In general, it is noted that separation is better for the amides in Table I1 in contrast to the esters in Table I. This result is probably due to the hydrogen bonding capacity of the amide hydrogen along with the rigidity imparted by the partial double bond character of the C-N bond in the amide linkage (IO). It is also probable that the use of different acids in Tables I and I1 contributes to separation differences for the esters and amides in this case. However, recent results in which the groups attached t o both asymmetric centers remain constant does reveal that diastereoisomeric amides separate better than esters (11). It is next of interest to compare the aromatic compounds 15, 19, and 21 with their corresponding alkyl derivatives 16, 20, and 22, respectively. While the retention times of the phenyl and cyclohexyl derivatives are approximately the same on the nonpolar QF-1 phase, substantial increases in retention occur for the aromatics relative to the alkyls o n the polar EGA phase, indicative of a solute-solvent interaction of the aromatic ring with the EGA ester phase. Comparison of the 01 and A(AGo) values for compounds 15 with 17 and 19 with 21 o n the polar EGA phase further indicates that separation is better the closer the aromatic ring is to the amine asymmetric center.
While much more work remains t o be done for an understanding of the separation of diastereoisomers, several conclusions may be drawn from the work of this paper and the previous ones ( I , 6, 7). In general, a polar phase, such as a polyester, produces better separation for diastereoisomeric esters and amides than a nonpolar phase. The resolving agent should be a readily available, low molecular weight optically pure compound containing a suitable functional group close to the asymmetric center. This last requirement results from the necessity of having the two asymmetric centers close t o one another in the diastereoisomeric molecule. If acyclic resolving agents are used, the three groups attached to the asymmetric center (along with the functional group) should have a large size differential. Alternatively, cyclic compounds with a functional group adjacent t o the asymmetric center, such as proline (7), also serve as excellent resolving agents. In general, the more rigid the diastereoisomeric molecule close t o the asymmetric centers, the larger will be the separation. Finally, if polar groups (along with the functional group) are attached t o the asymmetric center, it is difficult t o predict the effect o n separation.
(10) J. R. Dyer, “Applications of Absorption Spectroscopy of Organic Compounds,” Prentice-Hall, Englewood Cliffs, N. J., 1965, p 113. ( 1 1 ) B. L. Karger, R. L. Stern, S . Herliczek, unpublished results,
RECEIVED for review June 10, 1968. Accepted July 29, 1968. Work supported by National Aeronautics and Space Administration under Grant NsG 81 and National Science Foundation under Grant No. 8572.
Northeastern University, June 1968.
Spectrophotometric Determination of Some Organic Acids with Ferric 5-NitrosaIicylate Complex Kil Sang Lee and Dai Woon Lee Department of C/iemistry, Yonsei Unicersity, Seoul, Korea
Jae Young Hwang Chemical Research Laboratories, AGTL Inc., Natick, Mass.
A SPECTROPHOTOMETRIC STUDY of the colored complex produced by the interaction of a solution of ferric salt with salicylic acid and its derivatives, such as 5-chloro, 5-bromo-, 5-nitro-, and 3-nitrosalicylic acid, has been extensively made by many authors. The spectrophotometric determination of the stability constants of some ferric salicylates was reported by Ernst and Menashi ( I , 2). They also found that the complexes were formed by 1 mole of ferric ion and 1 mole of each reagent. The authors have found that the reddish-orange complex which is obtained by the interaction of Fe3+ion with 5-nitrosalicylic acid within the p H range of 2.5 to 3.0, can be used for the determination of iron within a concentration range of 5 t o 30 ppm. Molar absorptivity of ferric 5-nitrosalicylate was previously found to be 2253 mole-’ cm-ll-l at 492 mp (3). The absorption of the ferric 5-nitrosalicylate in a n aqueous or acetate medium at 492 mp is diminished considerably by the addition of small quantities of some water-soluble organic acids. This phenomenon is employed as the basis of a simple (1) 2. L. Ernst and J. Menashi, Trans. Faraday SOC.,59, 2838
spectrophotometric technique for the determination of organic acids. A properly diluted portion of the organic acid is added to a known quantity of the ferric 5-nitrosalicylate. The absorption is measured at 492 m p and compared to a previously prepared plot of absorbance cs. concentration of the organic acid. Kovalenko and Petrashen (4) determined some organic acids, such as citric, tartaric, and oxalic acid, by a spectrophotometric method with chromium (VI)-diphenylcarbazide. A method is described for oxidimetric determination of tartaric acid with potassium cupri-3-periodate in alkaline medium (5). This method could be used only for determination of tartaric acid in the presence of other plant acids after separation of tartaric acid by using ion exchanger Dowex 1-X 10. All methods which use permanganate in alkaline or acid medium have the disadvantage of spontaneous decomposition of permanganate a t elevated temperatures and after a prolonged reaction time. The present study, when compared with the methods mentioned above, is much less tedious, far simpler in the operation, and less time-consuming than the other methods.
(19631.
p 1794. (3) D. W. Lee, M. S. Thesis, Yonsei University, Seoul, Korea, (2jI&i,
1966.
~
(4) E. V. Kovalenko and V. I. Petrashen, Tr. Nooocherk, Politekhn, Znsf., 141, 59 (1964). (5) N. Velikonja, Arhiu Za Kemi/rr, 27, 161 (1955). VOL. 40, NO. 13, NOVEMBER 1968
0
2049
I
I
I
I
Iq0l
0.9
or 2.8 with the buffer solution and the contents of the flask were diluted to the mark. The final concentrations of unknown organic acid should not exceed its optimum range of the concentration in the calibration curves. All the solutions in the above series were examined with a spectrophotometer at 492 mp. PREPARATION OF CALIBRATION CURVES. The values of absorbances were plotted GS. micrograms of organic acids in above series. A straight-line of one organic acid differs from another in its slope and length. RESULTS AND DISCUSSION
0.4,
I
I
2.0
3.0
I
40
PH
Figure 1. Effect of pH (ferric 5-nitrosalicylate, 3.76 X 10-4M at 492 mp)
EXPERIMENTAL
Apparatus. All absorbance measurements were made with a Shimadzu Seisakusho Type QR-50 spectrophotometer with I-cm cells. All p H measurements were made with a Hitachi Horiba H-5 pH meter which was standardized at pH 7 and pH 4 with Hitachi Horiba standard pH buffers. Reagents. STANDARD IRONSOLUTION.The standard trivalent iron solution was prepared by dissolving weighed amounts of spectrometric Baker Analyzed Reagent iron wire in a minimum volume of 1 :1 nitric acid. The solution was heated to boiling on a hot plate to remove the oxides of nitrogen. Water was added periodically to keep the solution from evaporating to dryness. The iron solution thus prepared was transferred to a 1-liter volumetric flask and diluted to the mark with demineralized water. 5-NITROSALICYLIC ACID SOLUTION (1 ml O f the SOlUtiOn contains 2 mg of the acid). 0.2 gram of 5-nitrosalicylic acid (Eastman Organic Chemicals No. 1338) was dissolved in 100 ml of 307, methanol-HeO solution. CHLOROACETATE BUFFERSOLUTION.4.7 grams of monochloroacetic acid (Eastman Organic Chemicals No. 68) weredissolved in 500 ml of demineralized water and the pH was adjusted with 1N sodium hydroxide or 1 N hydrochloric acid. ORGANIC ACIDSSOLUTIONS (1 ml of the solution contains 1 mg of the organic acid). 0.1 gram of organic acid (analytical reagent quality) was dissolved in 100 ml demineralized water. The solution was transferred to a 300-ml polyethylene bottle and stored in a refrigerator. PREPARATION OF FERRIC 5-NITROSALICYLATE COMPLEX. 40 ml of 1.791 x 10-*M standard iron solution and 100 ml of 0.2% 5-nitrosalicylic acid were pipetted into a 1-liter volumetric flask, The pH was adjusted to 2.6 with the buffer solution and the contents of the volumetric flask were diluted to the mark with buffer solution and demineralized water. Procedure. With a microburet or a regular pipet, 5.0 ml of reddish orange ferric 5-nitrosalicylate solution were placed into a series of 10-ml volumetric flasks. Into these flasks, 0.1-ml increments of the standard organic acid solution were entered starting with 0.0 ml and ending with 2.0 ml. For each unknown organic acid solution to be determined, a 10-ml volumetric flask was prepared as above, by adding 5.0 ml of ferric 5-nitrosalicylate solution and a selected amount of the unknown organic acid. The pH was adjusted to 2.6 2050
0
ANALYTICAL CHEMISTRY
Effect of pH. The iron complex is quite sensitive to pH as shown in Figure 1. There is a small pH range, from 2.5 to 3.0, within which the absorbance values remain constant. At the higher pH values, the reaction is not complete, probably because of the presence of hydroxy complexes of the ferric ion. Again at pH values lower than 2.5, the absorbance decreases because the reagent does not seem to be at optimum conditions to form a complex with ferric ion in the strongly acidic solution. Therefore, the measurement of absorbance was made at a pH of either 2.6 or 2.7. Stability of Color. The absorbance of a complex which is used as a standard solution in the determination of the organic acids should remain unchanged for prolonged time. In the present experiment the color of the complex develops immediately at room temperature, and is independent of temperature within the range of 10 to 50 "C. The color intensity remains unchanged for a week. Nature of the Complex in the Solution. The complex formed when 5-nitrosalicylic acid reacts with ferric ion has maximum absorbance at 492 mp in the visible range, and the reagent shows no absorbance in this range. Therefore, the absorbance was also measured against a demineralized water blank. The empirical formula of the iron complex is established by two independent methods: the mole ratio method of Yoe and Jones (6), and the continuous variation method of Job, as modified by Vosburg and Cooper (7). The results indicate that a stable complex is formed between 1 mole of iron and 1 mole of reagent. In the study of the stability constant of the complex by Ernst and Menashi (2),the mole ratio was 1 : 1. Determination of Organic Acids. By the foregoing procedure, the water-soluble organic acids, such as oxalic, malonic, malic, citric, tartaric, tartronic, and 1-ascorbic acid which can react with ferric ion in the ferric 5-nitrosalicylate, diminish the absorption of the complex quantitatively within a certain range of the concentrations of acids. The present authors investigated the difference of reactivities between various organic acids and ferric 5-nitrosalicylate complex in order to explain this phenomenon. Ferrous salts in solutions of mineral acids react with the organic base CY, a'dipyridyl to give a soluble deep red, very stable complex cation, but ferric salts do not react under these conditions. On the other hand, after reduction, tervalent iron can also be detected by means of CY,a'-dipyridyl. 5-Nitrosalicylic acid does not react with ferrous salts under the condition in the case of ferric ion. Ferric 5-nitrosalicylate solution is decolorized by adding a few drops of organic acid solution. Then a red or pink color-
.,
f6) J. H. Yoe and A. L. Jones. IND.ENG. CHEW.ANAL.ED.. 16. 111 (1944). (7) W. C . Vosburg and G. R. Cooper, J. Amer. Chem. Soc., 63, 437 (1941). I
_
L
Table I. Precision Determination Absorbancea __ Run Oxalic acid no. added,pg/ml 10 20 30 1 0.686 0.601 0.515 2 0.685 0.600 0.516 3 0.690 0.595 0.520 0.598 0.517 4 0.691 0.595 0.521 5 0.687 Average 0.688 0.598 0.518 Std dev 2.6 x 10-3 2.7 x 10-3 2.6 x 10-3 Re1 std dev, 0.38 0.45 0.50 a Ferric 5-nitrosalicylate: 3.58 X 10-4M; pH: 2.6. ~
I 0
10
I
I
20
1
I
3 0 4 0 5 0
Organic a c I d
I
1
60
70
I
(re/ll)
Figure 2. Effect of various organic acids on absorbance of ferric 5-nitrosalicylate M : malic, Ta: tartaric, Ma: malonic, C: citric, 0: oxalic, T: tartronic, A: 1-ascorbic
ation appeared by adding a drop of CY, a’-dipyridyl reagent. The CY,a’-dipyridyl test for Fez+in the solution obtained after reaction between the complex and organic acid is especially sensitive in the case of 1-ascorbic and tartronic acid. This test for the rest of the organic acids is not sensitive, but positive. Consequently, the authors suggest that the degree of diminution of absorption by organic acid is dependent upon the different reducing powers of organic acids and the reactivities between organic acids and ferric ion in its complex. As shown in Figure 2 the limits of linearity for the curves of absorbance cs. micrograms of organic acids and the slopes differ from each other because of different abilities of the acids in reacting with the complex. A series of tests on replicate samples was run to determine the precision of this method for the oxalic acid (Table I). The results of the analysis of organic acids are summarized in Table 11. Interference. An attempt was made t o investigate the possible interferences from several metal ions, anions, and two coexisting organic acids on the determination of the various organic acids in the sample. The measurements are
Ion added
Mg2+ Zn2+ cu2+
Cd2+ Co2f A cNOzco32-
c1-
Maleic acid Glutaric acid
Added as MgACi ZnAcn CUAC~ CdAcz COAQ NaAc NaN02 Na,C03 KCI
Table 11. Summary of Analysis of Standard Organic Acid Optimum concentration Standard Acids range, Pg/ml deviation Oxalic 4-45 0.14 Citric 2-30 0.35 Tartaric 2-50 0.39 Malonic 6-40 0.15 Malic 10-50 1.22 Tartronic 2-20 0.09 1-Ascorbic 2-30 0.07
made with the prepared solution containing 100 p g per ml of diverse substance with 10-20 p g per ml of organic acids. The results are presented in Table 111. The presence of two organic acids (maleic, glutaric acid) for the determination of organic acids does not interfere with each other except for a slight interference found in the cases of tartronic acid. Cations that complex with 5-nitrosalicylic acid and organic acids such as Zn2+,Cu2+,Cd2+, and CoZf and anions such as CH3COO-, NOz-, COa2+, and C1- are investigated. The presence of Cu2+ion in the 1-ascorbic acid solution interferes seriously in the analysis because of oxidation of 1-ascorbic acid by Cu2+ion, which reacts as oxidation-reduction catalyst (8). This may also be applied in the case of oxalic acid. Ferric ion in its complex is reduced considerably by 1-ascorbic and tartronic acid. It is observed, by the CY,o’-dipyridyl test, that anion NOz- interferes with reduction of ferric ion by two organic acids. Interfering ions such as Cu2+ and NOzshould, therefore, be separated from the solution by an ionexchanger before the analysis. (8) A. Weissberger and J. E. Lu Valle, J . Amer. Clzem. SOC.,66,
700 (1944).
Table 111. Determination of Organic Acid in Diverse Material Amount Oxalic Citric Tartaric Malonic Malic added, (20 Pg) (20 Pg) (32 Pg) (32 Pg) (36 Pg) Pg/ml Found, pg/ml 200 20.0 20.0 32.0 32.0 36.0 100 20.0 20.0 32.0 32.0 36.0 100 13.0 20.0 32.0 32.0 36.0 100 20.0 19.0 32.0 32.0 36.0 100 20.0 18.5 29.5 32.0 36.0 100 20.0 20.0 32.0 32.0 36.0 100 20.0 18.5 32.0 32.0 36.0 100 20.0 19.0 32.0 32.0 36.0 40 20.0 18.5 32.0 32.0 36.0 100 20.0 20.0 32.0 32.0 36.0 100 20.0 20.0 32.0 32.0 36.0
1-Ascorbic (10 Pg)
Tartronic (10 P&
10.0
10.0
10.0
10.0
1.5
11.7 9.5
8.5
7.5 10.0 0.5 9.0
10.0 10.0 10.0
VOL. 40, NO. 13, NOVEMBER 1968
10.0
10.0 5.0 9.3 9.5 9.5 9.5 2051
Application. These organic acids in magnesium acetate medium (pH 4.5 or 6.8) can be determined without the interference by excessive of MgZf ion (Table 11). This method, therefore, could be used to study the application of magnesium acetate as an eluant in the separations of organic acids including oxalic acid by a strongly basic anion-exchanger (Dowex 1-XS). The ferric salicylate complex was also studied for the determination of some organic acids following the same procedures mentioned previously in the case of ferric 5-nitrosalicylate.
Ferric salicylate has reddish violet color within the pH range of 2.5 to 3.0 and has an absorbance maximum at 530 mp. The results of tartronic and I-ascorbic acids by ferric salicylate complex were much less accurate but malic acid and tartaric acid can be determined with good accuracy, compared to those by ferric 5-nitrosalicylate complex. RECEIVED for review January 15, 1968. Accepted July 11, 1968. Work supported by the Ministry of Education of Korea for which the authors express their gratitude.
Consecutive Determination of Alkali Metal Bromides and Thiocyanates in Mixtures James E. Burroughs and Alan I. Attia Bosg- Warner Cnrp., Roy C. Ingessoll Research Center, Des Plaines, 111. 60018 ONEof the most accurate volumetric methods for determining alkali metal thiocyanates in the presence of halogens is the one proposed by Schulek ( I ) . This technique, which is based on the use of bromine t o oxidize the thiocyanate t o bromine cyanide followed by iodometry, does not permit the direct determination of bromides and thiocyanates in a single sample. Kolthoff (2) reported that thiocyanates could be determined gravimetrically as cuprous thiocyanates, the details of which may be found in general analytical texts (3). A review of the literature failed to indicate any attempts to utilize this fact t o selectively separate thiocyanate from bromides in their mixtures. Thiocyanates and bromides have been successfully identified in their mixtures after destructive oxidation of the thiocyanates with peroxides ( 4 ) . The present studies resulted in the development of a rapid, accurate, and successive determination of bromide and thiocyanate in mixtures with a single standard solution of AgN03. A potentiometric end point permits the application of the method to highly colored solutions. EXPERIMENTAL
Apparatus. A sulfide specific ion electrode (Model 94-16, Orion Research) and a standard fiber junction calomel electrode in contact with a 1M N a N 0 3 salt bridge were used with a recording p H meter (Model EUW-301, Heath Co.) to determine end points. Alternately, a glass-Ag/AgCl electrode system could be utilized. Reagents. All reagents were of analytical grade quality or better, and used without additional treatment. To establish the accuracy of the method, standard stock solutions of KBr and NH4SCN were used. Procedure. A solid or liquid sample, containing no more than 100 mg of the alkali metal thiocyanate and 250 mg of the alkali metal bromide, was weighed by difference into a 150-nil beaker. The sample was dissolved in about 40 ml of deionized water and acidified with 1-2 drops of 6N “ 0 3 . After complete dissolution, about 20 ml of 0.1M (CH3(I) E. Schulek, Z. Anal. Ciiem., 62, 337 (1923). (2) I. M. Kolthoff and P. J. Elving, Eds., “Treatise on Analytical Chemistry, Analytical Chemistry of the Elements,” Part 11, Vol. 7 , Interscience, New York, N. Y . , 1961, p 90. (3) F. P. Treadwell and W. T. Hall, “AnaIytical Chemistry,” Vol. 11, 9th Ed., Wiley, New York, N. Y . , 1955, p 303. (4) I. M. Kolthoff and V. A. Stenger, “Volumetric Analysis,” Vol. 11, Interscience, New York, N. Y . , 1957.
2052
ANALYTICAL CHEMISTRY
COO)z.Cu H20, followed by 0.25 gram of L-ascorbic acid, were added. The solution was mixed well, after which 10 drops of 0 . 2 z (w/v) of a phenol red solution were added. Concentrated “,OH was added dropwise until one drop imparted a distinct color change of the solution from yellow to red. About 3 ml of the “?OH were added in excess. After about five minutes of mixing on a magnetic stirring apparatus, dilute HN03(6N) was added dropwise until the solution was definitely orange (pH 7-8). One half of a n ashless filter tablet was added and mixed well with the solution. After allowing five minutes for the precipitation to be completed, the solution was filtered through a n S&S White Label filter paper, the filtrate being caught in a 250-ml beaker. The precipitate was washed three times with 10-ml portions of deionized water. The volume of the filtrate was adjusted to 125-150 ml with deionized water. To this solution was added in succession, with mixing, 5 ml concd HNOI and 10 ml of 2 0 x (w/v) Fe(N03)3.9 H 2 0 solution. After mixing for five minutes, the bromide content was determined by titration with a standard 0.1N AgN03 solution, with the sulfide specific ion electrode to determine the end point. The filter paper containing the precipitated CuSCN, from above, was transferred to a 250-ml beaker and diluted with 125 ml of deionized water. About 5 ml concd H N 0 3 and 10 ml of 2 0 z (w/v) Fe(NOJ3.9H?O solution were added. After mixing for five minutes, the thiocyanate content was determined by the same procedure used for the bromide. DISCUSSIOh
This method takes advantage of the differences in the solubilities of CuBr and CuSCN in ammoniacal solutions. Cuprous bromide is soluble in excess NHiOH while cuprous thiocyanate is insoluble, affording a rapid and complete separation of the two anions. The analytical applications of ascorbic acid (which is not an acid at all, but a lactone) have been reviewed by Erdey and Svehla (5) and discussed by Belcher and Wilson ( 6 ) . In the present studies, ascorbic acid was employed for the generation in situ of cuprous ions from cupric acetate, resulting in the formation of a CuSCN precipitate which was more easily filtered. In order to circumvent the simultaneous precipitation ( 5 ) L. Erdey and G. Svehla, Chemist-Andyst, 52,24 (1963). (6) R. Belcher and C. L. Wilson, “New Methods of Analytical
Chemistry,” Chapman and Hall Ltd., London, 1966, Chap 111.
