236
ANALYTICAL EDITION
more than 3.0, 6. The highest fluoride found, 5.6 parts, was in a sample from a hot spring in California. If preferred, standards made by adding known quantities of fluoride, 0.10 to 0.40 mg., to 5 cc. of ferric chloride solution and 10 cc. of ammonium thiocyanate in the same volume might be used, preparing the samples as already described. This affords a direct determination of fluoride against fluoride standards, By using such fluoride standards, and by adding to the standards the quantities of sulfate and chloride (as sodium 'salts) contained in the volume of sample used, this method may be applied to the determination of fluoride in sea waters, brines, or any water containing more than 2500 parts per million of sulfate or more than 5000 parts of chloride. For such waters, however, 10 cc. of ferric chloride must be added to both sample and standard instead of the 5 cc. used when the effects of sulfate and chloride are counteracted by the addition of ferric chloride. This is necessary in order to have enough iron left for the fluoride after the sulfate and chloride have been satisfied.
Vol. 5, No. 4
This modification is recommended only for highly concentrated waters. It is not advisable to use it in preference to that in which the effects of these constituents are counteracted by the addition of ferric chloride for waters containing less than 2500 parts per million of sulfate or 5000 parts per million of chloride because (1) if 5 cc. of ferric chloride solution are used in sample and standard, the iron left for fluoride after the sulfate and chloride have been satisfied may be so reduced as greatly to restrict the fluoride range and sensitivity; (2) if 10 cc. are used the color in water containing low sulfate, chloride, and fluoride would be too deep to read; and (3) a different set of standards might be required for each sample. (1) Foster, M. D.,
LITERATURE CITED J. Am. Chem. Soc., 54, 4464 (1932).
RECEIVED March 7, 1933. Presented before the Division of Water, Sewage, and Sanitation Chemistry at the 85th Meeting of the American Chemical Society. Washington, I). C., March 26 to 31, 1933. Published by permission of the Direotor, U. 8. Geological Survey.
Determination of Fluorides in Illinois Waters C. S. BORUFF AND G. B. ABBOTT,State Water Survey, Urbana, Ill. An adaptation of the volumetric thorium preconcentration, drying, and disOTTLED tooth enamel has been found to be the most tillation of silicon fluoride from is caused by the concipitation dry residues in the presence of ingestion of reliable for the quantitative examination of sulfuric acid and silica (i7) with fluorides during the period the waters for fluorides. The fluoride content of a subsequent determination of the teeth are being formed (8, 18, soluble fluoride by one of many 19). Mottled teeth have been number of Illinois waters has been determined by m e t h o d s found in the literathe use of this method. found in p r a c t i c a l l y e v e r y F~~~ the data obtained it that the tUre, have been found tedious country. This condition is enand u n s a t i s f a c t o r y (23,27). demic over large areas in the waters containing the greater quantities of The technic is exacting. This southwestern part of the United fluorides are those which are found in and around m e t h o d w a s used in the first States and is frequently found in studies r e p o r t e d b y S m i t h , other s e c t i o n s (8, 11, 12, 1.4). the northeastern and northwestern part of the Lantz, and Smith (19). Illinois coal basin. All of these waters r u n high Studies already published have The method of Fairchild (9) largely l i m i t e d the p r i m a r y in nonincrustingsolids. T~ date the results has been modified and used by Sources Of thediet to have not shown fluorides to be characteristic of Churchill (r), Hale ( l i ) ,and that found in drinking waters (11-13, 18, 19). Bauxite, Ark., any one stratum since they are found in waters Smith (21). This method calls for the addition of a ferric salt and Oakley, I d a h o , have disassociated with rocks of Pennsyhanian, Silurian, which f o r m s ferric fluoride or and Ordovician age. carded abundant water supplies the complex FeF,j--- with the in favor of others containing less fluorides ( l a ) , and other cities fluorides present. The excess are considering similar action. The exact minimum concen- ferric ion is then caused to react with potassium iodide, and tration of fluorides that will cause mottled teeth is not known. the liberated iodine is titrated with thiosulfate. The reIt has been reported by some to be about 2.0 p. p. m, (21) liability of the method depends on the stability of the ironwhile others claim it to be about 1.0 p. p. m. (16). fluoride complex and on a quantitative reaction between Lack of agreement in results obtained in this and other ferric iron and potassium iodide. The reaction laboratories while analyzing waters for small quantities of 2Fe+++ 21- = 2Fe++ 1 2 fluorides, using methods reported in the literature, together with the fact that extremely little is known concerning the is reversible, and, if it is to be used in a quantitative defluoride content of Illinois waters, led the Water Survey to termination, the iodide must be added in large excess (26). initiate the following investigation. This is strictly a chemi- Fairchild and Churchill have not used sufficient potassium cal study; no attempt has been made by the authors to iodide. Upon testing the reliability of their methods, the correlate the presence or absence of mottled teeth with the present writers found that the amount of iodine liberated by fluoride content of the Illinois waters which have been ex- the reaction of 5.0 cc. of approximately 0.08 M ferric chloride with the 10 cc. of 5 per cent potassium iodide in distilled amined. Of the numerous methods found in the literature for the water, according to their procedures, required an amount of quantitative determination of fluorides, there seem to be only 0.025 N thiosulfate which varied from 13.7 to 15.2 cc. Others a few which might be utilized to determine the small quan- (6) have noted similar variations in this reaction. Increasing tities (0 to 15 mg. per 1.) found in waters. Gravimetric and quantities of potassium iodide (iodine-free) required incolorimetric methods are fraught with many difficulties and creasing volumes of thiosulfate to react with the liberated uncertainties (6, 9, 11, 16, 22, 24, 26, 27). Likewise, the iodine. Similar tests made by adding definite volumes (1
M
+
+
July 15, 1933
INDUSTRIAL AND ENGINEERING CHEMISTRY
to 20 cc.) of standard sodium fluoride (0.84 gram per 1.) to 100 cc. of distilled water gave discordant results. I n these tests the difference between the amounts of thiosulfate used in the final titrations and those used in the blanks varied from -0.03 to +0.37 cc. per cc. of sodium fluoride (0.84 mg.) added. The average of twenty determinations was +0.25 cc. per cc. of sodium fluoride. Tests in which a constant amount of standard sodium fluoride was added and which were titrated in accordance with the procedures of Fairchild and Churchill varied as much as 1.0 cc. of 0.025 N thiosulfate. The presence of iron in waters to be analyzed for fluorides further complicates matters. Additional modification of these methods by Smith and Smith (21) has not improved their reliability. (A water from the Colorado River a t Yuma, Ariz., which was reported by Smith to contain 4.0 p. p. m. of fluorides, but was not causing mottled teeth, was found by the present writers to contain only 0.4 p. p. m.) It is unfortunate that the Churchill method has been included in the appmdix of the new Standard Methods of Water Analysis (3). The method recently announced by Foster (10) calls for addition of excess ferric ion in order to obtain the ferric fluoride complex. The excess iron is determined by the thiocyanate method. Foster's data show no definite, constant relation between the ferric ion removed and the fluoride present. If the ferric fluoride or the FeFG--- complexes were formed, 25 mg. of iron should remove 26 and 52 mg. of fluorides, respectively. The analytical data do not correspond with either of these complexes. One of the basic requirements of any quantitative method is that a constant and definite relation must exist between the reactants. Thompson and Taylor (24) are using a modified CasaresDe Boer method (colorimetric) for the determination of fluorides in sea waters. The procedure seems very empirical. Churchill (7) has proposed the use of a spectrograph for the determination of fluorides in waters. The method would undoubtedly be reliable but the cost of the apparatus would make its use prohibitive in most water laboratories. Willard and Winter (27) have noted that soluble fluorides can be quantitatively distilled from acid solutions as hydrofluosilicic acid (H2SiFs). In this form and in the absence of large quantities of soluble salts (many interfere with direct determination), the fluorides can be determined by various methods. Allen and Furman (1, 2 ) have investigated the use of cerous nitrate as a titrating agent and triphenyltin as a precipitant for soluble fluorides. Each method has certain distinct disadvantages. Wichmann and Dahle (26) have investigated a modified Steiger-Merwin method. This procedure depends on a bleaching, by the fluorides present, of the yellow color of peroxidized titanium salts. The present writers have not had very much success in determining the fluoride content of distillates by comparing the depth of yellow color developed in them with that developed in similar standard 50-cc. Nessler tubes containing various known quantities of sodium fluoride. The use of thorium nitrate as a volumetric precipitant as outlined by Willard and Winter (97) seems promising.
RECOMMENDED METHOD OF ANALYSIS The present authors have found that, if waters containing small quantities of fluorides are made alkaline to phenolphthalein or litmus, they can be concentrated without loss of fluorine. This can best be accomplished by concentrating the sample in the distilling flask that is later to be used for volatilizing the hydrofluosilicic acid. Suficient water should be concentrated so that the final volume of about 50 cc. will contain a t least 0.2 mg. of fluorine. The following procedure is fundamentally that developed by Willard and Winter
237
(27). It has been modified and adapted for use in the analysis of waters. REAGENTS.(1) Zirconium nitrate, one gram of Zr(NOa)c5Hz0 in 250 cc. of water. (2) Alizarin red, one gram of sodium alizarin sulfonate in 100 cc. of ethyl alcohol. Filter off the undissolved residue and add 150 cc. of ethyl alcohol to the filtrate. Solutions 1 and 2 are kept in stock and mixed (three parts of 1 and two parts of 2) as needed. The color should be violet-red when mixed. (3) Thorium nitrate solution, 0.01 to 0.02 N , standardized against a standard fluoride solution. (4) Standard fluoride solution, lithium fluoride, 0.02 N (0.5188 gram/l.) or specially purified sodium fluoride ( I ) , 0.02 N (0.8400 gram per 1.). (5) Hydrochloric acid, about 0.2 N . (6) Sodium hydroxide solution, about 0.2 N . STANDARDIZATION OF THORIUMNITRATESOLUTION.Add small definite volumes of the 0.02 N fluoride solution from a buret to a small tall-form beaker or Erlenmeyer flask, add water t o make the volume approximately 20 cc., and then add an equal volume of ethyl alcohol. Add six drops (about 0.15 cc.) of the zirconium-alizarin mixture, then only enough dilute hydrochloric acid to destroy the color. Avoid excess acid. The present writers have confirmed the findings of Armstrong (4)-namely, that it is better to use only the sodium alizarin sulfonate as the indicator, instead of the zirconium-dye mixture, in titrating solutions containing less than 0.5 mg. of fluorides. Using the thorium nitrate solution, titrate over a white surface (in good light) to a faint permanent reap earance of color. Titrate slowly upon nearing the end point. k u n a blank titration on the indicator by determining the volume of standard (0.02 N ) fluoride solution necessary to cause disappearance of color in a slightly acid water-alcohol solution of six drops of the indicator mixture and compare this with the volume of standard thorium nitrate necessary to discharge the color. Calculate the strength of the thorium nitrate solution by use of the following equation: o.l cc. T h ( N 0 3 ) r
N F- soln. cc. of Th(N0a)r soh.
cc. of 0.02
I
0'38
A mg*Of F-
Make the water t o be examined ANALYTICALPROCEDUR~. alkaline t o litmus or phenolphthalein and concentrate a definite volume (to contain at least 0.2 mg. of fluorine) to approximately 50 cc. in a 250- or 500-cc. distilling flask. Add a few glass beads or silica, 20 cc. of concentrated sulfuric acid (60 per cent perchloric acid may be used) and sufficient water t o cause the solution t o boil at 110" C. Place the flask on an asbestos mat with an opening sufficiently large for about one-third of the flask to be exposed to the flame. Close with a two-hole rubber stopper through which passes a thermometer and capillary tube, both of which extend into the liquid. Connect a dropping funnel with the capillary tube so that water may be added during the process of distillation. Connect the flask with a water condenser. The distillate may be collected in an open container. Distill until the boiling point reaches 130' to 140' C. and hold at approximately this tern erature by adding water through the capillary tube until all t f e fluoride has been volatilized. This is usually accomplished by collecting 50 t o 75 cc. of distillate, but the present writers prefer to collect about 100 to 150 cc. and then concentrate this t o approximately 20 cc. on a hot plate after it has been made alkaline. Neutralize most of the alkali before proceeding. Add the six drops of indicator, dilute acid until the color of the indicator just disappears, and then an equal volume of neutral alcohol. Be sure the solution is only faintly acid. If no fluorides are present, the color will not be discharged. Redistill if a precipitate forms. Titrate at once in a tall-form beaker with the standardized thorium nitrate solution to the faint reappearance of the pink color. Report as parts per million of fluorides. Titration of known quantities of fluorides with the thorium nitrate standard always gave a definite and constant relation between the two. I n order to check further the above recommended procedure, the authors prepared a stock synthetic water which contained the following c . P. reagents: P. p .
P. 8 . m. CaC03, as Ca(HC0a)n MgSOr NaCl
170 160 30
FeCls xHaC1
77%.
1.0 3.0
To one-liter samples of this water were added various but known quantities of fluorides (0.5 to 5.0 p. p. m. as sodium fluoride). These samples were then concentrated and ana-
238
ANALYTICAL EDITION
Vol. 5, No. 4
LITERATURE CITED lyzed in accordance with the above recommended procedure. The analyses checked within +0.10 to -0.10 p. p. m. of Allen and Furman, J. Am. Chem. Soc., 54,4825 (1932). the amount of fluorides added. Duplicate analyses were Ibid., 55,90 (1933). also made of most of the natural waters examined; these Am. Pub. Health Assoc. N. Y . , Standard Methods of Water Analysis, 7th ed., 1933. also checked within satisfactory limits. The accuracy of the Armstrong, J . Am. Chem. Soc., 55, 1741 (1933). method, therefore, seems to fall within the limits of the Batchelden and Meloche, Zbid., 53,2131 (1931). periodic variations noted in the chemical make-up of natural Buswell and Gallaher, IND.ENG. CHEM.,15, 1186 (1923). waters. Churchill, Zbid., 23, 996 (1931). Dean, J. Am. Dental Assoc., 20, 319 (1933). The residue from waters containing in the order of 1000 Fairchild, J. Wash. Acad. Sci., 20, 141 (1930). p. p. m. of chloride should be distilled into sufficient alkali Foster, J. Am. Chem. Soc., 54,4464 (1932); IND. ENG.CHEM., in order to keep the distillate from becoming too acidic. If Anal. Ed., 5,234 (1933). Hale, Okla. Agr. and Mech. Eng. Expt. Sta., Bull. 3, No. 3, abnormal quantities of aluminum salts, gelatinous silica, or 35 (1932). borates are present in a water, a larger volume (about 200 cc.) Kempt and McKay, U. S. Pub. Health Service, Pub. Health of distillate should be collected in order to guarantee complete Repts., 45, 2923 (1930). recovery of the fluorides present. McCollum, Simmonds, and Becker, J . Biol. Chem., 63, 553 (1925). A number of Illinois waters (untreated) have been examined McKay, J. Dental Research, 10, 5, 561 (1930). for fluorides. Four lake-reservoir samples contained from McKay, paper presented before the Research and Biological 0 to 0.5 p. p. m. Five river samples contained from 0.2 to Section, Chicago Dental Society, Jan. 19, 1932. 0.5 p. p. m. Five shallow-drift, gravel, and limestone wells, Merwin, Am. J. Sci., 28, 119 (1909). Reynolds, Ross, and Jacob, J . Assoc. Agr. Chem., 11, 225 (1928). distributed about the state, contained up to 0.7 p. p. m. Sebrell, Dean, Elvove, and Breaux, U. 8. Pub. Health Service, of fluorides. One 497-foot limestone well in north central Pub. Health Repts. 48,437 (1933). Illinois produces a highly mineralized water containing 1.1 Smith, Lantz, and Smith, Ariz. Expt. Sta., Tech. Bull. 32, 253 p. p. m. of fluorides. Five deep St. Peter wells in and around (1931). Smith, Lantz, and Smith, Science, 74, 244 (1931). this same district produce highly mineralized waters containing Smith and Smith, Ariz. Expt. Sta., Tech. Bull. 43,213 (1932). from 1.0 to 2.0 p. p. m. of fluorides. One 404-foot Niagaran Steiger, J. Am. Chem. Soc., 30,219 (1908). well in this same district produces water containing 3.7 Tananaev, J. Applied Chem. (U.S. S. R.), 5, 834 (1932). p. p. m. of fluorides. Drainage from a fluospar mine conThompson and Taylor, IND.ENQ. CHEM., Anal. Ed., 5, 87 (1933). tained 1.7 p. p. m. of fluorides, while a 100-foot well nearby Treadwell and Hall, “Analytical Chemistry,” 8th ed., Vol. 11, contained only 0.1 p. p. m. p. 577, Wiley, 1924. Dean (8) has reported that some of the Monmouth resiWichmann and Dahle, J. Assoc. Oficial Agr. Chem., to be published. dents are affected with slight mottling of the teeth and that Willard and Winter, IND. ENG.CHEM.,Anal. Ed., 5, 7 (1933). 53 per cent of the children in Minonk and 66 per cent of those Winter, J . Assoc. Oficial Agr. Chem., 15, 505 (1932). in Fairbury have mottled teeth. Others (8) have reported mottled enamel in Milan and in and around Emington and RECEIVED May 20, 1933. Presented before the Division of Water, Sewage, Kempton. All of these waters have been found to contain and Sanitation Chemistry at the 86th Meeting of the American Chemical 1.0 p. p. m. or more of fluorides. Society, Washington, D. C., March 26 to 31, 1933.
Sources of Error in the Use in Water Analysis of Fairchild’s Method for Determination of Fluoride in Phosphate Rock MARGARET D. FOSTER, United States Geological Survey, Washington, D. C.
T
H E method developed by Fairchild (2) for the determination of fluoride in phosphate rock has been used for the determination of fluoride in water (1, 3). Studies of this method have shown that (1) with no interfering elements present the method usually gives high results and is accurate only to 0.2 to 0.4 mg. of fluoride, less than many waters contain in a liter, and (2) sulfate apparently withdraws some of the iron from reaction with potassium iodide and gives the same effect as fluoride. The amount withdrawn depends to some extent on the amount of sulfate present and to some extent on whether the sulfate is present as the sodium, calcium, or magnesium salt. The effect of sulfate may be practically eliminated by making the sample just acid to methyl orange, adding the 2.2 cc. of N hydrochloric acid which is usually added later, precipitating the sulfate with barium chloride (which has no effect on the method) and proceeding as before, but omitting further addition of acid. Chloride if present as the sodium salt does not appreciably affect the result. One gram of calcium or magnesium
chloride tends to give a slightly high result, giving an apparent fluoride of 0.3 mg. when none is present. Nitrates and borates, in amounts usually found in water, do not interfere. If small volumes of water are used for the determination, the possible error as indicated by the range in readings of the controls might be as great as the fluoride present or even greater. Large volumes, if used, would necessitate, in many waters, the removal of sulfate and concentration, with the probability of precipitation of calcium fluoride. The method does not seem well adapted to the determination of fluoride in water. LITERATURE CITED (1) Churchill, H. V., J. Am. Water Works Assoc., 23, 1399-1407 (1931). (2) Fairchild, J. G., J. Wash. Acad. Sci., 20, 141-6 (1930). (3) Smith, H. V., and Smith, M. C., Arizona Univ. College Agr., Tech. Bull. 43 (1932). RECEIVED March 7, 1933. Published by permission of the Director, U. S. Geological Survey. f
