Spectrophotometric Determination of Tetrathionate OSCAR A. NIETZEL
and
MICHAEL A. DESESA
Raw Materials Development Laboratory, National Lead Co., Inc., Winchester, Mass.
The spectrophotometric method of Robinson for the determination of tetrathionate has been modified to yield greater sensitivity and accuracy. The procedure depends on the production of thiocyanate equivalent to the tetrathionate and determination of thiocyanate with an excess of ferric iron. By developing the color in opaque cylinders, the rapid decomposition of the ferric thiocyanate color is eliminated. The use of ferric nitrate and nitric acid instead of ferric chloride and hydrochloric acid has decreased the absorbance of the reagent blanks. By measuring the final color at 460 mp instead of at 525 mp a twofold increase in sensitivity has been achieved.
T
HE A4nalyticalDepartment of this laboratory was presented with the problem of determining concentrations of tetrathionate ion as low as 0.005 gram per liter in the presence of amounts of sulfate on the order of 30 grams per liter. A literature survey was conducted for methods of determining tetrathionate. The published methods for macro amounts of polythionates have recently been reviewed by Jay ( Z ) , but none of these procedures approached the required sensitivity. Murayama ( 4 ) has reported a \vel1 defined polarographic reduction wave for tetrathionate ion in a supporting electrolyte of 0 . 1 S potassium chloride and 0.1-Ybarium chloride, and Furnees and Davies ( 1 ) reported a similar wave in phosphate buffers. A spectrophotometric method proposed by Robinson for the determination of low concentrations of tetrathionate ion (6) is based on the formation of thiocyanate from tetrathionate (and higher polythionates) by reaction with cyanide in an alkaline medium, and the subsequent formation of a red ferric thiocyanate complex with excess ferric chloride. Because of the need for a very sensitive procedure, this investigation was confined to an evaluation of possible polarographic and spectrophotometric techniques based on the methods mentioned above. The polarographic approach was abandoned after preliminary experiments indicated that the desired degree of sensitivity could not be obtained. An investigation of Robinson's spectrophotometric method proved to be productive. The present paper reports the results of this investigation and a modified spectrophotometric procedure. APPARATUS AND REAGENTS
Beckman Model D U spectrophotometer x i t h 1- and 10-cm. Corex cells. The bodies of the 10-cm. cells were covered with Scotch black electrical tape. One hundred-milliliter mixing cylinders coated with Krylon black acrylic spray, except for small vertical areas on opposite sides near the top graduation. All chemicals used were of analytical reagent grade purity. Sodium tetrathionate dihydrate was prepared according to the procedure of von Klobukoff as given by Mellor ( 3 ) . This salt was analyzed for tetrathionate ion by the volumetric procedure of J a y ( Z ) , and was found to contain 73.2% tetrathionate, which is the theoretical value. Stock solutions were prepared as required, since aqueous solutions of sodium tetrathionate are unstable. Amberlite IR-120 cation exchange resin. Columns of 15 ml. of resin in 25-ml. burets were employed. The resin was converted to the hydrogen form by repeated washes with 5y0 sulfuric acid. The exhausted resin was regenerated by washing with 100 ml. of 5% oxalic acid followed by 200 ml. of 4 V hydrochloric acid. RECOMMENDED PROCEDURE
Pass an aliquot of the sample, which contains 0.05 to 5 mg. of tetrathionate ion in no more than 25 ml., through 15 ml. of the hydrogen form of Amberlite IR-120 resin. Collect the effluent in
a 100-ml. beaker. Wash the column with 15 ml. of water, collecting the washings in the beaker. In order to develop the ferric thiocyanate color, neutralize the effluent and washings from the column to phenolphthalein with 1 to 10 ammonium hydroxide. Add 5 ml. of 5% sodium cyanide solution. After 15 minutes, add 10 ml. of 1to 1 nitric acid (in a fume hood). After the evolution of hydrocyanic acid has ceased, transfer the solution quantitatively into a Krylon-coated 100-ml. mixing cylinder, add 5 ml. of 21M ferric nitrate, dilute the solution to 100 ml., and mix well. For samples known to contain thiocyanate prepare a sample blank. Pass another equal aliquot of the sample over a fresh resin bed, and wash in the same manner. I n order not 'to convert the tetrathionate to thiocyanate, add the 10 ml. of 1 to 1 nitric acid first and then the 5 ml. of 5% sodium cyanide. After the evolution of hydrocyanic acid has ceased, transfer the solution quantitatively into a Krylon-coated 100-ml. mixing cylinder, add 5 ml. of 2,11 ferric nitrate, dilute the solution to 100 ml., and mix well. Transfer a portion of the sample and sample blank to either 1-cm. or IO-cm. absorption cells. Measure the absorbances of the sample and sample blank a t 460 mp (0.020-mm. slit) against a reagent blank. Subtract the absorbance reading obtained with the sample blank from t h a t obtained for the sample. Determine the amount of tetrathionate ion in the aliquot by reference to a calibration curve or by calculation using the average absorbance index. EXPERIMENTAL
Preliminary Work. The spectrophotometric approach of Robinson ( 5 )using the reaction S406--
+ 3 C S - + H20 = &Os-- + SO4-- + 2HCN + CNS-
followed by the colorimetric determination of the thiocyanate with ferric iron was examined. Employing the procedure of Robinson, a calibration curve was prepared. Beer's law was obeyed over the range of 0.5 to a t least 5.0 mg. of tetrathionate ion per 100 ml., and the average absorbance index was 0.0844 per mg. of tetrathionate ion per 100 ml. in 1-cm. cells. For the analysis of solutions, Robinson ( 5 )recommends taking a 25-ml. aliquot. I n order to get an absorbance reading of about 0.1 the sample would have to contain 0.045 gram of tetrathionate ion per liter. As this sensitivity was 10 times less than desired, an attempt was made to obtain the necessary increase by using IO-cm. cells. With the more dilute solutions, straightline calibration curves were obtained, but the absorbance index varied from 0.024 to 0.06. It was thought that in the extremely dilute solutions of tetrathionate being studied the formation of thiocyanate from tetrathionate was not stoichiometric. Attempts were made to drive the reaction t o completion by varying the cyanide concentration, pH, reaction time, and temperature. I n no case was the same absorbance index obtained with 10-cm. cells as with 1-cm. cells. In a few instances, when the hydrolysis reaction was allowed to take place overnight, the absorbance index was higher than that obtained in 1-cm. cells. This was shown to be caused by a secondary reaction of cyanide on the thiosulfate produced in the first decomposition CS-
+ SZO3--
=
CNS-
+ SO3--
During one experiment, it was noticed that the absorbance of a sample which was left in the sample compartment of the spectrophotometer for 10 minutes remained constant while the absorbance of a sample left in daylight for the same time decreased to one half the value of the unexposed sample. I t then became apparent that the instability of the ferric thiocyanate color to light was the cause of the inconsistent absorbance indices obtained with the most dilute solutions rather than an incomplete hydrolysis with cyanide. This fading of the ferric thiocyanate
1839
,
1840
ANALYTICAL CHEMISTRY
color has been recognized (6) as one of the limitations of the colorimetric thiocyanate method for iron. However, the fading observed in this experiment was about five times greater than previously reported. The fading is believed to be due to a reaction between the nitric acid and thiocyanate, and is considered to be promoted by the presence of ferric ions. This theory is supported by the authors' observation of an increased fading with high ferric ion concentration and low thiocyanate concentration. Stabilization of Fsrric Thiocyanate Color. I n an attempt to inivnsify the color of the ferric thiocyanate complex the samples were preparcad in 50% acetone solutions, but no pink color was apparent, indicating that very little ferric thiocyanate complex had formed. I t mas felt that the acetone was causing the ferric chloride, uhich ~ a used s as a source of ferric ions, to be undissociated in solution. Therefore, ferric nitrate and nitric acid were employed in the procedure as the source of ferric ions and acid, respectively. \Then the nitrate reagents were used in 50% acetone, the average absorbance index per milligram of tetrathionate ion per 100 ml. was 0.093 with 1-em. cells and 0.088 with 10-cm. cells. There was still a decrease in absorbance in a sample elposed to light for 10 minutes before being read, but the rate of color fading [vas not so fast as in the aqueous-chloride system.
1
0.6
1
Table I.
Effect of Acetone on Sensitivity and Stability
(All samples contained 0.50 m g . of S406-- per 100 ml. T h e color was developed in opaque mixing cylinders, and the absorbance was measured a t 460 mp in I-em. cells)
Hours
Acetone Concn.,
%
50 25 0
0.25
0 50
0.114
0.114 0.105 0.097
0.105 0.097
Table 11.
'
400 420 440 460 480 500 520 540 WAVELENGTH, my Figure 1. Absorption spectrum of ferric thiocyanate complex obtained from 2.93 mg. of letrathionate Color developed according to recommended procedure, and measured in 1-cm. absorption cells
The instability of the color to light was corrected by developing the complex in 100-ml. mixing cylinders coated with Krylon black acrylic spray. When the reagents were mixed in these vessels, the color was stable for a t least 4 hours, and in 18 hours the absorbance decreased by only 3%, as shown in Table I. Therefore by employing 50% acetone, nitrate reagents, and opaque mixing vessels, adequate color stability was obtained. Since the use of 50 ml. of acetone per 100 ml. of final volume limited the amount of sample which could be taken, a study was made to determine the minimum amount of acetone necessary to obtain a stable color. As the percentage of acetone was decreased from 50 to O%, only a 15% decrease in absorbance was observed with no decrease in stability, provided the color was
4
18
0.114 0.105 0.097
0.114 0.105 0.097
0.111 0,101 0.094
Calibration Data
Taken, Mg.
S406--
Absorbance 1-Cm. Cells 0.732 0.143 1.464 0.286 2.196 0.427 2,928 0.570 3.660 0.711 7.32 1.42 .4verage absorbance i n d e i Standard deviation Coefficient of variation, %
Absorbance Index 0.195 0.195 0.194 0.195 0.194 0.195 0,195 ?=0.00052 ~ 0 . 2 5
IO-Cm. Cells 0.050 0.100 0.150 0.200 0,230 0.500
0.097 0.194 0.294 0.390 0.488 0.973
Average absorbance index Standard deviation C'oefficient of variation, %
0.2
1
.4bsorbance
0,194 0.194 0.195 0.195 0.194 0.195
0.195-5 0 .O O O O O 10.28
developed in opaque cylinders, as shown in Table I. Since the slight decrease in sensitivity obtained iri the absence of acetone is more than compensated for by the ability to use larger sample aliquots, further use of acetone was discontinued. Acid Concentration. A study of the color development as a function of acid concentration revealed that the absorbance is constant over a final concentration of 0.35 to 1.O.V nitric acid, provided that the same amount of acid is present in the blank. Therefore, the use of 10 ml. of 1 to 1 nitric acid to give a final concentration of 0.8S nitric acid is recommended in the proposed procedure. Ferric Nitrate Concentration. A similar study revealed that the absorbance is constant over a final concentration of 0.1 to 0.3M ferric nitrate. Since excess ferric nitrate merely increases the bapkground absorbance, 5 ml. of 2 M ferric nitrate to give a final concentration of 0.1M ferric nitrate is recommended in the color-developing procedure. Optimum Wave Length. The spectrum of the complex, as shown in Figure 1, indicates that the optimum wave length is 458 to 462 mp instead of 525 mp as employed by Robinson ( 5 ) . Use of this lower wave length results in a twofold increase in sensitivity, without any change in cell length. Calibration Data. The calibration data, shown in Table 11, were obtained by developing the color on aliquots of a standard solution of sodium tetrathionate. Identical absorbance indices were obtained in the 1-em. and,lO-em. cells. Excellent reproducibility was achieved in each instance. A practical loR-er limit of about 0.05 mg. of tetrathionate ion per 100 ml. was attained, indicating that a sample containing as little as 0.002 gram of tetrathionate ion per liter can be analyzed by the recommended procedure. Interferences. KOformal study was made of interferences. Cations which form complexes or precipitates nith thiocyanate and could compete with the ferric ion for the available thiocyanate must be absent. Therefore, Robinson recommended passing such samples over a strong cation exchange resin in order to remove these interferences. This procedure has been retained.
V O L U M E 2 7 , NO. 11, N O V E M B E R 1 9 5 5
1841
I n addition to tetrathionate, the higher polythionates also react with cyanide to form thiocyanate.
+ 5CN- + H20 =
SsOs--
+ 4CN- + Hz0 =
SsOs--
S203-S203--
+ SO4-- + 2 H C S + 3CSS+ Sod-- + 2HCN + 2CSS-
Therefore the method as developed is suitable for the three polythionates, although each ion will interfere in the deterrnination of the other. I n practice, it is probable that all three ions nill exist to some degree in any sample suspected to contain polythionates. As any thiocyanate in the liquor is a serious interference] in the recommended procedure a sample blank is prepared by adding the reagents in such a manner that tetrathionate is not converted to thiocyanate. Robinson measured the sample against the sample blank, but this procedure required norking with a variable slit width. I n the recommended procedure both the sample and sample blank are measured against a reagent blank a t constant slit width. The difference in absorbance is proportional to the tetrathionate concentration Test of Method. Neither standard samples nor an independent method of analysis of sufficient sensitivity was available to test the method. Therefore, it was decided to add various amounts of a standard tetrathionate solution to a solution of unknown tetrathionate composition and attempt to determine the added tetrathionate accurately. Several 10-nil. aliquots of a liquor known to contain tetrathionate, with and without 0.05 to 2.5 nig. of added tetrathionate, were analyzed by the recommended procedure using both 1-cm. and IO-cm. absorption cells. The absorbance obtained for an aliquot containing no added tetrathionate was subtracted from the absorbance obtained on the other aliquots. From these corrected values the amount of tetrathionate ion found was calculated. In these calculations an absorbance index obtained by passing pure solutions over the ion exchange resin (1.4% lower than the value reported in Table 11) was employed. An average error of only 1% was observed between the amount of tetrathionate added and recovered (Table 111). The general
Table 111. Tetrathionate Added, M g .
Test of Method Error,
Tetrathionate Found, M g 1-Cm. Cells 0.501 1 01 1.51 2.00 2.48
0,500 1.00 1.50 2.00 2.50
70
+0.2 +1.0 +0.7 0.0 -0.8
10-Cm. Cells 0.0.50 0.100 0 . 150 0.200 0.250
0.049 0.097 0.150 0.197 0.245 .iY.
-2.0 -3.0 0.0 -1.5 -2.0 rl 0
tendency for the answers to be slightly low seems to indicate some holdup of the tetrathionate on the resin. If the calibration curve is obtained by passing pure solutions over the ion e.rchnTige columns in the same manner as the samples are to be treatetl, this holdup of tetrathionate should be automatically combensated for. LITERATURE CITED
Furness W. and Davies, W.C., Analynt, 77, 6 9 i (19.X). Jay, R.R., A N ~ LCHEM., . 25, 288 (1953). Mellor, J. W., ”Comprehensive Treatise on Inorganic and T1:ecretical Chemistry,” vol. 10, p. 617, Longmans, Green, L O I ~ L ! ~ ~ , 1930. hlurayama, T., J . Chem. SOC.Japan, Picre Chem. Sect., 7.!, ::A9 (1953). Robinson, R. E., unpublished data Sandell, E. B., “Colorimetric Determination of Traces of Sletals,” 2nd ed., p. 367, Interscience, New York, 1953. RECEIVED for review M a y 15, 1955. Accepted August 10, 1955. The T!aw Materials Development Laboratory is operated by t h e Kational Lead Co., Inc., for t h e Atomic Energy Commission. Work carried o u t under Contract S o . AT!49-6)-924.
CRYSTALLOGRAPHIC D A T A
100. Procaine Penicillin G H A R R Y A. ROSE, Lilly Research Laboratories, Indianapolis 6, Ind.
C&
CH3
CI~I-~-CII-COOH.CtHpN-CH*CH20C-
0 11
I I N \ /\
s
CH
C=O
\d Structural Formula for Procaine Penicillin G ROCAINE penicillin G is a salt of penicillin which is highly ‘insoluble in water. The compound has found great use in medicine, but the crystallography has been very incompletely described. Crystals suitable for crystallographic work may be obtained by slow evaporation of methanol-n-ater solutions (Figure 2). Good crystals may also be obtained by slow mixing of aqueous solutions of procaine hydrochloride and sodium penicillinate. The crystals obtained from methanol-water n-ere exception-
X-Ray Powder Diffraction Data d 13.83 9.49 8.38 7.38 7.03 56 6 .. 75 0 795 5.55 5.22 4.90 4.71 4.62 4.47 4.27 4.12 3.68 3.51 3.38 3.28 3.16 3.12 3.03 2.954 2,888 2.793 2.701 2.582 2.459 2.334 2.281 2.078 2.050 2.022
r/r1 0.33 0.27 0.33 0.13 0.27 0.13 0 . 05 73 0.20 0.67 0.20 0.97 1 .00 0.07 0.13 0.53 0.33 0.27 0.33 0.33 0.07 0.07 0.07 0.07 0.03 0.07 0.03 0.03 0.20 0.03 0.03 0.03 0.03 0.03
hkl 100 001 110 111 011 lOL
d(Ca1cd.) 13.92 9.54 8.37 7.3G 7.05 6.60
210 211 111 020 120 002 311, 500 302 221 312 3 2 i . 225 320 130 313
56 . 08 70 5.58 5.23 4.90 4.77 4.66,4.64 4.50 4.28 4.13 3.69,3.68 3.47 3.39 3.28
