Removal of Fluorides from Drinking Water RALPH H. MCKEE AND WILLIAM S. JOHNSTON, Columbia University, New York, N. Y.

HE first reference to mot- Attempts were first made to re- tled enamel of teeth was to be the move the fluorine with iron a d made in 1901 in a report aluminum sulfates, but neither America, England, Africa, and elsewhere. The mas of value, The amount of by Eager (7). Since then many

on the subject (1, 6 , 11, 12, mocingflriorides frompotable water by adsorption p. p. m. and the time of contact 16-19). However, it mas not by carbon. The Process, however, has the handi- ranged from 30 minutes to 48 until 1931, when Churchill (5) cap that the tirne of treaiment the water must hours. In the case of the alumi- indicated that fluorine might be be at a p H of num salt 83 p. p. m. mere added the cause of the trouble, that

Fluorides in drinking water have been shown of mottled enamel on teeth in T

a r t i c l e s h a v e been w r i t t e n present authors have developed a method of re- iron salt varied from 17 to 165

or less. this element was d e t e r m i n e d in water from endemic areas. The work of blcKay ( I S ) and the Smiths in Arizona (6, 16, 19) has definitely confirmed the fact that fluorine is the cause of the difficulty.

An article by Boruff ( 2 ) is the only one that has dealt with methods of removing fluorine from water. One criticism that may be made of his work is that no concentration in excess of 5.0 parts per million of fluorine was used, and even a t this concentration the removals were not satisfactory. Many waters from western endemic areas vary from 5 to 8 p. p. m., and cases are known in which it is as high as 16.8 p. p. m., so that a limit of 5.0 p. p. m. is felt t o be low.

METHODS OF ANALYSIS There are several suggested methods of analysis for small

amounts of fluorine (4, 10, 16), but each had some objection- able features which made it unsuitable for use in the present experiments. The method outlined by Fairchild (8) and modified by Churchill (6) was followed in the first part of this investigation and has proved fairly satisfactory.

Since this vork was started, other methods of analysis for small amounts of fluorine have appeared (3, 9, 20, 21).

In order to obtain satisfactory results with a fluorine range of 1 to 10 p. p. m., it was necessary to modify Churchill’s method slightly. I t was found that the ferric chloride solu- tion could be reduced to one-third its original strength- i. e., to 0.00533 M-and still give satisfactory results, but that no dilution could be made in the thiosulfate solution.

Churchill stated that the flasks should be heated for 30 minutes to complete the liberation of the iodine, but it was found that this was not important, as variations of from 30 to GO minutes in the time of heating a t 38” C. did not ap- preciably alter the results.

Since the waters from the American endemic areas are alka- line, the “synthetic water” used in this work was New York city water to which was added the desired amount of sodium fluoride and which was made alkaline with 0.272 gram of sodium acid carbonate per liter, which is an average figure found in analyses of such waters. This amount of bicarbonate produced a pH of 8.3 but had no effect on the titration, as the increased amount of sodium thiosulfate consumed was within the limit of error of the method, being less than 0.10 cc.

Portions (250 cc.) of the synthetic potable water containing fluorine were measured out, and the required amount of acid and adsorbent was added. This gave sufficient volume so that, on filtering off the adsorbent and rejecting the first portion of the filtrate, there would still be over 200 cc. left-enough to pipet the two 100-cc. portions required for the duplicate analyses. Blanks were similarly put through but without any adsorbent, and the difference between the two titrations gave the amount of fluorine adsorbed.

All determinations were run in duplicate.

and left in contact with thesolu- tion for 48 hours. In these ex-

periments the reduction of fluorine was only 0.25 p. p. in. which was within the limits of error of the method. Boruff (2) showed that with aluminum sulfate there was a reduction of slightly more than 1.0 p. p. m. of fluorine a t a pH of 8.0. The present lonTer removal may be explained by a different type of stirring. Boruff states that this is important. The authors used only the slightly alkaline water (pH, 8.3) for these salts, whereas Boruff varied his pH considerably. It was thought that, since ferric fluoride is nonionized, it might be more easily adsorbed. Therefore, a krric salt was added to an acid solu- tion of the water containing sodium fluoride. After a few minutes sufficient sodium carbonate was added to make the water alkaline again, the precipitate of ferric hydroxide was filtered off, and the filtrate was tested for fluorine, but there was no reduction.

The siliceous adsorbents used were two grades of Wyoming bentonite and one each of fuller’s earth, H. 8. C. Celite, and silica gel. To samples of the synthetic water 0.2 per cent by weight of adsorbent on the weight of the water was added. The material was left in the water for 16 hours and shaken from time to time. The adsorbent was then filtered off and the filtrate tested for fluorine. In no case was any of the fluorine removed. In these experiments the adsorbent was added to the water without changing its pH.

The next step was to add various amounts of acid and ad- sorbent. The pH o€ the original water was 8.3 which was reduced in small steps to 2.5. There was no removal of fluor- ine until the acidity had reached a pH of 3.6, and a t 2.5 ap- proximately one-half of the amount originally present was removed. While this \yas the first promising lead, it was far from the desired removal.

REMOVAL BY CARBON The following four different carbons were next tested:

(A) IYorit, a commercial carbon produced from the charcoal of the European pine; (B) the residual carbon discarded by the soda pulp industry; (C) the same as carbon B activated with acid (14); (D) a war gas type of adsorbent carbon which is an anthracite coal activated with hot carbon dioxide gas.

In fact, with carbon A the blank in all cases gave a higher thiosulfate consumption than the sample with the carbon. This indi- cated that something was leached out of this carbon which reacted with the thiosulfate.

Carbon B removed about half the fluorine a t a pH of less than 3.0.

Carbon C showed real promise, and the greater part of the work was devoted to its study. Here, also, there was little removal of fluorine until the pH had been reduced to 3.6, and between this and 2.5 the removal of fluorine increased from approximately 30 per cent to complete adsorption,

No fluorine was removed by the A or D carbon.

849

850 I N D U S T R I A L A N D E N G I N E E R I N G C H E M I S T R Y Vol. 26, No. 8

TABLE I. RUNS IN THE CONTINUOUS APPARATUS RUN 1, C CARBON RUN 2, C CARBON RUN 3, C CARBON RUN 4, B CARBON

p H P . p . m . pH P . p . m . pH P . p , m . pH P . p . m . p H P . p . m . p H P , p . m . p H P . p , m . RCN 5, E CARBON RUN 6, B CARBON RUN 7,a B CARBON

Wate r8 .3 8 . 0 5 . 7 8 . 0 3 . 1 8.0 3 . 0 8.0 5 . 2 8 . 0 4 . 4 8 . 0 5 . 4 8 . 0 FILTRATE

Liters 0-2 3-4 5-6 7-8 9-10

11-12 13-14 15-16 17-18

6 . 6 6 . 8 7 . 0 7 . 2 7 . 3 7 . 5 7 . 4 7 . 4 7 . 5

6 . 0

8 . 0

8 . 0

...

...

816 . . . 8 . 0

6 . 7 6 . 2 6 . 1 6 . 5 5 . 9 6 . 7 5 . 8 7 . 1 7 . 3

6 . 0 8.0 8.0

8 . 0

8 . 0

8 . 0

. . .

. . .

. . .

6 . 4 3 . 3 3 . 1 3 . 1 3 . 1 3 . 1 3 . 1 3 . 1 3 . 1

Ferric chloride (185 p. p. m.) was sdded to the water.

2 . 0

0 . 5

1 . 0

3 . 0

4 . 0

. . .

. . .

. . .

. . .

using the customary 0.2 per cent carbon based on water taken.

The efficiency of adsorption in relation to the pH of the solution is as follows, where 8.0 p, p. m. of fluorine was added:

PER CENT PH REMOVED

6 . 9 0 4 3 . 5 7 49 3 . 1 6 88 3 . 0 1 100

PER CENT PH REMOVED

2 . 5 3 99 2 . 3 8 99 2 . 2 5 100

Thus complete.remova1 is obtained a t a p H of 3.0 or less, as the 99 per cent values may be taken as complete removal of the fluorine.

The efficiency of fluorine removal by the carbon, when only 0.08 per cent by weight of carbon 04 the weight of water was used, is as follows :

FLUORINE FLUORINE FLCORIN. FLUORINE ADDED REMOVED ADDED REMOVED P . p . m. P . p . m. P . p . m. P . p . m.

10.0 8 . 4 5 1 0 . 0 7 . 5 0 7 . 5 7 . 0 5

5 . 0 4 . 6 0 2 . 5 2 . 7 5

This shows that 0.08 per cent carbon will remove practically all the fluorine from waters containing up to 7.5 p. p. m., and 80.0 per cent from water with 10.0 p. p. m. One part per mil- lion is the maximum that may be allowed in potable water if mottling of children’s teeth is to be avoided. Since the majority of waters analyzed from the endemic areas showed 8.0 p. p. m. or less, then 0.08 per cent of carbon C would seem to be sufficient in most cases. This was confirmed by using a water containing 8.0 p. p. m. at a pH of 2.25, where 0.08 per cent carbon removed all the fluorine, but 0.04 per cent removed 76 per cent, and 0.02 per cent only 57 per cent.

The next step was to develop some method of continuous treatment which might have a practical application:

For this purpose a glass tube 40 mm. in diameter and 750 mm. long was filled to a height of 620 mm. with carbon. For these tests a sieved material passing a 16-mesh but held on a 30-mesh screen was used. The water was run in at such a rate that a constant, slight hydrostatic head was maintained during the run. Owing to packing of the carbon, there resulted a slight decrease in the rate of flow during the run. In run 7 there was a marked drop in the rate of flow due to the iron salt added to the water.

In order to follow the rate of adsorption of the fluorine, the effluent water was collected in 2-liter samples which were ana- lyzed. This entailed a large number of analyses, and the Fair- child method would have been too slow. The method outlined by Thompson and Taylor (21) seemed satisfactory on test and was adopted with slight changes. The solutions were made up as follows: solution 1, 0.87 gram zirconium nitrate and 100.00 cc. water; solution 2, 0.14 gram alizarin red, 7.50 cc. normal sodium hydroxide solution, and 100.00 cc. water. The two solutions were mixed, allowed to stand for 3 hours, and then filtered. For the analysis 1.0 cc. of the reagent and 1.0 cc. of 5 N hydrochloric acid were added to 50.0 cc. of the water. Stand- ard fluorine solutions were made up by increments of 2.0 p. p. m. from none to lop. p. m., and 1.0 cc. each of the acid and indicator was added to 50.0 cc. of the standard. There was a range in color from pink with no fluorine to a straw yellow at 6.0 p. p. m. Above this there was no appreciable difference in color. If the unknowns were above 6 . 0 ~ . p. m., it was necessary to dilute them in order that they might be compared to the standards. Tests

6 . 6 5 . 7 4 . 8 4 . 5 4 . 2 3 . 8 3 . 5 3 . 3 3 . 3

4 . 0 0 .0 0 . 0 0 . 0

0 . 5 0 . 5

. . . . . .

...

6 . 4 6 . 1 6 . 0 5 . 9 5 . 9 5 . 9 5 . 9 5 . 9 5 . 9

8.0

8 . 0

6 . 0

5 . 0

6 . 0

...

...

. . .

...

6 . 7 6 . 3 6 . 2 6 . 1 6 . 1 6 . 1 6 . 1 6 . 0 5 . 9

8 . 0 5 . 0 5 . 0

5 . 0

5 . 0

5 . 0

. . .

. * .

. . .

6 . 7 6 . 5 6 . 3 6 . 3 6 . 1 6 . 1 6 . 3 6 . 5 6 . 4

8 . 0

6 . 0

5 . 0

7 . 0

6 . 0

. . .

. . .

. . .

. . .

showed that acid in excess of the 1.0 cc. added did not change the color of the standard solutions.

TABLE 11. RATES OF FLOW IN THE CONTINUOUS RUNS

FILTRATE 1 2 3 4 5 6 7 Litera I Minutes

3-4 7 . 2 5 7 . 7 5 3 . 5 3 . 7 5 4 3

9-10 s : i 7 : + 5 5 5 : 2 5 6 . 5 5-6 515 . . I . . . .

11-12 6 : 2 5 . . 4:+5 I . . . . 15-16 6 . 5 8:5 9 6 : 2 5 6 9 17-18 . . . . . . 5 : + 5 . . . . ...

In comparing Tables I and I1 with the earlier tables in which the small-scale batch tests were carried out, i t will be seen a t once that the results are of the same order in both cases-namely, that there is little removal of fluorine until the p H has been reduced to 3.0. In run 3 the poor removal of fluorine after the twelfth liter was run through was due to the fact that the limit of adsorption had been reached. The same 100 grams of carbon had been used for the first three runs of 20 liters each. In the fourth run fresh carbon (B) was taken, and this was effective for the entire 18 liters. As expected, the first few liters of the acidified water activated the carbon and made it effective.

In the five runs made on the acid side, in all cases the first 2 liters were always more alkaline than the rest, with resulting low fluorine removal. This is explained by the small amount of residual alkali present in the carbon resulting from its method of production.

The chief objection to this method of purification is that such a low pH is required before the flourine is removed. It was felt that the nonionized ferric fluoride might be more easily adsorbed. Several attempts were made in the small- scale tests to obviate this difficulty by adding ferric chloride to the water. In run 7 three times the amount of ferric chloride required to form ferric fluoride was added, and the water was allowed to stand for 5 hours before passing through the carbon. Once more, with the alkaline solution, almost no fluorine was removed.

SUMMARY

The results were quite inconclusive.

This work has indicated that it is possible to remove the fluorine from water by adsorption. Several different ma- terials have been tested; of these the carbons were the most promising and some were more efficient than others. The chief criticism of this work is that, even with the most efficient carbon, there is no removal of fluorine until the pH of the water has been reduced to about 3.0. Attempts to reduce to 1 p. p. m. or less the fluorine content of the water with a higher pH have been unsuccessful.

ACKNOWLEDGMENT

Frederick S. McKay for his suggestions.

LITERATURE CITED

The authors wish to express their grateful appreciation to

(1) Black and McKay, Dental Cosmos, 58, 1 3 2 4 6 (1916). (2) Boruff, IND. ENG. CHEM., 26, 69 (1934). (3) Boruff and Abhott, Ib id . , Snal. Ed., 5, 236 (1933).

August, 1934 I N D U'ST R EA L A N D E N G I N E E R I N G C H E M I S T R Y 85 1

(4) Carles, Compt. rend., 144, 37, 201 (1907). (5) Churchill, IXD. EXG. CHEM., 23, 996 (1931). (6) Dean, J. Am. Dental Assoc., 20, 319 (1933). (7) Eager, U . S. Public Health Rept., 16, 2576 (1901). (8) $airchild, J. Wash. Acad. Sci., 20, 141 (1930). (9) Foster, IND. ESG. CHEM., Anal. Ed., 5, 234 (1933).

(15) Papish, IND. ENG. CHEM., Anal. Ed., 2, 263 (1930). (16) Smith, M. C., and Lantz, E. L., Ariz. Agr. Expt. Sta., Tech.

(17) Smith, M. C., Lantz, E. L., and Smith, H. V., Ibid., 32 (1931). (18) Smith, M. C., Lantz, E. L., and Smith, H. V., Science, 74, 244

(10) Gautier and Clausmann, Compt. rend., 154, 1469, 1670, 1763 (19) Smith, M. C., and Smith, H. V., Ariz. Agr. Expt. Sta., Tech.

(11) Kempf and McKay, U . S. Public Health Rept . , 45, 2923 (1930). (20) Smith, 0. M., and Dutcher, H. A., IND. ENG. CHEM., Anal. Ed., (12) McKav. Dental Cosmos. 68. 847 11925). 6. 61 11934).

Bull. 45 (1932).

(1931).

(1912); 158, 1389, 1631 (1914). Bull. 45 (1932).

i13j McKay; J. Am. Dental Assoc., 20, 1137 (1933). (14) McKee, U. S. Patent 1,133,049 (1915).

(21) Thompson and Taylor, Ibid., 5, 87 (1933) RECEIVED April 7, 1934.

Formation of Gas Hydrates in Natural Gas Transmission Lines

E. G. HAMMERSCHMIDT, Texoma Natural Gas Company, Fritch, Texas

T HE p r e s e n c e of w a t e r vapor in natural gas has always been a source of

trouble to the n a t u r a l g a s in- dustry in the m e a s u r e m e n t and transportation of the gas. One of the chief difficulties has been the interruptions of ser- vice due to the liquefaction and s u b s e q u e n t f r e e z i n g of t h e water within the system. The solid matter which collects in the pipe line usually resembles ordinary snow in appearance. The movement of thegas through the pipe line tends to collect and compress the snow a t low spots until the line may become en- tirely plugged. The snow is h o n e y c o m b e d w i t h s m a l l channels through which the gas passes before the flow is entirely stopped. T h e c a u s e of t h i s freezing has usually been attrib- uted to subzero ground tem- peratures or to a combination of low temueratures with Dressure

Solid compounds, resembling snow or ice in appearance, are formed with methane, ethane, propane, and isobutane in the presence of water at elecated pressures and temperatures. The melt- ing point of these mixed hydrates in a natural gas mixture depends upon the pressure and Daries from about 34" F. at 110 pounds per square inch absolute to about 60' at 800 pounds.

The formation of gas hydrates in natural gas pipe lines depends primarily upon the pressure, temperature, and composition of the gas-water vapor mixture. After these primary conditions are fuljilled, the formation of the hydrates is ac- celerated by high velocities of the gas stream, pres- sure pulsations, or inoculation with a small crys- tal of the hydrate.

At equilibrium conditions the hydrates, be- cause of their lower vapor pressure, cause more water to be removed from the vapor phase than in the case of liquid water at the same temperature and pressure.

temperatures above the normal f r e e z i n g p o i n t of water is a phenomenon that has not been generally recognized by the gas industry.

S c h r o e d e r (7) has reviewed the history of t h e d i s c o v e r y of g a s hydrates: H u m p h r e y Davy, in 1810, discovered the first known gas hydrate, a crys- talline compound f o r m e d b y chlorine and water. Wroblew- ski, in 1882, reported a carbon dioxide hydrate. Cailletet, in 1878, reported a c e t y l e n e hy- drate and was the first to dis- cover that a sudden decrease in pressure aided in the forma- tion of these crystalline com- p o u n d s , Woehler , i n 1840, reported hydrogen sulfide hy- drate. Villard and de Forcrand have worked for m o r e t h a n 40 years on this class of com- pounds. Villard (6, 9) reported hydrates of methane, ethane,

fluctuations; the latter causes intermittent liquefaction and vaporization of the more volatile hydrocarbons, such as pro- pane and the butanes.

However, it was discovered, during a series of experiments where natural gas and water vapor were compressed to 800 pounds per square inch, that freezing occurred a t higher temperatures than would ordinarily be expected. Later the same observation was made on a commercial scale where the natural gas was compressed to about 600 pounds per square inch and cooled to 40" F. in a refrigeration unit (1, 2) which was designed to remove the excess moisture and oil from the compressed gas.

These observations suggested the possibility of another and probably more general cause of freezing in natural gas systems than had heretofore been recognized. Therefore, the object of this investigation was to determine the causes of freezing a t elevated temperatures and pressures together with such other factors as might affect the operation of a natural gas transportation system.

The combination of certain gases with water to form crys- talline compounds (hydrates) a t elevated pressures and a t

a c e t y l e n e , a n d e t h y l e n e . Schutzenberger reported the first double hydratehydrogen sulfide and carbon disulfide. Double hydrates are definite compounds having a definite melting point and are by no means a mixture of the single hydrates, since the decompo- sition temperature of the double hydrate may be entirely different from the decomposition temperature of either single hydrate. De Forcrand characterized that product, which was obtained from hydrogen sulfide and aqueous alcohol (Woehler, 1840) as a mixed hydrate of hydrogen sulfide and alcohol, and in addition discovered the great family of "sulf- hydrierten" hydrates whereby hydrogen sulfide could be united with a great number of halogen-substitution deriva- tives of the aliphatic series, in hydrate form. Cailletet and Bordet, in 1882, discovered the double hydrate of carbon dioxide and phosphine. De Forcrand and Sully Thomas, in 1897, found that acetylene and carbon tetrachloride form a double hydrate. They also reported double hydrates of acetylene, ethylene, sulfur dioxide, and carbon dioxide with the following: ethylene chloride, ethylene bromide, methyl iodide, methyl bromide, methylene chloride, and methylene iodide. Hempel and Seidel also reported similar compounds