Silica sol-treated beads can demonstrate adsorptive effects (IO). When coated with Of ’Ow Corning 200 silicone oil, sol-modified beads adsorb polar compounds such as dibutyl ether and methyl benzoate, causing tailing peaks. Nonpolar compounds, such as decane, elute as symmetrical peaks. With a 0.85% silicone oil loading, peak symmetries for all compounds in the test mixture are identical to those obtained with unmodified glass beads. Use of polar stationary liquid phases in smaller concentrations also results in the production of symmetrical peaks for polar compounds.
LITERATURE CITED (1) Bechtold, M. F., Snyder, 0. E. (to E. I. du Pont de Nemours & GO.), u. s. Patent 2,574,902 (Nov. 13, 1951). (2) Bohemen, J., Purnell, J. H., J . Chem. SOC.1961, 360.
matography 1960, Edinburgh,” R. P. W. Scott, ed., p. 139, Butterworths, London, 1960. (7) HalAsz, J., H o d t h , C., ANAL.CHEM. 35, 499 (1963). (8) Ibid., 36, 1178 (1964). (9) HalAsz, J., HorvAth, C., Nature 197, 171 (1963). (10) Kirkland, J. J., 5th International Symposium of Gas Chromatography, Brighton, England, September 1964. (11) Norem, s. D., ANAL. CHEM. 34, 40 (1962). w., Jojola, R.9 1bid.j 36, (12) Ohline, 1681 (1964).
(3) Dewar, R. A., Maier, C. E., J. Chromatog. 11, 296 (1963). (4) Giddings, J. c., ANAL. CHEM. 34, 458 (1962) ( 5 ) Ibid., 35,‘ 438 (1963). (6) Golay, M. J. E., “Gas Chro-
RECEIVEDfor review March 3, 1965. Accepted June 17, 1965. Division of Analytical Chemistry, Lab-Line Award Symposium, 149th Meeting, ACS, Detroit, Mich., April 1965.
ACKNOWLEDGMENT
I am indebted to Ralph K. Iler for his helpful suggestions and to Glenn J, Wallace for his assistance with the experimental work,
Hexa mmineco balt(lll) Tricarbona toco balta te(lll) A N e w Analytical Titrant
-
JAMES A. BAUR and CLARK E. BRICKER Department o f Chemistry, University o f Kansas, Lawrence, Kan.
b A new reagent for titrimetric oxidations is described and evaluated. is easy This reagent, CO(NH~)BCO(CO&, to prepare and, in the dry state, appears to be stable indefinitely. This reagent is a very weak oxidant in bicarbonate media but when added to an acid solution, the cobalt(lll) from the tricarbonato portion of the compound is released. The cobalt(ll1) thus generated is a very strong oxidant and is capable of reacting quantitatively with iron(ll), vanadium (IV), cerium(lll), and other reducing agents. The equivalence point of these titrations can be determined by redox indicators, by potentiometric methods, or by photometric detection.
T
POTENTIAL OF hydrated cobalt(II1) when it is reduced to cobalt(I1) is well known to be about 1.8 volts us. the standard hydrogen electrode ( 8 ) . Because cobalt(II1) reacts spontaneously with water to liberate oxygen, it is extremely difficult to keep solutions of this strong oxidant. Bricker and Loeffler (2) found that cobalt(II1) sulfate was stable in 18N sulfuric acid if the solution was stored at temperatures below 0’ C. With standardized solutions of cobalt(II1) sulfate, it was possible to oxidize quantitatively such weak reducing agents as cerium(II1). If a complexed cobalt(II1) salt could be prepared which would give stable solutions in neutral or alkaline media, it may be possible, on acidification of the solution, to generate the powerful oxidant in situ. Most complex cobalt (111) salts are either too stable to
HE HIGH
liberate the free cobalt(II1) ion on acidification or else the ligand or a moiety of the compound is a reducing agent. One exception is the tricarbonatocobalt(II1) anion which should liberate carbon dioxide and the hydrated cobalt(II1) cation upon acidification. Job (4) prepared a green compound from potassium carbonate and cobalt(II1) to which he assigned the formula, KaCoOa. Blanchetiere and Pirlot (1) used the green complex of cobalt(II1) and carbonate as an indirect colorimetric method for the determination of potassium after precipitating NaK2Co(NOz)6. I n a similar way, Willard and Ayres (9) used this same complex for the colorimetric determination of cobalt, and Laitinen and Burdett (6) developed an iodometric method for the determination of cobalt based on this green complex. McCutcheon and Schuele (6) prepared the compound Co(N&)&o(COs)3 and measured its solubility in water as well as some other physical properties. Mori, et al. ( 7 ) claim to have precipitated K&o(CO&. There is, however, no apparent reference in the literature where tricarbonatocobalt(II1) has been used as a titrant. This work describes the preparation of solid hexamminecobalt(II1) tricarbonatocobaltate(II1) and the use of this reagent as an analytical titrant. EXPERIMENTAL
Preparation of Co (NH3)&o( C o d 3 . T o 50 ml. of 1M cobalt(I1) chloride, add 100 ml. of water and sufficient
sodium bicarbonate to saturate the solution. Cool the solution in ice water and then, with constant stirring, add dropwise 10 ml. of 30y0 hydrogen peroxide. Continue to stir the solution for 5 to 10 minutes to allow any excess hydrogen peroxide to decompose. Filter the resulting mixture through a fritted glass funnel. T o the clear, dark green filtrate add solid hexamminecobalt (111) .chloride until the supernatant liquid develops a n orange tinge. The resulting solution is again filtered through a fritted glass crucible and the solid material is washed repeatedly with water until the washings are colorless. The solid is then kept in a vacuum desiccator over magnesium perchlorate until dry. The dry solid is then stored a t room temperature in a screw cap bottle. The hexamminecobalt (111) chloride is prepared essentially by the method described in reference (3). This method consists of mixing 73 grams of cobalt(I1) nitrate hexahydrate in 100 ml. of water with 80 grams of ammonium nitrate, 2 grams of activated charcoal, and 180 ml. of concentrated ammonia. The resulting solution is oxidized with hydrogen peroxide or by bubbling air through the solution until the brown precipitate ceases to form. The precipitate is filtered through paper, washed with water, and then dissolved in 1300 ml. of hot water containing sufficient hydrochloric acid to give an acid reaction. The hot solution is then filtered to remove the charcoal, and 400 ml. of concentrated hydrochloric acid are added to the filtrate. The solution is allowed to cool and the solid Co(NH3)&13 is removed by filtration, washed successively with 60% and 95% ethanol and then dried in a vacuum desiccator. VOL. 37, NO. 12, NOVEMBER 1965
1461
(2)
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Stability of solutions of Co("3)&o(CO&
140
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( 1 ) Solution exposed to light (2) Solution stored in dork
night. Similarly, if a carbonate-bicarPreparation and Stability of Solutions of C O ( ~ H ~ ) ~ C O ( Because C ~ ~ ) ~ . bonate buffer of p H 8.6 is used to dissolve the salt, the solution decomCO("3)&O(C03)3 is soluble t o the poses in less than 24 hours. extent of only 0.038 gram per 100 Solutions of Co(~H3)6Co(C03)3suitgrams of water ( 7 ) and, these aqueous able for analytical study are prepared solutions are quite unstable, a medium by adding approximately 3 grams of that would provide greater solubility the salt to a liter of water that has and increased stability was sought. previously been saturated with sodium Bicarbonate solutions of C0("3)6bicarbonate. After the solution has C O ( C O ~were ) ~ found to yield, by far, been stirred for 2 to 3 hours, it is the most stable solutions and the solufiltered through a fritted glass crucible bility of the salt in saturated sodium and the resulting green solution, which bicarbonate is approximately 10 times 5 X 10-~Jf,is ready that in water. If C O ( ~ H ~ ) ~ C O ( C Ois~approximately )~ for standardization. is dissolved in any medium that has a The stability of two bicarbonate p H less than 7, the tricarbonato comsolutions of CO(KH3) 6Co(c03)3 have plex decomposes and liberates free been studied over a perrod of 3 to 4 cobalt(II1). On the other hand, in months. One of these solutions was solutions more alkaline than about p H 8, kept in a clear, glass-stoppered bottle the green color disappears quite rapidly exposed to the fluorescent lights in the and a brown precipitate of cobalt(II1) laboratory, whereas the other solution hydroxide forms. The stability of C O ( ? ; H ~ ) ~ C O ( C ~ was ~ ) ~ stored in the dark. The initial concentration of the solution stored in in bicarbonate solution (DH 7.6) can the dark was approximately one half be explained from a cokideration of simple equilibria. It is known that the hexamminecobalt(II1) ion is quite stable and its rate of dissociation even in strong acid is very slow. On the other hand, the tricarbonatocobalt(II1) .60ion is in equilibrium with its component ions.
that of the solution exposed to light. The oxidizing normality of each of these solutions decreases nearly linearly with time. The data for these stability studies are shown in Figure 1. Standardization of Co(NH3)&o(co3)3 Solutions. Ferrous ethylenediammonium sulfate tetrahydrate was used as the primary standard but because the solutions t o be standardized were so dilute, a solution of known concentration of the primary standard was prepared. These solutions of the primary standard were 3 to 5 x 10-3M in iron(I1) and 1M in sulfuric acid. An accurately measured aliquot of the standard iron(I1) solution is diluted to approximately 125 ml. with 1.8M sulfuric acid. After two drops of ferroin indicator are added, the solution is stirred with a magnetic stirrer while being titrated with the cobalt(II1) solution until the color of the indicator changes sharply. The average deviation of triplicate standardizations from the average value should be within 2 parts per thousand. The volumes of the cobalt(II1) solution used in the standardizations must be corrected for the amount of oxidant used to titrate the indicator. For 2 drops of ferroin indicator, approximately 0.60 ml. of 5 X l O - 3 X cobalt(II1) solution is required. Determination of Cerium(II1). Because there are no indicators t h a t are suitable to detect the equivalence point in the titration of cerium(II1) with cobalt(III), photometric methods are the most convenient. Although it may be possible t o detect photometrically when a n excess of cobalt (111) has been added, it is more convenient and reliable to follow the formation of the cerium(1V) during these titrations. Cerium(1V) in sulfuric acid medium shows an increasing absorbance between 400 mp and 320 mp. I n order to follow the formation of cerium(1V) during a titration and not have the absorbance readings become too large, a wavelength not too sensitive for the detection of the cerium(1V)
r
cO(c0~)3-~ f 6HzO = C O ( H * O ) ~ + ~3C03-2
+
If acid is added to this system, the carbonate ion is destroyed. If base is added, cobalt(II1) hydroxide precipitates. The optimum stability is reached with a solution of bicarbonate which provides a sufficient carbonate ion concentration to suppress this equilibrium and a t the same time insufficient hydroxide ion concentration to precipitate the cobalt(II1) hydroxide. This explanation seems to be substantiated by the fact that a solution of C O ( N H ~ ) ~ cO(c03)3 prepared in a borate buffer of p H 7.7 decomposed on standing over1462
ANALYTICAL CHEMISTRY
w.40V
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Co(lll)
Spectrophotometric titration of cerium(ll1) with cobalt
(111) Wavelength = 400 rnp
is chosen. For this reason, a wavelength of 400 mp is used in these titrations. A11 measurements were made with a Bausch and Lomb Spectronic 20. The solution containing the cerium (111) is diluted to approximately 50 ml. with 0.5M sulfuric acid. Measured volumes of the cobalt(II1) solution are added while the analyte is being stirred vigorously with a magnetic stirrer. After each addition of titrant, the absorbance of the solution a t 400 mp is measured. The titration is continued until 3 or 4 consecutive absorbance readings show little or no change. The observed absorbance values are corrected for dilution by the titrant and then these corrected values are plotted us. the volume of cobalt (111) added. Such a graph is shown in Figure 2. Determination of Vanadium(1V). The procedure for the determination of vanadiuni(1T’) Lvith cobalt(II1) is very similar to t h a t described for cerium(II1). The only differences are that stronger sulfuric acid is used and a wavelength of 750 mp is used t o follow the decrease in the amount of vanadium(1V). The solution containing the vanadium(1V) wlfate is made 511f in sulfuric acid and then titrated with measured volumes of the cobalt(II1) solution. The titration is followed photometrically a t 750 mp and the equivalence point is detected in the same manner aq described for cerium (111). See Figure 3 for the results of one of thew titrations. RESULTS AND DISCUSSION
The precision of the cerium(II1) titrations vias shown by titrating five 1.00-ml. aliquots of a cerium(II1) sulfate solution. The average volume of cobalt(II1) solution required for each of these titrations was 7.75 =I= 0.03 ml. The precision for this titration is, therefore, ~ ! ~ 0 . 4 7 ~ . When eight 20.00-inl. aliquots of a vanadium(1V) sulfate solution were titrated with cobalt(II1) solution, an
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Figure 4. Potentiometric titration of iron(l1) with coba It(I1I)
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average of 23.7 f 0.4 ml. of titrant was used. The normality of the vanadium solution was, therefore, 3.80 x 10-3 with a precision of only &2%. When the same vanadium(1V) solution was titrated with standardized cerium(IV), the normality of the vanadium(1V) was found to be 3.78 X 10-3 and the precision in these titrations was comparable to that found with the cobalt(II1). Obviously, the agreement between the cobalt(II1) and cerium(1V) titrations of vanadium (IV) is good. The chief reason for the comparatively poor precision in both of these sets of titrations is the photometric end point which is not too satisfactory because of the low absorbance of vanadium(1V) (see Figure 3). I n addition to the spectrophotometric and indicator methods for determining the equivalence point of cobalt(II1) titrations, potentiometric methods were tried. A platinum wire
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Co(II1)
Spectrophotometric titration of Wavelength = 750 m p
vanadium(1V) with
6 ML.
8 of
IO 12 Co(ll1)
14
16
18
electrode served as the indicator electrode and a mercury, mercurous sulfate, saturated potassium sulfate electrode was the reference. All potential measurements were made with a Heathkit Model 1M-11 vacuum tube voltmeter. A typical titration curve for the titration of iron(I1) with cobalt(II1) is shown in Figure 4. The normality of the cobalt(II1) solution, calculated from four iron(I1) titrations followed potentiometrically, f 3 parts per was 2.812 X thousand. The normality of the same cobalt(II1) solution, calculated from similar titrations followed with ferroin indicator, was 2.801 X f 2 p.p.t. Similarly, titrations of vanadium(1V) with cobalt(II1) can be followed potentiometrically, but in these titrations the “break” in the titration curve is only about 200 mv. as compared to nearly 500 mv. shown in Figure 4 for the iron(I1) titration. The precision of the vanadium(1V) titrations, using a potentiometric end point, is somewhat better (Zk1.3 %) than with the spectrophotometric deHowever, if the tection (Zk2%). sample of vanadium(1V) containing 0.05 to 0.1 meq. was diluted to 100 or 150 ml. before the titration, very poor potentiometric titration curves were obtained. Because cobalt(II1) is such a strong oxidant, the possibility of this titrant reacting with water as well as with the reducing agent t o be determined always exists. For this reason, the rate of addition of titrant and the temperature of the solution during the titration are rather critical. For example, when a 25-ml. aliquot of an iron(I1) solution was titrated in about 30 seconds a t room temperature with approximately 17 ml. of cobalt(II1) solution, the normality of the cobalt(II1) was found to be 7.025 X =k 2 p.p.t. If this same titration was carried out slowly VOL. 37, NO. 12, NOVEMBER 1965
1463
at room temperature so that the titrant was added over a period of 13 to 14 minutes, the normality of the cobalt(II1) was 7.100 X lop3 f 2 parts per thousand. The 1% difference in these two sets of titrations must be attributed to the increase of the reaction of cobalt(II1) with water in the rapid titrations because of the high local concentration of titrant. If, on the other hand, this titration was performed rapidly a t 0’ C., a larger volume of titrant than used at room temperature was required, and the normality of the cobalt(II1) was calculated to be 6.798 x These results a t 0’ C. suggest that even more cobalt(II1)
must be consumed in oxidizing water than in those titrations carried out at room temperature. This explanation would appear to be valid only if the rate of the reaction between cobalt(II1) and \Yater decreases less rapidly with temperature than does the reaction between iron(I1) and cobalt(II1). A study of the rates of various oxidationreduction reactions as a function of temperature is now under investigation. LITERATURE CITED
(1) Blanchetiere, A,, Pirlot, J. M., Compt. Rend. Soc. Biol. 101, 858 (1929).
(2) Bricker, C. E., Loeffler, L. J., ANAL. CHEM.27, 1419 (1955).
(3) “Inorganic Syntheses,” vel. 11, P. 216, 1C.lcGraw-HilL New Y& (1946). (4) Job, A,, Ann. Chim. Phys. 20, 205 (1900). ( 5 ) Laitinen, H. A., Burdett, L. W., ANAL. CHEW 23, 1265 (1951). (6) RfcCutcheon, T. P.1 SchUele, W. J., J . Am. Chem. SOC.75, 1845 (1953). ( 7 ) hlori, RI., Shibata, M., Kyuno, E., Adachi, T., Bull. Chem. SOC. (Japan) 29, 883 (1956). ( 8 ) Noyes, A. A.1 Deahl, T. J.! J * Am. Chem. SOC.59, 1337 (1937). (9) Willard, H. H,, G. H., A ~ CHEM.12, 287 (1940).
RECEIVEDfor review August 2, 1965. Accepted September 21, 1965. Division of Analytical Chemistry, Lab-Line award Symposium, 149th Meeting, ACS, Detroit, Mich., April 1965.
Study of Organic Structure via Mercury-Sensitized Photolysis and Gas Chromatography Alcohols and Esters RICHARD S. JUVET, JR., and LUTHER
P. TURNER’
Department of Chemisfry and Chemical Engineering, University o f Illinois, Urbana, 111.
The controlled, mercury-photosensitized, liquid phase photolysis of some simple aliphatic alcohols and straight chain esters has been studied and the products have been analyzed by gasliquid chromatography. Each of the compounds studied yields a characteristic degradation pattern suitable for “fingerprint” identification of special structural features within the molecule. Within any of the series studied, the major mechanism of degradation is independent of chain length, which enables the positive identification of the irradiated structure through the formation of homologous or common products. Numerical constants involving the retention of these photolytic decomposition products using a linear programmed temperature gas chromatographic column have been tabulated, and these constants are characteristic of chain branching and the functional groups present in the irradiated material. The expression of retention data as retention indices allows the reproduction of these data in other laboratories.
T
he combination of pyrolysis and gas-liquid chromatography (GLC) has often been recommended for qualitative identification of materials as well as for extending GLC to the analysis of compounds of low volatility. Perry (12) has given a thorough review of the tech1464
0
ANALYTICAL CHEMISTRY
niques and potentialities of pyrolysisGLC. Unfortunately, the nature and relative amounts of pyrolysis products vary greatly with the experimental conditions and apparatus used ( 2 ) . The maximum temperature of the filament or pyrolysis chamber, the rate of heating, the sample size, the prior history of the filament or pyrolysis chamber and its composition all appear to influence the results. Thus, data obtained are difficult to reproduce from laboratory to laboratory. The effect of pyrolysis temperature on the amount and identity of the products formed is particularly striking and has been studied by several workers (2, 4, 11). Janak (3) found that rearrangements often occur in the pyrolysis of barbituric acid derivatives, leading to unexpected products. Keulemans and Perry (4) showed that rearrangement and secondary reactions are dependent upon sample size as well as pyrolysis temperature. Most of the recent work in this area has been concerned primarily with the development of experimental techniques for improving the reproducibility of pyrolysis. Pyrolysis-GLC has remained predominantly a technique in which identification depends upon the possession of previously obtained pyrolysis patterns with which the pattern of interest may be compared through careful control of all experi-
67803 mental parameters. Very little progress has been made concerning the basic modes of degradation which would allow this technique to be used as a means of qualitative structural determination for compounds previously unstudied. It is the purpose of this work to investigate the use of high-intensity monochromatic mercury resonance radiation for the controlled degradation of organic compounds, with subsequent analysis of the products by GLC. Ultraviolet radiation appears to have certain advantages which enhance its use as a degradative means for identification and structural determinations, The apparatus used in photolytic studies is relatively inexpensive and can be easily standardized. The conditions of degradation may be carefully controlled and accurately known. The product yields may be maintained a t very low values (< 2%), thus minimizing secondary reactions and changes in matrix. The number of products expected from degradations by mercury resonance radiation is much less than in the case of techniques such as pyrolysis in which the samples are subjected to more severe energetic conditions. The breakdown patterns are thus more simple and easy to interpret. I n order for a compound to undergo photolytic reaction it must absorb the 1 Present address, Department of Chemistry, University of Tennessee, Knoxville, Tenn.
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