1834
ANALYTICAL CHEMISTRY
work, to R. A. Anduze for analyzing some of the samples reported here, and to K. T. Cave for obtaining the infrared spectra.
Karush, F., Sonenberg, lI.,ASAL. CHEnf. 22, 175 (1950). Kuenteel, L. E., Ibid., 27, 301 (1955). Lewis, G. R., Herndon, L. K., Sewage and I n d . Wastes 24, 1456
LITERATURE CITED
Longwell, J., Maniece, W. D., Analyst 80, 167 (1955). Lundgren, H. P., J . Am. Chem. SOC.63, 2854 (1941). Lundgren, H. P., Elam, D. W., O'Connell, R. .I..J . B i d . Chem.
Barr, T., Oliver, J., Stubbings, 11'. V., J . SOC.Chem. I n d . (London) 67, 45 (1948).
Barton, -4.D., Young, L., J . Am. Chem. SOC.65, 294 (1943). Biffen, F. hl., Snell, F. D., IND. E N G .CHEX.,ANAL.ED. 7 , 234 (1935).
Brand, B. P., Johnson, P., Trans. Faraday SOC.52, 438 (1956). Brooks, F., Peters, E. D., Lykken, L., IND. E N G .CHEM.,Ax.4~. ED. 18, 544 (1946). Degens, P. K., Evans, H. C., Kornmer, J. D., Winsor, P. A , J . A p p l . Chem. (London) 3 , 54 (1953). Edwards, G. P., Ginn, 31. E.,Sewage and I n d . Wastes 26, 945 (1954).
Epton, S.R., Trans. Faradau SOC.44, 226 (1948). Evans, H. C., J . SOC.Chem. I n d . (London) 69, s76 (1950). Hammerton, C., J . A p p l . Chem. (London) 5 , 517 (1955). Haney, P. D., others, J . Am. Water Works Assoc. 46, 751 (1954).
Harker, R. P., Heaps, J. AI., Horner, J. L., S a t u r e 173, 634 (1954).
Harris, J. C., Short, F. R., Food Technol. 6, 275 (1952). House, R., Darragh, J. L., -4r;a~.CHEM.26, 1492 (1954). Jones, J. H., J . Assoc. O f l c . Agr. Chemists 28, 398 (1945).
(1952).
149, 183 (1943).
llarron, T. U., Schifferli, J., ISD. EX. CHEM.,d s a ~ED. . 18, 49 (1946).
Moore, W. A., Kolbeson, R. A, ~ N A L CHEM.. . 28, 161 (1956). Nukerjee, P., Ibid., 28, 870 (1956). Xorton, T. H., Otten, A. H., -47n.Chem. J . 10, 140 (1888). Sorton, T. H., Westenhoff, J . H . , Ibid., 10, 129 (1888). Reid, V. W., Alston, T., Young, B. IT., dnalyst 80, 682 (1955). Rosen, A. A,, Middleton, F. AI.,Taylor, X . W., J . Am. Water Works Assoc., in press. Sadtler, P., A S T M Bull. 190, 51 (1953). Sallee, E. If.,others, d x . 4 ~CHEM. . 28, 1822 (1956). Short, F. R., Good, G., Ibid., 28, 1504 (1956). Smith, E.L., Page, J . E., J . SOC.Chem. I n d . (London) 67, 48 (1948).
tTa11in, G. R., h A L . C H E h f . 22, 616 (1950). Yang, J. T., Foster, J . F., J . 4 m . C h e m SOC.75, 5560 (1953). Zin'kov, 2. E., Danyushevskii, Ya. L., Reinshtein, V.,Khomyakovskii, G. SI.,J . A p p l . Chem. ( r S S R ) 9, 1997 (1936). RECEIVED for review July 19, 1956.
Accepted October 5 , 1956.
Ninth Annual Summer Symposium-Analysis of Industrial Wastes
Direct Spectrophotometric Determination of Small Amounts of Chloride PHILIP W. WEST and HANS COLL Coates Chemical Laboratories, Louisiana State University, Baton Rouge, La.
There has long been a great need for a direct spectrophotometric method for the determination of small amounts of chloride. Such a procedure is now proposed, based on the use of iron(II1) perchlorate. Chloro complexes of iron(II1) exhibit an intense absorption band in the vicinity of 340 mp and lend themselves to quantitative measurement. The proposed procedure is accurate and relatively free from interferences. Large amounts of sulfate tend to interfere. However, the method can be used in the presence of the other halides and practically all other common ions.
in the spectral range suitable for the determination of chloride. Under the conditions employed in these studies, no decomposition of perchloric acid was evidenced a t any time. S o reduction t o chloride was ever observed nor were there any indications of other interfering or hazardous reactions. The procedures involved may be considered as perfectly safe (4). A spectrophotometric procedure for the determination of iron, presumably as the tetrachloroferrate(II1) complex, has been reported by Desesa and Rogers ( 2 ) . Their procedure requires a high concentration of hydrochloric acid and allows determination of iron in concentrations of the order of 10 mg. per liter. APPARATUS AYD REAGEVTS
F
OR some time it has been known that iron(II1) is not yellon,
as popularly believed, but instead is essentially colorless. The yelloly color attributed to ilon(II1) is actually due (3)to chloro or hydroxo complexes of iron. The light-absorption properties of the chloro complexes suggested the possibility of using chloride-free iron as a reagent for the colorimetric estimation of chloride and the present report summarizes the results of an investigation undertaken to ascertain the validity of a direct spectrophotometric procedure based on such an approach. Although iron( 111) chloro complexes absorb radiant energy in the visible portion of the spectrum, the greatest absorption is in the ultraviolet. It was found that the intensity of the absorption band exhibited by the complex around 350 mp depends largely on the amount of acid present. For reasons of sensitivity it was, therefore, deFirable to maintain an acid concentration as high as possible and work in the ultraviolet rather than in the visible portion of the spectrum. Among the strong acids available, perchloric acid appeared to he most suitable, for several reasons. The iron perchlorate complexes ( 1 , 6) have a lorn stability and they absorb little energy
A Beckman Model DU spectrophotometer Kith matched 1-cm. silica cells was used for most of the studies; a tungsten lamp served as the light source. It was established that a Model I3 spectrophotometer and Corex cells are likewise suitable for this work. Absorption spectra were recorded by means of a Beckman Model DK-I recording spectrophotometer. A Machlett closedsystem buret (10 ml.), connected with a large stock bottle by an all-glass conduit, was used for delivering exact volumes of perchloric acid. Drying tubes filled with anhydrous magnesium perchlorate protected the acid from moisture. The buret should have a close fitting glass stopcock to minimize leakage oi perchloric acid. All stopcock lubricants studied were found to be attacked by the acid upon prolonged contact and therefore the fit of the plug is of considerable importance. It was observed that the absorbancy of the iron chloro complex in perchloric acid stronger than 5iIr decreased when exposed t o bright light. The solutions m w e , therefore, protected from light by coating the volumetric flasks with black paint almost up t o the mark, and by covering the necks of the flasks with small bags of opaque paper. Low actinic glassware gave incomplete protection against bright light. Iron(II1) perchlorate solutions were prepared from reagent grade ferric perchlorate (nonyellow) as supplied b y the G. F. Smith Chemical Co. Trace impurities of chloride present in the iron perchlorate were removed by treatment of the salt with small
V O L U M E 28, NO. 1 2 , D E C E M B E R 1 9 5 6 portions of 70y0perchloric acid and decantation of the acid. The salt is only sparingly soluble in the acid, b u t the chlorides are readily dissolved and removed. I n preparing iron perchlorate solutions, an approximate amount of the purified salt was dissolved in water or in perchloric acid, and the solution was titrimetrically standardized by the potassium dichromate procedure. The solutions were then diluted to bring them to the desired strengths. Reagent grade perchloric acid (70 to 7 2 or 607,) supplied by Baker and Adamson and blallinckrodt, respectively, was found to be of satisfactory purity. No detectable effects due to chloride present in the acid have been observed. The perchloric acid always seemed to be of a higher degree of purity than indicated by the manufacturer’s assay (maximum limits of chloride impurity were given as 0.001%). A standard chloride solution was prepared by dissolving 1.314 grams of dried sodium chloride and diluting to 1 liter with distilled water. This solution had a chloride-titer of 0.8000 mg. and was further diluted to yield solutions of appropriate chloride concentrations. For evaluating the extent of interferences, solutions were prepared from reagent grade chemicals to contain known amounts of the ions listed in Table I. Solutions for anion interferences were prepared from sodium salts; those for cation interferences contained perchlorate or nitrate.
1835 first set contained 17.5 mg. of chloride per liter in 6N perchloric acid; the second contained 8.00 mg. of chloride per liter in 8h’ perchloric acid. The iron concentrations were varied from 0.001 to O.1M. Curves indicating the increase of background absorption with increasing concentration of iron are included in Figure n
L.
I t is evident from Figure 1 that increasing the concentration of perchloric acid beyond 8.5N reduces the precision of the measurements because of the increase in background without a corresponding increase in net sensitivity. A practical limit is further set by the concentration of the commercially available perchloric acid (approximately 12N). As indicated by Figure 2, a high concentration of iron(II1) seems to be desirable if the simultaneous increase of blank absorption with iron concentration is disregarded. If 70% perchloric acid is used, a limitation of iron(II1) concentration is given by the decreased solubility of iron perchlorate in the more concentrated acid. The stability, temperature dependency, and adherence to the Beer-Lambert law of the system were studied by preparing a reagent which contained 0.139 mole of ferric perchlorate in 1 liter
EXPERIMENTAL AND RESULTS
The effect of acid concentration on the absorbancy of solutions a t a given concentration of chloride and ferric ion was studied. Small volumes of iron and chloride solutions were transferred
OBO
to 10-ml. flasks, and varying amounts of 70y0 perchloric acid were added. The flasks were filled to the mark, shaken and alloxed to cool. At higher concentrations of acid the solitions were protected from light as described above. 9 study of absorption spectra of various solutions differing in acid concentration disclosed a pronounced increase of absorbancy with acidity, accompanied by a shift of the absorption bands towards the red end of the spectrum. The absorption peaks shifted from 338 mp ( 2 . 5 5 perchloric acid) to 348 mp (8.5N perchloric acid). A series of measurements was then carried out a t 350 mp with solutions which contained 8 mg. per liter of chloride ion, 0.02 mole per liter of iron perchlorate, and perchloric acid concentrations varying from 2 to 10-Y. Blanks of appropriate iron and acid concentrations were prepared, and the absorbancies of all solutions were measured with respect to water. The absorbancies due t o the presence of chloride ion were calculated by subtraction of the corresponding blank values (Figure 1). The influence of iron concrntration on the intensity of the abeorption band x a s next studied. Two sets of measurements were carried out, one a t 3-12 mbc,the other at 350 mp (Figure 2). The
t
-8
mrn. CI- (0.02M Fe)
0.60-
>
z
% m
040-
a
0.20-
I
0
I
I
I
I
1
2
4
6
8
10
N HCIOI
Figure 1.
Influence of perchloric acid concentration on absorbancy at 350 m p
Table I. Interferences Procedure 1 Procedure 2 Species Interference Reinarksn ’ Interference Remarksa Moderate (+) Sulfate Slight (+) Very slight (+) Nitrate Negligible (+) Negligible (+) Acetate None Decomposes Thiocyanate Heavy Moderate (+) ($1 Moderate (+) [ 111; can be corBromide Moderate (+) rected Slight (+) Iodide [ 7001; precipitates above 200 mg./liter Slight ( - ) Negligible ( - ) Fluoride (400) Slight (+) Phosphate Negligible ( - ) (850) Heavy ( - ) Must be absent lIercury(I1) Heavy ( Negligible (+) Chromium (111) Kegl/g!ble (+) (200) Kegligible (+) Cohalt(I1) Negligible (+) (300) iYeghgib1e ( - \ (500) Lead Negligible ( - ) Segligible (+) Kirkel Negligible (+) (GOO: Sone JIanganese (11) Negligible (+) (1000) Negligible (-) Aluminum Negligible ( - ) (> 1000) (200) Negligible ( - ) (>1000) Kegligible ( - ) Copper(I1) (>1000) None None Barium None Calcium Sone None None Sodium Precipitates above Precipitates above Potassium 400 mg./liter 300 mg./liter a Figures in brackets refer to mg./liter of interfering ion giving same absorbancy as 1 mg./liter of chloride. Figures in parenthesis refer to mg./liter of interfering ion causing a 3% error in the chloride value (these values were established a t chloride levels of 5 and 8 mg./liter for Procedures 1 and 2, respectively).
ANALYTICAL CHEMISTRY
1836 of 11.85 perchloric acid. Transfer of 7.2 ml. of this reagent t o a 10-ml. flask and dilution to the mark resulted in a solution 8.5;V n i t h respect to perchloric acid and 0.01.1f n i t h respect to iron. Because of the hygroscopic properties of the perchloric acid, this reagent had to be kept under anhydrous conditions. The specified volume of the reagent n a s delivered from the llachlett buret into flasks protected from light, and aliquots of standard sodium chloride solution nere added. .ilfter dilution to the mark with distilled water, the flasks were shaken and allon ed to cool. The absorbancies of the solutions were measured a t 353 mfi with respect to a blank \\ hich contained corresponding amounts of acid and iron. Ail linear relationship betneen absor bancy and chloride concentration n as established.
int,erference is expressed in terms of parts per million of int,erfering ions giving the same absorbancy as 1 p.p.m. of chloride. I n other cases the ext,ent of int,erferences is expressed as milligrams per liter of interfering ion causing an error of 3% at a given concent,ration of chloride ion. Consideration was also given to the presence of potassium ion, iron(III), and perchloric acid. Interferences arising from bromides and iodides were studied in special detail. The presence of bromide ion catised t,he appearance of an absorption band a t 420 mpL,probably due to the formation of a bromo complex with iron (Figure 3). The ratio of absorbancies a t 420 and 353 mfi, A420:A351, was found to be 3.6 and was reproducible. The same ratio for chloride absorption xvas calculated to be 0.082. Thus it is possible to correct for bromides present in solution, provided no interact'ion with the chloro complex takes place. Results obtained in this \\-a\. are presented in Table 11. By an analogous procedure the effect of iodides was studied. Iodide ion when added to the perchloric acid reagent is oxidized to iodine. The resulting absorption spectrum is shown in curve 1 of Figure 3. When the iodide concentration exceeded 200 mg. per liter, crystals of iodine were precipitat,ed. These were removed by extraction into carbon tetrachloride. The spectrum of the iodine remaining in the aqueous phase appears to I)e es-
V W
mM IRON(II1)
Figure 2.
z
t
40
mp
1 - 8.5N HClOi B r - 8.5N HClOa
4.
Influence of iron(II1) concentration on absorbancies
Measured a t 342 and 350
1.
2. 3. 5.
1 - 2 . 5 N HClOa 1 2.5N H C l O i after CCld extn. Br- 2.5N HClOa
for runs 1 and 2 , respectively
The useful concentration range mas found to be 1 to 12 mg. of chloride per liter (corresponding to absorbancies of 0.08 and 1.0 unit). Because only one fourth of the total volume is available for the chloride solution, these figures have to be quadrupled to represent the useful concentration range for an unknom n solution. The blank exhibited an absorbancy of approximately 0.06 unit with respect to distilled water. The influence of temperature was determined to be 1.4% per degree centigrade for the blank x i t h respect t o water, and -0.02% per degree centigrade for the sample a ith respect to the blank. The absorbancies of solutions protected from light were constant for 100 hours. Exposure to bright light in uncoated flasks after addition of chloride caused a rapid decrease of absorbancies. Subsequent storing in the dark s l o ~ l yincreased the absorbancies, although the original values xvere not again reached. I n order to test the reliability of the procedure, 21 solutions, each containing 5 mg. of chloride per liter, were measured with respect to blanks over a period of 1 week. The reagent as well as the volumetric equipment was the same in all experiments; corrections for changes in temperature were not applied. T h e values ranged from 0.428 to 0.449 absorbancy unit, the average being 0.440; the mean deviation was *0.86%. Interferences from a number of cations and anions were studied. Known amounts of potentially interfering ions were added to solutions containing 5 mg. of chloride per liter, and the difference from the theoretical value of absorbancy was calculated. I n studies on certain cation interferences where a substantial amount of nitrate ion u a s added to the solutions, corrections were applied for the presence of this anion. The remlts are summarized in Table I (Procedure 1 ) . Where positive deviations due to interaction of iron(II1) and interfering ions, independent from chloride concentration, could be assumed, the extent of
5
550
500
450
400
370
350
335
WAVE L E N G T H (millimicrons)
Figure 3.
Absorption spectra due to bromides and iodides
Table 11. Interferences from Bromide, Corrections Concentration, Absorbancies Mg./Liter of Chloride ChlorB ~ ~ -Absorbancies at 353 mp Error, ide mide 353 mfi 420 m r Calcd. Theor. % 2.5 16 0.332 0.434 0.209 0.213 - 1.9 2 5 2.5 2.5 5,0 5.0 5.0 5.0 5.0
16 40 40 8 16 32 40 40
0.355 0.504 0.498 0.475 0.541 0,649 0,722 0.710
0,420 1.095 1.11 0.220 0.438 0.838 1.10 1.10
0.245 0,207 0.195 0.426 0.432 0.427 0,428 0.416
+15 - 2.8 - 8.5 - 0.2 1.2 0.0 0.5 - 2.6
0.213 0.213 0.213 0.427 0.427 0.427 0.427 0.427
Values obtained by correction formula, A353 (corrected)
+ +
=
A353 X 3 . 6 - A m 3.5
Ratios of Absorbancies Chloride Bromide Iodide
= 0.082 -4azolAssr = 3 . 6 -44zo/A35a = 4 . 5 -4420/A353
AnrolAssa = 0.015 AllO/A353 = 2 . 3 Aa~o/Assa= 15
V O L U M E 28, N O . 1 2 , D E C E M B E R 1 9 5 6 sentially unaltered. It can be seen from Figure 3 that the absorbancy due to the presence of iodine drops to a low value around 353 mp. Thus no extensive interference from iodides in a determination of chloride should be espected. A correction for iodides similar to the one applied for bromides is possible but is not considered necessary. Corrections for t'he presence of iodides and bromides simultaneouslj- may be made by measuring the absorbancies of the solution a t 363, 420, and 470 mg. Constants expressing ratios of absorbanices a t different wave lengths are included in Table 11. The constants for iodide appeared to be less reproducible than those for bromide. The constants listed in Table I1 depend to a cert,ain estent on the concentrations of iron(II1) and perchloric acid present in solution. I n order to dispense with the necessit>-of storing the reagent under anhydrous conditions, an alternative procedure was investigated
1837 RECOMMENDED PROCEDURES
Procedure 1. Shake 20 t o 30 grams of reagent grade iron(II1) perchlorate with small portions of 70% perchloric acid in a glassstoppered flask. Decant the acid and repeat this operation until the acid layer no longer appears yellou. Mount a coarse frittedglass crucible of known weight on a suction flask, and transfer the purified iron perchlorate to the crucible. Press the crystals down n i t h a glass rod flattened on one end, and remove the perchloric acid by suction. Then req-eigh the crucible filled with iron perchlorate on a triple-beam balance; remove iron perchlorate from the crucible until approximately 8 grams are left. Then dissolve this quantity in a few milliliters of distilled water, and add 1 liter of 707, perchloric acid. Perchloric acid, like iron perchlorate, is extremely hygroscopic and should be exposed to air as little as possible.
I
Or
A reagent was prepared which contained 0.2 mole of iron(II1) perchlorate in 1 liter of 5 5 perchloric acid. Acid of this or lower concentration does not exhibit hygroscopic properties. The large amount of iron was chosen t o compensate for the decrease in sensitivity resulting from the use of lower acid concentrations. Five milliliters of this reagent-standardized with respect to acidity and iron concentration-were measured into 10-ml. flasks and the features of this procedure were studied by experiments analogous to the ones described for the first procedure. The results may be summarized as follows: The plot of absorbancy against chloride concentrat,ion, measured a t 348 mp, is linear; the workable range for 1-cm. cells is 4 to 40 mg. of chloride per lit,er, corresponding to a concentration of 8 t o 80 p,p,m. in t,he unknown solution t o be analyzed. The absorbancy of the blank amounted to approsimately 0.21 absorbancy unit. The temperature dependence was determined to be 395 per degree centigrade for the blank with respect to water, and 1.3% per degree centigrade for the sample with respect to the blaiik. A statistical study on 20 solutions of constant composition gave a spread from 0.387 to 0.424, averaging 0.401 with a mean deviation of &1.9%. If t,he absorbancy values were corrected for temperature (30" C.), the spread was 0.404 to 0.416, averaging 0.408 with a mean deviation &0.6%. The absorbancy readings for the blank as well as for the sample were practically constant with time. After 50 hours a decrease of approsimately 3% was observed. S o effect due to esposure to light could be detected. The result,s of interference studies using t,he second procedure are listed in Table I (Procedure 2). The absorption bands arising from the presence of bromide and iodide (Figure 3) appeared t o be ahifkd t,oivards the ultraviolet compared with the ones obtained by t'he first procedure; the iodine spectrum exhibits a serond band around 340 mp. Estraction of precipitated iodine into carbon tetrachloride seems to have an effect on the spectrum of the aqueous phase (Figure 3, curves 3 and 4). S o att'empt was made to determine chloride in the presence of the other two halides, but it should be possible to applj- a correction, as in the first procedure. The preparation of the reagents, involving standardization of iron perchlorate solution and of perchloric acid, can be simplified for practical work by preparing reagents of approximate strength, for which an individual chloride calibration curve can be establiahed. Perchloric acid can be assumed to be of the concentrtttion stated by t,he manufart,urer, and the appropriate amount of iron(II1) perchlorate can be weighed on a triple-beam balance. After this salt has been t,reated in the manner outlined below, almost twice the amount corresponding t,o t,he formula weight of iron perchlorate has to be t,aken to give solut,ions of desired concentrations, because of the large quantities of water bnd perchloric acid absorbed on the crystals. Occasional checks by titration t,o establish the concentration of the ferric perchloratr solutions are recommended.
1 - 1 2 p , ~ m . C I ' ( 0 0 1 M Fe and 8.5N "2104)
I
-.Blank
n
/
/
WAVE LENGTH (millimicrons)
Figure 4.
Absorption spectra of iron(II1) chloro complexes and of corresponding blanks
Transfer the reagent t o a container which is connected with a closed-system buret. The system is protected from moisture by means of magnesium perchlorate drying tubes. Deliver 7.4 ml. of the reagent accurately t o 10-ml. volumetric flasks which have been protected from light by a black coating almost up to the mark. Transfer 2.5 ml. of the unknown solution, to contain not more than 50 p.p.m. of chloride, t o the flasks, and bring the solutions to volume with distilled water. Shake the flasks to mix the solutions, and, after covering the necks of the flasks with a small bag of opaque paper, allow the sclutions to cool. Measure the absorbancies of the solutions at 353 mp with respect to a blank containing only reagent and distilled water. Determine the concentration of chloride by means of a calibration curve. Establish a new calibration curve every time a fresh reagent is prepared. If the reagent is t o be kept over an estended period of time, occasional checks on the calibration curve are recommended. Procedure 2. Xft,er purification of an appropriate amount of iron(II1) perchlorate in the way described above, dissolve approximately 120 grams of the salt in a mixture of 540 ml. of 60 perchloric acid and 460 ml. of distilled rvater. The reagent m , be stored in a glass container without protection from moist air. Transfer 4.9 ml. of this reagent to 10-nil. volumetric flasks, which need n3t be protected from light. Add 5 ml. of the unknown solution, to contain not more than 100 p.p.m. of chloride, and fill the flasks to the mark. Measure the absorbancies of the solutions a t 348 mp with respect to a blank. The temperature of t h t solutions should be recorded, hut it can be assumed that the temperature difference betn-een sample and blank is insignificant. After correcting the ahsorhancy readings to the temperature a t which the calibration curve had heen established, obtain the Concentration of chloride from the curve. .\pply the correction for temperature by adding t o the absorbancy reading a t temperature T the term (To - T ) x ; 1X~0.013, n-here To stands for the temperature a t which the calibration curve was set up: -47. designates the absorbancy measured a t temperatiire T . DISCUSSION AND COSCLUSION
Procedures 1 and 2 present condit,ions fol masimum sensitivity in a spectrophotometric determination of chloride h>-mean?
1838
ANALYTICAL CHEMISTRY
of iron(II1). Other procedures which may employ a l o n concentration of perchloric acid and of iron(II1) are conceivable. It is evident from the experimental results that differences between the systems of complexes present in Procedures 1 and 2 must exist, as exemplified by the difference in temperature coefficients, the varying extent of interferences, and the instability of solutions obtained by Procedure 1 when exposed to light. Figure 4 shows the spectral bands of sample and blank in the two procedures. I n determinations of chloride it was found desirable to measure absorbancies a t slightly higher wave lengths than correspond to the peaks of the respective absorption bands, because of the increase in background and extent of interference from certain substances with decrease in wave length. Because of higher background, appreciable temperature dependence, and more extensive interference from certain common ions, such as sulfate, Procedure 2 must be considered inferior to Procedure 1. Yet the advantages of Procedure 2-the fact that no protection from humidity is required and the insensitivity of the solutions towards light-should not be overlooked in practical work. With the exception of mercury(II), a rather uncommon constituent, other ions do not seriously impair the applicability of Procedure 1. Bromides and iodides, which account for heaviest interferences in most classical methods for the determination of chloride, may be tolerated in this procedure. Corrections for bromides are required, if present in quantities above 30% of the chloride content. Iodides may be tolerated up to a 20-fold excess with respect to chloride, provided no precipitate of iodine is formed. Sulfate ion, up to a sixfold excess, does not interfere (corresponding to a positive error of up to +3%). I n Procedure 2, on the other hand, interference from sulfate ion must be considered as serious, where sulfate present to the extent of 50% of the chloride concentration (by weight) causes a deviation of 3%. Larger quantities of sulfate must be removed by precipitation with barium perchlorate and subsequent centrifugation. Re-
moval of sulfate ion by precipitation after iron perchlorate reagent has been added was found to be incomplete. Likewise, an error will be introduced if the unknown solution contains appreciable amounts of acid and iron(II1) in addition to the quantities added as the iron perchlorate reagent. For a chloride concentration of 5 mg. per liter in Procedure 1 the error R-ill amount to 3% if the iron(II1) concentration is increased from 0.010 to 0.014M; the same effect is caused by an increase of perchloric acid concentration from 8.5 to 8.8S. I n Procedure 2, a t a chloride concentration of 8 mg. per liter, a 3% error is introduced by an increase of iron concentration from 0.100 to 0.102A1f, or bj- an increase of perchloric acid from 2 50 to 2.55N. These values represent a measure for the amounts of perchloric acid and iron(II1) which may be tolerated in an unknown to be analyzed for chloride. The main advantages of the methods described in this paper over other methods n o x in use for the determination of chloride lie in the simplicity of the procedures and their remarkable sensitivities. Generally speaking, only turbidimetric methods can be considered as more sensitive. Applications of this direct spectrophotometric method are anticipated in routine analysis of water, air, and certain physiological systems. ACKNOWLEDGMENT
The authors wish to express their appreciation to the Office of Ordnance Research for finanrial support of these investigations. LITERATURE CITED
(1) Bastian, R., Weberling, R., Palilla, F., - 4 x a ~ .CHEM.25, 284
(1953). (2) Desesa, 11.A , , Rogers, L. B., Anal. Chim. Acta 6,534 (1952). (3) Gamlen, G. B . , Jordan, D. O., J . Chem. SOC.1953, 1435. (4) Smith, G.F., Analyst 80, 16 (1955). (5) Sutton, J., Nature 169, 71 (1952). RECEIVED for review February 17, 1936.
Accepted June 20, 1966.
Ninth Annual Summer Symposium--Rapid Methods of Analysis
Instrumentation for Rapid Spectrochemical Analysis Optical and X-Ray Emission Monochromators and Polychromators J. W. KEMP Applied Research Laboratories, Glendale 8, Calif.
The success of direct-reading optical and x-ray emission techniques in providing high-speed analyses on a routine basis has led to the commercial availability of a variety of instruments for use in these fields. A review of these instruments and their capabilities should be of interest to those concerned with rapid routine analysis.
0
PTICAL spectrographic analysis has been recognized as a high-speed, routine control method for 20 years. The additional time savings afforded by the application of directreading techniques to the optical method have been generally used for the past 10 years. The x-ray fluorescence technique is no%going through its acceptance period as a routine control method. Briefly, a sample brought to a sufficiently high temperature by appropriate means emits optical line spectra of the elements present. Analysis and measurement of these spectra then provide information on the elemental composition of the sample. Thus, the optical method is destructive, as it is necessary to con-
sume a t least a small sample to provide an analysis. However, the x-ray method is nondestructive, Here, the sample is irradiated with high intensity x-rays. This causes the sample to fluoresce in the x-ray region, this fluorescence consisting of the x-ray line spectra of the elements present in the sample. Analysis of these spectra also gives information on elemental composition. Savings in both elapsed time and time per analysis are usually significant, and occasionally startling, with these techniques. I n order to replace the older, slower methods by the new techniques, ho\vever, a sizable capital investment is required. Certain limitations on sensitivity and accilracy must also be recognized. Because both optical and x-ray spectra are atomic phenomentt, these limitations can be correlated with the periodic table of the elements. T h e optical technique can be used for all metals and metalloids, but requires special equipment for the nonmetals and gases. The x-ray technique has a different $et of restrictions, being useful for all elements of atomic number greater than 19 nithout special equipment. Two classes of instruments are used: parallel and sequential. The parallel instruments give simultaneous determinations of
