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Water/Rock Interactions
Published in William J. Deutsch, Groundwater Geochemistry, 2020
In this case the presence of sulfate at 10−2.3 mol/L lowered the solubility of barite by over two orders of magnitude (from 10−4.99 to 10−7.68 mol/L). This is an example of the common-ton effect on solubility, where the sulfate ion common to both barite and gypsum lowers the solubility of barite. Barite solubility was decreased significantly because of the relatively high solubility of gypsum. If gypsum had been added to a solution in equilibrium with barite, the initial sulfate concentration (≈10−4.99 mol/L) would depress the solubility of gypsum, but only by a relatively small amount because the sulfate equilibrium concentration for gypsum (≈10−2.3 mol/L) is much greater than the initial sulfate concentration at barite equilibrium. In general, the impact of the common-ion effect is that the solubility of a mineral will be less if one of its components is present in solution when the solution contacts the mineral. The common-ion effect of calcium in calcite and gypsum on the solubilities of these minerals is discussed by Wigley.36 Wigley also describes the probable importance of concurrent solution and precipitation of calcite and gypsum in karst terrain.
Preformulation
Published in Sandeep Nema, John D. Ludwig, Parenteral Medications, 2019
N. Murti Vemuri, Abira Pyne Ramakrishnan
Common ion effect, or salting-out effect, is also depicted on Figure 5.1, representing the pH-solubility profile of a weakly basic drug. From the pH of maximum solubility, as one moves towards lower pH values, there is an increase in the concentration of the counterion (e.g., [Cl−]). Depending on the value of the solubility product (a function of the nature of the drug and the counterion), this increase may be compensated by a decrease in the concentration of the ionized drug molecule. This decrease occurs through a precipitation of the drug in its corresponding salt form. This phenomenon is known as “salting out,” or common-ion effect, and can be an important consideration in selecting salt forms or buffer systems for formulations.
Abiotic Processes
Published in Robert C. Knox, David A. Sabatini, Larry W. Canter, Subsurface Transport and Fate Processes, 1993
Robert C. Knox, David A. Sabatini, Larry W. Canter
The above discussion has centered on ground water systems where the dissolution/precipitation of discrete minerals is of concern. It is not uncommon in ground water systems for multiple minerals to be present and for common ions to exist amongst the minerals. The contribution of ions from one mineral will affect the solubility of other minerals containing the same ion (referred to as the “common ion effect”). This serves to further complicate the system and increase the difficulty in determining the solubility limits of various species. Algorithms have been developed to aid in this process. One example of such an algorithm is MINTEQA1, an equilibrium metal speciation model developed for the EPA (Brown and Allison, 1987). MINTEQA1 utilizes the thermodynamic expressions and constants necessary to assess the equilibrium speciation of metals in a variety of subsurface environments based on geochemical principles. A manual is available that explains the fundamentals of the model and the mechanics of using the model (Brown and Allison, 1987).
Aqueous solubility of beryllium(II) at physiological pH: effects of buffer composition and counterions
Published in Preparative Biochemistry & Biotechnology, 2020
Rebecca C. Lim, Bhagya De Silva, Ji Hye Park, Vernon F. Hodge, Ronald K. Gary
The solubility of beryllium in protein-free solutions is also of interest. The aqueous speciation of beryllium(II) is complex. At pH 4 and lower, Be2+ exists predominantly as the metal aquo complex [Be(H2O)4]2+, while the cyclic trimer [Be3(OH)3]3+ becomes increasingly abundant in the near-physiological pH range.[10] In general, increasing ionic strength increases the apparent solubility of a metal salt. In addition, specific counterions can have either of two effects. A specific counterion can decrease the apparent solubility via the “Common Ion Effect”. For a metal ion Mx+ whose insoluble form is depicted as (Mx+)a(Ay−)b, addition of anionic species Ay- decreases the observed solubility, a manifestation of Le Chatelier’s principle that is the basis for the Common Ion Effect. In the present case, the insoluble species expressed in this format is (Be2+)(OH–)2, and this phenomenon accounts for the pH-dependence of the solubility of beryllium ion in water. Alternatively, a specific counterion can increase the observed solubility via the “Diverse Ion Effect” (also known as the Uncommon Ion Effect). For the same metal ion Mx+ whose insoluble form is depicted as (Mx+)a(Ay−)b, addition of anionic species Bz− may increase the apparent solubility if a soluble complex (Mx+)c•(Bz−)d is formed. Ion pairing or similar interactions between the metal ion and anion explain the enhanced solubility observed with the Diverse Ion Effect. The term ion pairing includes ions that make direct contact, as well as those that interact while surrounded by complete or partial hydration shells.[11] We provide evidence that Be2+ participates in a Diverse Ion Effect relationship with SO42–. If H2SO4 is used for pH adjustment, the concentration of sulfate contributed from buffer preparation is, by itself, sufficient to increase solubility and to alter the results of binding studies monitored by isothermal titration calorimetry (ITC). This information can be useful when designing biochemical experiments that require Be2+ in buffered solutions of well-defined composition at biologically-relevant pH.