Solubility equilibrium data collection

Phase-Equilibrium Data Shaw, D. G., and A. Maczynski (eds.), IUPAC Solubility Data Series, Vol. 81 Hydrocarbons in Water and Seawater—Revised and Updated, published in 12 parts in /. Fhys. Chem. Ref. Data, 2005 and 2006 Gmehling, J., et al, Vapor-Liquid Equilibrium Data Collection Aqueous-Organic Systems, DECHEMA Chemistry Data Series, Vol. I, Part 1-ld, Schon Wetzel GmbH, Frankfurt/Main, Germany, 1988. 

On a ternary equilibrium diagram like that of Figure 14.1, the limits of mutual solubilities are marked by the binodal curve and the compositions of phases in equilibrium by tielines. The region within the dome is two-phase and that outside is one-phase. The most common systems are those with one pair (Type I, Fig. 14.1) and two pairs (Type II. Fig. 14.4) of partially miscible substances. For instance, of the approximately 1000 sets of data collected and analyzed by Sorensen and Arlt (1979), 75% are Type I and 20% are Type II. The remaining small percentage of systems exhibit a considerable variety of behaviors, a few of which appear in Figure 14.4. As some of these examples show, the effect of temperature on phase behavior of liquids often is very pronounced. 

Using the pKa and the estimated So, the DTT procedure simulates the entire titration curve before the assay commences. Figure 6.7 shows such a titration curve of propoxyphene. The simulated curve serves as a template for the instrument to collect individual pH measurements in the course of the titration. The pH domain containing precipitation is apparent from the simulation (filled points in Fig. 6.7). Titration of the sample suspension is done in the direction of dissolution (high to low pH in Fig. 6.7), eventually well past the point of complete dissolution (pH <7.3 in Fig. 6.7). The rate of dissolution of the solid, described by the classical Noyes-Whitney expression [37], depends on a number of factors, which the instrument takes into account. For example, the instrument slows down the rate of pH data taking as the point of complete dissolution approaches, where the time needed to dissolve additional solid substantially increases (between pH 9 and 7.3 in Fig. 6.7). Only after the precipitate completely dissolves, does the instalment collect the remainder of the data rapidly (unfilled circles in Fig. 6.7). Typically, 3-10 h is required for the entire equilibrium solubility data taking. The more insoluble the... 

Are the equilibrium constants for the important reactions in the thermodynamic dataset sufficiently accurate The collection of thermodynamic data is subject to error in the experiment, chemical analysis, and interpretation of the experimental results. Error margins, however, are seldom reported and never seem to appear in data compilations. Compiled data, furthermore, have generally been extrapolated from the temperature of measurement to that of interest (e.g., Helgeson, 1969). The stabilities of many aqueous species have been determined only at room temperature, for example, and mineral solubilities many times are measured at high temperatures where reactions approach equilibrium most rapidly. Evaluating the stabilities and sometimes even the stoichiometries of complex species is especially difficult and prone to inaccuracy. 

The pH dependence of dorzolamide solubility was also determined between pH 4.0 and 7.0, using acetate, citrate, and phosphate buffer solutions to set the desired pH. These data are collected in Table 2 (also plotted in Figure 4), and show a maximum solubility of approximately 40 mg/mL at pH 5.6. The equilibrium solubility decreases to approximately 13 mg/mL at pH 6 and 4 mg/mL at pH 7.0. These data indicate that in order to have a stable 2% solution for an ophthalmic formulation, the solution pH should be maintained below 5.8. At pH values exceeding 5.8, precipitation of the free base could occur. 

The measurements of water quality parameters (oxidation-reduction potential, pH, temperature, conductivity, dissolved oxygen, and turbidity) and the collection of field screening data with field portable instruments and test kits constitute a substantial portion of field work. Field measurements, such as pH, stand on their own as definitive data used for the calculations of solubility of chemical species and chemical equilibrium in water, whereas others serve as indicators of well stabilization or guide our decision-making in the field. Table 3.8 shows the diversity of field measurement... 

Barrie (1968) collected all the known data on water sorption. From these data it is possible to estimate the effect of the different structural groups on water sorption at different degrees of humidity. Table 18.14 presents the best possible approach to the sorptive capacity of polymers versus water, i.e. the amount of water per structural group at equilibrium, expressed as molar ratio. From these data the solubility (cm3 water vapour (STP) per cm3 of polymer) can be easily calculated. (The multiplication factor is 22.4 x 103/V, where V is the molar volume per structural polymer unit.)... 

In our effort to collect the appropriate data and develop the requisite understanding of geochemical processes, we have developed some adjunct computer programs. These include AACALC (Atomic Absorption and emission spectrometry CALCulation), EQLIST (Equilibrium computation LISTing), and EQPRPLOT (Equilibrium computation PRinting and PLOTing). AACALC (FORTRAN) reduces atomic absorption or emission spectrometry data to concentrations, EqLIST (PL/1) constructs tables from the WATEqZ (input) card file, and EqPRPLOT (FORTRAN) constructs ratio and scatter plots of dissolved constituents, activity products (AP), or activity product to solubility product ratios (AP/K) via computer terminal printer or tape-driven plotter. 

To determine the solubility of each form, one needs to monitor the solution concentration as a function of time more frequently. Enough data points need to be collected so the equilibrium concentration of each form can be assessed. Theoretically, a single experiment starting with the least stable form should generate solubility data for all the other forms. However, since the transition temperature is typically unknown initially, it is best to conduct the solubility experiment with each form. 

Models are often developed to explain certain kinds of data, ignoring other kinds that also might be pertinent. The initial development of Pitzer s equations (33.34) for activity coefficients in concentrated solutions was focused on explaining measurements of vapor pressure equilibrium and of electromotive force (emf). The data could be explained by assuming that the electrolytes examined were, at least in a formal sense, fully dissociated. Later work using these equations to explain solubility data required the formal adoption of a few ion pair species (30). Even so, no speciation/activity coefficient model based on Pitzer s equations is presently consistent with the picture of much more extensive ion-pairing based on other sources, such as Smith and Martell s (35) compilation of association constants. This compilation is a collective attempt to explain other kinds of data, such as electrical conductance, spectrophotometry, and acoustic absorption. 

The solubility products and reactions used here as a guideline to saturation states are given in Table VI. The results of the calculations for phosphate compounds are plotted as -log lAP (lAP is the ion activity product) as a function of depth at each station (Figs. 49 and 50). Only data from box cores collected during 1975-1976 and some selected horizons from the gravity cores are shown. Hydroxyapatite was supersaturated by a factor of lO -lO at all stations and is not plotted precipitation of this phase is known to be kinetically hindered in seawater (Martens and Harriss, 1970). Bray (1973) and Norvell (1974) inferred likely equilibrium of pore waters with whitlockite [Ca3(P04)2] in Chesapeake Bay and anoxic lake sediments, respectively. Long Island Sound pore waters also tend to have activity products close to those predicted for saturation with respect to whitlockite, although distinct undersaturation is found for most sediment intervals at NWC. 

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