Solubility Of Calcium Sulfate Research Articles

Solubility and Ksp of calcium sulfate: a general chemistry laboratory experiment

The experiment shared in this note can be used to accompany lectures on solubility or chemical equilibrium.

Composition of the solution in the hydration of cement

Tests on pure calcium hydroxide solutions have shown that solubility of calcium hydroxide decreases when alkali content increases. With addition of sulfate Ca 2+, concentration of the solution is higher than without it. Higher Ca 2+ concentration is not due to an oversaturation with Ca(OH) 2 but to the higher solubility of calcium sulfate. A similar result also was attained in solutions which are in contact with cement. In those cases Ca 2+ concentration in the solution decreased to the same extent the originally dissolved sulfate was precipitated as ettringite or monosulfate and finally it reached the value corresponding to the Ca(OH) 2 equilibrium. The solutions' pH value rose with increasing alkali content. This effect, however, clearly decreased from potassium to sodium to lithium. If lithium was added to a solution already containing potassium, the pH value decreased although OH-concentration of the solution continued to increase.

A thermodynamic model of CaSO 4 solubility in multicomponent aqueous solutions

A thermodynamic model is here developed to describe the solubility of calcium sulphate in salt water. A system of equations based on the ionic atmosphere theory of Debye and Hückel, Born model contribution, and local compositions of the nonrandom two-liquid (NRTL) model is used to describe isothermal activity coefficients for aqueous multicomponent solutions of partially or completely dissociated electrolytes. This approach is an extension of the model of Cruz and Renon for binary electrolyte solutions. Unlike previously published models, where the solubility behaviour of CaSO 4 depends only on the total dissolved solid (TDS) content, our model takes into account the nature and the amount of each ion which is present in the aqueous solution. Using only binary parameters, good agreement is obtained between the experimental and predicted values of calcium sulphate solubility in sea water and in natural and synthetic water (brackish brines), at 25 °C. By letting a limited set of binary parameters to be adjusted in the multicomponent case, we are able to calculate more accurately than previously the solubility behaviour of CaSO 4 over a wider range of salt compositions, including sea-water brines with magnesium augmentation.

HYDRAULIC DESIGN OPTIMIZATION OF ELECTRODIALYSIS FOR DEMINERALIZING BRACKISH WATERS

Abstract Conventional electrodialysis process design is largely empirical, requiring either experimental determination or previous experience with closely similar conditions. Formal cost minimization is not usually considered. This paper presents a non-linear programming model, with a posynomial cost function, subject to a number of hydraulic,-physicalchemical, and other constraints. The cost function includes terms proportioned to membrane area, D.C. power and pumping power. The hydraulic characteristics of electrodialysis spacers arc modelled as a packed-column, or Ergun-type expression, based on the summation of separate laminar and turbulent flow components. The laminar component is a modified Hagen-Poiseuille equation for flow through circular tubes; and the turbulent component is a modified Darcy-Weisbach equation for flow in closed conduits. Constraints arise from considerations of calcium sulfate solubility, limitations on available materials for spacers, and concentration (expressed in terms of l...

The solubility of calcium sulfate in aqueous solutions of sodium nitrate between 573 and 623 K

The solubility of calcium sulfate in aqueous solutions of sodium nitrate between 573 and 623 K

On the Equations Represented Solubility of Calcium Sulfate in Pure Water

On the Equations Represented Solubility of Calcium Sulfate in Pure Water

Calcium sulfate solubility in brackish water concentrates and applications to reverse osmosis processes; polyphosphate additives

The solubility of calcium sulfate dihydrate, which by precipitation can impede reverse osmosis (RO) processes, was determined at 25°C in three typical brackish water compositions and their concentrates. These waters corresponded to those used in U.S. Office of Saline Water (OSW) RO plants at Gillette, Wyoming, and Webster, South Dakota, and that of a post-irrigation water from the Wellton-Mohawk Canal in Arizona. Sodium hexametaphosphate was shown, when added in small amounts, to produce an apparent increase in the saturation concentration of CaSO 4. Later, this concentration decreased to the solubility in the absence of the additive. Calculations of maximum water recoveries and other parameters of direct application to RO processes are presented.

Solubility of calcium sulfate and association equilibria in CaSO 4 + Na 2SO 4 + NaClO 4 + H 2O at 273 to 623 K

The solubility of CaSO4·2H2O at 273 and 298 K and of CaSO4 at 423, 523, and 623 K has been determined in Na2SO4 + NaClO4 + H2O solutions at ionic strengths to 6 mol kg−1. From an interpretation of these results, various equilibrium quotients for CaSO4 have been obtained. These values are compared with similar quotients obtained previously in Na2SO4 + NaNO3 + H2O solutions over the same range of temperature and ionic strength and with values obtained in several mixed salt solutions at 298 K. Although a good correlation is observed for the solubilities in each system of mixed electrolytes, no easily discernible correlation among the various sets has yet been found.

Debye-Hückel correlated solubilities of calcium sulfate in water and in aqueous sodium nitrate and lithium nitrate solutions of molality 0 to 6 mol kg −1 and at temperatures from 398 to 623 K

In order to test further the applicability of the Debye-Hückel theory in describing solubility behavior at high temperatures, the solubility of calcium sulfate was determined in aqueous sodium nitrate and lithium nitrate solutions to 623 K. The solubilities are shown to adhere reasonably well to the extended Debye-Hückel theory at the experimental temperatures of 398 to 623 K to ionic strengths of 1 mol kg −1 , but with some divergences at higher ionic strengths. Comparative solubilities of calcium sulfate in water to 623 K are presented and discussed in an attempt to determine the correct values in water at zero ionic strength and at temperatures from 523 to 623 K.

Prediction of Deposition of Calcium Sulfate Scale Under Down-Hole Conditions

The thermodynamic method of calculating calcium sulfate scale has been improved by including the effects of complicated down-hole phase equilibria. Although the modified method has its limitations, it appears to provide more reliable results. Introduction Calcium sulfate (CaSO4) scale, especially gypsum, still causes unpredictable problems in many oil fields. This type of scale can cause severe plugging of equipment and producing formations; it creates the necessity of costly workovers and stimulation jobs. The best method for combating this scale seems to be the application of inhibitors - for example, by squeeze techniques - before the scale is formed. The reason is obvious: it is much more economical to prevent the formation of scale than to repair the damage caused by it. To apply the most suitable inhibition method, we have to know if and to what extent scaling occurs. At least some knowledge about the likely location of scale within the wall is equally important. The present methods of predicting scale in an oil field do not give all the answers, even though the literature contains enough data to develop and prove a more precise and reliable method. It is still common practice to wait for a production decline in the field, then pull the tubing and find - more or less accidentally - scale on the now-exposed metal surface. This scale sample is considered a proof for the occurrence of scale. This procedure can be very costly, and irreparable damage can be done before the scale is detected. Besides, this method fails frequently because scaling can occur even without the deposition of scale in the tubing. A better method for scale prediction is the use of water analyses and solubility data of CaSO4 compounds in the oilfield brines. There are many publications1–9,12,13,15,16 on the solubility of calcium sulfate compounds in water and brines at 1 atm. Most of the data in these publications are in good agreement. Together with the information from the water analyses, they are used to predict potential scale deposition. Diagrams, empirical equations, or thermodynamic relations are used to find the "maximum" solubility of calcium sulfate, mostly gypsum, in a water of similar composition.6–9,12–14 If the Ca++ and SO4-- concentrations in the actual water approximate the calculated "maximum" concentration of CaSO4, the field brine is said to have a "scaling tendency". We believe that this method can give only vague and unreliable information because it gives the solubility of CaSO4 or its hydrates at 1 atm only.

Diagenesis, Chemical Sediments, and the Mixing of Natural Waters

ABSTRACT Mixing of aqueous solutions is an important procedure in chemical laboratories. The chemical effects of mixing in the laboratory include changes in the concentration and electrical properties of a solution, shifts in homogeneous equilibria, and precipitation or dissolution of a solid phase. Similar chemical effects can be expected to occur in nature as a result of the mixing of natural waters, and in some instances these may result in diagenesis. A simple classification yields 14 distinct categories of natural waters with 91 possible combinations of mixing of pairs of such waters. Examples are known of the precipitation of hydroxides and carbonates as a result of the mixing of chemically dissimilar surface waters. Documented examples of subsurface mixing and diagenesis are few. However, because of the great variety of chemically dissimilar subsurface waters, plus the existence of several types of potential fields which cause the movement of subsurface waters, mixing and diagenesis in the subsurface may be a common and significant geologic process. The solubility of rock forming minerals is a non-linear function of such independent variables as salinity, partial pressure of gases, temperature, and so on. The dissolution of calcium carbonate as a result of Mischungskorrosion can be viewed in the light of the non-linearity of the solubility curve. Experimental data on the solubility of calcium carbonate and calcium sulfate as a function of added salts indicate that mixing of solutions which differ only in their content of dissolved electrolytes may cause either the precipitation or dissolution of the rock-forming minerals. Other common rock-forming minerals, with the exception of quartz at low temperatures, should exhibit similar behavior.

Moisture uptake by surface-applied superphosphate and movement of the phosphate and sulphate into the soil

Capillary uptake of soil water appeared to be more important than hygroscopic uptake of moisture by monocalcium phosphate in the wetting of surface-applied superphosphate particles. For hygroscopic uptake of water a relative humidity greater than 89% was found necessary, but particles on the soil surface are exposed to atmospheric humidity which is frequently well below 89%. This may prevent uptake of moisture from soils of low moisture content. The rate of moisture uptake from the vapour phase by superphosphate increased directly with increase in relative humidity above 89%. The rate of uptake was also influenced by temperature and surface area of the particle. Capillary uptake of soil water was influenced by soil moisture content, density of the particles, and the degree of contact between soil and particle. Movement of phosphate from the moist particle into the soil was rapid but the rate was affected by soil moisture, temperature, contact with the soil, and certain physical properties of the particle. In contrast to phosphate, movement of sulphate into the soil was slow because of the low solubility of calcium sulphate, and was greatly influenced by particle size. Leaching by rainfall may be an important factor in the movement of sulphate from coarse particles.

Effects of Brine Concentration and Pressure Drop on Gypsum Scaling in Oil Wells

Abstract Gypsum scale is a problem in oil wells that produce from reservoirs containing calcium sulfate. It becomes particularly troublesome in waterflood operations where large volumes of water are being produced. The solubility of gypsum or anhydrite in water increases with pressure due to a slight decrease in total volume as the salt dissolves. Upon release of pressure a supersaturated solution is obtained that deposits gypsum for a period of time. This is believed the most reasonable explanation for gypsum deposits at the bottom of producing wells. The study reveals that the amounts of scale formed at a given pressure drop and temperature depend on the amount of sodium chloride and other salts dissolved in the brine. The quantity of gypsum deposited increases with salt concentration to a maximum, then decreases until, with strong brines, no scale is formed. Introduction Secondary oil recovery from reservoirs, containing calcium sulfate often is hampered by gypsum scale. There are numerous causes of gypsum scale. Temperature changes can alter solubility and cause scale. Water incompatibility is also a common cause of gypsum deposition. If water containing calcium is mixed with another water containing sulfate, the solubility of calcium sulfate can be exceeded, resulting in the formation of gypsum scale. Another cause is water evaporation. When water evaporates the calcium sulfate concentration increases until the saturation concentration is reached; if the water continues to evaporate, the result is gypsum scale deposition. It has been suggested that the presence of large surface areas of solid calcium sulfate with relatively small volumes of water increases solubility, and thus causes calcium sulfate precipitation when the water flows into the wellbore. Gypsum scale also can be caused by pressure decreases, which cause a lowering of calcium sulfate solubility. Experiments show that there is a net loss of volume when anhydrite is dissolved in water (Appendix A). Calculations show that an increase in pressure would tend to shift the equilibrium toward an increase in anhydrite solubility (Appendix B). Experiments have confirmed this. Apparatus The experimental apparatus consisted of a 1-liter steel pressure bomb containing crushed anhydrite (mean diameter 0.5 in.). The bomb was kept in a drying oven for temperature control. The anhydrite was washed free of all fines with distilled water. The pressure was raised using a Blackhawk hand pump. Another steel bomb (0.5 liter) was placed between the pump and the bomb containing anhydrite to separate the pumping fluid (kerosene) from the water used in the experiments. This was done to prevent corrosion of the pump by water. A pressure gauge was placed between the separating reservoir and the pump. The pump and bombs were connected by 1/8-in. stainless steel tubing. After removing the water from the steel bomb following pressure release, the water was placed in 125-ml glass bottles that were covered to prevent evaporation. Small aliquots of water were removed from time to time to determine the amount of calcium sulfate remaining in solution. Procedure The bomb was filled with water, and the pressure was increased. During the first several hours after the pressure increase, the pressure decreased somewhat, and it was necessary to pump back to the desired pressure. This pressure loss was caused by the volume loss when the solid anhydrite dissolved in the water. After the first day the pressure remained constant, and it was not necessary to adjust for the remaining period at high pressure. The increased pressure was maintained for at least 5 days to permit the calcium sulfate solubility to reach equilibrium. Experiments showed that this was sufficient time for solubility maximum to be reached (Fig. 7). Anhydrite dissolves in water very fast initially, and continues to dissolve at a slowly increasing rate, reaching a maximum value in 5 to 7 days with no agitation (Fig. 7). Over 90 percent of the total anhydrite solubility is reached in 5 hours. The dissolution rate was ascertained by removing small amounts of water from the bomb and determining the amount of calcium sulfate in the water. EDTA (ethylene-diamine tetraacetic acid) titration and atomic absorption spectrophotometric procedures were used in these tests.

Prediction of Anhydrite Precipitation in Field Water-Heating Systems

Abstract A key step in feed water treatment for generating wet steam for thermal oil recovery is the removal of calcium and magnesium hardness by cation-exchange series softening. Knowing the solubility of any scale-forming salts in brines at elevated temperatures is necessary for fixing the level to which the feed water must be softened. Such calcium sulfate solubility data, previously not available above 392F, were determined by the authors in a flow equilibrium apparatus and will be reported elsewhere. These data were used to develop a method for predicting the solubility of anhydrite in hot water or steam droplets for saturated steam pressures as high as 2,000 psig (637F). (The calcium sulfate solubility product is represented by a combination of two factors, one reflecting the effects of ionic strength and the other accounting for the effects of complex ion formation in either calcium-magnesium-rich or sulfate-rich brines.) The method is applied to a calcium-magnesium-rich brine of moderately high salinity from a pilot hot-water flood, and to several sulfate-rich, low-salinity feed waters and blowdown (cooled steam droplets) samples from steam soak operations. The predicted calcium hardness levels corresponding to the calcium sulfate solubilities agreed reasonably well with the results of laboratory solubility determinations run on the field samples. Further testing of the method is needed for brines of other composition classes. Existing field cation exchange softeners in the cases tested are performing adequately since all the samples were found to be undersaturated with respect to calcium sulfate at their operating temperatures. Introduction Prevention of scaling caused by precipitation of calcium sulfate (anhydrite) is of considerable concern in connection with thermal recovery processes using wet steam or hot water. To avoid anhydrite precipitation in a heated system, an engineer must keep the product of the calcium and sulfate concentrations in the water or steam droplets below the value of the solubility product of anhydrite for the temperature and brine composition in question. Usually it is most practical to keep the concentration product lower than the solubility product by keeping calcium low in the presence of high sulfate, or by keeping sulfate low in the presence of high calcium. This can be done by a choice of combinations of natural waters and water treatment processes (such as series cation exchange softening to remove calcium). Until recently, few anhydrite solubility data, particularly for solutions containing other salts, were available for temperatures above 392F (211 psig steam); Marshall, Slusher and Jones studied the CaSO4-NaCl-H2O system up to 392F and surveyed the work of previous investigators. To model natural brines, one needs to study the solubility of anhydrite in aqueous solutions of sodium chloride. sodium sulfate, calcium chloride, magnesium chloride and their mixtures. Since steam pressures as high as 2,000 psig (637F) may be involved in thermal oil recovery projects. a solubility study was conducted between 482 and 617F. Discussed in this paper is the application of these data to the prediction of anhydrite precipitation in some practical steam soak and hot-water injection projects. Any simple method for predicting the solubility of an inorganic compound over a wide range of temperatures and solution compositions must be based on some assumptions, and therefore must yield approximate results. On the one hand, natural brines contain too many ionic species for all to be included in a simple scheme; on the other hand, there is no adequate theoretical basis for the exact prediction of solubility in even simple solutions of mixed electrolytes. However, it is possible at a given temperature to base a reasonable prediction scheme on two phenomena: the increase in solubility with increasing total concentrations of all ions (as measured by the ionic strength; see the Appendix). and an increase due to formation of complexes between calcium ions and sulfate ions and between sulfate and magnesium ions. Stiff and Davis developed such a scheme for predicting the solubility of gypsum (CASO4 * 2H2O) in brines at temperatures up to 212F. However, for the higher temperatures of present interest, the stable solid phase is anhydrite (CaSO4). This study involves a two-part method for predicting anhydrite solubility products. First, one predicts a value K, for a given ionic strength I and a given temperature T corresponding to mca/mso4= 1 from data of the CASO4- NaCl-H2O system. Second, one determines a group of factors F = F(Ca) * F(Mg) * F(SO4). where the individual factors account for increases in the solubility product due to complex formation by high concentrations, respectively, of calcium, magnesium and sulfate. Combining the two parts, one obtains the solubility product in molalities as

The solubility of calcium sulphate in sodium chloride and sea salt solutions

The results of experiments on the solubility of calcium sulphate in sodium chloride solutions and in simulated sea water solutions at temperatures of up to lOO°C are given. They are compared with those of similar experiments by other workers and found to be in satisfactory agreement. A tabulation of the best molal solubility products of CaSO 4 anhydrite in sea salt solutions for the temperature range from 50° to 200°C is given in the appendix.

Increased solubility of calcium sulphate in sea water containing hydrochloric acid

Data are presented which indicate that uptake of calcium sulphate by sea water brines can be increased from about 0.8% to 3 or 6% by aluminium chloride or hydrochloric acid additions respectively. The advantage of such increased solubility in postponement of calcium sulphate scaling and of descaling of sea water evaporation plants etc. is clearly indicated. A possible relationship between crystal structure and solubility is discussed and 'closed loop' descaling with the addition of hydrochloric acid is suggested as an alternative to mechanical chipping for removal of calcium sulphate deposits.

Calcium sulphate scale deposition in sea water evaporators

The importance of calcium sulphate scale in sea water evaporation is first described, both from an operational and from an economic point of view. A curve of the top temperatures of existing evaporators versus the corresponding concentration factor of the sea water is presented, below which safe operation of evaporators can be guaranteed. This practical curve is compared with the known solubility limits of calcium sulphate. Experiments are described in which the solubility of calcium sulphate in sea water is measured covering a range of temperatures between 250°F and 400°F and sea water concentration factors of 2 to 8. These new data are compared with information already available in the literature and extend the range of existing data. The experimental technique used in carrying out these measurements is also described. Experimental work designed to extend and clarify the practical scale limitation curve is then described. Scale build up was investigated in flash evaporators containing 3, 8 and 13 stages, multiple and single stage, operating over flash ranges between 3O°F and 130°F, concentrations between 2 and 8 times that of sea water, with and without seed re-cycle. The tendency of the evaporator heat transfer surface to scale is compared with the solubility data and the value of the scale limitation curve is confirmed with a seed free brine. Results with seed recycle show that the tolerable brine temperature tures can be greatly increased with supersaturated calcium sulphate deposited on the seeds in the condenser tubes and flash chambers. Seeding also inhibits scale build up by erosion of the tube walls but for seeding to be effective adequate residence time must be available. A theory of the effect of seed recycle is presented and an estimate of the reaction rate constant for the deposition of calcium sulphate on the seeds is given. The role of nucleation is mentioned. Finally, a description is given of a multiplicity of test rigs recently built to study the rate of deposition of calcium sulphate on heat transfer surfaces. These rigs are designed to circulate sea water brine in a closed loop under carefully controlled conditions and the scale deposition rate is measured on a spirally shaped heat surface which can be visually observed.

Equilibrium solubility of calcium sulphate hemihydrate in sodium chloride—magnesium sulphate solutions at elevated temperatures

AbstractResults are presented for the equilibrium solubility of calcium sulphate hemihydrate in water containing large quantities of dissolved sodium chloride and magnesium sulphate over the temperature range 200–300°F. The concentrations of sodium chloride and magnesium sulphate in the solution correspond to the quantities of these salts found in sea water.It has been found that the saturation solubility of calcium sulphate is dependent on the content of magnesium ions as well as the sulphate ions and dissolved sodium chloride in the solution. This information is of particular importance in the evaporation of sea water where the degree of prior precipitation of magnesium hydroxide will directly influence the later deposition of calcium sulphate on heat‐transfer surfaces. The use of the results in this connexion is discussed.

Ion strength, ion association, and solubility

This article describes two experiments in which attention is called to certain modifications of the ionic strength principle in applying the Debye-Huckel law expression to solubility data. The experiments are designed to demonstrate that the solubility of calcium sulfate increases in solutions containing certain 1-1 salts and to account for the increase by the prediction of theory and/or by the formation of complexes involving the calcium or sulfate ions.

An Improved Acid for Calcium Sulfate-Bearing Formations

Abstract An improved acid for the treatment of sulfate-containing limestones and dolomites is described. The acid is designed to reduce the reprecipitation of dissolved calcium sulfate and the possibility of plugging permeable flow channels. In addition, this improved acid has a much lower reaction rate than that of regular acid; the advantages of a "retarded" acid are obtainable. Field testing of the acid has shown it suitable for use in sulfate-containing formations. Substantial improvements in productivity generally resulted. Introduction Acid treatments of limestones, dolomites and other formations bearing carbonate deposits are frequently unsuccessful when the calcareous formation contains sulfate, either as anhydrite (CaSO4) or gypsum (CaSO4 2H2O). Preliminary dissolution in acid followed by redeposition of calcium sulfate appears to be a major factor contributing to poor well performance after acidizing. The precipitate is usually the gypsum form of calcium sulfate, but in higher temperature formations it may be anhydrite. The freshly precipitated crystals are nearly always very small and needlelike. They may occupy a gross volume many times that of the original anhydrite crystals and will obviously constitute an impediment to flow through newly enlarged flow channels. It is believed that the redeposition problem is most severe when anhydrite lines the fracture systems and large pores which supply the effective permeability of a formation. Microscopic inclusions of calcium sulfate also present large sulfate surface areas for dissolution in acid. In either case, great amounts of calcium sulfate may dissolve before the acid can be spent on formation carbonates. For regular spent acid (originally 15 per cent hydrochloric acid) the precipitate could be as much as 270 lb gypsum/1000 gal acid. Two techniques have been applied by the industry for reducing sulfate plugging during acidizing. The method commonly employed in the field is the attempted removal of a quantity of regular treating acid before it has reacted completely with the formation. This is practiced because the solubility of calcium sulfate is greater in a solution that is still acidic than in one which has been largely spent on the formation rock. The chance of precipitative plugging is therefore reduced if the withdrawal is successful. However, it is often impossible to get the acid out of the formation before precipitation occurs.