Determination of Hydrofluoric Acid in Nitric-H yd rofluoric Acid Mixtures S. A. LONG Bell Aircraft Corp,, Buffalo 5, N. Y. A procedure is described for the determination of hydrofluoric acid in nitric-hydrofluoric acid mixtures. After a specific gravity determination, the sample is treated with hydrogen peroxide, neutralized with sodium hydroxide, and titrated with standard aluminum chloride.
A
NUMBER of methods have been proposed for the deter-
mination of fluoride ion. Kolthoff ( 4 ) reported procedures for the determination of fluorides with aluminum chloride and cerous nitrate. The end point of the cerous nitrate procedure is delicate and its determination requires a considerable amount of experience. Furthermore, serious interference with Ce+++ and F- ions occurs and the end point is affected by high concentrations of sodium nitrate. Kurtenacker and Jurenka ( 5 ) reported that when a neutral fluoride solution is titrated with a solution of aluminum chloride, hydrolysis occurs upon addition of excess aluminum chloride and the system becomes acid. Saylor and Larkin (6) reported that Geyer ( 3 ) reviewed the current literature on methods for the determination of fluoride and indicated that the titration of fluoride ion with aluminum chloride, using methyl red indicator, was the best method. It was decided to investigate the aluminum chloride method and to develop a procedure applicable to the determination of hydrofluoric acid in nitric-hydrofluoric acid mixtures. EXPERIMENTAL
Aluminum chloride solution. The reaction of aluminum chloride with hydrofluoric acid in neutral solution may be assumed to occur according to the following equation: Al+++
+ 6F- % AlFs---
Upon addition of excess aluminum chloride solution, hydrolysis occurs according to the following equation: AIC&
+ HOH % ill (0H)a + 3HC1
Therefore, not only must the sample be neutral before analysis, but also the aluminum chloride must be checked for its acid content. Feigl and Krausz ( 2 ) suggest the addition of excess oxalate for subsequent titration with alkali. Small amounts of acid, however, cannot be determined in the usual way using this procedure (1). Another method ( 4 ) involves the addition of excess potassium fluoride previously neutralized to phenolphthalein. Since the complex fluoaluminate reacts neutral to phenolphthalein, the free acid may be titrated to that indicator with alkali. Several runs using the latter method indicated such a small acid content that it was decided to run a blank on reagents.
drops of methyl red indicator were then added and, using a microburet, the solution was titrated with 0.16M aluminum chloride solution, a t a temperature between 70" and 80' C., to the methyl red end point. The hydrofluoric acid titer was then calculated. Standardization against Hydrofluoric Acid. Five to six drops of pretiously analyzed 48%-by-weight hydrofluoric acid ( 7 ) were transferred to a waxed weighing bottle. After the weighing bottle was covered, a weight was taken. The contents of the weighing bottle were transferred quantitatively to a 350-ml. platinum dish containing 50 ml. of a 10% solution of nitric acid. The weighing bottle was reweighed and the amount of hydrofluoric acid added was obtained by difference. Five milliliters of 3% hydrogen peroxide were added to oxidize the dissolved oxides of nitrogen and the solution was neutralized with strong sodium hydroxide in order to minimize the volume of solution. The solution was then saturated with solid sodium chloride and 0.5N hydrochloric acid was added until the solution was just acid. After heating to between 70" and 80" C., the solution was titrated with 0.01N sodium hydroxide (carbonate-free) to the phenolphthalein end point. Two to three drops of methyl red were added and the resulting solution was titrated with standard aluminum chloride to the red end point. The blank on the reagents was found to be 0.15 ml. of aluminum chloride. PROCEDURE FOR HYDROFLUORIC ACID DETERWNATION
Determine the specific gravity of the sample by any convenient method. Transfer 5 ml. of the sample to a 250-ml. platinum dish containing 40 ml. of recently boiled distilled water containing 5 ml. of 3% hydrogen peroxide. Transfer the platinum dish with contents to a cool water bath and add 2 drops of phenolphthalein solution. Add 6.5 grams of sodium hydroxide, dissolved in a minimum amount of water, to the solution. Add dropwise, 0.5N sodium hydroxide until the indicator shows just pink, and follow with the addition, all a t once, of 30 grams of sodium chloride. Add dropwise, 0.5N hydrochloric acid until the pink just disappears. Heat to 70' to 80' C. and titrate to the first appearance of pink, using 0.01N sodium hydroxide. Add 2 to 3 drops of methyl red indicator. By. means of a microburet titrate the solution with standard aluminum chloride solution to the red end point. Keep the temperature of the sample in the range of 70" to 80" C. during titration. Calculations. Ml.of AlC13X H F titer Specific gravity X 5
Standardization against Sodium Fluoride. A known aliquot of the standard fluoride solution, containing 0.002 gram of sodium fluoride per milliliter of solution, was transferred to a 350. ml. latinum dish and saturated with sodium chloride to repress the lydrolysis of aluminum chloride. The solution was heated to between 70' and 80' C. and was neutralized to the phenolphthalein end point with 0.01N sodium hydroxide. Two or three
% HF
Table I. Standardization of Aluminum Chloride against Hydrofluoric Acid with Addition of 5.0 M1. of 3% Hydrogen Peroxide and Solution Saturated with Sodium Chloride 48.3% H F Added, Gram 0,3582 0.2762
STANDARD1 ZATION
In order to check the accuracy of standardization, the aluminum chloride solution was standardized against standard sodium fluoride solution and against known amounts of hydrofluoric acid in a nitric acid medium.
=
0.5282
0.4740
H F Titer Corrected 0.0189 0.0196 0.0194 0.0188
Blank, Reagent MI. 0.15
DISCUSSION
Attempts to standardize aluminum chloride against standard sodium fluoride solution without sodium chloride saturation resulted in poor precision. When the standardization was repeated, using sodium fluoride solution saturated with sodium chloride, excellent precision was obtained. Thirty per cent hydrogen peroxide solution was first used in standardizing aluminum chloride against standard hydrofluoric 1988
V O L U M E 26, NO. 12, D E C E M B E R 1 9 5 4 Table 11. Comparison of Hydrofluoric Acid Titers Obtained against Hydrofluoric Acid and Sodium Fluoride (Hydrogen peroxide, 3%, used in hydrofluoric acid standard: sodium chloride saturation used in both cases) Standard H F Titer Reagent Gram Corrected Required, ML 0.042 N a F 0.0190 0.05 0.4702 H F 0.0189
1989 to 0.62% by Reight hydrofluoric acid have been analyzed. The average mean deviation of the results was 0.005. ACKNOWLEDGMENT
The author expresses sincere appreciation to F. R. Piccirillo and W. H. Bergdorf, of this laboratory, for their helpful suggestions and discussions. LITERATURE CITED
Am. SOC.Testing Materials, “ASTM Standards,” Part 7 , Uesignation D 1179-51T (1952). Feigl, F., and Krausz, G., Ber., 58, 398 (1925). Geyer, R., Z.anorg. u.allgem. chem., 252, 50 (1948). Kolthoff, I. M., and Stenger, V. A., “Volumetric Analysis,” Vol. 11, Interscience Publishers, New York, 1947. Kurtenacker, H . , and Jurenka, W. I., ANAL. CHEM.,82, 210
acid i n nitric acid solution. The results were scattered because the phenolphthalein indicator seemed to bleach. Table I shows the results obtained when the standardization a a s repeated using 3% peroxide. Table I1 shows a comparison of the hydrofluoric acid titer obtained in standardizing aluminum chloride against standard sodium fluoride solution and standard hydrofluoric acid in nitric acid solution. Approximately 30 samples of mixed acids ranging from 0.03
(1930).
Saylor, J. H., and Larkin, 11.I?., ANAL.CHEbf., 20, 194 (1948). “Scott’s Standard Methods of Chemical Analysis,” 5th ed., Vol. 11, p. 2209, D. Van Nostrand Co., New York, 1939. RECEIVED for review April 20, 1954.
Accepted August 17, 1954.
Reactions of Some lewis Acids with a Series of Simple Basic Indicators in Aprotic Solvents DAVID L. HAWKE and JOSEPH STEIGMAN o f Chemistry, Polytechnic Institute o f Brooklyn,
Department
The reactions of a number of Lewis acids with a series of uncharged basic indicators were qualitatively examined in the solvents benzene, chlorobenzene, and chloroform. A qualitative order of acidity was established, which was similar to the order of catalytic activity reported for several Friedel-Crafts reactions. The reaction of the indicator anthraquinone with the acid aluminum bromide in chlorobenzene was examined spectrophotometrically. I t was concluded that the aluminum bromide could react with one or both oxygens in the anthraquinone.
T
H E present research is part of a program intended to establish a scale of the relative acid strengths of a number of Lewis acids. This work deals with the reactions of these acids with a selected series of indicators in the solvents benzene, chlorobenzene, and chloroform. A good deal of work has already been done on reactions of such acids with certain indicators in solvents like benzene, chlorinated hydrocarbons, and carbon disulfide (9). Initially, Lewis carried out titrations of bases like pyridine with acids such as boron trichloride and stannic chloride in organic solvents containing well-known indicators ( 7 ) . I n addition, he showed that the absorption spectrum of the acid form of methylene blue was the same whether acidification was performed by aqueous strong acids or by stannic chloride in chloroform (8). More recently a n investigation was reported of the color changes of a number of traditional indicators on the addition of acids and bases in aprotic solvents, and a relative order d acidity was established (12). For the present work, it was decided that the indicators should all be of the same charge type, reacting similarly with the various acids. They should possess known basicity constants in some solvent in relation to one another. Finally, the basicity constants should be sufficiently lar apart so that a wide range of acid strengths could be observed. The indicators which are ordinarily used in water do not fulfill the first and the last of these requirements. All three requirements are met by the simple un-
Brooklyn,
N. Y.
charged basic compounds invrstigateJ by Hammett in hie studies of the acidities of solvents ranging from water to concentrated sulfuric acid (4-6). These substances are of the same charge type, all adding one proton M hen dissolved in concentrated sulfuric acid, according to the equation:
B
+ Hi304
BH+
+ HSO4-
in which B represents the uncharged base, the BH+its conjugate acid. Secondly, the strengths of these indicator bases were measured in appropriate sulfuric acid-water solutions or perchloric acid solutions; they are expressed in terms of the acidity constants of the conjugate acids:
an equation in xhich activities are required, but in which a simple colorimetric measurement of concentrations will suffice because interionic effects are small in the sulfuric acid systems, and activity coefficient ratios are the same for the overlapping system of indicators which Hammett employed. Thirdly, the base strengths of these indicators cover a very wide range-and in a direction suitable for investigating stronger and stronger acids. Six of these indicators were selected for the present research. They are named in Table I, together with the acidity constants, the basicity constants, and the range of aqueous acid in which each indicator is useful.
Table I.
Code Letter 4 B C D E
F
Indicator p-Aminoazobenzene Benzeneazodiphenylamine p-Chloro-o-nitroaniline 2 ,CDinitroaniline Benaalacetophenone Anthraquinone
Indicators Approx. Ionization Acidity Constant Constant, as Base pKa in Water +2.80 fl.52 -0.85 -4.38 -5.61 -8.15
Acid Range
