V O L U M E 2 3 , NO, 4, A P R I L 1 9 5 1 procedures must, so far as the authors can see, follow the approach used here-namely, to exclude oxygen, and to attempt t o suppress the induced decomposition of the peroxide. Possible means of achieving this suppression would be to use solvents other than acetone, or to add compounds of strong suppressing effect,such as chloride or bromide ions, or maleic or fumaric acids (19).
For general use, where accurate results are desired, iodometric procedures are preferable to ferrous methods at present. However, where some special objection exists to the use of an iodometric procedure, a ferrous method may be of value. Two of the procedures developed above give fairly accurate results with cumene hydroperoxide and soap peroxides, so that investigation of these procedures with various other peroxides would seem desirable. For control work, where the result of the peroxide analysis is correlated with some other property such as rancidity or explosive tendency, and the true peroxide content is of secondary importance, peroxide determinations carried out by the ferrous method in the presence of air are of particular value, owing to the increased sensitivity resulting from the induced air oxidation of ferrous iron. Such methods include the titration procedure in air (above), and various colorimetric procedures dewribed elsewhere ( 1 , 3,20, 26). SUMMARY
The determination of organic hydroperoxides by reduction with an excess of ferrous iron is inherently much less accurate than iodometric methods. For control work ferrous iron methods may be very useful for the determination of traces of organic peroxides. When the reaction is carried out in the presence of an excess of oxygen, the results may be several times greater than the theoretical value, but reproducible results can be obtained. I n the present work the solvent used was acetone, which is a suppressor of the induced decomposition of the peroxide by the ferrous iron-peroxide reaction. Procedures are given in which the excess of ferrous iron is determined by amperometric titration viith dichromate or the ferric iron formed is determined colorimetrically after addition of thiocyanate. I n the absence of oxygen, accurate results are obtained with certain peroxides but not
603 with others. The sources of error of the various procedures are discussed in detail. LITERATURE CITED
(1) Bolland, J. L., Sundralingam, A,, Sutton, D. A , , and Tristram, G. R., Trans. Inst. Rubber Ind., 17, 29 (1941). (2) Dastur, iY.N., and Lea, C. H., Analyst, 66,90 (1941). (3) Golden, M. J., J . Am. Pharm. Assoc., 35, 76 (1946); 37, 234 (1948). (4) Hock, H., and Lang, S., Ber., 77B,257 (1944).
(5) Kharasch, M. R., private communication. (6) Kokatnur, V. R., and Jelling, hl., J . Ana. Chem. SOC.,63, 1432
(194 1). (7) Kolthoff, I. XI., and Harris, K.E., IND.ENG.CHEM.,ANAL. ED., 18, 161 (1946). (8) Kolthoff, I. RI., and Laitinen, H. A., private communication. (9) Kolthoff, I. IT.,and Medalia, A . I., J . Am. Chem. SOC., 71, 3777, 3784, 3789 (1949). (10) Kolthoff, I. M.,and Stenger, V. A., “Volumetric Analysis,” Vol. I, Kew York, Interscience Publishers, 1942. (11) Laitinen, H. d.,and Nelson, J. s.,IND.ENG.CHEM.,ANAL. ED., 18, 422 (1946). (12) Lea, C. H., J . SOC.Chem. Ind., 64, 106 (1945). (13) Ibid., 65,286 (1946). (14) Lea, C. H., Proc. Roy. SOC.,108B, 175 (1931). (15) Lea, C. H., “Rancidity in Edible Fats,” London, 1938. (16) Lips, A,, Chapman, R. A., and McFarlane, W. D., Oil & Soap, 20,240 (1943). (17) May, D. R., Ph.D. thesis, University of Minnesota, 1944. (18) Medalia, A4.I., and Kolthoff, I. M.. J . Poliimer Sci.. 4, 377 (1949). (19) Merz. J. H.. and Waters. W.A.. J . Chem. Soc.. 1949.S15. 120) Robey, R. F., and Wiese, H. K., ISD.ENG.CHEM.,ANAL. ED., 17, 425 (1945). (21) Tanner, E. M.,and Brown, T. F., J . Inst. Petroleum, 32, 341 (1946). (22) TTagner, C. D., Clever, H. L., and Peters, E. D., ANAL. CHEM.,19,980 (1947). (23) Wagner, C. D., Smith, R. H., and Peters, E. D., Ibid., 19,976 (1947). (24) Ibid., p. 982. (25) Woods, J. T., and Mellon, XI. G., ISD.ENQ.CHEM.,ANAL.ED., 13,551 (1941). (26) Young, C. A , , Vogt, R. R., and Nieuwland, J. A., Ibid., 8 , 198 (1930). (27) Yule. J. A. C.. and Tt-ilson. C. P.. Jr.. Ind. Ena. Chem.. 23. 1254 (1931). RECEIVED August 7, 1950. Work carried out under the sponsorship of the Office of Rubber Reserve, Reconstruction Finance Corp., in connection with the government synthetic rubber program.
Determination of Fluorine in Monofluorinated Organic Compounds in Air and Water JASON R I . SALSBURY’, JA-\IES W. COLE, JR., LYLE G. OVERHOLSERI, ALFRED R . ARMSTRONG$, AND JOHN H. YOE University of Virginia, Charlottesville, Vu.
I
T IS generally recognized that i t is more difficult to convert
covalently combined fluorine t o fluoride ion than the corresponding cases with the other halogens. A number of investigators have shon-n that the rupture of the chlorine t o carbon bond may be accomplished fairly readily by refluxing the compound withan excess of an alkalielement in ethyl alcohol (2,10,12,16,19, SS, 36). This approach has been successfully used with a few aryl fluorine compounds and fluorophosphates (21, 38). 3Iost of the reports, however, indicate that more drastic conditions usually requiring involved procedures are necessary for quantitative Present address, dmerican Cyanamid Co., Stamford, Conn. Present address, Carbide a n d Carbon Chemicals Corp., Oak Ridge, Tenn. a Present address, Department of Chemistry, College of William and M a r y , Williamsburg, Va. 1
2
conversion of fluorine in alkyl fluorides t o fluoride ion ( 4 , 1 7 , 2 2 ,Z,?, 24, 27, 28, 35, 42). Kilpatrick ( 1 6 ) has recently discussed the kinetics and mechanism of hydrolysis of some fluoro-organic compounds. When small quantities of tosic fluoro-organic compounds are dispersed in air or dissolved in water, i t is particularly desirable to have rapid and accurate methods for their estimation ( 5 , 6 ) . Several somewhat elaborate methods have been reported. I n gas studies, Cadenbach ( 7 ) , Drake (9),and \T7inter ( 4 0 ) converted combined fluorine t o fluoride ion by burning mixtures of the fluoro-organic conipound with hydrogen a t a jet; the hydrogen fluoride formed was absorbed in aqueous sodium carbonate. Henne (14) detccted organic fluorides in air by the formation of a white cloud when the niisture was passed over silica heated t o incandescence and the effluent gas \vas mised with ammonia.
604
ANALYTICAL CHEMISTRY
Hartley (13) treated air mixtures with chromosulfuric acid and ,estimated hydrogen fluoride from the extent of etching of the glass container; his group also used a pyrolysis lamp to convert the fluorine to hydrogen fluoride. Schumb and Radimer ( 3 0 ) measured hydrogen fluoride formed in a platinum tube a t 1100" C. Drake (9) estimated the concentration of fluoro-organic compound in air mixtures by noting the color change in a standardized test paper impregnated with thorium nitrate and Solachrome Brilliant Blue indicator when the t.est paper was exposed to the effluent gas mixture after it had contacted a hot platinum wire. R'Iilton, Liddell, and Chivers (25, 26) used aqueous ammonia in a gas scrubber to absorb methyl fluoroacetate and dialkyl fluorophosphate from air. Fluorine in the former compound was converted to fluoride ion by heating in a sealed tube, whereas with the latter compound fluoride ion was formed upon evaporation to dryness. LIuch attention has been given to the determination of fluoride ion in water (1, 3, 8, 11, 20, 29, S1, SQ), but only little work has been reported on converting the fluorine to fluoride ion when very small quantities of the organic compound are dissolved in water. Buswell (5, 6 ) estimated fluorine by evaporation of an alkaline mixture of the compound to dryness and heating the residue and also by mixing the aqueous solution with ethyl alcohol and burning it in an alcohol flame, Kimber (18)extended the techniques of RIilton (25) and reported complete decomposition of fluoroalcohols and fluoroacetates upon heating the ammoniacal mixture a t 170" to 175" C. for 2 hours. In the authors' work it was first found necessary to modify existing procedures for the volumetric (25, 26, 5'7, 59)and colorimetric (20, 34) determination of fluoride ion with thorium nitrate in order to adapt them to the problem of determining low concentrations of fluorine compounds in air and in water. Details of both a volumetric and of a colorimetric method are given because each is useful and the two serve as a check on each other. During the course of the work considerable attention was given to devising a new means of conveniently converting the fluorine in the compounds alone and in mixtures, to fluoride ion. Refluxing the representative fluorine compounds with sodium in anhydrous higher alcohols (6 to 8 carbon atoms) quantitatively converted the fluorine to fluoride ion which was subsequently extracted as sodium fluoride with water. These alcohols also served to absorb sniall quantities of the fluorine compounds from air mixtures. Fluorine in the compounds when originally dissolved i n water \-as readily converted to fluoride by refluxing the aqueous solution with a mixture of potassium metaperiodate, silver perchlorate, percloric acid, and glass wool. The fluoride was removed from the complex mixture by distilling the silicon tetrafluoride by a modification of the method of Elsworth and Barritt (11). S o interference v a s noted in the distillate Tyhen it was analyzed for fluoride ion with thorium nitrate. The compounds used \yere sodium monofluoroacetate, methyl nionofluoroacetate, 2-fluoroethanol, and diisopropyl monofluorophosphate, each of which was obtained through Officc, of Scimtific Research and 1)evelopment channels. FLUORINE DECOMPOSITION AND CONVERSION
Transfer an accurately weighed 10- to 20-mg. sample of the fluorine compound to a 50-ml. round-bottomed, ground-joint flask (hood). Add 0.2 gram of sodium metal and 15 ml. of :I higher alcohol, preferably hexyl alcohol. Reflux the mixture gently for 15 minutes under a water condenser of sufficient lengt,h to avoid loss through t'he top. Pour the hot mixture into a separatory funnel and extract the hot solution as soon as possible with two 10-ml. portions of water, each of which was used to rinse the flask. Dilute the aqueous extract ( l o w r layer) with water to 50 nil. in a volumetric flask (solut'ion A). VOLUMETRIC DETERMI3 4TIOh OF FLUORIDE 10%
Standard Fluoride Solution. Dissolve 0.221 gram of sodiuni fluoride, or 0.495 gram of potassium fluoride in'water and dilute to 1 liter. The solution contains 0.10 mg. of fluorine per nil. If
desired, check the concentration of the solution by the lead chloride fluoride method (32). Thorium Nitrate Solution. Dissolve 0.69 gram of thorium nitrate tetrahydrate, reagent grade, in water and dilute to 1 liter. The solution is approximately 0.005 S. Sodium Alizarin Sulfonate Indicator. Dissolve 0.5 gram of sodium alizarin monosulfonate, reagent grade, in 80 ml. of water; filter and dilute the filtrate to 100 ml. Solochrome Brilliant Blue BS Indicator. Prepare a 0.027, aqueous solution. Chrome Azurol S from the Geigy Co., 89 Barclay St., Few York, S . Y., is the same material as Solochromc Brilliant Blue BS except for a higher tinctorial value; a 0.015% solution of the Geigy product was found to be equivalent to a 0.0270 solution of Solochrome. A fresh solution should be prepared approximately every 3 iyeeks because a eignificant amount of a dark precipitate forms over longer periods. Buffer Solution. Dissolve 22.7 grams of C.P. chloroacetic acid and 4.8 grams of C.P. sodium hydroxide in enough water to make 1 liter of solution. Hexyl Alcohol. Eastman, practical grade, boiling point 153" to 156" C. Standardization of Thorium Nitrate Solution. The volume of thorium nitrate required in the titration of various quantities of standard fluoride solution is not quite a linear function of the fluoride ion; hence the fluoride ion equivalent of the thorium nitrate solution must be determined over the useful range and a calibration curve plotted. The acidity also must be carefully controlled. Add 3 drops of sodium alizarin sulfonate indicator to a solution containing n ml. of standard fluoride solution (where n is any integer from 0 to lo), (10 - n ) ml. of water, 80 ml. of 60% ethyl alcohol, and 1 drop of 1.2 N hydrochloric acid. -4dd 4 ml. of buffer solution and with a microburet titrate with 0.005 S thorium nitrate solution to the first permanent pink tint. The volume of thorium nitrate used when n = 0 is the indicator blank. This is about 0.6 ml. of 0.005 S thorium nitrate and should be determined each day. Subtract this volume from the volume of thorium nitrate used in the other titrations to obtain the true volume of thorium nitrate which is equivalent to the volume of standard fluoride solution used. Plot the volumes of thorium nitrate against milligrams of fluorine. Titration of Sample. Add a 10-ml. aliquot of solution A to a solution containing 30 ml. of 60y0ethyl alcohol, 15 drops of 1.2 S hydrochloric acid, and 3 drops of sodium alizarin sulfonate indi-
Table I. \-olumetric D e t e r m i n a t i o n of Fluoride Ion from Fluoro-organic C o m p o u n d s Reflux Wt. of AT.. % Sample, Deviation f r o m Time, No. of RIg. Theoretical Value Min. Analyses Sodiuni fluoroacetate (19.0% F ; reflux medium, hexyl alcohol) 1 10 +i 100 10 -3 1 30 +2.9 8-13 5 10 8-20 -1.7 7 5 lrIethyl fliioroacetate (20.6% F b.p. 103-104.5° C. a t 750 i i i t n . I I g ; reflux medium, hexyl alcohol) 60 1 12 --3 10 4 12 -3.4 10
(B.p. 103.5-101.2° C. a t 756 mm. H g ) 6 17
-3.3
2-Fluoroethanol (29.7% F,h.p. 103.0-103.3° C. a t 745 mni. I l g ; refliix Iiieililini, hexyl alcohol, b.p. 123-156O C.) 375 3 6 -11.3 1 9 - 10 30 15 10
1
8
9 6-20
-11
-11.'
(Reflux medium, octyl alcohol, b.p. 191-193' C.) 30
1
9
- 12
(Reflus medium, tetradecyl alcohol, b . p . 26.5' C.) 60 "0
2 3
85-112 6
-10.9
-13.1 (Reflux mediiim, 30% hexyl alcohol, 50% m-xylene)
1 14 - 13 60 (Refliix medium, 50% hexyl alcohol, 50% isoamyl ether) 1 9 - 10 30
~ i j ~ o p ~ fliiorophospiiate ~ p y l (10.3% F, b.p. 68.4-69.4' C . at 9 111111. M g ; reflux nledium, hexyl alcohol) -0.5 20 1 60 -2.2 20 2 15 20 + 0.9 4 10
V O L U M E 23, NO. 4, A P R I L 1 9 5 1 A procedure is presented for the determination of fluorine in four physiologically active monofluorinated organic compounds when the compound is present in low- concentrations in air or water. The compounds studied were sodium fluoroacetate, methyl fluoroacetate, 2-fluoroethanol, and diisopropyl fluorophosphate. The absorption of the compound from air is accomplished by scrubbing with a higher alcohol such as hexyl alcohol. This alcohol is also adapted to serve as the high temperature medium for converting the combined fluorine to fluoride, which is subsequently determined either
cator. Add 1.2 W hydrochloric acid dropwise until the indicator changes from purple t o yellow-green; then add 4 ml. of buffer solution and titrate with 0.005 N thorium nitrate to the first permanent pink tint. I n the titration the yellow-green solution changes first to orange and then to pink. Consistent results can be obtained only when the end point is taken as the first appearance of pink color. Table I contains a summary of the results employing the volumetric method for determination of fluoride ion.
Table 11.
Comparison of Colorimetric and Volumetric Determinations of Fluoride Ion
Source of Fluoride Potassium fluoride Sodium fluoroacetate
hlethyl fluoroacetate 2-Fluoroethanol Diisopropyl fluorophosphate
Fluoride in Sample, hlg. 1.86 2.8i 2.38 2.19 4.00 2.49 2.13 1.72 2.59 2 66 2.07
Fluoride Found, Mg. ColoriYolumetric metric 1.8i 1.89 2 81 2.87 2.42 2.40 2.18 2.07 3.83 3.87 2.49 2.53 2.02 2.08 1.65 1.71 2.48 2.44 2.41 2.36 1.96 2.01
COLORlMETRIC DETERMINATION OF FLUORIDE ION
Standard Fluoride Solution. Dissolve either 0.0221 gram of sodium fluoride or 0.0495 gram of potassium fluoride in water a n d dilute t o 1 liter. The solution contains 0.010 nig. of fluoride per ml. Thorium Nitrate Reagent. Dissolve 0.24 gram of thorium nitrate tetrahydrate and 61.8 grams of sodium sulfate in water. To the mixture add 39 nil. of formic acid (specific gravity, 1.2) and 17.4 grams of sodium hydroxide; dilute to 500 ml. The solution is 0.00087 III with respect to thorium nitrate and 0.87 -If each n.ith reppect t,o sodium sulfate, sodium formate, and formic acid. Sodium Alizarin Sulfonate. Prepare an aqueous solution containing 0.0744 gram of sodium alizarin sulfonate, reagent grade, per liter. Preparation of Standard Series. Prepare a series of color standards t o cover the range 0 to 0.12 mg. of fluoride ion at 0.01nig. intervds. Transfer the required volume of Ptandard fluoride solution t o a 100-ml. volunietric flask, add 5.0 nil. of sodium alizarin sulfonate, and dilute to about DO ml. with water. Add 5.0 ml. of the thorium nitrate solution, dilut,e t o the mark, and mix t,horoughly. Transfer to 50-ml. Sesslcr tuhes (220 mni.). Fresh standards should be prepared daily. Procedure. Decompose the fluoro-organic compound and extract the sodium fluoride n-ith t\!-o 10-ml. portions of water. Transfpr an aliquot containing about 1 nig. of sodium fluoride to a 100-nil. volumetric flask, add 5.0 ml. of sodium alizarin sulfonate, and carefully adjust the acidity by adding 0.3 S nitric acid by drops until the color of the solution just turns yellow. Dilute t o approximately 90 ml., add 5.0 ml. of the thorium nitrate solution, dilute t o the mark, and mix thoroughly. .Ifter 30 minutes match the unknown against the standards.
605 by a modified thorium nitrate titration or by a colorimetric technique. In water, fluoride is formed by refluxing with a mixture of potassium metaperiodate, silver perchlorate, and perchloric acid. A modified Willard-Winter separation is used before titration of fluoride. Fluorine may be quickly determined in quantities of the compounds down to 0.01 p.p.m. in a 10-liter sample of air and down to 2 p.p.m. in a 20-ml. sample of water. Diverse ions usually found in potable water do not interfere with the reaction. Factors influencing the determinations are discussed.
The details of the colorimetric procedure were worked out by Overholser using mainly standard solutions of sodium fluoride (41). Concentrations of this compound down to 0.1 p.p.m. were determined with an accuracy of *5%. However, the method was readily applicable to the organic compounds. The values in Table I1 represent results obtained by Salsbury with a series of unknown solutions. RESULTS W-ITH PURE COMPOUNDS
Sodium Fluoroacetate. .lnalyses of samples of sodium fluoroacetate from several sources were about 10Y0high in fluorine on the basis of 100% sodium fluoroacetate. Independent determination of total sodium and of ionizable fluorine in each sample showed 8.5% sodium fluoride. The absence of ammonium and potassium ions was established by qualitative tests. When the analytical results on the samples n-ere corrected for sodium fluoride, the fluorine content of sodium fluoroacetate was about *3% of the theoretical. Essentially the same results n-ere obtained with samples from several sources. Methyl Fluoroacetate. The methyl fluoroacetate, as obtained through Office of Scientific Research and Development channels, was about 157c Ion, in fluorine on the basis of 100qc methyl fluoroacetate. .kfter purif>-ing a sample by drying over anhydrous sodium sulfate and fractionation, the fraction boiling from 103" t o 104.5' C. a t 750 nim. of mercury showed values averaging 3% below the theoretical. A second sample vias dried by adding benzene and distilling. A distillate was also fractionated and the portion boiling from 103.5' to 104.2' C. a t 756 mm. of mercury \vas analyzed. Rtsults n.ith this sample were also about 370 IOW. 2-Fluoroethanol. The 2-fluoroethanol, obtained through the Office of Scientific Rcsearch and Development from the Department of Physiology, T*niversity of Chicago, was dried over anhydrous sodium sulfate and distilled. The fraction boiling from 103.0" to 103.2" C. a t 752 mni. of mercury was analyzed for fluorine. Another .ainple obtained from Xonsanto Chemical Co., .inni?ton, was dried by distilling with benzene. The distillate was fractionated twice, and the portion boiling from 103.0" to 103.3" C. a t 745 nim. of mercury was analyzed. Both samples slion.ed essentially the same fluorine content, the value3 averaging about 11% blow the theoretical. Increasing the rtflux time from 10 minutes to 6 hours and varying the reflux medium did not alter the results. It seems possible that this substance is an azcotropic mixture of 2-fluoroethanol and some nonhalogen compound. .l microanalysis (by ,James 31. Frcdericksen) of a purified sample showed 38.93% carbon and 6.25y0hydrogen; calculations for 2-fluoroethanol shored 37.58% carbon and 7.801, hydrogen. Diisopropyl Fluorophosphate. Diisopropyl fluorophosphate contains a fluorine t o phosphorus bond. The sample from the Office of Scientific Reararch and Development \vas dried over anhydrous sodium sulfitte arid distilled under reduced pressure. The fraction boiling from 68.4" to 69.4" C. a t 9 mm. of mercury
ANALYTICAL CHEMISTRY
606 was used. The fluorine values were within ahout 2% of the theoretical. The following alcohols were tried as reflux media: henayl, cyclohexyl, n-hexyl, isoamyl, isobutyl, %methyl eyclohexyl, n-octyl, and tetradecyl. Only in *hexyl, n-octyl, and tetradecyl alcohols was decomposition of the fluoro-organic compounds essentially complete. Sodium alcoholate seems to he the active constituent in the mixture because refluxing a sample of sodium fluoroacetate in a hexyl alcohol solution of sodium hexanalate gave the same results as were obtained by refluxing the fluoroorganic compound with sodium and the alcohol. The reflux time needed for rupture of the carbon to fluorine or carbon t o phosphorus bond is roughly an inverse function of the boiling point of the alcohol. In isobutyl dcohol, boiling point 107" to 108' C., 60 minutes are needed for complete decomposition of sodium fluoroacetate) in hexyl dcohol (boiling point 153' to 156" C.) all compounds were decomposed in 10 minutes. Less than 10 minutes refluxing gave low values, h u t increasing the time did not increase the yield. The amount of sodium does not appear to be significant, provided a large excess over the fluorine content is used. Another factor is the density of the reflux alcohol; the lower the density the more efficient is the aqueous extraction of fluoride. I n the volumetric method all titrations were carried out in dilute ethyl alcohol in order to maintain B constant condition and to minimize factors which slightly affect the calor of sodium aliaarin sulfonate. The samples vaned in fluorine content from 0.1 to 35 mg. without appreciable effect on either the time required t o split the bond, or the volume of water needed t o extract the sodium fluride. However. the alisuot of the sodium fluoride solution (ohtained from the'decompbsition of the fluoro-organic compound), when titrated nith nitrat,e solution using alizarin indicator, should contain between 0.1 and 0.9 mg. of fluoride ion. Samples containing less than 0.1 mg. of fluoride may he titrated with thorium nitrate using Soloehrome Brilliant Blue BS as indicator (34). Good agreement resulted when diquats of several solutions previously analyzed, using sodium alizarin sulfonzte BS indicator, were diluted tenfold and then titrated with thorium nitrate using Solochrome Brilliant Blue. BS as indicator. Titrations using Solochrome indicator, however, m u d he performed in low-form Nessler tubes. In a titration the thorium nitrate was added until the color just matched the hluepurple of a blank cont,ainingindicator and thorium nitrate. A calibration curve is also required when using this indicator; careful control of pH is essential.
The fluorc+organic compound ahsarhed was decomposed and analyzed essentially BS described previously. The fluoride ion was usually determined by titrating in a low-form Nessler cylinder an aliquot (not over 0.1 mg. of fluoride ion) with 0.005 N thorium nitrate. The p H was adjusted by making the aliquot just acid to phenolphthalein with dilute perchloric acid (1 drop of 0.2 N ) , adding 1 ml. of Soloohrome Brilliant Blue BS indicator, and then adding 0.5 ml. of chlorortoetio acid buffer. The solution was diluted to 50 ml. and titrated until the color just matched the blue-purple of a blank. Table 111 contains a summary of the analyses of known gas mixtures. ANALYSIS O F WATER CONTAINING FLUOR-ORGANIC COMPOUNDS
Figure 2 shows the apparatus used to decompose the compound and t o distill the silicon tetrafluoride. A 60-ml. West condenser was connected to the side arm fitting C, which was wrapped with asbestos cord. The apparatus was ooustructed from 24/40 standard taper fittings. The joint connecting A and B had a narrow nbhon of Lubriseal; the others were not lubricated. The well contained just enough mercury t o cover the hulh of the +).o"m,.ma+n.
Additional Reagents.
Potassium metmeriadate, Eastman
~~~, _,, add 0.2 &am of potassiu& metape&d.&e, 2 grams of sllver perchlorate, aud 15 ml. of 70% perchloric acid. Pour 20 ml. of the ~ , ~ O . ~ .
add 0.2 gram of glass wool (fluoride-free) and connect t h e i?e& condenser. Ileat the mixture gently a t first and then continue
ANALYSIS OF AIR CONTAINING FLUORO-ORGANIC COMPOUNDS
Gas mixtures of known composition were obtained by passing air through a system of stopcocks and flowmeters and over the volatile fluorine compound contained in a vaporieer whioh oould he weighed. The vaporieer was a. 40-cm. length of 7-mm. borosilicate glass tubing, sealed a t both ends, with two pieces of 2-mm. tubing joined near the top for the side arms. The unit was connected to the flowmeter system by pure gum ruhher tubing with the ends of the side arms fitted snugly inside of the ends of the glass tubing connected to the system. This precaution was necessary hecause rubber tubing is slightly permeable. While weighing the vaporizer, the side arms were closed by inserting a short piece of glass rod into each of the rubber connectors. Before weighing the vaporieer, it wzs wiped carefully with a piece of linen. The specially designed scrubber for the air-gas mixture shown in Figure 1was filled with 40 ml. of the hexyl or octyl alcohol and the unit was immersed in baths maintained in three series a t 20', 0",and -40" C., respectively. At the end of a run the ahsorhing liquid was drained into a 100-ml.ground-joint,round-bottomed flask, and the gas scrubber was flushed with an additional 10ml. portion of the absorbing liquid.
Figure 1. scrubber for Air-Gas Mixtures
V O L U M E 23, NO. 4, A P R I L 1 9 5 1 Table 111, Determination of Fluoro-organic Compounds in Air ( R a t e of flow through scrubber = 1.0 liter/min.) Concn. of Volume Recovery of Compound, No. of of -iir, Fluorine, Rlg./Liter Analyses Liters Av. % Methyl fluoroacetate (20.67, F, b.p. 103-104.2° C. a t 756 mm. HE: bath temperature, 20' C.) 1-i 4 5 85 * 3 3 10 88 * 1 0 . 1 -0.3 0 01-0.03 4 20 67 * 4 (Bath temperature, O o C.) 1..2 1 3 105 * 6 0 1 -0.3 10 10 90 * 3 0.01-0, OT 6 10-20 63 * 7 (Bath temperature, -40' C.) 3-i 3
