A Method of Analysis for Fluoride J
Application to Determination of Spray Residue on Food Products W. M. HOSKINS AND C. A. FERRIS, University of California, Berkeley, Calif.
T
HE widespread use of fluorine compounds as insecticides and the establishment by the Federal Government of a tolerance, or upper limit to the amount of added fluorine which may be present on or in foodstuffs, have combined to stimulate interest in the quantitative determination of minute amounts of this element. Consequently, several dozen publications on various methods of analysis have recently appeared. Work on the subject was undertaken in this laboratory in connection with a study of .the amounts of arsenic, lead, and fluorine which occur on fruit sprayed in various ways for the control of the codling moth in California. Preliminary experiments with several suggested methods indicated the superiority of the reaction between fluoride and thorium ions in the presence of an indicator sensitive to excess thorium. Accordingly, an intensive investigation has been made of Armstrong’s ( 1 ) modification of the Willard and Winter (14) method which involves the titration in 50 per cent alcohol of fluoride ion with standard thorium solution in the presence of sodium alizarin sulfonate. For comparison, the Sanchis (9) method adapted from that of Thompson and Taylor ( I d ) , which depends upon the bleaching effect of fluoride ion upon a zirconium-alizarin complex, has been used in certain cases.
solutions are allowed t o stand. This is undoubtedly due to slowness of one or more of the reactions which are involved in the equilibrium between T h + +,I?-, and sodium alizarin sulfonate. I n the somewhat analogous method used by Thompson and Taylor ( I d ) and by Sanchis (9)-i.e., the bleaching of red zirconium alizarin sulfonate by fluoridethe solutions are brought to boiling to hasten the establishment of equilibrium, which is said to require 48 hours in the cold. --+
Effect of Acidity and Use of Buffer Previous workers have paid scant attention to the acidity of the solution during titration. The usual directions have been to destroy the pink color with 1 to 50 hydrochloric acid solution (0.24 N ) before adding the thorium solution. For precise work, however, the acidity is important. If the solution is decolorized, but no excess acid is added, extremely low values are obtained. The addition of one drop of acid in excess usually is satisfactory, but larger amounts of acid lead to the use of too much thorium solution (Table I). Vanselow (13) of the Citrus Experiment Station has found the same effect for excess acid. While it is possible to control the acidity by careful use of dilute hydrochloric acid, the use of a buffer offers advantages. Preliminary experiments, using the indicator bromophenol blue, indicated that the most desirable p H was approximately 3.5. In order to give this p H a t half neutralization-i. e., a t maximum buffering powera buffer acid must have a paK of the same magnitude. The data of Jukes and Schmidt (6) show that paK for acetic acid is increased by about 0.7 in passing from water to 50 per cent alcohol. Hence, it seemed probable that the paK of the related acid, monochloroacetic, which is 2.8 in water, would be approximately 3.5 in a mixture containing equal parts of commercial 95 per cent alcohol and water-i. e., about 48per cent alcohol. The behavior of chloroacetic acid in 48 per cent alcohol was investigated by determining the pH of solutions containing 2.00 cc. of 1 M monochloroacetic acid solution and varying volumes of 1 M sodium hydroxide solution in a total volume of 100 cc. A hydrogen electrode was used and a glass electrode was standardized in the same solutions. The results are given in columns 1 and 2 of Table I.
Titration with Standard Thorium Solutions Standard fluoride solutions were prepared from recrystallized c. P. sodium fluoride. Their accuracy was checked by the standard method of precipitation as CaFz (IO). Stock solutions were made 0.05 M and dilutions t o lower concentrations as needed. Very dilute solutions of sodium fluoride, such as 0.001 M , decrease in strength if kept in glass vessels, but if the containers are lined with paraffin the concentration remains constant for long periods. Thorium solutions, approximately 0.05 M, were made from c. P. thorium nitrate and were standardized by precipitating as thorium oxalate, heating, and weighing as the dioxide. The solutions were then diluted to 0.01 M or 0.001 M for use. These dilute solutions were made at intervals of 2 t o 3 weeks, though they appear t o be stable for longer periods. Comparison of the dilute standard solutions of fluoride and thorium by the method to be described gave excellent agreement. The first experiments followed the procedure of Armstrong. It was soon found that three factors must be carefully controlled in order to obtain duplicable results: the volume of indicator added, the intensity of color in the blank and unknown, and the acidity. The first two are inseparably connected, for the more indicator used the smaller is the volume of thorium solution required to produce a given intensity of color in the blank. As judged by the eye, the pink color is most sensitive to changes in the amount of thorium added when its intensity is very low. It is not advisable to use large amounts of indicator in order t o reduce the thorium needed in the blank, because the pink tint is masked by the yellow color of the indicator where much is used in acid solution. Summarizing a large number of experiments, it may be said that when a volume of 50 cc. is used for the blank, 0.040 cc. of a 0.05 per cent aqueous solution of sodium alizarin sulfonatei. e., the concentration of the indicator in the solution is 4 X per cent-plus 0.040 t o 0.070 cc. of 0.001 M thorium solution gives a faint pink color which is very suitable for matching and is very sensitive to thorium ion. For other volumes in the blank, proportional amounts of the reagents should be used. It is advisable to run a fresh blank with each unknown, for the color slowly becomes more intense as the
TABLEI. PH OF CHLOROACETIC ACID-SODIUVHYDROXIDE BUFFERS (0.02 M ) IN 48 PERCENTALCOHOL Molecular Ratio Acid:Base 1:0.2 1:0.3 1:0.4
PH 2.95 3.39
1:0.6
1:0.6 1:0.7 1:o.s
3.73
4.17
Relative Volumes of Th + * + + Solution Blank and test solution a t same buffer Blank a t constant ratio, volume a t 0.5 buffer ratio 0.5 taken as unity 1.050 1.411 1.000 1,292 1.000 1.171 1.000 1,000 1,000 0,975 0.957 0.903 0.811 0.768
A plot of these results shows that a t half neutralization the p H is 3.55, which is accordingly the paK value of chloroacetic acid in equal parts by volume of water and commercial ethyl alcohol. The corresponding dissociation constant, Ka, is 2.8 x 10-4. It was necessary to prove that this buffer is able t o maintain a constant pH during the titration with tho6
JANUARY 15,1936
ANALYTICAL EDITION
rium solution. Typical titrations were made with varying amounts of fluoride ion present and the various buffer mixtures present a t 0.02 M concentration. The hydrogen electrode did not function in the presence of the indicator, but the glass electrode gave very steady potentials exactly equal in all cases to those previously obtained in the buffer solutions alone. Since it may be necessary to analyze samples containing considerable acid or base, the most desirable procedure is to bring the solution to the approximate transition point of the indicator with dilute acid or base-e. g., 0.05 hJhydrochloric acid or sodium hydroxide-before adding the buffer solutions. The third column of Table I shows how greatly the volume of thorium solution required to obtain an end point is altered a t various pH values of the test solution when the blank is held at a constant buffer ratio of 0.5. It is obviously not possible to obtain accurate results except over a very narrow pH range. I n the last column it is shown that when the blank and test solutions have the same pH a much wider range is allowable. The buffer ratio 0.5, a t which the pH is 3.55, is seen to be in the middle of the favorable range. The above results a t once raise the question of how greatly the pH of a 48 per cent alcohol solution varies when treated with a dilute solution of hydrochloric acid as in the ordinary adjustment of acidity according to the directions of Willard and Winter (14) and of Armstrong (1). For the sake of greater delicacy, tests were made in which 0.05 N hydrochloric acid was used instead of the 0.24 N (1 to 50) recommended by the above workers, so that the changes in pH were smaller than with corresponding volumes of stronger acid solution. The results are given in Table 11. A comparison of Tables I and I1 will show what very serious errors result from the inevitable inaccuracy involved in trying to add one drop of 0.24 A T acid in excess to an unbuffered solution.
7
to their procedure of standard fluoride solutions ranging in concentration from 0.0004 M to 0.004 M . The limiting factor seems to be the slowness with which equilibrium is reached near the end point. As mentioned before, the color of the blank slowly becomes more intense. For larger amounts of fluorine the uncertainty thus introduced is negligible, but with small amounts the unavoidable difference in age of the blank and test solutions may cause an appreciable error. This may be eliminated to some extent by using a second blank which is matched against the sample. The average of the volumes of thorium solution used in the two blanks is used. Attempts to speed up the attainment of equilibrium by running the titrations a t 60" C. were not successful, for the red thorium alizarin sulfonate was rapidly coagulated. There is no a priori reason for using a volume as large as 50 cc. for titration, so a series of experiments was made in which the volume was made up to 5 cc. Small matched test tubes were used and it was found easy to obtain the same accuracy with 6 to 90 y of fluoride as with the larger amounts in 50 cc. For purposes of comparison with an entirely different method of analysis two series of solutions from fruits and vegetables treated with fluoride sprays were tested by the present method and by that of Sanchis (9). Over the range from 4 to 400 y of fluorine in 25 cc. the latter method gave results which average 3.6 per cent higher than those obtained by titration. I n many cases agreement was very close, but occasionally samples run by the Sanchis method were quite different in color from the standards.
Interfering Substances
When fluoride is present in solutions containing possible interfering ions or compounds, the recommendation is usually made that it be isolated by distillation from perchloric or sulfuric acid solution. Under such conditions, however, the EFFECT OF PH UPON COLOR OF SODIUM ALIZARIN TABLE11. halides, nitrate, sulfite, and other rarer contaminants are also SULFONATE IN 50 PER CENTCOMMERCIAL ALCOHOL distilled and the possibility always exists that nonvolatile 0.085 0.146 0,030 0.045 0.075 Volume of Acid 0.000 4.8 4.1 6.3 5.2 8.8 7.2 PH substances will be carried into the distillate as spray. For Very Pale Yellow Pink Pale Very Color these reasons a study has been made of the effects of a number pink pale pale yellow of ions a t various concentrations in order to determine, a t pink yellow least approximately, the maximum amounts which may be tolerated in the fluoride solution. This depends, of course, In Table I1 is shown also the effect of pH upon the color of upon the amount of fluoride used and upon the accuracy desodium alizarin sulfonate in 50 per cent commercial alcohol as sired. The tests were made with 1.00 cc. of 0.003 iM sodium determined by adding 0.05 N hydrochloric acid in small increfluoride solution in a total volume of 50 cc. Varying amounts ments to 50 cc. of solution containing 0.04 cc. of 0.05 per cent of the sodium salts of each ion were added and the pH was adsolution of the indicator and made slightly basic a t the start. justed as described above. In Table I11 the concentrations The p H was determined with the glass electrode. Kolthoff are stated a t which an appreciable increase in volume of and Furman (6) state that the color changes occur in water thorium solution was required. The precise concentrations in the pH range 3.7 to 5.2. a t which interference first develops is accordingly a little It is obvious that the equilibrium between the acidic and lower than the figures given. If a stronger pink color is tnken salt forms of the indicator occurs in the alcoholic solution in a as the end point of the titration, interference occurs a t lower more alkaline range than in water, which is in agreement with concentrations. the behavior of most weak acids.
Accuracy and Delicacy of the Titration The procedure described above has been used many times in titrations of standard solutions of sodium fluoride. The range of concentration before addition of alcohol was from 0.00012 M to 0.0016 M-i. e., 2.3 to 30.4 p. p. m. as fluorine. The total fluoride in 25 cc. therefore varied from 57 to 760 y. The maximum errors were +3.7 and -2.7 per cent, with an average error of - 1.0 per cent. No distinct variation in percentage error occurred as the Concentration of fluoride was varied over this range. This is a somewhat smaller average error than that of - 1.7 per cent which may be calculated from the data of N'illard and Winter (14)for the titration according
TABLE111. E&ECT OF VARIOUSIONSUPON TITRATION OF FLUORIDE WITH THORIUM Ion Halogens
Nos-
ClOaSOa--
Concn. for Appreciable Effect 0.1 M
AsO8---
0.002 M
AsOa--Pod---
0.1 M 0.1 M
Ion
sop--
Concn. for
Appreciable Effect 1 x 10-4 M 1 X 10-6 M
x 10-0 M
