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INDUSTRIAL AND ENGINEERING CHEMISTRY NOMENCLATURE
Vol. 42, No. 1
LITERATURE CITED
A , B = coefficients in the van Laar activity coefficient equation expressed in the terms of the terminal values = fugacity of component i in homogeneous system fi = fugacity in the reference or standard state f: = the sth parameter in the relative volatility power series gC expression of component i with respect to j = the rth parameter for component i in the activity coef2; ficient, power series expression in terms of component j n = an integer represent,ing the total number of parameters in the relative volatility or activity coefficient power series expression P = pressure of system pi = partial pressure of component i in vapor phase pf = partial pressure in the standard state-i. e., the vapor pressure of pure component i r = an integer representing the number of a parameter in the activity coefficient power series expression s = an integer representing the number of a parameter in the relative volatility power series expression xi = 1/2 - xi zi = mole fraction of component i in homogeneous sytem, specifically the liquid phase yi = mole fraction of component i in vapor phase aij = relative volatility of component i compared to j as = aijforXi = 0 ai = ideal relative volatility = activity coefficient of component i in homogeneous sysyi tem = activity coefficient of component i in a binary system exyij pressed as a power series in terms of xi A = log aij - log a$
(1)
Carlson, H. C., and Colburn, A. P., IND.ENG.CHnhf., 34, 581
(1942). (2) Ferguson, J. B., J . Phys. Chem., 36, 1123 (1932). (3) Gibbs, J. W., “Collected Works,” 1701. 1, p. 55-92, New Haven, Yale University Press, 1948. (4) Gilmont, Roger, Ph. D. Thesis, Polytechnic Institute of Brooklyn, June 1947. (5) Horsley, L. H., IXD.END.CHEM., ANAL.ED.,19, 508 (1940). (6) Hovorka, F., et al., J . Am. Chem. Soc., 58, 2265 (1936). ( 7 ) Kirejew, V., ActaPhysicochim. U . R. S.S., 13,552 (1940). (8) Laar, J. J. van, 2. physik. Chem., 72, 723 (1910); 83, 599 (1913). (9) Margules, M., Ant. Akad. Wiss. Wien., Math. naturw. Klnsse, 104, 11, 1243 (1895). (10) Othmer, D. F., Iii-u. EKG.CHEM., 32, 841 (1940). (11) Othmer, D. F., and Gilmont, R., Ibid., 36, 858 (1944). (12) Scatchard, G., and Hamer, W. J., J . Am. Chem. Soc., 57, 1806 (1935). (13) Soatchard, G., and Raymond, C. L,, I b i d . , 60, 1278 (1938). (14) White, R. R., Trans. Am. Inst. Chem. Engrs., 41, 539 (1945). (15) Wohl, K., I b i d . , 42, 215 (1946). (16) Zawidski, J. v., 2. physik. Chem., 35, 129 (1900).
RECEIVED March 1, 1948. Presented before the Division of Physical and Inorganic Chemistry a t the 113th Meeting of the A h r E R I c A N CHEMICAL SOCIETY, Chicago, Ill. Based in part on a thesis submitted t o t h e faculty of the Polytechnic Institube of Brooklyn in partial fulfillment of the requirements for the degree of doctor of chemical engineering. F o r detailed paper (or extended version or material supplementary t o this article) order Document 2687 from American Documentation Institute, 1719 N St., N. W., Washington 6, D. C., remitting $0.50 for microfilm (images 1 inch high OD standard 35-mm+motion picture film) or S5.20 for photocopies (6 X 8 inches) readable without optical aid.
Scale Control in Sea Water Distillation Equipment CONTACT STABILIZATION W. F. LANGELIER, D. H. CAEDWELL, AND W. B. LAWRENCE Uniuersity of California, Berkeley, Calif.
C
C. H. SPAULDING
mainly by the employment of high purge, HEMICAL analyses of scale material, Fort Belvoir, Va. or overflow rates normally 50YG or more, removed from evaporators employed and by frequent chemical or mechanical in the distillation of sea water, indicate descaling operations. The necessity for descaling after 500 two distinct classifications-the scales are either predominantly hours of operation i s not unusual, and this constitutes an obcalcium carbonate or predominantly magnesium hydroxide and, jectionable feature of present methods of scale control. Boiler in both cases, calcium sulfate is usually a minor constituent. compounds, cornstarch, and other peptizing materials are used in Because of the generally uniform character of sea water from low pressure evaporators with favorable results in prolonging the whatever ocean i t is obtained, the differences in the type of scale intervals between descaling. Test runs, in thermocompression found in evaporators must be attributed to evaporator design or stills, by the Engineer Research and Development Laboratories to operating conditions. In general the calcium carbonate scales a t Fort Story, Va., beginning in 1943, indicated that greatly are found in low pressure evaporators and the magnesium hyincreased operating intervals could be obtained by removal of the droxide scales in thermocompression evaporators operating a t bicarbonate of the feedwater with strong acid; this treatment, atmospheric pressure. Boiling temperature rather than brine however, requires some degree of technical control and may be concentration appears to be the determining factor. Typical very inconvenient under field conditions. The studies mere not analyses of scale material removed from two types of evaporacarried sufficiently far to warrant the conclusion that corrosion tors are given in Table I. effects could be completely eliminated. Measures directed to overcoming the objectionable effects of A rational approach to the problem of scale control in sea water scale formation in sea water evaporators have of necessity varied evaporators should include consideration of (a)the limiting soluwidely from the methods successfully employed in fresh water bilities a t equilibrium for each of the three major scale composystems. Water softening, for example, is impractical because nents in sea water brines of varying concentration and a t several the total hardness of sea water is some sixty times greater than boiling temperatures, and ( b ) the rates of attaining these equithat of average fresh water, and further, in sea water only a libria. Because few of these data could be found in published relatively small portion of the total hardness constituents are literature, studies were initiated in this direction. rendered insoluble by moderate concentration through evaporaThis paper describes a method of scale control for sea water tion. Scale control in sea water evaporators, therefore, is effected
January 1950
I
INDUSTRIAL AND ENGINEERING CHEMISTRY
127
Compression evaporators employed in the distillation of sea water develop objectionable quantities of scale within 300 to 500 hours of operation. The scale is predominantly magnesium hydroxide. In the present study, solubility data were obtained for the three scale-forming salts in concentrated sea water brines a t 212' to 213' F. The data show that with respect to calcium carbonate and magnesium hydroxide the brines remain supersaturated over periods of time considerably longer than the normal detention periods in conventional evaporators. Based on these findings a process of scale control not requiring chemical treatment of the feedwater has been developed. In this process evaporator brine is continually circulated through a layer of stabilizing contact material
placed in a separate compartment thence back to the evaporator. The efficiency of the process has been demonstrated both in the laboratory and in the field. The field test which employed a standard thermocompression unit of 60-gallon per hour capacity was completed recently a t Fort Story, Va., under the direction of the Water Supply Branch of the Research and Development Laboratories, United States Army. In this test run of 1000 hours' duration, scale formation was reduced to a fraction of normal expectancy with corresponding salutary effects on the water to fuel ratio and capacity of the unit. The results are believed to be capable of marked improvement by modifications of the equipment in the light of accumulating data.
evaporators which requires no chemical treatment of the feedwater; it also describes briefly the laboratory findings which led to its development.
sulfate hemihydrate and the 120-hour sample to that of nearly pure anhydrite. Partridge and White ( 2 ) studied the solubilities of hemihydrate and anhydrite in the temperature range 100" to 200" C. and showed that hemihydrate exists as an unstable solid phase over the entire temperature range 0' to 200" C.; in the range 90" to 130" C. it is apparently stable for periods up to 48 hours. The authors' findings are generally in accordance with the conclusions of Partridge and White.
P
EQUILIBRATION DATA
~
El
LI
In Figure 1 are shown the results of equilibration data for calcium sulfate in concentrated sea water brines at 100' C. Samples of sea water, freed of carbonates, were prepared with adjusted ratios of calcium to sulfate and were reflux-boiled for varying intervals of time. The procedure involved analysis of the equilibrated samples and identification of the solid phase. Three series of tests were made. From the results of the first series, which are represented in the figure by circles, it became apparent that more than one type of sulfate was involved, depending upon the equilibration period-which in this series varied between 15 and 30 hours. Accordingly, two additional series were run, in one of which the equilibration period was terminated a t 17 hours and the other a t 120 hours. The results of the two latter series are represented by triangles and squares, respectively. Examination of the figure shows that all of the results of the 17-hour tests fall on the lower curve, while those of the 120-hour tests fall on the upper curve. Points representing the first series of tests are seen to be scattered about the area bounded by the two curves. The curves of Figure 1 indicate that under the conditions of the tests calcium sulfate precipitates in two distinct forms, depending upon the equilibration period. The two forms differ widely in their appearance and solubility. Examination of the precipitates under the microscope indicate that the first to form is calcium sulfate hemihydrate (CaS04.0.5H~O) and the second is the less soluble anhydrite (CaS04). In order further to verify this assumption, chemical analyses were made of composite samples of the precipitates from each series. The results showed the precipitates to be nearly pure calcium sulfate. However, the 17-hour composite sample lost 5.90% of its weight on ignition, while the 120-hour composite sample lost only 1.53%. The 17hour sample conforms closely to the expected analysis for calcium
I 'a
2.0
A
Second Series, 17 HOW Test$
m
Third
S a r i e ~ ,120 Hour T e s t s
I
These data enable calculation of the maximum permissible concentration of sea water without precipitation of calcium sulfate in either form. For this purpose, there has been plotted in Figure 1 the negative logarithm of the product of the molal concentrations of the calcium and sulfate ions for increasing concentrations of normal sea water. The point where the ionic product equals the solubility product is given by the intersection of the corresponding lines and represents the lowest concentrations where precipitation of the salt in question can be expected to occur. In the case of anhydrite, precipitation occurs when normal sea water is concentrated to two thirds of its original volume, but only after a boiling period of many hours. The more soluble hemihydrate salt, on the other hand, does not TABLEI. CHEMICALCOMPOSITION OF TYPICAL EVAPORATORprecipitate until a concentration to about one third-Le., conSCALES centration factor 3.1-has been exceeded. This corresponds to a Type of Evaporator required purge rate of 32% of the feed water, which appears to ThermoLow pressure compreasiona, double effect d, agree reasonably well with plant experience. Constituent - % % Precipitation of anhydrite should not be a problem in the presSilica 0.5 0.1 ent-day multiple-effect evaporators or thermocompression stills Iron and aluminum oxides 0.6 3.0 Calcium carbonate 1.6 87.7 operating on an overflow rate of 50% of the feed water, even Magnesium hydroxide 95.6 3.8 Calcium sulfate 1.1 5.6 though concentration effects are favorable. This is because the a Analysis by Engineer Research and Development Laboratories, Fort liquid detention time in these evaporators is much less than is Story Va. required for the induction of the reaction. Hemihydrate preb Analysis by U.8. Naval Engineering Experiment Station, Annapolis, Md. cipitation, however, does not require a long induction period and
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Vol. 42, No. 1
carried out in tightly stoppered 500-ml. bottles, submerged in a covered water bath maintained at 100" C. Provision for thc withdrawal of samples for analysis was the same as that described for the magnesium hydroxide tests. The samples were incubated for periods of 5 days or longer. Provision for continuous agitation was omitted, but the bottles were removed from the bath and shaken a t intervals of 8 to 16 hours. There were three series of samples. In the first series the alkalinity of the water was neutralized with hydrochloric acid, and brines of varying concentration were prepared by evaporation. To each sample of this series was added an excess of precipitated calcium carbonate powder. In the second and third series neutralized concentrated brines, prepared as in series I, were dosed with variable amounts of sodium bicarbonate and sodium carbonate, respectively. The filtered samples were tested for p H a t both 100" and 25" C. Calcium was determined by double or triple precipitation with ammonium oxalate, and alkalinity was determined by titration to a p H of 4.35, a glass electrode being used to indicate the end point. The data were plotted as in the case of magnesium hydroxide to show the relationship between brine concentration and the equilibrium constant for the reaction: Figure 2. Apparatus for Determining the Solubility of Magnesium Hydroxide in Boiling Sea Water Leeda and Yorthrup conductivity meter Powerstat voltage regulator C. Ikckman pH meter (Model H ) I). Glascol heating mantle E . %Liter boiling flabk F. 100' C. pH elevtrodes G . Conductivity cell H . Microburet J . Graham rondenser 4. B.
CaC03 (solid)
+ H + = Ca++ + HCOJ-
The equilibrium constant for this reaction involves the solubility product constant for calcium carbonate and the second dissociation constant for carbonic acid-Le., pK = pK, - pK2. It has been shown (3)that, in sea waters a t p H values lower than about 8.0, the bicarbonate concentration can be taken as equal to the total alkalinity, and therefore it is permissible in the equilibrium equation to substitute equivalents of total alkalinity for moles of bicarbonate ion and the equation can be written in the log form as :
pK = pCa
+ pAlk - p H
can occur when sea water evaporators are opeiated at overflow rates below 32%. Chemical analyses of scale samples, obtained from thermocompression stills operated in Arabia and using ground water high in calcium and sulfate ions, indicate that these scales were composed largely of calcium sulfate hemihydrate.
The pK values were found to vary only slightly with concentration of the brine, and the average value of -2.6 is taken as applicable for all concentrations u p to 3.0 when p H is measured a t 25" C.
~ I A G N E S I UHYDROXIDE M EQUILIBRIUhf. Determinations of the solubility of magnesium hydroxide in various concentrations of sea water were carried out in a 2-liter boiling flask with multiple openings for the insertion of a vent condenser, pH electrodes, a n electrical conductivity cell, a siphon with filter and cooling coil for withdrawal of samples, and a microburet for the periodic addition of sodium hydroxide. An illustration of the apparatus is shown in Figure 2. The procedure with each carbonate-free sample of brine consisted of periodically raising the hydroxyl ion concentration by addition of sodium hydroxide. The intervals between successive additions of hydroxide and sample collections varied between 1 and 20 hours, and in most cases the intervals were prolonged until repeated measurements of p H and conductivity indicated stable conditions. The withdrawal of filtered samples for magnesium and room temperature p H determinations was accomplished by siphonage through a steam-jacketed filter tube, operated under slight pressure to prevent flashing, and thence through a Graham condenser. Difficulties were experienced with clogging of the filter as well as in the analytical determination of small amounts of magnesium in the presence of large amounts of calcium. The data were plotted to show the relationship between brine concentrations and the equilibrium constant for the equation :
The experimental data reported can be most usefully employed if presented in the form of a chart or diagram showing
JIg(OH)t(solid)
+ 2"-
=
Mg++
STABILITY DIAGRAM
+ 2Hn0
The equilibrium constant for this equation involves the solubility product constant for magnesium hydroxide and the ionization constant for water-i.e., pK = 2pK, - pK, = 2pH pMg. Although several of the individual observations showed considerable deviation from the mean, there appeared to be only a slight variation due to the concentration of the brine. When utilizing p H measurements made a t room temperature, the mean pK value for all concentrations of brines was found to equal 14.6. Determination of magnesium hydroxide solubility has been a source of trouble for other investigators, as evidenced by the wide range of values given in the published literature. It seems probable that the discrepancies can be largely attributed to failure to attain consistently true equilibrium conditions. CALCIUMCARBONATE EQUILIBRIUM. Saturation of sea water brines of varying concentrations with calcium carbonate was
o H Meosured ot
25' C
Figure 3. Stability Diagram for Calcium Sulfate, Calcium Carbonate, and Magnesium Hydroxide in Concentrated Sea Water Brines at 100' C.
January 1950
b
I N D U S T R I A L A N D E N G I N E E R I N G CHEMISTRY
the limiting conditions of solubility of the three salts in question under specified conditions of brine concentration and temperature. Diagrams of this type have been used by Langelier (1) to indicate the pipe-scaling properties of fresh waters in terms of readily obtained analytical values. I n practice they are used to indicate departures from true equilibrium conditions, and are called stability diagrams. Such a diagram is shown in Figure 3. The scale values on the diagram are values obtained by analysis of the brine. If an analysis of the feedwater and the degree of concentration are known, the concentrations of calcium, magnesium, and sulfate in the brine can be calculated with sufficient accuracy on the assumption that the amounts precipitated will be negligibly small. The use of the diagram is not limited to concentrated sea water but, providing the total alkalinity is due only to bicarbonate, it is applicable without correction to similar saline waters, such as harbor waters, etc. The diagram is designed and can be used for waters of total mineral contents of between roughly 25,000 and 100,000mg. per liter. Equilibrium with respect to either form of calcium sulfate is a function only of calcium and sulfate ion concentrations. Accordingly, a single line for each form is shown. Plotted values of calcium and sulfate falling below and to the left of the calcium sulfate lines indicate undersaturation, those above and to the right indicate supersaturation with respect to these salts. Equilibrium with respect to magnesium hydroxide for the temperature and range of the salinities indicated varies only with p H and magnesium content. Plotted values of pH and magnesium to the left of the magnesium hydroxide equilibrium line indicate undersaturation, those to the right supersaturation. Equilibrium with respect to calcium carbonate is not only a function of pH and calcium concentration but also is dependent upon a third factor-namely, total alkalinity. Accordingly, separate lines must be used for different alkalinities; with this difference, the condition of the sample with respect to saturation with calcium carbonate is determined as in the case of magnesium hydroxide. A convenient manner of using the diagram for estimating the condition of the sample with respect to magnesium hydroxide and calcium carbonate saturation is to determine whether the saturation pH, as shown on the diagram, is lower or higher than the actual p H of the sample. The difference (pH - pH,) is referred to as the saturation index. If with respect to the particular constituent the index is negative, the sample is undersaturated; if positive, it is supersaturated. An example illustrating the use of the diagram is shown in Figure 4. In the normal operation of a thermocompression evaporator employed ia the distillation of sea water, the overflow rate is 50% of the raw-water feed rate. For these conditions, plant and laboratory data indicate average analyses of the brine as follows: calcium 2000; magnesium 10,000; and alkalinity, 100 ; each expressed as milligrams per liter in terms of equivalent calcium carbonate. The sulfate concentration averages 5700 mg. per liter; the pH averages 8.6. In Figure 4 these data have been plotted on the diagram. With respect to calcium sulfate, it will be noted that the brine is supersaturated with anhydrite but not with hemihydrate. With respect t o magnesium hydroxide, the plotted point is well to the right of the magnesium hydroxide line, which indicates supersaturation with this salt. "The saturation index is 4-0.8. With respect to calcium carbonate, two points have been plotted, To the left, calcium is ]plotted against total alkalinity and indicates the saturation 1pH of 7.0; to the right, calcium is plotted against the actual pH .of the brine, and the indication is that the saturation alkalinity i s approximately 5 mg. per liter. It will be seen then that the d c i u m carbonate saturation index is 4-1.6,which indicates a ,degree of supersaturation greater than for magnesium hydroxide. Because the scale samples removed from thermocompression evaporators normally contain magnesium hydroxide to the
129
pH Measured at 2 5 ' G
Figure 4. Stability Diagram with Plotted Analysis of Evaporator Brine under Normal Operating Conditions
almost complete exclusion of calcium carbonate, it is suspected that calcium carbonate requires a greater period of time €or precipitation from supersaturated solution than does magnesium hydroxide. INTERPRETATION OF EXPERIMENTAL DATA
In sea water evaporators the composition of the precipitate or scale cannot be predicted from equilibrium data alone but is dependent upon the characteristics and relative rates of precipitation of the salts from supersaturated solution. This conclusion is responsible for the idea that scale formation could be prevented by physical rather than chemical methods. The physical methods involve control of the detention period of the brine within the evaporator and can be secured by various means. The methods which have been considered involve ( a ) evaporator design to secure minimum liquid volume, ( b ) operation with high overflow rate, and (c) circulation of the evaporator brine through a contact unit to reduce continuously supersaturation. Only the latter is discussed in the present paper. CONTACT- STAB1LI ZATION METHOD
The contact-stabilization method of evaporator-scale control was developed for use with thermocompression evaporators operating a t 212" to 213' F. but should also be applicable to evaporators operating at lower temperatures. I n essence, the method is designed not to prevent scaling but rather to cause a major portion of the potential scale to deposit within a contact unit from which it can be easily removed without interrupting the operation of the evaporator. The method is based upon the principle that the rate and location of crystallization of salts from supersaturated solution can be controlled by contact with suitable mediums. More specifically, in the present method, brine from the evaporator, supersaturated with calcium carbonate and magnesium hydroxide, is caused to flow continuously through a compartment holding a layer of limestone granules or other suitable contact material, and thence back to the evaporator.
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Figure 5. Laboratory Sea Water Distillation Apparatus with Contact Stabilization L e f t . Exterually heated evaporator Right. Internally heated evaporator
Vol. 42, No. 1
The efficiency of the contact unit in removing saturation excess is dependent upon several factors including the degree and kind of saturation excess, the contact time, and the size, composition, and quantity of the contact material. In a given installation, the circulation rate fixes th% detention time of the volatile solids (bicarbonate plus carbonate) in the evaporator ahich in turn governs the pH increase. I n the authors’ studies, circulation rates of from six to ten times the raw water feed rate have given satisfactory results. The efficiency of the stabilization method has been demonstrated both in the laboratory and in the field. A photograph of the apparatus is shown in Figure 5 . Test runs using a 60gallon per hour thermocompression still with an improvised contact unit, a flow diagram of which is shown in Figure 6, were carried out a t the distillation test station a t Fort Story, Va. In the initial run of 1000 hours’ duration, scale formation wa5 reduced far below normal expectancy, with corresponding salutary effects upon the water-to-fuel ratio as !\ell as upon the capacity of the unit. Determination of copper and zinc content of the brine showed that corrosion in the evaporator was insignificant under this method of control. Subsequent runs of shorter duration have been made to evaluate the variable factors affecting the efficiency of the process and the design of the equipment. These tests have repeatedly shown the feasibility of scale reduction to less than 10% of normal expectancy. Correlation of laboratory test data with those of full scale operation is being continued, with a view to determirnrig optimum design factors. Analytical data from typical labolatory and full scale test runs, with and without contact stabilizhhn, are shown in Table 11.
Deposition of a portion of the saturation excess of each salt occurs on the contact surfaces. The efficiency of this method is subject to the maintenance of a correct balance between the rate at which supersaturation is built up in the evaporator and that at which deposition occurs in the contact unit. The formation of carbonate and hydroxyl ions which occurs in the evaporator is caused by the decomposition of bicarbonate ions through the evolution of free carbon dioxide into the stream. Feed Pump Sea Water
CONTACT STABILIZATION TEST TABLE 11. RESULTSO F TYPICAL RUNS (Laboratory and full scale) Laboratory Still 60-G./Hr. Still Without Tith Without With contact contact contact contaot unit unit unit unit 0.0021 0.0021 1.95 1.78 Feed water rate, gal./min Distillation rate, gal./min. 0.0010 0.0011 1.12 1.20 Concentration factor 2.0 2.0 2.36 3.11 Circulation ratio (circulation rate to feed rate) .. 10.1 ,* 11.0 Contact material .. Limestone Sand AYaAytical data
..
P I A
7.80 7.9 7.9 Feed water 7.85 Blowdown (contact unit in8.78 8.28 8.6 8.1 fluent) .. 8.0 Contact unit effluent 8.11 Alkalinity, m d l . as CaCOa 92 114 116 89 Feed waterBlowdown (contact unit in122 60 75 60 fluent) 52 54 Contact unit effluent Scale balance, lb./1000 gal. distillate= Potential scale in feed water 1,90 1.94 1.33 1.10 Scale deposited in evaporatof 0.87 0.093 0.84 0.112 Reduction due to contact unit . 0.78 .. 0.73 Percentage reduction due to contact unit 89.6 87.0 a Includes only calcium carbonate and magnesium hydroxide. Balance is based on t o t 3 alkalinity determinations expressed as equivalent calcium carbonate.
~.
.. . ..
..
..
I
Inlet I
1 Otsl llate
I , Heai Exchanaer
II I
rump
1
Brine Overflow
Figure 6. Flow Diagram of Thermocompression Distillation Process with Contact Unit ACKIVOWLEDGMENT
The work described in this paper has been done under a contract between the University of California and the Engineer Research and DeveIopment Laboratories of the Corps of Engineers, United States Army. Acknowledgment is made of the cooperation and material assistance of Charles R. Keatley, chief of the Water Supply Branch, Ernest H. Sieveka, chief of the Water Treatment and Laboratory Section, and other members of the staff. LITERATURE ClTED
(1) Langelier, J. Am. Wuter Works Assoc., 38,169 (1946). (2) Partridge, E.P., and White, A. H., J . Am. Chem. Soc., 51, 360 (1929). (3) Sverdrup, H. U., Johnson, M. W., and Fleming, R. H., “The Oceans,” p. 201, New York, Prentice-Hall, Inc., 1946.
RECEIVED May 23, 1949. Presented before the Division of Water, Sewage, CHEMICAL and Sanitation Chemistry at the 115th Meeting of the AXERICAN S O C I ~ T San Y , Francisco, Calif.
