for fluoride alone. By combining the electrode switch and the addition technique, several ions may be analyzed by a single addition which contains all the ions to be analyzed and with the speed and ease of a single ion measurement. Previously, multiple ion analysis meant preparing a calibration curve for each ion by plotting activity us. millivolt readings, o r by calibrating a logarithmic scale with two standards to bracket the unknown. The known addition obviates these shortcomings and prepares the way for multiple ion analysis with millivolt readings. Ion selective electrodes measure ion activity, but the method of known addition gives total ion concentration. The problems of pH adjustment, ionic strength differences between sample and standard, and complex formation which are critical in activity measurements can be used to advantage by the method of known addition.
Table 111. Comparison of Data Obtained by Known Addition for F- and CI- to an X-Ray Fluorescence for C1- and a Microdistillation, Colorimetric Method for F- in Calcium Halophosphate % F- by F- by C1- by C1- by std add. microdist std add. X-ray Sample 1 2.95 2.95 0.93 0.90 2 2.88 2.87 0.77 0.75 3 2.90 2.83 0.78 0.77 4 2.89 2.85 0.80 0.80 5 2.92 2.85 0.84 0.82 6 3.03 3.03 0.59 0.58 0.59 0.58 7 3.03 3.03 8 3.04 2.95 0.51 0.48 9 3.07 3.10 0.50 0.49 10 3.01 3.01 0.55 0.54
PRECISION AND ACCURACY
for fluoride and 0.8 for chloride. In Table 111, the accuracy of the method is checked against X-ray fluorescence and a microdistillation colorimetric method. In conclusion, although the theory of the method is somewhat complicated, the actual practice is rapid, simple, and precise and well-adapted for multiple ion analysis.
The following data were obtained both by controlled experiments and by including standard samples with routine samples for analysis. In Table 11, the precision of the addition method is compared to activity measurements. On replicate analysis of 10 samples by the known addition, a relative standard deviation of 0.7% for fluoride and 0.4% for chloride was obtained. On a production basis, a standard sample submitted six times with routine samples over a 2month period, gave a relative standard deviation of 1.1
RECEIVED for review May 4, 1970. Accepted October 22, 1970.
Polarographic Determination of Perbromate in the Presence of Bromate Bruno Jaselskis and J o h n L. Huston Department of Chemistry, Loyola Unicersity, Chicago, Ill. 60626 SINCETHE DISCOVERY of perbromates, several titrimetric methods for the determination of perbromate have been reported (/, 2). These methods are based on the reduction of bromate either by bromide in 12M HBr or by stannous chloride. In addition, a spectrophotometric method (3) for perbromate analogous to the perchlorate determination ( 4 ) has been described. All of these methods are accurate; however, they are tedious and long. A faster polarographic method has been developed and is described in this note. EXPERIMENTAL
Materials and Apparatus A stock solution of approximately 0.2M potassium perbromate was obtained from E. H. Appelman of the Argonne National Laboratory. Potassium perbromate was recrystallized and the crystals were dissolved to yield approximately 0.1M stock solution. The resulting solution was standardized by the iodometric method as described by Appelman. All supporting electrolytes were prepared from reagent grade chemicals. Polarographic determinations were carried out using a Sargent Model XXI polarograph and thermostated H-cell with a saturated calomel electrode(SCE). Appelman, J. Amev. Cl7ern. SOC.,90, 1900 (1968). Appelman, Zmrg. C/7em., 8, 223 (1969). (3) L. C. Brown and G. E. Boyd, ANAL.CHEM., 42,291 (1970). (4) S. Uchikawa, Bid/. Chem. Sac. Jup., 40,798 (1967). (1) E. H. (2) E. H.
Procedure Typical polarographic solutions of perbromate in the concentration range 5 X 10-5 to 5 X 10-4M were prepared by taking suitably sized aliquots, adding them to 10 ml of the supporting electrolyte stock solution, and diluting the contents to the 25-ml mark. Stock solutions of electrolytes were prepared by using weak acid and its conjugate, and sodium perchlorate to yield an ionic strength of approximately 0.1 upon dilution. The determination of perbromate alone or in the presence of bromate in amounts equal to or larger than perbromate was carried out in supporting electrolytes having pH > 3, such as phosphoric acid-phosphate, acetic acid-acetate, or sodium perchlorate-bicarbonate. The height of the perbromate reduction wave was measured by subtracting the current of the supporting electrolyte alone from the current in the presence of perbromate. The concentration of perbromate was then determined from the calibration curve. Both perbromate and bromate, if present in amounts less than perbromate, were determined by first recording the polarogram in 0.05M perchloric acid and then recording the polarogram after the neutralization of the same solution with solid sodium bicarbonate. In acid medium, only one reduction wave was observed, corresponding to the reduction of perbromate and bromate to bromide : In neutral or alkaline media, two reduction waves resulted, the first corresponding to the reduction of perbromate to bromate and the second due to the reduction of bromate to bromide. The perbromate concentration was determined from the calibration curve obtained in perchlorate-bicarbonate medium, while the bromate conANALYTICAL CHEMISTRY, VOL. 43, NO. 4, APRIL 1971
*
581
Table I. Effect of pH and Ionic Strength on the Reduction of 1.17 X 10-4M Perbromate to Bromate at the DME at 25 "C Diffusion E l / , DS. SCE current, pA Supporting electrolyte Calcd," _____ Br04- Br03Type p pH Obsd E",/? Br03- Br-* ncy5 Phosphate 0.16 3.20 -0.018 -0.078 0.90 2.80 0.366 Phosphate 0.16 6.00 -0.068 -0.080 0.87 , . . 0.420 Bicarbonate 0.16 10.30 -0.092 -0.116 0.88 . . . 0.520 Acetate 0.12 4.50 +0.017 ,., 0.89 , . . 0,440 Acetate 0.22 4.50 -0.018 .., 0.86 . . . 0,440 Acetate 0.51 4.50 -0,093 .,. 0.80 , , , 0,440 aThe €"I/? and cy for the reduction wave of perbromate to
c
~~
z
w
a
a
~
1.3.
v 3
i
bromate are obtained from the plot of E cs. log I
POTENTIAL
VS
* The diffusion current of Br03-
S.C.E.
__ _ _ InIn acetic 0.05M perchloric acid and 0.05M sodium perchlorate acid-sodium acetate buffer, pH 3.08 centration was calculated from the corrected diffusion current in perchloric acid - iBrO4-
(1)
and the calibration curve. RESULTS AND DISCUSSION
Perbromate is reduced irreversibly a t the dropping mercury electrode (dme) in a manner similar to periodate, as shown in Figure 1. The reduction of perbromate at p H > 3 proceeds in two steps: first the perbromate is reduced to bromate and then bromate is reduced to bromide. i n strongly acidic solutions, p H < 2, perbromate is reduced directly to bromide. The perbromate wave a t higher concentrations shows a retrograde current, which is affected by the electrolyte used, the dropping time of mercury, and the concentration of perbromate. The effect is minimized by using drop times longer than 5 sec, by using univalent electrolytes, and by limiting the concentration to less than 5 X 10-4M of perbromate. The reduction of perbromate is not affected by the hydrogen ion activity in the p H range 2 to 10, as seen from the El/, observed. In fact, the calculated Eo1/,values fall well within experimental errors. Insensitivity of perbromate reduction to hydrogen ion indicates that the rate determining steps d o not involve hydrogen ion. The reduction of bromate, on the other hand, is highly dependent on pH, as has been reported previously (5-8). Thus, by changing pH, the reduction waves of bromate can be moved over by one volt from the perbromate wave. The ionic strength affects not only the Ell2 but also the diffusion current observed. Since the perbromate reduction occurs on the positive side of the E,,,, E l , ? potential shifts toward more negative potentials with increasing ionic strength. The diffusion current for the reduction of perbromate is not affected by pH, but depends on ionic strength and height of mercury. The diffusion current decreases with in(5) E. F. Orleman and I. M. Kolthoff, J. Amer. Chem. SOC.,64
1044 (1942). (6) Zbid.,p 1070. (7) V. Cermak, Chem. L i s f y . , 50, 983 (1956). (8) V. I. Zykov, and S. I. Zhdanov, Zh. Fiz. Khim.,32, 644 (1958). 582
-
- 0.546 log t.
at pH's higher than 4.50 merges with the hydrogen discharge current.
Figure 1. Polarograms of perbromate in supporting electrolytes
ioorreoted = iabservod
Id
ANALYTICAL CHEMISTRY, VOL. 43, NO. 4, APRIL 1971
Table 11. Polarographic Determination of Perbromate at 25 "C and p = 0.1 Diffusion Molarity X lo4 current, pA Taken Found KBr04 KBr03 Acida Baseb K B r 0 4 KBr0; 1.73 0.43 0.680 . . . 0,680 i 0.008 2.30 0.57 0.905 0.905 i.0.008 4.66 1.15 1.900 1.900 i. 0.018 8.30 2.06 3.400 3.400 i. 0.03 1.82 . . . . . . 0.800 ... 0.800 f 0.010 3.66 . . . . . . 1.600 ... 1.600 i 0.016 7.35 . , , . . . 3.200 ... 3.200 =t0.022 3.93 0.55 0,905 0.800 0.900 i.0.012 0.76 i 0,013 5.42 1.12 1.800 0.400 1.820 i 0.025 0.41 f 0.011 a Diffusion current determined in 0.05M HC1O4 and 0.05M NaC104 supporting electrolyte. b Diffusion current determined in the supporting electrolyte which was obtained by neutralizing the above electrolyte with solid sodium bicarbonate.
creasing ionic strength and also is proportional to the square root of mercury as predicted for diffusion controlled processes. The diffusion current of bromate is highly dependent on p H as previously observed. Thus, in developing suitable analytical procedure, exact pH, ionic strength, and height of mercury must be controlled. The effects of p H ana ionic strength on the reduction wave are summarized in Table 1. The calculated ncr values and the shape of the reduction wave indicate that the reduction of perbromate is highly irreversible. Perbromate is in many respects similar to perchlorate; it is somewhat less stable but more difficult to prepare. Quantitative determination of perbromate alone or in the presence of small amounts of bromate can be carried out in dilute (0.05M) perchloric acid solution yielding one wave corresponding to the reduction of perbromate to bromide. Upon the addition of sodium bicarbonate to a point where methyl red changes color to yellow, only the reduction of perbromate to bromate is observed. Results of perbromate and perbromate-bromate mixture determinations are summarized in Table 11. Perbromate in the presence of bromate can be determined with a relative precision of better than 2%, while the relative precision for bromate in the mixtures is only somewhat better than 3.5 %. The error in the determination of bromate in the mixtures is greater than that for perbromate, since the diffu-
ACKNOWLEDGMENT
sion current corresponding to Br03- represents a difference, iBrOJ= iobserved - iBrOa-.In the presence of perbromate, the diffusion current must be measured by subtracting the charging current of the electrolyte alone from the current in the presence of bromate or perbromate. The polarographic method is rapid and quite suitable for the determination of perbromate in the presence of bromate.
The authors thank E. H. Appelman of the Argonne National Laboratory for the solution of potassium perbromate. RECEIVED for review August 6, 1970. Accepted November 25,1970.
Polarimetric Studies of Alkali Metal Ion Complexes of I-PropylenediaminetetraaceticAcid James D. Carr and D. G . Swartzfager Department of Chemistry, University of Nebraska, Lincoln, Neb. 68508 RECENTLY, SEVERAL REPORTS have appeared on the formation of weak 1 : 1 complexes between the alkali metal ions and the aminocarboxylate multidentate ligands: ethylenediaminetetraacetic acid ( I ) , d,I-propylenediaminetetraacetic acid (abbreviated d,l-PDTA) ( 2 , 3 ) and trans-l,2-diaminocyclohexaneN,N,N',N'-tetraacetic acid ( 4 ) . In the case of the potassium complex of d,l-PDTA, evidence of a chelated structure has been obtained via the NMR vicinal coupling constants of the ethyleneic protons. The stability constant of the complex was reported to be 1.5 at 100 "C and 5.7 at 35 "Cinextremely high ionic strength media ( p > 3) (3). Although such small interactions are usually quite difficult to observe under conditions where the information is most useful (lower temperature and ionic strength), this is not the case when optically active PDTA is employed. It has been well established that the optical rotatory properties of dissymmetric molecules are, in many cases, extremely sensitive to changes in conformation, solvation, and changes of a chemical nature. In the case of I-propylenediaminetetraacetic acid (abbreviated here as IPDTA or L), changes which occur upon complex formation are accompanied by a large change in the optical rotation. This provides a convenient and accurate method of monitoring the formation of these weak complexes. EXPERIMENTAL Apparatus. Polarimetric measurements were performed at 365 mp in a 10-cm cell thermostated at 25 'C in a PerkinElmer Model 141 polarimeter. All p H measurements were made with a Corning Model 12 expanded scale p H meter. Reagents. All solutions were prepared with deionized water and stored in polyethylene bottles. The anhydrous I-propylenediaminetetraacetic acid was prepared by the method of Dwyer and Garvan ( 5 ) . A 0.5% aqueous solution of the active acid gave a specific rotation of -50.0" at the sodium D line. Stock solutions of the cations sodium, potassium, cesium (Alfa Inorganics), and tetramethylammonium (Southwestern Analytical) were prepared from their respective hydroxides. Solutions of lithium and tetraethylammonium (1) G. Anderegg, Hela. Chim. Acta, 50, 2333 (1967). (2) J. L. Sudmeier and A. J. Senzel, ANAL.Cmw, 40, 1693 (1968). (3) J. L. Sudmeier and A. J. Senzel, J. Amer. Chem. SOC.,90, 6860
hydroxide were prepared via ion exchange of their chloride salts. All cation stock solutions were checked for the presence of sodium and potassium impurities by flame photomwere etry. Corrections, which were never greater than made for these impurities whenever necessary. Procedure. Working solutions were prepared volumetrically from the stock solutions. The initial concentration of I-PDTA was about 2.0 X 10-2 molar. After the alkali metal ion was added (in the hydroxide form), the ionic strength was adjusted to 0.5 with tetramethylammonium hydroxide which gave an initial p H of about 13.4. The optical rotation of the solution was determined as the pH was lowered by the addition of concentrated HCI. The observed molar rotation was then calculated according to Equation 1 after the total concentration of the active species was corrected for dilution.
5z,
[Cl'lobs = where
=
&be
b
=
c =
Cl'obs/bC
the observed rotation (degrees) cell length (centimeters) concentration (moles per liter) RESULTS AND DISCUSSION
The effect of a large excess of each of the cations on the observed molar rotation over the pH range from 1.5 to 13.5 is illustrated in Figure 1. Since the observed molar rotations in the presence of both tetramethylammonium and tetraethylammonium were identical within experimental error over the entire pH range, it is assumed that these cations do not interact with 1-PDTA. The behavior in the high pH region (9 t o 13.5) is generally consistent with the known complexation of lithium, sodium, and potassium; however, there is some indication that cesium is also being complexed or is in some way interacting with the active ligand. This interaction proved to be too small to measure quantitatively. The effect of varying the potassium ion concentration on the observed molar rotation in the high pH region is shown in Figure 2. Assuming that the species KHLZ- is nonexistent, the. constraints on the system are represented by Equations 2-5.
(1968). (4) J. D. Carr and D. G. Swartzfager,ANAL.CHEM., 42,1238 (1970). ( 5 ) F. P. Dwyer and F. L. Garvan, J. Amer. Chem. SOC.,81, 2955
(3)
(1959). ANALYTICAL CHEMISTRY, VOL. 43, NO. 4, APRIL 1971
583
