0.2 scale division. As each division 0 is equivalent to approximately 0.005 C. (see below), this corresponds to a standard deviation of 1 X 10_3° C., which is the same as the value reported by McMullan and Corbett {10). If only the first six solutions of Table I are considered, the data up to 0.06 molal can be represented by the linear
equation m
=
1.046 X
”3 Ad
(3)
with
an estimated standard deviation of 0.4 scale division or 2 X 10~3° C. It is from the slope of Equation 3 (molality per scale division) and the cryoscopic constant reported by Kraus and coworkers {1, 9) for benzene (5.07° C. per molality) that the estimate of 0.005° C. per scale division was obtained for this work. In practice the cryoscopic constant for each new bottle of benzene was calculated from the freezing point depression (expressed in scale divisions) measured for an approximately 0.05 molal solution of triphenylmethane in
benzene. A 0.05 molal solution of triphenylmethane in benzene has a freezing point about 50 scale divisions below that of
the pure solvent (Equation 3). At this concentration the estimated standard deviation, or error, of 0.4 division corresponds to a relative error of 0.8%. Assuming a negligible error in the weights of solvent and solute used for a single determination, the estimated relative error for the determination of the cryoscopic constant from the ratio of molality to freezing point depression should also be 0.8%. This was confirmed by sets of replicate determina-
tions of the cryoscopic constant which were made over a period of 10 months The value on various lots of benzene. of 1.4% so obtained, and based upon 23 degrees of freedom, is not significantly different from the estimate of 0.8% based upon 5 degrees of freedom. The data presented here are for an ideal solute-solvent system and were obtained solely to illustrate the performance characteristics of the equipment. When cryoscopic measurements are to be made on nonideal systems, appropriate extrapolations of the data must be employed to correct for the variation with composition of the apparent molecular weight of the solute. ADVANTAGES OF AUTOMATIC CRYOSCOPY
The following advantages characterize the apparatus described in this report. Once the solute and solvent have been weighed into the cryoscopic cell and the cell has been assembled in the freezing bath, no further operator attention is
required. A permanent record of the cooling curve is automatically obtained. The extrapolations required for the determination of the freezing points of solutions are more easily made from a continuous curve than from one plotted manually from discrete temperature readings. ACKNOWLEDGMENT
The author gratefully acknowledges the assistance of Harley W. Middleton, Marie DeVito, and Joyce H. Northrop in the construction and testing of the equipment.
LITERATURE CITED
(1) Batson, F. M., Kraus, C. A., J. Am. Chem. Soc. 56, 2017 (1934). (2) Beck, A., J. Sci. Instr. 33, 16 (1956). (3) Daniels, F., Mathews, J. H., Williams, J. W., Bender, R., Alberty, R. S., “Experimental Physical
Chemistry,” 5th ed., p. 68, McGraw-Hill, New York, 1956. (4) Giguére, P. A., Secco, E. A., Can. J. Chem. 32, 550 (1954).
(5) Glasgow, A. R., Jr., Ross, G., J. Research Natl. Bur. Standards 57, 137(1956). (6) Glasgow, A. R., Jr., Streiff, A. J., Rossini, F. ~D.,Ibid., 35, 355(1945). (7) Herington, E. F. G., Handley, R., J. Chem. Soc. 1950, 199. (8) Herington, E. F. G., Handley, R., J. Sci. Instr. 25, 434 (1948). (9) Kraus, C. A., Vingee, R. A., J. Am. Chem. Soc. 56, 511 (1934). (10) McMullan, R. K., Corbett, J. D., J. Chem. Educ. 33, 313 (1956). (11) Mikhkelson, V. Ya., J. Anal. Chem. U.S.S.R., 9, 21 (1954). (12) Müller, R. H., Stolten, H. J., Anal.
Chem. 25, 1103 (1953). (13) Newman, M. S., Kuivila, H. G., Garrett, A. B., J. Am. Chem. Soc. 67, 704 (1945). (14) Richards, L. A., Campbell, R. B., Soil Sci. 65, 429 (1948). (15) Spooner, D. C., J. Sci. Instr. 29, 96 (1952). (16) Stull, D. R., Ind. Eng. Chem., Anal. Ed. 18, 234 (1946). (17) Stull, D. R., Rev. Sci. Instr. 16, 318 (1945). (18) Witschonke, C. R., Anal. Chem. 24, 350 (1952). (19) Zeffert, B. M., Hormats, S., Ibid., 21, 1420 (1949). (20) Zeffert, B. M., Witherspoon, R. R., Ibid., 28, 1701 (1956). (21) Zemany, P. D., Ibid., 24, 348 (1952).
Received
for review May 9, 1957. Accepted November 23, 1957. Division of Analytical Chemistry, 131st meeting, ACS, Miami, Fla., April 1957. _
Drying and Decomposition of Sodium Carbonate ARTHUR
E.
NEWKIRK and IFIGENIA ALIFERIS
General Electric Research Laboratory, Schenectady, N. Y.
Thermobalance curves are given for the drying and decomposition of sodium carbonate in the temperature range from 25° to 1040° C. using different crucible materials and atmospheres. The reaction of sodium carbonate and silica resulted in a weight loss at temperatures as low as 500° C. It is recommended that sodium carbonate for analytical use be dried in platinum to prevent errors due to its reaction with silica and silicates.
It
is more than a coincidence that the accepted maximum temperature for drying sodium carbonate for use as a
982
·
ANALYTICAL CHEMISTRY
primary standard, and the lower limit for the initiation of the reaction of sodium carbonate and silica are the same —i.e., 300° C. This fact is not generally recognized. Directions for drying sodium carbonate usually specify porcelain or platinum containers, but glass weighing bottles are also used. In this laboratory, satisfactory results have been obtained by drying in glass at 250° C. {1). The recommended temperature range from 250° to 300° C. is the result of many studies (6), but the upper temperature limit is surprisingly low in view of the several reliable reports of the thermal stability of
sodium carbonate at temperatures close to its melting point, 851 ° C. Richards and Hoover {8) found that sodium carbonate heated under carbon dioxide for a long time just below the fusion point lost on fusion only about 0.003% in weight. Duval {2) reported that sodium carbonate, when heated in a thermobalance, is stable in the range from 100° to 840° C. Motzfeldt (7), who made a careful study of this decomposition in a platinum cell, concluded that there is no chance for significant decomposition at any temperatures below 800° C. at least. He attributes previous discordant results to
’
reaction of the sodium carbonate with the moisture of the air (or the water content of the initial sodium bicarbonate) to form sodium hydroxide and carbon dioxide. A recent study by Easterbrook (3) with the aid of a thermobalance showed that sodium carbonate prepared by heating sodium reached constant sesquicarbonate weight at 210° C. in 2 hours; there was no further change in weight after 19 hours at 270° C. or an additional hour at 340° C. Easterbrook concluded that the slightly high titer of the product (maximum calculated purity, 100.023%) resulted from the formation of oxide or hydroxide produced by reaction with water during the decomposition of the sesquicarbonate. The authors studied the reaction between sodium carbonate and other substances at elevated temperatures, and heated sodium carbonate and sodium carbonate plus silica in a Chevenard thermobalance using a number of different crucible materials and atmospheres. The results and additional data from the literature suggest that another factor causing the discordant results may be reaction of sodium carbonate with glass or porcelain used as a container.
due to water in the sodium carbonate as pointed out by Duval (0). In runs 8 through 11 the sodium carbonate
dried by heating in weighing bottles of borosilicate glass for 16 hours at 350° C., and hence showed no initial water loss. The large losses in runs 1 through 4, starting in the neighborhood of 800° C., are chiefly attributed to the reaction was
Table I. Weight Losses in the Reaction of Sodium Carbonate with Porcelain and Silica Caled. Loss, Gram
Run No.
Obsd. Loss,
Gram 0.1927 0.2028 0.2083 0.2087
0.2076 0.2076 0.2076 0.2076
1
2 9 10
APPARATUS AND MATERIALS
This work was done with a penrecording Chevenard thermobalance, whose construction and performance have been described (9). Experiments in air were performed with the balance as received, but experiments in controlled atmospheres were performed using a quartz liner tube for the furnace fitted with a loose cap at the bottom and a tube for the introduction of gas at the top. The porcelain crucibles were Coors regular laboratory grade. The alumina crucible was Triangle RR grade supplied by Morgan, England. The gold crucible was spun from pure gold sheet, and the platinum crucibles were of regular laboratory grade. The sodium carbonate (General Chemical Co., anhydrous reagent grade) was stored over Drierite (W. A. Hammond Drierite Co., Yellow Springs, Ohio) except during weighing of the samples. Nitrogen from the laboratory supply was dried by passage through a trap packed with glass wool and cooled in liquid nitrogen. It was humidified by passage through water in a sinteredglass bubbler at 25° C. The carbon dioxide (Matheson Co., “bone dry” grade) was used as received. EXPERIMENTAL RESULTS
each case.
10
'---^
experimental thermobalance curves are reproduced in Figure 1. The explanation for the shape of these curves seems straightforward. The small loss below 100° C. is probably
DISCUSSION
:
j
;
L
|
'-i™'
Í
1040 '
C"
'
400
2Ó0
TEMP.
’
6Ó0
Í-
!
‘C,
TIME HR.
Figure 1. Thermobalance sodium carbonate"
runs
with
Atmos-
Curve
Crucible
9 10
Porcelain Porcelain Porcelain Alumina Platinum Platinum Platinum Platinum Platinum Platinum
11
Gold
1
2
3
4 5 6
7 8
phere Air Air Dry Air Air Dry
9 and 10.
N2C
N2e
Wet
N2c
C02c Dry N2C Dry N2c
Dry
“Heating rate 300° 6
The
of sodium carbonate with the porcelain alumina crucible. When the heating is carried out for a long period of time, the weight loss is approximately quantitative (Table I). When the sodium carbonate is heated in platinum or gold, the weight loss is much less rapid and is probably due to the decomposition of sodium carbonate to form sodium oxide and carbon dioxide as proposed by Motzfeldt (7) (runs 5, 6, 7, and 11). As would be expected, the presence of a positive stream of gas to flush out the carbon dioxide resulted in a faster rate of weight loss (run 6 vs. run 5). However, when water was present in the gas stream, the observed rate of weight loss was less (run 7 vs. run 6). A sample of the sodium carbonate dried at 350° C. showed no significant change in weight on further heating 12 hours at 600° C. and 4 hours at 650° C. in a platinum crucible in air. The decomposition of sodium carbonate at temperatures up to 1040° C. is completely suppressed by an atmosphere of carbon dioxide (run 8). The reaction of sodium carbonate with coarse silica sand occurs rapidly at 800° to 850° C. (run 9); the reaction temperature is lowered by grinding the silica (run 10). In this latter run, the first evidence of reaction as shown by weight loss is at about 500° C. The weight loss in these runs wras quantitatively equivalent per formula weight to the carbonate present (Table I), there being an excess of silica or silicate in or
N2c
Sample Na2C03 Na2C03fc
Na2CO* Na2COs Na2C03 Na2COg Na2COj Na2COs Na2COj " Si02 Na2C08 + Si02d Na2C03e
C. per hour, except runs
Crucible covered. Gas flow rate 250 ml. per minute. Heating rate 300° C. per hour to 520° C, then 50° C. per hour. e Maximum temperature 922° C., but sample cooled and held 1 hour at 915° C. after reaching 922° C. c
d
The above results show that sodium carbonate-silica mixtures will lose weight at temperatures as low as 500° C. Howarth, Maskill, and Turner (4, 5) have shown that the reaction between these compounds can occur even at 300° C. Unfortunately, reports of decomposition of sodium carbonate below 800° C. do not specify the material of the vessel: Waldbauer, McCann, and Tuleen (11) report decomposition at 482° C., and Smith and Croad (10) report decomposition above 300° C. but neither mentions the nature of the container. It is likely that they, and others, were observing a chemical reaction with the container rather than the thermal decomposition of sodium carbonate. The authors suggest that sodium carbonate for analytical use be dried by heating in dry air or carbon dioxide using platinum or other inert material as a container. Under these conditions, any temperature in the range of 250° to at least 700° C. should be satisfactory. For work of the highest accuracy it may be necessary to just fuse the VOL. 30, NO. 5, MAY
1
958
·
983
material in
an atmosphere of carbon dioxide as proposed by Richards and Hoover (8). LITERATURE CITED
(1) Balis, E. W., Bronk, L. B., Liebhafsky, . A., Pfeiffer, H. G., Anal. Chem. 27, 1173 (1955). (2) Duval, C., Anal. Chim. Acta 13, 32 (1955).
(3) Easterbrook, W. C., Analyst 82, 383 (1957).
(8) Richards, T. W., Hoover, C. R., J. Am. Chem. Soc. 37, 95 (1915). (9) Simons, E. W., Newkirk, A. E., Aliferis, Ifigenia, Anal. Chem.
(4) Howarth, J. T., Maskill, W., Turner, W. E. S., J. Soc. Glass Technol. 17, 25T (1933). (5) Howarth, J. T., Turner, W. E. S., Ibid., 14, 402T (1930). (6) Kolthoff, I. M., Stenger, V. A., “Volumetric Anatysis,” Vol. II, p. 80, Interscience, New York,
(10) Smith, G. F., Croad, G. F., Ind. Eng. Chem., Anal. Ed. 9, 141 (1937). (11) Waldbauer, L., McCann, D. C., Tuleen, L. F., Ibid., 6, 336 (1934).
(7) Motzfeldt, K., J. Phys. Chem. 59,
Received for review August
1947.
139 (1955).
29,48(1957).
Accepted December 28, 1957.
9,
1957.
Titrimetric Determination of Fluorine Particularly in Aluminum Fluoride LAWRENCE V. HAFF
General Chemical Division, Allied Chemical & Dye Corp., Morristown, N. J. C. P. BUTLER and J. D. BISSO
General Chemical Division, Allied Chemical & Dye Corp., Port Chicago, Calif.
Discordant results for total fluorine reported from different laboratories analyzing identical samples of aluminum fluoride. Serious losses of fluorine occur in distilling and titrating large amounts of fluorine by conventional methods. Recoveries are especially low and erratic when borosilicate glassware is employed. Details of a rapid and accurate procedure for pyrohydrolytic assay of aluminum fluoride were worked Routine analyses of aluminum out. fluoride are greatly expedited by the procedure, which, however, requires a standard sample of alumiIn assaying the standnum fluoride. ard sample, appropriate correction must be made for the losses incurred in the distillation and subsequent concentration of the fluorine. The number of replicate samples required to establish the correction and the magnitude of the correction can be greatly reduced by avoiding use of borosilicate apparatus. were
A
of conventional volumetric methods for determining fluorine in aluminum fluorine indicates that recoveries of large amounts of fluorine from the distillation procedure are only about 98% complete and further losses are incurred when glassware containing boron is used. Under standardized conditions recoveries are reproducible. Analysis by a pyrohydrolytic method of samples standardized by conventional methods established that pyrohydrolysis offers a rapid, accurate, and convenient procedure for assaying alu-
984
study
·
ANALYTICAL CHEMISTRY
minum fluoride. Recoveries were reproducible and about 99% complete. LOSSES
IN
CONVENTIONAL
PROCEDURES
For a number of years two procedures for determining relatively large amounts of fluorides have been used throughout this company. As applied to aluminum fluoride, both involve fusion of the sample with sodium carbonate, followed by acidification and steam distillation
described by Willard and Winter (22). The analysis may be concluded by Armstrong’s procedure (1, 2), as modified by Rowley and Churchill (20). Practically, this titration is limited to determination of small amounts of fluorine, and the small aliquot which must be titrated involves serious magnification of small errors. Alternatively the fluorine in the distillate may be converted quantitatively to fluosilicate, precipitated as potassium fluosilicate, filtered off, and titrated. This procedure is essentially a macromethod and all of the fluoride distilled from a 0.2-gram sample may be titrated conveniently. Unlike the preceding titration, this procedure is insensitive to the interferences of small amounts of phosphates and sulfates and the end point is familiar to most analysts. Accordingly it was designated the standard procedure for the analysis of aluminum fluoride. As rather lengthy evaporation and filtration procedures are involved, manufacturing locations were allowed the option of using the thorium as
nitrate titration samples.
on
all but critical
An investigation of the procedures was initiated when discrepancies appeared
in reported assays. It quickly became apparent that recovery of fluorine was never complete by either method. While the relative percentage of fluorine lost would not be important in microanalysis, it assumed serious proportions in assay of a material containing some 50% of fluorine. Because some of will affect almost all these errors macrodeterminations of fluorine, some details of this work are presented. Standard Samples. No standard fluoride sample was available to the authors at the time this investigation was initiated. Cryolite has been suggested as a standard (7), but its purity is problematical. Sodium fluoride, used by Hoskins and Ferris (8), Kimball and Tufts (9), Reynolds and Hill (17), and Matuszak and Brown (12), is of uncertain assay. Kimball and Tufts (9) reported fluorspar unsatisfactory as a standard and suggested analyzed lead chlorofluoride. The presence of large amounts of either calcium or lead in samples intended to simulate sodium carbonate fusions of aluminum fluoride wTas considered undesirable.
In this work, the exact weight of fluorine involved in each analysis was determined by titrating about 0.2 gram of 48% hydrofluoric acid to the phenolphthalein end point with standard O.lzV sodium hydroxide solution. The titrations were performed in platinum dishes and concluded at the boiling point to eliminate interference by fluosilicic acid in the hydrofluoric acid or dissolved silica in the standard base. The titrated solution, after cooling, is transferred to the still, acidified, and distilled. The distillate is then analyzed by one or both of the above methods to determine the weight of fluorine re-
