be required to determine an accurate estimate of what the actual increase would be. The lifetime for the l/e-inch mesh Cu screen should compensate for the initial investment. Copper sulfate soaks are being used once a week in an intake structure adjacent to the one used in these field studies. A copper ion concentration of 53 p.p.m. produced essentially the same results as those obtained in the above field studies. The cost of the copper sulfate used was approximately $20 per week. A major disadvantage to this technique, however, is that the intake structure must be shut down for an 8-hour period to apply the copper sulfate properly. The labor involved amounts to a minimum of $200 (direct labor costs) per application. If the pump in the auxiliary intake structure fails during application of the copper sulfate, the entire power unit will then be out until the unit receiving the soak can be put into operation. Chlorination also is being used by the power company that sponsored this study. Two tons of chlorine per week are used in intake structures identical in size to the one described in the example. This costs $10,400 per year for chlorine alone. The treatment is successful in killing hydroids and bryozoans, but has relatively little effect upon barnacles and oysters. Much higher concentrations of chlorine would violate local water standards. The copper electrolysis technique for prevention of marine fouling would probably be most useful to industries in areas whose water standards prohibit the required concentrations of antifouling chemicals. An improvement in this technique might be the use of a Cu alloy screen, the alloy being more noble than Cu, causing the Cu to corrode. Also, it is possible
to obtain the desired Cu dissolution rate by coupling a Cu screen to a screen more noble than it in the field water of interest. Acknowledgment The authors express appreciation to Gary Jacobs, New Orleans Public Service, for assisting with the field studies. Literature Cited Butler, G., Ison, H. C. K., “Corrosion and its Prevention in Water,” p. 26, Reinhold, New York, N. Y . , 1966. Crisp, D. J., in “Ecology and the Industrial Society,” G. T. Goodman, R. W. Edwards, and J. M. Lambert, Eds., pp. 99-117, Wiley, New York, N. Y . , 1965. Dethier, V. G., in “Copper Metabolism: A Symposium on Animal, Plant and Soil Relationships,” W. D. McElroy and B. Glass, Eds.. DU. _ - 154-73. Johns Houkins Press. Baltimore, Md., 1950. Galtsoff, P. S., “The American Oyster, Crassostrea cirginica Gmelin,” U. S . Department of Interior. Washington. - D. C.. Fishery Bull., Val.* 64, pp. 21-31 (1964). Iselin, C . O’D., Ketchum, B. H., Redfield, A. C., Eds., “Marine Fouling and Its Prevention,” pp. 1-388, U. S. Naval Institute, Prepared by Woods Hole Oceanographic Institution, Woods Hole, Mass., George Banta Publishing Co., Menasha, Wis., 1952. Pourbaix, H., “Atlas of Electrochemical Equilibrium,” p. 387, Pergamon Press, Long Island City, N. Y . , 1966. Tate, M. W., Clelland, R. C., “Nonparametric and Shortcut Statistics in the Social, Biological, and Medical Sciences,” pp. 62-4, 93-5, Interstate Printers and Publishers, Inc., Danville, Ill., 1957. ,
Receiced for reciew December 29, 1967. Accepted February 17, 1969. Most of this work was supported bj. fiindsprocided by New Orleans Public Serrice. Inc.
Calcium Sulfate Hemihydrate Scaling Thresholds in Natural Waters from 100” to 150 “ C . Julius Glater and Kochy Fung Department of Engineering, University of California, Los Angeles, Calif. 90024
Calcium sulfate hemihydrate, an important form of mineral scale, may deposit in saline water evaporators. This metastable crystal modification is more soluble than anhydrite (the stable phase above 100 “C.) but solubility of both forms decreases sharply with increasing temperature. Pure solid phase hemihydrate can remain in contact with its dissolved ions for about 17 hours at 100 “C., but the rate of conversion to anhydrite increases with the temperature. Conversion rates are also faster in solutions with high concentrations of background ions. Because of the relatively short “life” of hemihydrate in saline water backgrounds at elevated temperatures, it is difficult to measure its solubility by conventional methods of equilibration. An experimental technique for hemihydrate solubility measurement in saline water solutions at conditions approaching equilibrium is described. The method is based on visual observation of incipient crystal formation in a glass pressure vessel at various temperatures and pressures just above the saturation vapor pressure. Measurements were carried out on sea water and Roswell water over a range of 100” to 150 O C . Experimental data on scaling threshold are compared with literature calculations based on Debye-Huckel Theory. The rate of conversion of hemihydrate to anhydrite has also been studied. 580
Environmental Science & Technology
T
he rapid expansion of distillation technology has resulted in a recent focus of attention on the chemistry of calcium sulfate scale. Potential deposition of both hemihydrate and anhydrite scales places a severe limitation on evaporator performance in the high temperature range. Two research areas of manifold interest are studies of calcium sulfate solubility in saline water solutions and studies of phase transition kinetics between the major crystal modifications of this compound. There have been numerous publications on the solubility of gypsum and anhydrite in various salt solutions. Gypsum and anhydrite are the stable modifications at low and high temperatures, respectively. Since the classic work of Langelier, Caldwell, et al. (1950), however, little experimental work has been reported on hemihydrate which is metastable at all temperatures. Theoretical studies on the variation of cilcium sulfate solubility with temperature and ionic background have been reported by Marshall and Slusher (1968), Templeton and Rodgers (1967), and Power, Fabuss, et al. (1966). Outstanding among these is the elegant work of Marshall and Slusher. By application of extended DebyeHuckel theory, Marshall and Slusher have presented a scheme for calculating solubility limits for any form of calcium sulfate in all natural water backgrounds from 0 ” to 200 “C. Both hemihydrate and anhydrite scales have been observed in distilling plants under a variety of operating conditions.
The scale form depends on temperature, water composition, and heat transfer, whether it takes place under boiling or nonboiling conditions. According to Langelier, hemihydrate would be the expected scale form since it is metastable for about 17 hours at 100 “C. This is well beyond the expected residence time in any distilling plant. In practice, however, anhydrite has been observed in high temperature plants, and there is some concern about variation in the rate of phase transition of hemihydrate to anhydrite at higher temperatures. Power, Fabuss, er a / . (1966) and Johnson (1967) have shown that phase changes between calcium sulfate modifications take place more rapidly in solutions of high ionic strength. Johnson (1967) and Partridge (1930) observed that phase transition kinetics is strongly temperature dependent. Studies preliminary to the present work support these findings and show that transition of hemihydrate to anhydrite takes place approximately 40 times faster at 130 “C. than at 100 “C. Because of hemihydrate instability, especially at elevated temperatures, few investigators have chosen to study the equilibrium solubility of this important, hut elusive, scaling compound. Some experimental data in synthetic sea water has been reported by Lu and Fabuss (1968) and Marshall and Slusher (1968) over a limited temperature range. The authors’ experiments have been carried out with natural saline waters. Work with Marineland sea water over a period of years has consistently given reproducible experimental results. Marineland water is nearly identical in composition to sea water described by Svedrnp, Johnson, et ai. (1942). Good experimental data on hemihydrate solubility can be obtained by a semi-equilibrium visual method described in a previous paper (Glater, Ssutu, et a[., 1967). A visual device for calcium sulfate crystallization studies has also been reported by Shaffer and Knight (1967). The visual method provides fast reproducible results on scaling threshold. This improved lsboratnry technique can be carried out directly on feed brine samples and should provide useful data for engineering p-irposes. Apparatus and Experimental Procedure
The experimental portion of this study is divided into two parts. In the first, saline water is heated at constant volume and constant concentration factor (CF) until visible hemihydrate crystals appear in the bulk solution. In the second group of experiments, sea water solutions saturated with hemihydrate are maintained at constant temperature and concentration factor while samples are removed for chemical analysis. All experimental runs were carried out in a specially designed pressure bomb illustrated in Figure l . The glass vessel
Figure 1. Pressure vessel, head, Range, and gasket assemblies
was constructed from standard 2-inch borosilicate glass pipe approximately 5 inches long. The flared end was ground smooth and clamped by wing nuts to a a//s-inchstainless steel head. A special gasket composed of a compressed asbestos spacer and silicone rubber O-ring provided a good seal with finger pressure alone. The head was equipped with a pressure gage, pressurization port, release valve, thermocouple well, and sample removal port. The sampling line was fitted with a fine fritted glass filter and discharged 1hrough a stainless steel condenser. All metal parts contacting Ilot brine solutions were coated with silicone rubber for protection against corrosion. t* G < ..^: vd pa.l.6. The glass bomb has withstood press,.,, .” and appears to be satisfactory for solubility studies at temperatures up to 150 “C. Measurements of hemihydrate scaling thresholds were carried out by adjusting water samples to pH 4 with bydrochloric acid to prevent alkaline scale formation. Samples were then preconcentrated to some specific CF value less than 3.5 and filtered through a fine glass frit. A 150-ml. aliquot of each concentrate was transferrmI to the bomb. In prac tice, a stock CF 3.5 solution can be prepared and lower CF vdues can be obtained bv dilution with di stilled water. The bomlJ is eauinoed e o A ir with a magnetic stirring b;.s- nnrl +he h.Iuyu ,> fir mly secured with three wing nuts. Assembled apparatus is shown in Figure 2. The device was pressurized to about 10 p.s.i.g. with a nitrogen tank and checked for leaks. If pressure is maintained, a run may he started by stirring the sample and heating rapidly at first, then slowly (1 to 2 “C.per minute) as the crystallization end point is approached. When the threshold precipitation temperature for any C F is reached, a fine slurry of hemihydrate needles will be visible with the naked eye under strong illumination. The temperature is then recorded. Magnification or polarized light used in Thiele tube experiments (Glater, Ssutu, et a/., 1967) is not necessary for these high temperature runs. Sharp reproducible end-points were consistently obtained by this method. Several solid samples isolated from these determinations were identified as pure hemihydrate by x-ray diffraction. In a second series of experiments, samples of sea water at different concentration factors were saturated with reagent made gypsum at room temperature. The samples were filtered id 150-ml. aliquots transferred to the bomb which was then .essurized with nitrogen to prevent boiling. Heat was applied
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Figure 2. Pressure bomb assembly for visual end point Volume 3, Number 6,June 1969 581
through a glass-col heating mantle (Figure 3). The sample was stirred continuously and temperature was raised as rapidly as possible until a specified setting was reached. The temperature was then maintained constant ( + l "C.) by predetermined variac settings. When equilibrium was established with solid hemihydrate, filtered samples were withdrawn (via positive pressure) through the condenser jacketed with ice water. The time required for equilibration varies with temperature and CF. Coilversion of hemihydrate to anhydrite is very rapid at higher temperatures. Time values were roughly determined by x-ray analysis of selected samples held at constant temperature for varying time intervals. Hemihydrate solubility was calculated from analysis for calcium ion. Calcium was determined by standard EDTA titration at pH 13 using Cal-Blue indicator. A fluorescent end point was obtained under ultraviolet light. Adsorption of calcium ion on the magnesium hydroxide precipitate can be minimized by use of polyvinyl alcohol according to the method of Lott and Cheng (1959). By selecting certain concentration factors and analyzing for calcium at various temperatures, threshold limits for hemihydrate precipitation were established. A third series of experiments was run with samples of gypsum saturated sea water heated for long periods at constant temperature. The course of phase transformation between hemihydrate and anhydrite was followed by plotting calcium concentration DS. time. Equilibrium with hemihydrate occurs in the short horizontal region in the upper portion of Figure 7. Threshold precipitation limits for hemihydrate were calculated from these data. Results and Discussion Figure 4 is a phase diagram for calcium sulfate hemihydrate in Marineland sea water. Scale-free operation can be expected in the area below this line whereas precipitation of hemihydrate will occur above the line. Circles represent experimental data from visual end points. Square and triangular data points were derived from chemical analysis of solutions saturated with hemihydrate. This graph predicts a threshold C F value of 3.7 at the atmospheric boiling point, varying
PRESSURE INDICATOR-
-THERMOCOUPLE
considerably from the 3.2 value reported in an earlier paper (Glater, Ssutu, et al., 1967). This discrepancy can be explained by the different manner in which each set of experiments was run. The lower value, determined by Thiele tube, was obtained in a boiling sample which was progressivelyconcentrated, whereas the value reported in this work was derived at constant volume and CF. Lu and Fabuss (1968) have shown that calcium sulfate deposition depends strongly on operating conditions. They consistently observed lower threshold CF values in experiments involving evaporation than in heating runs. However, hemihydrate data shown in Figure 4 are most applicable for setting design limits in flash evaporators where heat transfer occurs under nonboiling conditions, assuming that residence times are short enough to avoid transition to anhydrite scale. The lower line in Figure 4 represents the often quoted work of Standiford and Sinek (1961). This plot was developed from fragmentary industrial data and places too severe a restriction on the operation of distilling plants. The other line was calculated by Marshall and Slusher (1968) from experimental data on sodium chloride and synthetic sea water solutions. This plot is remarkably parallel to the present experimental line. In their development, Marshall and Slusher assumed that increased solubility of calcium sulfate in sea water over sodium chloride solutions is due to the existence of stable weakly dissociated magnesium sulfate and calcium sulfate complexes (although the effect of neutral calcium sulfate species is absorbed into the solubility product constant). Sulfate ion may thus be tied up in species such as MgS04O and CaSOd0. This causes more calcium sulfate to dissolve to satisfy the solubility product constant. According to Templeton and Rodgers (1967), there are other important weakly dissociated species which may increase the solubility of calcium sulfate. They have used data in the 250 to 325 "C. temperature range to calculate association coefficients for CaS040, MgS04O, and Ca2SO4+2in the presence of high calcium concentrations and Ca(SO&-* in the presence of high sulfate concentrations. Complexing is most pronounced in solutions containing excess sulfate ion. This type of situation is represented by sea water which has a mole ratio of sulfate to calcium of about 3 to 1. Perhaps better agreement with Marshall and Slusher would have been achieved had they taken this factor into consideration. There are indeed many equilibria which must be considered in the very complex electrolyte solution represented by sea water.
.N PRESSURIZING \' 5 L I L EYG "CINT
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STAINLESS STEEL HEAD
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Figure 3. Details of pressure bomb assembly for equilibration studies 582 Environmental Science & Technology
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TEMPERATURE
Figure 4. Hemihydrate scaling thresholds of Marineland sea water concentrates
Figure 5 shows the scaling threshold line for hemihydrate in natural Roswell brackish water. All experimental data were obtained by visual end point. In this case, there is better agreement with the calculations of Marshall and Slusher. Roswell water has a much lower ionic strength than sea water, but the ion product of calcium sulfate is about the same in both solutions. The magnesium ion concentration of Roswell water is only about one-fourth of the sea water value, and calcium and sulfate occur in approximately equimolar quantities. Because of these differences, Roswell water has a lower scaling threshold and shows m9rz nearly ideal behavior with respect to Debye-Huckel Theory. Table I shows a comparison of the concentrations of scale forming ions in sea water and Roswell water. The chemical data points shown in Figure 4 were derived from analysis of sea water solutions saturated with hemihydrate. It was necessary to exercise caution in these experiments since the life of solid hemihydrate in contact with its aqueous solutions decreases very rapidly with temperature. As a guideline for this work, several samples of sold hemihydrate were equilibrated in sea water solutions for varying periods of time at 130 "C. The solid phase was then identified by x-ray diffraction. At this temperature, hemihydrate was stable for about 25 minutes. Samples collected beyond this time contained some anhydrite. Complete conversion to anhydrite was accomplished within three hours. Experimental data points for the constant C F runs were derived by preconcentrating four samples of sea water to C F values between 1.5 and 3. Each sample was saturated with gypsum at room temperature, and filtered. Portions were transferred to the bomb. As the temperature was raised, each solution became saturated with hemihydrate, and increasing amounts of solid phase were precipitated. Calcium analysis of filtered samples was carried out at several temperatures and compared with the original calcium concentration. At lower temperatures, the samples contained more calcium than the ambient calcium content of a given preconcentrated water. They are said to have a negative per cent precipitation. When the calcium content is greater than ambient, the sample is said to have a positive per cent precipitation. This technique is similar to that employed by Lu and Fabuss (1968), except that they measured positive degrees of precipitation only. The arbitrary levels of precipitation are based on augmentation or depletion of the calcium content of some ambient preconcentrated water. Per cent precipitation may be calculated from the following relationship :
Table I. Comparison of Scale Forming Ions in Sea Water and Roswell Water
Total dissolved solids, p.p.m. Calcium ion Magnesium ion Sulfate ion Bicarbonate ion CaS04 ion product (molar) SOa-2:Ca+zratio
Sea Water
Roswell Water
34,483 400 1,272 2,649 142
15,860 488 303 1,510 120
2.75 X 2.76:l
1.93 X 1.33:l
each plot with the 0% precipitation line. The four points represent a saturated solution of hemihydrate at the given temperature and CF. These points are replotted as triangles in Figure 4. The following equilibration times were used for each C F value: Equilibration Time, Hrs. 1.50
CF 3.0 2.5 2.0
1 .oo 0.50 0.25
1.5
To illustrate this method, consider C F 2 sea water. The initial concentration of calcium ion (prior to gypsum saturation) is 0.010 X 2 = 0.020molar. This concentration is taken as 0% precipitation. At 132 "C., the concentration of calcium in equilibrium with solid hemihydrate is found to be 0.022M. The solution contains a higher concentration of calcium than initially present, therefore, the per cent precipitation may be considered negative.
70 PPt.
0.020 - 0.022 o,020 x 100
=
=
-1070
This experimental method avoids the necessity of working with pure hemihydrate which is difficult to prepare and decomposes on standing. The method is also advantageous in the high temperature range where the life of hemihydrate is
CF 2 5
CF 3
C F 1.5
CF 2
Results of these experiments are illustrated in Figure 6. Threshold scaling temperature is read at the intersection of
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Figure 5. Hemihydrate scaling thresholds of Roswell brackish water concentrates
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Figure 6. Precipitation of hemihydrate from Marineland sea water concentrates saturated with calcium sulfate Volume 3, Number 6, June 1969 583
of the requisite assumptions mandatory in a theoretical evaluation of solubility. This work has stressed the need for accurate kinetic data on the transition of hemihydrate to anhydrite. The course of this transformation can be readily followed by simple chemical analysis for calcium ion. These experiments are being continued in an effort to gain further insight into phase transition kinetics.
CF 2 TEMP 135’C
Acknowledgment
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Figure 7. Transformation of hemihydrate to anhydrite in Marineland sea water saturated with calcium sulfate
The authors acknowledge the talents of E. K. Selover and D. J. Albright in making the experimental apparatus a reality. They are also grateful to Leo Morgan for x-ray diffraction studies and to J. W. McCutchan for his continuing support in this work. Thanks are due also to A. F. Pillsbury of the University of California Water Resources Center. Literature Cited
very short. Threshold values can be derived by extrapolation to the 0 % precipitation line from data obtained at lower temperatures. Long runs were also carried out with gypsum saturated solutions heated at constant temperature. In this series of experiments, an effort was made to match C F and temperature to certain regions of the visual hemihydrate threshold curve. Samples were withdrawn and analyzed for calcium at periodic intervals. A representative plot may be seen in Figure 7. The short horizontal region in the upper portion of the curve shows equilibrium with hemihydrate, whereas the lower horizontal line corresponds with anhydrite equilibrium. The identity of each solid phase was confirmed by x-ray diffraction. Threshold C F values were obtained by dividing the observed calcium content in the hemihydrate region by the calcium content of the preconcentrated water sample. Data points are plotted as squares in Figure 4. Conclusions
Good agreement between visual and chemical data have established clearly the reliability of this visual method for determination of hemihydrate scaling thresholds. Measurements can be carried out on any saline water within the temperature range described. This experimental method is free
584 Environmental Science & Technology
Glater, J., Ssutu, L., McCutchan, J. W., ENVIRON. SCI. TECHNOL. 1, 41 (1967). Johnson, D. R., Ph.D. thesis, University of Michigan, Ann Arbor, Mich., ORA Project C8967, November 1967. Langelier, W. F., Caldwell, D. H., Lawrence, W. B., Znd. Eng. Chem. 42, 126 (1950). Lott, P. F., Cheng, K. L., Chemist-Analyst 48, 13 (1959). Lu. C. H.. Fabuss. B. M.. Znd. Ene. Chem. Process Desien ” Develop.’ 7, 206 (1968). ’ Marshall, W. L., Slusher, R., J. Chem. Eng. Data 13, 83 (1968). Partridge. E. P.. Universitv of Michigan. Ann Arbor. Mich.. Research Bud. No. 15, i930. Power, W. H., Fabuss, B. M., Satterfield, C. N., J . Chem. Eng. Data 11, 149 (1966). Shaffer, L. H., Knight, R. A., ENVIRON.SCI. TECHNOL.1, 661 (1967). Standiford, F. C., Sinek, J. R., Chem. Eng. Progr. 57, 58 (1961). Svedrup, H. U., Johnson, M. W., Fleming, R. H., “The Oceans,” p. 166, Prentice-Hall, Englewood Cliffs, N. J., 1942. Templeton, C. C., Rodgers, J. C., J. Chem. Eng. Data 12, 536 (1967).
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Received for review August 5, 1968. Accepted January 23, 1969. This work was supported by State of California saline water research funds. These funds are administered by the Water Resources Center at the“University of California, Los Angeles, Calif. 90024
