ANALYTICAL CHEMISTRY
1538 been assigned to the H-COstructure. Alcohols and acids absorb at several wave lengths near 2.8 microns but in different relative intensities. It appears that the hydrogen-bonding vibrations involving -OH, -CO-, -CHO-, -COOH, and -COOR absorb at several frequencies in this range, and that identification of structure must involve consideration of the relative intensities. Acids absorb predominantly at 2.72 and 2.83 microns and alcohols at 2.75 microns. In caprinoin (11hydroxy-10-eicosanone, curve 57) the neighboring -OH and -COgroups absorb in a unique couplet at 2.05 and 2.12 microns. Esters. Methyl esters have sharp maxima at 2.14 and 2.26 microns and weak ones a t 1.90 and 1.95 microns (curves 41,42). A maximum at or near 2.28 microns is also found in the spectra of substances having dominant CHI groups. However, the bands at 2.14, 1.95, and 1.90 microns’ are found only in esters. Ethyl esters (curves 9 and 22) exhibit similar absorptions. Conjugations of an ester with double bonds causes multiple bands to appear near the maxima at 2.14, 1.95, and 1.90 microns (curves 33 and 34). Ethyl acetate (not shown) has a strong band at 2.26 microns, a weaker band at 2.14 microns, and a still weaker one at 2.13 microns. The absorption at 2.30 microns, which is absent in methyl acetate, is clearly expressed by a strong band in ethyl acetate. In the glycerol esters the band at 2.26 microns is absent and the one at 2.14 microns is obscure. However, acetic and propionic acids exhibit absorption at 2.26 microns suggesting that the CHa on a short-chain compound absorbs similarly to the methyl and ethyl esters. The 1,3-dipalmitin spectrum (curve 11) shows hydroxyl absorption at 1.42 and 2.77 microns. DISCUSSION AND CONCLUSIOIVS
Kear-infrared spectra (1 to 3 microns) offer some advantages over conventional infrared spectra (2 to 15 microns) in structure
identification. For example, the cis unsaturation has specific absorption in the near-infrared which is easily resolved from nearby bands, and which should lend itself to quantitative measurement. However, the fundamental band at 3.30 microns is not so easily resolved and detected without the use of the lithium fluoride prism in conventional equipment. Hydroperoxide groups are easily distinguishable by their near-infrared spectra, whereas infrared spectra (2.8 microns) do not distinguish these structures. On the other hand, trans double bonds can be detected best by the infrared spectra (10.3 microns). The near-infrared spectral region allows detection of many structures whose fundamental vibrations lie in the conventional infrared region. From this study, which was aimed a t characterization of structures encountered in studies of lipides, it has been amply demonstrated that near-infrared spectra are valuable in studies of aliphatic substances, and it appears that their usefulness may be extended t o studies of other organic structures. ACKNOWLEDGMENT
This work was aided by grants-in-aid from the Atomic Energy Commission (Contract AT 11-1-108), the Office of Naval Research (contract S8onr 66218) and the Yational Dairy Council. The authors are indebted to J. R. Chipault for valuable discussions concerning this problem. LITERATURE CITED
Kaye, W., Spectrochim. Acta6, 257 (1954). (2) Wheeler, D. H., “Progress in the Chemistry of Fats and Other Lipids,” vol. 11, p. 268, Pergamon Press, London, 1954. (3) Wulf, 0. R., Liddel, U., J . Am. Chem. SOC.57, 1464 (1935).
(1)
RECEIVED for review March 22, 1956. Accepted June 28, 1956. Institute publication No. 141.
Hormel
Determination of Water in Fuming Nitric Acid by Near-Infrared Absorption LOCKE WHITE, JR., and WILLIAM J. BARRETT Southern Research Institute, Birmingham 5, A l a .
Water in fuming nitric acid can be determined by its optical absorption at 1.423 microns. A simple modification of the Beclrman DU spectrophotometer is convenient for the determination. In addition to the water that can be chemically determined, there is a small amount of water formed by the self-dissociation of the acid. Because this self-dissociation of the acid is suppressed by nitrate ion from the dissociation of nitrogen dioxide, nitrogen dioxide has a small infiuence on the absorption due to water. Corrections for this effect are given; unless both water and nitrogen dioxide contents are low, the correction is negligibly small.
T
HE near-infrared absorption of fuming nitric acid is an almost specific measure of its water content. This statement is true for water contents up to at least 6y0 and nitrogen dioxide contents up to about 20’33, at a wave length of 1.423 microns.
BACKGROUND
Dalmon and Freymann ( 1 ) found that the addition of water to pure nitric acid caused a weak absorption band to develop at 0.97 micron. The strength of the band increased with increasing concentrations of water. Kinsey and Ellis ( 7 , 8 ) observed that the absorption bands of pure liquid water at 1.92 and 3.0 microns occurred also in 95% nitric acid, almost unaffected except for a slight shift in wave length and some sharpening. Liquid water has been reported (2) to have absorption bands at approximately the following wave lengths in microns: 0.98, 1.18, 1.45, 1.74, 1.79, 1.96, and 2.79. XOabsorption bands which can be attributed to nitrogen dioxide have been reported between 1 and 2 microns. This is important because fuming nitric acid often contains nitrogen dioxide (which name is usually used collectively in this paper to include also nitrogen tetroxide and all the products of its ionization). APPARATUS
The basic optical element used in this work is the Beckman Model D U quartz spectrophotorrleter. To achieve adequate
V O L U M E 28, NO. 10, O C T O B E R 1 9 5 6
1539
sensitivity in the near-infi ared region, the phototube housing and amplifier of the spectrophotometer were replaced with an interchangeable unit consisting essentially of a 300-cycle-persecond chopper, a lead sulfide photoconductive cell, and a tuned amplifier. The new phototube housing contains the chopper, located just after the cell compartment; an off-axis paraboloid, identical with that in the lamp housing, about 8 inches from the exit slit, to form a reduced image of the exit slit on the phototube; a Continental Electric CE-GiOl-C3 photoconductive cell; and a preamplifier. Reducing the image of the exit slit permits closer spacing of the electrodes in the lead sulfide cell, and this improves the signal-to-noise ratio. A small separate cabinet contains a tuned amplifier, a meter, and appropriate controls. The meter is recalibrated in logarithmic units, and a scale switch increases the gain of the amplifier in steps of two, so that the meter can always be used in the upper half of its scale. An uncalibrated gain control is used to adjust the meter to full scale under reference conditions. During the course of the work a variety of chopping frequencies and of amplifiers has been used with essentially identical results. However, six instruments were modified in the particular way described above; several of them have performed satisfactorily in other laboratories. Complete circuit diagrams and fairly complete sketches for the chopper-phototube-preamplifier housing are available from the authors. The wave-length scale of the Beckman spectrophotometer may be significantly in error around 1.4 microns. I t has been shown that a wave-length error of 0.006 micron can make an apparent difference of almost 0.3% water in a sample containing 5y0 water. Therefore, it is important that the wave-length scale of the instrument be checked, or that the apparent wave length of maximum absorption be carefully determined and used. This wave length was estimated, in terms of nearby mercury lines, to be 1.423 i0.002 microns. The original exploratory work for this method was done at 1.92 microns. At that wave length, however, sample cells suitable for water concentrations up to 5% could be no thicker than about 0.5 mm. All the work reported here was performed a t 1.423 microns, where easily reproducible 5-mm. cells are appropriate. Because the vapors from the samples are so corrosive, it is almost essential to use glass-stoppered cells. Phoenix Precision Instrument Co., Philadelphia, Pa., has supplied cells on special order that are almost entirely satisfactory for this use. However, one difficulty arises. The phototube is some distance from the sample cell, and the image ultimately focused on the phototube is small. If the windows of the sample cell are not truly parallel, the cell acts as a prism and shifts the position of the focused image on the phototube. As a result, the image may move to a region of different sensitivity. About half the sample cells received from Phoenix showed significant differences in their apparent transmittances because of prism effect when turned through 180" in the cell holder. Approximately half of the Beckman 5-mm. cells tested also showed the same effect. Thus far it has seemed preferable simply to reject cells that show this effect rather than to attempt to manufacture them to closer parallelism. In order for the cells to fit into the regular cell holder the ground-glass joints must be small. A glass hypodermic syringe with a tip drawn to a capillary, or preferably a Teflon syringe, is convenient for filling the cells through these narrow openings. As a reference for absorption spectrophotometry, pure nitric acid is both too unstable and too inconvenient. In the mork described here a bar of plate glass was used as reference. EQUILIBRIA IN FUMING NITRIC ACID
If water content of nitric acid is plotted as a function of absorbance at 1.423 microns, two conclusions result: The data show unmistakable curvature, as in Figure 2, and the results are slightly dependent upon nitrogen dioxide concentration, as in Figure 4. These two phenomena are explained by two equilibria which exist in nitric acid: 2HNO3 = liOz+
+ xO3- + HzO
(1)
and
N \ T ~=O ~2x02 = NO+
+ NOa-
(2)
The water from self-dissociation of nitric acid cannot be determined chemically, but it would be expected to be just as effective in absorbing infrared radiation as the water that, because it can be determined chemically, is normally considered to be the water content. And because both of the above equilibria involve
nitrate ion, the nitrogen dioxide content would be expected to influence the infrared absorption due to water. A third equilibrium that exists in nitric acid is
H20
+ "03
=
H30+
+ N03-
(3)
The formation of hydronium ions raises the question as to whether part of the optical absorbance observed might not be due to the ion, H30+, as well as to water. However, it has been shown ( 3 ) that Reaction 3 proceeds to a limited extent, and that the concentration of hydronium ions relative to that of water molecules (or solvated water) is very small.
A
t
\I/ I
-d 2 5 m m . TEST TUBE
I
1
I
12-mrn. O.D. TUBING
STOPPER 2 4 / 4 0 JOINT I N N E R PART
I-mm. CAPILLARV TUBING
ABSORPTION CHAMBER PATH L E N G T H , 10 mrn.
0 Figure 1.
Special absorption cell
Ingold and his associates have estimated the value of the equilibrium constant for Reaction 1 by cryoscopic measurements a t -40" C. ( 3 )and by Raman spectral measurements a t 15' C. (6). The results by the two methods agree within their experimental uncertainty, but the uncertainty in the Raman measurements was fairly large, corresponding to zk 20% of each of the dissociation products. The literature contains no relevant measurements at higher temperatures. For Reaction 2, Ingold's group ( 4 ) showed by Raman spectroscopy a t -10' C. that almost none of nitrogen dioxide in nitric acid existed as nitrogen tetroxide and that the material was "chiefly" ionized. They did not estimate quantitatively the degree of ionization. Lynn, Mason, and Sage (9) estimated from spectrophotometric observations that nitrogen dioxide was about i O % ionized at 0" C. An independent estimate of the extent of self-dissociation of nitric acid was obtained during this work. Water was added in weighed increments to a known weight of nitric acid containing excess nitrogen pentoxide. After each addition the absorbance of the resulting mixture was determined at 1.423 microns. The resulting empirical curve of absorbance as a function of p-eight of added water can be interpreted, as explained below, to yield both a value of the self-dissociation constant and a calibration of the spectrophotometric method for water, entirely independent of any chemical analyses for water.
-
The nitrogen entoxide was distilled directly into the special absorption vessef)(Figure I). When enough nitrogen pentoxide
1540
ANALYTICAL CHEMISTRY
r
SPECIES WATER, %
Figure 3. Correction to be subtracted from species water to yield chemical water K = 5 X 10-5 EXCESS
N2O5-
-EXCESS
H20
C H E M I C A L WATER, % (OPEN C I R C L E S 1 I
I
v~~iupper).--Same~crllEulati~n-f6;.~= IO -4 V (1ou.er). Same calculation for K = 2.5 X 10-5 - - - Relationship between chemical water and species water that would exist if no ionization effects were observed-if chemical equals species water
had condensed in the cell, about 50 to 60 grams of acid, very low in water content, and at a temperature just above the freezing point, was added. The solution was brought to 25" C. and the nhqnrhnnw nt. 1 423 miornnq
W'RS
mensiirerl
hv nlioinv the
rectangular cross-section part oi the absorption vessel in the cell compartment of the spectrophotometer. Water was added from a weighed hypodermic syringe through the capillary in increments of 0.10 to 0.15 gram. Mixing was accomplished by pouring the acid from one leg of the apparatus to the other without wetting the stopper. Thus, the addition of the water and the mixing were carried out with no exposure of the acid to the atmosphere. Prior to the first addition of water a small sample of acid was withdrawn for nitrogen dioxide determination. After the last measurement of absorbance another sample was withdrawn. That decomposition was negligible was evident from the fact that the actual decrease in concentration of nitrogen dioxide agreed with the calculated decrease. The interpretation is based on the assumption that the absorbance of the sample is a linear function of its species water content. (Species water content is the sum of the water content from selfdissociation and the water content determined by chemical analysis. The latter is hereafter called chemical water.) The interpretation assumes also that there is negligible species water originally, that the density of the sample does not change over the entire range of compositions, and that the nitrogen dioxide is totally ionized. The last two assumptions simplify the calculations. They could be refined, but the results seem to be precise enough without the complications resulting from refinement. With the underlying assumptions just stated, one attempts to construct a theoretical curve of absorbance u8. chemical water content that matches the empirically observed curve. This can be done by a trial-and-error estimate of a combination of two constants: (1) the slope of the calibration line relating absorbance and species water content, and (2) the self-dissociation constant of the acid. If any combination of these two constants produces a theoretical curve essentially identical with the observed curve, then these two constants-and the underlying assumptions-are presumably correct. This procedure was followed with two different initial samples.
After the water had been added, these samples contained 0.7 and 5.7% nitrogen dioxide. In each series about 20 increments of water were added, and the total increase in weight in each series agreed with the sum of the increments within f1% of the weight of added water, or f O . O 5 % of the total weight of acid. In neither case was the agreement b e b e e n theoretical and empirical curves perfect. However, theoretical curves which bracketed each of the empirical curves were calculated. Figure 2 is a comparison between one of the empirical curves and one of the theoretical. The range of uncertainty corresponding to the bracketing theoretical curves can be expressed as follows: The species water content throughout the curves was established within =kO.l%, and the equilibrium constant (4) was shown to be considerably closer to 5 X 10-5 than to 2.5 X lo-' or to (concentrations expressed in weight per cent). This is a sufficiently accurate determination of the equilibrium constant for the spectrophotometric calibration. Construction of Theoretical and Experimental Curves. I n order to clarify the procedure by which the best value for the equilibrium constant was obtained, the construction of Figure 2 is described in some detail. The experimental data were plotted as absorbance us. weight of added water. The data from one run are shown as solid circles. The starting material for this run was 60.834 grams of nitric acid mixture containing 0.69% nitrogen dioxide and excess nitrogen pentoxide. At the completion of the run the nitrogen dioxide concentration was 0.677,; the calculated value, based on the weight of water added, was 0.65%. This small difference was neglected in the calculations that follow. First, the linear relation between absorbance and concentration of species water was approximated. To obtain the intercept, the species water was assumed to be zero for the first few points, those with the least added water; their absorbance was essentially constant at 0.228, independent of the amount of added water. To obtain the slope, a straight line was drawn through the last few points, those with the most added water, where self-dissociation is largely suppressed. The intersection of this line with the scale of grams-of-water-added approximated the experimental composition at which the species water content was zero. The intersection was located at about 1.0 gram. Thus, it was possible to use this approximation in conjunction with the experimental weight data to calculate species water concentrations in per cent units for several of the higher points. These values for per cent species water and the corresponding absorbances yielded the follo-xing approximate equation: Absorbance = 0.228
+ 0.428 (% species water)
(5)
V O L U M E 28, NO. 10, O C T O B E R 1 9 5 6 Up to this point, self-dissociation has been disregarded, arid no assumptions dependent on it have been made. However, the curvature of the experimental line (solid circles) indicates that some dissociation does occur and that it affects the absorbance in the region near 1.0 gram. ilt this point the experimental value of the absorbance, by graphical interpolation, was 0.332, corresponding by Equation 5 to a species water content of 0.24%. This figure, with the known concentrations of nitrogen dioxide and nitric acid, gave, as a first appro\-imation to the equilibrium constant, 2.0 X The estimated composition of the final acid, obtained as described above, made it possible t o coiivert the scale of gramsof-water-added into an estimated scale of per cent chemical FT ater. At the higher points, the chemical water concentration is nearly equal t o the species water concentration. The plot of otiserved absorbance us. estimated per cent chemical water was compared with the theoretical curve of absorbance us. chemical water calculated from Equation 5 and the estimated equilibrium constant. The nature of any discrepancy between the theoretical curve and the observed points showed how t o improve the estimates of the equilibrium constant and the slope for the next approximation. For the final approximation the slope was 0.44 and the equilibrium constant was 5.0 X These values are represented in Figure 2 by the open circleq. CALIBRATIOS
Except for the “absolute” measurements just described, thc calibration of the method must depend on chemical analysis of the acid. I n the analysis, water is determined by difference after titrations for total acidity and nitrogen dioxide. By careful refinement of procedure, analyses believed to be accurate within 10.05y0 of water were achieved for the calibration. Agreement between the two methods of calibration fell within the limits of the experimental error of the chemical analysis. On the basis of 56 samples analyzed, the chemical water contents having been corrected t o species water contents with the equilibrium constant quoted above, the following straight line defines the relationship between species mater content and absorbance: % species water = 4.74 (absorbance) - 0.58 (6) These samples contained up to 105; nitrogen dioxide and up to 6% water.
1541 Figuie 3 gives in graphical form the correction to be subtracted from species water content to yield chemical water content. The correction is obviously a function of the unknown nitrogen dioxide content as well as the water content. Unpublished work in this laboratory has shown that nitrogen dioxide can also be determined spectrophotometrically (at 450 to 500 mp, depending on concentration). Unfortunately, temperature has an important influence on that determination. Howver, unless the acid is low in both water and nitrogen dioxide, the correction is rather small and constant. Therefore, it is not usually necessary t o h o w the nitrogen dioxide content precisely. One evidence of the validity of the corrcctioii for nitrogen dioxide content is shown in Figure 4. Figure 4, A , shows chemical water content as a function of absorbance, ~ i t h o u tany allowance for nitrogen dioxide content. Figure 4, B , allows for nitrogen dioxide content; it modifies the data of Figure 4, A , by first including the appropriate corrections from Figure 3 and then plotting species water instead of chemical water. The reduction in the spread of the data in Figure 4, B, as compared with Figure 4, A, is evident. The data of these two figures cover a narrow range of concentration in which the nitrogen dioxide effects are most pronounced. At higher water contents the correction for nitrogen dioxide is much less important, as Figure 3 shows. Calibration Equation 6 and the nitrogen dioxide corrections of Figure 3 were deduced from data on samples containing up to 10% nitrogen dioxide. Less extensive measurements have been made on samples containing 12 to 17y0 nitrogen dioxide. For these samples a slightly different calibration equation, k n o m with somewhat less precision, applies:
% chemical water
=
5.00 (absorbance) - 0.67
(7)
In the presence of so much nitrogen dioxide, self-ionization of nitric acid must be negligible. Therefore, the phenomena previously discussed here mould predict that Equations 6 and 7 should be identical. On the contrary, theie is small but significant difference between them. Among phenomena not pieviously allowed for, addition of nitrogen dioxide a t constant water content incieases the density of the mixture (6). Because this effect concentrates more water in the optical beam, it would shift the calibration toward lower weight content of water for a given absorbaiicc. But this tendency is in the wrong direction t o account for the difference between Equations 6 and 7. However, if nitrogen dioxide reacted with water to produce nitrous and nitric
1542
ANALYTICAL CHEMISTRY
acids, as it does in dilute aqueous solutions, that reaction could account for the difference in the calibration line at high nitrogen dioxide contents. At one stage of the work the absorbances at 1.423 microns of about 30 samples containing up to 5% water were determined on each of four modified Beckman DU spectrophotometers. The variations among the instruments never exceeded 0.27, water equivalent, and the average spread was only about 0.1% water. Over the range of normal laboratory temperatures, the spectrophotometric determination of water in nitric acid is not significantly influenced by temperature. Early in the work, before the techniques were precise enough to show the nitrogen dioxide effect, the influence of dissolved nitrates of iron, nickel, and chromium was briefly investigated. KO influence on the water determination was detected. It is assumed, therefore, that dissolved nitrates do no more than suppress the extent of the self-dissociation of the acid. The method has been in use in several laboratories for 2 years or more.
assisted in developing the instrumentation and in making preliminary measurements. Twenty-three of the analyses used in establishing the final calibration were performed by the Kava1 Air Rocket Test Station, Dover, N. J., through the courtesy of J. D. Clark and H. G. Streim. The probable existence of the effects of ionization equilibria was predicted to the authors by H. E. Higbie of M. W. Kellogg Co., New York, K. Y.
ACKNOWLEDGMENT
RECEIVED for review September 28, 1955. Accepted June 21, 1956, This paper represents a part of the work done under Contract No. AF 18(600)-53 with the Air Research and Development Command, Wright Air Development Center, Wright-Patterson Air Force Base, Ohio.
Ruby James made the bulk of the measurements on which this paper is based, and most of the analyses. Walter B. Wade
LITERATURE CITED
(1) Dalmon, R., Freymann, R., Compt. rend. 211,472 (1940). (2) Ellis, J. W., Phys. Rev. 38, 693 (1931). (3) Gillespie, R. J . , Hughes, E. D., Ingold, C. K., J. Chem. SOC. 1950, 2552. (4) Goulden. J. D. S., Millen, D. J.,Ibid., 2620. (5) Ingold, C. K., hlillen, D. J., Ibid.. 2612. (6) “International Critical Tables,” vol. 111, p. 133, XlcGraw-Hill,
New York. 1933. (7) Kinsey,W. L., Ellis, J. W., Phys. Rev. 36, 603 (1930). (8) Ibid., 51, 1074 (1937).
(9) Lynn, S., Mason, D. M., Sage, B. H., Ind. Eng. Chem. 46, 1953 (1954).
Spectrophotometric Determination of Potassium with Sodium Tetraphenylborate RONALD T. PFLAUM and LESTER C. HOWICK Department o f Chemistry, State University o f lowa, lowa City, lowa
Potassium tetraphenylborate was investigated spectrophotometrically in an acetonitrile-water system. The tetraphenylborate ion shows absorption maxima at 266 and 274 m,u, with molar absorptivities of 3225 and 2100, respectively. Beer’s law is obeyed over a concentration range from 5 X 10-6 t o 7.5 X 10-1 M . The results obtained on the determination of potassium in selected samples indicate the feasibility of employing the described system for such determinations. Spectrophotometric evaluations of the solubilities of tetraphenylborate salts in aqueous solutions were obtained.
Table I.
Summary of Potassium Determinations
Table 11.
ITHIX the past 5 years, sodium tetraphenylborate has come into prominence as a precipitant for potassium. Various methods for the determination of potassium based on the insolubility of the potassium salt in aqueous solution and its solubility in certain organic solvents have been advanced. Gravimetric (2, 7 ) , titrimetric (6, 9 ) , turbidimetric (8),conductometric ( 7 ) , and voltammetric ( 1 ) methods have been proposed. A spectrophotometric method is a logical and useful extension to this existing list of measurements. Tetraphenylborate salts are soluble in certain organic solvents. Dissolution in acetonitrile leads to a solvent medium that is especially well suited for spectrophotometric measurement. This work is concerned primarily with an investigation of the potassium salt in such an acetonitrile medium. I t was undertaken in order to elucidate the feasibility of a spectrophotometric determination of potassium. APPARATUS AND REAGENTS
All spectrophotometric measyements were made at room temperature (approximately 25 C.) with a Cary Model 11
Potassium, Mg. Present Found
Sample
Solubilities of Tetraphenylborate Salts in Pure Water at 25” C. Salt “4
cs
K Rb T1
Solubility, M X 106 Experimental Literature 10.7 2.79 17.8 2.33 5.29
3 . 2 8 (20’ C.) (4) 1 8 . 2 (IO) 4 . 4 1 (20’ C.) ( 4 ) 2 . 9 (5)
recording spectrophotometer, using 1-cm. matched silica cells. A Beckman Model G pH meter was used for all pH values. Sodium tetraphenylborate was obtained from the J. T. Baker Chemical Co. It was used as received for all work except for obtaining the absorption curve of the reagent. I n this case, reagent recrystallized from an acetone-hexane mixture was used. Crystalline salts of ammonia, cesium, potassium, rubidium, and thallium(1) were prepared by reaction of the respective chlorides with the reagent in aqueous solution. Recrystallization of the precipitated material was effected from an acetonitrile-water system. Acetonitrile was obtained from the Matheson, Coleman & Bell Division of the Matheson Co. Purification was effected by treating with cold saturated potassium hydroxide, drying over
