Table I. Determination of PG in the Presence of BHA, BHT, and TBHQa Actual, Pg/mI
Sample 1 Sample 2 Average Average error, %
0.52 0.52 0.52
...
Calibration curve, Wml
Standard addition, !Jg/ml
0.55 0.57 0.56 7.7%
0.52 0.55 0.53 1.9%
a BHA, BHT, and TBHQ at a concentration of 0.5 pg/ml.
tion was made of the analytical potential to determine PG in the presence of BHA, BHT, and TBHQ using chloroform as a solvent. Table I compares the results of the determination of PG in the presence of the other three antioxidants in which the calibration curve method and standard addition method were used. The results indicate the standard addition method is more accurate. The ratio of the slopes of the calibration curve to the standard addition curve is 0.97. This indicates that there is little influence on the fluorescence intensity of P G with BHA, B H T , and TBHQ present. A more detailed study would have to be undertaken to determine the effect of different amounts of BHA, BHT, and TBHQ on the fluorescence intensity of PG. However, previous results using the standard addition method show that PG can be determined very accurately in the presence of BHA and B H T over a wide concentration range of BHA and B H T (2). Also, chloroform solutions of TBHQ varying in concentration from 0.50-2.0 pg/ml showed no fluorescence a t the maximum excitation and emission wavelengths for P G a t the sensitivity of the instrument used to determine PG. For the same T B H Q solutions, a very slight fluorescence appeared a t the greatest sensitivity setting of the instrument that decreased with in-
creasing concentration of TBHQ. This sensitivity setting was approximately 30 times that used to determine PG. I t was not possible without further purification of TBHQ to decide if the fluorescence was from an impurity in TBHQ or fluorescence due to TBHQ. However, the results in Table I indicate that this fluorescence is not significant a t the 0.5 pg/ml level. Also, the calculated relative fluorescence intensity of 0.5 pglml of PG in chloroform based on the highest sensitivity setting of the instrument is 653. The relative fluorescence intensity of 0.5 pg/ml TBHQ in chloroform is 3 a t the highest sensitivity setting of the instrument. Thus, the relative fluorescence intensity of P G in chloroform is 218 times that of a TBHQ chloroform solution of equal concentration indicating the contribution of TBHQ chloroform solution to the total fluorescence is minimal. Work is continuing on the potential of chloroform to quench the fluorescence of other phenols. The analytical possibilities of selective quenching are great. In this work, P G was determined accurately a t the ppm level in the presence of three other phenols. No separation steps were needed. In the separation and analysis of complex mixtures containing phenols, chloroform might be used to selectively quench the fluorescence of a large number of phenols or certain phenols with special structure features. An approach of this type will aid in the identification of phenols or eliminate fluorescence interference from phenols.
LITERATURE CITED (1)C. A. Parker, "Photoluminescence of Solutions", Elsevier Publishing Company, New York, N.Y., 1968. (2)H. W. Latz and R. J. Hurtubise, J. Agric. Food Cbem., 17, 352 (1969). (3)E. Sawicki, T. W. Stanley, and W. C. Elbert, Taianta, 11, 1433 (1964). (4)J. B. F. Lloyd, Ana/yst(London),99, 729 (1974). (5) M. Zander, Z.Anal. Cbem., 263, 19 (1973). (6)fed. Regist., 37, 25356 (1972). (7)J. D. Winefordner and M. Tin, Anal. Cbim. Acta, 31, 239 (1964).
RECEIVEDfor review July 25, 1975. Accepted September 11, 1975.
Spectrophotometric Determination of Iodide and Iodine at Parts-per-Million Level as a Thiocyanate Complex Bruce R. Clark' Analytical Chemistry Division, Oak Ridge National Laboratory, Oak Ridge, Tenn. 37830
Douglas A. Skoog Department of Chemistry, Stanford University, Stanford, Calif. 94305
A simple method to determine iodide or iodine in aqueous solution is afforded by the generation and photometric measurement of an iodine thiocyanate complex, I(SCN)z-, in the presence of excess hydrogen peroxide. Because other halides do not produce serious interferences when present in excesses of several thousandfold, separation from these anions can be avoided in many natural samples. Numerous methods for the determination of aqueous iodide and iodine solutions have been summarized in recent reviews (1-3). A number of methods prescribe some preliminary separation scheme, e.g., solvent extraction (4, 5 ) . Author to whom correspondence should be sent. 2458
Many modifications and improvements have been made on the method first reported by Sandell and Kolthoff (6) in which the rate of the iodide catalyzed reduction of Ce4+ by As3+ is dependent on the iodide concentration. This method remains one of the most accurate and sensitive and has few interferences, but standardized solutions and careful rate determinations are necessary. A few methods have shown a high sensitivity for iodide in the presence of large excesses of chloride. Schnepfe (7) has determined iodide as well as iodate in sea water in the submicrogram range, but this procedure is quite lengthy. A recent paper (8) describes a very simple procedure for the determination of iodine in the -1, 0, or +1 oxidation
ANALYTICAL CHEMISTRY, VOL. 47, NO. 14, DECEMBER 1975
Table 11. Effects of Several Extraneous Ions upon the Absorbance of I(SCN),- Solutionsa Excess
Table I. Absorbance of the I(SCN),- Complex as a Function of Iodide Concentration Concn I-, ppma
A, 302 n m
Std dev
Measured concn I-, ppmb
% ErrorC
0.378 0.123 ... 0.378 0.0 0.385 0.125 0.004 0.378 -1.8 0.385 0.123 0.385 0.123 0.385 0.125 0.385 0.116 -2.3 0.785 0.248 0.003 0.767 0.785 0.248 0.785 0.252 0.785 0.245 0.958 0.303 ... 0.936 -2.2 -1.8 ... 1.07 1.09 0.345 1.12 0.360 ... 1.11 -0.9 +3.1 1.27 0.425 ... 1.31 +0.7 0.001 1.36 1.35 0.439 1.35 0.438 1.35 0.440 1.35 0.439 ... 1.43 -0.7 1.43 0.463 1.61 0.528 ... 1.63 +1.2 1.71 0.545 0.002 1.68 -1.8 1.71 0.542 1.71 0.542 1.71 0.540 1.71 0.545 1.89 0.619 ... 1.91 +1.1 2.66 0.858 0.002 2.65 -0.4 2.66 0.858 2.66 0.860 2.66 0.862 a Other concentrations: H,O,, 1.0 x 10d4M; NaSCN, 0.205M; HClO,, 1.017M. Temperature: 8 "C. b Based upon E = 41100 1. mol-' cm-* and average absorbance for replicate measurements. (' Difference between nominal value and measured value divided by the nominal value. state a t the parts-per-billion level. Bromide, however, is a n interference with this method.
EXPERIMENTAL Apparatus. A Beckman Model DU Spectrophotometer was used for all measurements at a set wavelength. A Bausch and Lomb Spectronic 505 was used for scanning. Three matched 1.000-cm silica cells provided the capability for measuring two solutions and a blank in parallel. The Beckman DU cell compartment was surrounded with hollow blocks for the circulation of coolant from a constant temperature bath; cell temperatures could be maintained to within fO.l "C over a range of 2-25 "C. Reagents and Solutions. Reagent grade chemicals and deionized water were used t o prepare the solutions. Standard iodide solutions were prepared from sodium iodide dried overnight at 110 "C. Solutions were freshly prepared before each experiment using the dried salt as a primary standard. An assay indicated 99.98 f 0.13% purity (as NaI), justifying its use as a primary standard. Dilute solutions were prepared from the successive dilutions of 10-3M stock solution. Sodium thiocyanate contained iodine consuming impurities and was recrystallized from methanol. The purified salt required 0.08 ml of 0.01N triiodide to titrate 5.0 g of the solid salt. It should be noted that this purification step is probably not essential for the determination of iodine by the'method outlined in this work. Justification for this stems from the likelihood that sulfide is the iodine consuming impurity, reducing iodine to iodide, but this would not interfere since an excess of hydrogen peroxide is present and would oxidize all the iodide and iodine to form the complex, I(SCN)2-. Hydrogen peroxide solutions were freshly prepared each day from 28-32% commercial stock solution. In studies requiring accurate values of concentrations, the hydrogen peroxide solutions were standardized vs. a standard potassium permanganate solution. Perchloric acid solutions were prepared from the dilution of
Ion
Fc1-
c1-
(approximate)b
2x 1x 3x 1x 3x
% A A , 302 nrn - 10.0 + 2.0 + 2.0
104 104 104 104
- 4.0 104 - 6.8 1 x 104 - 1.7 so423 x 104 - 5.1 s0,z3 x 104 - 2.5 H,P04 + 6.0 Acetate 1 x 104 i x 104 - 2.3 Citrate (c) i x 104 + 2.8 Ca2+ 100 +315 Fe3+ 50 + 19.3 co2+ ( c ) 500 +160 co2+ ( c ) 100 + 50.5 cuz+ ( c ) 1 x 103 - 4.5 ZnZ+( c ) 1 x 103 - 4.5 Cdzf ( c ) a Concentrations: H,O,, 1.0 x 10-4M; NaSCN, 0.205M, HClO,, 1.017M; NaI, 8.61 X 10-6M. Temperature: 8 "C. b Excess of nominal iodide concentration. C NaI concentration: 1.060 x 10-5M. BrBr-
commercial 60% acid solutions. Standardization with primary standard grade sodium carbonate was done for these studies, but is not necessary for the application of this method. Salts used for interference studies were: NazS04, NazHP04, NaC1, NaBr, NaF, FeNH.@04)~.12H20, CaC12, NaN03, CdC12, CuSO4, ZnClz, NaC2H302 (acetate), and Na~C&07.2M20 (citrate). Analytical Procedure. Stock solutions of 5 X 10-4M hydrogen peroxide, 1.OM sodium thiocyanate, and 5M perchloric acid are cooled in an ice slurry. To a 25-ml volumetric flask, add 5.0 ml each of sodium thiocyanate, perchloric acid, and hydrogen peroxide followed by an amount of cooled sample containing sufficient iodide or iodine to produce an absorbance between 0.1-0.9 (0.3-3.0 ppm I-). At the same time, a blank solution is prepared in the same manner. Dilute the flasks to volume with deionized ice water and fill the cells which have been precooled in the cell compartment to the thermostated temperature (0-10"C). Measure the absorbance at 302 nm vs. the blank solution. Maximum absorbance is reached in about 20 minutes and remains constant for several hours. Use L = 4.11 X lo4 1. mol-' cm-' to compute the concentration as iodide.
RESULTS AND DISCUSSION Aqueous solutions of iodine and thiocyanate form a rather stable charge transfer complex, I2SCN (9). In the presence of hydrogen peroxide, acidic solutions of iodide or iodine and thiocyanate produce the complex, I(SCN)2-, with the stoichiometries ( I O ) :
+ I- + PSCN- + 2H+ Hz02 + 1 2 + 4SCN- + 2H+ Hz02
+ 2H20 2I(SCN)2- + 2H20
-+
+
I(SCN)2-
(1) (2)
Several studies were performed to ascertain the optimum physical and chemical conditions for quantitatively generating a stable complex. Earlier data ( I O ) indicate that I(SCN)2- is unstable at room temperature and no manipulation of chemical compositions in these studies was successful in improving the stability. Absorbances reached a maximum after one or two minutes and diminished by 50% in about one hour. T h e decomposition reaction is presumably the oxidation of thiocyanate t o sulfate and other nonabsorbing products ( I O ) . A t lower temperatures (0-10 "C), solutions were prepared which exhibited a constant absorption for several hours. Various solution compositions were studied at 8 "C. Too large a n excess of hydrogen peroxide (a few millimolar) caused decomposition of the complex, but solutions 0.1-
ANALYTICAL CHEMISTRY, VOL. 47, NO. 14, DECEMBER 1975
2459
0.2mM were stable and sufficiently concentrated to give quantitative results. The thiocyanate concentration was found to be important in driving the reaction to completion and affected the rate of complex formation. Low absorbances observed a t concentrations less than about 0.2M might, in part, result from the decomposition of I(SCN)zto form the non-absorbing ISCN (IO). Hydrogen ion concentrations were about 1M in all cases. Lower concentrations caused a slight decrease in stability and higher concentrations (above 2-3M) led to the decomposition of thiocyanic acid (9). Absorbance measurements on thirty separately prepared solutions are shown in Table I. Precision of replicate measurements is good and the overall accuracy has a range of about f 3 % . A linear least squares fit to the data gives a slope of 41100 f 200 1. mol-' cm-l and a y-intercept of -0.003. Possible interferences from extraneous ions were examined. Table I1 summarizes these observations. In each case, two solutions of identical iodide concentrations were compared with an extraneous ion added to one. These were carried through the procedure in parallel. Clearly, chloride presents no serious interference a t the levels examined and bromide lowers the absorbance only a few percent. The fact that the chloride interference does not change with chloride concentration suggests that the 2% difference in absorbance may be a measurement error rather than an interference. Sulfate, a frequent interference in iodide analyses, causes a small decrease in absorbance. The transition metal ions interfered to varying extents. Since thiocyanate complexes form with these cations, interferences are to be expected.
Nitrate ion cannot be tolerated much beyond trace levels. I t has an absorption maximum near 295 nm in this system. In the application of this method to natural samples, pretreatment with a cation exchange resin could remove interfering cations to acceptable levels. For accurate analyses of iodide in the presence of large fluoride or bromide excesses, a standard addition technique with a series of solutions would not be too time consuming. Although the procedure described here requires a thermostated cell compartment, an alternate approach would be to keep the solutions in an ice slurry after mixing until the reaction is complete (about one hour a t 0 OC). Rapid measurement in a room temperature cell compartment should not affect the results significantly. Fogging of the cells could be prevented by purging the cell compartment with dry nitrogen.
LITERATURE CITED (1) (2) (3) (4) (5) (6) (7) (8) (9) (10)
M. J. Fishrnan and B. P. Robinson, Anal. Chem., 41, 323R (1969). M. J. Fishman and D. E. Erdmann. Anal. Chem., 43, 356R (1971). M. J. Fishrnan and D. E. Erdmann, Anal. Chem., 45, 361R (1973). T. C. Chadwick. Anal. Chem., 47, 933 (1975). 0. A. Tataev and M. S. Bezhaev, Sb. Nach. Soobshch., Dagestan. Univ., Kafedra Khim.. 1967, 79-82; Chem. Abstr., 69, 21821h (1968). E. E. Sandell and I. M. KoRhoff, J. Am. Chem. Soc.. 56, 1462 (1934). M. M. Schnepfe, Anal. Chim. Acta, 58, 83 (1972). J. L. Lambert, G. L. Hatch, and B. Mosier, Anal. Chem., 47, 915 (1975). C. Lewis and D. A. Skoog, J. Am. Chem. Soc., 84, 1101 (1962). C. Long andD. A. Skoog, horg. Chem., 5, 206 (1966).
RECEIVEDfor review June 20, 1975. Accepted August 14, 1975. Oak Ridge National Laboratory is operated by the Union Carbide Corporation for the U.S. Energy Research and Development Administration.
Filter Sampling Method for Atmospheric Sulfur Dioxide at Background Concentrations Herman D. Axelrod and Steven G. Hansen' National Center for Atmospheric Research, Boulder, Colo. 80303
The traditional wet chemical technique for atmospheric sulfur dioxide measurements was developed by West and Gaeke ( I ) . This method uses a bubbler trapping system containing a tetrachloromercurate(I1) (HgC12 2NaC1) solution. The analysis is spectrophotometric, involving the dye pararosaniline. Further refinements have been made by Scaringelli e t al. ( 2 ) . The West-Gaeke procedure has been successfully used to measure concentrations as low as 5 ppbv (10.5 fig/m3 a t 600 Torr). However, this level is still well above background or base-line concentrations estimated to be in the 0.1-1.0 ppbv range. The main limitation of using the WestGaeke procedure for very low levels is the inability to sample large quantities of air with a bubbler or impinger because of sample solution evaporation and flow rate limitations. Therefore, if an improved sampling method can be developed, measurements of SO2 a t base-line levels can be accomplished.
+
Present address, Department of Chemistry, University of California, Berkeley, Calif. 2460
An alternative to bubbler sampling techniques is the use of impregnated filters. The basic advantages of this system have been previously reported ( 3 ) . Furthermore, Huygen ( 4 ) impregnated filter paper with KHC03, as well as KOH in combination with glycerol or triethanolamine, to capture SO*. However, this technique was not useful a t relative humidities below 30% nor was the technique used for baseline measurements. The work presented here demonstrates that Whatman No. 17 filter paper impregnated with tetrachloromercurate(I1) solution can be used to sample for SO2 a t levels as low as 0.05 ppbv a t sampling flow rates as high as 6.5 l./min for as long as several hours. The analysis is accomplished by a modification of the West-Gaeke procedure (1).
EXPERIMENTAL Apparatus. Spectrophotometric measurements were performed with a Coleman Model 124 double beam spectrophotometer at 560 nm. All measurements were made at 23 O C in an air conditioned
room.
ANALYTICAL CHEMISTRY, VOL. 47, NO. 14, DECEMBER 1975
