Effect of Temperature on pH of Alkaline Waters WATER§ CONTAINING CARBONATE, BICARBONATE, AND HYDROXIDE ALKALINITY JEROME GREEN National Aluminate Corporation, Chicago, I l l . Beckman glass electrode used in conjunction with a Leeds & Northrup Standard 1199-19 saturated calomel electrode. Where necessary, corrections were applied for the sodium ion error of the glass electrode through the use of information furnished by the manufacturer. Such corrections did not exceed 0.07 p H unit. A Beckman Model G.pH mctcr was used for the electrometric measuremcnts. The pH measuring system was Calibrated with buffer solutions repared from pH standard materials obtained from the National gureau of Standards. One solution was 0.025 equimolar in potassium dihydrogen phosphate and disodium hydrogen phosphate, the other 0.01 M in borax. The pH of these buffer solutions is certified up to 60" C. Extrapolation of the data ives a pH a t 90" C. of 6.90 for the phosphate buffer and 8.91 for the borax buffer. It is estimated (14) that these values are in error by not more than 0.05 p H unit. The disagreement between the known and observed pH differences of the buffer solutions was not more than a few hundreths of a pH unit a t any of the experimental temperatures. When a Beckman pH 10 buffer solution, certified to 80" C., wasused, a greater difference, as large as 0.07 p H unit, was observed a t the elevated temperatures between its pH and that of the borax buffer than was indicated by the accepted pH values. All p H measurements reported here are referred to the borax buffer as standard. Waters of different p H values and containing varying amounts of carbonate, bicarbonate, and hydroxide alkalinity were prepared by adding to distilled water the proper amounts of stock solutions of sodium carbonate and sodium bicarbonate prepared from C.P. chemicals. Total alkalinity was dctcrmined by titration with standardized 0.02 N sulfuric acid, using the bromocresol green-methyl red indicator developed by Cooper (2). I n each of two constant tem erature ( k 0 . 3 " C.) baths were immersed flasks containing bu8er solutions, storage vessels for the electrodes, and a flask for the water whose p H was t o be measured. For the high temperature measurements this latter flask was a 2-liter, 3-necked vessel provided with a sam ling line and a condenser whose outlet was loosely plugged witR cotton. The sampling line went to a glass coil situated in a cold water bath and thence to a lOO-ml., 3-necked flask immersed in the 25 C. bath.
The pH of waters containing carbonate, bicarbonate, and hydroxide alkalinity was measured with glass electrodes at 25", 60", and 90" C. as a €unction of total allcalinity. From these data graphical representations are made of the change of pH with temperature in waters of different total alkalinity. A comparison is made between the experimental results and published values calculated from theoretical considerations. Except at high allcalinity, where agreement with one set of calculations is good, the experimental data show a somewhat smaller effect of temperature on pH than that indicated by calculations. Possible reasons for the discrepancies are indicated, and the applicability of the data to practical problems is discussed briefly.
THE
application of equilibrium calculations to water treating problems concerned with the solubility of slightly soluble compounds involves the use of p H as one of the variables (4, 6, 7, 11, 12, IS). In order to apply such calculations properly to conditions a t elevated temperatures, it is necessary to know, among other factors, the p H existing a t the temperature in question. For waters containing carbonates and bicarbonates as the only salts of weak electrolytes, recourse has been had to theoretical calculations which permit the estimation of pH a t temperatures other than that a t which it is measured ( I , 6, 1 0 , l l ) . Two such sets of calculations which have been reported recently (6, 10) are not in complete agreement, and insufficient experimental data are available to serve as corroborating evidence of the results reported. The work reported here was undertaken to provide an experimental check on the validity of the calculations. The experimental data are intended to serve as a guide in the use of the published tables or curves in order to indicate the possible magnitudes of the errors involved or corrections to be applied. 4
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EXPERIMENTAL PROCEDURE \ %
APPARATUS AND MATERIALS
A Leeds & Northrup Standard 1199-22 glass electrode was used for pH measurements a t the higher temperatures. This electrode is stated to be serviceable in the range 30 to 90' C. Data furnished by the manufacturer (8),led to the inference that, for the experimental conditions under which it was used, its sodium ion error is negligible. The high temperature electrode was used in conjunction with a reference electrode composed of the element from a Leeds & Northrup Standard 1199-23 calomel electrode immersed in an approximately 3.3 N solution of potassium chloride. The liquid junction was made through an asbestos fiber sealed in a glass tube. The very low rate of leakage permitted the junction to be maintained for considerable periods without apreciable contamination by potassium chloride of the solution eing used. The calomel element was sufficiently distant from the heated bath so that it was only slightly above room temperature. The pH measurements a t 25" C. were made with a No. 1190
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In' the experimental procedure the amounts of sodium carbonate and sodium bicarbonate required to obtain the desired pH and alkalinity were added to about 2 liters of distilled water contained in the sample flask immersed in the high temperature bath. The water was agitated occasionally over a period of about 2 hours by inserting a high speed electric stirrer through a neck of the flask. The glass electrode and calomel electrode, which were calibrated just prior to use, were inserted in the smaller necks of the flask and the pH recorded. A sample of the water was then withdrawn through the cooling coil into the flask a t 25" C. by applying suction t o the latter vessel. This flask contained the electrodes used for p H measurements a t 25' C., which had been calibrated reviously, a thermometer, and a sample withdrawal line. Wgen thermal equilibrium was established the pH of the water a t 25" C. was measured. The water was then siphoned from the flask and its total alkalinity was determined. Additional water and alkalinity were introduced into the sample flask in the high temperature bath and the procedure was repeated. I n this manner the alkalinity was progressively in-
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4 00 600 so0 PPM T o m / b / - K A i i n i / r v Figure 1. Example of P r i m a r y Experimental D a t a
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Measurements made on sodium carbonate solution
creased from about 17 p.p.m. to approximately 85, 340, and 850 p.p.m. expressed as equivalent calcium carbonate. For each pH and alkalinity a t 25" C. two experiments were made on separate samples, one in which the water was cooled from 60' C., the other in which it was cooled from 90 O C. Efforts were made to minimize errors in p H measurements a t the higher temperat'ures by allowing suificient time for the attainment of equilibrium in the electrode system. Between successive pH measurements the glass electrode was soaked for a short time in hot 0.1 S hydrochloric acid. Throughout a given series of high temperature measurements the calibration of the electrode system remained within a few hundreths of a pH unit of its initial value. Between pH measurements the electrodes were stored in distilled water contained in vessels immersed in the baths. The buffer solutions mere exposed as little as possible t o the atmosphere and were renewed at, intervals. Precautions were taken to obtain a representative sample of the water for pH measurements at 25 C. by flushing the flask and sampling line with a portion of the water before collection of the fina,l sample. Measurements of p H made a t higher temperatures after different time intervals and on water samples which were allowed to stand overnight at 25 C. indicated that equilibrium n-as reached in the experiments.
EXPERfMENTA L L A N G E L I E R f6) POWiFLL E T AL
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100 PPM. TO TAL AL KALINf T Y I
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RESULTS
Data from a typical series of experiments are presented in Figure 1which shows the p H a t 25 ' and 90' C. of a sodium carbonate solution of increasing concentration. A similar pair of curves \vas obtained for experiments in which the pH of water with the same type of alkalinity was measured at 60" and 26" C. From these two sets of curves pH values a t constant alkalinity were obtained for 25 ', 60 ', and 90" C. In like manner data for other ITaters of different' pH were obtained. The pH values at 25" C. were not usually identical for the two series of experiments, especially a t The lower alkalinities. I n drawing the curves in Figures 2 and 3, which show the variation in pH with kmperature at constant alkalinity, the data for 60' or 90' C. were shifted slightly to cause the pH values a t 25" C. to coincide. It is believed that no appreciable error was introduced by this procedure. I t is difficult to estimate the absolute accuracy of the ekperimental data. The uncertainty in the pH of the borax buffer a t '90" C., possible variation in the liquid junction potenrial, and rhe dificultics associated with pH measurements on slight,ly buffered solutions (corresponding to vaters of lo^ alkalinity) conibine to indicate that under unfavorable conditiolis errors of 0.1 pH unit or more niay exist. For most, of the measurements at)25' and 60' C. on waters containing more than 100 p.p.ni. total alkalinity, it is probable that the uncertainty is less than this. The only experimental data of a similar nature appear to be those of Kuentzel, Henslep, and Bacon (3) who measured the pH a t 25 ', 40 and 60 ' C. of sodium carbonate and sodium bicarbonate solutions having alkalinities of about, 100 and 1000 p,p.m., expressed as equivalent calcium carbonate. Their results are in fair agreement with the present ones and Kith the calculated values at the higher alkalinity but diverge apprrciably a t the lower alkalinO,
ity,
TEMPERA TURE
Figure 2.
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Effect of Temperature on p€I
THEORETICAL CALCULATIONS
A method of calculating the effect of temperature on pH of waters containing carbonate, bicarbonate, and hydroxide alkalinity was developed by Amorosi and McDermet ( 1 ) following suggestions made by JlcKinney (9). These calculations require a knowledge of the pH a t one temperature and the total carbon dioxide content of the xater. Since total alkalinity is more commonly determined in mater analysis than is total carbon dioxide, Powell, Bacon, and Lill (11) used this variable in their calculations, which wcre stated to be based on the method of Amorosi and XcDermet,. Recent,ly, Powell, Bacon, and Knoedler (10) have revised the earlier calculations on the basis of more accurate values for the ionization constants of carbonic acid. In neither of the papers by Powell and eo-workers was an explicit st,atement made as to the manner in which the transformation of variablcs from carbon dioxide to alkalinity v a s made. A different method, partly graphical in nature, was used by Langelier (6) to calculate the effect of temperature on pH as a function of alkalinity. An accuracy of better than 0.1 pH unit was not sought. The discrepancies betiwen his calculations and
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the earlier ones of Powell et a(.were attributed by Langelier to the use of different values for the ionization constants of carbonic acid. Since the revised calculations of Powell et al. are based on the same values of these constants (up to 50" C.) used by Langelier, it is apparent that other factors must account for the disagreement which still exists between the two sets of calculations. In the absence of intormation as to the values of the constants used by Powell et al. for temperatures above 50" C., i t is not possible to state whether the discrepancies in the calculations above this temperature can still be attributed to this source. Both methods of calculation are based on the assumptions that no gain or loss of carbon dioxide occurs on change in temperature and that in the equilibrium expressions concentration may be equated to activity. It is believed that the experimental method used satisfied the first requirement and that any error introduced by nonconformity with the second stipulation is considerably less than the experimental uncertainty.
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500 PPM TOTAL A L K A L I N I T Y
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EXPERIMENTAL L A N G E L I E R (6) I
COMPARISON O F EXPERIMENTAL AND THEORETICAL RESULTS
A comparison of the theoretical calculations with the experimental data is made in Figures 2 and 3. Since the calculated curves have been arbitraily selected, failure of the curves in a given set to coincide does not necessarily indicate disagreement between them. Rather, it is the degree to which the curves in a set are parallel that is the criterion to be employed for this purpose. It should be realized that the theoretical calculations are based on extrapolated values for the ionization constants of carbonic acid above 50" C. The agreement between the results is closer at the higher alkalinities. This may be due in part to a greater experimental error at the lower alkalinities. In general a smaller effect of temperature on pH is shown by the experimental results than that indicated by the calculations. The experimental data are in better agreement with the calculations of Langelier than with those of Powell et al. Within the limits of experimental error and of the uncertainty assigned by Langelier to his calculations - the experimental data substantiate the validity of Langelier's values.
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APPLICATION
I n applying the calculated data to actual problems i t is important to realize the conditions which must be satisfied. The water must contain no salts of weak bases or of weak acids other than carbonic acid. This restriction eliminates the applicability of the data, for example, to waters high in silicates, phosphates, sulfides, ammonium salts, weak organic acids or the salts thereof. If the amounts of such compounds are small compared with that of the carbonates and bicarbonates it is probable that no great error will be introduced by their presence. No gain or loss of alkalinity should occur on a change in temperature. This means that precipitation of carbonates or hydroxides may not occur. Loss of water by evaporation and absorption or liberation of carbon dioxide are also precluded. The assumption that concentration may be equated to activity restricts the data to waters which do not contain high concentrations of dissolved solids. I n practice i t would be advisable to obtain the p H and total alkalinity of a sample of water cooled from the higher temperature and use these data to estimate the p H at the elevated temperature. I n applying p H data to calculations involving the solubility of calcium carbonate i t is important to bear in mind that the sohbility product of this compound has not been determined at elevated temperatures and hence for such temperatures the calculations are based on extrapolated values of this constant. I n order to utilize more fully data on the variation of p H with temperature and alkalinity, experimental work on the calcium carbonate equilibrium at higher temperatures is necessary.
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Figure 3.
60 80 TEMPERATURE, 'C.
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Effect of Temperature on pH LITERATURE CITED
(1) Amorosi, A. M.,and McDermet, J. R., Am. SOC. Testing Materials Proc., 39, 1204-18 (1939). (2) Cooper, S. S., IND. ENG.CHEW,AKAL.ED.,13,466-70 (1941). (3) Kuentzel, L. E., Hensley, J. W., and Bacon, L. R., IND.Eso. CHEW,35, 1286-9 (1943). (4) Langelier, W. F., J . Am. Water W o r k s dssoc., 30,1500-21 (1936). (5) Zbid., 38, 169-78 (1946). (6) I b i d . , 179-85 (1946). (7) Larson, T. E., and Buswell, A. M., Zbid., 34, 1667-78 (1942). (8) Leeds & Northrup Co., private communication. (9) McKinney, D. S.,Am. SOC.Testing Materials, Proc., 39, 11911203 (1939). (10) Powell, S. T., Bacon, H. E., and Knoedler, E. L.,IND.ENG. CHEM.,40, 453-7 (1948). (11) Powell, S. T., Bacon, H. E., and Lill, J. R., Ibid., 37, 842-6 (1945). (12) Ryznar, J. W., J . Am. Water Works Assoc., 36, 472-83 (1944). (13) Ryznar, J. W., Green, J., and Winterstein, M. G., IND. ERG. CHEW.,38, 1057-61 (1946). (14) Smith, E. R., Natl. Bur. Standards, private communication. RECEIVED September 16, 1948. Presented before the Division of Water, Sewage, and Sanitation Chemistry a t the 114th Meeting of the AXERICIS CHENICAL SOCIETY, St. Louis. Mo.
