azine phosphate solutions, and [OH-] is, of course, even smaller. Hence, [BHz2+] = [A2-] =
[BH+] = [HA-I
=
ion (8),respectively, it can be shown that the value of [BH+]/ [BHzz+] and [HA-]/[A*] is 8.55 at 25" C. The total ionic strength of the solution is therefore 1.314m. The ratio is 5.59 at 0" C and 13.9 at 50" C. The equation [H+Iz = klkz predicts that the hydrogen ion concentration of the buffer solution should be dependent on the concentration of piperazine phosphate only because activity coefficients have been neglected. If activity coefficients are introduced and the thermodynamic dissociation constants (K) are used, it can be shown that (aH+)z =
[BH+]/@3H22+]= [HA-]/[A2-] = d k x
YBH~ YBH+YA* The symmetrical arrangement of the activity coefficients indicates that the change in paH value with concentration should be small.
for piperazinium ion Introducing the kl value (4.64 X (7) and the kz value (6.34 X lo-*) for dihydrogen phosphate
RECEIVED for review October 30, 1967. Accepted December 13, 1967.
(7) H. B. Hetzer, R. A. Robinson, and R. G. Bates, J. Phys. Chem.,
(8) R. G. Bates and S . F. Acree, J. Res. Natl. Bur. Std., 30, 129
[H+] = d k i k z and
in press.
(1943).
Infrared Determination of Trace Amounts of Polyatomic Inorganic Ions James R. Lawson and Ralph L. Barnett, Jr.' Department of Physics,
Fisk University, Nashville, Tenn.
COMPILATIONS of the spectra of inorganic compounds in the 2- to 15-micron region ( I ) and the 15- to 36-micron region (2) have demonstrated the feasibility of qualitative analysis by infrared spectrometry. Although not competitive with the higher accuracy and precision of standard gravimetric techniques, infrared methods offer rapid quantitative determinations with an accuracy to roughly 1%. Such determinations are of particular value at very low concentrations of sample. The contour and intensity of the absorption bands in the spectrum of an inorganic ion or compound are functions of the temperature and of the environment. It is possible to change appreciably the environment, and thus the absorption characteristics, by putting the ion or molecule into a solid solution. The advantages of this technique were first shown by Hiebert and Hornig in their study of HC1 (3), and have been amply illustrated at low temperatures by more recent papers (4). Large ions, or polyvalent ions, may not form solid solutions. In these cases the solute is uniformly distributed as very small crystals in the solute matrix. It has been observed (5) that an increased absorption results from smaller crystal size. The following characteristics indicate the existence of a solid solution rather than a mixture of host and solute compounds: frequency shifts are observed for the solute ion in 1 Present address, Florissant Valley Community College, Ferguson, Mo.
(1) F. A. Miller and C . H. Wilkins, ANAL.CHEM., 24, 1253 (1952). (2) F. A. Miller, G. L. Carlson, F. F. Bentley, and W. H. Jones, Spectrochim. Acta, 16, 135 (1960). (3) G. L. Hiebert and D. F. Hornig, J. Ckem. Pkys., 27, 752 (1957). (4) G . C . Pimentel, Spectrochim. Acta, 12, 94 (1958). ( 5 ) J. Bonhomme, Ibid., 7 , 32 (1955).
636
ANALYTICAL CHEMISTRY
different solvents ; selection rules arising from the existence of the unit cell in the pure solute crystal no longer hold in solid solution; and at high concentrations of solute, either the solubility will be exceeded and a spectrum characteristic of the pure solute will be obtained, or deviations from Beer's law due to the solute-solute interaction will be observed. Solid solutions in alkali halide crystals are stable at room temperature. They may be formed by the process of adding a small quantity of the appropriate inorganic salt to a molten alkali halide and cooling to form crystals, or a single crystal, in which the solute ions are trapped in the host lattice (6, 7). The use of molten alkali halides requires temperatures of 700" to 800" C, at which many ionic species decompose or react with traces of oxygen. Thus, cyanide becomes cyanate ion and nitrate partially decomposes to the nitrite ion. Solid solutions have also been produced by heating a mixture of salts under pressure (8), and by warming disks pressed from mixtures (9), relying upon diffusion in the solid state. These methods allow observation of species not stable at higher temperatures. In the majority of observed spectra, the fusion technique led to appreciable increases in the absorption peak heights of certain fundamental bands. This indicated that the sensitivity of inorganic infrared analysis might be increased. Three ions (OCN-, Sod2-,and COa2- were chosen for study in KBr, KC1, and NaC1. These represent neither the best nor (6) A. Maki and J. C . Decius, J. Chem. Pkys., 31, 772 (1959). (7) H. W. Morgan, Symposium on Molecular Structure and
Spectroscopy, Ohio State University, June 1957. (8) J. A. A. Ketelaar, and J. van der Elsken, J. Ckem. Phys., 30, 336 (1959). (9) W. C. Price, W. F. Sherman, and G. R. Wilkinson, Proc. Royal SOC.(London),,4255, 5 (1960).
1600
1400
I900
WAVENUMBER, CM -1 1000 666
(00 500
400
90
80 70
Y z
P f
$0 50
Z
2 6
7
0
9 10 11 15 WAVELENGTH, MICRONS
PO
40
95
30
Figure 1. Infrared spectra of sod2A . 0.05 mg B. 1 mg K2SO4fused in 400 mg KBr C. 1 mg &SO4 ground
20 (0
- --.
I
GROUND I
I
'
the worst to be expected in sensitivity, and were studied only to illustrate the accuracy, precision, and increase in sensitivity of a typical inorganic analysis. EXPERIMENTAL TECHNIQUES
Sample Preparation. Milligram quantities of solute were weighed on a Cahn electro-balance, Model M-10, and mixed with 1-gram quantities of the alkali halide solvents. Lesser concentrations were obtained by dilution. The mixed solute and solvent were fused, and quenched by pouring on the clean surface of a mullite mortar. The crystallized solid was ground by hand and 300-mg quantities were taken for the preparation of disks. A 0.5-inch-diameter vacuum die was used, the disks being formed under pressure of ca. 150,000 psi applied for 5 minutes. The purity of each alkali halide was checked by preparing a disk from the pure salt and examining its infrared spectrum in the regions of analytical interest. Solute concentrations varied from approximately 0.13 pmole to 12.5 pnole per gram of solvent. KBr, KC1, and NaCl were used as solvents. Solid solutions of cyanate ion were obtained by heating KCN and the appropriate solvent, conversion to the cyanate by atmospheric oxygen being quantitative at the fusion temperature. Samples containing the sulfate and carbonate ions were made by melting weighed amounts of K2SO4,K2C03,and C a C 0 3with the alkali halides. A series of disks were prepared without fusion to obtain peak absorbances for comparison. Weighed amounts of the solutes were mixed with each of the three solvents, ground with mortar and pestle for 12 minutes, and pressed as described above. These disks are referred to as "ground samples" in this paper. Samples containing the cyanate ion were prepared with KOCN. Measurements. All infrared spectra were recorded on a Beckman IR-5 spectrophotometer, at NaCl resolution. Spectra of the sulfate ion were recorded from 7 to 11 microns, with a mechanical slit width of 0.400 mm; spectra of the carbonate ion from 5 to 9 microns with a 0.245-mm slit; and the spectra of the cyanate ion from 4 to 10 microns with a 0.180-mm slit. All spectra were recorded at a speed of 0.5 micron per minute, Bands chosen for measurement of peak heights were the S042- degenerate stretching vibration at 9 microns, the COS" asymmetric stretching mode at 7 microns, and the C=N stretching vibration at 4.6 microns. The base line technique was used in determining peak heights. The pressed disks were weighed, and the absorbance data corrected to disks containing 300 mg of solvent. These values were then
plotted against the solute ion concentration, expressed as micromoles of solute per gram of solvent, to provide analytical working curves. RESULTS
The fused sample spectra indicate that the cyanate ion is in solid solution. Frequencies of 2169 cm-l, 2184 cm-1, and -4 cm-l) were observed in KBr, KCl, and 2208 cm-l(each NaCl lattices. The peak intensities also show a marked dependence upon the lattice. It is assumed that the sulfate and carbonate ions exist as small crystals distributed through the
+
Table I. Ion Concentrations for Peak Absorbance of 0.16 WaveConcentration, length, pmoles/gram Ion Compound microns Matrix Ground Fused 7.9 2.0 OCNKOCN 4.6 KBr 7.4 3.0 KCl NaCI 10.0 2.8 2.1 0.9 8.9 KBr SO *SrSOa 1.6 0.9 KCl 1.4 1.1 NaCl 8.5 1.6 7.0 KBr COa2- Na2C03 4.1 2.6 KCl 3.0 5.4 NaCl Table 11. Limits of Detectability in Solid Solutions Minimum detectable quantity (pg)
Ion OCN-
so4 'coa 2-
Compound KOCN
KBr 6
4 BaS04 SrSOr NazCOa CaCOa
VOL. 40,
8 7 7 8 19
NO. 3,
Solvents) KC1 9 4 11
7 11 17 24
MARCH 1968
NaCl 9 8 14 8 22 28 26
637
Table 111. Precision of the Solid Solution Method Strontium sulfate concentration: 0.200 mg in 300 mg KBr
Group
Mean absorbance
Group range
Mean deviation
Range
46) B(6) c(6)
0.688 0.697 0.657
0.010
0.6% 1.2% 1. O %
0.060
0.023 0.016
Combined summary Mean Mean deviation 0.680
253Z
Sodium carbonate concentration: 0.200 mg in 300 mg KBr
Group 46) B(6) C(6)
Mean absorbance 0.628 0.645 0.650
Group range 0.025
0.029 0.010
Mean deviation 0.4%
Range 0.036
Combined summary Mean Mean deviation 0.642 1.7z
1 . 7 ~
0.6Z
Potassium cyanate concentration: 0.150 mg in 300 mg KBr
Group
Mean absorbance
Group range
Mean deviation
Range
46) B(6) C(6)
0.541
0.021 0.017 0.010
1.3%
0.053
0.570
0.578
alkali halide. The concentrations required to produce a peak absorbance of 0.16 with a 300-mg disk indicate the increased sensitivity of the fusion method, and are given in Table I. The peak intensities for the cyanate ion in solid solution are dependent upon the spectrometer slit width, inasmuch as the band has a width of only a few wavenumbers. The bands observed were identical to those reported by Maki and Decius (6). Beer’s law plots were made for each of the seven solutes, ground and fused with the three alkali halides. We define the minimum detectable concentration as that sample giving a peak absorbance of 0.02, and thus a 10: 1 signal to noise ratio. The values determined for fused samples are given in Table 11. In every case, maximum sensitivity is obtained with KBr as solvent. For the carbonate compounds, there exist cases in which fusion decreases the intensity of the band at 9 microns. The minimum detectable quantity of cyanate ion is 6 pg, using a 0.180-mm slit. With scale expansion or higher spectral resolution, one tenth of this amount can be easily detected. Both ground and fused sample spectra of K S 0 4 are shown in Figure 1. The strong absorption at 9 microns is the asymmetric stretching mode, while the less intense band near 10 microns which appears in the fused sample spectrum is the symmetric stretching mode. Under the criteria described, the limit of detectability for the sulfate ion is 4 pg. The spectra of ground and fused Na2C03in KBr are shown in Figure 2. For the carbonate ion, 7 pg is determined to be the minimum detectable sample. The precision of this fusion method for very low solute concentration depends upon several factors, including the uniform distribution of solute through the solvent, the photometric accuracy and stability of the spectrometer, and errors in weighing. The reproducibility was studied to estimate the reliability of a typical analytical application. The results are summarized in Table 111. Samples of each of the three solutes in KBr were prepared at three different times, indicated by A, B, and C. Each sample, after fusion and grinding, was divided into six 300-mg portions. Disks were prepared and the reproducibility of the absorbance values determined. 638
ANALYTICAL CHEMISTRY
Combined summary Mean Mean deviation 0.563
2.7Z
1.3%
0.6%
COMMENTS
Several anions may be quantitatively determined in one sample, with certain precautions about the cations present. The effects of different cations are shown in Table 11. The cations definitely affect the environment of the anion and produce different peak absorbances. As an example, KzCO3 fused in KC1, to which a small quantity of CaClz has been added, shows an absorbance identical to that obtained by fusing CaC03in KBr. Recrystallization at the midplane of an alkali halide pressed disk, one of the principal factors in the decrease of transmission with time, produces changes in the spectra of fused samples. The peak absorbances decrease slowly to values between those for ground and fused samples. All measurements reported here were made within two hours after formation of the disks. In normal analytical procedures such measurements are made immediately, and slow variations are not a disadvantage. The increase in peak absorbance obtained through the formation of solid solutions has not been generally recognized. Reagent grade alkali halides, if fused and pressed into disks, will frequently show weak sharp absorptions from ionic impurities, not observable unless a solid solution is formed. Such bands have been found also in alkali halide windows supplied for the infrared region (IO). As mentioned previously, grinding and mixing under certain conditions may lead to the formation of solid solutions. The spectra reported by Milkey (11) as arising from KBr upon lengthy grinding in a mortar and pestle, are due to CN-, COa2- and S04+ ions, presumably impurities in the KBr. Such an effect indicates that highly reproducible grinding techniques are important where KSCN (12), PbS04 (13), and NaN3 (14) are employed as internal standards in analyses by the pressed disk technique. (10) R. S. McDonald, Symposium on Molecular Structure and
Spectroscopy, Ohio State University, June 1959. 30, 1931 (1958). (11) R. G. Milkey, ANAL.CHEM., (12) S. E. Wiberley, J. W. Sprague, and J. R. Campbell, Ibid., 29, 210 (1957). (13) N. Wright, Appl. Spectry, 9, 105 (1955). 31, 1602 (1959). (14) R. T. M. Fraser, ANAL.CHEM.,
The technique of freeze-drying has been applied to the determination of the SO4*- ion by Tai and Underwood (15). From the spectra obtained from nine different sulfates they chose KzS04as the most suitable for quantitative measurements. The band contour obtained by freeze-drying KBr and K2SOris similar to that obtained by the fusion technique, and (15) H. Tai and A. L. Underwood, ANAL.CHEM., 29,1430 (1957).
the Beer's law plots obtained by the two methods are almost identical. Though more time consuming, freeze-drying should have application for inorganic ions which decompose at the temperatures required for the fusion method. RECEIVED for review February 28, 1966. Resubmitted and accepted November 28, 1967. Work supported by U. S . Atomic Energy Commission Grant N. AT-(40-1)-2760.
Spectrophotometric Determination of the Perchlorate Ion N. L. Trautweinl and J. C. Guyon Department of Chemistry, University of Missouri, Columbia,Mo.
SIMPLE,direct spectrophotometric methods for perchlorate ions do not exist. Present methods involve a precipitation or an extraction step prior to measurement. Nobar and Ramachandran ( I ) precipitated perchlorate 'with excess methylene blue, filtered the precipitate, and determined the excess methylene blue. Swasaki, Utsumi, and Kang (2) eliminated the problem of the solubility of the precipitate in the methylene blue procedure by extraction of the methylene blue-perchlorate complex into 1,2-dichloroethane. Other precipitation methods include the precipitation of perchlorate by ferroin (3) and salts of copper tetrapyridine complex (4). Golosnitskaya and Petraschen (5, 6) complexed perchlorate ion with brilliant green and malachite green and extracted the complexes. Zatko determined perchlorate by the oxidation of VSOl to (VO)*+, and measurement of the (VO)2+ formed (7). Fritz, Abbink, and Campbell (8) extracted a ferroinperchlorate complex with n-butyronitrile. This paper presents a novel approach to the determination of the perchlorate ion. The resulting method is simple, selective, and has good color stability. The sensitivity of the technique is comparable to that of the ferroin method (8) and superior to that of other techniques. The approach presented here has a simple procedure, involving no precipitation or extraction steps. EXPERIMENTAL
Apparatus. All spectral measurements were made on a Cary Model 12 automatic recording spectrophotometer using matched 1.000 9 0.002 cm fused quartz cells. Temperatures were controlled to 90.3" C using a water bath with a Sargent Model 3554 thermoregulating unit. 1 Present address, Department of Chemistry, Arkansas State University, State University, Ark.
(1) G. M. Nobar and C. R. Ramachandran, ANAL. CHEM.,31, 263 (1959). (2) I. Swasaki, S. Utsumi, and C. Kang, Bull. Chem. SOC.Japan, 36, 325 (1963). (3) Z. Greaorowicz, F. Buhl, and Z . Klima, Mikrochim. Ichnoanal. Acta, 1963, p 116. (4) W. Bodenheimer and H. Weiler. ANAL.CHEM., 27.1293 (19551. ( 5 ) V. A. Golosnitskaya and V. I. Petraschen, Zh. Analif. Khim,, 17, 878 (1962); A n d . Absf., 10, 2264 (1963). (6) V. A. Golosnitskaya and V. I. Petraschen, Tr. Nouocherk. Politekhn. Znsf., 141, 73 (1964); CA, 64, 1351e (1966). (7) D. A. Zatko, ANAL.CHEM., 37, 1560 (1965). (8) J. Fritz, J. E. Abbink, and P. A. Campbell, Zbid., 36, 2123 (1964).
1.200
0.700
'! 4"
0.600
r
1
t
J
2
0.500
0.400 0.300 0,200
o.ino 0
1
2
3
4
5
6
8
i
nil of 100.0 g/l SnC12
9
10
11
12
. 2H-0
Figure 1. Effect of SnCI, concentration (1) No added Clod(2) 10 ppm Clod(3) A absorbance, 1-2
Reagents. All reagents used were of the highest quality obtainable. The a-furildioxime was purchased from Eastman Kodak Co. and used without further purification, A solution of this ligand was prepared by dissolving 7.0000 grams in methanol and diluting to 1 liter with methanol. The potassium perrhenate was obtained from the University of Tennessee and used without further purification. Recommended General Procedure. Prepare a calibration curve by adding up to 0.75 mg of perchlorate ion to 50-ml volumetric flasks and bringing to a total volume of 6 ml with deionized water. To each flask add 5.00 ml of aqueous 0.0950 gram/liter potassium perrhenate solution, 4.70 ml of 1 :1 hydrochloric acid, 17.00 ml of methanol, and 7.00 ml of 8 . 5 x tin(I1) chloride. Mix and transfer each flask to a 55" C water bath for 30 minutes. Add 5.00 ml of the afurildioxime solution, dilute to 50 ml with deionized water, mix well, and return to the 55" C water bath. After exactly 30 minutes, read the absorbance at 532 mp. The calibration curve of log absorbance us. concentration is linear within experimental error. VOL. 40, NO. 3, MARCH 1968
639
