June 14, 2007
Interactions Between Chloride and Sulfate or Silica Removals From Wastewater Using an Advanced Lime-Aluminum Softening Process: Equilibrium Modeling
By Abdel-Wahab, Ahmed Batchelor, Bill
ABSTRACT: Interactions among chloride, sulfate, and silica removals from recycled industrial wastewater using an ultra-high lime with aluminum process (UHLA) were studied. An equilibrium model that is able to accurately predict the chemical behavior and interactions between chloride and sulfate or silica with UHLA at various initial conditions and chemical reagents was developed. X- ray diffraction (XRD) analysis was conducted to identify the precipitated solids formed in the UHLA process. Model predictions indicated that simultaneous removal of sulfate and chloride can be best described by the formation of a solid solution containing calcium chloroaluminate, calcium sulfoaluminate (ettringite), calcium monosulfate, tricalcium hydroxyaluminate, and tetracalcium hydroxyaluminate. However, simultaneous removal of silica and chloride can be best described by precipitation of calcium silicate and calcium aluminosilicate in addition to a solid solution containing calcium chloroaluminate, tricalcium hydroxyaluminate, and tetracalcium hydroxyaluminate. The XRD results indicated the presence of the same solids assumed by the equilibrium model. Water Environ. Res., 79 528 (2007). KEYWORDS: sulfate removal, silica removal, chloride removal, industrial wastewater, softening, solid solution, cooling water, brine, equilibrium modeling.
(ProQuest-CSA LLC: ... denotes formulae omitted.)
The ultra-high lime with aluminum process (UHLA) is an advanced softening process that has excellent potential for improving industrial water use efficiency and minimizing wastewater discharge in water and wastewater treatments systems, such as recycled cooling water, pretreatment before membrane desalination, or brine treatment. It has the ability to remove sulfate, silica, and chloride, in addition to the majority of scale-forming compounds, such as calcium (Ca2+), magnesium (Mg2+), carbonate (CO32"), and phosphate (PO43"), which limit the extent of recycling (Abdel-Wahab and Batchelor, 2002, 2003; Abdel-Wahab et al., 2002; Batchelor et al., 1985).
Equilibrium experiments were conducted to study interactions between chloride and sulfate or silica removals with the UHLA process using lime [Ca(OH)2] and sodium aluminate (NaAlO2) as chemical reagents, while using sodium chloride (NaCl), sodium sulfate (Na2SO4), and sodium silicate (Na2SiO3) as chloride, sulfate, and silica sources, respectively. The results of these experiments indicated that the formation of calcium sulfoaluminate and calcium monosulfate is more favorable than the formation of calcium chloroaluminate. Therefore, chloride concentration was found to have a negligible effect on sulfate removal with UHLA. On the other hand, increased sulfate concentrations resulted in decreasing the removal efficiency of chloride, even at high stoichiometric doses of lime and sodium aluminate with respect to (initial chloride concentration + initial sulfate concentration) in the solution. This was explained by an increase in the activity of calcium chloroaluminate solid in the solid solution in the presence of sulfate. Silica was found to have a small effect on chloride removal with UHLA. Similarly, chloride was found to have a negligible effect on silica removal with UHLA.
The presence of calcium and aluminum and high pH values found in the UHLA system allow for the formation of aluminoferrite mono (AFm) and aluminoferrite tri (AFt) solid phases. The AFm and AFt phases are calcium derivatives of a big family of layered materials called layered double hydroxides (LDHs) and are composed of positively charged main layers of composition calcium aluminum hydroxide [Ca2Al (OH)e]+ and negatively charged interlayers of composition [X~ and c number of water molecules, nH2O], where X" is a mineral anion, such as hydroxide (OH~), chloride (Cl~), bromide (Br"), iodide (A), nitrate (NO3~), sulfate (SO42""), and carbonate (CO32") (Birnin- Yauri and Classer, 1998; Classer et al., 1999; Rapin et al., 2002; Renaudin and Francois, 1999; Renaudin etal, 2000; Renaudin, Francois, and Evrard, 1999; Renaudin, Kubel, Rivera, and Francois, 1999). The interlayer anions in an AFm or Aft phase are loosely held by electrostatic forces. Hence, it is possible for anions in the interlayer to exchange with anions in the aqueous solution (Birnin- Yauri, 1993), and a solid solution containing a mixture of solids is formed (Birnin-Yauri, 1993; Bimin-Yauri and Classer, 1998; Classer et al., 1999; Pollmann, 1986; Stronach, 1996). The removal of anions in UHLA process is not additive, but depends on the relative affinity of these anions to bind in the interlayer space of the AFm/ Aft solids and the possibility of a solid solution formation.
In an effort to further understand the interactions between chloride and sulfate or silica removals and develop a tool to describe the chemical behavior in the UHLA process, a fundamental chemical equilibrium model of the chemical processes was developed. The purpose of this paper is to describe this equilibrium model. Expectations are that final concentrations and chemical behavior in treated water or wastewater can be predicted knowing the chemical doses and initial concentrations in the feed water.
Materials and Methods. Equilibrium experiments to evaluate interactions among the processes removing chloride, sulfate, and silica were conducted by adding dry lime [Ca(OH)2] and dry sodium aluminate (NaAlO2) to solutions of sodium chloride (NaCl), sodium sulfate (Na2SO4), sodium silicate (Na2SiO3), NaCl + Na2SO4, or NaCl + Na2SiO3 in sealed plastic bottles that were used as completely mixed batch reactors. All solutions were prepared with decarbonated deionized water purified with a Barnstead Nanopure system (Barnstead International, Dubuque, Iowa) to greater than 18 IU. After the addition of chemical reagents, the reactors were rapidly closed, tightly sealed, and mixed by continuous shaking for 2 days at room temperature (23 to 250C) using a controlledtemperature shaker. Samples were taken and filtered through 0.450 -iec membrane filters. The filtrates were analyzed for calcium using atomic absorption spectrophotometry, for aluminum with UV spectrophotometry (eriochrome cyanine method; APHA et al., 1995), for chloride and sulfate using ion chromatography, and for silica with UV spectrophotometry (molybdosilicate method; APHA et al., 1995). The pH of a sample was determined before filtration using a pH meter with a combination glass electrode standardized with pH 10.00 and pH 13.00 buffers.
X-ray diffraction (XRD) spectroscopy was used to identify the solid phases formed in precipitation experiments. The solids were allowed to precipitate using the same procedures used in equilibrium experiments. The experimental conditions that were used in the equilibrium experiments to form the solids at different combinations or initial concentrations of chemical compounds (chloride, calcium, aluminum, sulfate, and silicate) are listed in Table 1. The precipitated solids were collected by centrifugation. Then, the separated solids were dried at room temperature (23 to 250C) in a carbon-dioxide-free atmosphere. The solids were scanned between O and 80 degrees 2E with a scan speed of 2 degrees/min by a Rigaku automated diffractometer using copper potassium alpha (CuKa) radiation (Rigaku, Woodlands, Texas).
Model Development A fundamental model of the chemical processes in UHLA was developed to predict final concentrations and chemical behavior in treated water using information on the chemical doses and initial concentrations in the feed water. Precipitation was assumed to be the mechanism that controls the solubility of the species in the system. The model was based on the geochemical modeling software, PHREEQC (Parkhurst and Appelo, 1999). The PHREEQC is a computer program written in the C programming language that is designed to perform a wide variety of low-temperature aqueous geochemical calculations. It is based on an ion-association aqueous model and has capabilities for (1) speciation and saturation-index calculations; (2) batch-reaction and onedimensional transport calculations involving reversible reactions, which include aqueous, mineral, gas, solid-solution, surfacecomplexation, and ion-exchange equilibria, and irreversible reactions, which include specified mole transfers of reactants, kinetically controlled reactions, mixing of solutions, and temperature changes; and (3) inverse modeling, which finds sets of mineral and gas mole transfers that account for differences in composition between waters, within specified compositional uncertainty limits. The flexible format of the PHREEQC input file allows models to be built and used to simulate a wide variety of aqueous-based scenarios.
Total initial concentrations and chemical doses for each experimental condition were used as input data in the PHREEQC input file. The database of PHREEQC was modified to include new solids and their solubility products, new aqueous species, and solidsolution formations. Table 2 shows the solubility products (Ksp) for new solids that were included in PHREEQC database using the following stoichiometric reactions: ... (1)
Assumptions that pure solids were formed and that solid solutions were formed were tested. To evaluate the equilibrium model, the model was used to predict final concentrations with PHREEQC for each hypothesis generated using initial solution compositions and equilibrium constants. Different hypotheses describing the chemical behavior in the UHLA process were generated and tested. The process of hypotheses generation was based on available information in the literature and preliminary analysis of experimental results.
Results and Discussion
Equilibrium Modeling for Sulfate-Hydroxide System. Three initial concentrations of sulfate (10, 50, and 100 mM) were investigated at each of three ratios of chemical doses to initial sulfate concentrations. Ratios of lime dose to initial sulfate concentrations were 1:1, 2:1, and 3:1, while the aluminum dose was chosen to be equal to 50% of the lime dose for all experiments. Different hypothesis were developed and tested. The standard error in the model-predicted concentrations for the dependent variables was compared for each hypothesis. Error analysis indicated that the experimental results can be best described by assuming a solid solution of calcium sulfoaluminate (ettringite), calcium monosulfate, tricalcium hydroxyaluminate, and tetracalcium hydroxyaluminate. The solution was found to be undersaturated, with respect to calcium sulfate, for all data points. A comparison between measured and model-predicted sulfate concentrations is shown in Figure 1 and indicates that the model succeeded in adequately predicting final sulfate concentrations.
Fractions of each solid in the solid solution were calculated for each data point using PHREEQC and are shown in Figure 2 as functions of lime and sodium aluminate doses. The fraction of monosulfate in the solid solution increased with increasing lime and sodium aluminate doses, while the fraction of ettringite decreased. The addition of lime and sodium aluminate resulted in increasing calcium, aluminum, and hydroxide concentrations in the solution, with respect to sulfate. This favors precipitation of calcium chloroaluminate over ettringite. Figure 2 shows that the fractions of tricalcium hydroxyaluminate and tetracalcium hydroxyaluminate are negligible for all data points. This indicates that precipitation of solids containing sulfate in the UHLA process is more favorable than that of solids containing hydroxide. These results agree with previous research on LDHs, which indicated that divalent anions, in general, have greater affinities to bind in the interlayer space of the LDHs than monovalent anions (De Roy et al., 1992; Rives, 2001).
The ratio of lime dose to aluminum dose is also important in controlling the fractions of solids in the solid solution and in the removal efficiency of sulfate. Ettringite contains a high ratio of calcium to aluminum. Therefore, if the ratio of lime dose to aluminum dose increased above the stoichiometric ratio of 2, this would thermodynamically favor the precipitation of ettringite over monosulfate and would result in increasing the fraction of ettringite in the solid solution, thus increasing the sulfate removal efficiency. The equilibrium model was used to simulate the effect of the ratio of lime dose to aluminum dose on sulfate removal. The PHREEQC was used with aluminum doses that varied and with lime doses that were maintained at a fixed ratio to the initial sulfate concentration. This ratio was set at the stoichiometric value for ettringite precipitation (i.e., 2). Figure 3 shows the effect of the ratio of lime dose to aluminum dose on the fractions of solids in the solid solution. The fraction of ettringite increased, and the fraction of monosulfate decreased, with increasing ratio of lime dose to aluminum dose. The fraction of monosulfate becomes negligible with respect to the fraction of ettringite, when the stoichiometric ratio of lime dose to aluminum dose approached the stoichiometric ratio of ettringite precipitation (i.e., 3.0). Therefore, to maximize the fraction of ettringite and maximize sulfate removal efficiency, the ratio of lime dose to aluminum dose should be at least 3.0.
Equilibrium Modeling of Chloride-Sulfate-Hydroxide System. Three initial concentrations of sulfate (10, 50, and 100 mM) were investigated at three initial concentrations of chloride (10, 50, and 100 mM). Each chloride-sulfate (Cl-SO4) combination was investigated at three molar ratios (1:1, 2:1, and 3:1) of lime dose to the sum of the initial concentrations of chloride and sulfate (Cl + SO4). Aluminum doses were chosen to be 50% of lime doses for all experiments. The same database file that was used to model the sulfate-hydroxide system was used to model the chloridesulfate- hydroxide system. Error analysis indicated that the results could be best described by assuming the formation of a solid solution containing calcium chloroaluminate, ettringite, monosulfate, tricalcium hydroxyaluminate, and tetracalcium hydroxyaluminate. Measured concentrations and model-predicted concentrations of sulfate, chloride, calcium, aluminum, and final pH are shown in Figures 4, 5, 6, 7, and 8, respectively. Figures 4 to 8 indicated that the model succeeded in adequately predicting chemical behavior during the simultaneous removal of chloride and sulfate in UHLA. However, the model moderately underestimated final chloride concentrations. The hypothesis was made that the increase of final chloride concentrations is a result of the increase of the activity of calcium chloroaluminate solid in the solid-solution formation in the presence of sulfate-containing solids. Stumm and Morgan (1996) reported that the solubility of a constituent in the solidsolution formation is greatly reduced, and its activity increases when it becomes a minor constituent of a solid-solution phase. The observed solubility product of calcium chloroaluminate solid was calculated by Gauss-Newton nonlinear regression routine using new program called INVRS K (Schwantez, 2002), which integrates the modeling power of PHREEQC with a nonlinear regression routine to calculate values of unknown or poorly defined chemical equilibrium and kinetic constants. Measured final chloride concentrations in the chloride- sulfate-hydroxide system were used as dependent variables. The observed solubility product of calcium chloroaluminate was obtained to be 10~26'42 compared with the value of 10~27- 10 reported by Bin- Yauri and Classer (1998) and was used in the original model. The new value of the solubility product of calcium chloroaluminate was used in the modified model to more accurately predict final chloride concentrations (Figure 9).
X-Ray Diffraction Results for Solids Formed in ChlorldeSulfate- Hydroxide System. Figure 10 shows diffractograms of solids precipitated in the chloride-sulfate-hydroxide system (XRDSO4). The XRD data for solids obtained in the chloridesulfate-hydroxide system were compared with Joint Committee for Powder Diffraction Studies (JCPDS) cards data (JCPDS, 1990) (Table 3) and indicated the presence of the same solid phases used in the model development (ettringite, monosulfate, calcium chloroaluminate, tricalcium hydroxyaluminate, and tetracalcium hydroxyaluminate). However, the XRD patterns also showed the presence of another solid phase, called Kuzel salt [Ca4Al2Cl (SO4)0.5(OH)i2] (Classer et al, 1999), which was not included in the equilibrium model. The solubility product of Kuzel salt was calculated to be 1(T2720 using data from Classer et al. (1999) and the following reaction:
Kuzel salt was included in the equilibrium model, but no significant changes were observed in the predicted final concentrations, even though the modified model predictions showed that Kuzel salt was formed as part of the solid solution. This was because the formula for Kuzel salt is an alternative expression of a solid solution formed by the combination of calcium chloroaluminate and monosulfate, according to the following equation:
The negligible changes in final predicted concentrations were the result of using a value for the solubility product of Kuzel salt (1CT27- 20) that was similar to the value that can be derived from the solubility products of calcium chloroaluminate and calcium monosulfate. The following equations can be derived using eqs 1, 5, and 8 and the definition of the solubility products:
Where Ksp_mos, Kv^uaiare the solubility products of calcium monosulfate and Kuzel salt, respectively. The following two equations can be derived by algebraically manipulating eq 12 with eqs 10 and 11:
By algebraically manipulating eqs 13 and 14, the solubility product of Kuzel salt can be calculated from the solubility products of calcium chloroaluminate and calcium monosulfate, as follows:
The difference between the two values of the solubility product of Kuzel salt could be a result of the effect of the solid-solution mechanism on the solubility of the solid.
Equilibrium Modeling of Silicate-Hydroxide-Chloride System. In the silicate-hydroxide-chloride system, three initial concentrations of silica (O, 1.5, and 3.0 mM) were investigated at four initial concentrations of chloride (O, 10, 50, and 100 mM). Each chloride- silicate (Cl-SiO3) combination was investigated at three molar ratios (1:1, 2:1, and 3:1 ) of lime dose to the sum of the initial concentrations of chloride and silicate (Cl + SiO3). Aluminum doses were chosen to be 50% of lime dose for all experiments. The hypothesis was made that the silicate removal could be described by the precipitation of calcium silicate and calcium aluminosilicate solids. The ion activity products of calcium silicate and calcium aluminosilicate were calculated from the following reactions, which are based on the asumption that calcium silicate and calcium aluminosilicate are the proper formulas for calcium silicate and calcium aluminosilicate solids, respectively: ... (16)
The activities of species were calculated from the measured total concentrations of calcium, aluminum, silicon, and pH by PHREEQC using the Davies Equation. The measured concentrations were used as input values to PHREEQC, and no precipitates were allowed to form. An average value of the values of the ion activity products (IAP) of HT8- 02*0- 5 for calcium silicate and eo23- 09^- 8 for calcium aluminosilicate were obtained. The average IAP of calcium silicate is consistent with the value of 10~7'8 obtained by Batchelor and McDevitt (1984). To test the hypothesis that silica removal can be described as the precipitation of both calcium silicate and calcium aluminosilicate and to model silica removal with UHLA, the chemical equilibrium model was modified to include silica reactions. The average values of the IAPs of calcium silicate and calcium aluminosilicate were used as their solubility products, and the model was used to predict final concentrations using the total initial concentrations of calcium, aluminum, silicon, chloride, hydroxide, and sodium as input. Figure 11 compares the measured and modelpredicted final silica concentrations over a range of chemical doses and initial chloride concentrations. These results indicate that the model succeeded in adequately predicting measured silica concentrations, which supports the hypothesis that silica removal with UHLA can be described as the precipitation of calcium silicate and calcium aluminosilicate.
X-Ray Diffraction Results for Solids Formed in SilicateHydroxide- Chloride System. Figure 12 shows diffractograms of solids precipitated in the silicate-hydroxide-chloride system (XRDSi). The XRD patterns show at least two peaks of calcium aluminosilicate (at d = 12.34 and 4.16 A), in addition to calcium chloroaluminate, tricalcium hydroxyaluminate, and tetracalcium hydroxyaluminate. The comparison of XRD data for calcium aluminosilicate with JCPDS cards data is shown in Table 4. Detection of aluminosilicate solid agrees with the hypothesis that calcium aluminosilicate is formed during UHLA. However, the presence of calcium silicate could not be confirmed. There were no peaks that matched the JCPDS cards for calcium silicate solids. This could be because, at such high chemical doses (60 mM lime and 30 mM sodium aluminate), the concentration of calcium silicate is negligible, and the dominant silica-containing solid phase is calcium aluminosilicate.
Interactions between chloride and sulfate or silica removals using the UHLA process were investigated. An equilibrium model was developed, and it succeeded in predicting the multicomponent removal and chemical behavior in the UHLA process at various initial concentrations and doses of reagents. The XRD analysis for precipitated solids in the UHLA process showed the presence of the same solids that were assumed by the equilibrium model. Equilibrium modeling and XRD results indicated that the chemical behavior of simultaneous removal of both chloride and sulfate can be best described as the formation of a solid solution containing calcium chloroaluminate, calcium sulfoaluminate, calcium monosulfate, tricalcium hydroxyaluminate, and tetracalcium hydroxyaluminate. The XRD patterns also showed the presence of another solid phase, called Kuzel salt, which was not included in the equilibrium model. Simultaneous removal of chloride and silica was found to be best described by the precipitation of calcium silicate and calcium aluminosilicate, in addition to the formation of a solid solution of calcium chloroaluminate, tricalcium hydroxyaluminate, and tetracalcium hydroxyaluminate. However, there were no peaks that matched the JCPDS cards for calcium silicate solids.
This project was funded, in part, with funds from the State of Texas as part of the program of the Texas Hazardous Waste Research Center. The contents do not necessarily reflect the views and policies of the sponsor, nor does mention of trade names or commercial products constitute endorsement or recommendations for use.
Submitted for publication October 7, 2005; revised manuscript submitted May 8, 2006; accepted for publication May 16, 2006.
The deadline to submit Discussions of this paper is August 15, 2007.
Abdel-Wahab, A.; Batchelor, B. (2002) Chloride Removal from Recycled Cooling Water Using Ultra-High Lime with Aluminum Process. Water Environ. Res., 74, 256.
Abdel-Wahab, A; Batchelor, B. (2003) Effects of Water Quality, pH, and Temperature on Chloride Precipitation with Ultra-High Lime with Aluminum Process. Proceedings of the 76th Annual Water Environment Federation Technical Exposition and Conference, Los Angeles, California; Oct. 11-15; Water Environment Federation: Alexandria, Virginia.
Abdel-Wahab, A.; Batchelor, B.; Schwantes, J. (2002) An Equilibrium Model for Chloride Removal from Recycled Cooling Water Using Ultra-High Lime with Aluminum Process. Proceedings of the 75th Annual Water Environment Federation Technical Exposition and Conference, Chicago, Illinois, Sept. 28-Oct. 2; Water Environment Federation: Alexandria, Virginia.
American Public Health Association; American Water Works Association; Water Environment Federation (1995) Standard Methods for the Examination of Water and Wastewater, 19th ed.; American Public Health Association: Washington, D.C.
Batchelor, B.; McDevitt, M. (1984) An Innovative Process for Treating Recycled Cooling Water. / Water PoIlM. Control Fed., 56 (10), 1110.
Batchelor, B.; McDevitt, M.; Chan, D. (1985) Removal of Sulfate from Recycled Cooling Water by the Ultra-High Lime Process. Proceedings of American Water Works Association Water Reuse Symposium 111, San Diego, California, Aug. 26-31, 1984; American Water Works Association Research Foundation: Denver, Colorado.
Birnin-Yauri, U. A. (1993) Chloride in Cement: Study of the System CaO-Al2O3-CaCl2-H2O, Ph.D. Thesis, University of Aberdeen, Scotland, United Kingdom.
Birnin-Yauri, U. A.; Classer, F. P. (1998) Friedel's Salt, Ca2Al(OH)6(Cl1OH) 2H2O: Its Solid-Solutions and Their Role in Chloride Binding, Cem. Concr.Res.,2S,n.
Damidot, D.; Classer, F. P. (1993) Thermodynamic Investigation of the CaO-Al2O3-CaSO4-H2O at 250C and the Influence of Na2O. Cem. Concr. Res., 23, 221.
De Roy, A.; Forano, C.; El Malki, K.; Besse, J. P. (1992) Anionic clays: Trends in Pillaring Chemistry. In Expanded clay and Other Micmporous Solids, VoI [Eth], Occelli, M. L., Robson, H. E. (Eds.); Van Nostrand Reinhold: New York, n 108-169.
Classer, F. P.; Kindness, A.; Stronach, S. A. (1999) Stability and Solubility Relationships in AFm Phases Part I. Chloride, Sulfate, and Hydroxide. Cem. Conor. Res.. 29, 861.
Joint Committee for Powder Diffraction Studies (1990) Powder Diffraction File, Inorganic Volume; Joint Committee for Powder Diffraction Studies: Swarthmore, Pennsylvania.
Nacken, R.; Mosebach, R. (1936) Untersuchungen an den Vierstoffsystemen CaO-Al2O3-CaCl2-H2O und CaO-SiO2-CaCl2-H2O (Investigations of Systems with Four Components of CaO-Al2O3-CaCl2- H2O and CaOSiO2-CaCl2-H2O). Zeitschriflfur anorganische und allgemeine Chemie (J. lnorg. Gen. Chem.) 228, 19 (in German).
Parkhrust, D. L.; Appelo, C. A. J. (1999) User's Guide to PHREEQC (Version 2)- A Computer Program for Speciation, Batch-Reaction, One- Dimensional Trasport, and Inverse Geochemical Calculations, U.S. Geological Survey Water-Resources Investigation Report 994259; U.S. Geological Survey: Denver, Colorado.
Pollmann, H. (1986) Solid Solution of Complex Calcium Aluminate Hydrates Containing Cl", OH" and CO32" Anions. Proceedings of the 8th International Symposium on the Chemistry of Cements, Vol. E[Eth], Rio de Janeiro, Brazil, Sept. 22-27; Brazil Association of Portland Cement: Rio de Janeiro, Brazil, 300.
Rapin, J. P.; Renaudin, G.; Elkaim, E.; Francois, M. (2002) Structural Transition of Friedel's Salt 3Ca).Al2O3.CaCl2.10H2O Studied by Synchrotron Powder Diffraction. Cem. Concr. Res., 32, 513.
Renaudin, G.; Francois, M. (1999) The Lamellar Double Hydroxide (LDH) Compound with Composition 3CaO Al2 O3 Ca(NO3) 1OH2O. Ada Crystallogr,, CSS, 838.
Renaudin, M.; Francois, G.; Evrard, O. (1999) Order and Disorder in the Lamellar Hydrated Tetracalcium Monocarboaluminate Compound. Cem. Concr. Res., 29, 63.
Renaudin, G.; Kubel, F.; Rivera, J. P.; Francois, M. (1999) Structural Phase Transition and High Temperature Phase Structure of Friedel's Salt, 3CaO.Al2O3.CaCl2.10H2O. Cem. Concr. Res., 29, 1937.
Renaudin, G.; Rapin, J. P.; Humbert, B.; Francois, M. (2000) Thermal Behavior of the Nitrated AFm Phase Ca4 Al2 (OH)i2 (NO3)2 . 4H2O and Structure Determination of the Intermediate Hydrate Ca4 Al2 (OH)12 (NO3)2 . 2H2O. Cem. Concr. Res., 30, 307.
Rives, V. (2001) Layered Double Hydroxides: Present and Future. Nova Science Publisher, Inc.: Huntington, New York.
Schwantez, J. M. (2002) Equilibrium, Kinetic, and Reactive Transport Models for Plutonium, Ph.D. Dissertation, Texas A&M University, College Station, Texas.
Stronach, S. A. (1996) Thermodynamic Modeling and Phase Relations of Cementitious Systems, Ph.D. Thesis, University of Aberdeen, Aberdeen, United Kingdom.
Stumm, W.; Morgan, J. J. (1996) Aquatic Chemistry: An Introduction Emphasizing Chemical Equilibria in Natural Waters, 3rd ed.; John Wiley & Sons, Inc.: New York.
Ahmed Abdel-Wahab1*, Bill Batchelor2
1 Visiting Assistant Professor, Texas A&M University at Qatar, College Station, Texas.
2 Professor, Civil Engineering Department, Texas A&M University.
* Texas A&M University at Qatar, P.O. Box B-6, College Station, TX 77844; e-mail: ahmed-abdelwaha
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