acidic medium, the prcwnce of heavy metal ions does not interfere in the subsequent measurements. Silver ion consumes reagent through a redox reaction and vanadium reacts to give a colored coordination compound. Results on Selected Samples. T h e results of determinations of cobalt and nickel on selected samples are presented in Table 11. The synthetic samples were prepared from mixtures of t h e perchlorate salts. Kational Bureau of Standards sample 161 n.as selected for its relatively low percentage of iron. Cobalt and nickel can be determined in varying ratios with the usual spectrophotometric accuracy of 1 2 7 , . Cominon anions and other metal ions, n-ith th(. pxception of those described above, do not offer serious interferencc in the concentration range of 500 to 2000 p.p .m. ACKNOWLEDGMENT
The authors thank the Research Corp.
Table 11.
Summary of Determinations
Cobalt, Present 6.25 SI 2.60 S? 1.25 Sa S4(20 nig. Fe 20 mg. Zn) 2.75 Sj(10 mg. AI 10 mg. Cu 20 mg. Fe) 5.80 S6 (10 mg. Cu 20 mg. Fe 20 mg. Zn) 2.60 NBS 161 Ni-Cr casting alloy 0.23 Sample
++
+
++
for the financial assistance which made this work possible. The authors also thank the hlallinckrodt Chemical Works for the generous sample of dithiooxamide. LITERATURE CITED
(1) Chatterjee, R., J. Indian Chem. Soc. 15. 608 (1938). (2) Fische;, E.,'Ber. 2 2 , 1931 (1889). (3) Gupta, H. K. L., Sogani, N. C., ANAL. CHEM.31, 918 (1959). (4) Jonassen, H. B., Chamblin, V. C., Wagner, V. L., Jr.,lbid., 30,1660 (1958).
Mg. Found 6.19 2.60 1.24 2.72 5.86
2.53 0.22
Kickel, Mg. Present
Found
1.25 2.60 6.25 4.55 5.80
1.26 2.54 6.11 4.50 5.80
2.60
31.50
"57
31 , 1 5
(5) Kraus, K. -4.,Moore, G. E., J . -4m. Chem. SOC.75, 1460 (1953). (6) McDowell, B. L., Meyer, A. S.,Jr., Feathers, R. E.,
ANAL.CHEM.31, 931 ( (7) Pearse, G. A., Jr., Pflac J . Am. Chem. Soc., in press. ( 8 ) Pflaum, R. T.. Dieter. L. H.. Proc. Iowa Acdd. Sci. 64, 235 (1957). ' (9) Wheatley, R. D., Colgate, S. O., ANAL.CHEM.28, 1897 (1956). '
RECEIVED for review March 17, 1959. Accepted November 5, 1959. Division of Analytical Chemistry, 135th Meeting, .4CS, Boston, Mass., April 1959.
Determination of Iron and Iron-Aluminum Mixtures by Titration with EDTA DONALD G. DAVIS' and WILLIAM R. JACOBSEN2 School o f Chemistry, Georgia institute o f Technology, Atlanta 7 3, Ga.
b Iron in the presence of aluminurn has been titrated with EDTA. To avoid the interference of aluminum, a pH of 1.0 was used. The end poin', was determined spectrophotometrically with 5-sulfosalicylic acid acting a the indicator. Results were accurate to 0.3%, even though the amount of aluminum was twice that of iron, if the iron concentration was not much less than 10-'M in the solution being titrated.
S
methods for the direct titration of iron with (ethylenedinitrilo) tetraacetic acid (EDTA) have been proposed ( I O ) . Although Cheng, Bray, and Kurtz (3) have reported that aluminum does not interfere with the titration of iron at p H 2 t o 3 except in "large amounts," Sweetser and Bricker ('7) found t h a t the presence of aluminum caused noticeable EVERAL
1 Present address, Louisiana State University in New Orleans, New Orleans, La. Present address, Savannah River Plant, E. I. du Pont de Nemours & Go., Inc., Aiken, S. C.
positive errors in the p H range between 1.7 and 2.3. Because the findings of Sbveetser and Bricker were confirmed in this laboratory, a theoretical and experimental study was undertaken to establish conditions under which the titration of iron could be performed free from the interference of aluminum. CHEMICALS AND APPARATUS
,411 chemicals used were the best obtainable, usually reagent grade. Standard solutions of iron and aluminum rvere prepared by dissolving a weighed amount of the pure metal in the appropriate acid and diluting t o exactly 1 liter. The standardization of the iron solution was checked by titration with potassium dichromate solution according t o the classical procedure. The titrant solutions of EDTA were prepared by dissolving the reagent grade disodium salt of (ethylenedinitri1o)tetraacetic acid in distilled water. These solutions were stored in a polyethylene bottle and standardized against a standard bismuth solution at p H 1.2 t o 1.4 (9). Bismuth was chosen as a standard so that the EDTA could be standardized at approximately the pH at which it was t o be used. This procedure eliminates slight errors due to
the presence of extraneous metal ions in the distilled water, such as calcium and copper. Ascorbic acid was added t o mask iron impurities in the bismuth. The indicator solution was a 2 weight yo solution of Eastman White Label 5-sulfosalicylic acid dissolved in mater. It was stable indefinitely. The spectrophotometric titrations were performed using a Beckman Model D U spectrophotometer. The cell compartment of this instrument was replaced by a 3.75 X 3.75 X 5.50 inch black painted aluminum box, with holes in the side to alloiy the passage of light from the optical system to the phototube. The buret tip protruded through the lid of the box into the titrate, which was contained in a 100-ml. tall-form beaker held in the light path. The solution was stirred between additions of titrant by a magnetic stirrer placed beneath the titration compartment. Calibration volumetric glassware \vas used throughout this work. p H was adjusted by using a Beckman Zeromatic p H meter equipped with the usual glass and calomel electrodes. PROCEDURE
An accurately measured portion of VOL. 32, NO. 2, FEBRUARY 1960
215
standard iron solution was diluted to about 60 to 80 ml. and the pH adjusted to the desired value (usually 1.0 to 1.1) with a concentrated solution of sodium acetate. The addition of a base was necessary, because the solutions of iron and aluminum were on the order of 1M in hydrogen ion. (Most practical samples would also have been brought into solution by treatment with acid.) When the p H adjustment was completed, 8 t o 10 drops of indicator solution were added and the titration was performed either visually or spectrophotometrically. Visual titrations were, however, difficult: especially a t p H 1.0. The color change was very gradual and indistinct to the eye. If the disappearance of the iron-sulfosalicylic acid complex was followed a t 510 mM (1) using the spectrophotometer, sharp and accurate titration curves were recorded (Figure 1). The p H of the solution titrated did not change measurably during the titration. DISCUSSION AND RESULTS
On the basis of information available in the literature (2, 6) it is possible to calculate the effectiveness of any proposed titration (4). The absolute stability constants of metal-EDTA complexes are defined as:
where M is the metal ion of interest and Y is the anion of EDTS. (Charges are omitted for simplicity.) In any practical case an apparent stability constant must be calculated which takes into account the pH and the presence of complexing agents other than EDTA. The apparent stability constant is related to the absolute stability constant by: K (apparent) = K (absolute) ffP
where cy and @ are, respectively, a measure of the effect of pH on the dissociation of EDTA and a factor to account for the complexing of 11 by the
Table I.
0.75
0.65
2
z
0.55
I--d 0.2M L
0.45
.
Figure 1 Spectrophotometric titration curve of 0.1 F iron alone a t pH 1.O
indicator and/or the buffer. (Y and p were derived in the usual way (4). The logarithms of the absolute stability constants for the EDTA complexes of ferric iron and aluminum are 25.1 and 16.13, respectively. At pH 1.0, log QI was calculated to be 17.1 but 0 was only slightly greater than 1. [The @ factor was calculated using the value of 1014.42 (1) for the stability constant of the iron-5-sulfosalicylate acid complex.] From Equation 2 it can easily be calculated that the apparent stability constant for the iron EDTA complex is 8.0. while that for aluminum is - 1.0. As a general rule, the minimum log K apparent for a satisfactory titration in the presence of an indicator is considered to be 6-log C, where C is the metal ion concentration after dilution. The difference between the logarithms of apparent stability constants of two different metal ions present in a mixture must be greater than 4-log C for a differential titration to be feasible. This rule is computed with the idea that either no indicator or one which forms a very weak metal ion complex is used. Because the apparent
Determination of Iron Alone and in Mixtures with Aluminum a t pH 2 to 3 b y Titration with EDTA
Method Visual
Spectrophotometric 'Visual Spectrophotometric
216
0.85
A1
Toernp.,
Present,
Fe Taken,
Fe Found,
Error,
C. 40-50
hlg.
Mg. 55.87
07
25
...
25 40-50 25
51 51 51 102
Mg. 55.96 56.03 55.80 56.91 55.85 5.61 5.64 65.9 65.1 z5.1 15.3
ANALYTICAL CHEMISTRY
5.59 60.9 74.5
/O
+ O . 16 +0.29 -0.13 $0.07 -0.04 $0.4 +0.9 $8.3 +6.9 +0.8
fl.1
stability complex of the iron-5-sulfosalicylic acid complex is about 102 a t p H 1.0, this last rule should be used. These considerations indicate that it is theoretically possible to titrate iron in the presence of aluminum a t pH's down to 1.0 using 5-sulfosalicylic acid as an indicator. Experimentally it was found that two important limitations must be considered: The end point becomes difficult to determine visually if pH is controlled beIow 2.0, and the interference of aluminum is noticeable above p H 2.0. The results in Table I exemplify the errors associated with the determination of iron-aluminum mixtures a t pH 2 to 3 by various methods. Good results are obtained for iron alone in this pH range even when the end point was determined visually. The use of spectrophotometric end point detection seems to lessen but not eliminate the interference of aluminum. Visual end points were very unsatisfactory a t p H 1.0, but good results could be obtained spectrophotometrically. The iron-indicator complex is so weak a t p H 1.0 that its color fades greatly during the titration but is not distinctly extinguished at the end point. Although a metal-indicator complex with an apparent stability constant of only 100 mould not be expected to give good results visually, the spectrophotometric end points were actually found by extrapolation of straight lines established from measurements made a t some distance from the end point, where an excess of iron or of EDTA would cause the equilibrium to be shifted essentially completely in one direction or another. No interference due to aluminum was found, as is shown in Table 11, except at rather low iron concentrations. These positive errors are due to the fact that the apparent stability constant of the iron-EDT-4 complex has been reduced to a value of about lo8 or possibly IO7. If the concentration of iron in the solution being titrated is then lowered belolv about O . O l X , a measurable excess of EDTA will be used up in the complexation of aluminum before the iron is completely removed from its indicator complex. Actually this phenomenon expresses itself in the spectrophotometer titration curves by making the intersection more obtuse and rounded. The best results are obtained only with samples containing 50 mg. or more of iron per 100 ml. of solution if aluminum is present. This effect of concentration agrees approximately with the rule presented above. Although it would seem that increasing the pH should have no effect on the difference between the apparent stability constants for iron and aluminum, this titration was subject to
noticeable positive errors even a t p H 2.0 (compare Tables I and 11). This might be thought to be due to a n increase in the B factor for iron. However, as this increase is small and accurate titrations for iron alone a t p H 2 to 3 are easily accomplished, the blame for erroneous iron results must rest on the aluminum. Although the concentration of alumi-, num-hydroxy complexes is believed to be negligible below pH 3 (8),the establishment of equilibria is notoriously slow. I n addition, mixed complexes of aluminum, hydrogen ion, and EDTA can esist in the p H range under consideration ( 2 ) . For these reasons aluminum is apparently able to compete with the iron to a slight extent for complexation above pH 2 but not a t p H 1. Flaschka and Abdine (5) have devised a procedure for the determination of aluminum alone or the sum of iron and aluminum by titrating a boiling solution of the metal ions with EDTA to a copper-PAX end point a t p H 3.0. This method may be applied to the determination of aluminum directly following the spectrophotometric titration of iron. Ammonium acetate vias used to adjust the pH to 3.0. Several
Table II. Determination of Iron-Aluminum Mixtures a t pH 1.0 b y Titration with EDTA
(End point found spectrophotometrically) -41 PresFe Fe ent, Present, Found, Error, Mg. Ng. Mg. 70 0 60.9 60.Ta -0.3” 5-~1
60.7
-0.3
51 102 51 0 50 100 150 0 5 25 50
60.5 60.9 121.2 55 80 55.77 55.83 55.77 5.64 5.74 5.76 5.79
-0.7 0 -0.6 -0.12 -0.18 -0.07 -0.18 $0.9 +2.7
(1
121.8 55.87
5.59
+3.0 $3.6
iiverage of 9 determinations.
typical iron-aluminum mixtures were determined by the combined procedures with an accuracy to 0.3%. Slightly more than enough EDTA to complex all the iron must be added to establish the base line of the spectrophotometric titration curve. If the amount of aluminum is smaller than the excess, a back-titration method would
be necessary rather than the direct titration of Flaschka and Abdine. ACKNOWLEDGMENT
The authors thank Hermann Flaschka for his many helpful discussions and suggestions. LITERATURE CITED
(1) Agren, Allen, Acta Chem. Scand. 8 , 266 (1954). (2) Bjerrum, Jannik, Schwarzenbach, G., Sillin, L. G., “Stability Constants,” Part I, Chemical Society, London, 1957. (3) Cheng, K. L.,Bray, R. H., Kurtz, T., ANAL.CHEM.25,347 (1953). (4) Flaschka, H.A., “EDTA Titrations,” Pergamon Press, London, 1959. (5) Flaschka, H.A , , Abdine, H., 2.anal. Chem. 152, 77 (1956). (6) Schwarzenbach, Gerold, “Complexo-
metric Titrations.” Interscience. New York. 1957. (7) Sweetser, P.B.,Bricker, C. E., ANAL. CHEM.25,253 (1953). (8) Wanninen, Erkki, Ringbom, A,, Anal. Chim. Acta 12,308 (1955). (9)Welcher, F. J., “Analytical Uses of Ethylenediaminetetraacetic Acid,” p. 209, Van Nostrand, Princeton, N. J., 1958. (10) Ibid., p. 222.
RECEIVEDfor review July 29, 1959. Accepted November 6,1959.
Carbon-Hydrogen Stretching Frequencies STEPHEN E. WIBERLEY, STANLEY C. BUNCE, and WALTER H. BAUER Department o f Chemistry, Rensselaer Polytechnic Instifufe, Troy, N. Y. b A correlation chart of the carbonhydrogen stretching region is presented and discussed. Spectra of 10 cyclobutyl compounds and 60 aromatic compounds in the region of 2700 to 3 1 00 cm.-l are shown. This information should prove of value in spectrastructure correlation studies of related materials.
length region are summarized in Figure 1. The range for the various groups is shown, the average value for the group being indicated by a short line over the range. The maximum and minimum values are based on current spectra and undoubtedly as more data become available these ranges will broaden. ALIPHATIC C-H COMPOUNDS
B
lithium fluoride prisms free of the hydrogen fluoride absorption band are now generally available, and with the advent of commercial infrared spectrometers which have a grating attachment for a sodium chloride prism, it seems pertinent to review the information which can be obtained by studying spectra measured with such a prism or a grating attachment. In addition, new data on cyclopropyl, cyclobutyl, and 60 aromatic compounds are presented. In particular the region involving carbon-hydrogen stretching frequencies -namely, from 2700 to 3100 cm.-’-is covered. The correlations in this wave ECAIJSE
Bellamy ( 2 ) has reviewed thoroughly the work of Fox and Martin (6) on hydrocarbons and Pozefsky and Coggeshall (22) on compounds containing sulfur and oxygen. As shown in Figure 1, the CH, bands occur a t 2960 and 2870 cm.-’ and the CH2 bands a t 2930 and 2860 cm.-l. In many cases only one band in the region of 2860 to 2870 cm.-’ can be resolved. In a given straight-chain homologous series such as the fatty acid series, the CHBband a t 2960 is stronger than the CH2 band a t 2930 when the ratio of methylene to methyl groups is 4 to 1 or less (8). With larger ratios the 2930 band is more intense. TT7ith branched-
chain acids the ratio need only be 3 to 1 before the relative intensities reverse. However, exceptions to this rule occur when the methyl group is adjacent to the carboxyl group, as is shown by the comparison of 2-methylhexanoic acid with 3-methylhexanoic acid (9). In Figure 1 the frequency range for the methyl and methylene groups in compounds containing oxygen and sulfur has been determined from the data of Pozefsky and Coggeshall ( 2 2 ) . For oxygenated and sulfur-containing compounds the methyl and methylene bands are approximately 7 cm.-l higher and the extinction coefficients are greater than for the corresponding hydrocarbons. .A single methyl group attached to-a nitrogen atom gives rise to a band in the 2780- to 2805-cm.-l range when the group is in an aliphatic or a nonaromatic heterocyclic system ( I S ) . In an aromatic system the band is between 2810 and 2820 cm.-l. Two methyl groups attached to a nitrogen atom yield two bands: one between 2810 and72825 crn.-l and the other between 2765 and VOL. 32, NO. 2, FEBRUARY 1960
217
