Solubility product constant predicting precipitates

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. 

In this section, you determined the solubility product constant, Kgp, based on solubility data. You obtained your own solubility data and used these data to calculate a value for Kgp. You determined the molar solubility of ionic solutions in pure water and in solutions of common ions, based on their Ksp values. In section 9.3, you will further explore the implications of Le Chatelier s principle. You will use a reaction quotient, Qsp, to predict whether a precipitate forms. As well, you will learn how selective precipitation can be used to identify ions in solution. 

In Chapter 9, as in most of Unit 4, you learned about equilibrium reactions. In this section, you analyzed precipitation reactions. You mainly examined double-displacement reactions—reactions in which two soluble ionic compounds react to form a precipitate. You used the solubility product constant, Ksp, to predict whether or not a precipitate would form for given concentrations of ions. In Unit 5, you will learn about a class of reactions that will probably be new to you. You will see how these reactions interconvert chemical and electrical energy. 

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. 

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. 

It is interesting to compare these values and to predict what happens if, to a solution which contains the complex ion, a reagent is added which, under normal circumstances, would form a precipitate with the central ion. It is obvious that the higher the value of the instability constant, the higher the concentration of free central ion (metal ion) in the solution, and therefore the more probable it is that the product of ion concentrations in the solution will exceed the value of the solubility product of the precipitate and hence the precipitate... 

Predicting precipitates The solubility product constant expression can also be used to predict whether a precipitate will form when two solutions of ionic compounds are mixed. The molar concentrations of the ions in a solution are used to calculate the ion product, Qsp- If 0sp > sp a precipitate will form, reducing the ion concentrations until the system reaches equilibrium and the solution is saturated. If Qsp < sp precipitate forms. The following example problem demonstrates how to use and to determine whether a precipitate will form. 

Sometimes it is important to know under what conditions a precipitate will form. For example, if we are analyzing a mixture of ions, we may want to precipitate only one type of ion to separate it from the mixture. In Section 9.5, we saw how to predict the direction in which a reaction will take place by comparing the values of J, the reaction quotient, and K, the equilibrium constant. Exactly the same techniques can be used to decide whether a precipitate is likely to form when two electrolyte solutions are mixed. In this case, the equilibrium constant is the solubility product, Ksp, and the reaction quotient is denoted Qsp. Precipitation occurs when Qsp is greater than Ksp (Fig. 11.17). 

Solubility products can be used to predict the stability of a mineral by comparing the observed ion product, [A (aq)][B (aq)], to the mineral s K. K the ion product is greater than K, the solution is supersaturated with respect to that mineral. In this case, precipitation should proceed spontaneously until the ion concentrations are decreased to levels that lower the ion product to the value dictated by the K. Conversely, if the ion product is less than the K, the solution is undersaturated with respect to that mineral. Dissolution should proceed spontaneously until the ion concentrations are increased to levels that raise the ion product to a value equal to the K. At equilibrium, where the ion product has a value equal to that of the K, the rate of mineral dissolution is equal to the rate of precipitation, so the ion concentrations remain constant over time. 

The direction of a reaction can be assessed straightforwardly by comparing the equilibrium constant (Keq) and the ratio of the product solubility to the substrate solubility (Zsat) [39]. In the case of the zwitterionic product amoxicillin, the ratio of the equilibrium constant and the saturated mass action ratio for the formation of the antibiotic was evaluated [40]. It was found that, at every pH, Zsat (the ratio of solubilities, called Rs in that paper) was about one order of magnitude greater in value than the experimental equilibrium constant (Zsat > Keq), and hence product precipitation was not expected and also not observed experimentally in a reaction with suspended substrates. The pH profile of all the compounds involved in the reaction (the activated acyl substrate, the free acid by-product, the antibiotic nucleus, and the product) could be predicted with reasonable accuracy, based only on charge and mass balance equations in combination with enzyme kinetic parameters [40]. 

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