(22) T. J. Rohm and G. G. Guilbault, Anal. Chem.. 46, 590 (1974). (231 . . A . Hulanicki. R . Lewandowski. and M. Mai. Anal. Chim. Acta. 69. 409 (1974). (24) Instruction Manuals, 93 Series Nitrate and Perchlorate Electrodes, Orion Research, Cambridge, Mass., 1975. (25) H. S. Frank and W.-Y. Wen, Discuss. faraday SOC.,24, 133 (1967). (26) "Solute-Solvent Interactions", C. Ritchie and J. Coetzee, Ed., Marcel Dekker, New York, 1969. (27) R. M. Diamond and D. C. Whitney, in "Ion Exchange", Vol. I, J. A. Marinsky, Ed.. Marcel Dekker, New York. 1968, Chap. 8. (28) A. Leo, C. Hansch. and D. Elkins, Chem. Rev., 71, 525 (1971). (29) J. J. James, G. P. Carmack, and H. Freiser. Anal. Chem., 44, 853 (1972). (30) S. Back and J. Sandblom. Anal. Chem., 45, 1680 (1973).
(31) K. W. Bunzl, J. Phys. Chem., 71, 1358 (1967). (32) . . L. P. Hammett. "Phvsical Oraanic Chemistrv". 2nd ed.. McGraw-Hill. New York, 1970, pp 369ff. (33) J. Steigman and D. Sussman, J. Am. Chem. Soc., 89, 6406 (1967). (34) "Biological Correlations-The Hansch Approach", W. Van Valkenburg. Ed., Adv. Chem. Ser., 114 (1972).
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RECEIVEDfor review September 16, 1975. Accepted December 8, 1975. Presented in part a t the 164 National Meeting, American Chemical Society, New York, August 1972. The authors appreciate the financial assistance given GMO by Johnson and Johnson.
Determination of Aluminum in Paper Machine White Water by Potent iometric Titration with FIuor ide Ion A. Homola and R. 0. James* Department of Physical Chemistry, University of Melbourne, Parkville, Victoria 3052, Australia
Aluminum in white water may be determined by potentiometric fluoride titration with less than 2 % relative standard deviation provided calcium, which may be a source of interference, is below 200 mg/l.
Table I presents the complex formation reactions and their stability constants ( 1 1 ) ,and ( 3 - A )and ( 3 - B ) . The emf of a fluoride ion selective electrode-reference electrode cell; LaF31 F-, AlF;-", A13+, Al(OH);-YIIRE
Aluminum salts, for example, aluminum sulfate or papermaker's alum, A12(S04)3m18 HzO; potash alum, KAl(SO& 24 HzO and sodium aluminate, NaAlOz are used in the papermaking industry to control certain reactions which often depend on colloid stability. One of the principal functions of this addition of aluminum salts is to enhance the retention of wet end additives (2). Also because A13+ is a weak multiprotic acid which ionizes around p H 5, its surface active hydrolysis products provide some buffer capacity and p H control ( 3 , 4 ) . Some papermaking reactions, notably rosin sizing, are quite dependent on the amount of aluminum added and particularly on the correct proportioning of rosin and alum ( 5 ) . It is important, therefore, that a simple and reliable method for the determination of aluminum in mill water be available to plant personnel. Colorimetric methods for A1 determination do exist (3, 6, 7); however, these often require 1) careful timing because color intensity is time-dependent, and 2) pH control. In some mills, aluminum is sometimes "determined" as a weak acid simply by the total acidity titration or by potentiometric titration with strong base (8). However, this method is generally unsatisfactory because 1) there may be other sources of acidity unrelated to the presence of aqueous aluminum, 2) partial hydrolysis of A13+ in mill water means that only a fraction of the aluminum can be titrated, and 3) the mechanism of hydrolysis of A1 varies with time, concentration, and temperature (3, 9, I O ) . Hence, apart from "approximate" estimates of aluminum levels, this technique leaves much to be desired. The present paper describes a potentiometric titration method suitable for the determination of soluble forms of aluminum. However, i t can also be applied in measuring the total content of aluminum in the presence of fibers and other solids. The method is based on the formation of aluminum fluoride complexes in A1 solutions to which fluoride is added. 776
ANALYTICAL CHEMISTRY, VOL. 48, NO. 4, APRIL 1976
is measured as a function of the volume of sodium fluoride added to the A1 solution. The emf is proportional to logloF-. At a given value of F- or emf, the stoichiometry of the Al-FL3-"' complex is fixed. Hence, a t constant emf the difference in volume between fluoride titrations in the absence and presence of aluminum is equivalent to the amount of fluoride bound in AI-F complexes. Hence, by measuring the titration volume difference between "unknown" and "blank" solutions a t a fixed cell emf for which the average stoichiometry of complexes is known or can be measured, the concentration of aluminum can also be measured. A potentiometric fluoride titration method for A1 has been described by Baumann ( 1 2 ) ;however, this method relies on the formation of AlFi- in ethanol solvent on the addition of excess F- and the precipitation of the insoluble cryolite, Na3AlFe. T o illustrate the equilibria involved in our method, we have calculated the equilibrium distribution of Al, F, and so4 among complex and hydrolyzed species for several combinations of concentrations as a function of pH using a computer program COMICS/COPS modified from an earlier program COMICS devised by Perrin and Sayce (13).The results of these calculations for total concentrations A ~ T= FT = and SO1 = mol dm-3 are shown in Figure 1. The stoichiometry for each reaction and the formation constants are given in Table I. For F-, approximol dm-3, the average ratio of fluoride mately 5 X bound per aluminum is slightly more than 3 and is approximately independent of p H in the p H range 4 to 5.5.
EXPERIMENTAL Apparatus. T h e cell was comprised o f a n O r i o n model 94-96A fluoride i o n electrode and a double j u n c t i o n calomel reference elect r o d e in a fluoride solution buffered t o pH 4. Reagents and Solutions. Analytical reagent grade chemicals
- , - - - - -F . -----soq AI I
..*
-
=‘10-3
--e-
=
S O ~= ICY
:*‘
Complex
SO,’-
No.
y
F-
AI3
OH-
2
V
X
Log,,
P
1 1 1 -5.3 2 1 -9.9 2 3 2 -8.0 2 4 1 -15.6 3 5 1 -23.0 4 6 8 -72.9 20 7 1 -1 2.0 8 1 1 2.16 9 2 1 3.62 10 1 -1 3.7 11 1 1 6.6 2 1 11.2 12 16.0 3 1 13 14 4 1 18.7 15 5 1 20.0 16 6 1 20.4 3 -10.5 17 Al(OH), solid 1 Some of the constants have been corrected to concentration or mixed constants a t ionic strength 0.1 N.
6 40 c
Table 11. Calibration Results 3
0.1
Concentration of Ala (mol dm-3)
Volume of mol dm-3 NaF
at EMF = +10 m V (ml)
Volume difference AT/ (ml)
1.00 2.40 3.32 4.05 4.95 6.15 7.96 10.00
0 1.40 2.32 3.05 3.95 5.15 6.96 9.00
Flgure 1.
Distribution of AI-F-S04 species among complexes as a percent of the added concentration of each vs. pH
0.0 0.46 x 0.82 x (--) AI complexes, (- - -) F, (- - -) SO4. The coefficients indicate the stoi1.03 x chiometry used in calculatingthe percentage distribution of each component 1.26 x 1.74 x 2.30 x I ’
10-3
10-3 10-3 10-3
10-3
10-3 3.00 x 10-3
a
Sample volume = 100 ml.
Table 111. Effect of Interfering Ions on Aluminum Determination 4
Substance added, mg/l. Ca, 20 Ca, 100 Ca, 200 Ca, 300 Mg, 30 Fe(III), 50 SO,, 1500
0
1C
Volume
0,l mol dm3
rnl. YoF
Figure 2.
Titration curves of aluminum with F- in sodium acetate buffer, pH 4. Dotted line corresponds to the F / A I ratio equal to 3
were used to prepare all solutions. Standard fluoride solution, 0.1 mol dm-3, was prepared by dissolving 4.20 g of sodium fluoride, 2.97 g of sodium acetate anhydrate, and 10.11 mol of glacial acetic acid in distilled water and diluting to 1 1. Standard alum solution, mol dm-3 or 100 ppm Al, was prepared by dissolving 3.7 X 1.234 g of .&(S04)3.18 H20 in 1 1. of distilled water. Sodium acetate buffer pH 4 was prepared by dissolving 62.3 g of sodium acetate anhydrate and 212.40 mol of glacial acetic acid in distilled water and diluting to 1 1. (16).
AI(II1)
AI(II1)
added, mg/l.
fo&d, mg/l.
30 40
30.3 41.6 83.6 35.6 50.4 15.5 60.9
80
35 50 15 60
Procedure. The standard titration procedure is as follows: add 5 ml of sodium acetate buffer to 100 ml of sample (either A1 solution or water for “blank”). Titrate with sodium fluoride solution until the emf of the F- electrode is at least t 10 mV relative to the calomel cell. This corresponds to approximately 5 X mol dm-3 F-. The calibration curve is prepared by titration of A1 standards and of a blank solution without Al. The difference in volume of added fluoride between the standards and the blank solution at emf = +10 mV is then plotted as a function of the aluminum concentration. In subsequent determinations of unknown Al, the volume of 0.1 M F- titrated to emf = +10 mV is corrected for the “blank” titration and the aluminum concentration is read from the calibration curve. For low levels of Al, below 10 mg/dm3, it may be more convenient to titrate with more dilute NaF, e.g., 0.01 mol dm-3, or to use a microburet. ANALYTICAL CHEMISTRY, VOL. 48, NO. 4, APRIL 1976
777
RESULTS AND DISCUSSION When fluoride ion is added in excess to aluminum in slightly acidic solutions, p H 4 to 5 , t h e formation of the series of fluoro complexes reduces the free A13+ concentration and, hence, prevents precipitation of aluminum hydroxides and also reduces the extent of formation of polymeric hyFT droxo complexes. For example, for A ~ Tand precipitation does not occur below p H 6 and the polymeric species, e.g., Als(OH):,f form at most 1%of t h e total aluminum Concentration. Consequently, under t h e conditions of our method, all aluminum complexes are thought to be monomeric. This means that, for a fixed activity of F-, t h e ratio of F bound t o aluminum is independent of the total concentration of Al. This is confirmed by the results in Figure 2. If the amount of fluoride bound to aluminum, as calculated from t h e difference in volume a t constant emf between the aluminum standards and the blank solution, is plotted vs. aluminum concentration over the range -20 mV t o +20 mV, linear lines are obtained. T h e slope of these lines is the ratio of F- bound per Al. At -15 mV F-/Al N 2.6 and a t +10 mV F-/Al N 3.0. T h e constancy of the F- bound per A1 is demonstrated by the calibration results in Table 11. From the slope of the A V vs. A1 concentration plot, the bound F- per A1 = 3.0. Several ions commonly present in mill waters, such as Ca2+, Fe", and Mg2+, form complexes or insoluble solids with fluoride and may therefore be a source of interference with this ion selective electrode method. Consequently, samples containing varying amounts of aluminum were prepared with these interfering ions in excess of concentrations frequently found in practice. Sulfate ions were also included because of possible complexation with A1 according t o Table I. The results of these tests as given in Table
I11 indicate t h a t calcium is the only ion which does cause interference; however, its effect is small provided t h e concentration of Ca is below 200 mg/l. The accuracy of the determination is excellent and t h e relative standard deviation as determined on repeated (7) titrations of samples containing 10 mgh. aluminum was 1.1%. LITERATURE CITED (1) R . M. K. Cobb and D. V. Lowe, TAPPI, 38, 49 (1955). (2) A. W. McKenzie, V. Balodis, and A. Milgrom, Appita, 23,40 (1969). (3) J. D. Hem et al., "Chemistry of Aluminum in Natural Water", U.S. Geological Survey Papers 1827-A (1967), -B (1968), -C (1969), -D (1972), -E (1973). U S Govt. Printing Office, Washington, D.C. 20402. (4) P. L. Hayden and A. J. Rubin, in "Aqueous-Environmental Chemistry of Metals", A. J. Rubin, Ed., Ann Arbor Science, Ann Arbor, Mich., Chap. 9. (5)J. P. Casey, "Pulp and Paper", Vol. 11, Interscience, New York, 1960, p 1056. (6) K. E. Shull and G. R . Guthan, J. Am. Water Works Assoc., 59, 1456 (1967). (7) R. C. Turner Can. J. Chern., 47, 2521-2527 (1969). (8) S. H. Watkins, TAPPI, 45, 216A (1962). (9) R. E. Mesrner and C. F. Baes inorg. Chern., I O , 2290 (1971). (IO) R . E. Mesrner and C. F. Baes, "Hydrolytic Behavior of Toxic Metals", Atomic Energy Commission Oak Ridge, Tenn.. ORNL-DWG 731441 (1974). (11) L. G. Sillen and A. E. Martell, "Stability Constants", Chemical Society, London, Special Publication No. 17 (1964) and No. 25. (12) E. W. Baumann, Anal. Chern., 42, 110 (1970). (13) D. D. Perrin and I. G. Sayce, Taianta, 14, 833 (1967). (14) R. 0. James and E. Bate, unpublished results. (15) C. Brosset and J. Orring, Sven. Kern. Tidskr., 55, 101 (1943). (16) D. D. Perrin and B. Dempsey. "Buffers for pH and Metal Control", Chapman & Hall, London, 1974.
RECEIVEDfor review September 16, 1975. Accepted December 16, 1975. R.O.J. wishes to thank the Australian Government for a n ARGC Research Fellowship, and A.H. thanks Australian Paper Manufacturers P t y L t d for a Postdoctoral Fellowship.
~~
~
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CORRESPONDENCE
Refractive Index Anomalies in Stopped-Flow Measurements Sir: T h e purpose of this communication is to alert users of certain designs of stopped-flow instruments to a potential source of error in measurements made a t temperatures other than ambient. Chattopadhyay and Coetzee ( I ) have warned of a n error arising from measurements on solutions originating in the Kel-F valve block of a Durrum Model D-110 stopped-flow spectrophotometer. The poor thermal contact between such solutions and t h e thermostating bath gives rise to systematic errors in the temperature a t which measurements are made. Therefore, it has been recommended ( I ) that a minimum volume of 0.34 ml be dispensed from each drive syringe in each mixing experiment. We have encountered a different source of error on using this instrument a t temperatures other than ambient with 0.375 ml dispensed per drive syringe. In the course of kinetic studies in 1,2-dichloroethane (DCE) solvent a t temperatures above ambient, an apparent formation and disappearance of a strongly absorbing intermediate species was observed. Further studies revealed t h a t this phenomenon was an artifact which appeared a t 778
*
ANALYTICAL CHEMISTRY, VOL. 48, NO. 4, APRIL 1976
any visible wavelength on mixing DCE with DCE. A similar effect had been noted previously by Imamura ( 2 )in dichloromethane solvent. Stynes and James (3) have also mentioned an anomaly observed in toluene solvent. The upper curve in Figure 1 shows the absorbance vs. time trace observed a t 621 nm on mixing DCE with DCE a t 35 "C using the Durrum stopped-flow spectrophotometer equipped with Kel-F stopped-flow cell. The lower curve in Figure 1 shows the trace obtained without mixing both before and 1 min after the mixing experiment. Several further characteristics of this phenomenon were noted. A very similar phenomenon was observed on mixing DCE with DCE a t 45°C using the Kel-F dual-purpose stopped-flow temperature-jump cell. No change in absorbance was observed on mixing DCE with DCE a t 45 "C using the stainless steel cell, however. The magnitude of the maximum absorbance change during this phenomenon was found to be roughly in direct proportion to the difference between the thermostat temperature and the ambient temperature. The phenomenon was found to invert, giving apparent negative absorb-
