solutions that do not contain any other source of the ions directly involved in the solubility equilibrium. These expressions must be modified when salts containing common ions are present in solution. 

Identify the formula for the salt, and write the chemical equation that represents the solubility equilibrium. Write the solubility product constant expression based on the chemical equation. 

The procedure for the calculation of salt solubilities in aqueous solutions can be extended to organic solvents or solvent mixtures, starting from the condition that the fugacity (chemical potential) of a precipitated salt in phase equilibrium is identical in water, an organic solvent or the aqueous solution (r ee Figure 8.16). This means that the already available standard thermodynamic properties in the aqueous phase given in Table 8.1 can be used to determine the salt solubility in organic solvents or solvent mixtures [10]. 

An ion added to a solution already containing that ion is called a common ion. When a common ion is added to an equilibrium solution of a weak electrolyte or a slightly soluble salt, the equilibrium shifts according to Le ChStelier s principle. For example, when silver nitrate (AgNOs) is added to a saturated solution of AgCI,... 

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