Chemical equilibrium solubility product constant

ELECTROLYTES, EME, AND CHEMICAL EQUILIBRIUM TABLE 8.6 Solubility Product Constants Continued)... 

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 ... 

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. 

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. 

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. 

A suite of both oxidized and reduced iron minerals has been found as efflorescences and precipitates in or near the acid mine water of Iron Mountain. The dominant minerals tend to be melan-terite (or one of its dehydration products), copiapite, jarosite and iron hydroxide. These minerals and their chemical formulae are listed in Table III from the most ferrous-rich at the top to the most ferric-rich at the bottom. These minerals were collected in air-tight containers and identified by X-ray diffractometry. It was also possible to check the mineral saturation indices (log Q(AP/K), where AP = activity product and K = solubility product constant)of the mine waters with the field occurrences of the same minerals. By continual checking of the saturation index (S.I.) with actual mineralogic occurrences, inaccuracies in chemical models such as WATEQ2 can be discovered, evaluated and corrected (19), provided that these occurrences can be assumed to be an approach towards equilibrium. 

Use the following terms to create a concept map chemical equilibrium, equilibrium constant, solubility product constant, reversible reactions, and Le Chdtelier s principle. 

The solubility product constant expression is the product of the concentrations of the ions each raised to the power equal to the coefficient of the ion in the chemical equation. The small value of indicates that products are not favored at equilibrium. Thus, few barium ions are present at equilibrium (1.0 X lO M) and a patient can safely ingest a barium sulfate solution to obtain a clear X ray like the one shown in Figure 18-14. 

Write the balanced chemical equation for the solubility equilibrium and the solubility product constant expression. 

For aqueous systems, a unit activity is expected for the solid species (i.e., we assume that the chemical reactivity of a solid in water is unchanging as long as there is solid in equilibrium with the solution). Also, for dilute concentrations, we assume that the activities are equal to the concentrations of the species. With these assumptions, we can reduce the solubility product constant equation to... 

Not all of the elements selected for study are expected to behave in the same way, based on general chemical properties, when buried in a reducing, sulfide-rich, marine environment. As a first approximation, the expected concentrations of dissolved metals can be estimated from published solubility-product constants, by assuming a free-sulfide concentration and neglecting complex formation, as has been done in Table XII. The calculated equilibrium concentrations are extremely low, except for Mn and possibly Fe, which suggests that Cu, Zn, and Pb should be fixed as their insoluble sulfides, whereas Mn might be chemically mobile. 

The principle of equilibrium can also be applied when an excess of a solid is added to water to form a saturated solution. The solubility product (K p) is an equilibrium constant defined by the law of chemical equilibrium. Solubility is an equilibrium position, and the K p value of a solid can be determined by measuring its solubility. Conversely, the solubility of a solid can be determined if its K p value is known. 

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. 

As with any equilibrium, the extent to which this dissolution reaction occurs is expressed by the magnitude of its equilibrium constant. Because fliis equilibrium equation describes the dissolution of a solid, the equilibrium constant indicates how soluble tire solid is in water and is referred to as the solubility-product constant (or simply tire solubility product). It is denoted Kgp, where sp stands for solubility product. The equilibrium-constant expression for this process is written according to tire same rules as those that apply to any equilibrium-constant expression. That is, Ihe concentration terms of the products are multiplied togefli-er, and each is raised to the power of its stoichiometric coefficient in the balanced chemical equation, and these are divided by the concentration terms of the reactants multiplied together, and each raised to the power of its stoichiometric coefficient. Solids, liquids, and solvents do not appear in the equilibrium-constant expressions for heterogeneous equilibria (Section 15.3), however, so the solubility product equals the product of the concentration (f the ions involved in the equilibrium, each raised to the poioercfits coefficient in the equilibrium equation. Thus, the solubility-product expression for the equilibrium expressed in Equation 17.15 is... 

The equilibrium expression for a chemical equation that represents the dissolving of an ionic compound is the solubility-product constant (Kgp). For Cap2, the solubility-product constant is ... 

The Solubility-Product Constant, Ksp The solubility-product constant of an ionic compound is the equilibrium constant for the chemical equation that describes the dissolving of the compound. 

The ionization of an ionic salt, such as NaCl (see Equation 11.10), poses a special problem. The un-ionized chemical species in this equilibrium is not dissolved. This is an example of what is termed a heterogeneous equilibrium (i.e., two or more of the equilibrium participants are present in different phases [solid, liquid, gas, or dissolved]). In this case, we have a solid that is not dissolved while everything else is dissolved. Since the solid is undissolved, it does not make sense to refer to its molar concentration. An interesting and important fact about this undissolved solid is that regardless of how much is present, the concentrations of the dissolved ions are constants (at a given temperature). Thus, while the molar concentrations of the ions are real numbers, the molar concentration of the un-ionized species is a nonsensical term. In this case, a special equilibrium constant is defined which uses only the molar concentrations of the dissolved ions in its definition. This special equilibrium constant is called the solubility product constant (K p), which is defined as the mathematical product of the molar concentrations of the ions raised to the power of their balancing... 

Having introduced matters pertaining to the electrochemical series earlier, it is only relevant that an appraisal is given on some of its applications. The coverage hereunder describes different examples which include aspects of spontaneity of a galvanic cell reaction, feasibility of different species for reaction, criterion of choice of electrodes to form galvanic cells, sacrificial protection, cementation, concentration and tempera lure effects on emf of electrochemical cells, clues on chemical reaction, caution notes on the use of electrochemical series, and finally determination of equilibrium constants and solubility products. 

Note The above potentials, E, are for pH 7 at equal concentrations of oxidised and reduced species. These equilibrium values are as important as stability constants and solubility products for an understanding of cellular chemical systems. These are free energy changes in volts, E, and where n3E is in kilocalories. The [Fe3+]/[Fe2+] is related to an equilibrium constant, K (see Section 4.17). 

One of the most basic requirements in analytical chemistry is the ability to make up solutions to the required strength, and to be able to interpret the various ways of defining concentration in solution and solids. For solution-based methods, it is vital to be able to accurately prepare known-strength solutions in order to calibrate analytical instruments. By way of background to this, we introduce some elementary chemical thermodynamics - the equilibrium constant of a reversible reaction, and the solubility and solubility product of compounds. More information, and considerably more detail, on this topic can be found in Garrels and Christ (1965), as well as many more recent geochemistry texts. We then give some worked examples to show how... 

Solubility Equilibria The Solubility Product Principle.—It was seen on page 133 that the chemical potential of a solid is constant at a definite temperature and pressure consequently, when a solution is saturated with a given salt Mv A, the chemical potential of the latter in the solution must also be constant, since the chemical potential of any substance present in two phases at equilibrium must be the same in each phase. It is immaterial whether this conclusion is applied to the undissociated molecules of the salt or to the ions, for the chemical potential is given by... 

Keep in mind that all precipitation events are due to the fact that the concentrations of the soluble species (Me and S=, for example) have exceeded the solubility of the solid species. For the sake of discussion, let s look at a solution containing 1 g/L copper as cupric ions (Cu+2). The solubility product for Cu(0H)2 is defined as the equilibrium constant for the dissolution of copper hydroxide to form the individual ions of copper and hydroxide. Chemically, it is defined by the following equilibrium ... 

Tables of this sort are extremely useful, because they feature much chemical and electrical information condensed into quite a small space. A few electrode potentials can characterize quite a number of cells and reactions. Since the potentials are really indices of free energies, they are also ready means for evaluating equilibrium constants, complex-ation constants, and solubility products. Also, they can be taken in linear combinations to supply electrochemical information about additional half-reactions. One can tell from a glance at an ordered list of potentials whether or not a given redox process will proceed spontaneously.

The available thermodynamic data are of two types stabihty constants, enthalpy and entropy of reaction for the formation of soluble complexes Th(S04) " " and solubihty data for various solid phases. The two sources are linked because the solubility of the solid phases depends on the chemical speciation, i.e., the sulphate complexes present in the aqueous phase. The analysis of the experimental stability constants has been made using the SIT model however, this method cannot be used to describe the often very high solubility of the solid sulphate phases. In order to describe these data the present review has selected a set of equilibrium constants for the formation of Th(S04) and Th(S04)2(aq) at zero ionic strength based on the SIT model and then used these as constants in a Gibbs energy minimisation code (NONLINT-SIT) for modelling experimental data to determine equilibrium constants for the formation of Th(S04)3 and the solubility products of different thorium sulphate solids phases. 

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