ANALYTICAL EDITION
336
Vol. 6, No. 5
potassium bisulfate may be used to absorb ammonia, but no quantitative results are available. In using sulfuric acid in sintered glass beads, it has been stated by others using this method that the acid may spread over the surface of the mercury. The authors, using sintered beads made from 100- to 150-mesh soft glass and following the procedure previously outlined (%),have not experienced this difficulty. However, if too much acid is left on the bead and spreads over the mercury, it can be removed by the use of a follow-up bead of potassium hydroxide.
The methods for the microanalysis of gases described in this series of articles have been developed to the point where they may be used successfully to determine quantitatively an appreciable number of the more common gases. In Table V may be found summarized the gases which have been studied by the authors in this way, along with the reagents and methods employed. By absorbing the gases in the proper sequence, almost any compatible mixture of these gases may be analyzed. For example, the following nine gases could be systematically separated by removing them in the order named: water, carbon dioxide, ethylene, oxygen, carbon . TABLEV. SUMMARY OF GASES ANALYZEDAND REAGENTS USED monoxide, hydrogen, methane, ethane, and nitrogen. It REAQENTS GASES should be possible to accomplish this separation, using an Water vapor Potassium hydroxide or phosphorus pentoxide ( 1 ) original sample of 100 cu. mm., with the same degree of acCarbon dioxide, hydrogen chloride, Potassium hydroxide (I) curacy as by standard macroanalytical methods using an and other acid- roducing gases Acetaldehyde, a n f vapors of similar Potassium hydroxide (6) original sample of 100 cc. compounds
Oxygen Ammonia and other base-producing gases Acetylene
Ethylene and other unsaturated hydrocarbons Carbon monoxide Hydrogen Methane and ethane or any two readily combustibte gases Nitrogen and other inert gases
Phosphorus ( f ) Phosphorus entoxide or conoentrated sulAric acid in sintered glass Cuprous chloride and potassium hydroxide (B) Concentrated sulfuric acid (fuming, for ethylene) ( 2 ) Silver oxide (8j, cupric oxide and potassium hydroxide, or explosion (1j Cupric oxide and potassium hydroxide, or explosion (1) Explosion (1 j
By difference
The dark blue bead for the oxidation of hydrogen and carbon monoxide probably consists in part of a solid solution of cupric oxide in potassium hydroxide (7). The color is a good indication of the effectiveness of the bead as an absorbing agent. Care should be taken not to use a large excess of potassium hydroxide and this hydroxide should be relatively pure. The presence of an appreciable amount of potassium carbonate seems to reduce the reactivity of the absorbent. It is evident that this cupric oxide-potassium hydroxide reagent cannot be used to effect a separation of carbon monoxide and hydrogen. The silver oxide method does this very well with the limitations previously described ( 2 ) .
ACKNOWLEDGMENT The authors wish to acknowledge their indebtedness to P. A. Leighton for his kindly interest and many suggestions in regard to this work.
LITERATURE CITED Blacet and Leighton, IND. ENG.CHBM.,Anal. Ed., 3,266 (1931). Blacet, MacDonald, and Leighton, Ibid., 5,272 (1933). Dennis and Nichols, “Gas Analysis,” p. 157, Macmillan Co., N. Y.. 1929. (4) International Critical Tables, Vol. IV, p. 434,McGraw-Hill Book Co.,N. Y.,1928. (5) JLger, J. Gasbeleucht., 41, 764 (1898). (6) Leighton and Blacet, J. Am. Chem. SOC.,54, 3165 (1932); 55, 1766 (1933). (7) Mellor, “Comprehensive Treatise,” Vol. 3, p. 138, Longmans, Green & Co.,N. Y.,1928. (8) Pierce, Friedenwald, and Freeman, Am. J. Physiol., 104, 553 (1933). (9) Swearingen, Gerbes, and Ellis, IND. ENG.CHEM.,Anal. Ed., 5. 369 (1933). RIOCEIVED May 9, 1934. Presented in part before the Division of Physical and Inorganic Chemistry at the 84th Meeting of the American Chemical Society, Denver, colo., August 22 t o 28, 1932.
Anhydrous Sodium Carbonate for Standardization LOUISWALDBAUER, D. C. MCCANN,AND LAWRENCE F. TULEEN T h e State University of Iowa, Iowa City, Iowa
S
ODIUM carbonate has been used as a standard for many years in neutralization methods of analysis. Many textbooks quote Lunge (4) and require that, in the preparation of sodium carbonate from sodium bicarbonate, the sample should not be heated over 300’ C. Some state that sodium carbonate decomposes above this temperature; some, that it becomes hygroscopic. It was the purpose of the authors to investigate thoroughly the whole matter of preparation, as well as the mass of conflicting statements.
EXPERIMENTAL REAGENTS.Stock c. P. sodium bicarbonate, such as is available in any laboratory, was used as a starting material for sodium carbonate preparation, In terms of sodium bicarbonate this material analyzed 100.45 .per cent. The hydrochloric acid solution was prepared from re-
distilled hydrochloric acid and water redistilled from alkaline potassium permanganate solution, The acid was standardized against calcite, benzoic acid, and potassium acid phthalate, making suitable corrections for indicators. The error in standardization was 1 part in 2500. The sodium hydroxide solution was prepared carbonatefree by the method of Cornog (1). PREPARATION O F SODIUM CARBONATE. Sodium bicarbonate decomposes to sodium carbonate a t any temperature above 102” C. A sample heated a t this temperature for one week analyzed 100.20 per cent in terms of sodium carbonate. When heated to higher temperatures a well-stirred sample remained a t 140’ for some time, Samples heated to 150” C. and cooled a t once were found to be completely converted. It is obvious therefore that any sample heated above 150” C. is completely converted, regardless of the length of time of heating.
INDUSTRIAL AND ENGINEERING CHEMISTRY
September 15,1934
OF HEATING SODIUM BICARBONATE TABLE I. RESULTS
PERCENTIN
TIME
TEMPERATURE
TERMS OF NazCOa
OF
OF
CONVBRSION
HEATINQ
c.
Hours 1 w-eek 1 week 24 24 12 24 4s 24 24
102 115 165 200 260 260 260 300 375 375
100.20 100.04 100.10 100.15 100.07 100,09 100.17 100.09 100.13 100.10
4s
TABLE11.
T~MPERATURE OF HEATINQ
c. 100 375
INSThBILITY O F SODIUM CARBONllTE
TEMPERATURB OF
CONVBRSION
EXPOEURH
375 260 375
NazC03 %
Moisture Carbon dioxide and moisture Carbon dioxide and moisture
99.86 78.50 97.85
Table I11 illustrates the results obtained under the same conditions as in Table 11, but using pure sodium carbonate, prepared by washing with ethyl alcohol, recrystallizing, and drying a t 110" C. TABLE 111.
INSTABILITY OF PURE SODIUM C.4RBONATE
EXPOSURE
NazCOa
%
Carbon dioxide and moisture Moisture
91.62 98.36
It is obvious that both pure and impure sodium carbonate are subject to change on exposure, regardless of the temperature a t which they have been heated. No other conclusions should be drawn from Tables I1 and 111. IMPURITIES. The nature of the impurities causing high results on the samples of sodium carbonate and bicarbonate was not easily discovered. Carbon dioxide analyses were made on several samples but the results could not be made sufficiently precise to warrant any conclusions. It is significant, however, that a sample of sodium carbonate, washed with ethyl alcohol and dried, analyzed, by neutralization methods, 100.00 per cent within experimental error. This would lead one to suspect sodium oxide or hydroxide, since sodium hydroxide is soluble in ethyl alcohol. Table IV illustrates some results of washing sodium bicarbonate alcohol before conversion to sodium carbonate, and washing sodium carbonate after its preparation. The temperature of preparation of sodium carbonate was 260" C. TABLE
Iv. RESULTSO F
Sodium bicarbonate washed before conversion
{
AV.
%
;::t: ]
HE.4TING PURIFIED SODIUM CARBONATE
TIME Houvs 24 24 '
NazCOa
% 100.04 100.05
VAPORPRESSURE. Johnson (3) has found that a t 940" C. sodium carbonate exerts a pressure of 5 mm. of mercury and a t 1100", 10 mm. of mercury. Measurements were made in this laboratory a t temperatures up to 940" C., using an evacuated system and a calibrated McLeod gage. Precautions were taken by prolonged pumping and heating to drive off all adsorbed gases inside the system before a determination was made, These precautions were the result of the work of Hariba and Baba ( 2 ) on the vapor pressure of sodium and potassium chloride. The pumps were sealed from the system and the pressure was checked for 6 hours before heating the sample. Pressures were read a t 10" intervals during heating, allowing 20 minutes at constant temperature for each reading. The first observable pressure due to decomposition was shown a t 482" C. The average of several trials was 499" C. This eliminates any possibility of decomposition below 450" C., a temperature far above any needed for preparation of sodium carbonate from sodium bicarbonate. CRYSTALSTRUCTURE.It was thought that a change in crystal structure might be responsible for the previously mentioned theories of decomposition. X-ray photographs of samples a t 25", 260', 300", and 375" C. were made by the powder method, using a heater to maintain the sample a t the above temperatures. In no case was any change in structure found, other than the expected shift of lines due to expansion. The authors were unable to prepare crystals large enough for a complete crystal analysis, but were able to check the petrographic results of Winchell ( 5 ) ,who states that sodium carbonate crystals from fusion are biaxial, negative, with 2V = 34" =t 3". The authors, using crystals grown from fused sodium hydroxide, also noted that the crystals appear to have three good cleavages and show marked lamellar twinning. Extinction angles with the cleavage were found varying from zero to 30". From the above evidence, it appears that anhydrous sodium carbonate may be monoclinic. CONCLUSIONS Washing so-called c. P. anhydrous sodium carbonate with ethyl alcohol and drying it a t 110' C. render it satisfactory as a volumetric standard. Anhydrous sodium carbonate does not decompose below 450", nor does it change its crystal structure below 375' C. The impurities leading to high results are very probably due to sodium oxide and (or) hydroxide and come from the decomposition of sodium bicarbonate during its conversion to sodium carbonate. LITERATURE CITED
WASHING
Na?COa
to carbonate regardless of temperature, since all previous samples analyzed high. Further evidence that no other decomposition takes place is given by Table V in which samples of the above purified sodium carbonate were heated. TABLE v. RESULTSO F
It would seem from the analyses in Table I that temperature and time of heating made very little difference. There is no indication of decomposition a t temperatures as high as 375" c. INSTABILITY. Instability in the presence of moist air and moist, carbon dioxide-laden air is rather difficult to measure. The results in Table I1 should be taken only qualitatively. The samples were those of Table I (of 24 hours' heating) and were in all cases exposed for 6 hours.
337
100.17
Bbove sodium carbonate washed after conversion
{ ::::::I
100.03
Sodium carbonate prepared directly from c. P. sodium bicarbonate and washed after conversion
{ 100.02 100.04
100'03
100.03
It is obvious from Table IV that some sodium oxide or hydroxide is formed in the conversion of sodium bicarbonate
(1) Cornog, J. Am. Chem. Soc., 43,2574 (1923). ' (2) Hariba and Baba, Bull. Chem. SOC.J a p a n , 3, 13 (1928). (3) Johnson, 2. phys. Chem., 62, 330 (1908). (4) Lunge, 2. angew. Chem., 17, 231 (1904). (5) ,W'inchell, -4.N., "Artificial Inorganic Solid Substances and Artificial Minerals," p. 199, John Wley & Sons, N. Y., 1931. RECEIVEID Aprd 6, 1934, Extract from the thesis submitted by D. C. McCann for the degree of Ph.D., June, 1934. The work on vapor pressure measurements of the decomposition of sodium carbonate is an extract from the M.S. theais of Lawrence F. Tuleen, submitted in August, 1932.
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