point. Hon-ever, poor mixing or slowness of reaction could raise the concentration of unreacted titrant abnormally a t the electrodes and thereby shift the end point of the titration curve ahead of the true equivalence point. These opposing errors tend to cancel each other. I n any case, the remaining error is cancelled in the standardization. Acids were titrated in both aqueous and nonaqueous (ASTM D664) solvents at the normal speed and a t the slowest speed setting to estimate the amount of overshooting The normal speed titration took 2.5 minutes; the slow speed titration took 8 minutes. With the aqueous solution the fast titration consumed 0.16% more titrant than did the slow titration. With the nonaqueous solvent the fast titration used 0.20y0 more titrant than did the slow titration. The precision of the instrument was checked by replicate titrations of 50-ml. portions of aqueous 0.05M potassium acid phthalate with 0.1N KOH titrant. Each of the titrations required 2.5 minutes. The standard deviation of 26 determinations was 0.064%. A similar series of 8 determinations of cyclohexanepropionic acid in nonaqueous ASTM D664 titration solvent
gave a standard deviation of 0.08%. Thus, the instrument performs about as well with the nonaqueous as with the aqueous solution. The current drawn by the grid of the input tube of the instrument was measured by connectins a 100,000megohm resistor between the two electrode sockets. The grid current was calculated from the shift in the pH reading that occurred when this resistor was shorted: picoamperes
=
pH shift X 5.8 X 1O’O
ohms
The current drawn by this instrument was 0.05 picoampere. Six popular commercial pH measuring instruments drew from 5 to 200 times as much current. The low grid current (high input impedance) makes the instrument particularly suitable for titrations in solvents of low conductivity. The balanced arrangement of the input tubes and the use of regulated direct current for the filaments stabilize the circuit. After a warm-up time of 30 minutes the drift rate is about 0.1 pH unit per hour. Checking a t the test points and adjustment of R2 is normally done monthly. The output of
this type of amplifier is not strictly linear; the maximum error is about 0.1 pH unit at the extremes of the scale. This is negligible for titration work. The instrument has been used to titrate 5800 samples over a period of 14 months. In this period the contacts of switch S-1 required cleaning once and one 12AX7 tube in the power supply failed. It is planned to install O.lFf. condensers across the switch contacts to minimize arcing. ACKNOWLEDGMENT
The authors thank S. L. Duncan for his design of the regulated power supply. LITERATURE CITED
(1) ASTM Standards on Petroleum Prod-
ucts and Lubricants Method D66458, “Neutralization Number by Potentiometric Titration,” Am. SOC.Testing Materials, Philadelphia, Pa , 1960. ( 2 ) Kelley, M. T., Fisher, D. J., Wagner, E. B., ANAL.CHEM.32, 61 (1960). (3) Ph;flips, J. P., “Automatic Titrators, pp. 55-65, Academic Press, Sew York, 1959. (4) Robinson, H. A,, Trans. Electrochem. SOC.92, 445 (1947). RECEIVEDfor review October 6, 1960. Accepted January 4, 1961.
Determination of Acridine by Potentiometric Titration in Acetic Acid P. RAMACHANDRA NAIDU and V. R. KRISHNAN Chemical laboratories, Sri Venkateswara University, Tirupati, South India
b Acridine dissolved in glacial acetic acid has been determined potentiometrically b y titration with perchloric acid. The method is rapid and more convenient than other methods described. Only the presence of bases of comparable strength interferes with the titration.
A
is of interest as a coal-tar component, spot-test reagent, fluorescent indicator, and the parent substance of a number of pharmacologically active compounds (1). Very few methods are available for its estimation. In the gravimetric method, it is weighed as the bisulfite compound (5). I n volumetric estimations, it is precipitated as the picrate and the product titrated with methylene blue (a), or an alcoholic solution of the compound is titrated directly with sulfuric acid (4). Titrations carried out in our laboratory by the second method gave results which were not reproducible. CRIDINE
Weak bases show an increase in strength when dissolved in acid solvents; therefore acridine (pK 5.6) should exhibit a greater relative base strength when dissolved in acetic acid. The volumetric estimation of acridine has been investigated by potentiometric titration with perchloric acid in glacial acetic acid medium. EXPERIMENTAL
Reagents. Glacial Acetic Acid. Reagent grade glacial acetic acid was refluxed with the amount of acetic anhydride required to react with the water present for about 8 hours. Refluxing was continued for another 6 hours after the addition of 2% by weight of chromic anhydride. The acid was distilled rapidly using a fractionating column, and the fraction boiling a t 115-16” C. was collected. Perchloric Acid. A solution in acetic was prepared by a modification of the method of Hall and Conant (3). A 60% aqueous solution was added slowly to the requisite amount of
chilled acetic anhydride to react with the water present, and diluted further to the desired concentration with the glacial acetic acid described above. This solution of perchloric acid was standardized by potentiometric titration with a standard solution of sodium acetate in acetic acid (3). Chloranil (Eastman Kodak). Hydrochloranil prepared from chlorani1 by reduction with phosphorus and iodine. Acridine. A British Drug Houses laboratory reagent sample, melting point 110’ C. Procedure. Samples of acridine were weighed out and dissolved in glacial acetic acid t o yield ca. 0.1N solutions. A mixture of chlorani1 (0.6 gram) and hydrochloranil (0.8 gram) was added and a platinum rod, inserted in the solution, served as the indicator electrode; a saturated calomel electrode (S.C.E.) was used for reference and a saturated solution of lithium nitrate in glacial acetic acid functioned as the salt bridge. The solution was titrated rapidly with standard perchloric acid and the change VOL 33, NO. 4, APRIL 1961
497
Table 1.
KO. 1 2 3 4 5 6
Titration of Acridine with Perchloric Acid
Acridine, Mg. Taken Found 182 I57 98 49 45 27
2 8 6 5 8 8
181 8 157 4 98 G 49 2 45 5 273
DISCUSSION
0.65 r
The estimation of acridine by potentiometric titration with perchloric acid in glacial acetic acid medium is accurate within 0.5 mg. on quantities of acridine ranging from 28 to 180 mg. Furthermore, the method is more rapid thac sther available methods. Estimatioc of acridine in the presence of other bases of similar strength was found to be not feasible by this method, although nonbasic additives did not interfere. Excess of acetic anhydride had to be avoided in the glacial acetic acid used as solvent, as its presence tended to give a lower value for acridine found.
Error, Mg. -0 4 -0 4
0 0
-0 3
-0 3 -0 5
in e.m.f. during the course of the titration was measured using a Mullard potentiometer unit. Equilibrium was quickly established after the addition of each lot of perchloric acid and the titration could be completed in 10 minutes. RESULTS
Figure 1 shows the e.m.f. (us. S.C.E.) plotted against milliliters of perchloric acid (0.0098N) in a typical case. The exact end point was located by computing the volume a t which the second derivative of the increments of e.m.f. with respect to volume increments, A2E/AVZ, becomes zero (6). The ratio of e.m.f. increments to volume increments was calculated for each 0.1
0
2
-
Z
$0.4
!i
L
0.3 50
10
20
30
d0
ML.OF PERCHLORIC ACID Figure 1. Titration of acridine with perchloric acid
ml. of reagent added and this ratio, AEIAV, was plotted against volume of the reagent to obtain the second derivative. The results obtained in the series of experiments are given in Table I.
LITERATURE CITED
(1) Albert, A., "The Acridines," p. 229, Edward Arnold & Co., London, 1951. (2) Bolliger, A., Analyst 64, 416 (1939). (3) Hall, N. F., Conant, J. B., J . Am. Chem. SOC.49, 3047 (1927). (4) Khmelevskif, V. I., Ovchinnikova, I. I., Org. Chem. Ind. (U.S.S.R.) 7, 626 (1940). (5) KhmelevskiI, V. I., PostovskiI, I. Ya., J . Aaal. Chem. (U.S.S.R.) 17., 463 , (1944j (6) Lingane, J. J., "Electroanalytical Chemistry," p. 70, Interscience, New York, 1953.
:
RECEIVEDfor review July 29, 1960. Accepted November 21, 1960.
Apparatus for High Pressure Polarography with Application to Polarography in Liquid Ammonia at 25' C. WARD B. SCHAAP, ROBERT F. CONLEY,' and F. C. SCHMIDT Deparfmenf of Chemistry, Indiana University, Bloomington, Ind.
b Apparatus is described for use in polarographic studies with a dropping mercury electrode a t elevated pressures. The apparatus was tested in liquid ammonia at 25" C. and 10-atm. pressure. The apparatus can b e used, in general, in the study of polarography in solvents a t temperatures above their normal boiling points and allows polarographic and derived thermodynamic data in different solvents to b e compared a t a standard temperature.
T
HE use of the dropping mercury electrode for polarographic studies is restricted by the physical properties of mercury to temperatures ranging from -39" C., its freezing point, to nearly 356' C., its boiling point. In actual practice, the dropping mercury electrode has been used in liquid am-
498
ANALYTICAL CHEMISTRY
monia a t -36" C. (3-5, 7) and in fused salt media at 160' to 220" C. (8). At higher temperatures solid microelectrodes have generally been used in polarographic studies. iittempts by Steinberg and Nachtrieb to retain the advantages of a dropping electrode by replacing mercury with a higher boiling metal, such as lead, bismuth, or silver, were unsuccessful (8). Recently, however, Heus and Egan reported using a dropping bismuth electrode successfully a t 450" C in a fused LiC1-KC1 eutectic melt ( 1 ) . At temperatures below the freezing point of mercury, which may be encountered in nonaqueous polarography in low boiling solvents, solid microelectrodes are again a possibility. In these latter cases, however, it is also possible to raise the boiling point of the solvent above the freezing point of
mercury by allowing the pressure of the system to increase sufficiently. In this paper, apparatus is described which allows polarographic studies with a dropping mercury electrode to be carried out a t elevated pressures. The apparatus was tested in polarographic studies in liquid ammonia at 25" C. At this temperature, nearly 60" above its normal boiling point, liquid ammonia develops a pressure of approximately 10 atm. The accepted standard temperature for thermodynamics of aqueous systems is 25' C. Jolly (6) has advanced reasons for the choice of 25' C. as the standard temperature for thermodynamics in liquid ammonia. the principal
1
Present address, Georgia Kaolin Co.,
433 Yorth Broad St., Elizabeth, N. J.
