Salts solubility product constants

Acid dissociation Aquated metal ions Amphoteric behaviour Solubilities of salts Solubility product constants... 

It is important to note that the solubility product relation applies with sufficient accuracy for purposes of quantitative analysis only to saturated solutions of slightly soluble electrolytes and with small additions of other salts. In the presence of moderate concentrations of salts, the ionic concentration, and therefore the ionic strength of the solution, will increase. This will, in general, lower the activity coefficients of both ions, and consequently the ionic concentrations (and therefore the solubility) must increase in order to maintain the solubility product constant. This effect, which is most marked when the added electrolyte does not possess an ion in common with the sparingly soluble salt, is termed the salt effect. 

If S moles of CaCC>3 dissolve in a liter of water, then S moles each of calcium ion and carbonate ion form. With these ion concentrations equal to S, the solubility of CaCC>3 is calculated as 9.3 x 10 5 M. The higher solubility of magnesium carbonate in water, 6.3 x 10 3 M, results from the larger solubility product constant. Nevertheless, both of these carbonate salts are rather insoluble, and the excess carbonate anions provided by the sodium carbonate effectively precipitate the calcium and magnesium ions from solution. 

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. 

K-SALT 1 (KspA) Solubility Product constant - KspA 12.81... 

K-SALT 2 (KspB) Solubility product constant - KspB -5296... 

For this particular salt the numerical value of A)is 1.6 x 10 8 at 25°C. Note that the Pb2+ and SO2- ions are formed in equal amounts, so the right-hand side of the equation could be represented as [x 2. If the numerical value of the solubility product constant is known, then the concentration of the ions can be determined. And if one of the ion concentrations can be determined, then /C, can be calculated. 

The solubility product constant, Ksp, is the equilibrium constant expression for sparingly soluble salts. It is the product of the ionic concentration of the ions, each raised to the power of the coefficient of the balanced chemical equation. 

Chemists use a quantity called the solubility product constant, or to compare the solubilities of salts. AT p is calculated in much the same way as an equilibrium constant (K, see Chapter 14). The product concentrations are multiplied together, each raised to the power of its coefficient in the balanced dissociation equation. There s one key difference, however, between a and a is a quantity specific to a saturated solution of salt, so the con-... 

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 ... 

If the salts are not of the same kind, molar solubilities of each salt must be calculated to compare their solubilities. For example, although the solubility product constant of Ag2Cr04 is smaller than that of AgCl, the former is more soluble than the latter. 

Moeller et al. [294, 295] have thoroughly studied the pH values at which the rare earth hydroxides are precipitated from various salt solutions, and also the solubility and solubility product constants of the hydroxides. Their results are summarized in Table 18. 

The use of radiotracers is an excellent technique for measuring the solubility product constant of sparingly soluble salts or for making other studies of substances present in low concentrations. Another very important and classic example of the use of radiotracers is that of studying the occurrence and properties of isotopic exchange reactions—reactions of the type... 

FIGURE 11.19 The relative magnitudes of the solubility quotient, Qsp, and the solubility product constant, fCsp, are used to decide whether a salt will precipitate (left) or dissolve (right). 

The solubility product constant allows us to predict the degree of completeness of precipitation reactions. Whenever the product of the concentrations (each raised to the appropriate power) exceeds Ksp, the salt will precipitate until the concentration product equals Ksp. 

Solubility Product Constant and Solubility of Sparingly Soluble Salts... 

In this chapter, we will extend the concepts of equilibrium that have been discussed in previous chapters. In Chapter 10 we discussed the concept of equilibrium in relation to saturated solutions in which an equilibrium was established between solvated ions and undissolved solute. In Chapter 11 we discussed the solubility of different salts when we looked at the formation of precipitates. In this chapter you will see the connection between these two ideas with the introduction of the solubility product constant, Ksp, which is a quantitative means of describing solubility equilibria. This measure helps to predict and explain the precipitation of different salts from solution. You will also see how the common-ion effect, temperature, and pH affect solubility. 

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. 

Recall from previous chapters that solubility is the amount of a solid that it takes to saturate the solution that it is in. Also remember that not every salt is completely soluble in water. The solubility product constant helps determine how soluble a salt is in solution at a particular temperature. The values for these constants can be found in Appendix 4, Reference Tables, in the back of this book. 

One salt to examine is PbS04. Lead sulfate dissolves according to the following equation PbS04(s) <—> Pb2+(aq) + S042 (aq). The solubility product constant for lead sulfate can be written as KsP = [Pb2+][S042 ] /1. Notice that the solid was not included in the equilibrium constant expression. What concentration of lead and sulfate ions will you find in a saturated solution of lead sulfate at 298 K The solubility product constant for lead sulfate is 1.6 x 10-8. This means that the concentration of lead ions and sulfate ions is each 1.26 x 10-4M. 

PROBLEM Magnesium hydroxide has a solubility product constant of 1.8 x 10-11 at 298 K. Write the equilibrium constant expression for this salt. What is the concentration of magnesium and hydroxide ions in a saturated... 

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. 

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

A friend tells you The constant Ksp of a salt is called the solubility product constant and is calculated from the concentrations of ions in the solution. Thus, if salt A dissolves to a greater extent than salt B, salt A must have a higher Ksp than salt B. Do you agree with your friend Explain. 

If a saturated solution of silver chloride contains an AgCl concentration of 1.34 X 10 M, confirm that the solubility product constant of this salt has the value shown in Table 3. 

Copper(I) chloride has a solubility product constant of 1.2 x 10 " and dissolves according to the equation below. Calculate the solubility of this salt in ocean water in which the [CF] = 0.55. 

Though slightly soluble hydroxides are not salts, they have solubility product constants. Magnesium hydroxide is an example. 

The equilibrium constant for the dissolution of a slightly soluble salt is the solubility product constant Ksp. 

The molar solubility of a salt in water is not the same as its solubility product constant, but a simple relation often exists between them. For example, let s define S... 

Solubility product constants (like solubilities) can be sensitive to temperature. At 100°C the Ksp for silver chloride is 2.2 X 10 hot water dissolves about 12 times as much silver chloride as does water at 25°C. Refer to Table 16.2 for the solubility product constants at 25°C of a number of important sparingly soluble salts. 

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