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Precipitation equilibrium during evaporation of water containing high dissolved salts
Desalination, 60 (1986) l-8 Elsevier Science Publishers B.V., Amsterdam - Printed in The Netherlands
PRECIPITATION EQUILIBRIUM DURING EVAPORATION CONTAINING HIGH DISSOLVED SALTS*
1
OF WATER
JASBIR S. GILL and RICHARD G. VARSANIK Calgon Corporation, Tel. 412-7777-8000;
P.O. Box 1346, Pittsburgh, PA 15230 (U.S.A.) telex 866439 CALGON PGH
(Received October 3,1985) SUMMARY
The kinetics of mixed precipitates have been studied using an evaporative cycling-up technique. In mixed precipitates with one common scale-forming ion, the effects of one mineral precipitation on the supersaturation, nucleation, and crystal growth of other subsequent precipitates have been discussed. The data indicate that in evaporative systems, precipitation of sparingly soluble minerals is not solely controlled by supersaturation, but also by the time involved in attaining the critical supersaturation. Thermodynamic equilibria should be used with caution for predicting a chemical treatment of systems encountering mixed scales. Keywords: Crystal growth, Kinetics, Mixed precipitates, scale/deposit. INTRODUCTION
The precipitation and adherence of scale-forming minerals, such as calcium carbonate, calcium phosphate, and calcium sulfate continues to present performance limits for evaporative desalination, boilers, and cooling tower systems. In natural water systems, attempts to describe mixed precipitation reactions using an equilibrium model have had varying degrees of success [ 1, 21. The difficulties are compounded, since conditions encountered in the real world are most likely removed from thermodynamic equilibria. In fact, true equilibrium cannot even be simulated in the laboratory in some cases. Still, chemical treatment of waters is often prescribed based on equilibrium calculations. Discrepancies between the observed (dynamic) and the calculated (equilibrium) composition of the water can become the cause of failure of the chemical treatment. Furthermore, the failure of many heterogeneous water treatment processes to reach equilibrium demonstrates the desirability of a kinetic analysis, rather than an equilibrium approach. The use of lower quality water, brackish makeup and higher cycles of *Presented in part at the 1985 Cooling Tower Institute Annual Meeting, January 20-23, New Orleans, Louisiana, U.S.A. OOll-9164/86/$03.50
0
1986 Elsevier Science Publishers B.V.
2
concentration is making the chemical treatment of cooling waters increasingly difficult [3]. The formation of mixed scales in a real system cannot be predicted with certainty using only thermodynamic equilibria. In an evaporative desalination plant and a high cycle cooling tower, water can be supersaturated with respect to several minerals simultaneously; however, the actual scale composition will be dependent on the relative supersaturations and kinetics of individual minerals. The mineral which precipitates first serves as a nucleus [4] for subsequent mineralization which could begin at a lower supersaturation than without the presence of such seeds. This paper presents the initial results of a study simulating a zero blowdown condition. A metastable makeup water is evaporated at fixed temperature and evaporation rate, while the volume of the concentrate is kept constant. An attempt is made to investigate the relative importance of kinetics of precipitation and supersaturation in terms of cycles of concentration calculated based on chemical analyses of the water.
EXPERIMENTAL
Throughout this study, distilled deionized water, analyzed reagent grade chemicals and Class A Pyrex glassware were used. A two-liter polypropylene flask with a side arm is placed in a tank of thermostated water to maintain the temperature of the evaporating solution. Temperature of the solution can also be maintained by immersing an electrically heated heat exchanger into the polypropylene flask. The temperature or heat flux is controlled by using a voltage regulator. This heat exchanger can provide an additional advantage of’ monitoring scale/deposit adherence on to any metal surface. A refractive index liquid level sensor is placed in the side arm to maintain a constant volume in the main flask by controlling a solenoid valve on the makeup water reservoir. Evaporation is achieved by passing filtered, dry air at a regulated and measured rate through a fine f&ted, teflon tube placed TA33LE I MAKEUP WATER Ion
Concentration
Calcium Magnesium Chloride Sulfate Bicarbonate Sodium Orthophosphate Lithium
240 40 568 4000 300 2033 5 5
(mg/l)
3 at the bottom of the flask. The makeup water can be concentrated at various rates by controlling the rate of aeration. The precise control of the rate of evaporation helps in the kinetics study by being able to reach a critical supersaturation in different time intervals. The makeup water shown in Table I was selected so that it is stable at room temperature and gives a sufficient induction time to establish the concentration process before any mineral precipitates. Aliquots are withdrawn at various time intervals, filtered and analyzed for pH, lithium, calcium, total alkalinity, orthophosphate, and chloride. Trace lithium concentration is measured to determine the cycles of evaporation. Carbon planchets, placed at the bottom of the flask prior to start of the experiments, are removed at various time intervals for scanning electron microscopy (SEM) and elemental analysis of the precipitates. Samples were also collected for X-ray diffraction by aborting runs at different times. RESULTS AND DISCUSSIONS
In Fig. 1 concentration of calcium ions in solution is plotted against the number of cycles of concentration, based on the lithium concentration in solution. The evaporation rate in these experiments was 0.850 l/24 h (L/D). Curve 1 is the calcium ion concentration experimentally determined in solution by analyzing a filtered aliquot of the sample. Curve 2 presents the
Fig. 1. Precipitation of the mixed scales at different cycles of concentration: calcium phosphate at 1.2 cycles, calcium carbonate at 1.6 cycles and calcium sulfate at 3.5 cycles. (1) Experimental total calcium in solution; (2) total calcium expected for no precipitation; (3) conditional equilibrium ionic product for calcite at different cycles.
expected calcium ion concentration in the solution in the absence of any precipitation. When the makeup water calcium is 240 mg/l, at two cycles of concentration, calcium level should be at 480 mg/l and so on. Thus, if the actual curve (Curve i) departs from Curve 2, it indicates precipitation. Curve 3 is the calculated equilibrium curve for calcium carbonate precipitation, which reflects the increase in equilibrium solubility [ 51 of calcium carbonate due to the increase in the total ionic strength of the solution during cycling-up. The upward swing for calcium in Curve 1, after the drop due to precipitation, could be explained on the basis of the increase in calcium ion pseudo-equilibrium concentration for an overall increase in ionic strength. In these experiments the major precipitant, up to 3.5 cycles, is calcite which has the greatest impact on calcium ion concentration in solution. The solution also precipitates calcium phosphate (Fig. 2), as determined by the phosphate concentration in solution. Calcium phosphate precipitation, which begins at much lower cycles, has a very small impact on the calcium ion concentration in solution. Curve 2, in Fig. 2, shows the calculated level of orthophosphate expected in solution in the absence of any calcium phosphate precipitation. Precipitation of calcium sulfate dihydrate (gypsum) did not start until after 3.5 cycles. Characterization of minerals precipitated and their morphology was done using SEM, EDXA, and X-ray diffraction. This analysis requires a sufficient amount of pure samples, especially without surface contamination. As pointed out earlier, under our experimental conditions, the precipitation of calcium phosphate is immediately followed by calcium carbonate precipitation. However, the relative amount of calcium phosphate precipitate is much smaller than the calcium carbonate precipitate. Thus, in order to determine the morphology and exact timing of calcium phosphate precipita-
No. of
CyCkS
Fig. 2. Precipitation of calcium phosphate as a result of cycling of the mixed system makeup water at 1.2 cycles.
5
tion, solid samples should be withdrawn before the initiation of calcium carbonate growth. Subsequent growth of calcium carbonate will cover the calcium phosphate crystals, masking them from detection. Carbon planchets are pulIed out at different times and analyzed for mineral identification and morphology characterization. Once calcium carbonate precipitation starts, it is almost impossible to find calcium phosphate, either by SEM or by elemental analysis of the solids by EDXA. In a mixed scale system, analysis of different precipitants is not the only problem. In fact, the entire processes of nucleation and crystal growth are altered. Crystal growth or precipitation from solution requires three steps [6] : a. supersaturation, b. nucleation, and c. growth of the nuclei. In mixed scale, especially when one of the scale-forming ions is common, all the three processes are altered compared to single component systems. Let us discuss first supersaturation for a mixed system of calcium phosphate, calcium carbonate, and calcium sulfate with calcium as the common ion. As in our experiments, the water becomes supersaturated with respect to all three scales, but calcium phosphate precipitates out first, followed by calcium carbonate and calcium sulfate, respectively. As soon as calcium phosphate starts precipitating, a small amount of calcium, phosphate, and hydrogen or hydroxyl ions are removed from solution as a solid phase (calcium phosphate). This changes the pH, ionic strength, and ionic product of the scale-forming ions and thus supersaturation for calcium carbonate [ 5, 71. Now, when calcium carbonate begins precipitating it encounters two different conditions for precipitation: nucleation/seeding, provided by the calcium phosphate precipitate; and a lower supersaturation as a result of calcium phosphate precipitation. Changes in these conditions can result in a marked difference in the precipitation rate and morphology of the growing crystals. Since the overall Gibb’s energy changes associated with the formation of critical nuclei for the heterogeneous nucleation are always smaller than those associated with homogeneous nucleation [S] , the precipitation of calcium carbonate in the presence of calcium phosphate seed crystals begins at lower supersaturation than in the absence of these seed crystals. Similarly, calcium carbonate precipitation affects the rate of calcium phosphate precipitation by lowering its effective supersaturation due to the depletion of calcium and carbonate ions to form calcium carbonate solids. When another scale mineral, with a common ion such as calcium in calcium sulfate, begins precipitating, its precipitation is affected by the previous precipitates. The kinetics of precipitation under the new conditions becomes the determining factor in controlling the scale, rather than the predictions based exclusively on thermodynamic equilibrium. Results shown in Fig. 3 are for a pure calcium carbonate system. In these
6
experiments both sulfate and orthophosphate ions in the makeup water were replaced by nitrate ions to maintain the same ionic strength as for the data in Fig. 1. To maintain approximately constant calcium ion activity, the calcium concentration had to be reduced in this experiment by an amount equal to that which was complexed by phosphate and sulfate in the earlier experiments. The effect of the seeding from calcium phosphate precipitation on calcium carbonate precipitation can be noticed by comparing Figs. 1 and 3. In Fig. 3, precipitation of calcium carbonate is observed after longer induction times and higher cycles of concentration than in Fig. 1. Thus for mixed scale systems, the process is more complex than can be solved by predicting the potential for scale formation using thermodynamic equilibria alone. For waters with high dissolved solids and high cycles of concentration through evaporation, crystallization phase diagrams provide a more accurate prediction of scaling or nonscaling conditions. Equilibrium predictive models are further plagued by phase transformations of various scaling minerals. Such transformations could result from the presence of impurities [9] or the experimental conditions pertaining to pH [lo], temperature [ll],ionic strength [12], and or different precursors [ 131. Many phases of the same mineral can give different characteristics, such as solubility [ 111, adherence to heat exchange surfaces, and response to inhibitors [141The kinetics of precipitation by cycling-up makeup water was evaluated using controlled evaporation at different rates. Three evaporation rates investigated were 0.85, 0.72, 0.55 L/D,. It was found that at lower evapora-
1
1
1
1.5
2.0
’ 2.5
‘3.0
No. of Cycles
Fig. 3. Precipitation of calcium carbonate as a result of cycling of the makeup water, containing only calcium carbonate scale forming ions, at 1.8 cycles compared to 1.6 cycles for the mixed system.
S
S
c
C
p
p Evaporaton
0.85
0.72
RateL/D
Fig. 4. Precipitation of the mixed scales for different evaporation rates; P, calcium phosphate, C, calcium carbonate, S, calcium sulfate.
tion rates precipitation began at lower supersaturations, but after longer induction times (Fig. 4) than for higher evaporation rates. At 0.85 L/D evaporation rate of the makeup water in Table I, calcium phosphate precipitated at 1.2 cycles of concentration (1.2 times the makeup water composition), calcium carbonate at 1.6 cycles of concentration, and calcium sulfate at 3.5 cycles of concentration. While at a slower evaporation rate of 0.72 L/D, calcium phosphate precipitate started at slightly lower cycles (1.18), however, after a longer induction time (9 h compared to 8.45 h at 0.85 L/D), and similar behavior was followed by calcium carbonate and calcium sulfate. These data indicate that precipitation of these sparingly soluble minerals is not entirely controlled by supersaturation, but also by the time involved in attaining the critical supersaturation. The data confirm predictions based on crystallization theory [ 151. The results presented in this paper demonstrate the importance of using phase diagram modeling for predicting precipitation in mixed scale systems. Thermodynamic equilibria should be used with great care for predicting a chemical treatment for systems encountering mixed scales. Prediction of scale formation using solely equilibrium calculations can be misleading in evaporative desalination processes and high cycle cooling waters, where mixed scales are often encountered. The study initiated in this paper, upon completion, will improve the prediction of chemical treatment for these applications.
ACKNOWLEDGEMENT
The authors are grateful to Ms. Cheryl Anderson for carrying out the experiments.
8 REFERENCES 1 A.E. Hill and J.H. Wills, J. Am. Chem. Sot., 60 (1938) 1647. 2 L.B. Yeats and W.L. Marshall, J. Chem. Eng. Data, 17 (1972) 163. 3 J.C. Cowan and D.J. Weintritt, Water-Formed Scale Deposits, Gulf, Houston, TX, 1976 pp. l-39. 4 J.S. Gill and G.H. Nancollas, Desalination, 29 (1979) 247. 5 T.F. Kazmierczak, Kinetics of Precipitation of Calcium Carbonate, Ph.D. Dissertation, State University of New York at Buffalo, New York, 1978, pp. 77-78. 6 J.S. Gill, Proceedings 41st International Water Conference, Pittsburgh, PA, (1980) pp. 191. 7 W. Stumm and J.J. Morgan, Aquatic Chemistry, Wiley, New York, 1970. 8 J.W. Mullin, Crystallization, 2nd edn., Butterworths, London, 197.2. 9 R.F. Reitmeier and T.F. Buebrek, J. Phys. Chem., 44 (1940) 535. 10 J.L. Meyer and C.C. Weatherall, J. Colloid Interface Sci., 89 (1982) 257. 11 G.H. Nancollas, A.E. Eralp and J.S. Gill, J. Sot. Pet. Eng., 18 (1978) 133. 12 Y. Kitano, Bull. Chem. Sot. Japan, 35 (1962) 1973. 13 B. Dickens and L.W. Schroeder, J. Res. NBS, 85 (1980) 347. 14 S. Sarig, F. Kohana and R. Leshem, Desalination, 17 (1975) 215. 15 A.E. Nielsen, Kinetics of Precipitation, Pergamon, Oxford, 1964.
PRECIPITATION EQUILIBRIUM DURING EVAPORATION CONTAINING HIGH DISSOLVED SALTS*
1
OF WATER
JASBIR S. GILL and RICHARD G. VARSANIK Calgon Corporation, Tel. 412-7777-8000;
P.O. Box 1346, Pittsburgh, PA 15230 (U.S.A.) telex 866439 CALGON PGH
(Received October 3,1985) SUMMARY
The kinetics of mixed precipitates have been studied using an evaporative cycling-up technique. In mixed precipitates with one common scale-forming ion, the effects of one mineral precipitation on the supersaturation, nucleation, and crystal growth of other subsequent precipitates have been discussed. The data indicate that in evaporative systems, precipitation of sparingly soluble minerals is not solely controlled by supersaturation, but also by the time involved in attaining the critical supersaturation. Thermodynamic equilibria should be used with caution for predicting a chemical treatment of systems encountering mixed scales. Keywords: Crystal growth, Kinetics, Mixed precipitates, scale/deposit. INTRODUCTION
The precipitation and adherence of scale-forming minerals, such as calcium carbonate, calcium phosphate, and calcium sulfate continues to present performance limits for evaporative desalination, boilers, and cooling tower systems. In natural water systems, attempts to describe mixed precipitation reactions using an equilibrium model have had varying degrees of success [ 1, 21. The difficulties are compounded, since conditions encountered in the real world are most likely removed from thermodynamic equilibria. In fact, true equilibrium cannot even be simulated in the laboratory in some cases. Still, chemical treatment of waters is often prescribed based on equilibrium calculations. Discrepancies between the observed (dynamic) and the calculated (equilibrium) composition of the water can become the cause of failure of the chemical treatment. Furthermore, the failure of many heterogeneous water treatment processes to reach equilibrium demonstrates the desirability of a kinetic analysis, rather than an equilibrium approach. The use of lower quality water, brackish makeup and higher cycles of *Presented in part at the 1985 Cooling Tower Institute Annual Meeting, January 20-23, New Orleans, Louisiana, U.S.A. OOll-9164/86/$03.50
0
1986 Elsevier Science Publishers B.V.
2
concentration is making the chemical treatment of cooling waters increasingly difficult [3]. The formation of mixed scales in a real system cannot be predicted with certainty using only thermodynamic equilibria. In an evaporative desalination plant and a high cycle cooling tower, water can be supersaturated with respect to several minerals simultaneously; however, the actual scale composition will be dependent on the relative supersaturations and kinetics of individual minerals. The mineral which precipitates first serves as a nucleus [4] for subsequent mineralization which could begin at a lower supersaturation than without the presence of such seeds. This paper presents the initial results of a study simulating a zero blowdown condition. A metastable makeup water is evaporated at fixed temperature and evaporation rate, while the volume of the concentrate is kept constant. An attempt is made to investigate the relative importance of kinetics of precipitation and supersaturation in terms of cycles of concentration calculated based on chemical analyses of the water.
EXPERIMENTAL
Throughout this study, distilled deionized water, analyzed reagent grade chemicals and Class A Pyrex glassware were used. A two-liter polypropylene flask with a side arm is placed in a tank of thermostated water to maintain the temperature of the evaporating solution. Temperature of the solution can also be maintained by immersing an electrically heated heat exchanger into the polypropylene flask. The temperature or heat flux is controlled by using a voltage regulator. This heat exchanger can provide an additional advantage of’ monitoring scale/deposit adherence on to any metal surface. A refractive index liquid level sensor is placed in the side arm to maintain a constant volume in the main flask by controlling a solenoid valve on the makeup water reservoir. Evaporation is achieved by passing filtered, dry air at a regulated and measured rate through a fine f&ted, teflon tube placed TA33LE I MAKEUP WATER Ion
Concentration
Calcium Magnesium Chloride Sulfate Bicarbonate Sodium Orthophosphate Lithium
240 40 568 4000 300 2033 5 5
(mg/l)
3 at the bottom of the flask. The makeup water can be concentrated at various rates by controlling the rate of aeration. The precise control of the rate of evaporation helps in the kinetics study by being able to reach a critical supersaturation in different time intervals. The makeup water shown in Table I was selected so that it is stable at room temperature and gives a sufficient induction time to establish the concentration process before any mineral precipitates. Aliquots are withdrawn at various time intervals, filtered and analyzed for pH, lithium, calcium, total alkalinity, orthophosphate, and chloride. Trace lithium concentration is measured to determine the cycles of evaporation. Carbon planchets, placed at the bottom of the flask prior to start of the experiments, are removed at various time intervals for scanning electron microscopy (SEM) and elemental analysis of the precipitates. Samples were also collected for X-ray diffraction by aborting runs at different times. RESULTS AND DISCUSSIONS
In Fig. 1 concentration of calcium ions in solution is plotted against the number of cycles of concentration, based on the lithium concentration in solution. The evaporation rate in these experiments was 0.850 l/24 h (L/D). Curve 1 is the calcium ion concentration experimentally determined in solution by analyzing a filtered aliquot of the sample. Curve 2 presents the
Fig. 1. Precipitation of the mixed scales at different cycles of concentration: calcium phosphate at 1.2 cycles, calcium carbonate at 1.6 cycles and calcium sulfate at 3.5 cycles. (1) Experimental total calcium in solution; (2) total calcium expected for no precipitation; (3) conditional equilibrium ionic product for calcite at different cycles.
expected calcium ion concentration in the solution in the absence of any precipitation. When the makeup water calcium is 240 mg/l, at two cycles of concentration, calcium level should be at 480 mg/l and so on. Thus, if the actual curve (Curve i) departs from Curve 2, it indicates precipitation. Curve 3 is the calculated equilibrium curve for calcium carbonate precipitation, which reflects the increase in equilibrium solubility [ 51 of calcium carbonate due to the increase in the total ionic strength of the solution during cycling-up. The upward swing for calcium in Curve 1, after the drop due to precipitation, could be explained on the basis of the increase in calcium ion pseudo-equilibrium concentration for an overall increase in ionic strength. In these experiments the major precipitant, up to 3.5 cycles, is calcite which has the greatest impact on calcium ion concentration in solution. The solution also precipitates calcium phosphate (Fig. 2), as determined by the phosphate concentration in solution. Calcium phosphate precipitation, which begins at much lower cycles, has a very small impact on the calcium ion concentration in solution. Curve 2, in Fig. 2, shows the calculated level of orthophosphate expected in solution in the absence of any calcium phosphate precipitation. Precipitation of calcium sulfate dihydrate (gypsum) did not start until after 3.5 cycles. Characterization of minerals precipitated and their morphology was done using SEM, EDXA, and X-ray diffraction. This analysis requires a sufficient amount of pure samples, especially without surface contamination. As pointed out earlier, under our experimental conditions, the precipitation of calcium phosphate is immediately followed by calcium carbonate precipitation. However, the relative amount of calcium phosphate precipitate is much smaller than the calcium carbonate precipitate. Thus, in order to determine the morphology and exact timing of calcium phosphate precipita-
No. of
CyCkS
Fig. 2. Precipitation of calcium phosphate as a result of cycling of the mixed system makeup water at 1.2 cycles.
5
tion, solid samples should be withdrawn before the initiation of calcium carbonate growth. Subsequent growth of calcium carbonate will cover the calcium phosphate crystals, masking them from detection. Carbon planchets are pulIed out at different times and analyzed for mineral identification and morphology characterization. Once calcium carbonate precipitation starts, it is almost impossible to find calcium phosphate, either by SEM or by elemental analysis of the solids by EDXA. In a mixed scale system, analysis of different precipitants is not the only problem. In fact, the entire processes of nucleation and crystal growth are altered. Crystal growth or precipitation from solution requires three steps [6] : a. supersaturation, b. nucleation, and c. growth of the nuclei. In mixed scale, especially when one of the scale-forming ions is common, all the three processes are altered compared to single component systems. Let us discuss first supersaturation for a mixed system of calcium phosphate, calcium carbonate, and calcium sulfate with calcium as the common ion. As in our experiments, the water becomes supersaturated with respect to all three scales, but calcium phosphate precipitates out first, followed by calcium carbonate and calcium sulfate, respectively. As soon as calcium phosphate starts precipitating, a small amount of calcium, phosphate, and hydrogen or hydroxyl ions are removed from solution as a solid phase (calcium phosphate). This changes the pH, ionic strength, and ionic product of the scale-forming ions and thus supersaturation for calcium carbonate [ 5, 71. Now, when calcium carbonate begins precipitating it encounters two different conditions for precipitation: nucleation/seeding, provided by the calcium phosphate precipitate; and a lower supersaturation as a result of calcium phosphate precipitation. Changes in these conditions can result in a marked difference in the precipitation rate and morphology of the growing crystals. Since the overall Gibb’s energy changes associated with the formation of critical nuclei for the heterogeneous nucleation are always smaller than those associated with homogeneous nucleation [S] , the precipitation of calcium carbonate in the presence of calcium phosphate seed crystals begins at lower supersaturation than in the absence of these seed crystals. Similarly, calcium carbonate precipitation affects the rate of calcium phosphate precipitation by lowering its effective supersaturation due to the depletion of calcium and carbonate ions to form calcium carbonate solids. When another scale mineral, with a common ion such as calcium in calcium sulfate, begins precipitating, its precipitation is affected by the previous precipitates. The kinetics of precipitation under the new conditions becomes the determining factor in controlling the scale, rather than the predictions based exclusively on thermodynamic equilibrium. Results shown in Fig. 3 are for a pure calcium carbonate system. In these
6
experiments both sulfate and orthophosphate ions in the makeup water were replaced by nitrate ions to maintain the same ionic strength as for the data in Fig. 1. To maintain approximately constant calcium ion activity, the calcium concentration had to be reduced in this experiment by an amount equal to that which was complexed by phosphate and sulfate in the earlier experiments. The effect of the seeding from calcium phosphate precipitation on calcium carbonate precipitation can be noticed by comparing Figs. 1 and 3. In Fig. 3, precipitation of calcium carbonate is observed after longer induction times and higher cycles of concentration than in Fig. 1. Thus for mixed scale systems, the process is more complex than can be solved by predicting the potential for scale formation using thermodynamic equilibria alone. For waters with high dissolved solids and high cycles of concentration through evaporation, crystallization phase diagrams provide a more accurate prediction of scaling or nonscaling conditions. Equilibrium predictive models are further plagued by phase transformations of various scaling minerals. Such transformations could result from the presence of impurities [9] or the experimental conditions pertaining to pH [lo], temperature [ll],ionic strength [12], and or different precursors [ 131. Many phases of the same mineral can give different characteristics, such as solubility [ 111, adherence to heat exchange surfaces, and response to inhibitors [141The kinetics of precipitation by cycling-up makeup water was evaluated using controlled evaporation at different rates. Three evaporation rates investigated were 0.85, 0.72, 0.55 L/D,. It was found that at lower evapora-
1
1
1
1.5
2.0
’ 2.5
‘3.0
No. of Cycles
Fig. 3. Precipitation of calcium carbonate as a result of cycling of the makeup water, containing only calcium carbonate scale forming ions, at 1.8 cycles compared to 1.6 cycles for the mixed system.
S
S
c
C
p
p Evaporaton
0.85
0.72
RateL/D
Fig. 4. Precipitation of the mixed scales for different evaporation rates; P, calcium phosphate, C, calcium carbonate, S, calcium sulfate.
tion rates precipitation began at lower supersaturations, but after longer induction times (Fig. 4) than for higher evaporation rates. At 0.85 L/D evaporation rate of the makeup water in Table I, calcium phosphate precipitated at 1.2 cycles of concentration (1.2 times the makeup water composition), calcium carbonate at 1.6 cycles of concentration, and calcium sulfate at 3.5 cycles of concentration. While at a slower evaporation rate of 0.72 L/D, calcium phosphate precipitate started at slightly lower cycles (1.18), however, after a longer induction time (9 h compared to 8.45 h at 0.85 L/D), and similar behavior was followed by calcium carbonate and calcium sulfate. These data indicate that precipitation of these sparingly soluble minerals is not entirely controlled by supersaturation, but also by the time involved in attaining the critical supersaturation. The data confirm predictions based on crystallization theory [ 151. The results presented in this paper demonstrate the importance of using phase diagram modeling for predicting precipitation in mixed scale systems. Thermodynamic equilibria should be used with great care for predicting a chemical treatment for systems encountering mixed scales. Prediction of scale formation using solely equilibrium calculations can be misleading in evaporative desalination processes and high cycle cooling waters, where mixed scales are often encountered. The study initiated in this paper, upon completion, will improve the prediction of chemical treatment for these applications.
ACKNOWLEDGEMENT
The authors are grateful to Ms. Cheryl Anderson for carrying out the experiments.
8 REFERENCES 1 A.E. Hill and J.H. Wills, J. Am. Chem. Sot., 60 (1938) 1647. 2 L.B. Yeats and W.L. Marshall, J. Chem. Eng. Data, 17 (1972) 163. 3 J.C. Cowan and D.J. Weintritt, Water-Formed Scale Deposits, Gulf, Houston, TX, 1976 pp. l-39. 4 J.S. Gill and G.H. Nancollas, Desalination, 29 (1979) 247. 5 T.F. Kazmierczak, Kinetics of Precipitation of Calcium Carbonate, Ph.D. Dissertation, State University of New York at Buffalo, New York, 1978, pp. 77-78. 6 J.S. Gill, Proceedings 41st International Water Conference, Pittsburgh, PA, (1980) pp. 191. 7 W. Stumm and J.J. Morgan, Aquatic Chemistry, Wiley, New York, 1970. 8 J.W. Mullin, Crystallization, 2nd edn., Butterworths, London, 197.2. 9 R.F. Reitmeier and T.F. Buebrek, J. Phys. Chem., 44 (1940) 535. 10 J.L. Meyer and C.C. Weatherall, J. Colloid Interface Sci., 89 (1982) 257. 11 G.H. Nancollas, A.E. Eralp and J.S. Gill, J. Sot. Pet. Eng., 18 (1978) 133. 12 Y. Kitano, Bull. Chem. Sot. Japan, 35 (1962) 1973. 13 B. Dickens and L.W. Schroeder, J. Res. NBS, 85 (1980) 347. 14 S. Sarig, F. Kohana and R. Leshem, Desalination, 17 (1975) 215. 15 A.E. Nielsen, Kinetics of Precipitation, Pergamon, Oxford, 1964.