LITERATURE CITED
Figure 5. Portion of a microwax spectrum at fast scan showing presence of CIHZn-l* peaks at m / e 170 and 184
Table IV.
Peak 225 226
Measured mass 225.2582 225,1637 225.0782 226,2604 226.1732 226.0867
Mass Measurements on Ozokerite
Eauivalent hydrocarbon Formula Mass CiBHaa 225.2582 Ci&a 225.2582 C1&33 225.2582 Cl6Ha.r 226.2661 CiaHa4 226.2661 CiaH34 226.2661
about half a n hour and peaks had to be retuned a t stages of a quarter of an hour to new ion repeller and beamcentering positions. However, the presence of C,H2,-2 ions is confirmed
AM
Assignment
0.0945 0.1800 0.0057 0.0929 0.1794
C-Hii S&Hi6 C"-CH C-Hi2 St-CaHis
Peak C16H33 Ci7Hzi CizHi7Sz CisC'aHaa CuHn CizHi8Sz
by mass measurement of the doublets at m/e 170 and 184 (Figure 5). Future work is planned to provide a detailed study of the complex high resolution spectra of microwaxes.
(1) Beynon, J. H., "Mass Spectrometry
and Its Applications to Organic Chemistry," Elsevier, Amsterdam, 1960. (2) Carlson, E. H., Paulissen, G. T., Hunt, R. H., O'Neal, M. J., Jr., ANAL. CHEM.32, 1489 (1960). (3) Clerc, R. J., Hood, A., O'Neal, M. J., Jr., Zbid., 27, 868 (1955). (4) Elliott, R. M., Craig, R. D., Errock, G. A., Proceedings of Fifth International Instruments and Measurements Conference, Vol. I, p. 271, Academic Press, New York, 1961. (5) Levy, E. J., Galbraith, F. J., Melpolder, F.,W., "Advances in Mass S ectrometry, R. M. Elliott, ed., Vof 2, p. 395, Pergamon, London, 1963. (6) Lumpkin, H. E., ANAL. CHEM. 36, 2399 (1964). (7) Minchin, J., J. Znst. Petrol. 34, 542 (1948). (8) O'Neal, M. J., Jr., Weir, T. P., Jr., ANAL.CHEM.23, 830 (1951). (9) Reid, W. K., Xead, W. L., Bowen, K . hl., Mass Spectrometry Conference, Paris, IP/ASTM/GAMS, September 1964. (10) Reid, W. K., Mead, W. L., West, A. R., ANAL. CHEM.36, 1140 (1964). (11) Thornton, E., West, A. R., 2.Anal. Chem. 170, 348 (1959). (12) Tunnicliff, D. D., Wadsworth, P. A., Schissler, D. O., ANAL. CHEM.37, 543 (1965). (13) Van Katwijk, J., Appl. Spectros. 18, 102 (1964). RECEIVEDfor review August 16, 1965. Accepted November 29, 1965.
Estimation of Solubility of Bismuth Compounds in Liquid Ammonia ANNIE G. SMELLEY,' FRANCIS E. BRANTLEY,2 and ARTHUR F. FINDEW U. S. Bureau of Mines, University, Ala.
b The solubilities of six bismuth compounds in liquid ammonia are reported. The compounds were dissolved in liquid ammonia and equilibrated for 2 hours at - 3 3 " C. An aliquot of the solution was then analyzed for bismuth. The bismuth was determined polarographically using a dropping mercury electrode with 1M HzS04 as a supporting electrolyte, Triton X-100 as a maximum suppressor, and a mercury pool reference electrode. Numerical results are given for the solubility of the triiodide, trichloride, tribromide, nitrate, sulfide, and lactate of bismuth.
P
to a n investigation of electrodeposition of bismuth from liquid ammonia, a literature search was made to obtain solubilities of bismuth salts which were being considered for use in the study. I n spite of the RELIMIKARY
large amount of work reported in the literature concerning solubilities in liquid ammonia, numerical values for solubilities are cited for only a small number of compounds. Solubilities for bismuth compounds are listed in descriptive phrases such as soluble, very soluble, slightly soluble, and insoluble (3, 4, 6). Much information is reported concerning the reactions of bismuth halides with ammonia, but only general statements concerning solubilities have been made. For these reasons this work was started to obtain quantitative solubility data. Several methods are described in the literature for determining solubilities in liquid ammonia. These are either elaborate and time consuming for precision measurement (IO), or rather simple techniques yielding qualitative statements of solubilities (5). The method and apparatus used for studying the solubility of the bismuth salt was a
modification of those described in the literature. A conventional polarographic method of analysis was used to determine the amount of bismuth dissolved in the liquid ammonia. It was decided that for the work reported here, any additional accuracy that might be gained by the use of special procedures in drying ammonia, purifying chemicals used, and establishing equilibrium conditions would not be justified. Errors are introduced when these factors are disregarded; however, the objective was to establish a n order of magnitude rather than expend the time necessary for obtaining more accurate results. Chemist, Bureau of Mines, U. S. Department of the Interior, Tuscaloosa, Ala. Present address, Division of Minerals, Bureau of Mines, U. S. Department of the Interior, Washington, D. C. a School of Chemistry, University of Alabama, University, Ala. VOL. 38, NO. 3, MARCH 1966
449
fitted with a two-hole neoprene stopper, This was attached by a flexible hose to a 190-ml. solubility tube containing a stainless steel filter stick of fine porosKO. Solubility, ity which extended nearly to the of Grams Salt er Salt detns. 100 Grams $€IBbottom of the tube. The ammonia solution was sampled through the filter Bismuth stick and an aliquot was collected in a triiodide 10 0.022 f 0,006 graduated cylinder immersed in liquid Bismuth nitrogen. -4s the apparatus was open nitrate, to the atmosphere the solubility deDentahvterminations were made a t atmospheric drate 8 0.023 f 0.006 pressure. Bismuth trichloride A Fisher Elecdropode and a Sargent 6 0.010 f 0.002 Bismuth Model XV Polarograph were used to sulfide 5 0.010 f 0.003 record the polarographic waves. A Bismuth dropping mercury electrode was used lac tat e 5 0.12 f 0.01 with a mercury pool reference in a Bismuth Heyrovsky cell obtained from E. H. tribromide 5 0.00011 f 0.00004 Sargent and Co. Procedure. Transfers of liquid ammonia were made from the cylinder directly to the dewar. The dewar EXPERIMENTAL was then incorporated into the apparatus and transfers of liquid Reagents. The liquid ammonia ammonia were made by inverting the was obtained from T h e Matheson dewar and pouring the ammonia Co., Inc. (Trade names and manuthrough the flexible tube into the tube facturer's names are used for informacontaining the bismuth salt. The samtion only and endorsement by the ple tube contained a large plasticBureau of Nines is not implied.) coated magnetic stirring bar for agitaPurified grade bismuth nitrate pentation of the solution. The mixture was hydrate and bismuth trichloride, anstirred for 2 hours with additional hydrous, were obtained from J. T. ammonia being added a t intervals to Baker Chemical Co. Bismuth trikeep the liquid level constant a t all bromide, anhydrous, was obtained from times. At the end of the 2-hour mixing City Chemical Corp. Bismuth triperiod, the volume was adjusted to iodide was obtained from K &. K Labthe calibration mark in the sample tube, oratories, Inc. Bismuth sulfide and vigorously stirred for an interval of bismuth lactate were obtained from several minutes, and an aliquot of the A. D. Mackay, Inc. The bismuth saturated solution was taken. This metal was analytical reagent grade was done by placing a clean, dry gradufrom The Nallinckrodt Chemical Works. ated cylinder under the tube leading Triton X-100 was obtained from Rohm from the stainless steel filter stick and & Haas Co. X-ray diffraction patclosing all other outlets to the system. terns of each of the compounds were obThe increase in vapor pressure in the tained and were compared with the system forced the saturated liquid ASTM compilation ( I ) . A11 diffracammonia solution through the filter tion patterns agreed with published into the graduated cylinder which was values except bismuth trichloride and cooled by liquid nitrogen. The volume bismuth lactate. The patterns for of aliquot was usually around 15 ml. bismuth trichloride as obtained from and was measured to the nearest 0.1 two different sources, however, were ml. The aliquot was then transferred nearly identical. The sample which to a 250-ml. beaker and evaporated to was used for the solubility determinadryness. Interferences from anions in tions was analyzed for bismuth and the polarographic determination were chlorine and found to be a stoichioremoved by evaporating the so!ution metric bismuth trichloride. A controwith sulfuric and nitric acids and fumversy about the x-ray diffraction pattern ing to dryness. The residue was disof bismuth trichloride has been noted solved in 20 ml. of 2.5M sulfuric acid previously (12). The pattern for the and diluted to 50 ml. in a volumetric sample of bismuth lactate that was used flask. The concentration of the bisfor the solubility determination did not muth ions in this solution was then agree with the heptahydrate form, determined polarographically (7) using which was the only pattern given for a dropping mercury electrode with a this material in the ASTM compilation. mercury pool reference. Maxima in h standard bismuth solution was prethe polarographic waves were removed pared by dissolving a weighed amount of by adding 4 drops of a 0.2% Triton bismuth metal in nitric acid. The soluX-100 solution. The solutions were tion was evaporated to dryness and degassed with helium, and the helium taken u p in 0.5N H2S04. This solution was allowed to flow over the surface was used as standard for the polaroof the solution while the polarogram graphic analyses. A 0.2% solution of was recorded. Triton X-100 mas used as a maximum The bismuth wave occurred with a suppressor. Triple distilled mercury was half wave potential of -0.42 volt us. used in the dropping mercury electrode a mercury pool. The wave height was assembly and helium gas was used to measured and the concentration calcudegas the polarographic cells. lated using a calibration curve prepared Apparatus. The apparatus used under identical circumstances but with in this study is similar to t h a t destandard bismuth solutions. These scribed previously (11). A dewar was Table 1. Solubility of Bismuth Salts in Liquid Ammonia at 33" C.
450
ANALYTICAL CHEMISTRY
solutions were prepared from the standard bismuth stock solution. Initial polarograms were made on standard bismuth solutions using 1M H804 as the supporting electrolyte and methyl red-bromcresol green as the maximum suppressor. A persistent maximum was obtained at the beginning of the wave which could not be suppressed with u p to 20 times the normal amount of methyl red-bromcresol green solution. After Triton X100 was used as the maximum suppressor, the waves contained no maxima and were well formed. RESULTS AND DISCUSSION
Values are reported in Table I for the solubility and the average deviation as determined for each of the six bismuth salts in liquid ammonia at its boiling point, -33" C. Hunt and Broncyk (6) stated that bismuth trichloride and bismuth nitrate appeared to be insoluble while Franklin and Kraus (3) reported a table of solubilities listing bismuth triiodide as being slightly soluble and bismuth nitrate as being soluble on the addition of ammonium nitrate. Estimates by Gore (4) were "moderately soluble" for bismuth trichloride and "insoluble" for bismuth sulfide. Bismuth bromide is reported to be more soluble than bismuth chloride (8). Bismuth iodide is characterized as rapidly soluble (2) or easily soluble (8). The usual trend in solubilities for the bismuth halides in water is in the order C1 > Br > I. This trend has not been observed in this work in liquid ammonia. Because of the fact that bismuth halides have been reported to form addition compounds with ammonia, it is likely that the solubility of the addition compounds has been measured ( 8 ) . The tabulations are based on the bismuth in solution and are calculated for the solubility of the simple salts. Because no numerical values for the solubilities of these bismuth compounds have been reported, there was no basis for comparison with previous work. Solubility determinations were made with sodium chloride for which data have been reported (6). Using a procedure similar to that for the bismuth salts, an average of five runs in this apparatus yielded a value of 5.2 f 0.2 gram of NaC1 per 100 grams of NHI. This is higher than the interpolated value of 3.43 grams of NaCl per 100 grams of saturated solution which was obtained using solubilities of 4.0 grams of NaCl per 100 grams of saturated solution at -30' C. and 2.1 grams of NaCl per 100 grams of saturated solution at -40" C. (9)* LITERATURE CITED
(1) American Society for Testing and
Materials. Powder Diffraction File. Philadelphia, Pa. Bismuth lactate heptahydrate No. 1 Set (1945), Card 0075; Bismuth sulfide No. 6 Set (1955),
Card 0333; Bismuth triiodide S o . 7 Set (1957), Card 269; Bismuth trichloride S o . 11 Set (1961), Card 600, Bismuth tribromide KO. 11 Set(1961), Card 601, Bismuth nitrate pentahydrate No. 12 Set (1962), Card 148. ( 2 ) Franklin, E. C., J . Am. Chem. SOC. 2?, 820 (1905). (3) Franklin, E. C., Kraus, C. A., Am. Chem. J . 20. 820 (1898). (4) Gore, G.,’Proc.~Roy. SOC.(London) 21, 140 (1873). ( 5 ) Hunt, H., J . Am. Chem. SOC.54, 3509 (1932).
(6) Hunt, H., Broncyk, L., Zbid., 5 5 , 3528 (1933). ( 7 ) Kolthoff, I. M., Lingane, J. J., “Polarography,” pp. 550-1, Interscience, New York, 1952. (8) Schwarz, R., Striebich, H Z . Anorg. illlgem. Chem. 223, 339 (193k). (9) Seidell, A., “Solubilities of Inorganic and Metal Organic Com ounds,” 3rd ed., p. 1248, D. Van gostrand Co., Xew York, X. Y., 1940. (10) Watt, G. W., Jenkins, W. A,, Robertson, C. V., ANAL. CHEM. 22, 330 (1950).
(11) Watt, G. W.,Moore, T. E., J . Am. Chem. SOC.70, 1197 (1948). (12) Wolten, G. M.,hlaver. S. W., Acta. dryst. 11, ‘739 (1958). ”
RECEIVED for review February 1, 1966.
Resubmitted December 7 , 1966. Accepted December 30, 1965. The work upon which this report is based was done under a cooperative agreement between the Bureau of Mines, U.S. Department of the Interior, and the University of Alabama.
Titrimetric and Equilibrium Studies Using Indicators Related to Nile Blue A MARION MACLEAN DAVIS and HANNAH B. HETZER National Bureau of Standards, Washington, D. C.
b The behavior of Nile Blue A and several closely related dyes in nonaqueous spectrophotometry and titrimetry was explored. Nile Blue A was used as the salt, as the anhydro-base, and as the oxazone. The other dyes were a salt or anhydro-base having an alkyl, aralkyl, or aryl group attached to the 5-amino nitrogen in place of hydrogen. The arylamino derivatives, though less basic than Nile Blue A, are much stronger than Methyl Yellow. They are useful reference bases in aprotic solvents. Their applicability in determining stoichiometry and the relative strengths of strong and moderately strong acids is discussed. The strengths of the indicators relative to sulfonephthaleins and other common acid-base indicators are different in aqueous and nonaqueous solvents, pointing to a useful role in nonaqueous differentiating titrations.
T
acid-base behavior and perform acid-base titrations in anhydrous solvents have been hampered by the lack of strong, crystalline bases. Tribenzylamine would be a n admirable reference base if it were not so weak. 1,3-Diphenylguanidine has proved very useful but does not meet every need. A logical place to seek strong reference bases is among compounds in which proton addition is favored because the resulting cation is stabilized by resonance. Such stabilization accounts for the relatively great basicity of diphenylguanidine and other guanidine derivatives. =inhydrobases derived from cationic acid-base indicator dyes are another group of compounds in which proton addition is favored by resonance. Because of their high molar absorptivity they are exceptionally sensitive reagents. HOSE WISHING TO STUDY
20234
Our work with dyes of the Xile Blue group was first undertaken in the early part of World War I1 during an effort to find sensitive, and also stable, basic indicator standards for the detection and semiquantitative estimation of undesirable acids in aprotic media such as airplane motor fuels, lubricants, transformer oils, and used drycleaning solvents (“Stoddard Solvent,” halogenated hydrocarbons). Improved indicators mere also needed for the quantitative estimation of acids by titration. Sile Blue X and the related cationic indicator dyes discussed in this paper belong to the group known as “oxazines” or, more specifically, as “benzophenoxazines” (‘7, SO). Oxazines contain two aromatic rings joined by a nitrogen and an oxygen atom, as shown in Figure 1. Vsually they are prepared in the form of salts (I), the anion being chloride or sulfate. Addition of alkali converts the salt into an anhydro-base (11), whereas heating the salt in an acidified aqueous solution brings about its hydrolysis to form the oxazone (111). Obviously anhydro-bases can be formed only from salts having a t least one proton a b tached to the &amino nitrogen. It is also evident that any series of salts having the same substituents (R‘) attached to the 9-amino nitrogen will yield the same oxazone. Benzophenoxazines have been synthesized and studied more extensively than the analogous phenoxazines (dyes which resemble compounds I to I11 except for the absence of the ring attached at position a ) . Early in the investigations it became apparent that satisfactory practical tests could not be developed without first performing systematic studies of acid-base behavior in aprotic solvents. Inadequacies in the dyes then available were also disclosed. When new compounds of the Nile Blue group became
available supplemental studies were made. This paper summarizes exploratory investigations of two kinds: (1) Spectrophotometric studies of acidbase equilibria in various nonaqueous solvents (mainly benzene or acetonitrile). The oxazone of X l e Blue A (formula 111, R’ = Et) and several anhydro-bases were used as reference bases in combination with various acids.
+ I
TI
A m Figure 1. General formulas of a benzo [a]phenazoxonium salt (I), and the corresponding anhydro-base (benzophenoxazine) (11) and oxazone (benzophenoxazone) (Ill) In the salt, anhydro-base, and oxozone of Nile Blue A, R = H and R ’ = Et. Commonly, X - = CI- or 1 1 2 504-2
VOL. 38, NO. 3, MARCH 1966
451
