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STATIC-DYNAMI SURFACE TENSIONS
DV./CM.
Vol. 23, KO.11
A n interesting point in chemical philosophy should also be raised-namely, does this theory of film formation preclude all others? May there not be some other property that predominates to such an extent in some few substances that it determines the film formation? Such a property would be an unusual toughness of film accompanying a high degree of positive adsorption. Saponin and perhaps soap may belong in such a class. Saponin, for example, has been studied in connection with other work in boiler-water foaming and i t seems a t times to exhibit anomalous behavior. The effectiveness of castor oil in destroying or inhibiting foam seems also to be out of proportion to the concentration used. Literature Cited
! Figure '-static and Dynamic SurfaceFigure 6-Static and Dynamic SurfaceCane-Sugar Differences Solutions and Foam-Zone Heights Tension Differences and Foam-Zone Heights of Tension of Sulfuric Acid Solutions
sulfate shows a much higher percentage of film formation than sodium chloride. If, however, the comparison be made on the basis of equal concentration of sodium, it will be seen that the behavior of the two salts is almost identical, but the same concentration of sodium in the sulfocyanate does not give the same results. I n this case, however, the surface adsorption has the opposite sign.
(1) Foulk, IND. ENG.CHEM.,21, 815 (1929). (2) Foulk, I b i d . , 16, 1121 (1924) (3) Foulk Trans. Am. Sac. Mech. Ens., 50, FSP-5046,219 (1928). (4) Hansley, Doctor's Dissertation, Dept. Chem., Ohio State Univ., 1928. ( 5 ) Harkins, J . A m . Chem. s8*228 (1916). (6) Iredale, Phi'. Mas.3 46, (1923)' (7) McBain and Davies, J . A m . Chem. Soc., 49, 2230 (1927) ( 8 ) McBain and DuBois, I b i d , , 61, 3434 (1929). (9) Schmidt, z.physik. Chem., 39, 1108 (1912). (10) Stocker, Ibid., 94, 149 (1920). (11) Talmud and Suchowolskaja, I b i d . , 164,277 (1931).
Removal of Traces of Iron from Aqueous Solutions of Sulfates' T. W. Richmond and F. K. Cameron UNIVERSITY OF NORTH CAROLINA, CHAPELHILL, N. C.
HERE are a number of possible raw materials which might become available for the commercial production of sulfates of aluminum, sodium, magnesium, etc., if there were also available a practical way of removing from the salts small quantities of iron which contaminate the final product. Various attempts to solve this problem may be cited profitably. Zabicki (9) reports that, after two crystallizations of ammonia alum, an 86 per cent yield was obtained, containing but 0.002 per cent ferric oxide. Cooling and stirring were so controlled as to form very small crystals only. He describes an electrolytic method where a high overvoltage with hydrogen on a mercury cathode causes immediate solution of iron in the mercury, until the aqueous solution contains but 0.003 gram ferric oxide t o 100 grams aluminum oxide. Hultman and Linblad ( 4 ) propose adding to a neutral solution of aluminum sulfate, the hydroxide, carbonate, or sulfide of an alkali metal or ammonia, in quantity just sufficient to precipitate a semi-insoluble (?) basic aluminum sulfate. The liquid residue, heated under pressure, yields a precipitate, which has a composition of approximately the formula A120a.1.5S03, and which is very low in iron. Exactly opposite in principle is a proposal ( 2 ) to dissolve the raw material containing aluminum and iron in a large excess
T
1 Received
June 6, 1931.
of hot concentrated sulfuric acid. On cooling, an aluminum sulfate is obtained in this way, containing about 0.03 per cent iron. A complete removal of the iron is obtained by redissolving in hot sulfuric acid, 35-45' Baum6, and cooling, when pure aluminum sulfate crystallizes from the solution. Vittorf (8), investigating the possibility of commercially separating ferric sulfate and aluminum sulfate by dissolving in 90 per cent alcohol, reports that, on a first crystallization, the product contains about 0.1 per cent iron; but, on three recrystallizations, a 97 per cent recovery of aluminum sulfate is effected, containing 0.005 per cent iron. Adsorption methods have been tried. Thus, Fodor and Rosenberg (1) report that kaolin adsorbs ferric hydroxide from a solution of ferric chloride, but not completely, since the mineral does not neutralize the acid. Talc, however, is reported to neutralize the acid; hence the absorption of iron from the solution is complete. The authors have tried, but failed, to confirm these latter observations. Another proposal ( 7 ) is to treat solutions of aluminum sulfate containing iron, preferably when heated, with kaolin, clay, or bauxite. Part of the iron is removed by the mineral. More is removed by a little precipitated aluminum hydroxide, and the last traces by adding also some ferrocyanide, the blue prey cipitate being adsorbed on the aluminum oxide. Besides adsorption, base exchange suggests itself as a pos-
November, 1931
INDUSTRIAL AND ENGINEERING CHEMISTRY
sible procedure. Magistad (6) reports that the iron in ferrous sulfate exchanges with calcium in aqueous contact with soil zeolites, while ferric iron exchanges with aluminum. in studying the action of permutite in softening Kolb (j), waters, found that the product enters into reversible reactions with soluble salts except in the case of manganese, which does not combine with the zeolitic residue but is precipitated as manganese dioxide. This observation is of importance, since manganese, as well as iron, is often a Contamination in the manufacture of sulfates. It was deemed advisable to investigate further the possibilities of adsorption and base exchange. Adsorption and Base Exchange
A number of solutions of potassium alum were made of from one per cent concentration to saturation. From these, series were prepared containing 0.05 per cent, or less, iron as ferric sulfate. Fifty-cubic centimeter samples of these solutions in standard 8-ounce nursing bottles were vigorously agitated for 12 hours in a mechanical shaker after introduction of a weighed amount of an absorbent. After the shaking, the contents of the bottles were allowed to settle for about 2 hours, and a portion of the supernatant clear liquid withdrawn for iron determination. This was done by measuring the color developed when potassium or ammonium sulfocyanate was added, following, with slight modifications, the procedure of Griffin (3). A number of North Carolina clays and various pulverized minerals showed only negligible absorption of iron or removal of it from the liquid phase. Charcoals, however, were more or less effective. A saturated solution of alum with animal charcoal in the proportion of a 50-cc. solution to 1.75 grams of char showed a reduction of from 0.001548 to 0.0002 gram of iron; i. e., 86.5 per cent of the iron was removed from solution. This same charcoal, after steaming, was somewhat more effective. Nuchar was less than half as effective, while Darco and G. A. D., well-known commercial chars, were relatively ineffective. Adsorption of iron was the more efficient, the less concentrated the solution of alum. Considerable quantities of alumina were always adsorbed with the iron. Addition of excess sulfuric acid, sufficient to keep the alumina in the liquid phase, prevented also the removal of iron. Consequently, the procedure holds little promise of commercial application. Since the data in themselves are of no practical importance, they are not given except for the case of the saturated solution of alum treated with the steamed animal charcoal. These data are assembled in Table I. In the first column are shown the grams of char used to treat 50 cc. of the solution. In the second column are shown the grams iron present after treatment. In the third column are shown the grams of iron in the solution when sulfuric acid was added after treatment with the char. In all cases where chars were employed, a chart of the absorption data is a parabolic curve, indicating that the results were due to adsorption. This explanation was confirmed by testing the data with the Freundlich equation in the form: XLM = KCn
where X M
C K and n
= = = =
amount of iron removed weight of adsorbent concentration of iron constants
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product was described by the manufacturers as being of glauconite type, and a partial analysis gave 46.32 per cent SiO,, 19.15 per cent Fez03, and 9.22 per cent A1203, in harmony with the ideal formula KFeSizOO of the mineral glauconite. A typical set of results is shown by the data in Table I1 for three solutions, each containing 1 per cent alum, and 0.0004, 0.00425, and 0.0425 per cent iron, respectively, before treatment with the permutite. The coldmn headed X / M calculated from the data in the third column shows a steady decrease, indicating that the removal of iron from the solution is not an adsorption effect. That this ratio is not a constant but steadily diminished with increase in the amount of permutite employed, was shown to be the case in a number of experiments. It was also found that, if the iron in the solution be sufficiently low initially, then iron would dissolve from the permutite. There is, therefore, a limit to which the iron can be reduced by permutite, controlled by the relative concentrations of the several solutes. With 5 per cent glum solutions, the content of iron was reduced to 0.008 gram per 100 grams of alum. With 10 per cent alum solutions, the limit attainable by treatment with this particular glauconite appears to be about 0.004 per cent iron. Table I-Concentration of Iron in Alum Solutions with Steamed Animal Charcoal (50 cc. of saturated alum solution) CHAR -IRONNeutral Acid Grams Gram Gram 0.0017 0.0017 0.0 0.25 0.0016 0.00035 0.50 0.0016 0.00018 0.0016 0.00012 0.75 0.00008 0.0016 1.00 0.00005 1.25 0.0015 0.00004 1.50 0.0015 0.00003 0.0013 1.75 0,0012 0.00002 2.00 Table 11-Concentration
X/M
0,0004
0.0002 0.0002 0.0002 0.0002 0.0001 0.0002 0.0003
of Iron in Solutlons after Contact with Permutite
(SO cc. of 1 per cent alum solution)
PERMUTITE Grams 0.0 0.2 0.5 1.0 2.0
after Contact
Gram 0.0004 0.00045 0.0003 0.0002 0.00008
IRON Gram 0.00425 0.00325 0.00225 0.00095 0.0002
.
X/M
Gram 0.045 0.037 0.032 0.021 0.013
0.005 0.004 0.003 0.002
Solutions of magnesium sulfate and sodium sulfate were investigated in the same manner as the alum solutions. As typical of the results generally found, Table I11 shows the data obtained by treating 5 per cent solutions of sodium and magnesium sulfates] respectively, with the permutite. They are of the same character as the results obtained by treating alum solutions. Noteworthy was the failure to detect dissolved alumina in the solutions after treatment with permutite, indicating that the base exchange is between the iron in the liquid phase and the alkali metal (or metals) in the solid. The amounts involved are so small, relatively, to the total bases in liquid and solid that analysis before and after treatment is too uncertain to show whether the iron exchanges with alumina or with alkali metals. Tahle 111-Concentration of Iron in Solutions of Sodium Sulfate and Magnesium Sulfate after Contact with Permutite (50 cc. of 5 per cent solutions of sodium and magnesium sulfate) PERMUTITZ
Grams
-
Gram
IRON Gram
-
Gram
SODIUM S U L F A T E
C being a constant for this case, the ratio X / M should be constant as shown by the figures in the fourth column of Table I. A number of alum solutions, containing from 1 per cent salt to saturation and with various amounts of ferric iron present, were treated with a pulverized permutite. This
0.0 0.5 1.0 2.0
0.0005 0.00007 0.00007 0.00007
0.005 0,0036 0.0024 0.0006
0,0250 0,0175 0.0150 0.0130
MAGNESIUM S U L F A T E
0.0 0.5 1.0 2.0
0.0005 0.0004
0.0003 0.00006
0.005 0.0032 0.0021 0.0007
0.026 0.0175 0.0158 0.0135
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Summary
Clays, kaolin, and pulverized talc have been found inefficient in absorbing iron from aqueous solutions of sulfates. Various forms of charcoal show wide variations in effectiveness as adsorbers of iron from sulfate solutions. Animal charcoals are much more effective than vegetable chars. With alum solutions, alumina is also adsorbed in amounts sufficient to prove impracticable the using of chars. A permutite of the glauconite type is reasonably effective in removing iron from sulfate solutions. From the nature of the reactions involved, the removal is never complete. Working with diluted solutions, the remaining concentration of iron may be brought to a very low figure. Too low concentrations, however, may effect removal of iron from the per-
Vol. 23, No. 11
mutite with consequent contamination of the solution by iron. Literature Cited (1) Fodor and Rosenberg, Kolloid-Z., 46, 91 (1928). (2) German Patent 232,563 (1908), Chem. Fabr.-Griesheim-Elektron, Frankfurt; C. A . . 5, 2709 (1911). (3) Griffin, "Technical Methods of Analysis," McGraw-Hill, 1927. (4) Hultman and Linblad, Swedish Patent 53,134 (1922); C. A . , 17, 3407 (1923). (5) Kolb, Chem.-Zfg., 35, 1393 (1911). (6) Magistad, Arizona .4gr. Expt. Sta., Tech. Bull. 18, 445 (1928); C. A . , 22, 3595 (1928). (7) Norske Aktieselskab for Electrokemisk Ind. Norsk Industri-Hypotekbank. British Patent 123,720 (1918); C. A . 13, 1624 (1919). (8) Vittorf, Trans. Inst. Econ. Mineral. Pelrog. (Moscow), 1934, No. 8, 1-16; C. A , , 20, 1497 (1926). (9) Zabicki, Prsernysl Chem., 12, 77 (1928); C. A , , 22, 2641 (1928).
Thermodynamic Properties of Dichlorodifluoromethane, a New Refrigerant' IV-Specific Heat of Liquid and Vapor and Latent Heat of Vaporization Ralph M. Buffington and Joseph Fleischer FRIGIDAIRE CORPORATION, DAYTON,OHIO
T
HIS paper covers the
experimental determination of some thermal properties of dichlorodifluoromethane and completes the report of the experimental data necessary for the construction of tables of thermodynamic properties of this substance. Specific Heat of Vapor
The specific heat of dichlorodifluoromethane vapor at atmospheric pressure was measured in a flow calorimeter at 0', 25.S0, and 49.9"C., the experimental data fitting the equation, C, (molal) = 17.0 0.0279t (tOC.) The heat capacity of the liquid was measured as 30.3 calories per mole per C. at 17" C., using the method of mixtures, and as 25.4 at -43' C., using an electrical heating method. The latent heat of vaporization was determined by an electrical heating method as 4880 calories per mole at the boiling point, -29.8' C.; 4100 at 23" C.; and 3960 at 28" C. The ratio C,/C, for the vapor at 25' C. and atmospheric pressure was determined by the Kundt method as 1.139.
+
O
The specific heat of the vapor a t atmospheric pressure was measured a t three temDeratures in a flow calorimeier. The apparatus, which is illustrated in Figure 1, consisted of a 500-cc. Dewar flask fitted with a tight rubber stopper carrying a vapor inlet and a vapor outlet tube. The outlet tube was thermally insulated from the remainder of the apparatus by means of a silvered vacuum jacket. A 115-ohm nichrome heating coil was inserted in the lower end of the outlet tube; directly above it were placed several disks of copper gauze for the purpose of insuring temperature equilibrium in the vapor. As a further precaution against heat loss, the lower end of the outlet tube was inserted in a glass cup, which caused the cold gas to flow around the outside of the heater tube. Copper-constantan thermocouple junctions were placed in the inlet and outlet tubes, each junction being soldered to a disk of copper gauze. Both junctions and also the heating coil were enclosed in radiation shields of aluminum foil. The calorimeter was totally immersed in a thermostat, the temperature of which was controlled to *0.05" C. by a thermoregulator at the two higher temperatures; closer temperature control was obtained a t 0" C. by thorough stirring of an ice-and-water mixture. The vapor was passed through 12 meters of flattened copper coils immersed in the thermostat 1
Received July 13, 1931.
before entering the calorimeter. A steady flow of vapor was maintained by boiling liquid contained in a 1-liter Dewar flask with a constant heating current; different rates of vapor flow were obtained by varying the current. Measurements of the rate were obtained by noting the weight of the container at measured time intervals. A capillary flowmeter, attached to the outlet tube of the calorimeter, served to give a visual indication of the constancy of flow. The calorimeter heating current was obtained from lead storage cells; various heating rates were obtained by varying the number of cells used. The input of electrical energy was determined by measuring the voltage across the heater with a voltmeter and the current through the heater by the potential drop, measured with a potentiometer, across a standard 1ohm resistance in series with the heater. The temperature rise produced in the vapor was measured by means of the differential couple, the e. m. f. developed being measured with a Leeds and Northrup type K potentiometer. The thermocouple wire had been carefully calibrated, so that the temperature coefficient of its e. m. f. was known over a wide temperature range. The thermocouple junction in the inlet tube also served as the hot junction of a second couple, the other junction of which was kept a t 0" C., which was used to measure the temperature of the inlet vapor. The experimental data given in Table I were obtained by averaging the measurements taken over periods of at least l/2 hour, after conditions had become steady. It was found that, a t rates of vapor flow below 400 grams per hour, the measured specific heat became larger the slower the rate of flow, while
