1544
ANALYTICAL CHEMISTRY
Factors which did influence the R, values significantly included temperature and distance between solvent surface and the initial spot (or length of exposed wick). I t is these latter two factors which should be controlled for accurate work. Factors such as p H of reagents and quality of paper used, which are known to influence R,, were kept constant for all the runs. ACKNOWLEDGMENT
The authors wish to acknowledge the aid and encouragement given them by Bruno W. Volk, director of laboratories, throughout the course of these studies, and to express appreciation to Renee Eisner for the typing and editing of the manuscript.
(9) Giri, K. V., Radhakrishnan, A. li., and Vaidyanathan, C . S. A N ~ LCHEM., . 24, 1677 (1952). (10) Giri, K. V., and Rao. N. -4.N.,A-uture, 169, 923 (1952). (11) Ishii, S., and Ando, T., Bull. ('hem. 9oc. J a p a n , 2 3 , 172 (1950). (12) Jermyn, hZ. A., and Isherwood, F. A., Riochem. J . , 44, 402 (1949). (13) Jirgensons, B., University of Texas Publication, R e p t . 5109, 56 (May 1951). 24, 643 (14) Kowkabany, G. N., and Cassidy, H. G.. ANAL.CHEW., (1952). (15) McFarren, E. F., I b i d . , 23, 168 (1951). (16) RIcFarren, E. F., and ~ l i l l sJ. , A , , I b i d . , 24, 650 (1952). (17) Marchal, J. G., and llittwer, T.. Compt. rend. SOC. bid,,145, 417 (1951). (18) IIarchal, J. G., and lIittwer, T.. Kofrinkl. S e d . A k a d . Wetenschap. Proc.. 54C, 391 (1951). (19) lIeredith, P., and Sammons, H. G., Biochem. J . , 4 9 , lxii (1951). 21, 1429 (1949). (20) Sluller. R. H.. and Clem D. L., .&NIL. CHEM., (21) Ibid., 23, 396 (1951). (22) Ibid., p. 403. (23) Ibid., p. 408. (24) Novelle, L., Sature, 166, 1000 (1950). (25) Rockland, L., and Dunn, AI., Science, 109, 5.39 (1949). (26) Roseheek, S.,Chem. Weekblad, 46, 813 (1950). (27) Rutter, L., Analyst, 75, 37 (1950). (28) Rutter, L., Nature, 161, 436 (1948). (29) Toennies, G., and Kolh, J. J., ANAL.CHEM.,23, 823 (1951). (30) Winsten, W. A., Science. 107, 605 (1948). ~~
LITERATURE CITED
Bate-Smith, E. C.,,Biochem. Sac. Symposia, 3 , 62 (1950). Block, R. J., ANALCHEM.,22, 1327 (1950). Burma, D. P., Nature, 168, 565 (1951). Consden, R., Gordon, A . H., and Martin, A. J. P., Biochem. J., 3 8 , 224 (1944).
Draper, 0. J., and Pollard, A . L., Science, 1 0 9 , 4 4 8 (1949). Giri, K. V., Current Sci. ( I n d i a ) ,2 0 , 296 (1951). Giri, K. V., Krishnamurthy, K., and Venkataaubrarnanian, T. A , , Ibid., 2 1 , 11 (1952). Ibid., pp. 21, 44.
RECEIVED for review August 21, 1952. -4rcepted .Jiily 29, 1953.
Rapid Precipitation of Barium Sulfate ROBERT B. FISCHER AND T. BEN HHINEHAMMER' D e p a r t m e n t of C h e m i s t r y , Indiana University, Bloomington, f n c l .
This study was undertaken to determine by means of light microscope and electron microscope techniques the influence of various conditions upon the morphology and size of particles of precipitated barium sulfate, and to develop and test a procedure for the rapid gravimetric determination of sulfate as barium sulfate, in which the reagents may be poured together rapidly and in which little or no digestion time is required prior to filtration and washing. Conditions investigated microscopically include
M
ANY workershave investigated the precipitation of barium
sulfate, and the subject is treated in detail in both elementary and advanced textbooks of quantitative analysis. Nevertheless, gravimetric determinations of sulfate as barium sulfate are very time-consuming and still cause considerable difficulty both in student work and in analytical laboratories. The present study was undertaken for a twofold purpose: to determine by means of light microscope and electron microscope techniques the influence of various conditions upon the morphology and size of particles of precipitated barium sulfate, and t o develop and test a rapid procedure for the determination of sulfate as barium sulfate. The particle size and morphological characteristics of any gravimetric precipitate are of significance from the standpoints of rate of settling, ease of filtration, and surface coprecipitation of impurities. Popoff and Neuman (7) employed the light microscope to study several factors involved in a reverse method for the precipitation of barium sulfate. The electron microscope has been used to ascertain the influence of several variables upon the morphology of barium sulfate crystals, covering principally wide ranges of con-
' Present address, Mound Laboratory, burg, Ohio.
Momanto Chemical Co , Miamis-
sulfate ion concentration, barium ion concentration, quantity of barium ion in excess, pH, temperature, and presence of foreign ions. The quantitative procedure developed yields results that are both precise and accurate within 1 part per thousand, even in the presence of high concentrations of alkali ions. This procedure possesses no real disadvantage in comparison to conventional procedures, and it has great advantages from the standpoints of speed and ease of handling, filtration, and washing.
centrations, digestion of the precipitate, and coprecipitation with crystals formed very slowly ( 3 , s ) . Hintz and Weber ( 5 )reported a procedure for the rapid precipitation of barium sulfate, and Rieman and Hagen (8)found that the rapid precipitation method yielded better results in the presence of high concentrations of alkali chloride than did slow precipitation procedures. However, no rapid method of precipitating barium sulfate has yet found very Ridespread acceptance. EXPERIMENTAL PROCEDURES
Twa general methods mere employed for the precipitations.
Method A. Barium chloride solution, 0.023 M ,is prepared from reagent grade salt and the solution is allowed to stand a t room temperature for 30 days prior to use. Potassium sulfate solution, 0.015 M , is prepared from reagent grade salt. Total volume of reagents is 6 ml., with 2 to 4 ml. of each reagent used as required in each precipitation. The pH of the sulfate solution is adjusted to 1 with hydrochloric acid. For observations of morphology, temperature is 95" to 100" C. for both reagents; for size measurements, temperature is 60' C. for sulfate solution and 30" C. for barium chloride solution. Reagents are mixed rapidly by pouring barium chloride solution into sulfate solution. A 10 to 20% excess of barium chloride solution is included a t the time of initial mixing.
V O L U M E 2 5 , NO. 10, O C T O B E R 1 9 5 3 Method B. Barium chloride solution, 0.02 M , is prepared from reagent grade salt and the solution is allowed to stand a t room temperature a t least 24 hours prior to use. Sulfate solution is 0.006 to 0.16 M , preferably about 0.01 M . The pH of the sulfate solution is adjusted to 1 with hydrochloric arid, unless alkali salts are present in high concentration, in which case p H is adjusted to 5 (see discussion below). Total volume of reagents is 150 to 200 ml. Temperature is 95" C. for both solutions. Barium chloride solution is poured rapid1 into sulfate solution. An initial 10 to 20% excess of barium chgride solution is used, although more may be added just prior to filtration if desired. Vigorous hand stirring is employed for 2 or 3 minutes. N o digestion period is necessary other than a few minutes to allow the beaker to cool. Filtration (9 rapid, only very brief washing is necessary, and ignition over burner is for 20 minutes a t 600" to 800' C. Finally, the precipitate is weighed.
1545
the electron microscope, each precipitate was washed prior to mounting for observation in this instrument; conventional procedures for mounting particulate specimens were employed. EXPERIMENTAL OBSERVATIONS AND DISCUSSION
Effect of Sulfate Ion Concentration. Method A was employed, except that the sulfate ion concentration was varied over the range 0.0026 to 0.026 M . As this range is equivalent t o 0.06 to 0.6 gram of barium sulfate in 100 ml. of solution, it includes the range normally encountered in routine laboratory practice. The size data are as follow: K ~ S O I Molarity ,
4 .0926MOLAR K,SQ
L .010 MOLAR K,SO+
.016MouR K*SO+
.OPI MOUR K,SO+
.026YoLpR K,SO,
Figure 1.
Size, Microns
The observed trend is contrary to expectations from the von n'eimarn law, but it is in keeping with earlier reports that, over a wider range of concentrations, the size decreases as the concentration is varied above and below some optimum value. Initial pore sizes of common filter media are of the order of several microns, so a sulfate molarity of 0.006 may be considered as an approximate minimum for gravimetric determinations. Morphological characteristics are shown in Figure 1. The natural habit of barium sulfate is orthorhombic, and ideal rrystals are rectangular in cross section (Figure 1, upper right). The degree of perfection of the crystals is excellent up to 0.016 M; it decreases somewhat a t 0.021 M and markedly a t 0.026 M . Precipitates consisting of well-perfected crystals settle much more readily than those consisting of irregular ones.
Effect of Sulfate Ion Concentration upon Barium Sulfate . 0 2 YOLAR
Method A was employed, with individual modifications as indicated below, for all microscopic observations. The conditions listed are comparable to those commonly encountered in analysis, except for the small total volume, which was selected for economy and convenience, particularly in view of the fact that no stoichiometrically quantitative data were to be derived with Method A. Method B was used for the complete quantitative determinations of sulfate as barium sulfate. The sulfate solution represents the "unknown." The conditions specified for this procedure were actually derived from the results of the experiments with Method A. The light microscope was employed for all measurements of particle size. For each measurement, the precipitation was repeated in quintuplicate, and a t least 20 crystals from each of the five precipitates were measured lengthwise and averaged. Each sample was mounted for observation by placing 1 drop of the well-shaken suspension of the precipitate on a clean glass slide and covering it with a cover glass to ensure uniform distribution of crystals throughout the field of view. Morphological characteristics of the precipitated particles were observed principally with the electron microscope, although they were checked with the light microscope. The two instruments yielded identical results in so far as available resolution permits comparisons. All micrographs presented in this report are electron micrographs. Ae only dry specimens could be observed in
.OS MOLAR
% -5 .IO MOLAR
.OIMU.AR
IOOIEXCESS
.IOYOLAR
IOOIEXQSS
Figure 2. Effect of Barium Chloride Concentration and 100% Excess Barium Chloride upon Barium Sulfate
Effect of Barium Ion Concentration. Method A was employed except for adjustments of the barium ion concentration over the range 0.01 t o 0.1 M . This range includes the values
1546
ANALYTICAL CHEMISTRY
recommended in most of the common textbooks of quantitative analysis. The size data are as follows: KpSOa, Molarity 0.100 0.082
Size, Microns 16.2 15.5 15.3 14.4 14.4 14.4 13.6 12.5 9.0 6.9
0.066 0.049 0.041 0.036 0.030 0.023 0.016 0.010
Thus the size is fairly constant above about 0.025 M but drops off rapidly below that value. The left micrographs of Figure 2 reveal that the crystals are reasonably uniform a t concentrations of about 0.02 M , but that the degree of perfection decreases markedly a t higher concentrations of barium chloride. Unfortunately, it is not feasible to include micrographs of the larger variety of grotesque shapes obtained a t the 0.1 M concentration. Effect of Excess of Barium Chloride. For the micrographs of Figure 2, Method A was employed, except that the concentration of barium chloride was varied; in each case the volume was selected so that the stated excess would be provided. For the micrographs of Figure 3, Method A was employed, except that the concentration of sulfate ions was varied; in each case the volume of standard barium chloride was selected to provide the stated excess. With the possible exception of the case in which the sulfate ion concentration is a t its lowest value, the 100% excess of barium chloride markedly decreases the perfection. Both flaring and splitting of the ends occur, and extremely ragged surfaces develop. The ragged particles agglomerated much more rapidly than did the more perfected crystals, thus enhancing the danger of ropreeipitation by occlusion. Effect of pH. Method A was employed, except that the p H was adjusted in half-unit steps from 0.5 t o 5.0. Size data are as fol-
smaller, more perfected crystals. The large, irregular crystals exhibited a tendency to break into smaller fragments upon handling. Effect of Temperature. Method A was employed, except that the temperatures of the reagent solutions were adjusted from 25' to 100" C. As the temperature was increased, the crystals became larger and more perfected. The tendency for barium sulfate to assume its natural, orthorhombic habit of dense, rectangular crystals is thus enhanced a t elevated temperatures. Turbidimetric measurements indicated that thc particles precipitated from hot solutions settled much more rapidly than did those from colder reagents. The method and rate of mixing of the two solutions were more critical at high than at low temperatures.
,010 yam
,016 Y U R
lows: pH
Size, Microns
0.5 1.0 1.5 2.0 2.5 3.0 3.5 4.0 5.0
9.0 13.3 18.0 22.4 27.1 31.4 37.3 43.1 51.8
A plot of these data in terms of size us. p H consists of a straight line, with the size increasing as the p H increases. This trend is apparently contrary to the von Weimarn law, in view of the increased solubility of barium sulfate a t low p H values. H o w e v e r , Figure 4 reveals that the crystals formed at the lower p H values are decidedly more perfected than those formed in less acidic media. The dense, compact, and well-perfected crystals formed a t low p H values settle very rapidly, as determined quantitatively by turbidimetric measurements. Even though the crystals formed a t higher pH values are larger, they are fluffy and very buoyant and probably exhibit greater specific surface area than do the
.02I YOLbR
.026 WOCAR
Figure 3.
Effect of 100% Excess Barium Chloride upon Barium Sulfate
Table I.
Sulfate Determinations by Rapid Precipitation, Method B
pH of Sulfate Neutralization prior= Line No. to Filtration Solution 1 1 No 2 3 4 5 6 7 8 9 10 11 12 13 14 15 16
1 2 4 5 63 1 5.1 1 5.1 1 2 4 5 6
1
Foreign
Electrolyte None Eone None None None None 1 . 2 M KC1 1 . 2 M KCI Sone None 1.2M , NaCl 1 . 2 iM NaCl 1 . 2 M NaCl 1 . 2 ill XaC1 1 , 2 M NaCl
YeS
Yes No NO No No No No No No No No No No Yes
17 18
NO NO
19
No
0.012 M
NaXOa
0.012M FeCl; Versene Fe-3 (30 ml./l.) Versene Fe-3 (630 m1.A) and 0 012 M
Error, Partsb per Thousand - 8 , -8, -8, -8, -8, -8, - -88, - 8 , - 9
-8, -8, -8, -8 - 8 , - 7 , -8, - 8 - 1 , - 1 . 0. 0. - 1 , - 1 t l . -1, + 1 , 0, + I , - 1 -21, -22, -22, -22 - 1 . 0, 0, - 1 , - 1 , + 1 - 1 , -1, - 1 , - 1 , -1, - 1 - 1 , 0, + I , - 1 , 0, + l - 7 , - 8 . - 8 , -8, - 7 - 7 . - 8 . - 8 , -8, - 7 . - 7 -7, -7. -7. -7, -7, -7 - 3 , 0 , - 1 , 0, -1, +1 + 1 , 0, + 1 , - 1 , +1, + 1 +12, + 1 1 , +12. +12, +12, +12 -15, -16, -15, -1.5, -163, -16 -1,
- 1 , 0, - 1 ,
-1
FeCla 0,+ I , + 1 , + I , + 1 Versene Fe-3 (60 m M . ) and 0.024 ,lf FeCh +2, +2, +2, +2, + 2 +2 Some precipitate suspensions were neutralized prior to filtration. b Each sample contained enough sulfate to yield about 0.3000 gram of barium sulfate. Errors are expressed in terms of parts per thousand with a negative error indicating that the "found" weight was less than the "calculated" weight. The several error values on each line arose from replicate determinations. C Versene Fe-3 is ethylenediaminetetraacetic acid. used in the form commercially available from the Bersworth Chemical Co. 20
1
No
NapSOc
V O L U M E 25, NO. 10, O C T O B E R 1 9 5 3 Effect of Foreign Ions. Microscopic observations were made of precipitates formed by Method A with the inclusion of various foreign ions in the sulfate solution a t the time of precipitation. The many observations may be summarized as follows: The foreign ions frequently cause an increase in crystal perfection, an increase in size uniformity and, a t higher concentrations, a decrease in crystal size; the ions tested may be listed in order of decreasing effect upon crystal morphology as: nitrate (most imperfection), phosphate, chlorate, ferric, aluminum, and potassium (most perfection). The relative magnitudes of the effects of the several coprecipitants may be explained for the most part on the basis of a solid solution type of coprecipitation. The nitrate ion is much smaller than the sulfate ion, while the chlorate ion is almost the same size as is the sulfate ion; thus, as nitrate does not fit so well as the chlorate in place of a sulfate, nitrate would be expected to cause more imperfection on the basis of a solid solution mechanism of coprecipitation. Likewise, a potassium ion fits very well into the space normally occupied by a barium ion, and the influence of potassium ion upon crystal morphology is very slight. Both iron and aluminum ions are much smaller than barium ions, and the coprecipitation with these two substances may involve some sort of adsorption mechanism.
W0.s
98%
pH 1.5
08.C
WLO
SS'C
I
OH 3.0
30%
Figure 4. Effect of pH upon Size and Perfection of Barium Sulfate
I n an earlier report from this laboratory ( d ) , coprecipitanta were shown to cause surface protuberances upon crystals formed very slowly by diffusion mixing of the reagents. Although some similar bumps were seen on the rapidly precipitated barium sulfate, the general observation was not one of decreased perfection. Therefore, the action of coprecipitants must be somewhat dependent upon the method of mixing the reagent solutions. LaMer reported that the addition of a neutral salt such as potassium nitrate decreases the supersaturation ratio (6). Therefore, such electrolytes would be expected to increase the crystal size. However, all observed effects of foreign ions upon crystal size were in the direction of decreasing the size. Apparently the solid solution development and/or adsorption during crystal development retard the growth along one or more axes.
1547 Effect of Aging Barium Chloride Solution Prior to Use. An unexpected aging effect was noted in the present study. Particles of barium sulfate precipitated from a fresh barium chloride solution were decidedly smaller than those formed from a barium chloride solution which had stood a t room temperature for some time. The crystal size differential was severalfold. As this phenomenon has apparently not been reported in the literature, it was investigated in some detail in this laboratory. [Subsequent to the original submission of this paper the author's attention was drawn to a similar report by Bogan ( I ) . ] As yet it is not possible to propose a valid explanation of this very perplexing, but very real, aging effect. Therefore, no details of the investigation need be given here, although the phenomenon is discussed briefly in a different connection in another report from this laboratory
(4). Quantitative Tests. From the results of the microscopic tests described in the preceding paragraphs, the procedure listed as Method B was devised for complete, rapid determinations of sulfate as barium sulfate. The value or range of values for each condition was selected to yield crystals that were large and reasonably well perfected, compromising between these two desiderata where necessary. Large crystals are desirable for several reasons: low specific surface area for adsorption of impurities, ready filtration through common filter media, and ease of washing. Well-perfected crystals are desirable for rapid settling (irregular ones are relatively buoyant), prevention of agglomeration which would enhance the occlusion of mother liquid, low specific surface area for adsorption of impurities, and easy and rapid washing. Many quantitative determinations were made upon samples containing knoan amounts of sulfate by Method B. I n each determination the volume of sulfate solution was about 100 ml., and the corresponding weight of barium sulfate was very close to 0.3000 gram. Although only a 10 to ZOY0 excess of barium chloride was used a t the time of the initial precipitation, an amount sufficient to bring the excess up to about 40% was added just prior to filtration. Use of this additional barium chloride may not be necessary in most cases; however, it is a desirable safeguard in the analysis of samples of which even the approximate sulfate content is not known. Complete rinsing of the beaker in which the precipitation was conducted appeared to be the only washing necessary. A water aspirator was used in the filtrations. Porous porcelain crucibles were used for the filtrations, and each was placed in a common glazed porcelain crucible for the ignition over a burner. Typical quantitative data are included in Table I. All determinations were made by Method B except as indicated in Table I. Determinations upon potassium sulfate a t a pH of 1 are low by 8 parts per thousand, while similar determinations upon sodium sulfate are precise and accurate within 1 part per thousand (lines 1 and 9). Ionic radii are 1.41 A. for potassium, 1.35 A. for barium, and 1.07 A. for sodium. Therefore potassium could coprecipitate with barium sulfate in the form of a substitutional solid solution much more readily than sodium, as discussed above. I n order to preserve electrical neutrality, either two potassium ions must replace one barium ion, which is highly improbable, or else the substitution of one potassium ion for one barium ion must be accompanied by substitution of one bisulfate ion for one sulfate ion. That the latter is the case is proved by the data of lines 5 and 6 ; a t a p H value of 5 or 6, the bisulfate ion concentration is very small, and the coprecipitation of potassium is of no consequence. Even the coprecipitation resulting a t pH 1 from a hundredfold excess of potassium chloride (line 7 ) is effectively eliminated a t p H 5 (line 8). Apparently the bisulfate ion concentrations a t pH values of 2 and 4 are not low enough to eliminate the error, but a p H of 5 or 6 is satisfactory (lines 3 to 6). Although the sodium coprecipitation error is much less than that of potassium, it may become significant a t very high sodium ion concentrations (lines 11, 12, and 13), but the error is eliminated at a p H
ANALYTICAL CHEMISTRY
1548
of 5 or 6 (lines 14 and 15). I n the presence of high concentrations of the alkali ions, the pH a t the time of precipitation should he made 5 instead of 1 as otherwise recommended. The crystals are admittedly less perfect and settle less readily at the higher pH, but the compromise seems both necessary and acceptable. The potassium error is a real one, such as solid solution formation, and the error at pH l is not merely one of solubility of harium sulfate (lines 1 and 2). Coprecipitations of nitrate and ferric ion are fully as serious in the rapid precipitation method as in the noImal slow precipitation methods (lines 16 and 17). Apparently nitrate ion, if present in high concentration, must be removed-e.g., by evnporation-prior to precipitation of barium sulfate if the results are to be accurate. However, as the precision is good (line 16), perhaps an empirical correction factor could be applied in the calculations The interference from ferric ion in reasonable concentrations, however, can be eliminated with ethylenediaminetetraacetir acid (lines 19 and 20 compared to line 17). Ethylenediaminetetraacetic acid by itself causes no interference (line 18), unless it is present in considerable excess; use of the commercially available ethylenediaminetetraacetic acid reagent at a rate of approximately 200 ml. of the commercial preparation per mole of ferric ion is recommended.
CONCLUSIONS
On the basis of many quantitative tests, data on some of which are presented in Table I, it appears that the rapid precipitation procedure of Method B possesses no real diszdvantage in comparison to the more conventional procedures, yet it offers greater speed and ease of handling, filtration, and washing. Results obtained with this rapid precipitation method are both precise and accurate within 1 part per thousand, even in the presence of high concentrations of alkali ions. LITERATURE CITED (1) Bogan, E. *J., thesis, Ohio State University, 1949. (2) Fischer, R. B . , . ~ A LCHEX., . 23,1667 (1951). (3) Fischer, R. B., J . Chem. Educ., 24,484 (1947). (4) Fischer, R. B., and Rhinehammer, T. B., Ax.4~.CHEM., t o be pub-
lished. ( 5 ) Hints, E., and Weber, H.. 2. cct~al.Chem., 45,31 (1906). (6) LaMer, V. K., and Dinegar, R. H., .I. A m . Chpm. SOC.,73, 380
.
(1951) (7) Popoff, S.,and Neuman, E. W., ISD. EXG.CHEM.,ANAL.ED., 2 , 4 5 (1930). (8) Rieman, W., and Hagen, G., Ibid., 14, 150 (1942).
RECEIVED for review March 27, 1953. Accepted July 27, 1953. Publication N o . 586, Chemistry Department of Indiana Cnirersity.
Estimation of Cholesterol and Triterpenols in Unsaponifiable Fraction of Wool Wax HEINZ DUEWELL Division of Industrial Cheniistry, C o m m o n w e a l t h Scienti3c and Industrial Research Organiaation, Melbourne, Australia VETHOD
is described for the simultaneous determination of
A-cholesterol and the triterpenols in the unsaponifiable portion
of wool wax, which makes use of the color devcloped with thcs Liebermann-Burchard rcagent (acetic anhydride and concentrated sulfuric acid). This is a modification of Lederer and Tchen’s method for the estimation of triterpenols, but by using a spectrophotometer and standard solutions of cholesterol and triterpenols the necessity for temperature control IP avoided and the results are accurate within f3yO. The investigation of wool was fractions in progress in these laboratories demanded a rapid and relatively accurate method for the determination of the proportions of the chief groups of wax alcohols present. namely cholesterol. triterprnols. and aliphatic alcohols. Cholesterol and thc triterprnols give a bluishgreen and yellowish-green color, respectively, when treated with the Liebermann-Burehard reagent. Schoenheinier and Sperry (6) based a method for the determination of blood cholrsterol on this color reaction by measuring the absorbance at 620 mp after a given time. Lederer and Tchen ( 4 ) found that the triterpenols showed a narrow band with a maximum a t 458 mp, whereas cholcsterol produced a very broad band with a maximum a t 620 mp. The intensity of absorption of triterpenols a t 460 m p differs little from the absorption a t 458 mp and was about ten times that of cholesterol a t 620 mp, Lederer and Tchen realized that this difference in absorption spectra would allow the estimation of both cholesterol and triterpenols, but directed their main interest to the determination of the latter. They showed that no great difference existed between the absorbance produced by the individual wool wax triterpenols, so that they could all be measured together. Delsal(1) in a paper on the hydrolysis and estimation of cholesteryl fatty esters found that under some conditions higher absorbances were produced by the esters in solutions of equivalent cholesterol content than by cholesterol itself. However, by decreasing the amount of sulfuric acid in the Liebermann-Burchard reagent from lOyo (v./v.) to 5%, he was able to get identical absorbances for free cholesterol and its esters after 30 minutes Lederer and Tchen used the method to determine the percentage
of triterpenols in n 001 wax. They gave no account of any detailed investigation of the method and when used here, it was not found possible to determine triterpenols in the unsaponified wax although reliable results could be obtained after saponification. Thus direct measurements on a crude way sample gave 11.9% of triterpenols, whereas measurement on the unsaponifiables shoived only ‘3.7% of triterpenols present in the original wax. The present paper describes a modification of Lederer and Tchen’s procedure ( 4 ) incorporating several improvements of particular importance and adapting it to allow the simultaneous determination of cholesterol and of triterpenols. It was found advisable to use standard cholesterol and triterpenol solutions for each run of measurements, rather than rely on calibration curves which apply strictly a t only one temperature and require very careful control of the composition of the Liebermann-Burchard reagent. The mixture of acetic anhydride and concentrated sulfuric acid was found to be unstable and was not used after 10 hours. Evaporation to dryness as used by Lederer and Tchen (at room temperature to lessen oxidation of triterpenols) is avoided in the present procedure, for equal volumes of solution were used in the samples and the standards. The accuracy of the results was increased by using a spectrophotometer in place of a photometer fitted with filters. As was shown by Lederer and Tchen, the absorbance measured a t 460 mp must be corrected for the contribution of eholesterol before calculating the amount of triterpenol present. This contribution is of the order of 55% of the absorbance measured a t 620 mp. Lederer and Tchen considered that for the estimation of triterpenols alone, the absorption a t 620 mp could be supposed to result entirely from the presence of cholesterol. This supposition leads only to an error of about 0.5% in the triterpene alcohol estimation and being within the limits of accuracy of the method, is justified. But if this simplification is used when cholesterol is also determined, then large errors are introduced as can be seen from a comparison of columns 7 and 10 of Table I. REAGENTS AND EQUIPMENT
Cholesterol and Triterpenol Solutions. For the preparation of solutions, alcohol-free chloroform, stored over sodium hydrox-
