LITERATURE CITED
Table 11. -_______--
Calcd. 149.6 214.4 212 2 143.9 263.0 3.5 142 2 58 6 61 4
a
Effect of AI/F Ratio on Silica Recovery from HzSiF6-H3P04-CaO-HF Systems"
Si02
Foiund 120 3 202.1 17!3. 2 11'7.9 220 5 12 131 4 513.6 61.8
Recovery
%
80 ~~
94 85 82 84 34 92 100 101
100 5(i.2 56 2 101 58.7 58 2 101 57.3 57.0 226.9 100 228.0 104 8.2 Na2CO3-Na2BLhfusion method. Reagent grade 49'; HF.
Al/F
HF,bmg.
... 1 1
1 1.5 1.5 3.0 3.0 3.0
3.0 3.0 3.0 4.0 4.0
Remarks K O Al
, . .
...
480 166 986
... , . . I
.
.
~o borate
... ... ...
... 184 967
(1) Berzelius, J. J., J . f i i r Cheinih und Physik 16, 426 (1816). (2) Brabson, J. A . , I)uncan, R. D.,
Murphy, I. J., ANAL.CHEM.35, 1102 (1963). (3) Cronkite, W.A,, Ibid., 35, 766 (1963). (4) Furman, N. H., "Scott's Standard Methods of Chemical Analvsis." 6th ed., 1701. 1, pp. 950-75, Van -Nostrand, New York, 1962. ( 5 ) Harel, S., Herman, E . R., Jalnii, Al., *qNAL. CHE.V. 27, 1144 (1955). (6) Hazel, IT. h l . , Ibid., 24, 196 (1952). (7) Hoffman, J . I., Lundell, (>. E. T., J . Res. S a t l . Bur. Std. 22,465 (1939). ( 8 ) Kolthoff, I. >I., Elving, P. J., "Treatise on dnalyticlal Chemistry," Part 11, Vol. 2 , pp. 107-202, Interscienre, Sew York, 1962. (9) Schrenk, W. T., Ode, W. H., ISD. ENG.CHEM... ~ X V A LEd. . 1.201 11929). (10) Shell, H. R., A u m . C ~ E M27, . 2006 (195.51 \~._.,
poor silica recovery when borate was not added to the flus. A series of mihtures was analyzed by the same procedure with variations in the aluminum to fluoride ratio and in the use of borate. Results in Table I1 indicate an optimum AI to F ratio in the fusion of about 3 to 1 for good silica
recoveries, and again demonstrate the beneficial effect of the borate addition. The Yational nureau of Standards sample S o . 56b, reportedly containing 10.1% SiOz, was analyzed using the above procedure. The SiOz found was 10.25y0 and the standard deviation was 0.0915.
(11) Shell, H. R., Craig, R. L., Ibid., 26,998 (1954).
FRED W.CZECH TEDP. HRYCYSHYN ROBERT J. FVCHS FMC Corporation Inorganic Research and Development Department Carteret, N. J.
Reinterpretation of Chronopotentiometric Study of Kinetics of a Chemical Reaction Coupled between Two Charge Transfer Reactions SIR: Alberts and Shain (2) have recently reported a n investigation of the kinetics of a chemical reaction coupled between two charge transfer reactions (the Electrical-Chemical-Electrical mechanism) utilizkg potentiostatic and chronopotentiometric techniques. I n the electrochemical system studied, p-nitrosophenol is first reduced to phydroxylaminophenol, which is not electroactive in the potential range of interest but which can split out water to form p-benzoquinoneimine which is reducible to p-aminophenol. The relative enhancement of the potentiostatic current or the chronopotentiometric transition time perniits an evaluation of the rate constant for the interposed chemical reaction, the dehydration of p-hy ciroxylaminophenol. The potentiostatic method produced values of the rate constant nhich were almost independent of the time of electrolysn. Hoaevcsr, the rate constanti evaluated fiom the chronopotrntionietric data wrre not conitant brit increased with decreasing tranqition time. Two w t i i c e q of thii behavior u. ere wggested. First, hccauqe the measurements \\we made with a hanging drop inercury electrode, the theorrtical equation ielating the rate constant to the oh-erietl tiawltion
time would not apply because it pertains only to a planar electrode (6). An attempt was made to correct the data for the sphericity effect but this did not markedly improve the constancy of the results. The second factor affecting the chronopotentiometric results is the concomitant reduction of adsorbed reactant and the charging of the electrical double layer. Though this factor was discussed by Alberts and Shain (b), no attempt was made to apply a correction. I n this communication the chronopotentiometric results are corrected for this effect. I t would be difficult to derive a correction for the effect of charging the double layer and/or reducing adsorbed reactant. -in empirical correction, first propoqed by Lingane ( J ) , has proved successful for correcting ordinary chronopotentiometric data for such paraqitic proceqses as oxide film formation (or reduction), charging of the double layer, and reduction (or oxidation) of adsorbed reactant. I n this mrthod the observed current is corrected by the equation
icoTc = iohs - Q / T (1) where Q is the quantity of electricity coniumed by all paraqitic reactions and T i i the tramition time. Equa-
tion l has also been proposed in a different form by Bard (3) who discusses its underlying assumptions. Because the fraction of the current invoh-ed in parasitic processes is not constant through the duration of the electrolysis, currents corrected by Equation l represent time average currents. Equation 1 was applied to the chronopotentiometric data for the reduction of p-nitrosophenol contained in Table I1 of reference 2. .1 value of Q equal to 1.7 pcoulombs yielded recalculated values of the rate constant which were almost independent of transition time. The recalculated results are wmniarized in Table I. The quantity i r m l is simply the value of i+ 2 which would be observed if the rate of the chemical reaction mere very large. ilberts and Shain (2j calculated irml'* from the transition time equation for a planar electrode, wing the known value of the diffusion coefficient of p-nitrosophenol and setting n q u i 1 to 4. The recalculated values of k were obtained by linear interpolation of the values of k r and i~,"*'zr1 * in an exten4ve table provided by Alhei ts (f j . The simple coirection embodied in Equation 1 greatly imln-oves the conqtanry of the value- of k . Furthermore, the average recalculated value, 0.72 VOL. 36, NO. 10, SEPTEMBER 1964
2027
Table I. Recalculation of Chronopotentiometric Determination of Rate Constant for Homogeneous Chemical Reaction Coupled to Reduction of p-Nitrosophenol
(10-3.11 p-nitrosophenol, pH
=
4.8,20% alcohol and 0.005% gelatin. i,,,, Q = 1 7 pcoulombs.)
= iobs - Q / r .
i r ml 1 2 / i r l 1 2 2,
Obs. 20 25 30 40 50 60
pa.
Corr. 19 62 24 31 28 81 37 59 45 95 53 95
T,
4 2 1
0 0 0
see. 50 45 43 706 420 281
calculated from plane Uncorr. Corr. 1 28 1 30 1 42 1 38 1 57 1 51 1 71 1 61 1 82 1 67 1 89 1 70
second-', is in good agreement with the value from the potentiostatic technique, 0.59 or 0.65 second-', depending upon the choice of i, ( 2 ) . The chronopotentiometric results should also be corrected for sphericity effects, but no rigorous method of achieving this is available. Such a correction would further decreaie all values of k . A value of Q equal to 1.7 pcoulombs brings about the greatest improvement. However, any value of Q between 1 and 2 pcoulombs will markedly improve the results. If Q is less than 1 pcoulomb,
k , sec.-1 Gncorr. Corr. 0 77 0 69 0 86 0 73 0 89 0 71 1 19 0 80 1 57 0 75 2 05 0 64 Av. = 0 . 7 2 Std. dev. = 0 . 0 5
the k's for short transition times are not sufficiently corrected, and if Q is more than 2 pcoulombs, the k for the shortest transition time is over-corrected. Ideally, Q should be the quantity of electricity required to charge the double layer from the initial potential to the potential at the tranpition time plus the quantity of electricity required to reduce ad5orbed reactant. ;Ilberts and Shain ( 2 ) report that the area of the hanging drop mercury electrodes used was typically about 0.07
square cni. If an average double layer capacitance of 20 pf. per square em. and a potential interval of 1 volt, between initial potential and transition potential are assumed, Q for charging the double layer would be 1.4 pcoulombs. To this must be added a quantity corresponding to reduction of any adsorbed reactant. I n other words, a Q of 1 to 2 pcoulombs is not a t all unreasonable. This simple correct'ion brings the chronopotentiometric result's int'o good agreement with the potentiostat'ic results and thus suggests that charging of the electrical double layer Fas the major cause of the inconstancy of the rate constants calculated by hlberts and Shain. LITERATURE CITED
(1) Alberts, G. S., Ph.D. thesis, m i vereity of Wisconsin, Madison, Wis., 1963. ( 2 ) Alberts, G. S., Shain, I., ANAL. CHEM.35, 1859 (1963). (3) Bard, A. J.;Tbid., p. 340. (4) Lingane, J. J., J . Electroanal. Chem. 1, 379 (1960). ( 5 ) Testa, -4. C., Reinniuth, W. H., ANAL. CHEM.33, 1320 (1961).
DES"
H.EVANS
Department of Chemistry Harvard University Cambridge, Mass.
Colorimetric Determination of Bromate with o-Aminobenzoic Acid SIR: Recently Macdonald and Yoe (I) described a spectrophotometric method for the determination of bromate by its reaction with o-arsanilic acid. Though the limit of identification of bromate is 0.5 fig.! the maximum esists for about 2 minutes and a large amount of nitric acid is required for the production of color. Chlorate, iodate, and many inorganic ions interfere with the determinations ( 1 ) . We found that o-aminobenzoic acid also reacts with bromate in presence of acid to produce a reddish-brown color whlch obeys Beer's law. The visual limit of identification of bromate is 5 pg. 11er ml. The intensity of the color depends upon the p H of the solution and remains constant for more than 72 hours. Chlorate, iodate, inorganic ions, and organic compounds do not interfere. The method is specific and is suitable for the determination of bromate in loa concentrations. The mechanism of the color reaction is not very clear.
in a test tube followed by 1 ml. of 1% o-aminobenzoic acid and 1 ml. of 0.15A' hydrochloric acid. The total volume was made 5.5 ml. with distilled water and solution was heated by dipping the tube in water a t 80-90' C. till a reddish-brown color appeared. The solution was diluted to 10 ml. and the colorimetric measurements were made by a Beckman photoelectric colorimeter model C using a blue filter. The experiment was repeated with different volumes of bromate solution and a curve wa5 drawn between the absorbancy and concentration of bromate. The amount of bromate from the un-
9
Y
EXPERIMENTAL
Procedure. Potassium bromate solution (0.1 ml.) containing approximately 20-25 p g . of bromate was taken 2028
ANALYTICAL CHEMISTRY
2
3
PH Figure 1 . of color
Effect of pH on intensity
known sample may then be directly determined from the curve by measuring the absorbance after bromate-o-aminobenzoic acid reaction. The effect of pH on the intensity of color produced by bromate with oaminobenzoic acid was studied by varying the pH of the solution by the addition of different volumes of hydrobefore heating chloric acid (O.l?~~v) the solution RESULTS A N D DISCUSSION
The results in Figure 1 show that intensity of color decreases with increase of pH. For accurate results, it is essential that the pH of the solution should lie between 1 to 2 and remain the same throughout the experiment. Any change in 1)H after the color was produced does not alter the color intensity. Experiments made a t pH 1.03 to 2.64 indicated that the limit of identification of bromate decreased with increase of pH. For example, 5 pg. per ml. of bromate at pH 1.03 and 125 pug. per ml. a t pH 2.64 was visible to the eye. The effect of excess o-aminobenzoic acid was also studied and it was found that about 50-fold excess of o-aminobenzoic acid does not affect the color intensity. However. n-hen a very large
