V O L U M E 23, NO. 10, O C T O B E R 1951 times the molybdenum concentration) it does not interfere and ensures that molybdenum is not in the trivalent form, The only difficulties with overreduction of molybdenum if no perchloric acid is added during reduction with zinc occurred with the sample containing 7% molybdenum, and with a sample containing 11% chromium. Chromous ion seems to hinder the reoxidation of molybdenum(II1). The presence of perchloric acid during the reduction prevented low results in each case. Experiments with the reduction of molybdenum with zinc in the presence of perchloric acid indicate that molybdenum(II1) is not produced until the perchlorate is virtually all reduced to chloride. Thus if the solution is removed from the zinc as soon as the iron appears reduced, very little perchloric acid will be attacked and it is certain that molybdenum is not reduced beyond the pentavalent state. The concentrations of perchlorate and sulfate were selected arbitrarily for convenience. The analvst may use any
1507 concentrations he chooses, so long a3 they are carefully regulated. It is recommended, however, that the concentration of perchlorate be at least 0.1 M a n d that it be the predominant anion in the solution. Work is in progress in this laboratory to deternine whether this effect can be applied to the determination of chlorates and perchlorates as well as molybdenum, and to learn the mechanism of this reduction, which seems to be very specific for molybdenum. LITERATURE CITED
(1) Bredig and Michel, 2. p h y s i k . Chem., 100, 124 (1922j. (2) Haight, G. P., and Sager, W. F., unpublished work. (3) Holtje and Geyer, 2. anorg. Chem., 246, 265 (1941). (4) Lingane and Laitinen, I s u . Exvn. C H t x . . ANAL. ED., 11, 504
(1939). R E C E I V EAlignst D 2 2 , 19.50.
Ceric Sulfate- Arsenious Acid Reaction in Microdetermination of Iodine ALLEN LEIN AND NEEYh SCHWARTZ Yorthwestern University Medical School, Chicago, Zll.
HE reaction between ceric sulfate and arsenious acid was first Tapplied to the determination of iodides by Sandell and Kolthoff (11, 12) and subsequently adapted to the microdetermination of iodine in biological materials (1-4, 7-10, 13, 14). [In acid solution, ceric sulfate forms sulfatoceric acid with the cerium in the anion (6). I n this report, however, cerium in both oxidized and reduced forms is considered cationic.] The basic reaction, catalyzed by soluble iodides, involves the reduction of ceric ion by arsenite ion in an acid medium:
The reduction of the yellow ceric to colorless cerous ion is usually followed photometrically, and a series of empirical standardization curves relating photometer readings to known quantities of iodide a t predetermined reaction times is employed. Attempts to use these procedures in this laboratory for the determination of serum iodine yielded erratic results, indicating that a reinvestigation of the method was required. As none of the previously published methods takes advantage of the velocity constant as a measure of catalysis by iodide, this possibility was investigated, and studies on the light-absorption characteristics of ceric sulfatp solutions and on the kinetics of the oxidation-reduction reaction were required.
As acidity and temperature were found to influence the optical density of ceric sulfate solutions, these factors were controlled at the levels used in studying the reaction. Accordingly, a series of ceric sulfate solutions was prepared: To each cuvette were added 2 ml. of 2 N sodium carbonate] 2 ml. of 7 N sulfuric acid, various quantities of ceric sulfate, and sufficient water to bring the total volume to 10 ml. (This method is intended for use in the determination of serum protein-bound iodine: the sodium carbonate represents the residue from an alkaline dry-ashing procedure. An equivalent excess of sulfuric acid is also included.) The cuvettes were then placed in a constant temperature water bath a t 40' C., and finally read in the spectrophotometer a t 420 m p . The results disclosed a curvilinear relationship between optical density and concentration, confirming the reported ( 3 ) deviation from Beer's law. The curve describing this relationship is a hyperbola, rectifiable by plotting the reciprocals of both variahles.
LIGHT-ABSORPTION CHARACTERISTICS O F CERIC SULFATE SOLUTIONS
Ceric sulfate solutions absorb maximally in the ultraviolet range a t about 315 mp, and in this region nbsorption seems to behave in accordance with Beer's law ( 5 , 15). However, as most photometers are not useful at this wave length, light a t the blue end of the visible spectrum is employed. Chaney ( 3 )presents a curve which indicates that a t 420 mp the absorption characteristics of ceric sulfate solutions deviate considerably from Beer's law. Salter (S-fO), a t the same wave length of light, uses Klett photometer readings (directly proportional to optical density) to express concentration of ceric solutions and thereby implies conformity with Beer's law.
X Coleman Universal spectrophotometer with light a t a wave length of 420 mp was used with cuvettes consisting of 20 X 150 nim. borosilicate glass tubes, selected for uniform optical properties. Selection was made by measuring a t 420 mp the perM ceric centage light transmittance of a solution of 5 X sulfate [Ce(HSOa)r] in 3.6 N sulfuric acid: cuvettes that did not agree within 0.5% transmittance were rejected.
I -Lo
-01
-06
I '04
I -02
I 00
I 02
LOG (CE+++*), MILLIMOLES/LITER
Figure 1. Relationship between Log Concentration of Ceric Sulfate and Percentage Light Transmittance at 420 mp
As in the measurement of velocity constants the logarithm of concentration of reactants is required, it was found convenient to plot percentage light transmittance against log concentration of ceric sulfate. This resulted in an S curve, the central portion of which was linear (Figure 1). In order to exclude the nonlinear portion of the curve, and to remain vr-ell within the accurate range of the photometer, transmittance readings were subse-
ANALYTICAL CHEMISTRY
1508 quently confined to a range of 30 to 65%; between these limits, the regression equation was:
T = 20.0 - 63.2 log Ice'+++]
(2)
where T = per cent light transmittance. Then, solving for log [Ce++++I, Log [Ce++++]= 0.316 - 0.0158T (3)
If only changes in log concentration of ceric sulfate are required, the intercept may be neglected, yielding the equation: &log [Ce++++]= 0.0158 AT
(4)
The standard error of estimate of log concentration of ceric was found to be 0.011 for the linear portion of the curve; this represents a standard deviation of 52.570 for the estimation of ceric concentrations from transmittance values.
mine whether they are drawn empirically froni data or calculated from the equation. -d [Ce ++-+I Salter (8) presents an equation, = K [I-], which dt stipulates a direct proportion between the rate of reduction of the cerium and iodide concentration, and implies zero-order kinetics for the reaction with respect to cerium. This equation is contradicted by a nomogram purporting to show a linear relation-hip between the logarithm of ceric concentration and reaction time. However, interpretation of this work is coniplicated by the use of N e t t photometer readings for expressing ceric concentrations. without any correction for deviation from Beer's law.
05
-
.0.
-
KINETICS OF CERIC SULFATE-ARSENIOUS ACID REACTION
As Equation 1 indicates, this reaction is multimolecular with necessarily complex reaction kinetics. One method for expressing the velocity constant in the face of such complexities is to use the reciprocal of the time required for a given fraction of one of the reactants to react. This figure is proportional to the appropriate velocity constant of the reaction, regardless of the order of its kinetic behavior. Sandell and Kolthoff ( l a ) measured the t h e Eequired for the reaction to go to completion and found a linear relationship between the reciprocal of that time and quantity of iodide present. Their findings therefore imply a linear effect of iodide on the velocity constant of thr reaction.
~
5 W CL
03-
i v)
% .02-
L
01
I
I
Multimolecular reactions follow (pseudo) first-order kiiirtics ivhen all the reactants but one are present in high concentration; catalyst concentration, being constant, may influence velocity constant u-ithout altering the kinetic order of a reaction. .kccordingly, using a constant iodide concentration, the rate of the ceric sulfate-arsenious acid reaction was determined with arsenious acid concentrations ranging from 0.5 to 20 times the molar concentration (1 to 40 times the equivalent concentration) of ceric sulfate used.
0.2
1 o
I IO
eo
ao
40
so
eo
TIME, MINUTES
Figure 2. Hate of Change in Log Concentration of Ceric Sulfate at Yarious Initial Concentrations of Arsenious Acid &umber of curve is ratio of initial molar concentrations of arsenious arid to ceric sulfate
This procedure is feasible with the relatively large amounts of iodide used by Sandell and Kolthoff, as the reaction proceeds so rapidly as to yield a quick and sharp end point, but is not practical for smaller quantities of iodide. Measuring the time required for a given fraction of the ceric sulfate to react is also inexpedient. An alternative procedure involves measuring the quantity of cerium reduced in any given time; under these conditions, however, an expression for velocity constant requires a consideration of reaction kinetics. Chaney (3)presents a monomolecular reaction equation for the catalytic reaction but provides no esperimental data t o support it. His published standardization curves include no experimental points, and it is not possible to deter-
A series of reaction mixtures was prepared by adding the following to each of eight photometer cuvettes: 2.0 ml. of 2 S eodium hydroxide, 1.0 ml. of 14 -Vsulfuric acid, 0.10 microgram of iodide, various quantities of 0.100 -11' arsenious acid, and sufficient water to bring the total volume to 9.0 ml. The cuvettes and reagents were brought to and maintained a t a temperature of 40" C., and a t a given time, 1.0 ml. of 0.010 11.1 ceric sulfate solution was added and transinittarice readings were made every 10 minutes for 1 hour. The results ryere plotted (Figure 2 ) as the logarithm of the ceric concentration against time, which should yield linear curves for first-order kinetics, as shown by the appropriate equation, Log [Ce-++-]
=
K
log [Ce+7+-I0 - 2.3 -x t
(5)
in which [Ce++-+] = concentration of celic sulfate a t reaction time = t , [Cer+++Ia = initial concentration of ceric sulfate (at t = 0), and K = velocity constant. By statistical methods the straight lines shown in the figure were fitted to the data and except for the lowest concentration of aisenious acid, there n as no detectable deviation from linearity. From the slopes of the fitted equations, the velocity constants, K , were calculated, and \?ere plotted as a function of the initial molar ratio of arsenious acid to ceric sulfate (Figure 3). The results reveal that a t relatively low concentrations of arsenious acid, the first-order velocity constant becomes independent of changes in the ratio of reactants which necessarily occur during the reaction (linearity of curves of Figure 2 ) and a t
V O L U M E 2 3 , NO. 10, O C T O B E R 1 9 5 1
1509
high concentration tends to become independent of the initial iatio of the reactants (plateau of Figure 3). A molar concentration ratio of arsenite to ceric much above 20 is inconvenient because of the limited solubility of arsenious acid. Hoirever, at R iatio of 20, the reaction rate for a given quantity of iodictr I C elevated, thereby increasing the sensitivity of the procedure, arid the first-order velocity constant is relatively independent of variations in the initial arsenious acid concentration. FIRST-ORDER VELOCITY CONSTANT A S FUNCTION O F IODIDE CONCENTRATION
+ 0.360 I-
(6)
= 2.78 A log [Ce"+++]/t - 0.0244
(7)
[Ce+'++]/t = 0.00878
A log
This equation was solved for I-, yielding:
I-
Substituting from Equation 4, 0.0439 A T l t - 0.0244
I-
4'E
(8)
I
111 oider to determine the effect of iodide concentration on the velocity constjarit, a series of reaction mixtures was run as before b u t n i t h a 20 to 1 molar ratio of arsenious acid to ceric sulfate, and with quantities of iodide ranging from 0 to 0.12 microgram. As the slope of a firet-order curve (Figure 2) is directly proportional to velocity constant, A log [Ce'++C] was calculated (Equation 4) and thc velocity constant expressed as A log [Ce+-++]li.
0
I
I
I
1
5
10
20
30
10
NdCl IN REACTION MIXTURE, MG.
Figure 5 . Effect of Sodium Chloride Concentration on First-Order Velocity Constant
Using this procedure, the standard error of estimate of iodide was found to be =k0.0017 microgram. OUTLINE O F AY4LYTICAL METHOD I
I
I
I
I
02
.OI
06
.On
10
I
I2
IODINE IN REACTION MIXTURE, MICROGRAMS
Figure 4.
Effect of Iodide Concentration on F i r s t Order Velocity Constant
Tiirl reaction followed (pseudo) first-order kinetics at each iodide (,oncentration, and the first-order velocity constant increased nith the quantity of iodide (Figure 4). Contrary to expectation, the effect of the iodide was not linear, but it appeared possible that this was due to the presence of a contaminant capable of reacting with iodide and thus interfering with the catalytic action of thr latter. On the basis that the effect of such a contaminant might be suppressed by providing an excess of chloride ion for it to react with, a series of reactions was run in which the iodide was kept constant, but various quantities of sodium chloride were added. In Figure 5 , the resulting velocity constants are plotted against sodium chloride concentration ; the curve shows that the velocity constant is increased by adding sodium chloride, but seems to approach a maximum asymptotically. However, because iodide is a common contaminant of sodium chloride, it is doubtful that a sodium chloride level can be attained at which further increments would fail further to elevate the reaction rate. On the other hand, the shape of the curve clearly indicates that the chloride, and not contaminating iodide, is primarily responsible for the effect, by whatever mechanism may be concerned. Selecting two quantities of sodium chloride, 5 and 100 mg., experiments for determining the effect of iodide concentration on velocity constant were repeated. With the lower quantity, slight curvilinearity was still noted, but at the 100-mg. level of sodium chloride, there was no observable departure from linearity (Figure 4). These results are qualitatively in accord with the findings of Barker ( 1 ) who reported that the use of 3 mg. of sodium chloride rendered the reaction more sensitive to small quantities of iodide without appreciably altering the response to large quantities. The equation for the regression line obtained with the use of 100 mg. of sodium chloride (Figure 4)was
Reaction mixtures are prepared in the selected cuvettes and contain the following: The unknown iodide (or known iodides when preparing a standardization curve). Base, 4 me. (2.0 ml. of 2.0 S sodium carbonate). This represents the alkaline residue from a contemplated dry-ashing process. Sulfuric acid, 14 me. and sodium chloride, 100 mg., added as 2.0 ml. of 7.0 A' sulfuric acid, containing 50 nig. of sodium chloride per ml. 0.10 M arsenious acid, 2 ml. This reagent is prepared by dissolving 9.9 grams of arsenic trioxide in 60 ml. of hot 1 S sodium hydroxide, diluting LTith water to about 250 ml., acidifying with sulfuric acid till just acid t o litmus, and then diluting with water t o 1000 ml. A4sa t the arsenite-ceric ratio employed, the velocity constant of the reaction is relatively independent of the exact concentration of arsenious acid, the directions do not stipulate the precision required by other methods. 0.010 M ceric sulfate, 1 ml., in 3.6 S sulfuric acid, prepared by diluting 53 grams of the salt to 1000 ml. with the acid. As only changes in ceric concentration are measured, it is unnecessary to standardize this solution. Sufficient water t o give a final volume of 10 ml. All glassware is cleaned with dichromate-sulfuric acid cleaning solution, and rinsed in tap m t e r , distilled water, and finally distilled water redistilled from alkaline permanganate in a borosilicate glass still. Solutions are prepared with the doubledistilled water, using C.P. reagent grade chemicals. The ceric sulfate solution is added to the reaction mixture last, after the cuvette and its contents are brought to 40" C., and serves to initiate the reaction. After about 5 minutes, when time, t, is considered to be = 0, a transmittance reading, T , is made and then repeated a t convenient intervals. Using only transmittance values falling between 30 and 65%, A T / t is calculated and introduced into Equation 8, which is then solved for I-. The following advantages over previous adaptations of the ceric sulfate-arsenious acid reaction are provided by the procedure described: The solutions of arsenious acid and ceric sulfate need not be standardized. The exact time a t which the reaction is initiated need not be known; only A T values for known time intervals are required. This advantage is also provided by a recording photometer recently described by Chaney (2).
ANALYTICAL CHEMISTRY
1510 As the standardization curve is linear, iodide concentration is easily and conveniently calculated. With the increased quantity of arsenious acid used, the reaction rate is elevated and the method is more sensitive and practical for small quantities of iodide. The method is reliable, measuring iodides in a range of 0 to 0.r2 microgram with a standard error for duplicates of 0.0012 microgram. LITERATURE CITED
Barker, S. B., J . B i d . Chem., 173,715 (1948). Chaney, A. L., ANAL.CHEM., 2 2 , 9 3 9 (1950). Chaney, A. L., IND.ENG.CHEM.,ANAL.ED.,12, 179 (1940). Connor, A. C., Swenson, R. E., Park, C. W., Gangloff, E. C.. Liberman, R., and Curtis, G. M., Surgery, 25, 510 (1949). ( 5 ) Freedman, A. J., and Hume, D. E., ANAL.CHEM., 22, 932 (1950). (1) (2) (3) (4)
(6) (7) (8) (9)
Jones, E. G., and Soper, F. G., J . C h a . SOC.,1935, pt. I, 802. Lein, A., EndocrinoE., 29, 905 (1941). Salter, W. T., W e s t . J. Surg., Obstet., Gynecol., 55, 15 (1947). Salter, W. T., and Johnston, M. W., J. Clin. E n d o c r i n d . , 8 ,
911 (1948). (10) Salter, W. T., and McKay, E. A , Endocrinol., 35, 380 (1944). (11) Sandell, E. B., and Kolthoff, I. M., J . Am. Chem. Soc., 56, 1426 (1934). (12) Sandell, E. B., and Kolthoff, I . M., Mikrochim. A c t a , 1, 9 (1937). (13) Sappington, T. S., Halpern, N., and Salter, W. T., J . P h a r m . E z p t l . Thm., 8 1 , 3 3 1 (1944). (14) Taurog, A., and Chaikoff, I . L., J . Biol. Chem., 163, 313 (1946). ANAL,CHEM., (15) Weybrew, J. A., Matrone, G., and Boxley, H. M., 2 0 , 7 5 9 (1948). RECEIVED August 10, 1950.
Infrared Spectroscopic Analysis of Mixtures of Bromochlorobenzenes LLOYD N. FERGUSON AND ALMA J. LEVANT' Howard University, Wushington, D . C .
S A tool for studying orientation in aromatic bromination, a
A procedure has been worked out for analyzing mixtures of
chlorobenzene, and o-, m-, and p-bromochlorobenzene. This method waa chosen because it was desired to have a single analytical technique for examining the reaction products from the bromination of various monosubstituted benzenes. The mean relative error for ten synthetic four-component mixtures was =k0.3%, Considering the inertness of the type compounds studied here, these results would undoubtedly be difficult to obtain on a small male by other methods.
A Perkin-Elmer spectrophotometer, Model 12B, was used. As a solvent that would not appreciably react with bromine nor interfere with the carbon-halogen light absorption carbon disulfide was selected. The spectra, in carbon disulfide, of the substances studied are pictured in Figure 1. For analyses, optical densities were measured a t 13.46, 13.3, 12.89, and 1 2 . 2 5 ~with slit widths of 0.3 to 0.5 mm. Concentrations were always adjusted, by dilution, so that optical densities were read between 0.15 and 0.35. At these concentrations (
