Minerals solubility product constants

Table 3.9. Solubility product constants of common compounds and minerals of selected major and trace elements at 25 °Ca...

Using the solubility product constants for calcite and strontianite and assuming a calcium activity of 1.6 mmol/L, a distribution coefficient of 0.8 for strontium and 0.98 for calcite, and a ratio of 50 1 (=0.02) in the solid-solution mineral, the following equation gives the activity of strontium ... 

Mineral solubility is dependent on the type of ions composing a particular mineral. The parameter used to predict mineral solubility is the solubility product constant or... 

The quantity pi sp is called solubility product constant or simply solubility product . It indicates the extent of dissolution of a particular chemical species in a given solution. Appendix C gives values of pi sp for many useful oxides and minerals. 

A thermodynamic model of dissolution is presented in this chapter, which relates the solubility product constant to the thermodynamic potentials and measurable parameters, such as temperature and pressure of the solution. The resulting relations allow us to develop conditions in which CBPCs are likely to form by reactions of various oxides (or minerals) with phosphate solutions. Thus, the model predicts formation of CBPCs. 

The appendices contain the thermodynamic data, the solubility product constants that are relevant to CBPC formation or their durability, and formulae of minerals that were discussed in the text. The thermodynamic data of phosphates is difhcult to find in the common literature. Some excellent sources such as Phosphate Minerals by Nriagu and Moore and Inorganic Phosphate Materials by Kanazawa are out of print. The most commonly used data books such as CRC Handbook of Physics and Chemistry do not contain data on most phosphate compounds. For this reason, these appendices are provided to facilitate the discussion in the text and also for the benefit of those who wish to pursue further research in CBPCs. 

The data from different sources does not often match exactly, while it has been necessary to use the data on oxides from one source and phosphates of the same elements from another. To avoid any confusion resulting from this, we have used a certain order in using these sources. Pourbaix s Atlas of Electrochemical Equilibria [1] is the first source that we have used for the Gibb s free energy of oxides and ions, and for the solubility product constants. This is prompted by the fact that much of the formulation discussed in this book is hinged to Pourbaix s treatment, and to be consistent, Pourbaix s data is preferred over others. The CRC Handbook of Chemistry and Physics [2] is the next source from which much of the enthalpy and specific heat of oxides and ions are taken. The data on phosphates comes from The Phosphate Minerals [3], while the mineral formulae are from Dana s Mineralogy Handbook [4] and also from The Phosphate Minerals [3]. These detailed references and additional ones [5,6] useful for further development of CBPC materials are given below. 

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. 

Solubility Product Constants for Soil Clay Minerals... 

After speciation and activities have been calculated for all the free ions, ion pairs, triplets, etc., a mineral saturation index can be computed. The saturation index, SI, is defined as the logarithm of the ratio of the ion-activity product, lAP, to the solubility product constant,... 

Most parts of the oceans are nearly saturated with CaF2. The mineral fluorite, CaF2, may precipitate when ocean water evaporates. A saturated solution of CaF2 at 25°C has a solubility of 3.4 x 10 " M. Calculate the solubility product constant for CaF2. 

For many of the more abundant elements, such as Al, Fe, and Mn, precipitation of mineral forms is common and may greatly influence or even control their solubility. For most trace elements, direct precipitation from solution through homogeneous nucleation appears to be less likely than adsorption-desorption, by virtue of the low concentration of these metals and metalloids in soil solutions in well-aerated dryland soils. When soils become heavily polluted, metal solubility may reach a level to satisfy the solubility product to cause precipitation. Precipitation may also occur in the immediate vicinity of the phosphate fertilizer zone, where the concentration of heavy metals and metalloids present as impurities may be sufficiently high. Precipitation of trace metals as sulfides may have a significant role in metal transformation in reduced environments where the solution sulfide concentration is sufficiently high to satisfy the solubility product constants of metal sulfides (Robert and Berthelin, 1986). 

It may also be argued that the simulation of water-rock interactions should allow for solubility equilibria involving feldspars, micas, etc. For such studies the choice of solublity product constants and free energies must and should be made by the investigators. We cannot propose such values here when an enormous range of values and properties (solid-solutions, interlayering, defects, surface areas, etc.) is known to exist for these minerals and reversible solubility behavior has not been demonstrated. 

In this equation lAP denotes the Ion Activity Product (in the example of gypsum or anhydrite these would be ([Ca " ] [SO/ ]). Kj,p is the solubility product constant of the respective mineral. A saturation index SI = 0 describes the condition in which the solution of the corresponding mineral is just saturated, SI > 0 describes the condition of supersaturation of the solution, SI < 0 its undersaturation. The activity [A] of a substance is calculated according to the equation ... 

Solubility product Solubility product constant, iQp, reflects the relationship between dissolved species and precipitated species. Each ionic compound has its own solubility limit, which is the maximum amount of the compound that can remain in solution. IQp is commonly used in solubility calculations to determine the precipitation potential of mineral salts. Certain combinations of cations and anions form sparingly soluble salts in water, and scaling in RO/NF may occur when the salts are concentrated beyond their solubility limits. See Table 6.10. 

The mineral fluorite is calcium fluoride, Cap2. Calculate the solubility (in grams per liter) of calcium fluoride in water from the solubility product constant... 

EXERCISE 18.4 Anhydrite is a calcium sulfate mineral deposited when seawater evaporates. What is the solubility of calcium sulfate, in grams per liter Table 18.1 gives the solubility product constant for calcium sulfate. 

Strontianite (strontium carbonate) is an important mineral of strontium. Calculate the solubility of strontium carbonate, SrC03, from the solubility product constant (see Table 18.1). 

Magnesite (magnesium carbonate, MgC03) is a common magnesium mineral. From the solubility product constant (Table 18.1), find the solubility of magnesium carbonate in grams per liter of water. 

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. 

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. 

It appears to be the case that most animals maintain the concentration of mineral ions at constant levels in their extracellular fluids. Perturbations with various forms of acidosis usually result in the animal re-establishing an equilibrium between its body fluids and the apparent solubility product of some mineral. Two important conclusions follow from this. First, it provides a theoretical basis for defining calcification. When there is a change of phase in the total extracellular fluids (i. e., mineralization occurs) then the fluids re-equilibrate to make good the ions which have been lost as minerals. 

However, the most common sources of different results are both based on the approach used for the calculation of the activity coefficient (chapter 1.1.2.6) and the thermodynamic data sets themselves (chapter 2.1.4), which provide the respective program with the fundamental geochemical information of each single species. The thermodynamic databases available partly use severely differing data with different solubility products, different species, minerals and reaction equations. Nordstrom et al (1979, 1990), Nordstrom Munoz (1994), Nordstrom (1996, 2004) discuss this inconsistency of thermodynamic datasets in detail. For some species, for which stability constants have been published, not even the existence of the respective species has been proved doubtless, as can been shown in the following example. 

Parentheses denote activity and brackets denote concentration of the species. The concentration of the Al(OH)3 species represents approximately the lowest possible solubility point of the mineral and it is the product of two constants (K -K ). Thus, its magnitude is not in any way related to pH. Mineral solubility increases as pH increases above the solution pH of zero net charge because of increasing complexa-tion effects, and mineral solubility also increases at pH values below the solution pH of zero net charge because of diminishing common-ion effects (Fig. 2A). All minerals are subject to the common-ion effect and many minerals are subject to the complexation or ion-pairing effect (Fig. 2B). 

The equilibrium constant for a reaction between a solid and its saturated solution is known as the solubility product and is usually given the notation Ksp. Solubility products have been calculated for many minerals, usually using pure water under standard conditions (1 atm pressure, 25°C temperature). 

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