CHAPTER 3

Analytical Chemistry of Fluorine and Fluorine-containing Compounds

BY PHILIP J. EL VINO University of Michigan CHARLES A. HORTON Carbide & Carbon Chemicals Company, Plant Oak Ridge, Tennessee

AND

HOBART H. WILLARD University of Michigan Page Introduction 53 I. Sampling of Fluorine-containing Materials 54 II. Analysis of Gaseous Samples 55 A. Determination of Fluorine 55 B. Determination of Hydrogen Fluoride 56 C. Determination of Other Gaseous and Volatile Inorganic Fluorides 57 D. Determination of Volatile Organic Fluorides 58 E. Analysis of Fluorine Gas 59 F. Analysis of Electrolytic Cell Gases 61 1. Gas Analysis 62 2. Molecular Weight Determination 65 3. Fluorocarbon Analysis 67 III. Separation and Isolation of Fluorine 67 A. Decomposition, Dissolution, and Other Preliminary Treatment of In­ organic Materials 68 1. Ashing Procedures 68 2. Fusion Procedures 71 3. Evaporation Procedures 73 B. Decomposition of Fluorocarbons and Organic Compounds 74 1. Oxidation Methods 76 2. Reduction Methods 78 3. Methods Involving Alkaline Fusion 81 4. Methods Involving Reaction with Silicon Dioxide 81 5. Hydrolytic Methods 83 C. Isolation of Fluoride by Volatilization 83 1. Distillation As Fluorosilicic Acid 83 2. Pyrohydrolysis: Evolution As Hydrofluoric Acid 89 3. Miscellaneous Volatilization Methods 90 51 52 PHILIP J. ELVING, CHARLES A. HORTON AND HOBART H. WILLARD

Page IV. Qualitative Detection and Identification of Fluorine 90 A. Fluoride Ion and Fluorine-containing Compounds 92 1. Etching and Hanging Drop Tests 92 2. Bleaching of Zirconium-alizarin and Similar Lakes 94 3. Miscellaneous Color and Fluorescence Tests 96 4. Precipitation Tests 97 5. Microscopic Tests 98 6. Miscellaneous Tests 99 B. Fluorine in Organic Compounds 100 C. Fluorine in Gaseous Samples 102 V. Determination of Fluoride Ion 102 A. Precipitation and Gravimetric Determination 102 1. As Lead Chlorofluoride 102 2. As Calcium Fluoride ! 104 3. As Rare Earth Metal and Other Metal Fluorides 108 4. Miscellaneous Gravimetric Methods 109 B. Titrimetric Methods: Precipitation and Complexation 110 1. Thorium Titration HO 2. Zirconium Titration 117 3. Titration with Iron (III) and Aluminum (III) 118 4. Titration with Cerium (III) and Rare Earth Metal Ions 119 5. Indirect Titration of Lead Chlorofluoride and Calcium Fluoride 121 C. Titrimetric Methods: Neutralization 123 1. Reactions Involving Fluorosilicic Acid and Potassium Fluorosilicate 123 2. Miscellaneous Neutralization Titrations 127 D. Electrometric Methods 128 1. Potentiometric Titration 128 2. Conductometric Titration 131 3. Amperometric Titration 132 4. High-frequency Oscillator Titration 133 5. Polarography 134 E. Photometric Methods 134 1. Colorimetric Methods Involving Bleaching 134 2. Miscellaneous Colorimetric Methods 147 3. Fluorometric Methods 149 4. Nephelometric Methods 150 5. Emission Spectroscopy 152 F. Miscellaneous Methods 155 1. Enzymatic Activity Inhibition 155 2. Etching and Wettability of Glass 156 3. Catalytic Activity 157 VI. Fluorine Compounds 157 A. Spectrophotometric Technics 157 1. Infrared Absorption and Raman Scattering 157 2. Ultraviolet Absorption 159 B. X-Ray Diffraction Patterns 16° C. Miscellaneous Methods Based on Measurement of Physical Properties. 161 D. Analysis of Fluorocarbons 163 1. Decomposition to Obtain Fluoride Ion 163 ANALYTICAL CHEMISTRY OF FLUORINE 53

Page 2. Determination of Constituents Other Than Fluorine 164 3. Determination of Fluorocarbons As Compounds 168 E. Assay and Analysis of Hydrofluoric Acid and Hydrogen Fluoride 170 VII. Determination of Fluorine in Specific Materials 171 A. Biological Samples 172 1. Plants 172 2. Animals 173 B. Fertilizers, Phosphates, and Phosphate Rocks 173 C. Foods and Beverages 174 D. Rocks, Minerals, and Ores 174 E. Water 175 F. Air 176 G. Miscellaneous 176 Bibliography 177

Introduction The analytical chemistry of fluorine is very different from that of the other halogens. This is emphasized by the differences in comparative solubilities of metallic salts, e.g., silver fluoride is very soluble while cal­ cium fluoride is sparingly soluble. As might be expected, the analytical chemistry of fluorine is dominated by those properties of fluorine which differentiate it from the halogens and other elements. An important factor in the detection, separation, and determination of fluorine is the volatility of silicon tetrafluoride and its ready formation in the presence of dehydrating acids. The unique reaction between hydrofluoric acid, and silica and silicates, resulting in the property of hydrofluoric acid of etching glass, is also of prime importance in the detec­ tion and determination of fluorine. Other factors of importance in the development of analytical methods for fluorine include the stable com­ plex species formed by fluoride with aluminum, iron, thorium, titanium, and zirconium ; the volatility of many fluorine-containing compounds such as the fluorides of boron, hydrogen, and silicon; and the usually amor­ phous or gelatinous nature of the comparatively few insoluble inorganic fluorine compounds. Fluorine is usually best separated from other possibly interfering ele­ ments by steam distillation as fluorosilicic acid from a perchloric or sul­ furic acid solution. In many cases the sample must be ashed or subjected to a fusion process as a necessary preliminary to the distillation step. The recovery of fluoride ion from organic compounds is usually difficult, owing to the great stability of the carbon to fluorine bond when additional fluorine or chlorine atoms are linked to the same carbon atom. The decomposition of the organic compound may require such drastic attack as high temperature fusion with an alkali metal. 54 PHILIP J. ELVING, CHARLES A. HORTON AND HOBART H. WILLARD

The analytical chemistry of fluorine has received tremendous impetus in recent years as a consequence, to cite only a few causes, of the demands of the atomic energy program for both inorganic and organic fluorine- containing compounds, the role of fluorine in dental caries with the conse­ quent demand for the determination of minute amounts of fluorine in water, and the importance of fluorine in insecticides and in agricultural raw materials and products in general. A moderate number of reviews of the analytical chemistry of fluorine or of special subdivisions of it have appeared. The latter are usually referred to in the appropriate section of the discussion. The most compre­ hensive reviews which have appeared recently are probably those of Hernler and Pfeningberger in 1938 (H58), Wilson in 1944 (W63), Rinck in 1948 (R37), Kurtenacker in 1949 and 1951 (K73 and K74), Element in 1950 (K39), the report published in 1950 on the analytical chemistry of the Manhattan Project of World War II (B122, M28), and McKenna in 1951 (M27). Other reviews and bibliographies are given in references B3, B15, E14, E21, F66, F67, F68, F72, G33, M28, M67, M100, N18, Rl, and W27.

I. Sampling of Fluorine-containing Materials The technics and precautions necessary in sampling various types of fluorine-containing materials can be readily located from the references given in Section VII for the different specific classes of substances as well as in other sections, e.g., electrolytic cell gases in Section II-F. The main factor to be normally considered in the sampling, handling, and preliminary treatment of fluorine-containing substances is the possi­ ble loss of fluorine through the formation of hydrofluoric acid, which would volatilize as HF or react with any silica or silicate present, such as the glass of the container, to form volatile silicon tetrafluoride. The volatility of boron and other fluorides is another possible source of fluorine loss. In any preliminary treatment involving heating or evaporation of an acidic solution, the possibility of HF loss should be considered. In the preliminary separation of the R203 group by precipitation in ammoniacal solution fluorine may be lost if calcium is present due to the coprecipita- tion of calcium fluoride (F73). The drying of the sample itself may result in some fluorine loss if the water can react with the sample at the drying temperature to form volatile fluorides. Chapman, Marvin, and Tyree (C47) evaporated at 200° mixed hydro­ fluoric and perchloric acid solutions containing compounds of some thirty-seven different elements. Losses due to volatilization of the fluorides of the elements were found in varying degree for boron, silicon, ANALYTICAL CHEMISTRY OF FLUORINE 55 germanium, arsenic, antimony, chromium, selenium, manganese, and rhenium. No loss was observed for compounds of the metals of the first two groups: lanthanum, cerium, titanium, thorium, tin, lead, vanadium, bismuth, molybdenum, tungsten, uranium, iron, cobalt, and nickel.

II. Analysis of Gaseous Samples

A. DETERMINATION OF FLUORINE The assay and analysis of fluorine gas as well as the analysis of fluorine-rich gaseous mixtures are discussed in subsequent subsections of the present section which deal with the analysis of fluorine gas and of the cell gases produced in the electrochemical process for making fluoro­ carbons. The detection of fluorine in gaseous samples is considered in Section IV. The detection and determination of fluorine and fluorine compounds in air is discussed in Section VII-F. Air samples are generally taken so as to separate particulate fluorides from the gaseous fluorides. Generally, electrostatic precipitators are used to remove the solids and the gases are absorbed in caustic solution using an impinger or gas scrubber (A18, B110, L55, W81, Zll). Paper has also been used to remove the dust (W54). A potentially general method for the determination of fluorine in gaseous samples is based on the stoichiometric displacement of bromine from a bromide over which the sample is passed and measurement of the bromine produced. Staple, Schaffner, and Wiggin (Si 13) described a procedure for continuous analysis of a gas in which the sample is led over sodium bromide maintained at 230 to 250°, followed by photoelectric measurement of the light absorption of the resulting bromine-containing mixture, using a narrow spectral band filter at 435 đŔě. The precision in the range of 0 to 30% fluorine is ±1.0% (relative) and in the concentra­ tion range of 30 to 40%, ±3% (relative). Nash (NI, N2) has described a similar procedure for the determination of fluorine in air or nitrogen, using sodium bromide heated to 150°. For samples containing 0 to 10% fluorine, the absorption is measured at 425 đŔě where the accuracy is within 0.5%. Samples in the concentration ranges of 10 to 20% are best measured at 525 đŔě. Oxygen does not react with the bromide at 150°. The sample is passed over sodium fluoride after the bromide process to remove hydrogen fluoride which might interfere by etching the photom­ eter cell. Sample rates of 20 to 100 ml. per minute were satisfactory. The results can be recorded. Staple and Grilly (SI 12) have used an analogous principle for the con­ tinuous analysis of fluorine-nitrogen mixtures by thermal conductivity. The latter technic is not satisfactory for the direct measurement of 56 PHILIP J. ELVING, CHARLES A. HORTON AND HOBART H. WILLARD fluorine-nitrogen samples, but passage of the sample over sodium fluoride to remove hydrogen fluoride and then over sodium chloride kept at 200° produces a chlorine-nitrogen mixture which is readily measured. The sensitivity is ±0.06% F2 and the reproducibility is ±0.12% F2. Other workers (H29, K23, T58) have used a similar method and have chemically determined the chlorine formed by the reaction of the fluorine with sodium chloride. Oxygen and inert gases were estimated by standard technics after absorbing the chlorine. In one case, oxygen has been sepa­ rated from fluorine by adsorbing the oxygen on charcoal on —180° (K5) ; this is not a very safe procedure. Determination of fluorine by measurement of the iodine produced on contact with iodide solution is described in Section II-E (see also H29). Obviously, any component in the sample which is capable of oxidizing iodide, bromide, or chloride to iodine, bromine, or chlorine under the experimental conditions will interfere in the methods just described. A number of absorption methods have been proposed for fluorine. The mercury absorption method (B52, B122, M81, S25), which is good to 0.5% for the range of 0.5 to 100% F2, is discussed in Section II-E. Sodium sulfite solution containing a minute amount of p-aminophenol will absorb fluorine but not oxygen; Orsat technic is less satisfactory than allowing the solution to flow into the gas buret (W53). Fluorine has also been determined by the increase in weight of silver (and of the copper tubing in the apparatus) (B66) ; the difficulties in the use of silver absorption are discussed in reference S25. Liquid fluorine shows no visible (M100) or infrared absorption spec­ trum; one reference (B67) claims, apparently incorrectly, that fluorine shows an ultraviolet absorption peak at 2820 A.

B. DETERMINATION OF HYDROGEN FLUORIDE

The determination of hydrogen fluoride, evolved as a result of pyro- hydrolysis or of the decomposition of organic fluorine-containing com­ pounds and of fluorocarbons, is discussed in Sections II-D, III-B, III-C, IV-B, and VI-D. The assay of liquefied hydrogen fluoride is discussed in Section VI-E. The analysis of aqueous solutions of hydrogen fluoride is discussed throughout the present chapter but particularly in Sections V-C and V-D. Manometric and other methods for the determination of hydrogen fluoride in fluorine are described in Sections II-E and II-F. Its detection and determination in air are discussed in Section VII-F; a superior method for concentrations as low as 1 p.p.m. is the ferrisal procedure described by Flagg (F40, VI1). ANALYTICAL CHEMISTRY OP FLUORINE 57

Kirslis and coworkers (K28) have described a differential flow analyzer for hydrogen fluoride-nitrogen mixtures which may also be applicable to the determination of hydrogen fluoride in samples containing fluorine. The sample stream is split into two streams, one of which is passed over sodium fluoride which removes the hydrogen fluoride. The resulting pres­ sure difference between the two streams is a measure of the hydrogen fluoride content. Hydrogen fluoride shows no ultraviolet absorption spectrum (S2) but has a characteristic infrared absorption pattern (see Section VI-A-1).

C. DETERMINATION OF OTHER GASEOUS AND VOLATILE INORGANIC FLUORIDES The analytical chemistry of boron trifluoride has been summarized by Booth and Martin (B79). They have described the procedure developed by Swinehart, Bumblis, and Flisik (S130) for the sampling and analysis of boron trifluoride, which includes the following procedures: (a) Sampling water-soluble and water-insoluble gases, and sample measurement. (6) Determination of air as material insoluble in 30% sodium chloride solu­ tion, (c) Determination of sulfur dioxide by addition of excess standard iodate-iodide solution and back-titration with standard thiosulfate solu­ tion, (d) Determination of silicon tetrafluoride by a complex alkalimetric titration, (e) Determination of sulfur trioxide by precipitation as barium sulfate. (/) Assay for boron trifluoride by differential neutralization titra­ tions for the various acidic constituents. Boron trifluoride can be determined when present in admixture with silicon tetrafluoride by its reaction with and absorption in nickel fluoride (M25) or acetyl fluoride (M27). The latter procedure was developed for use in nuclear cross-section studies employing BF3 as a reference gas. The boron trifluoride content of liquid organic ethers has been determined by heating the sample with sodium fluoride, evaporation to dryness, and weighing the sodium fluoborate formed (W9). The precision was ±0.5% (relative). Silicon tetrafluoride has been estimated by formation of a 1:1 triethyl- amine complex, whose dissociation pressures are 45 mm. at 25° and 1 mm. at —78°. By selective freezing, it was possible to analyze boron trifluoride- silicon tetrafluoride mixtures with an accuracy of 0.2% (G38, G39). Silicon tetrafluoride has also been estimated gasometrically by absorption in water (S45). It is possible to separate silicon tetrafluoride from hydro­ gen fluoride by adsorption on sodium hydrogen fluoride (NaHF2) at 80 to 250° according to one patent (R80). It may be noted that it is impossi­ ble to separate hydrogen fluoride from water, fluorosilicic acid, or both, by distillation, because of the formation of azeotropes (A4). 58 PHILIP J. ELVING, CHARLES A. HORTON AND HOBART H. WILLARD

Granular activated carbon retains all the silicon tetrafluoride con­ tained in air passed through it (M43). Carbon dioxide is adsorbed but is removed by continued passage of air. Hydrogen chloride dislodges the adsorbed silicon tetrafluoride. The sensitivity of the spectral determination of silicon tetrafluoride in nitrogen by a spark discharge has been investigated for varying experi­ mental conditions (R41). Sulfur hexafluoride and gaseous fluorocarbons have been separated by use of differences in their vapor pressures (N4). The determination of oxygen fluoride in electrolytic cell gases is described in Section II-F. A possible basis for an analytical method for oxygen fluoride in solution based on the oxidation of iodide is also de­ scribed by Dennis and Rochow (D44).

D. DETERMINATION OF VOLATILE ORGANIC FLUORIDES In general, the methods used in decomposing and determining volatile organic fluorine-containing compounds are identical with those used for the determination of less volatile compounds which can be, nevertheless, vaporized by temperature and gas sweeping. Accordingly, the section on the decomposition of organic compounds includes the procedures used in handling volatile organic fluorine compounds. The most successful of these methods for volatile compounds have included combustion of the gas in moist oxygen or hydrogen to form HF or over silica to form SiF4. As little as 0.025% carbon tetrafluoride in hydrogen can be detected by burning the gas and observing the etching of glass by the flame due to the formation of hydrogen fluoride (T52). A halide meter manufactured by Davis Instruments (D24, M70) can be used to determine many halogen-containing (fluorine, chlorine, or bromine) hydrocarbons in air in the concentration range of 0 to 500 p.p.m. with an accuracy of 10%. A photoelectric photometer is used to measure continuously the intensity of the blue lines of the copper spec­ trum produced in an electric arc between two electrodes. Halide vapor coming in contact with the hot tip of the copper electrode reacts to form a copper halide which vaporizes at the temperature of the electrode and is carried into the arc. The intensity of the blue region of the spectrum is proportional to the halide vapor present. The microammeter reading has to be calibrated in parts per million for the particular compound involved. A leak tester for industrial high vacuum manufactured by Distillation Products (D54) uses Freon-12 (CF2C12) as the probe gas. The sensitive element is a platinum diode in which the presence of a halogen catalyzes the positive emission of the anode; the emission, after amplification, is registered on a microammeter. Ten parts per million of halogen in air ANALYTICAL CHEMISTRY OF FLUORINE 59 can be detected. A halide torch for leak detection also employing Freon has been described (S23) which uses the fact that a trace of a halide gas in the air intake of a Bunsen-type burner will color the flame which impinges on a copper plate a bright green. Difluorochloromethane in air has been estimated by absorption in 20% sodium hydroxide followed by treatment with pyridine; the intensity of the red color formed in one layer of this mixture was compared to stand­ ards (K17). An ultraviolet spectrophotometric method has been used to determine the concentration of monohalogenated benzenes in air. The dried air was absorbed in ethanol at —50°, and the transmittancy meas­ ured at 260 đŔě in the case of fluorobenzene (A38). Kirslis and Staple (K29) estimated fluorocarbons in the air by pyrolysis of the gaseous sample in a platinum tube filled with platinum contacts and heated to 950°. The fluoride evolved was estimated by use of thorium or zirconium alizarin test papers. Zirconium-p-dimethylazo- phenylarsonate papers (H22), although sensitive to traces of fluoride, had a color change which was difficult to observe. A modified method in which the liberated dye was extracted in acetone was more successful (R2). Thermal decomposition over white-hot silica (H52) or with red-hot calcium oxide in a bomb (H53) has also been used in decomposing volatile organic fluorine compounds. Stable volatile fluorocarbons, such as the Freons, have been decom­ posed by burning the sample with oxygen in a silica-packed silica tube heated to 900°. The fluoride formed was absorbed in dilute caustic. Afterwards oxygen, nitrogen, hydrogen, and finally nitrogen again were passed through the tube to ensure complete recovery.

E. ANALYSIS OF FLUORINE GAS Early attempts to analyze fluorine gas have been listed by Bigelow (B52). A safe, routine method of analysis of fluorine gas has been de­ scribed by the duPont Company (T58). This method is suitable for con­ trolling the quality of the fluorine (a) as produced directly by electrolytic cells, (6) in the fluorine purification system (removal of hydrogen fluoride), (c) in plant pipelines, and (d) when packaged in cylinders after purifica­ tion and compression under pressures as high as 400 lb. per square inch. For samples containing over 50% fluorine, the gas sample is analyzed for fluorine, oxygen, hydrogen fluoride, and "inert" residue which is mostly nitrogen. The gas sample is passed through an analysis train con­ taining the following units which perform the functions indicated: (a) sodium fluoride pellets in a copper or nickel tube which remove hydrogen fluoride by the formation of the acid fluoride, HNaF2; (b) anhydrous sodium chloride which converts the fluorine to chlorine; (c) removal of a 60 PHILIP J. ELVING, CHARLES A. HORTON AND HOBART H. WILLARD sample which presumably now contains chlorine, oxygen, and inert gas; (d) cold 2 Ν sodium hydroxide solution which absorbs the chlorine to form sodium hypochlorite. The solution from step (d) is treated with potassium iodide and acetic acid to form iodine which is titrated with standard thio- sulfate solution. The chlorine in the gas sample removed in step (c) is determined by absorption in caustic; the oxygen is then determined by absorption in alkaline pyrogallol solution; the residual gas represents the "inerts." The total quantity of fluorine present in the original sample used is the sum of that found by the indirect thiosulfate titration and that found by the caustic absorption on the gas sample removed in step (c). The hydrogen fluoride is determined by maceration of the sodium fluoride pellets with cold neutral potassium nitrate solution and subse­ quent titration with silicate-free sodium hydroxide solution. The absolute error in the HF determination is less than 0.15%. The estimated over-all error in fluorine purity is about 0.4%. An analogous procedure has been described (K23) which permits the handling of 4 to 5 1. of fluorine from which perhaps only 10 to 50 ml. of residual gas is collected for identification and determination of impurities present in small amount. Hydrogen fluoride is first removed by anhydrous sodium fluoride; the gas sample is then led over sodium chloride for the replacement of fluorine by chlorine. The resulting product is then ana­ lyzed by a precision method developed for the complete analysis of chlorine. A large sample is absorbed in alkaline arsenite solution; the absorbed chlorine is determined by titration of the chloride ion using the Volhard procedure, and the absorbed carbon dioxide is determined by an evolution procedure involving removal from acidified solution and neutralization titration after absorption in dilute alkaline solution. The residual gas is analyzed in an Orsat apparatus for oxygen, carbon monox­ ide, hydrogen, and inerts. In one case the molecular weight of the latter was found to be 93 and was largely CF4 (molecular weight of 88) ; OF2, if present, will be absorbed in the arsenite solution. The HF is determined by crushing the sodium fluoride granules under standard sodium hydrox­ ide solution and back-titrating with standard acid. The hydrogen fluoride content of fluorine can also be determined by condensing a known quantity of fluorine at liquid nitrogen temperature, pumping off the fluorine, warming the container and residual hydrogen fluoride to room temperature, and measuring the pressue due to the HF (B122, T58). The error due to adsorption of hydrogen fluoride by metal parts of the apparatus is of the order of 0.1%; other errors may be caused by the presence of condensable silicon fluorides. The method has been applied to fluorine samples containing about 1 mole % HF. Miller and Bigelow (M81) described a complicated glass apparatus ANALYTICAL CHEMISTRY OF FLUORINE 61 for the analysis of free fluorine gas which depends on the reaction of fluorine with mercury to form mercury fluoride and uses manometric measurement. The residual gases are analyzed for carbon dioxide, oxygen, and inerts; hydrogen fluoride is removed from the gas but is not deter­ mined; OF2, if present, introduces serious errors. The method requires skilled operation and is apparently not too well suited for general work (K23). Bigelow (B52) in a review of the procedure claims that the prin­ cipal error is that involved in the pressure readings which is approxi­ mately ±0.15% and is at least partially compensating, so that the final values for volume % of fluorine are accurate to 0.1%. Fluorine, about 97.5% pure and containing mostly oxygen as impurity and some nitrogen, was analyzed for fluorine by shaking with mercury in a quartz buret and measuring the decrease in volume (S25). A correc­ tion had to be made for the adsorption of the residual gas on the reaction product. The method based on the absorption of fluorine by powdered silver is unreliable due to incomplete fluorine removal. Fluorine purity may thus be determined to within 0.5% over the range of 0.5 to 100% fluorine (B122); the residual gas may be analyzed by the usual Orsat procedure. In mixtures of inert gases and up to 10% fluorine, the latterô can be determined by passing the sample into dilute hydriodic acid in 1 Ë acetic acid and measuring the iodine produced (M81, T58). The methods developed by Simons and coworkers (S71) for the chemi­ cal analysis of the cell gases produced in the electrolytic process for pro­ ducing fluorocarbons can often be applied to the determination of the impurities in fluorine gas; these methods are described in the subsequent section on the analysis of cell gases.

F. ANALYSIS OF ELECTROLYTIC CELL GASES Simons and coworkers (S74) have modified and developed methods which are suitable for the analysis of the cell gases produced in the elec­ trochemical process for making fluorocarbons. Particular attention was given in the various methods to avoid interference by the other compo­ nents present. Procedures are described for (a) the determination of carbon dioxide, hydrogen fluoride, oxygen fluoride, oxygen, and carbon monoxide; (b) molecular weight determination; and (c) fluorocarbon analysis. Samples are withdrawn from the metal system through J4-inch or Y\ 6-inch tubing, although after the removal of hydrogen fluoride, }>£-mch. copper tubing is satisfactory. Where gases must be led below the surface of an absorbing liquid, J^-inch Saran tubing is used to avoid the corrosion found with iron or copper. The needle valves controlling sample with­ drawal are subject to extremely corrosive conditions. The valve which has 62 PHILIP J. ELVING, CHARLES A. HORTON AND HOBART H. WILLARD proven most satisfactory is the Hoke Model 341, modified by the manu­ facturer for use with hydrogen fluoride. The body, housing, and packing gland are of Monel, and the spindle and retaining ring of Z-nickel. The packing can be of the ordinary hydrogen fluoride-resistant type or of Teflon. These valves can be used on lines which must be evacuated.

1. Gas Analysis

The difficulties in determining the carbon dioxide content of the gases arise from the presence of hydrogen fluoride. The procedure in general is to remove the hydrogen fluoride by a specific reagent and then determine the remaining acidic constituents, regarded as C02. The customary titra­ tion for carbonate in the presence of mineral acids, involving the difference between the phenolphthalein and methyl red end points, is not satisfac­ tory in the presence of hydrogen fluoride since sodium fluoride is signifi­ cantly hydrolyzed at the acid end point. The indicator makes no sharp color change, and the value obtained varies with the concentration of the sodium fluoride. Hydrogen fluoride reacts with aqueous calcium acetate to form insolu­ ble calcium fluoride and acetic acid; carbon dioxide is substantially insoluble in such a medium. Although calcium fluoride precipitates as a gelatinous colloid which tends to foam badly, the addition of a few grams of ç-butyl alcohol to the absorption flask prevents this from being a serious problem. In addition, the absorption flask is constructed of a 50-ml. round-bottom flask with a neck of 20-mm. tubing, about 6 inches long, §o that sufficient height for disentrainment is provided but the size of the vapor space is kept small. The apparatus is shown in Fig. 1. Item 1 represents the calcium acetate absorber; 2, the 100-cc. gas buret with acidulated water as confining fluid; 3, the comparison buret for pressure equalization; 4, the leveling bulb for manipulation of the gases; 5, the carbon dioxide absorber, filled with 20% aqueous potassium hydroxide solution, and, if desired, glass tubing to increase the absorption surface; and 6, the reservoir for receiving displaced base from the absorber. Item 7 is the absorption flask used to replace 1 in the HF and OF2 determinations. In practice, the gas sample is drawn into the buret through the calcium acetate absorber at the rate of 10 to 20 bubbles per minute. The first part of the gas is discarded through the auxiliary stopcock until the connecting lines and vapor spaces are purged. The sample is then taken and standard absorption procedures followed. The results are in terms of volume of C02 per volume of acid-free gas. Hydrogen fluoride content of the gas stream is determined by absorp­ tion in a basic solution and titration of the excess base. The difficulties lie ANALYTICAL CHEMISTRY OF FLUORINE 63 in obtaining a true sample. Hydrogen fluoride reacts with and is strongly adsorbed by metals and packing material. Traces of moisture in the lines absorb large amounts and may lead to the deposition of aqueous acid. In addition, the low surface tension of hydrogen fluoride leads to entrain- ment in the gas stream so that the gas sample taken may be far lower in hydrogen fluoride than the main stream. In consequence, unless extreme care is taken, the analytical results show much less hydrogen fluoride removed than that which must be added to maintain the liquid level in the cell.

FIG. 1. Hydrogen fluoride absorption apparatus.

The apparatus shown in Fig. 1 is usee}, container 7 being substituted for 1. The absorbing liquid is 35 ml. of 0-1 Ν potassium hydroxide, which is usually sufficient for a 75-cc. gas sample. The gas passes through the absorber and is measured in the gas buret. The solution is then titrated with 0.1 Ν hydrochloric acid to the phenolphthalein end-point. After correction for the C02 content (one equivalent per mole) previously determined, the results are expressed in grams of hydrogen fluoride per volume of acid-free gas. Oxygen fluoride, OF2, is readily and quantitatively reduced by acidic potassium iodide solution. If the solution is only slightly acid, iodide is quantitatively oxidized to I2 by OF2, while the reaction with any oxygen 64 PHILIP J. ELVING, CHARLES A. HORTON AND HOBART H. WILLARD in the gas stream is negligibly slow. In more strongly acid solution the reaction with oxygen may give high results, and the etching by hydro­ fluoric acid on the glassware may be severe. In basic solution iodide is oxidized partially to iodate, and the analytical results may be low. Sodium dihydrogen phosphate was selected as a buffer to maintain the desired pH; enough should be used to take care of the hydrogen fluoride in the gas stream. For samples low in hydrogen fluoride but high in OF2, adjust­ ment may be necessary for the two moles of base produced in the reduc­ tion of each mole of OF2: OF2 + 4KI + H20 = 2I2 + 2KOH + 2KF The sample is taken as described for hydrogen fluoride analysis, absorber 7 being filled with a monosodium phosphate solution, 0.2 Í in potassium iodide. After absorption, the solution is titrated with thiosul- fate using starch as an indicator. Since the extent to which carbon dioxide will be absorbed is quite sensitive to pH in this region, in accurate work it is desirable to free the collected gases of C02 by absorption in the potassium hydroxide bubbler before the final reading is taken. Results are expressed as grams of OF2 per volume of OF2-free, acid-free gas. For comparison purposes, conversion of the weight figures to the volume equivalent is frequently more effective. Oxygen is determined by absorption in alkaline pyrogallol in the standard Orsat apparatus. Since oxygen fluoride would interfere, its removal before analysis is necessary. As mentioned in the preceding section, the action of oxygen on alkaline iodine solution is slow, while that of OF2 is rapid. The sample is therefore bubbled slowly through a basic solution of sodium sulfite, 0.05 Ν in potassium iodide. The iodide func­ tions as a catalyst for reduction of OF2 by the sulfite. In some cases where the absorbent was allowed to become acidic, as much as 75% of the oxygen was reduced by the iodide solution. The procedure in the Orsat absorption is standard. Results are expressed as volume of oxygen per volume of OF2-free, acid-free gas. Carbon monoxide is determined on the residual gases after removal of oxygen. An acid cuprous chloride solution has been used in the usual Orsat, but more reliable and faster reagents are available. Since pyrogallol produces carbon monoxide under some conditions, the use of an oxygen absorbent other than pyrogallol would be desirable when carbon monoxide analysis is to follow. At present, the residual gases are assumed to be only hydrogen and fluorocarbon. Procedures may have to be developed for determining nitro­ gen and nitrogen trifluoride where they are believed present in appreciable amount. ANALYTICAL CHEMISTRY OF FLUORINE 65

2. Molecular Weight Determination It is convenient for control purposes to determine the amount of hydrogen, the amount of fluorocarbons, and the composition of the fluoro­ carbons in the gas stream. These methods have not been as carefully studied as the chemical procedures and are described in more general terms. It is frequently necessary to determine the ratio of fluorocarbons to hydrogen. If the average molecular weight of the fluorocarbons is known, the vapor density of the purified gas is sufficient. In other cases, par­ ticularly where the fluorocarbon content is high, it may be more desirable

FIG. 2. Molecular weight apparatus. to determine the volume of hydrogen associated with a given weight or volume of condensable material. After chemical purification of the gases to remove hydrogen fluoride, carbon dioxide, and oxygen fluoride, there remains fluorocarbon, hydro­ gen, and at times nitrogen, nitrogen trifluoride, and oxygen. Special procedures applicable to the cases where the last are present in significant amounts may have to be developed; at present, the residual gases are assumed to be only hydrogen and fluorocarbon. Molecular weights are determined in the apparatus shown in Fig. 2. The evacuating system includes a mercury diffusion pump backed by an efficient mechanical pump. The gases are introduced through the phos­ phorous pentoxide drying tube 1 with the condensing trap 2 immersed in liquid air. The exact procedure will depend on the approximate ratio of condensable to noncondensable material. In general, the gases passing 66 PHILIP J. ELVING, CHARLES A. HORTON AND HOBART H. WILLARD

through the trap are allowed to collect until the pressure in the system is almost one atmosphere. Bulb 3, of 12-1. capacity, is used only when the fluorocarbon content is quite low. When the pressure has reached the desired value, the stopcock between the diying tube and trap 2 is closed, and the pumps are connected to remove the hydrogen. The pumps are disconnected, the trap is allowed to warm until a pressure of several centimeters is reached, and liquid air is again applied. If a significant pressure remains, it is advisable to cool trap 5 with liquid air while the residual gas is being removed. The process is repeated until no appreciable pressure remains when the sample is cooled in liquid air. While CF4 has a low vapor pressure in liquid air, prolonged pumping will cause some loss and a consequent error in the analysis. For most purposes a sample of sufficient size can be obtained so that negligible error is introduced from a residual pressure of 1 mm. The condensate is then allowed to expand into the same volume as that occupied by the noncondensable material and its pressure determined. The pressures will be in the same ratio as the original volume percentages. If the molecular weight of the fluorocarbons is desired, further pre­ cautions are needed to obtain a true sample, since simple vaporization will result in segregation of low- and high-boiling portions. In many cases the sample is small enough to permit use of trap 5 as a mixing chamber. The entire condensate is transferred to the trap and the two stopcocks closed. The material is allowed to vaporize and stand for several minutes to allow uniform gas distribution. The stopcock to the system is then opened, and the lines to bulb 4 are filled. For large samples, bulb 3 is used as a mixing chamber. Bulb 4, used to obtain the vapor density, is a round-bottom bulb of about 100-cc. volume, to which has been sealed a stopcock and the male half of a ground joint. The female half of the joint is connected through a second stopcock to the system mainfold. The volume of the bulb is deter­ mined by finding the weight of water it contains. The weight of the evacuated bulb is determined before it is connected to the system, and a careful routine must be adopted to assure reproducible weights after the joint lubricant is removed. When the system is filled with gas of uniform composition, the stop­ cocks are opened to the bulb, and the pressure and temperature of the gas noted. If the sample is to be recovered, the system contents can be condensed after the bulb stopcock is closed. The joint is carefully cleaned of grease, dried, and the bulb reweighed. For many control purposes, the molecular weight of the gas stream, after removal of acids and OF2, is needed. In such cases the purified sample is introduced to the evacuated system without condensation, and ANALYTICAL CHEMISTRY OF FLUORINE 67 the weight determined in the usual manner. Because of the low density of hydrogen, this method is inaccurate for low fluorocarbon content.

8. Fluorocarbon Analysis

Hydrogen-free material of high fluorine content, boiling below about 100°, can be analyzed by a modification of the sodium fusion method as described by Simons and Block (S72), which involved the decomposition of the material by heated sodium. Carbon is determined by weighing the water-insoluble residue and may be confirmed by the loss of weight on oxidation, substantially as described in the article. Fluoride may be determined in the water-soluble fraction by any standard procedure. For higher boiling fluorine-containing organic material, sodium peroxide decomposition in a Parr bomb is very satisfactory. For resins or fluorocarbons the normal procedure is satisfactory, but for material which might react with cold sodium peroxide, such as acids, the gelatin capsule technique is necessary (P23). For fluorocarbon carboxylic acids it has been found necessary to coat the capsules with paraffin to prevent their dissolving. The fusion mixture is dissolved in water and neutralized for the subsequent determination of fluorine. The thorium nitrate pro­ cedure is usually particularly useful here, since the alkali salt concentra­ tion is so high that frequently no precipitate at all forms in the lead chlorofluoride method. The determination of hydrogen in hydrogen-containing fluorocarbons involves the decomposition of the material over heated magnesium to liberate hydrogen, which is converted to water by hot copper oxide and weighed (P22, S73).

III. Separation and Isolation of Fluorine

Before the determination of fluorine as fluoride ion by any of the methods reviewed later in this chapter, the samples generally must be decomposed or dissolved in some manner, or both decomposed and dis­ solved. In general, these preliminary steps are different for inorganic fluorine compounds, organic samples containing ionizable fluorine, and organic materials containing nonionizable fluorine. Decomposition of the first two types of sample is considered in the first section which follows, and methods for other fluorine-containing organic compounds are described in the second section. A third section deals with methods for separating fluoride ion by volatilization from other ions which interfere with its determination. The first section also includes other methods for the separation of fluoride ion or fluorine compounds from interferences. 68 PHILIP J. ELVING, CHARLES A. HORTON AND HOBART H. WILLARD

A. DECOMPOSITION, DISSOLUTION, AND OTHER PRELIMINARY TREATMENT OF INORGANIC MATERIALS Ashing, evaporation, fusion, and other procedures used in the pre­ liminary treatment of samples prior to the actual isolation and determina­ tion of fluoride are described in this section.

1. Ashing Procedures Before dissolution of the sample and separation of the fluoride ion from other substances, samples which contain considerable organic matter such as foodstuffs, soils, and animal and plant tissues must usually be ashed to remove such organic material. Ashing also serves to concentrate the inorganic fluoride present, which is an important factor in the deter­ mination of trace amounts of fluorine. Ashing must be done in the pres­ ence of a fixative, i.e., a substance which will tie up all the fluoride in a nonvolatile form. Satisfactory ashing is probably the most difficult step in the various methods for the determination of fluoride ion in samples containing organic matter. The authors feel that the troubles experienced by analysts in many methods for the determination of fluoride may be largely due to loss of fluorine in the ashing process rather than to diffi­ culties in the subsequent volatilization or final determination of fluoride. Although varied ashing techniques and a considerable number of fixatives have been extensively studied, more critical work on the removal of organic matter is desirable. Ashing was necessary preceding the older method of separation of fluoride as silicon tetrafluoride, as in baking powder (P17) or biological material (Al), as well as prior to separation as fluorosilicic acid by the Willard-Winter Technique. Direct distillation of samples from sulfuric acid, first probably noted by Rosanow (R47), is a fairly recent develop­ ment which helps avoid loss during ashing. Maclntire and Palmer (M22) recommended only drying soil samples at 160° without a fixative, followed by direct distillation of fluoride from sulfuric acid at a solution temperature of 165°. The distillate was evapo­ rated and redistilled from perchloric acid. The Association of Official Agri­ cultural Chemists uses a direct distillation method in the presence of potassium permanganate and sulfuric acid for the separation of fluorine in insecticides (L33, L34). Wulle (W80) used no fixative in ashing blood but took care that the sample did not sinter in the gold crucible used. Wilson has discussed and reviewed the various ashing agents in use until 1944 (W63). ANALYTICAL CHEMISTRY OF FLUORINE 69

Calcium Compounds. The Aluminum Company of America ashes with lime, CaO, at 600° or less (A18). In one study Dahle (D3) recommended a calcium oxide fixative and ashing under 600°. The calcium oxide was purified to reduce its intrinsic fluoride impurity content in either of two ways. In one method, the oxide was treated with perchloric acid, fluoride was distilled off, calcium carbonate was precipitated from the perchlorate solution, and the carbonate was ignited to the oxide. In the second method, calcium oxalate was precipitated and ignited to the oxide (D3). Dahle has noted trouble in obtaining homogeneity with the fixative in certain types of samples such as cottonseed meal, oats, and corn (D9). Hoskins and Ferris (H73) ashed for 20 minutes at 720° when using cal­ cium oxide. Calcium oxide has been used as a fixative for biological-type samples (S37). The mixture was dried, the fat burned off, and the residue ashed at 600° in a silica dish; teeth were ashed 18 hours at 600 to 700°, and bones 12 to 24 hours at 650 to 700°. Other materials which have been ashed, using calcium oxide (as fixative), include food at 600° in platinum (G9), wine at 525° (R19), soil (M16), plants at 600° (C93), and gelatin at 600° (M95). Clifford (C72) used calcium hydroxide as a fixative and ashed in platinum at 550 to 600° for 1 to 2 hours. Maclntire and coworkers (M20, M22) recommended using 1 part of calcium hydroxide for 1 part of soil and ashing at 500° for only 5 minutes. Losses were noted for lengthy heating at 500° or for shorter periods at 900°. The amount of calcium fixative needed in ashing foods depends on the sample acidity, the amount of fluoride present, and the proportion of organic matter (D3). Blood has been ashed with calcium hydroxide for 1 hour at 450° in platinum dishes (M62), but only 89 to 92% of known amounts of fluoride was recovered. Similar recovery was found in ashing peach leaves with calcium hydroxide (S62). Wichmann (W36) also had trouble in avoiding losses with this fixative for samples containing carbohydrates. Winter (W67a) recom­ mended ashing plants with calcium hydroxide for 7 days. Food has been ashed, using two 10-minute periods at 650° and, as a fixative, copper acetate treated with calcium hydroxide until the mixture was basic (S22). The A.O.A.C. (L33) recommended calcium hydroxide or peroxide as a fixative for ashing soils first at less than 500° and then at 900° for 30 minutes; subsequently, they recommended 500° throughout (L34). For flours, pyrethrum, etc., they covered the sample with additional fixative. Maclntire in his earlier studies (M14, M16) found that calcium and magnesium peroxides were satisfactory as fixatives for soils or plant ash except for soils containing calcium silicate, for which he preferred cal­ cium hydroxide and an ashing temperature of 550°. Later (M15, M22), he found that extended heating of the soil with calcium peroxide at 900° is 70 PHILIP J. ELVING, CHARLES A. HORTON AND HOBART H. WILLARD unnecessary and that calcium hydroxide is unsatisfactory at that tem­ perature. Calcium peroxide has been recommended as a fixative in ashing superphosphate (M19). In a comparison (M18) of calcium and magnesium peroxides as fixatives, the calcium compound was preferred for samples of low organic and high silicate content; the magnesium compound was preferred for samples of high organic and low silicate content. Smith and Gardner (S86) found that direct ashing of blood at 350 to 750° with potassium hydroxide, sodium carbonate, calcium oxide, or calcium carbonate as fixative resulted in erratic low recovery of known amounts of fluoride. Synthetic mixes containing sodium fluoride, ferric chloride, and sodium carbonate were analyzed with a recovery of only 32% of the fluoride. They finally adopted a direct distillation treatment of the blood from sulfuric acid, followed by evaporation, ashing in the presence of calcium oxide at 575° for 15 to 16 hours, and final redistillation from perchloric acid. Others (L20) recommended similar treatment for blood. Calcium oxide or hydroxide for use as a fixativeN can be readily prepared by treatment of pure calcium metal with water. Also used as fixative agents have been calcium carbonate (F27), calcium acetate with chromic acetate and quartz for wood (12), calcium acetate for wood (C95), copper and calcium carbonates (W69), and calcium hydroxide (Bll, W66, W70, Z5). Maclntire, Jones, and Hardin (M21) found that fluoride was lost in the calcination of soils. Large proportions of calcium oxide or ashing at 900° caused decreased recovery of fluoride (see also reference M15). They again recommended a double distillation for the soil sample rather than ashing in the presence of calcium oxide. Others (L20) also favor similar decomposition treatment of soils. It must be noted that in a treatment of calcium fluoride and quartz at 500 or 900° for 50 minutes, only 40% of the fluoride was recovered (M22). Evaporation in the presence of calcium oxide without ashing has been used for some samples (B86, G10). Magnesium Compounds. McClure (M4) ashed at 500° using mag­ nesium acetate. Shewsbury (S59) used magnesium acetate for samples of limestone, rock phosphate, tooth, and bone, and recovered 95 to 99% of known amounts of fluoride after the Willard-Winter separation. His over­ all precision was ±3.6% of the fluoride present. In the case of grass, only 95 to 97% of the fluoride could be recovered when magnesium acetate was used and ashing was done at 500° (W68). For minerals, about 3% loss occurred when using magnesium acetate (G48). Other materials ashed using magnesium acetate include plants (D20, D21) and milk at 500° (M4). Ashing with a mixture of magnesium oxide and acetate at a dull red heat in silica dishes has also been used (B88). Magnesium acetate ANALYTICAL CHEMISTRY OF FLUORINE 71

was Tound to be a good fixative for ashing at 570°, whereas magnesium peroxide was only fair (C93). Magnesium peroxide has been found an unsatisfactory fixative for soils limed with calcium silicate (M14). Further study (M20) showed fluoride was not completely recovered from soils which were ashed first at 500° and then at 900° using magnesium peroxide or nitrate, fused with alkali metal carbonates, or heated at 900° with calcium hydroxide. However, full recovery was obtained in ashing organic composition samples with magnesium peroxide. A mixture of 1 g. magnesium oxide, 7 g. magnesium acetate, and 1 g. calcium oxide has been used as a fixative for ashing at 500° (Wl). Largent (L19) ashed food with magnesium peroxide in a nickel vessel at 570° for 6 to 12 hours. It was found unneces­ sary to add magnesium peroxide in ashing fat-free bones if they were heated for 3 hours at 600° (M8). Aluminum Compounds. Dahle (D4) has used aluminum nitrate as a fixative to prevent loss of fluoride during ashing to remove organic matter. This compound gave satisfactory results for samples containing organic acids but was unsatisfactory for eggs. Wichmann (W43, W36) also found aluminum nitrate to be a good fixative for ashing foods; how­ ever, the aluminum caused some trouble in the subsequent distillation by the Willard-Winter technique (W52). For a constant amount of aluminum added, increasing amounts of organic matter, fluoride, or organic acids caused loss of fluoride (W43). Increased aluminum content eliminated the losses. Ashing at 600 to 650° was satisfactory, but complete destruc­ tion of the organic matter was unnecessary when the titanium colorimetric method for fluoride ion was used. Troubles were also encountered in getting a homogeneous mixture of the sample with the fixative (W43). Winter (W66) used a small amount of aluminum silicate and more calcium hydroxide for plant and biological samples, and ashed at a dull red heat. Losses of fluoride using calcium oxide have been noted at 900 to 925° (Jl); in other studies only 90 to 98% of the fluoride has been recov­ ered (W67, W70). Miscellaneous Compounds. Copper metal or copper carbonate has been recommended as a good fixative (W69). Samples have also been ashed with sodium carbonate (C38, C90, C91, D63, E19, G54, H70). McClendon and Foster (M3, W40) use a closed platinum system for ashing and burning the sample in oxygen ; the evolved gases were collected in a solution of sodium hydroxide and azide. 2. Fusion Procedures Since the time of Berzelius (B49), refractory types of samples for determination of fluoride have been fused with an alkaline flux to convert 72 PHILIP J. ELVING, CHARLES A. HORTON AND HOBART H. WILLARD the fluoride to a soluble form. Such fusion methods are necessary for compounds which are not readily decomposed by sulfuric or perchloric acids in the usual Willard-Winter (W52) separation of fluoride from interfering substances. Such compounds include opal glass, certain silicates and borofluorides, and, often, cryolite, fluorite, beryllium fluoride, or complex metal fluorides in ores or minerals. Fusion procedures have also been used for zinc blende (K20), coal (C91), soil (M7), meat (N21), wood (W16), and basic slags (W15). The usual fusion flux has been sodium potassium carbonate with silica, as quartz, often added (B49, B76, C91, D67, H41, L24, L33, L34, L47, M7, M20, R27, S42, S77, S102, S103, S124, T46, W15, W16, W63, W66). The silica converts some of the refractory metal fluorides to soluble sodium fluorosilicate and also acts as a fixative to prevent loss of hydrofluoric acid during the fusion. Other fusion mixtures used in decom­ positions for subsequent determination of fluoride include sodium potas­ sium carbonate with titanium dioxide (T57), sodium hydroxide (K58, P52, Z9), sodium hydroxide with potassium nitrate (B62), sodium perox­ ide (K20), sodium carbonate with potassium nitrate (L47), sodium hydroxide with calcium oxide (N21), and magnesium carbonate (M20). In all cases where silica was in the flux or sample it was necessary to remove silicate before proceeding with the determination in most cases. In methods using the indirect volumetric titration of chloride after precipitation of lead chlorofluoride (K40, S103), or where silica was not present or added (see K20, L33, N21, S42, S102, T46), the silicate separa­ tion is unnecessary. The silicate has usually been removed by use of ammoniacal zinc carbonate (B49, B76, C52, C91, H66, H67, M14, R27, S77) or by boiling with ammonium carbonate (B49, K37, K58, P52, T57, W15) or other alkali (P23). One unusual separation involved pre­ cipitation with a reagent consisting of molybdate, 8-hydroxyquinoline, and hydrochloric acid (U2); another depended on insolubility of such contaminants in strongly alcoholic media (M7). In removal of silica with zinc oxide or carbonate, entrainment of fluoride in the precipitate must be avoided by thorough washing of the precipitate (D63). Until the advent of the Willard-Winter separation, it was also neces­ sary to remove phosphate from the aqueous extracts obtained after such alkaline fusions. Chancel (C45) probably was first to remove the phos­ phate as insoluble silver phosphate; silver fluoride is very soluble. This technique has been used since that time (C49, M7, W15, W16). Silver also removes arsenate and chromate if present in the fusion aqueous extract (W16). Other separations used after fusion have included removal of gallium (HI7) or aluminum (Z9) by precipitation as the hydroxides. ANALYTICAL CHEMISTRY OF FLUORINE 73

3. Evaporation Procedures It has been commonly assumed that evaporation of solutions, even to dryness, causes no appreciable loss in the original fluoride content of the solutions. Such evaporations are generally performed in ordinary boro- silicate glassware. Some work has indicated that these practices are not completely safe. An evaporation is often necessary, however, either to reduce the volume in order to facilitate separation of fluoride by the Willard-Winter technique, or to concentrate the fluoride in order to bring its concentration within the range of the colorimetric or volumetric method which it is desired to use. Loss of small amounts of fluoride ion during evaporation of slightly alkaline solutions was probably first reported by Reynolds and Hill (R25). They evaporated in glassware two-thirds of an original 150-ml. volume which was alkaline to phenolphthalein and noted losses of the order of 1%, with negligible losses below 1 mg. fluoride. McClure (M4) noted relatively greater losses for smaller amounts of fluoride (10 to 50 μg.) and found a difference in the magnitude of the loss depending on the type of vessel used for the evaporation. For borosilicate glass 76 to 99%, for porcelain 70 to 92%, and for platinum 93 to 104% of the known amounts of fluoride were recovered after evaporation of 150 ml. to less than 10 ml. Rinck (R37) noted similar errors due to evaporation, the effect being influenced to some extent by the final analytical method employed. He always noted slight losses using borosilicate vessels and recommended the use of platinum dishes or Jena glassware and evapora­ tion at a pH maintained between 6 and 8. Matuszak and Brown (M52) have also noted losses which they attribute to sodium silicate dissolved during the evaporation. Von Fellenberg (F28) reports that evaporation or boiling of water samples of pH less than 7 causes no loss of small traces of fluoride when the hardness (or calcium carbonate) content is low. Losses due to adsorption on the glassware were noted when the hard­ ness was greater than 14 grains per gallon. Evaporation has also been carried out in nickel (LI9) or in vitreous silica vessels as well as in the platinum dishes often recommended (B14, C38, C71). The A.O.A.C. per­ mits the use of either platinum or porcelain vessels for the concentration of water samples for analysis (L34). Obviously, evaporation to dryness in the presence of a fluoride fixative is often necessary before ashing to destroy organic matter, as discussed earlier in this section. The present authors have had other private communications relating to slight losses during evaporation but cannot agree with Rinck's counsel (R37) : "Avoid as much as possible all evaporation of solutions to be used 74 PHILIP J. ELVING, CHARLES A. HORTON AND HOBART H. WILLARD for the determination of fluoride." The actual volatilization of fluoride from alkaline solution seems doubtful, and the reported volatilization may actually be due to mechanical loss or interaction, including adsorp­ tion, with, the container used.

B. DECOMPOSITION OF FLUOROCARBONS AND ORGANIC COMPOUNDS

At the present time, research in fluorine chemistry is being carried on in many university and industrial laboratories on the preparation, prop­ erties, and uses of organic compounds containing fluorine, as well as of the class of compounds called fluorocarbons. The research on these com­ pounds and their increasingly widespread industrial application has called attention to the desirability of simple methods for determining fluorine, particularly in compounds where very stable carbon-fluorine bonds exist. The lack of an entirely adequate analytical method often presents a serious obstacle to research and industrial development. The unsatisfac­ tory nature of the methods formerly used in the analysis of these stable fluorine compounds is illustrated by the fact that one method which used a complicated apparatus took 4.5 hours to decompose a single sample. Other investigators noted that it took a week of heating at a high tempera­ ture to decompose some of these fluorine compounds so that they could be analyzed. A statement made a decade ago in an article (H54) dealing with the preparation of organic fluorides will illustrate this point: "Fluo­ rine analyses, which are difficult and tedious, were performed at crucial points only. In general, the analytical results have a tendency to be slightly low, and this is attributed to the great difficulty of obtaining a complete decomposition of these extremely stable compounds." However, as is usually the case, the need of better methods for analyzing compounds containing fluorine resulted in great activity in this area. In particular, attention was focused on halogen-containing hydrocarbons and fluorocarbons because of their importance to the work of the Manhattan District project and its successors. The determination of halogen, including fluorine, in organic com­ pounds consists in two principal steps: (1) decomposition of the sample to yield ionizable halogen, and (2) determination of the halide ion. Several satisfactory procedures exist for the decomposition of organic compounds containing halogen other than fluorine, including the Carius method of heating with nitric acid in a closed tube, oxidation by sodium peroxide in a Parr Bomb, combustion in oxygen, or hydrogénation in the presence of a suitable catalyst. Each of these decomposition methods has its advantages and its limitations. After decomposition of the chloro, ANALYTICAL CHEMISTRY OF FLUORINE 75 bromo, or iodo compound, various standard methods provide for the separation and determination of the halide ion. Difficulties are introduced in the analysis of aromatic fluorine com­ pounds or aliphatic compounds containing high percentages of fluorine because of the stability of the carbon to fluorine bond. Meyer states (M67) that the other halogens in halofluoro compounds can be determined by the Carius method or by lime fusion in glass without splitting out the fluorine. The presence of fluorine also may increase the difficulty of determining chlorine in chlorofluoro compounds; the chlorine in certain chlorofluoropropanes could be determined only by the Carius method, which required a whole week of continuous heating at 250 to 300° to yield quantitative results (H52). The Carius method, even if sufficiently drastic to decompose fluoro compounds, is unsatisfactory because the resulting hydrogen fluoride or hydrofluoric acid attacks the glass of the reaction tube. The older (i.e., before 1941) methods described in the literature, with few exceptions, either do not work for the more stable compounds or are described in insufficient detail and without sufficient experimental verification to permit their use. Most of the fluorine compounds for which analyses were reported are of comparatively low fluorine content, com­ monly having only a single fluorine atom in a molecule of high molecular weight. This choice of compounds made the methods appear better than they actually are for two reasons: (1) a relatively large error on the basis of fluorine present appears as a small absolute percentage error; and (2) applicability of the method to the more stable compounds containing two or three fluorine atoms on a single carbon atom is not tested. Since the standard texts and references on organic analysis still fail in most cases even to mention the topic, it is worth while to summarize the methods which have been proposed in the literature for the decomposi­ tion of organic fluoro compounds and the recovery of fluoride ion in a determinable form. Of the earlier literature Meyer (M67) and Bocke- muller (B63, B64) give the most extensive discussions of the topic, but the principle of the methods described is not clearly indicated nor is the method of determining fluoride ion shown; moreover, the treatment is incomplete and noncritical. Comparative data on three of the older methods are given by Huckabay, Busey, and Metier (II) for three mono- fluoro aromatic compounds (Table I). Brief reviews of the determination of fluorine in fluorocarbons and organic compounds have been given by Elving and Ligett (E21), Mc- Kenna (M27), McKenna, Priest, and Staple (M28), and Nikolaev (N18). Many of the procedures developed or described since 1941 have been per­ formed on micro (1- to 5-mg. samples), semimicro (5- to 50-mg. samples), 76 PHILIP .J ELVING, CHARLES A. HORTON AND HOBART H. WILLARD and macro (sample size over 05 mg.) scales. Table II classifies the methods which have been suggested ni the literature, and which are discussed ni the following paragraphs; this classification si based upon one used by Elving and Ligett (E20, E21).

1. Oxidation Methods Many fluorine-containing organic compounds sa well sa fluorocarbons can eb decomposed by passage ni a stream fo oxygen through a platinum tube usually heated ot 900° ro higher. Water vapor si commonly added to aid ni forming HF, which si absorbed ni water ro ni alkaline solution. Gaseous and volatile liquid fluorocarbons are decomposed by passage with nitrogen and moist oxygen through a platinum tube ta 1000° ; the HF produced si absorbed ni water and determined yb the colorimetric titanium procedure (S34). nI view fo the relatively low precision fo the latter method, its use ni determining fluorine ni the range fo 20% may eb questionable. Compounds containing C, H, O, and F and boiling above 60° have been similarly decomposed ta 1250°; the HF was determined alkalimetrically after absorption ni water (M84). Carbon can eb absorbed from the dried gas ni the usual manner.

TABLE I Comparative Determination of Fluorine (H76)

% Fluorine

B.p., Etch Peroxide Na-NH3 Compound °C method (B94) bomb V8 Calcd.

Fluorobenzene 85 19.5 * * 19.8 p-Fluorotoluene 117 17.2 17.4 * a-Fluoronaphthalene 212 13.5 13.1 13.2 13.0

* Method not suited for this compound.

Traces fo fluorine (0.001 ot 0.25%) ni organic compounds have been recovered yb vaporizing the sample ni a stream fo methane, burning the latter mixture ta a jet, passing the resulting mixture with oxygen over platinum heated ni a combustion tube ot 600 ot 800°, and absorbing the HF formed ni water (H76). The fluoride si determined by a modified Willard and Winter method. For samples containing 0.26 ot 0.01 % F, the average deviation si 4%; for 0.005 ot 0.001% F, 6%. Of the older methods ni the literature, fluorine ni alkyl fluorides was determined (M98) by combustion ni oxygen ni a copper tube containing a mixture fo cupric oxide and lead monoxide. Bockemùller (B63) burned ANALYTICAL CHEMISTRY OF FLUORINE 77

TABLE II Decomposition of Fluorochemicals and Organic Compounds Containing Fluorine I. Oxidation methods A. Combustion in oxygen B. Fusion with sodium peroxide C. Alkaline oxidation II. Reduction methods A. Ignition in hydrogen B. Treatment with sodium in liquid ammonia C. Treatment with alkali metal in or­ ganic solvent D. Alkali metal fusion III. Methods involving alkaline fusion A. Fusion with calcium oxide B. Fusion with sodium carbonate or hydroxide IV. Methods involving reaction with A. Corrosive action on siliceous material silicon dioxide B. Combustion over silicon dioxide using oxygen and hydrogen V. Hydrolytic methods samples of organic fluorides mixed with calcium carbonate in a platinum boat in a platinum tube; the fluorine was converted into calcium fluoride; the results were good to ±0.3 absolute %. It has been suggested (M65) that the fluorine in gaseous organic fluorides be determined by burning the compound in oxygen on a hot platinum wire to form hydrofluoric acid. Alkyl fluorides have been burned in air or oxygen in a platinum tube to give HF which was absorbed in alkaline solution and determined by thorium titration (G72). Fluorine has been determined by combustion under 25 atmospheres pressure in a calorimetric bomb containing potassium iodide and iodate; iodine equivalent to the fluorine was liberated and determined by titra­ tion with standard thiosulfate solution (PI6). In one of the more recent oxidation procedures (C66) which employs standard microchemical combustion apparatus and technic, the sample (3 to 5 mg.) is burned in a platinum-packed tube kept at 900°, and the exit gases are passed through a Grote-type receiver containing water. The acid content of the latter solution is determined by titration with 0.01 Ν carbonate- and boron-free sodium hydroxide solution. Sulfur and other halogens in the sample are removed by silver in the combustion tube filling. Metals in the sample interfere by the retention of fluorine; nitrogen and phosphorus may form acids which are measured with the hydrofluoric acid. It is emphasized that the products from the micro- combustion of fluorine compounds in oxygen may on absorption form HBF3OH due to the presence of boron in the quartz or borosilicate apparatus ; this may result in low fluoride values, especially on neutraliza­ tion titration. The error can be corrected by titration of the boric acid 78 PHILIP J. ELVING, CHARLES A. HORTON AND HOBART H. WILLARD after the addition of mannitol; each mole of boric acid measured is equivalent to three Ă or HF. This method gives very satisfactory results. Since HBF4 is a stronger acid than HF, the present authors believe that it would not cause low results in a neutralization titration and that the error would not be corrected by titration in the presence of mannitol. The determination of fluorine by oxidation of the organic fluoride with sodium peroxide in a Parr bomb seems to be successful with many compounds (C48, D51, F35, H8, L30, L44, M31, M80, N7, P12, P39). The filling commonly used in the bomb consists of 0.5 g. sample (not over 0.1 g. fluorine), 0.5 g. sugar or starch, and 14 g. sodium peroxide. In the case of fluorocarbons, a mixture of 8 to 9 g. sodium peroxide and 1.2 g. of a mixture (3 parts potassium nitrate to 1 part sucrose) is suggested (M28). Organic fluorosilanes have been decomposed using 0.3 to 0.4 g. sample, 15 g. sodium peroxide, and 1 g. starch (P23). Erratic results have been reported in some cases for the peroxide method (M27). The concentration of salts obtained in a peroxide fusion procedure is apparently too high to permit use of PbCIF precipitation or thorium titration except with calibration and the use of a correction. Certain fluoro compounds explode spontaneously when mixed with the peroxide in an open vessel (S57). Kuster and Neunhoffer (K76) claimed to be able to decompose organic fluorides by ignition in a nickel crucible with calcium peroxide or with a mixture of sodium carbonate and sodium peroxide. The fluorine content of fluorine derivatives of proteins was determined by fusion with sodium hydroxide and potassium nitrate in a nickel dish, and precipitation of barium fluoride along with barium sulfate from the resulting extract (B62). The appreciable solubility of BaF2 is a drawback. Yoe (Y8, Y12) has suggested decomposing fluoroacetates and fluorophos- phates by refluxing with a mixture of potassium periodate, silver per­ chlorate, and glass wool in perchloric acid solution.

2. Reduction Methods The reduction methods for converting fluorine in organic compounds to fluoride ion use either hydrogen gas or alkali metal. Ignition in hydro­ gen yields HF, which can be absorbed in water or alkaline solution and subsequently determined. Treatment with sodium or potassium either as such, or dissolved in liquid ammonia or an organic solvent, results in formation of alkali metal fluoride. In either case the fluoride ion can then be determined by various methods: titrimetric, gravimetric, or colori­ metric. In addition, HF can be measured in some cases by neutralimetric titration. Cadenbach (B86) suggested vaporizing the fluorine compound in a ANALYTICAL CHEMISTRY OF FLUORINE 79 stream of hydrogen and burning the resulting mixture; a lamp method has been similarly used to determine small amounts of fluorine (0.1%) and sulfur in hydrocarbons (M52). Volatile compounds containing fluo­ rine and sulfur have been burned in hydrogen in a transparent and resistant plastic apparatus (N3). A simple lamp apparatus has been described (W71) in which the sample is vaporized into a combustible gas (H2); the mixture is then burned at a jet in air. The apparatus could apparently be adapted to fluorine determination by simply altering con­ structional details. Tetrafluoromethane has been detected by burning the suspected gas with hydrogen and noting the etching action of the flame (T52). Fluorine and chlorine in methane and benzene derivatives were determined by burning the sample with hydrogen and a little air in a narrow platinum capillary and analyzing the resulting condensate for hydrofluoric and hydrochloric acids. Although Chablay (C41) described in 1914 the use of sodium in liquid ammonia for decomposing organic halogen-containing compounds, it was not until 1931 that Vaughan and Nieuwland (V8) reported that fluoride ion could be recovered from organic fluorides by treatment of the liquid ammonia solution of the compound with metallic sodium. This method was also used by Govaert (G59). McBee and coworkers (M76) suggested using a sealed tube technic which would permit shaking the reaction mixture for any period of time necessary to decompose the compound; fluoride is determined by precipitation as PbCIF, followed by Volhard determination of the chloride in the precipitate. The precision and accuracy vary from 0.2 to 1 relative %. Organosilicon fluorides have been reported as not reacting with sodium in liquid ammonia; they were de­ composed by peroxide fusion (P23). Nonvolatile organic fluorides have been decomposed using sodium and liquid ammonia in the presence of an inert solvent such as diethyl ether or tributylamine (M31). Fluorine in fluorobenzene was recovered by keeping at 100° for several days a sealed tube containing ecjual volumes of the compound and dry benzene, and sodium wire (W8). Whearty (W30) refluxed alcoholic solutions of fluorine-containing chlorobenzenes with metallic potassium and determined the fluoride ion formed by precipitation as calcium fluoride; the results reported for fluorine fitted a probable formula in only one of the three analyses reported. More recently, o-fluorodiphenyl was decomposed by refluxing with sodium in isopropyl or sec-butyl alcohol (U3). Fluorine in minute amounts of air-borne organic compounds was determined by absorption of the compounds in anhydrous C6 to C8 alcohol and refluxing with sodium (S5) ; recoveries varied from 50 to 105%. The latter reference describes briefly several methods 80 PHILIP J. ELVING, CHARLES A. HORTON AND HOBART H. WILLARD developed for decomposing small amounts of organic compounds which are mentioned only in reports under the Office of Scientific Research and Development and the National Defense Research Committee (Y9, Yll, Y10). Xylene has also been used as a medium for the decomposition of fluorine compounds by reflux with metallic sodium (R72). The presence of fluorine had been shown by fusion with sodium and evolution of hydrofluoric acid from the resulting properly prepared residue (L48). Ruff and coworkers (R61, R63, R64) reported decomposing fluoromethane, fluoroform, and trifluoronitromethane by allowing the gases to stream into an evacuated quartz tube containing heated sodium in an iron boat; complete destruction was apparently not secured even after repeated passage over the hot sodium metal, and the amount of residual gas had to be determined. Simons and Block (S72) were able to decompose fluorocarbons by an analog< u J technic with fairly good results. Piccard and Buffat (P41) analyzed fluorobenzene for fluorine by heating the compound to 400° with metallic potassium in an evacuated tube: the fluoride ion formed was determined by conductometric titration with calcium chloride. The results on three samples of fluorobenzene were ±0.4 absolute % of the correct value. Potassium fusion has been used to detect the presence of fluorine in aromatic compounds (B17). The most thorough investigation of the alkali metal reduction technic for the determination of fluorine in organic compounds was made by Elving and Ligett (E20). The procedure which they developed depends upon decomposition of the compound by heating with sodium or potas­ sium metal in an evacuated sealed glass tube at a moderately elevated temperature, i.e., 350 to 450°, and determination of the resulting alkali fluoride by standard methods. The most applicable procedure for the latter seems to be precipitation of the fluoride ion as lead chlorofluoride and determination of the chloride content of the precipitate by titration with standard silver nitrate solution (F73, H41), using Volhard's method with the addition of nitrobenzene to avoid the necessity of filtering the precipitated silver chloride (C4). The same technic served for the analysis of chloro, bromo, and iodo compounds and has several advantages over the procedures commonly employed in these determinations. Gaseous, liquid, and solid compounds from the aliphatic, aromatic, alicyclic, and heterocyclic series with fluorine content as high as 60% and two or three fluorine atoms on the same carbon atom were successfully analyzed. The procedure permits the simultaneous determination of chlorine and fluo­ rine in chlorofluoro organic compounds. A considerable number of modifications of the alkali metal fusion procedure have appeared, usually from the viewpoint of avoiding the use of a glass tube and of using higher temperatures. Potassium or sodium ANALYTICAL CHEMISTRY OF FLUORINE 81 fusion in nickel bombs such as those used in the Parr peroxide procedure have given very satisfactory results (E16). Recommended practice in­ cludes (a) heating at 500 to 550° for 2 hours, followed by a Willard and Winter distillation and thorium or aluminum titration (G37, K23); (b) heating at 550 to 700° for 1 hour, followed by precipitation of PbCIF (B18); (c) heating at the maximum temperature obtainable with a Bunsen burner for 30 minutes, followed by thorium titration (B36).

3. Methods Involving Alkaline Fusion

Many investigators have endeavored to apply the Liebig lime fusion method to organic fluorine-containing compounds. Attempts to carry out the fusion in glass resulted in the formation of some silicon tetra­ fluoride and in partial loss of sample (S24). The use of a platinum tube improved the results, but the latter were still in error by 1 or more absolute % (B31, D47, D49). Meyer and Hub (M68) asserted that to remove fluorine linked to an aromatic nucleus it is necessary to decompose the compound at 1000°, which they did by heating the sample mixed with calcium oxide in an open nickel tube to yellow heat for 2 hours. The results on o-fluorobenzoic acid and o-fluorobenzamide were appreci­ ably low. Better results \vere later reported (F33, S80) in analyzing the three fluorobenzoic acids and analogous compounds by this method. Ruff and Keim (R65) found the correct fluorine to chlorine ratio on passing trichlorofluoromethane mixed with steam and nitrogen over hot calcium oxide in a platinum tube. Henne and Renoll (H53) analyzed chlorofluoro- propanes by decomposition over red-hot calcium oxide in a small steel bomb. Unfortunately, no experimental details were given except the chlorine to fluorine ratios found, the accuracy being ±0.05 for calculated ratios of 0.75 to 1.67. In a paper (M30) applying this method to the determination of chlorine in organic compounds, no mention is made concerning the details of the fluorine determination. Fluoroacetyl sugars are decomposed by careful ignition with sodium carbonate (B93) ; some other compounds have been decomposed by heat­ ing with sodium carbonate at 500 to 550° (C38). Ethyl phenylfluoro- acetate was decomposed by prolonged fusion with sodium hydroxide (F61).

4. Methods Involving Reaction with Silicon Dioxide The ready formation and volatility of silicon tetrafluoride and fluoro­ silicic acid under suitable conditions has been utilized in various methodr s proposed for the determination of fluorine in organic compounds. Se\ eral procedures have been used to determine the amount of fluorine from the 82 PHILIP J. ELVING, CHARLES A. HORTON AND HOBART H. WILLARD empirically calibrated loss in weight of a glass plate exposed to the hydro­ fluoric acid formed when the compound is suitably treated with sulfuric acid, or by the similar loss in the glass vessel in which the reaction occurs (D15, H46). In the most recent procedure (B94) of this type, the sample is treated with potassium nitrate and sulfuric acid in a Pyrex flask, 1 mg. loss in the weight of the flask corresponding to 1.2 mg. of fluorine. The reaction between alkforyl silicon halides and silicon or silicon copper alloy at 500 to 1000° has been suggested as having possibilities for remov­ ing fluorine and other halogens from volatile fluorocarbons (110). Minute amounts of fluorine compounds in water were readily con­ verted to fluoride ion by refluxing the solution with potassium meta- periodate, silver perchlorate, perchloric acid, and glass wool; the fluoride was removed by distillation as SiF4 (S5); recoveries of 64 to 105% were obtained for 0.002 to 0.2 mg. compound per milliliter. Hubbard and Henné (H75) described a method for the simultaneous determination of chlorine and fluorine in aliphatic compounds which consisted of passing a measured amount of the gaseous or vaporized com­ pound mixed with air or oxygen through a quartz tube heated to 900° and packed with crushed silica. The tube is then swept successively with oxygen, nitrogen, hydrogen (to remove absorbed silicon fluoride and to reduce silicon oxyfluoride), and nitrogen. The exit gases are passed through an absorbent solution containing sodium hydroxide and hydrogen peroxide, and the fluoride and chloride formed are determined in separate portions of the solution. It was later mentioned that only the gaseous and highly volatile chlorofluoropropanes could be analyzed by decomposition over white-hot silica (H52). Bigelow and coworkers (C6) described a procedure based on Hubbard and Henne's method; a single combustion took 4.5 hours, exclusive of time needed to prepare and weigh the sample and to determine the halides in the absorbent solution. Volatile organic fluorides yield SiF4 on ignition with oxygen in a silica tube packed with crushed silica (M31). In 1938, a patent (H49) was issued for detecting volatile organic fluorine derivatives in air, which was based on bringing the suspected air into contact with incandescent silica and passing the resulting gas into contact with ammonia to form a white cloud or precipitate. Halocarbons free of hydrogen can be decomposed by passage in an oxygen stream over platinum and quartz powder at 1000°. The SiF4 formed is adsorbed on alumina at 175° (T25); fluorine is calculated from the gain in weight. The precision is poor. Chlorine or bromine are absorbed by metallic silver at 295°; carbon dioxide by Ascarite. Hydrogen forms water which is also adsorbed on the alumina. It has been claimed that decomposition of fluorine-containing compounds is possible in oxygen ANALYTICAL CHEMISTRY OF FLUORINE 83 and quartz at 450 to 900°; the SiF4 was absorbed in KF solution (N17). Carbon and hydrogen are also determined in the same combustion.

5. Hydrolytic Methods The stability of the carbon to fluorine bond is lessened, and the fluorine can be removed by hydrolysis, if oxygen is linked to the carbon in a carbonyl linkage such as in acid fluorides; d-glucosyl fluoride lost all its fluorine on 10 minutes heating in acid solution (H45). Acetyl fluoride was decomposed by heating with an aqueous solution of calcium chloride (M64). Swarts (S128) asserted that fluorine in the side chain of aromatic compounds could be removed by heating to 200° with concentrated sulfuric acid or by heating to 150° with water in a sealed tube. The fluorine in alkyl fluorides is converted to silicon tetrafluoride on standing for 7 or 8 days with sulfuric acid in a buret over mercury (M99). Boron 0-diketone fluorides are decomposed by refluxing with alcoholic or aqueous potassium hydroxide (M109). The quantitative decomposition of gaseous halocarbons by steam has been reported (C77) ; on the basis of this fact, fluorine in solid halocar­ bons which may also contain hydrogen, oxygen, chlorine, bromine, and nitrogen can be determined by decomposition on a micro scale (1- to 10-mg. samples) at 1100° in steam over platinum in a combustion tube made of quartz, platinized quartz, or platinum (R34). The fluoride is determined by its bleaching effect on ferric ion-salicylic acid solution which is used to absorb the decomposition products. The average devia­ tion was 0.1 to 0.5% for 10 to 49% F. Liquid halocarbons distill too rapidly through the combustion system. Sulfur interferes by forming species which reduce the ferric ion in the measurement system, and phosphorus by forming phosphoric acid which complexes ferric ion. Alkali and alkaline earth metals interfere by forming fluorides which are not completely decomposed by pyrohydrolysis. The pyrohydrolytic decomposition of inorganic fluorides is discussed elsewhere in this chapter. Organic fluorosilanes are not hydrolyzed on prolonged refluxing with base except for methyl compounds (K61).

C. ISOLATION OF FLUORIDE BY VOLATILIZATION

1. Distillation As Fluorosilicic Acid Until twenty years ago fluoride was usually separated from interfering substances as silicon tetrafluoride. Before 1900, Wohler (W75), Penfield (P27, P28), Fresenius (F59), Offermann (D2), and others (B34, C21, C20, L20, Sill, T6, W22) had described the separation of fluoride as silicon 84 PHILIP J. ELVING, CHARLES A. HORTON AND HOBART H. WILLARD tetrafluoride. Their work was followed by the studies of Treadwell and Koch (T49), Wagner and Ross (W4), Reynolds and coworkers (R27, R28, R31), Shuey (S61), and others (Bll, C32, K2, K3, L2, L24, M57, MHO, M112, M113, M115, P17, P36, P57, R40, R47, R48, R49, R50, S60, T46, T48) on the same method for volatilizing the fluoride. Since 1940, only Bien (B51) has used this technic. This method of separation has been superseded by separation of the fluoride as fluorosilicic acid and thus will not be discussed in detail here. For contrast with the method generally used at present, the apparatus and procedure used by Bien (B51) will be discussed briefly later in this section. In 1933, Willard and Winter (W52) developed the method for sepa­ rating fluoride from interferences as fluorosilicic acid. A similar method was proposed independently at the same time by Tananaev (T7). In this method the fluorine is liberated as gaseous fluorosilicic acid by the action of perchloric (or sulfuric) acid in the presence of silica, as quartz, glass, or porous plate. The fluorosilicic acid is carried out of the apparatus with water vapor, while the temperature of the boiling solution is held at 135° by addition of water (or steam), and regulation of the heat applied to the solution. Willard and Winter (W52) found that all the fluoride was usually volatilized in the first 50 to 75 ml. of distillate. They also found that the presence of boric acid, aluminum salts, or gelatinous silica in the sample retarded volatilization of the fluoride. The method is applicable to all inorganic fluorides and samples which can be decomposed by perchloric or sulfuric acid and is suitable for plant ash, from which fluoride cannot be recovered quantitatively by volatiliza­ tion as silicon tetrafluoride. Opal glass and other materials which are not decomposed by perchloric acid digestion must be fused with alkaline fluxes as discussed elsewhere in this chapter. The use of perchloric acid naturally requires previous destruction of organic matter, in most cases when such matter is present. Shuey (S64) first demonstrated that strong sulfuric acid is fairly satisfactory for the direct distillation of samples containing small amounts of organic matter. Winter (W68) and Shuey (S63), in other early work on this separation method, indicated that recovery of fluoride from plant ash may be incomplete, whereas recovery is quantitative from aqueous solutions of water-insoluble, perchlorate- soluble fluorides. Rinck (R37) recovered over 99.5% of the fluorine from fluorite and cryolite in the first 100 ml. of distillate in the absence of retarding ions. Reynolds (R23, R24) showed that 100 to 150 ml. of distillate must be collected for quantitative evolution of fluoride from phosphatic samples. In some types of such phosphatic samples, a coating of silica accumulated in the distillation flask after a time, retarding evolution of the fluoride. He removed such deposits by boiling the flasks ANALYTICAL CHEMISTRY OF FLUORINE 85 in strong caustic. A similar cleaning technic is often used at present in laboratories using the Willard-Winter separation. Reynolds (R24) also investigated distillation from phosphoric acid solution; sufficient phos­ phate came over in the distillate to give erratic, high results, unless the distillate was concentrated (alkaline) and redistilled from perchloric acid. A later paper by Reynolds and Hill (R25) further considered determina­ tion of fluorine in phosphatic materials. They recommended distillation from perchloric acid for phosphate rock. Distillation in the presence of excess potassium permanganate improved separation of fluoride from pyritic sulfur and organic matter. Churchill, Bridges, and Rowley (C64) recommended double distillation, from sulfuric and then from perchloric acid, for samples containing phosphate. Dahle and Wichmann (D12, D14), in an extensive study of the sepa­ ration as fluorosilicic acid, showed that, within limits, the recovery of fluoride is inversely proportional to the flask input volume and directly proportional to the temperature of the solution between 125 and 155° for sulfuric acid, and that variations in recovery are greater for distillation from perchloric or phosphoric acids than from sulfuric acid. They found that the fluorine recovery varied with the distillate volume collected according to the equation: ) (c 0.020945 = Κ = \ log , . (1) ι (c — X) where Κ = a constant, t = milliliters of distillate collected, c = original fluoride concentration in flask, and (c — x) = fluoride concentration in flask when t ml. have been distilled. Markova (M39) and Dahle and Wichmann (D13) also found that large amounts of aluminum ion retarded distillation of fluoride. For com­ plete recovery, a higher distillation temperature must be employed or a larger volume of distillate collected or both. Recovery was better from sulfuric than from perchloric acid under similar conditions, so they recom­ mended a first distillation of 300 to 350 ml. at a flask solution temperature of 162 ± 2°. The distillate was made alkaline and evaporated to 5 ml. in porcelain dishes and then redistilled from sulfuric acid at 137 ± 5°, col­ lecting 1 1. of distillate. The first distillate was concentrated to speed recovery in the second distillation. The second distillation was necessary because too much sulfate came over at the higher temperature used. Dahle and Wichmann (D14) in further studies of the distillation of fluorosilicic acid from perchloric, sulfuric, or phosphoric acid solutions found that the addition of soluble salts of nonvolatile acids caused a greater decrease in rate of fluoride recovery than would be expected and that recovery from perchloric or phosphoric acid solutions varies 86 PHILIP J. ELVING, CHARLES A. HORTON AND HOBART H. WILLARD

with the amount of distillate collected according to the logarithmic relation, equation 1, found for distillation from sulfuric acid. Dahle (D7) studied the "blank" fluoride due to the equipment used in this distillation. The "blank" was smaller when Vitreosil rather than glass beads was used to supply the silica. They noted that 20 μg. of fluoride always seems to be retained by the apparatus. However, many others have been able to recover as little as 5 Mg- of fluoride quantitatively. Later, Dahle (D8) recommended two distillations at 135 ± 3°, first from perchloric and second from sulfuric acid, for fluoride in baking powders. The first distillate was made alkaline and evaporated before the second distillation. He also noted that acid in the distillate causes slightly high results in the volumetric method using thorium. To reduce acidity in the distillate, Dahle, Bonnar, and Wichmann (Dll) kept the flask temperature under 140° and shielded it to avoid superheating. Kraft and May (K61) distilled ashed blood once from sulfuric acid at 130 to 140°. However, Wulle (W80) found that a second or even third redistillation was necessary if fluoride was high. Hartmann, Chytrek, and Ammon (H24) added peroxide to the sulfuric acid to aid in decomposition of the blood. McClure (M4) recommended additions of a soluble silver salt, the sulfate or perchlorate, to the distillation flask to avoid distillation of hydrochloric acid and to fix chloride, bromide, and iodide in an insoluble form. His distillation blank, including the thorium titration blank, was only 0.5 Mg. of fluoride. Maclntire and Hammond (M18) compared the addition of steam to that of water in the distillation of fluoride from perchloric acid. Steam required half as much time as water for the distilla­ tion of a given volume and also decreased bumping of the solution. Further studies of the use of a silver salt and steam were made by Mix (M91, M92), Donovan (D62), and van der Merwe (M62). Donovan (D62) compared four methods for the determination of fluoride. Clifford (C72) made a further study of the distillation of fluoride from perchloric acid solution. He found that the apparatus "blank" could not be completely eliminated and that 100% recovery is almost never ob­ tained for up to 200 Mg. of fluoride when 150 ml. of distillate are collected using a temperature of 135°. Somewhat better results were obtained using sulfuric acid. He again noted (C75) the retarding effect of gelatinous silica, aluminum, and boron compounds. For good recovery of fluorides from borates, Kazarinova-Oknino (K15) distilled fluorosilicic acid from sulfuric acid solution at 155°. Workers at the Atomic Energy Commission Projects (B122, VI1) have generally separated fluoride by single distillation as fluorosilicic acid from perchloric acid in a glass apparatus. However, they have also ANALYTICAL CHEMISTRY OF FLUORINE 87

employed gold or platinum apparatus. For blood, a direct distillation from sulfuric acid, followed by concentration, ashing, and redistillation from perchloric acid, has been used (S85). Recently, Maclntire, Hardin, and Jones (M20) used a direct double distillation for the determination of fluorine in soils rather than an ashing step followed by a single distillation. They found that some fluoride may be lost in the ashing, even in the presence of calcium hydroxide as a fluoride fixative. They first distill the soil at 165° from sulfuric acid, then concentrate the distillate and redistill from perchloric acid at 135°. Koehler (K46) has presented a mathematical treatment to determine how many distillations from perchloric acid are necessary for quantitative recovery of fluoride. The most recent edition of the A.O.A.C. book, " Official Methods of Analysis of the Association of Official Agricultural Chemists" (L34), follows their earlier practice (L33) and recommends a double distillation, first from sulfuric acid and second from perchloric acid at 130 ± 5°, for separation of fluoride in analysis of soils or economic poisons; a single distillation from sulfuric acid for fluorosilicate insecticides when boron, aluminum, or gelatinous silica are absent, with permanganate also pres­ ent when traces of organic matter are present; a single perchloric acid distillation for food, food preservative, and water samples. Several types of apparatus have been designed to control the distilla­ tion of fluoride as fluorosilicic acid from interferences. Klement (K38) introduced the use of a Claisen-type flask to reduce spattering. Gilkey, Rohs, and Hansen (G38) designed an apparatus heated by tetrachloro- ethane (b.p. 146°) under reflux, while Huckabay, Welch, and Metier (H77) proposed a simpler apparatus using the same principle and heating compound. Reynolds, Kershaw, and Jacob (P52) designed apparatus for six simultaneous determinations using steam and electrical heating for the solution in the flasks. Harris and Christiansen (H21) use a Vigreaux- type column to keep acid out of the distillate by means of the column reflux. Richter (R33) describes an electrically heated apparatus for fluoride determinations using a pressurized steam bath chamber to con­ trol the temperature. Churchill (A18, C63) also designed a multiple still arrangement using steam and gas heat. Willard, Toribara, and Holland (W51) developed an electronically controlled apparatus which controls the distillation temperature by the proper addition of water to the flask. One of the authors' laboratories (C.A.H.) uses a similar electronic control on the electric flask heaters, providing water by addition of electrically produced steam at a constant rate. Kilian (K20) supplies both water and steam to an electrically heated flask and collects the distillate under a slight vacuum in a Wolff-type receiver. Snowdon and Petrillo (S99) stir 8 PHILIP J. ELVING, CHARLESA.HORTONANDHOBARTHWILLAR

TYPICAL Y HUCKABA R RICHTE CLAISEN TYP E (H77) (R33 )

FIG . 3. Apparatus used for separation of fluorine by volatilization . ANALYTICAL CHEMISTRY OF FLUORINE 89 the liquid in a special flask with a magnetic stirrer to reduce bumping and supply heat with an infrared lamp and "Glas-Col" heater. For textiles they decompose the cloth and separate the fluoride directly using sulfuric acid with addition of water to keep the solution at 135 to 140°. Typical types of apparatus are shown in Fig. 3.

2. Pyrohydrolysis: Evolution As Hydrofluoric Acid One of the notable analytical developments of the atomic energy program in methods for the determination of fluoride was the separation of fluoride in solids from interferences by pyrohydrolysis, which has only been described briefly by Warf (W14). The method depends on the hydrolysis of heated hydrolyzable halides by steam. The steam also carries the volatile hydrogen halide out to a condenser whereupon it is absorbed in water and titrated with standard base. The furnace tempera­ ture for the solid is usually about 1000°. The method works well for fluorides and chlorides such as UF4, ThF4, A1F3, BiF3, MgF2, ZnF2, rare earth metal fluorides, and the corresponding chlorides. The metal ion is converted to the oxide in the process. Alkali and alkaline earth metal fluorides, chlorides, and beryllium fluoride, which hydrolyze with difficulty, can only be decomposed by hydrolysis in the presence of excess U308, A1203, Cr203, or sometimes V20B, which accelerate the process. Complex fluoro salts such as the alkali fluorosilicates, fluoroborates, and fluoroniobates also hydrolyze at 1000° in the presence of U308. A platinum reactor is necessary for decom­ position of fluorides if quantitative recovery of hydrofluoric acid from the sample is desired. Davis (D25) used a similar technic for fluoride in zirconium metal. He oxidized the metal in a stream of moist oxygen in a platinum tube at 700 to 900°. The hydrofluoric acid formed was collected in the condensed water vapor and titrated with thorium nitrate. The method should work on other hydrolyzable metals. Cline and Westbrook (C77) used a similar pyrohydrolytic technic for the determination of volatile fluorocarbons in air. This pyrohydrolysis technic should be tried for analysis of soils, minerals, fluorosilicates, fluoroborates, and other types of samples which require a long Willard- Winter (W52) distillation for quantitative recovery of fluoride. In some of these cases a hydrolysis promotor such as U308 or A1203 will surely be necessary. A pyrolysis type of decomposition for organic fluorine compounds which yields, in most cases, hydrofluoric acid is a part of several methods (C2, C66, M65, M84, N3, R34) discussed in another portion of this chapter. 90 PHILIP J. ELVING, CHARLES A. HORTON AND HOBART H. WILLARD

Ginsburg (G43) and Zschacke (Z10) separated fluoride from aqueous sulfuric acid solution by distillation as hydrofluoric acid.

8. Miscellaneous Volatilization Methods Willard and Winter (W52) first studied volatilization of fluorine as hydrofluoroboric acid (HBF4) from sulfuric or perchloric acid, as suggested by Schwerin (S36). The resulting recovery of known amounts of fluoride under varied conditions was very erratic. The best results were obtained by prolonged distillation of a sample with 1 g. boric acid, 10 ml. 1:1 sulfuric acid, and water, at a temperature below 130°. Recovery of 4.2 mg. fluoride ran from 3.3 to 4.2 mg. in eight runs. Later Fasano (F10, F9) suggested distillation of fluoride from interferences as boron trifluoride. Interferences in the distillate were removed by ether extraction from an aliquot made strongly acidic with hydrochloric acid. Recently, Revinson and Harley (R21) recommended separation of the fluoride as fluoroboric acid and claimed quantitative recovery when the amount of boric acid used is kept at a minimum. Bien (B50) published a method for determination of micro amounts of fluoride which separates the fluoride as silicon tetrafluoride in a very simple apparatus. The sample is mixed with quartz sand in a modified Feigl distillation flask and covered with concentrated sulfuric acid. On heating, the evolved silicon tetrafluoride is caught in a drop of water on the protruding thorn of a joint which closes the flask. The present authors do not believe this procedure would be very satisfactory for most samples. Schulek and Rozsa (S31) separated first boron and then fluoride from water by volatilization. The boron was removed as methyl borate from a solution containing hydrochloric acid, zinc chloride, and methyl alcohol at 100°. The fluoride was steam-distilled from the residual hydrochloric acid-zinc chloride solution at 160 to 180°. This is the only case known, to the authors where fluoride has been volatilized from a hydrochloric acid solution. IV. Qualitative Detection and Identification of Fluorine The detection and identification of fluorine has largely depended on two general reactions of fluoride ion. The first involved the reaction of fluorides with sulfuric acid to evolve hydrogen fluoride, whose etching action as hydrofluoric acid made the test positive or which reacted as hydrofluoric acid with silica to evolve silicon tetrafluoride, which in turn reacted with water to form a mixture of silicic and fluorosilicic acids. The second reaction depended on the bleaching action of fluoride ion on colored metal-dye complexes due to the formation of stable metal-fluoride ANALYTICAL CHEMISTRY OF FLUORINE 91 complexes. The gelatinous nature of so many of the insoluble metal fluorides has caused precipitation tests to be much less extensively used than the solubilities of the metal fluorides might warrant. It is obvious that many of the methods used for the separation and for the determination of fluorine in both inorganic and organic compounds can be used as the basis for tests for the detection as well as the identifica­ tion of fluorine. As is true of the analytical chemistry of all the elements, qualitative tests can often be developed into quantitative methods, and the reverse process can also take place. In particular, methods based on colorimetric and spectrophotometric measurement technics can serve for qualitative analysis. In the succeeding sections, particular attention will be paid to the most commonly used methods without attempting to describe all the methods suggested for the detection of fluoride ion; as already indicated, a large fraction of the methods suggested for the quantitative determination of fluoride ion have also been advocated as being suitable for its identification. Reference is made to typical pro­ cedures without any attempt to include all possible references. Procedures for the detection of fluoride in various particular types of material can also be located through the references given in Section VII. Suitable collections of laboratory procedures for the detection and identification of fluoride will be found in the reference works of Tread- well-Hall (T47), Furman-Scott (F73), and Feigl (F14, F12, F15). Infor­ mation about qualitative tests for fluoride will also be found in the reports of the International Committee on New Analytical Reactions and Reagents of the International Union of Chemistry (N14, N15, N16, W25). These reports are critical evaluations of reagents proposed in the litera­ ture and give the principle of the method for each reagent, the sensitivity in micrograms per milliliter, and, usually, some indication of possible interferences. Hernler and Pfeningberger (H58) have reviewed the litera­ ture on the detection of fluorine on a micro scale up to 1936. Insoluble fluorides and silicates are often best first decomposed by fusion with sodium carbonate, followed by water extraction. Calcium fluoride is not decomposed entirely by sodium carbonate fusion unless some silica or silicate is added to the melt. The resulting aqueous fluoride solution can be used for testing for fluoride ion. Fluorosilicates are decomposed on heating with sulfuric acid, liberat­ ing HF and SiF4; the etching and hanging drop tests can be used. On heating them, only SiF4 is evolved (use hanging drop test); the residue gives normal fluoride reactions. Soluble fluorosilicates form a crystalline precipitate, BaSiF6, with barium chloride (solubility, 0.27 mg. per milli­ liter) and a gelatinous precipitate with potassium chloride (the solubility, 1.2 mg. K2SiF6 per milliliter, is much decreased in the presence of excess 92 PHILIP J. ELVING, CHARLES A. HORTON AND HOBART H. WILLARD potassium chloride or alcohol) (T47). Addition of a base causes decom­ position to fluoride and silicate; ammonia causes precipitation of silicic acid. Fluorosilicic acid is not decomposed by water. Fluoride can be removed from many interfering substances by distilla­ tion from borate-sulfuric acid solution as BF3, removal of the boron by ether extraction, and colorimetric determination of the fluoride by ferric- thiocyanate or zirconium-alizarin methods (F10). Two of the most important compounds involved in the detection of fluorine are its compounds with hydrogen and silicon. Hydrofluoric acid is a moderately strong acid (pKa = 3.15) and can be distinguished from all other acids by its unique ability to dissolve silica, silicic acid, and silicates which is practically observed as its power of etching glass. The rate of the reaction between silica and hydrofluoric acid is depend­ ent on the specific nature and degree of subdivision (fineness) of the silicious material. Silicon tetrafluoride is a colorless gas with a penetrating odor; it reacts with water to form fluorosilicic acid, H2SiF6, and silica; it reacts with hydrogen fluoride to form fluorosilicic acid.

A. FLUORIDE ION AND FLUORINE-CONTAINING COMPOUNDS

1. Etching and Hanging Drop Tests Most fluorides when warmed or heated with concentrated (or prefer­ ably 90% H2S04) sulfuric acid will produce free hydrofluoric acid. If the reaction is carried out in glass, the latter will be attacked by the product with the formation of volatile silicon tetrafluoride and nonvolatile fluorosilicates of the sodium, calcium, or other metals present in the glass. These fluorosilicates are usually converted by the sulfuric acid to volatile hydrogen and silicon fluorides and nonvolatile sulfates. The presence of fluoride can then be confirmed by the action of either hydrogen fluoride (hydrofluoric acid) or silicon tetrafluoride. The former in its very characteristic reaction with silicates will etch glass, whereas the latter will cause water to become turbid as a result of its hydrolysis to form gelatinous silicic acid and fluorosilicic acid. To facilitate the tests, silica is often added to the mixture of fluoride and sulfuric acid. These tests are generally successful except when (a) the fluoride is mixed with a large excess of a form of silica, which is very readily reactive with hydrofluoric acid so that a stable oxyfluoride (SiOF2) may be formed (D19); or (b) the test is applied to certain refractory fluorine-containing minerals such as topaz and tourmaline (T47). The test is best made in a platinum vessel using a sample containing a relatively large amount of fluoride and comparatively little of an amorphous form of silica. Quartz ANALYTICAL CHEMISTRY OF FLUORINE 93 is not readily attacked by hydrofluoric acid. Boron apparently interferes in both of these tests, e.g., cf. reference M66. These tests are preferably applied to the powdered solid sample although they can be used on fluoride solutions with some loss in sensi­ tivity; evaporation of the neutral or alkaline solution to dryness and treatment of the residue seems warranted. As little as 0.1 Mg. is readily detected under optimum conditions (W57). Silicates and other compounds not readily attacked by sulfuric acid should first be fused in platinum with a large (five to ten times by weight) excess of sodium carbonate, followed by extraction of the melt with water. The silica can then be precipitated by the addition of ammonium carbonate and removed. Calcium fluoride, preferably with calcium car­ bonate as carrier, is then precipitated, the calcium carbonate in the precipitate removed by dilute acetic acid extraction, and the residue used for the test. The etching test is preferably made by heating the powdered sample with concentrated sulfuric acid in a platinum or lead crucible covered with a polished scratch-free glass plate. The inner side of the plate is covered with a uniform layer of wax with a small mark being made through the wax to expose a small area of the glass surface. To prevent the wax from melting, the outer surface of the plate is generally cooled by a condenser made of an Erlenmeyer flask, a beaker of ice water, or a piece of ice (the latter is usually too messy a procedure). After exposure for an hour more or less, depending on the nature of the sample and its fluorine content, the wax is removed with hot water, and the clean, dry plate is examined by reflected light for signs of etching. The occasional relative insensitivity of the test has been ascribed to various factors, such as large internal volume of the apparatus, leakage of evolved hydrogen fluoride, failure to expose the glass on scraping the wax away, use of too low a temperature, and insufficient time of exposure. Various workers have suggested improvements in the technic (e.g., B18, F12, K44, M2, W79). Caley and Ferrer (C5) described an easily machined lead apparatus which enables as little as 25 Mg- of fluoride to be detected on only 30 minutes of heating. Fetkenheuer (F32) found that quantities of fluoride too small to produce a visible etching of glass can be detected by the following test. If the sample is warmed with sulfuric acid in a test tube, the latter acid will no longer flow smoothly over the wall of the tube on tilting the latter but will coalesce in drops similar to water on a waxed surface. The effect is undoubtedly due to some type of alteration of the glass surface (Flla). The microchemical detection of fluoride by changes in the wettability of 94 PHILIP J. ELVING, CHARLES A. HORTON AND HOBART H. WILLARD glass has been thoroughly investigated by Hagen (H7) and by Dubnikov and Tikhomirov (D72). The latter have evaluated the conditions for the detection of fractions of a microgram of fluoride as well as the possible effect of a variety of other ions on the sensitivity of the test. The limit6 of identification is given as 0.5 μg., and the concentration limit as 10~ (F14). Typical procedures for the etching test under various conditions and for various types of samples will be found in references A17, B12, Bill, C5, C96, D46 (foods and alcoholic beverages), D79, E29, F14 (rocks and mineral waters), F73, G17 (waters, minerals, tissues), G18, G32 (body fluids and tissues), Ml, N22 (fluorosilicates), P15, T47, and W57. In the hanging drop test the sample is mixed with silica and heated after the addition of sulfuric acid. If a crucible is used, as in the etching test, a drop of water is placed on the inner surface of the cover. If a test tube is used, the tube is closed with a grooved, one-hole rubber stopper through which passes a short glass rod or tube ; a drop of water is placed on the inner end of the rod (the end may be blackened with asphalt paint to enhance contrast), or several drops are placed in the tube. The appearance of a gelatinous deposit of silicic acid in the water indicates the presence of fluoride. This test is apparently particularly successful for the detection of fluorine in silicates. Daniel (D19, T47), who has critically studied the method, developed a procedure sensitive to 0.1 mg. CaF2. Browning (B109, F73) described an interesting variation in which the mixture of sample, silica, and sulfuric acid is heated in a lead-covered cup which has a small hole in the cover. The hole is covered with a piece of moistened black filter paper. A white deposit appears on the paper in the presence of 1 mg. of CaF2 or 5 mg. of Na3AlFe. In another variation, the SiF4 is allowed to collect in a drop of water on a glass surface, e.g., a microscope slide; evaporation of the water results in formation of a ring of silica ("persistent ring" test) (M44, M45). Typical procedures are given in references B60, B79 (fluoroborates), B109 (silicates and fluorosilicates), C33, F12 (rocks and mineral waters), F73, G36, M1, N22 (fluorosilicates and mixtures with organic substances), R70, S9, S27 (zinc ores), and T47. 2. Bleaching of Zirconium-alizarin and Similar Lakes DeBoer (B68, B69) was apparently the first to develop, or at least to popularize, the detection of fluoride ion by its bleaching action on a zirconium lake. This test has won wide acceptance and has resulted in a voluminous literature. The test is based on the fact that zirconium chloride or other soluble zirconium salt reacts with the bright yellow ANALYTICAL CHEMISTRY OF FLUORINE 95 sodium alizarin sulfonate, Ci4H704S03Na, to form a dark reddish-violet lake. The addition of even a minute amount of fluoride ions causes decolorization, i.e., conversion from the dark color to a light yellow due to the formation of colorless complex ZrF6 anions and the liberation of the free dye. Fluorosilicate, boron trifluoride, and other complex com­ pounds of fluorine behave similarly. A popular embodiment of this bleaching test involves the use of strips of test paper impregnated with the zirconium-alizarin complex. To test for the presence of fluoride, a drop of the neutral test solution is added to the moist red spot formed by placing a drop of dilute hydrochloric or acetic acid solution on a piece of the test paper; fluoride turns the spot yellow. In the presence of small amounts of fluoride, the test strip is best heated in steam. The 5limit of sensitivity is about 1 μg. with a concentra­ tion limit of 2 X 10" . The test has been extensively applied by the spot plate technic with a claimed limit of sensitivity of 0.2 Mg. Stone (S124) found that alizarin is a more sensitive reagent than the alizarin sulfonic acid or its sodium salt and is less likely to react with interfering substances. Representative studies and descriptions of the test as applied in solu­ tion, on paper, and on the spot plate include references B71, F9, F10, F12, F13, F14, F17, F21, F45, G28, H22, K29, K56, K57, L28, M38, P21, S124, T47, W24, W25, and W74. Procedures particularly applicable to silicates, rocks, and insoluble fluorides are given in All, F12, F14, F17, and G28. Extraction of insoluble compounds with hydrochloric acid often dissolves enough fluoride for the test. Other metallic ions form violet colorations with the alizarin sulfonic acid, but their color is discharged on acidification with hydrochloric acid. Possible sources of interference are any anions forming insoluble zirco­ nium compounds or stable, soluble zirconium complexes, since their presence will also result in the formation of free dye; e.g., appreciably large amounts of oxalate, phosphate, arsenate, sulfate, and thiosulfate interfere with the test. Oxalate can often be removed by heating the origi­ nal solid sample. Sulfate interference is minimized by adding benzidine hydrochloride to the test solution and then adding a drop of the resulting suspension to the paper. Oxidizing agents such as chlorate, bromate, and iodate oxidize the chloride ion in the acidic solution to free chlorine, which in turn bleaches the reagent. This interference is eliminated by the addi­ tion of sulfite. Interferences are often avoided by applying the test after separation of the fluoride by volatilization as HF, BF3, or SiF4. It is possible to decrease interference still further by ether extraction of the distillate (F9). Prior volatilization can generally be applied to isolating the fluorine 96 PHILIP J. ELVING, CHARLES A. HORTON AND HOBART H. WILLARD

in insoluble samples. The fluoride can often be recovered in a solution suitable for the test by shaking the insoluble compound with dilute hydrochloric acid (P21). Substitutes for the zirconium, the alizarin, or both in forming the colored lake have been frequently proposed, which have been claimed to give superior results and to avoid certain interferences. For example, zirconium ion forms an insoluble brown salt with p-dimethylaminoazo- benzenearsonic acid (F14, F21), whose color is converted to the red of the dye by fluoride ion. There is lessened interference by aluminum, sulfate, and arsenate; phosphate interferes if it exceeds the fluoride present (H22). The equivalent of 0.4 μg. of CaF2 can be detected in the presence of 16 mg. of A1203. Other anthraquinone dyes proposed as substitutes for the alizarin include purpurin (K54), anthrapurpurin, flavopurpurin, anthragallol, rufigallic acid (B69), and 1,2,4,5,8-alizarincyanine (S66). Shvedov (S66) claimed that the last named was the most sensitive fluoride indicator of the hydroxyanthraquinon8 e dyes, with a maximum sensitivity of 0.01 mg. F per liter (10~ ), as well as requiring less time for the bleaching action. Other suggested lakes include thorium-alizarin (B39, K29) and titanium-dihydroxymaleic acid (W24: Vol. II, p. 107).

3. Miscellaneous Color and Fluorescence Tests Several delicate tests for the detection of fluoride ion are based upon its ability to quench the fluorescence of aluminum complexes by the formation of the aluminum fluoride complex ion. Typical procedures involve the reddish-orange fluorescence of aluminum and 2-hydroxy- naphthylazo-2-naphthol-4-sulfonic acid or Chrome-blue Í (B85) in acetate buffered solution; fluorescence of the complex of aluminum and 8-hydroxyquinoline (F16), whic6 h has a sensitivity limit of 0.05 Mg. F and a concentration limit of 10~ ; and fluorescence of the reaction product of aluminum and morin (B57), whose sensitivity is comparable to the preceding test (see also reference G58). Many of the bleaching tests for the determination of fluoride which depend upon the stable complex ions formed by fluoride ion with certain metallic ions can be used for the detection of fluoride, e.g., the fading of the yellow color of titanium in sulfuric acid solution produced on the addition of hydrogen peroxide (K57) and the decolorization of ferric ion in concentrated bromide solution (E28). Details for such tests are often given in the references for the quantitative procedures which are cited in the appropriate sections, e.g., Section V-E-l. Feigl (Fll) has reviewed the interference of fluoride ion in the color tests for vanadium, molyb­ denum, and tungsten. These interference phenomena might well serve ANALYTICAL CHEMISTRY OF FLUORINE 97 as the basis for tests for the detection and colorimetric estimation of fluoride. A recent rapid test (C87) for fluoride is based upon the color change of an acidimétrie indicator due to the reactio3n Al(OH), + 6F- = AlFe~ + 30H~ A drop of sample solution (neutralized to the chosen indicator) is added at the intersection of two paper strips, one impregnated with an indicator, e.g., methyl red, and the other with aluminum acetate. The presence of 0.01 % or more NaF in the drop will cause a noticeable color change. The fluorides of calcium, strontium, and barium can be identified and differentiated by the color changes on the adsorption of pH indicators by the solids (K71). Synthetic and natural CaF2 can be differentiated as can fresh and aged CaF2. An indirect colorimetric test for fluoride uses the molybdenum and benzidine blue tests (F14, F18, F19, G28, L28). The sample is warmed with sulfuric acid and silica sand, and the SiF4 is collected in a drop of water. The latter is tested for the presence of silicic acid by the addition of molybdate to form molybdosilicate, followed by the addition of benzidine to produce a blue color or precipitate. The test has been applied to plants and soils by examining the residue obtained on ignition in an oxygen calorimetric bomb (R15), to minerals and rocks (F12, F20, L28), and to mineral waters (F12). An identification limit of 1 μg. is possible (F12). 4. Precipitation Tests Fluoride forms fairly insoluble salts with calcium, strontium, and lanthanum, slightly soluble salts with lithium, copper, barium, lead, and ferric iron, and slightly soluble fluorosilicates with sodium, potassium, and barium. Most, if not all, of these precipitates as ordinarily obtained are gelatinous. On aging or on coprecipitation with a carrier salt, the physical nature of the precipitate may improve. Further details concern­ ing some of the insoluble calcium salts discussed in this section can be found in Section V-A on precipitation and gravimetric determination. Lanthanum acetate is apparently the best precipitant for the qualita­ tive detection of fluoride; the gelatinous precipitate which first appears slowly becomes crystalline. The precipitate, although not readily soluble in dilute acids, is gradually dissolved by strong mineral acids (F37, G35, M69, T47). The sensitivity of the test is increased by the addition of a dye which is adsorbable by the precipitate. Addition of eosin, which has proved most suitable, to the acetate buffered solution results in a red precipitate with a sensitivity of 1 μg. F (F37). Phosphate, chromate, molybdate, sulfite, and large concentrations of alkali metal salts interfere. 98 PHILIP J. ELVING, CHARLES A. HORTON AND HOBART H. WILLARD

Cerous nitrate also precipitates cerous fluoride from a solution acidi­ fied with acetic acid (S76). Calcium ion forms a white slimy precipitate, insoluble in weak acids such as acetic. It is best precipitated along with calcium carbonate; after ignition and extraction with acetic acid, a more dense residue is obtained which can be used for the etching or hanging drop tests. Barium acetate, on addition to the boiling sample solution to which sulfate ion has been added, forms a precipitate of barium fluoride and sulfate; the latter is then used for the etching test (W79). Barium fluoride is soluble in nitric acid and ammonium salts. The literature on the precipitation of calcium fluoride has been re­ viewed by Mahr (M33). The test can be applied to minerals after car­ bonate fusion and silica removal, by adjustment to pH 9 and addition of CaCl2 (C96). The identity of the precipitate is best verified by the etching test. Calcium nitrate can be used on the solution obtained from the potas­ sium fusion of organic compounds (B17). Other procedures are given in references D79, F42, T17, and W18. As little as 0.04 mg. F in 10 ml. will form a precipitate in acetate buffered solution, heated to 50°, on the addition of a benzidine-mercury succinimide reagent (M72, P30). Oxidizing agents and sulfuric and phos­ phoric acids must be absent. Fluorosilicic acid forms slightly soluble salts with barium and potas­ sium ions (T47). 5. Microscopic Tests The relative paucity of well-defined crystalline fluorides and related compounds has restricted the detection of fluorine by means of micro­ scopic observation. However, a few tests have been described which offer some advantage in identifying not only fluoride ion but certain fluorine- containing compounds. In several cases it has been suggested that the usual tests for the detection of fluoride be confirmed by microscopic observation. For example, it has been advocated that, in testing by the evolution of SiF4, the latter be collected in a drop of water on a micro­ scopic slide. The white ring can be confirmed by the addition of a drop of 0.5% barium acetate solution and observation of the typical crystals of BaSiFe formed on evaporation to dryness at 30 to 40° (M44). A considerable number of tests for the microscopic identification rely on observation of Na2SiFe (C7, G28, G54, G61, P48, T54) or BaSiFe (C7, G28, C31, C33, M44). Campos (C7) found that concentrations of 0.6 to 1.2 p.p.m. fluoride can be identified by the characteristic crystal forms of Na2SiFe (hexagonal prisms), CaSiFe (rods), or CaSiF6 (hexagonal). Other fluorosilicates which can be identified are K2SiF6 (octahedrons), BaDiFe (crossed crystals), HgSiF6 (hexagonal double pyramids), FeSiF6 (hexa- ANALYTICAL CHEMISTRY OF FLUORINE 99 gonal), Ag2SiF«-4H20 (monoclinic prisms), Li2SiF6 (monoclinic prisms), (NH4)2SiFe (hexagonal leaflets), and ZnSiFe (hexagonal prisms) (C7). These tests can be applied in a variety of ways. Absorption of HF vapors in a concentrated solution of sodium silicate causes the formation of easily recognizable Na2SiF6 crystals (G61). The use of a crystalline derivative to confirm the silica ring or deposit test has been indicated. The test can be made quantitative by counting the Na2SiFe crystals (P48). The specific applications to many materials such as bones (C33) and mineral waters (C21, P48) have been discussed (see Section V). Mercurous nitrate has been advocated as a reagent for producing characteristic crystals with fluoride which are suitable for identification (D42, R54). Benzidine acetate forms well-defined microcrystalline rosettes with fluoride (C7). A ferrous bipyridyl complex has been investigated as a fluoride reagent (P49). Hynes and Yanowski (H81, H82) obtained characteristic crystalline products suitable for identification by the reactions of hydrogen fluoride (bifluoride) ion with xanthocobaltic chloride (nitropentamminocobaltic chloride) (H81), or 1,5-carbonatotetramminocobaltic nitrate (H82). These are useful in the presence of other anions. The reaction of some 80 anions, including fluoride, bifluoride, and fluorosilicate, with the various complex cobalt ammines are summarized by these authors (Y3). Phosphoryl fluoride, POF3, is converted in cold water to POF2OH, which gives a distinguishing crystalline salt with nitron acetate (LI4). Fluoroborate, BF4~, also forms a nitron salt (L13). Mason and DeLa- Mater (M50) have discussed micro reactions suitable for the identifica­ tion of fluorophosphates (M50), e.g., reaction of dialkyl fluorophosphates with thiourea. They also described the microscopical properties of 0- fluoroethyl-3,5-dinitrobenzoate, 0-fluoroethyl-N-(a-naphthyl)-carbamate and fluoroacetophenylhydrazide (M49). Crystallographically, the simple fluorosulfonates resemble the perchlorates (L14). Fluoroborate reacts with potassium permanganate to form mixed colored orthorhombic crystals of KBF4—KMn04; perchlorates and sul­ fates give similar crystals (B79). Booth and Martin (B79) consider the best test for fluoroborate to be the nitron precipitate mentioned in the previous paragraph. The nitron fluoroborate forms light green needles (m.p. 225°). Similar nitron precipitates are given by nitrate, fluorophos- phate, perchlorate, perrhenate, and tungstate (L13).

6. Miscellaneous Tests

Only a few tests which seem to be generally overlooked or which merit further consideration are included in this section. 100 PHILIP J. ELVING, CHARLES A. HORTON AND HOBART H. WILLARD

Fluoride can be detected by the gradually vanishing yellow-green color produced by heating in a colorless flame a mixture of the sample, borax, and potassium hydrogen sulfate, fused to a bead in a platinum loop (Ml). A transient green flame coloration due to BF3 is obtained when a powdered sample containing fluorine is heating with 3 parts of a flux consisting of 2 parts potassium hydrogen sulfate and 1 part boro­ silicate (powdered Pyrex glass) (W61). Gatterer (G12) has described the spectroscopic detection of fluorine by excitation of the sample in an evacuated tube placed in a high fre­ quency field; intense spectral emission is obtained in the region of 4500 to 7000 A. The sensitivity limit on a sample of 10 to 20 mg. is 0.001%. Fluorine can also be detected by the band spectra produced in the pres­ ence of alkaline earth metals (see section on emission spectroscopy and reference P3). Apparently fluoride ion forms a volatile chromium compound under the conditions of the standard chromyl chloride test for chloride ion, and this test can also be used to identify fluorine (M46). A possible basis for a fluorine test is the observation of Steinherz (SI 17) that on increasing anodic polarization of tin in solutions of hydro­ chloric, hydrobromic, hydroiodic, fluorosilicic, sulfuric, nitric, iodic, perchloric, and acetic acids, the current initially increases according to Ohm's law, reaches a maximum at a potential characteristic of the acid, and then decreases.

B. FLUORINE IN ORGANIC COMPOUNDS Many of the technics used for converting fluorine bound to carbon to fluoride ion for quantitative determination (Sections III-B and VI-D) have been used to isolate fluoride ion preliminary to detection. Thus, fusion with sodium metal of even such stable compounds as those con­ taining trifluoromethyl groups gives a fluoride solution in which fluoride is detected by precipitation with cerous ion (S76). Fusions with metallic potassium followed by the zirconium alizarin test is quite sensitive (F14). Burning in hydrogen or decomposition over hot platinum gives HF detect­ able by the bleaching of zirconium or thorium lakes (test-paper strips) (D68, K29). The peroxide bomb is suitable to prepare a solution for the thorium-alizarin test (S129). In compounds boiling above 60°, as little as 1% fluorine in a 1-mg. sample can be detected as HF after wet combustion with a mixture of iodic, chromic, sulfuric, and phosphoric acids (B39). On heating a sample containing a few tenths of a milligram of fluorine with chromate and sulfuric acid in the presence of glass, the sulfuric acid will no longer wet the glass. The silicic acid turbidity will be observed on suspending above the solution a drop of water on the end of a rod (B63, D51). ANALYTICAL CHEMISTRY OF FLUORINE 101

The Beilstein copper flame coloration for halogen is generally not very satisfactory for fluorine compounds (B63, Ă14). Fluorine-containing compounds in air can be collected on silica gel in tubes, then desorbed by heat, and pyrolyzed by passage over hot plati­ nized silica gel; the HF formed is detected by passage over zirconium- or thorium-alizarin paper (W74). Fluorocarbons in air as low as 1 p.p.m. are detectable by pyrolysis in a platinum tube at 950° and detection of the fluoride ion (K29). Methods for the detection of fluorine-containing compounds in gases are further discussed in the next subsection. Goldman, Flannery, Arent, Hoag, and Buswell (G53) have described the application of a large variety of decomposition and detection methods for the detection of fluorine-containing organic compounds in water. A considerable number of studies made for American and British govern­ mental agencies are summarized and cited. Sanchis (S8) has discussed the critical factors in a Chrome Azurol-thorium spot test for detecting organic fluorine compounds in water, using potassium persulfate as an oxidizing agent for the recovery of fluoride ion. Little is generally available on the characterization of organic com­ pounds containing fluorine, owing in large part to their relative scarcity until recent years. The chemical inertness of the carbon-fluorine bond has always been a deterrent to organic qualitative analysis. Some avail­ able methods are scattered through the literature on the preparation and properties of organic fluorine-containing compounds, e.g., the use of fluoroacetamide to characterize methyl fluoroacetate (M6). The chemical inertness of the carbon-fluorine bond, e.g., resistance to fission by sodium hydroxide, could be used to differentiate fluorine-containing compounds from those containing other halogens. The peculiar physical properties of fluorine-containing organic compounds, e.g., low refractive index, should be useful in distinguishing and characterizing such compounds. In the subsection on microscopic identification, the feasibility of the use of crystal properties in the qualitative analysis of organic fluorine-con­ taining compounds is indicated. Ramsey and Patterson (Rll) developed a qualitative test for mono- fluoroacetic acid which depends on formation of red thioindigo: FCHoCOOH + Ο SH S SC

COOH C C S II ο Only α-halogen acetic acid types react, and 20 Mg. gives a positive test. 102 PHILIP J. ELVING, CHARLES A. HORTON AND HOBART H. WILLARD

C. FLUORINE IN GASEOUS SAMPLES A number of methods have been proposed for the detection of fluorine in gaseous samples of specific types which are of general applicability. Many of these, as in those cited in the previous subsection for detecting organic compounds in air, depend upon conversion to HF and detection of the latter by the bleaching of a thorium or zirconium lake. A patent (M38) has been issued for the detection of hydrofluoric acid vapor in air by passing the latter through a solid porous medium impreg­ nated with a zirconium-alizarin lake (color change from light pink to bright yellow). As little as 0.025% CF4 in hydrogen can be detected by burning the gas and observing the etching of glass by the flame due to the presence of HF (T52). A patent (H49) covers the detection of volatile, organic fluorine-containing compounds such as the Freons in the gas phase by passage of the sample over incandescent silica to form SiF4, and then into contact with ammonia to form a white cloud. Another patent (H79) covers the detection of fluorine-containing gases by flame coloration. Boron trifluoride also forms a white solid, NH3BF3, with ammonia (B79). As little as 0.2% SiF4 in boron trifluoride can be detected as a residual gas by contact with acetyl fluoride, which reacts with the boron trifluoride to form white solid fluoroborate (F43). It is obvious that any of the infrared spectrophotometric and mass spectrometric technics used in gas analysis could be applied to the detec­ tion and identification of fluorine compounds in the gas phase. References to the available infrared data are given in another section of this chapter. Mass spectrometer cracking patterns could be used similarly. V. Determination of Fluoride Ion A. PRECIPITATION AND GRAVIMETRIC DETERMINATION The various precipitation and gravimetric methods for the determina­ tion of fluoride described in this section have in some cases been adapted for direct or indirect volumetric titration methods, which are discussed in a subsequent section. The discussion is divided on the basis of the compound precipitated. 1. As Lead Chlorofluoride The gravimetric determination of lead chlorofluoride, PbCIF, was first suggested by Berzelius (B49). Although the solubility of lead chloro­ fluoride in water at 25° is 33.8 (H32) to 37.0 (S114) mg. per 100 ml., the solubility is much less in the presence of lead or chloride ions (A3), de­ creasing with increasing concentration of these common ions. For exam- ANALYTICAL CHEMISTRY OF FLUORINE 103 pie, the solubility of lead chlorofluoride in 0.02 and 0.04 Ν lead chloride is 0.8 and 0.5 mg. per 100 ml., respectively, at 25°, and less at lower tem­ peratures. Tananae9 v (T8) gives a solubility product for lead chlorofluoride as 2.3 X 10~ , from which the fluoride unprecipitated in a7 solution con­ taining 0.1 M excess chloride (or lead) will be 2.3 X 10~~ M ; Haul and Griess (H32) reported a value of similar magnitude. Starck (Si 14) precipitated lead chlorofluoride from a sample solution neutralized to phenolphthalein, using saturated lead chloride as precipi­ tant. The precipitate was washed with lead chloride and then water, and dried at 140 to 150°. The deviation on pure sodium fluoride ran from +0.46 to -0.30% of the fluoride taken in 17 determinations (S114). Hammond (H15) precipitated lead chlorofluoride from a solution be­ tween pH 4.6 to 5.3 adjusted using bromcresol purple; the other condi­ tions were those specified by Starck (SI 14). Hawley (H41) precipitated lead chlorofluoride from an acetic acid solution using lead acetate as precipitant, and a wash solution of water saturated with lead chlorofluo­ ride, finishing the determination volumetrically as discussed elsewhere. Fischer and Peisker (F38) used dilute lead chloride or lead chloroni- trate solution to precipitate lead chlorofluoride in a solution made just acidic to methyl orange with nitric acid. After standing, most of the super- nate was decanted through the filter, and the precipitate was treated with 4 ml. of water and washed onto the filter with a minimum of water saturated with lead chlorofluoride. The precipitate was dried 30 minutes at 150°. Winkler (W64) precipitated lead chlorofluoride at pH 4.0, allowing the precipitate to settle overnight before filtration, then washing once with water, twice with a PbCIF saturated solution, once more with water, and then drying. Donovan (D67) recommended cooling the solu­ tion in ice after precipitation and before filtration of lead chlorofluoride in order to reduce the solubility of the compound. Ergen and Heath (E27) recommended the use of medium porosity fritted-glass, Gooch- type crucibles for the filtration. Treadwell and Hall (T48) note that for the precipitation of lead chlorofluoride between pH 3.5 to 5.6, the presence of more than 0.5 mg. aluminum, 50 mg. boron, 0.5 g. ammonium ion, or 10 g. alkali metal ions causes low results. In the determination of fluoride in organo fluorophos- phate (K24), it was found that ammonium and carbonate ions should be removed and acetic acid rather than nitric acid used for neutralization; Hammond's (H15) procedure was preferred to that of Hoffman and Lundell (H66) among the volumetric methods. Kaufman (K13), using the precipitation part of the Hoffman-Lundell (H66) procedure, obtained quantitative results for about 20 mg. F between pH 4.40 and 4.75, lower results to pH 5, and higher above pH 5. Results were erratic for 10 mg. F 104 PHILIP J. ELVING, CHARLES A. HORTON AND HOBART H. WILLARD and less. Other studies on similar fluorine compounds (C43) indicated solutions should be heated at 80° for 0.5 hour after addition of reagents for lead chlorofluoride precipitation and cooled in ice for 10 minutes before filtration. Presence of over 5% alcohol from decomposition of the organic compound caused coprecipitation of lead chloride. Delgery (D33) in conductometric and thermal studies of mixed lead halide salts found evidence for 4PbF2-PbCl2, 4PbF2PbBr2, and 4PbF2- Pbl2, in addition to the better-known compounds PbCIF, PbBrF, and PblF. Many workers (F36, E17, G29, H17,16 P23, R51, R64, S31, S77, S78, S102, V2, V6, W16) have used methods depending on formation of lead chlorofluoride, followed by either gravimetric or titrimetric measurement of the chloride in the precipitate or remaining in solution when a known amount of chloride is added. Elving and Ligett (E21) suggested use of methods based on this principle for determination of the larger amounts of fluoride obtained in decomposition of organic fluorine compounds. In recent work (S16) the volumetric methods depending on formation of lead chlorofluoride originally suggested by Kapfenberger (K9), Hawley (H41), and Hoffman and Lundell (H66), a gravimetric modification of the latter, and other modifications involving formation of lead chloro­ fluoride were studied. It was found that only under controlled conditions will the precipitate of "lead chlorofluoride'' have the known composition, and that Kapfenberger's (K9) volumetric procedure was the best. For the various methods the composition of the precipitate, in molar ratios, ran from Pb/Cl = 0.912 ± 0.010 to 1.203 ± 0.003, and Pb/F = 0.948 ± 0.010 to 1.010 ± 0.008. In each method tried the precision was better than the accuracy. All results indicated that neither the gravimetric nor volumetric methods for fluoride, depending on formation of lead chlorofluoride, can ever give results better than 5 parts per thousand. By use of the volu­ metric method, discussed in another section, some of the interferences of the gravimetric method are reduced.

2. As Calcium Fluoride The first gravimetric method for the determination of fluoride was the precipitation of calcium fluoride from aqueous solution proposed by Berzelius (B49). The precipitate is generally very gelatinous, difficult to filter, apt to become colloidal (K7, M59, M60), and, as indicated subse­ quently, not completely insoluble in water. In other early work Heintz (H44) decreased these difficulties by precipitating calcium phosphate along with the fluoride and obtaining the weight of calcium fluoride by difference after weighing the phosphate. Rose (R52, R53, R59) improved the filterability of calcium fluoride by precipitation with calcium car- ANALYTICAL CHEMISTRY OF FLUORINE 105 bonate. The ignited mixed precipitate was treated with dilute acetic acid, which dissolved the calcium carbonate, yet left the calcium fluoride in a dense form which filtered satisfactorily. Others used similar methods (C46, C67, D55, G3, Gil, G60, J6, K59, P29, R16, W20). These modifi­ cations are subsequently discussed. As early as 1893 Kohlrausch (K48) showed that the solubility of cal­ cium fluoride is 1.6 mg. per 100 g. of water at 18°; Carter (C25) gave a value of 4 mg. per 100 ml. of saturated solution; Mougnaud (M120), 1.83 mg. per 100 ml. of water at 18°. Others (D76) found a solubility of 15.3, 17.5, and 19.2 mg. of calcium fluoride per milliliter in 0.5, 1.0, and 2.0 Ν acetic acid at 40°. The use of acetic acid to dissolve carrier agents in the precipitation of calcium fluoride must be considered in the light of such data. Moreover, Meyer and Schultz (M69) indicated the calcium fluoride may adsorb acetate giving results 1 to 2% high. Adolph (A3) found that in drying and ashing calcium fluoride precipitated \vith calcium carbonate there was a loss of 1.5 mg. of the fluoride for each 10 ml. of 1.5 Ν acetic acid used to remove the calcium oxide. Mazzucchelli (M59, M60) and Kandilarov (K7) studied the colloidal properties of calcium fluoride. Kandilarov (K6) assured retention of any colloidal calcium fluoride in his gravimetric method by filtration through a membrane filter. He ignited the precipitate at 500° after drying in a vacuum desiccator. Mikhailova (M71) added some 10% gelatin solution after precipitation of calcium fluoride to help coagulate it for filtration. Geyer (G34) was able to obtain a more dense and filterable calcium fluoride by slow addition of calcium chloride to a hot solution buffered with acetate. The good results obtained were due to a compensation of errors, as the fluoride was not completely insoluble, but the precipitate adsorbed calcium to yield the proper weight of precipitate. , Dupuis and Duval (D77) reported on the temperatures to which various fluorides, including calcium fluoride, should be heated or ignited for constant weight without decomposition. Their results are summarized in Table III. Carrière and coworkers (C22, C23, C24), and Hart (H23) obtained better results ni precipitation fo calcium fluoride with carbonate ni na ammoniacal solution than ni a dilute acetic acid solution. The precipitate contains some calcium carbonate which compensates for the loss due ot the solubility fo calcium fluoride (M118, M119). Karasinski (Kll), no the other hand, reported better results for precipitation from a dilute acetic acid solution than from na alkaline carbonate solution. Mougnaud (M120), who centrifuged the calcium fluoride precipitate and washed ti with water saturated with the same compound, obtained equally good results for precipitation from acidic, neutral, ro alkaline solution. Bonis 106 PHILIP J. ELVING, CHARLES A. HORTON AND HOBART H. WILLARD

TABLE III Heating Temperatures for Precipitates Obtained in Gravimetric Determination of Fluoride (D77)*

Compound Compound Heating temp. Precipitant precipitated weighed limits, °C

CaCl2 CaF2 CaF2 400-950 La(N03)3 LaF3La203 475-946 Bi(N03)3 BiF3 BiF3 50-93 Th(N03)4 ThF4xH20 Th(OH)4 242-475 PbClo PbCIF PbCIF 66-538 U(S04),> UF4 U308 812-945 KC1 K2SiF6 K2SiF6 60-410 BaCl 2 BaSiFe BaSiF6 100-345 BaCl2 BaSiF6 BaF> 542-946 BaCl2 BaF2 BaF2 542-946 Triphenyltin chloride Triphenyltin fluoride <158

* These authors were unable to filter thorium fluorosilicate successfully.

(B76) also precipitated fluoride as calcium fluoride with calcium car­ bonate but dissolved out the latter with dilute acetic acid before drying. Shuey (S60) found that precipitation of fluoride as calcium fluoride in the presence of carbonate gave low results owing to solubility of calcium fluoride in water and dilute acetic acid. Lisitsyn and Volkov (L40) analyzed calcium fluoride samples by dissolution in boiling aluminum chloride solution, addition of a four- to sixfold excess of acetic acid to complex the aluminum, and then precipita­ tion of calcium fluoride after neutralization with ammonia to the methyl orange end point. A correction was applied for the solubility of calcium fluoride in the acetic acid. Krause (K65) precipitated calcium fluoride from solution, using calcium hydroxide and dissolving most of the excess hydroxide with acetic acid. After ignition of the impure calcium fluoride precipitate, the precipitate \vas suspended in water and titrated with standard hydrochloric acid in order to correct the weight for the calcium oxide present. Deussen and coworkers (D47, D48, D49) fused their samples with calcium oxide in a platinum crucible, fixing the fluoride. The excess lime was slaked with water and dissolved in dilute acetic acid; after aging, the calcium fluoride was filtered off from a 10% alcoholic solution. Results were generally 0.8 % low. Starck and Thorin (SI 15) precipitated calcium fluoride with a known amount of oxalate from a solution slightly acidified with acetic acid. They subtracted the known weight of calcium oxalate from the total ANALYTICAL CHEMISTRY OF FLUORINE 107 weight of dried precipitate to calculate the fluoride. Scott (S42) suggested addition of some potassium hydroxide to help coagulate the precipitate of calcium oxalate and fluoride. Dubiel (D71) also used precipitation of calcium fluoride in the presence of oxalate and later converted the precipi­ tate to calcium sulfate for weighing. Presence of other sulfates would give erroneous results (B92). Dinwiddie (D52) precipitated calcium fluoride with calcium sulfate, washing the precipitate with water saturated with these salts. The mixed precipitate was dried at 200° and weighed. After treatment with sulfuric acid, evaporation to dryness, and heating at 200° again, the change in weight can be calculated as fluoride. Results on sodium fluoride analyses were satisfactory. Calcium fluoride has also been precipitated with cal­ cium sulfate in the presence of gelatin to obtain a better precipitate (M71, W44). The weighed mixed precipitate was heated with boric and perchloric acids to volatilize off the fluoride as boron trifluoride, as sug­ gested by Schwerin (S36). The residue was fumed and cooled, perchlorates removed using 50% methyl alcohol, the residual calcium sulfate weighed, and the fluoride calculated by difference. Ryss (R76) analyzed fluoroborates gravimetrically as calcium fluoride. The boric acid liberated in the reaction was neutralized by addition of potassium chlorate and iodide with potassium vanadate as catalyst, so that the calcium fluoride would precipitate in a filterable form. Wamser (W10) used a similar method. Ruff and coworkers (R61, R62, R63, R64, R65), as well as many others (B39, G63, M68, P23, P39, S57, T52, V8, W30), have determined fluorine in organic compounds gravimetrically as calcium fluoride. On fairly large amounts of fluoride, results are usually quite satisfactory, n^ainly because of compensation of errors. However, Whearty (W30) and some others have reported low results. Vaughn and Nieuwland (V8) corrected their results for solubility in the wash water used. Gautier (G 15) noted that calcium fluoride may be converted to another salt with loss of fluoride as hydrogen fluoride on evaporation in acidic solutions. Roper and Prideaux (R46) recommended alkalimetric titration rather than gravimetric determination as calcium fluoride for determination of fluoride in commercial bifluorides. Travers (T44) indi­ cated that the formation of a complex with aluminum by fluoride may cause low results in the gravimetric determination of fluoride as calcium fluoride. The presence of other ions, such as iron (III), beryllium, zir­ conium, or boron, which form complexes with fluoride ion, may also cause low results, although fluoroborates have been analyzed by the gravimetric calcium fluoride method (A24). Although fluoride is still occasionally determined gravimetrically as 108 PHILIP J. ELVING, CHARLES A. HORTON AND HOBART H. WILLARD calcium fluoride, the gravimetric determination as lead chlorofluoride, discussed earlier, is superior and is more widely accepted. For those who wish to determine fluoride gravimetrically as calcium fluoride, Treadwell and Hall (T48) have a suitable procedure. Workers not mentioned in the present section who have precipitated calcium fluoride include those listed in references B72, C80, G17, K37, K41, L27, L46, P57, S55, T38, and W16. In 1933, Mahr (M33) reviewed the method. It must be remem­ bered that phosphate, arsenate, tungstate, molybdate, chromate, vana­ date, antimony, titanium, zirconium, and aluminum ions must be re­ moved (H66) before precipitation of calcium fluoride.

3. As Rare Earth Metal and Other Metal Fluorides Meyer and Schultz (M69) suggested the detection and determination of fluoride by precipitation as lanthanum fluoride from an acetate buffered solution. After drying at 110°, the composition of the precipitate was LaF3 + La(OAc)3; after ignition, LaF3La203. The fluoride content was calculated from the difference in the two weights. Fischer (F37) had trouble using the above method quantitatively because of the poor precipitation qualities of the lanthanum fluoride and its tendency to adsorb other ions. Fischer (F37) and others (G32, G33, M58) used pre­ cipitation as lanthanum fluoride for a qualitative test for fluoride. Giammarino (G35) used a procedure similar to that of Meyer and Schultz (M69) with satisfactory results down to 5 mg. F per milliliter. For samples containing 0.03 to 0.5% fluoride, he (G35) recommended a nephelometric method depending on formation of colloidal lanthanum fluoride. None of the other rare earth metals have been used for the precipita­ tion of fluoride in gravimetric determinations, although precipitation of yttrium fluoride is the basis of a volumetric method (B32, F57) for fluoride. Gabriel (G3) suggested gravimetric determination of fluoride as thorium fluoride, and a method was soon developed by Deladrier (D32). Adolph (A3), who studied this technic, obtained erratic results which he attributed to formation of Na2ThF6. Pisani (P44) indicated that the original precipitate is thorium fluoride tetrahydrate, which converts first to the anhydrous thorium fluoride and then to thorium dioxide on heating. Gooch and Kobayashi (G55) precipitated thorium fluoride tetrahydrate from a solution 0.02 to 0.2 Ν in acetic acid and weighed the precipitate as thorium dioxide. They (G55) also proposed an indirect volumetric method based on precipitation of fluoride as thorium fluoride. Wadhani (W3) has recently restudied the precipitation of thorium fluoride. Domange (D60, D61) recommended precipitation of fluoride as ANALYTICAL CHEMISTRY OF FLUORINE 109 bismuth (III) fluoride, since it filters well, is less soluble than calcium fluoride, and has a high molecular weight. However, the other halide ions, phosphate, and sulfate also form slightly soluble salts with bismuth in dilute acetic acid solutions and thus interfere. The method has not been used extensively. Balavoine (B12) and Karaoglanov (K10) used precipitation of barium fluoride for the qualitative detection of fluoride. Blum and Vaubel (B62) precipitated barium sulfate and fluoride together from an acetic acid solution. After ignition and weighing of the mixed precipitate, the fluoride was volatilized by heating with sulfuric acid, the residue ignited and weighed. Fluoride was calculated from the change in weight. Karaoglanov (K10) and Gautier and Clausmann (G17, G18) also used precipitation as calcium or lead fluoride for qualitative tests. Zarin and Dubnikov (Z3) considered determination gravimetrically as calcium, lead, and strontium fluoride, but adopted instead an alkalimetric titration depending on precipitation of barium fluorosilicate. The heat stability of the fluorides used in gravimetric methods was studied by Dupuis and Duval (D77), whose results have been summarized in Table III. 4. Miscellaneous Gravimetric Methods Krause and Becker (K64) prepared various aryl tin fluorides, including diphenyltin difluoride, m.p. over 360°; tri-p-toluyltin fluoride, m.p. 305°; tri-ra-toluyltin fluoride, m.p. 205°; tri-p-cresyltin fluoride, m.p. 242.5°; and triphenyltin fluoride, m.p. 357°. The latter is sufficiently insoluble in cold alcohol, ether, or water that they suggested the quantitative precipi­ tation of fluoride as this aryltin fluoride. Allen and Furman (A13) de­ scribed a procedure for quantitative precipitation of triphenyltin fluoride using triphenyltin chloride as the precipitant. They added 95% ethanol to an aqueous fluoride solution to give a final alcohol concentration of 60 to 70% by volume, heated to boiling, and added, with rapid stirring, twice the calculated amount of triphenyltin chloride in hot alcohol. The solution was left overnight and was filtered through a glass frit after cooling in ice. The precipitate was washed with 95% alcohol saturated with triphenyltin fluoride and dried at 110° for 30 minutes. For 95 Mg. F taken, the average error was ±7.6 Mg.; for 48 Mg., ±8.1 Mg- The original fluoride solution should be at pH 7 to 9 for correct results. The other halide ions, nitrate, and sulfate do not interfere; however, carbonate, silicate, and phosphate must be previously removed. The precipitate filters and washes easily and need not be ignited ; the gravimetric factor is favorable. Disadvantages are unavailability and expense of the precipitat­ ing reagent and its rather low solubility, limitation to less than 40 mg. 110 PHILIP J. ELVING, CHARLES A. HORTON AND HOBART H. WILLARD fluoride, and possible entrainment of reagent in the precipitate. Although the method has much to recommend it, the procedure has not been used extensively. Pertusi (P30) and Miller (M72) precipitate fluoride as the complex, dibenzidine-tetrahydrofluoride-mercuric fluoride, i.e., (benzidine)2-(HF)4- HgF2. This complex fluoride is precipitated by the addition of benzidine in dilute acetic acid and 0.02 Í mercuric succinimide to a solution con­ taining fluoride made barely acidic with acetic acid and heated to 50°. The precipitate is washed with cold water, dried over concentrated sul­ furic acid, and weighed. Oxidizing agents, sulfate, and phosphate inter­ fere (W24). The weight loss of a glass vessel due to reaction of fluoride with the glass to evolve silicon tetrafluoride was one of the first methods (C20, C21, L42, P60, S43, S123, T6) to be used for the gravimetric de­ termination of fluoride and even has been used fairly recently (P57) for an approximation of the fluoride content of fluorspar. Other methods depend­ ing on the attack of glass by fluoride in acidic solution are discussed in another section of the chapter. Several early workers (F59, P27, R40), after evolution of fluoride as silicon tetrafluoride, collected the gas on powdered pumice and calculated fluoride from the gain in weight. Others (S54, T28) adsorbed it in sodium fluoride, forming sodium silicofluoride, and weighed the resultant mix­ ture. The evolved silicon tetrafluoride has also been hydrolyzed in water to form silica, which was weighed (B34, Dl, L21). A few early workers (Cil, C12, C13, Sill, W22, W28, W62, W75) precipitated fluoride as potassium silicofluoride, usually from a 50% alcoholic solution. Bodenstein and Jockusch (B66) determined elemental fluorine gravi­ metrically by the increase in weight of silver packed in a copper tube. Fluoroborate has been determined gravimetrically by precipitation of nitron fluoroborate (B79, W10). Nitrate, perchlorate, perrhenate, tung- state, and fluorophosphate interfere. Fluoroborate also has been precipi­ tated as a nickel hexammine difluoroborate, (Ni(NH3)e)(BF4)2 (B79). Perchlorate also precipitates with this complex nickel ion. White (W32) has devised gravimetric methods for difluorophosphoric and hexafluorophosphoric acids using nitron as the precipitant; and for difluorophosphoric acid as silver difluorophosphate from 80% alcoholic solution. B. TITRIMETRIC METHODS: PRECIPITATION AND COMPLEXATION 1. Thorium Titration Thorium nitrate is the most commonly used titrant for the volumetric determination of both very small and larger amounts of fluoride. A visual ANALYTICAL CHEMISTRY OF FLUORINE 111 indicator which changes color in the presence of excess thorium ion due to complex or lake formation is generally used to detect the equivalence point. The instrumental methods using this titrant are discussed in other sections. The method depends on the fact that thorium fluoride is very insoluble; thus, thorium is effectively removed until all the fluoride has reacted. Willard and Winter (W52) first used thorium nitrate as a titrant with a zirconium-alizarinsulfonate lake as the indicator. This indicator is yellow in slightly acidic solution when a trace, or more, of fluoride is present to complex the zirconium, thus liberating the free dye. The indicator turns to the pink or magenta color of the zirconium lake with the first excess of thorium ion. The initial pH was adjusted to the yellow form of this indi­ cator (pH 3.5 to 4.0), and the solution contained 50% of alcohol, in which thorium fluoride is completely insoluble. Coupled with the method of these authors (W52) for separation of fluoride as fluorosilicic acid from numerous interfering ions, it is the best procedure for the determination of fluoride according to Rinck (R37) and others. Since this titrant has been used very widely, the various factors involved in the use of this method will be discussed in some detail, although it will be impossible to cite all the slight variations employed. For additional discussion the reader is referred to Rinck's excellent review paper (R37). Armstrong (A34) and others soon found that sodium alizarinsulfonate alone, commonly called Alizarin Red S, is suitable as an indicator. In this case, the first slight excess of thorium changes the yellow acidic color of this dihydroxyanthraquinone to a pink or magneta color due to complex or lake formation with thorium. Effect of pH. The hydrogen ion activity is one of the most important factors which require control in this method, although there are consider­ able differences of opinion as to what is the optimum pH. It has been definitely established, however, that the pH must be controlled within narrow limits. Naturally, the hydrolytic tendencies of thorium limit con­ sideration to acidic solutions. Both the type of medium used, aqueous or semialcoholic, and the ionic strength affect the actual hydrogen ion ac­ tivity, which may not be truly indicated by indicators or a pH meter. The acidity also affects the dissociation of the fluorosilicate ion and the HF molecule as shown in equations 1 and 2. 2 SiF6~ + 40H- ;=± 6F- + H2Si03 + 2H20 (1) HF ^ F" + H+ (2)

The sharpness of the indicator color change at the end point varies with the pH of the medium, the indicator used, and alcohol content. The 112 PHILIP J. ELVING, CHARLES A. HORTON AND HOBART H. WILLARD change in acidity due to addition of the thorium solution, which is always somewhat acidic, must be also taken into account. Hoskins and Ferris (H73) favored pH 3.5, established and maintained by use of a monochloroacetate buffer, for titration with Alizarin Red S indicator in 50% ethanol. Rowley and Churchill (R57) and Armstrong (A33) favored pH 3.0 for the same buffer and indicator in aqueous solu­ tion. Milton, Liddel, and Chivers (M85, M88) also used the same buffer at pH 3.0 in aqueous solution wit,h Chrome , Azurol S or Sollochrome Brilliant Blue BS indicator (3 -sulfo-2',6 -dichlorohydroxydimethyl- fuchsondicarboxylic acid). Nichols and Kindt (N5, N6) used the mono­ chloroacetate buffer at pH 3.15 in aqueous medium containing 0.015% starch with Alizarin Red S as indicator for titrations in which the equiva­ lence point was detected using a photometer. Many others (B18, B58, B122, G33, G72, H70, H79, Jll, K22, K46, M26, P4, S71, S120) have used a similar monochloroacetate buffer. Dahle, Bonnar, and Wichmann (Dll) and Allen (see Dll) prefer to use a pH of 2.75 established by addition of sufficient acid to give a total acidity equal to 2.0 ml. of 0.05 ËĂ HC1 per 40 ml. The original acidity is determined on a separate aliquot using the same Alizarin Red S indicator. This method of pH control has been adopted by the A.O.A.C. (A37, L33) and many others. Use of this acidity control eliminates the salt errors, discussed later, caused by the use of a buffer. Carter (C27) found that solutions buffered at pH 3 gave results 20 % lower than at pH 1.6, which he recommended. Above pH 3.4 he obtained lower, and below pH 1.4 considerably higher, results in titrations with thorium. Matuszak and Brown (M52) and Kazarinova-Oknino (K15) recom­ mended a pH of 3.3 established by addition of 1 ml. of 0.4 Ν acetic acid to 50 ml. of neutralized sample for less than 1 mg. F, and a lower pH for larger amounts. To decrease the pH towards this end, they titrated with a solution of 0.05 Ν thorium in 1.2 Ν acetic acid and claimed propor­ tionality between the titrant and fluoride content up to 50 mg. per 50 ml., but Rinck (R37) obtained proportionality only up to 1 mg. per 50 ml., using similar conditions. Williams (W56) recommended pH 2.7, obtained by neutralizing an aliquot to the 2,5-dinitrophenol end point, and adding to the aliquot sodium chloride and 2.0 ml. of a solution containing 0.02% Alizarin Red S plus thorium equivalent to 2.5 Mg. F, about 0.07 Ν in HC1. The mixture was titrated with thorium nitrate in 0.0072 N HC1. The use of thorium in the indicator eliminated the "indicator blank" and allowed use of twice as much indicator as in other methods; the introduction of sodium chloride helped control the effect of other ions and give a constant ionic strength. Use of the slightly acidic thorium solution kept the pH ANALYTICAL CHEMISTRY OF FLUORINE 113 constant during the titration. Smith and Gardner (S85, S86, VI1) used a modified form of the same method and obtained percentage deviations less than half of those using the A.O.A.C. back-titration procedure (L33). Rinck (R37) recommended pH 3.6 to 3.9 for titration in 50% alcohol, used for less than 150 Mg. F, and pH 3.25 to 3.45 for titration in aqueous solution, used for larger amounts of fluoride. Stevens (S120) recommended titration at either pH 2.8 or 3.3 for water samples, using monochloroacetate or acetate buffers, respectively. Rickson (R36) recently recommended gallocyanine as an adsorption type of indicator, for which optimum results were obtained at pH 5.3. Dissociation of thorium fluoride was decreased by titration in 50% alcohol. Willard and Horton (H72, W47) recommended various pH ranges for the five best visual bicolor and two best fluorescent indicators which they found in a comprehensive study of indicators for this titration. A sum­ mary is given in Table IV of their findings for titration of less than 0.6 mg. F. In general, it should be considered that while chloroacetic acid (pKa = 2.86) serves best as basis for a buffer system in the pH range of 1.9 to 3.9, acetic acid {pKa = 4.76) serves similarly in the pH range of 3.8 to 5.8. Standardization of the Thorium Solution. Hammond and Maclntire (H16) compared the fluoride found as calculated from the thorium normality (established by intercomparison with ammonium oxalate) to

TABLE IV Optima Indicator Concentrations, Buffer pH, and Percentages of Alcohol for Visual Titration of 0.5 to δ mg. of Fluoride with Thorium Nitrate {H72, W47)

Average Indicator concn. % Alcohol Buffer compound mg./50 ml. recommended pH*

Bicolor Purpurin sulfonate 0.15 0 3.3 Alizarin Red S 0.20 0-50 f 3.0 Eriochromcyanin R 0.70 0 4.2 (or 3.6) 2,3-Dicyanoquinizarin 0.60 10-20 3.8 Chrome Azurol S 0.20 0 3.1 Monocolor, fluorescent Quercetin 0.50 50 3.5 Morin 0.4-1.0 50 3.5-3.8

* pH of buffer as measured in aqueous solution, t 50 % alcohol recommended for microgram range. 114 PHILIP J. ELVING, CHARLES A. HORTON AND HOBART H. WILLARD the fluoride added for titrations both,in aqueous and 48% alcoholic media, adjusted in both cases to pH 3.0 ± 0.2 by neutralization and either addi­ tion of a fixed amount of hydrochloric acid or monochloroacetate buffer for both microgram and milligram amounts of fluoride. In the milligram range both methods of adjusting acidity and both media gave results which were fairly concordant with the known amounts added; however, the mean deviation was least for the alcoholic solution adjusted in pH with the buffer. In the microgram range with 48% alcohol, better pre­ cision and accuracy were achieved by acidity adjustment using hydro­ chloric acid; with aqueous media, results were 50% of the value too high using hydrochloric acid, and 70% of the value too high using chloro- acetate buffer, on the average, as compared to the amounts taken. As indicated by this work (H16), the best practice is to standardize the thorium solution against pure sodium fluoride (or fluorosilicic acid ob­ tained by distillation of standard fluoride) using the same conditions as for unknowns. Unfortunately, many workers have used some more indirect standardization of the thorium nitrate, depending on stoichio­ metric reaction and strict proportionality between the thorium and the fluoride (or fluorosilicic acid). 4 Nature of the Medium. In the main, two media have been used for the thorium titration of fluoride ion: aqueous and 50% ethanol solutions. Armstrong (A34) was the first to show that a purely aqueous solution could be used successfully in this method instead of the semialcoholic medium used by Willard and Winter (W52). The work of Rowley and Churchill (R57) indicated that the end point color change was somewhat sharper in aqueous solution. Reynolds and Hill (R25) also investigated the effect of the type of medium on this titration, as did Hammond and Maclntire (HI6) in work discussed earlier in this section. Dahle, Bonnar, and Wichmann (Dll) preferred an aqueous solution, since they found no proportionality between the thorium used and fluoride present for less than 125 Mg. F; consequently, the A.O.A.C. official method (L33) uses aqueous solutions. Rinck (R37) found aqueous solutions superior for titration of larger amounts, inasmuch as the color change was sharper using Alizarin Red S, but obtained more precise results for less than 100 Mg. using 50% alcoholic solution. Willard and Horton (H72, W47) studied the effect of varying proportions of alcohol on the sharpness of the color change with many indicators. They found that the amount of alcohol required, if any, varied with the indicator used. In their fluorometric method (W48) using thorium as titrant, they noted a minimum alcohol concentration of 30% below which less than 1 mg. F could not be titrated. However, in 50% alcohol the method extended down to less than 50 Mg- F. ANALYTICAL CHEMISTRY OF FLUORINE 115

A few other modifications in the medium have also been used in the thorium titration method. Doherty and Retzsch (D59), as well as Williams (W56), titrated in an aqueous solution containing sodium chloride to maintain a constant ionic strength. Ellis and Masgrave (E16), Nichols and Kindt (N6) and, earlier, Stross (S126) titrated in aqueous solution containing a certain amount of starch. This starch helped to keep the thorium fluoride formed colloidally dispersed in the solution. Agar, pectin, gum arabic, and analogous substances have been tried for a similar purpose. Stevens (SI20) added sodium sulfate to his blanks to compensate for the sulfate in the solutions titrated. Kazarinova- Oknina (K15) titrated in 30 or 60% alcoholic solution, while Lockwood (L45) used a 50% glycerol medium. The latter claimed that a glycerol solution has a better end point than an alcoholic solution. At times, mild reducing agents have been added to the distilled fluorosilicic acid solution usually used in the thorium titration to remove traces of free chlorine. The chlorine or hypochlorite, if present, would oxidize the organic indicator used in the titration. Sodium nitrite (L45), sodium azide (M3), and hydroxylamine hydrochloride (C72, L33, M92, W50) have been used for this purpose. Attention should be called here to the fact that Rinck (R37) and others have noted that the monochloroacetate buffer often used in this titration may hydrolyze slowly, forming hydrogen chloride and decreasing the pH of the buffer. However, Horton (H72) did not experience this difficulty. It seems more likely that this effect may be due to impurities in the monochloroacetic acid. It is surprising, in view of such observations, that a formate-formic acid buffer has not been tried in this titration. It could buffer in the preferred range (p/f« = 3.75) and also would reduce traces of free chlorine in the aliquot titrated. However, the only time such a buffer has been used was in the colorimetric method for fluoride using thorium and Alizarin Red S described by Salsbury, Cole, Over- holser, Armstrong, and Yoe (S5, Y9, Y10). Indicators. Several other indicators have been used beside sodium alizarin-sulfonate. These include the zirconium-quinalizarin lake (F38), an Alizarin Cyanine R lake, which was preferred to many other zirconium- hydroxyanthraquinone lakes (S65), methyl red (Zl), the zirconium purpurin lake (C73, K54), a mixture of Alizarin Red S and 1,2,5,8-tetra- hydroxyanthraquinone (E19), Chrome Azurol S or Sollochrome Brilliant Blue (sodium sulfodichlorohydroxydimethylfuchsondicarboxylate) (M6, M85, M86, M88, S5, Y8, Y9, Y10, Yll), Gallocyanine, Sollochrome Alizarin B150 and Sollochrome Cyanine R200 (R36), other dyes admixed with Alizarin Red S (Y8, Y9), and Alizarin Red S with methylene blue (B14, B15). Willard and Horton, who studied more than three hundred 116 PHILIP J. ELVING, CHARLES A. HORTON AND HOBART H. WILLARD

different indicators, preferred purpurinsulfonate as a visual indicator (H72, W47). They found only one new indicator, 2,3-dicyanoquinizarin, which compared favorably to the previously mentioned indicators. In addition to the visual indicators, they found that morin and quercetin are good fluorescent indicators (W48). Concentrations, alcohol content, and pH recommended for the best ones are given in Table IV. With most of the indicators used in the thorium titration, some prac­ tice is required before the analyst can judge the end point color correctly. In the case of Alizarin Red S, some wçrkers have never been able to see the correct light pinkish buff color of the equivalence point or even the pale pink which follows with an extra drop of dilute thorium solution. Owing to these difficulties, attempts have been made to devise artificial color standard solutions, such as standards prepared from cobalt nitrate and potassium chromate solutions (E5, M52). Kolthoff and Stansby (K54) used a similar mixture for comparison with a zirconium titrant and purpurin indicator. Others (S37) have had difficulty with the zirconium- Alizarin Red S lake indicator. Taylor (T22) noted that the latter mixed indicator should be freshly prepared. This is probably duc to a slow coagulation of the lake which then does not disperse well in the solution titrated. Miscellaneous Experimental Conditions. Most titrations are carried out in Nessler tubes, although some have been done in Erlenmeyer flasks or beakers. The usual volume titrated is 40 to 50 ml. However, a volume of I to 10 ml. (H73, VI1) has been used successfully. The use of 100 ml. (G9) is not recommended. Temperature variations of the titration do not influence the precision of the titration (R37). Dahle, Bonnar, and Wich­ mann (Dll) found that the rate of titration has no effect on the precision of the determination. This holds true in the present authors' experience if the last few drops of thorium nitrate solution are added with color comparison between each drop. The present authors recommend the use of Alizarin Red S in 50% alcoholic solution buffered with formate-formic acid (pH 3.0 in water) for the titration of less than 50 Mg. of fluoride using thorium as titrant; pur­ purinsulfonate in aqueous solution buffered at pH 3.3 for 50 to 400 μζ. fluoride; and the photofluorometric volumetric method (W48) discussed in Section V-E-3 for titration of more than 0.5 mg. Use of a formate buffer is suggested both to take care of traces of oxidants in the solution and for better stability. If the thorium is standardized under the same conditions as are used for analysis of samples, any "salt errors" due to the buffer will cancel out. Use of a buffer will give a more constant hydrogen ion activity and ionic strength than adjustment to a certain acidity as now recommended by the A.O.A.C. ANALYTICAL CHEMISTRY OF FLUORINE 117

Interfering Substances. There are many interfering ions in the thorium titration of fluoride, as is true for all methods for the determination of this element. All ions which complex or precipitate thorium in a weakly acidic solution; all ions which give colors with or oxidize the indicator; and all ions which precipitate easily with fluoride in acidic solution inter­ fere quite seriously. After separation of the fluoride as fluorosilicic acid, the only substances which need to be considered from the previous groups are sulfate and phosphate ions and free chlorine. The free chlorine which can occur by decomposition of perchloric acid or carryover in the distilla­ tion has been removed by the use of sodium nitrite (L45), sodium azide (M3), hydroxlyamine hydrochloride (C72), or a formate buffer. Many other substances interfere very slightly. Among these are the alkali metal chlorides, nitrates, perchlorates, acetates, and monochloroacetates. It is because of this that Clifford (C72), Hammond and Maclntire (H16), and the A.O.A.C. (L33) adjust to a constant acidity instead of using a buffer. Nevertheless, by standardizing the thorium solution with the same amounts of these ions present, their slight effect can be minimized. The back-titration procedure specified by the A.O.A.C. (L33) helps, to minimize the variations which are caused by the alkali salts present, gives a comparison solution, and eliminates the indicator blank.

2. Zirconium Titration Although zirconium forms a very stable insoluble fluoride, it has not been used extensively as a volumetric titrant. Kolthoff and Stansby (K54) developed and others (C37, G52, K55, K61, S22, W80) have used zirconium oxychloride in 10 Ν hydrochloric acid as titrant and purpurin as indicator in a medium which was 8.5 Ν in hydrochloric acid. The method was unsatisfactory below 0.5 mg. F but could be used in the presence of aluminum or boron whose fluoride complexes are less stable than the complex with zirconium at high acidities. The end point color was matched to a K2Cr207—Co(N03)2 standard tint. Without distillation of the fluoride, Schloemer (S22) determined 0.01 to 0.1 % F in ashed food with a precision of ±10 to 15% of the actual content. Ugnychev and Bilenko (U2) determined fluoride in apatite by titration with zirconium nitrate in a strongly acidic aqueous solution (over 8 N) using Alizarin Red S as the indicator. Specht (S101) used a similar method as a check on the lead chlorofluoride method for the 1- to 35-mg. range. Nolke (N23) determined fluoride by using the unknown in acetic acid as titrant into a solution containing a known amount of zirconium and a trace of Alizarin Red S indicator, frequently stirred with a layer of amyl alcohol. At the end point the alizarin sulfonic acid liberated from the metal com­ plex goes into the amyl alcohol layer, giving a yellow tint which is more 118 PHILIP J. ELVING, CHARLES A. HORTON AND HOBART H. WILLARD easily detected than in the aqueous layer. Von Zeppelin (Z7, Z8) has also determined fluoride in a microscale by titration with zirconyl chloride in strongly acidic solution.

3. Titration with Iron (III) and Aluminum (III) Methods for the volumetric estimation of fluoride using ferric or aluminum solutions as titrants depend on precipitation of the insoluble sodium fluoro complexes, Na3FeF6 and Na3AlF6, respectively. Another portion of this chapter shows that FeFe and A1F6 are not the only fluoro complexes formed by these two metal ions. Thus, even if the sodium salts precipitate, a small amount of lower fluoro complexes may remain in solution and vitiate the results slightly. The visual titrimetric method for fluoride using ferric ion was de­ veloped by Greeff (G64), who indicated that Na3FeF6 precipitated stoichiometrically in a neutral, semialcoholic solution, saturated with sodium chloride. Thiocyanate ion was used as an indicator, and the first excess of ferric ion gave a blue color in a supernatent layer of ether. An accuracy of about 0.2% was reported (G64) for the titration of 250-mg. samples of alkali fluoride. Fairchild (F7) modified the method by addition of an excess of standard ferric solution and iodometric back-titration of the excess. Similar methods have been used by others (B44, D52, F51, G44, N21, S89, S90, S92, S97, S108, U4, VIO, W66, W67). Uebel (Ul) titrated an aqueous solution saturated with sodium chloride, detecting the end point using thiocyanate in a supernatent layer of amyl alcohol Treadwell and Kohl (T50) showed that, in spite of modifications, Greeff's method (G64) is not applicable for less than 20 mg. of fluoride. All these workers used ferric chloride as the titrant. Bellucci (B37) was able to use the visual method described above only down to 200 mg. F, and he found that the optimum amount of sodium chloride necessary varies with the amount of fluoride being determined and the sample size. Others (B44) indicated that the test solution must be exactly neutral for satisfactory results. Foster (F51), Smith (S89, S90), and Urech (U4) all found that the Greeff-type method gives high results. Visintin (V10) used a modifica­ tion of the method for the milligram range of fluoride, but Giordani (G44) indicated that the modification gives inaccurate results. Treadwell and coworkers (T50, T51, T52, T53) have also suggested potentiometric methods for fluoride using ferric ion as the titrant. The potentiometric methods and a concentration cell method (L51) are discussed in a subse­ quent section. The visual titrimetric determination of fluoride using ferric chloride as the titrant is not recommended by Rinck (R37) or the present authors. Kurtenacker and Jurenka (K75) proposed the titration of fluoride with ANALYTICAL CHEMISTRY OF FLUORINE 119 neutral aluminum chloride as the titrant, based on formation of insoluble sodium fluoroaluminate. They titrated in a hot solution (75°), using methyl red as the indicator, which changes to the red acidic form with excess aluminum. The method (K75) was satisfactory for less than 20 μg. of fluoride; the precision was ±0.2 mg. at the 3-mg. fluoride level. Phos­ phate, borate, and other ions forming insoluble aluminum salts interfered. Geyer (G34) used a similar method after separation of the fluoride as hydrofluorosilicic acid. Rinck (R37) obtained satisfactory results with the method for the 20- to 500-mg. F range and found that chloride and nitrate ions did not interfere, whereas sulfate made the end point indis­ tinct and the results high. Saylor and Larkin (S17) used aluminum chloride as titrant into a solution neutralized to the phenolphthalein end point, saturated with sodium chloride and kept at 90°. Eriochrom- cyanine R (sodium o'-sulfohydroxydimethylfuchsondicarboxylate) was used as the indicator for excess aluminum; lead, nickel, chromium, car­ bonate, silicate, sulfide, and sulfate ions interfere. One difficulty experi­ enced in titration with neutral aluminum chloride solution is that the titrant changes its titer with time. Fuchs (F70) proposed titration of fluoride with a basic solution of aluminum, which should retain a constant titer and only require phenol­ phthalein as an indicator. The end point occurs when the red basic color of the indicator disappears after reaction of the hydroxyl groups, liberated by precipitation of Na3AlFe, with aluminum ions in the titrant. Rinck (R37) did not obtain satisfactory results using this procedure. Conductometric and fluorometric methods using aluminum salts as a titrant for the determination of fluoride are discussed elsewhere in this chapter. Tananaev and Levina (T16) have used an inverse modification of the methods discussed above for the titrimetric determination of aluminum. They titrated a neutral solution, saturated with sodium chloride and containing ammonium thiocyanate and a trace of ferric ion, with a stand­ ard sodium fluoride solution. The equivalence point was detected by the decolorization of a supernatent layer of t-butanol. The accuracy was only ±15%; cobalt, nickel, manganese, zinc, and copper ions did not interfere, but magnesium and iron did.

4. Titration with Cerium (III) and Rare Earth Metal Ions Of the rare earth metals, only cerium (III) and yttrium have been used in volumetric methods for fluoride. The methods are based on the formation of insoluble rare earth metal fluorides. Kurtenacker and Jurenka (K75) apparently were the first to titrate fluoride with cerous nitrate, using methyl red as indicator. The cerium 120 PHILIP J. ELVING, CHARLES A. HORTON AND HOBART H. WILLARD solution had to be standardized empirically using standard fluoride. The method is sensitive to the presence of sulfate and chloride ions. Batchelder and Meloche (B26) first used an indirect procedure for this type of reac­ tion ; excess standard cerous nitrate was added to the neutral solution of fluoride, the precipitate of cerous fluoride was coagulated by addition of zinc oxide and heating to 80°, and the excess cerous ion was back-titrated with standard potassium permanganate. To avoid low results due to adsorption of permanganate ion on the precipitate, they found it neces­ sary to filter the solution just before the end point was reached. The method was applicable to the 1- to 100-mg. F range with an accuracy of about ±0.3 mg., but even the authors admit that the method is difficult to perform. Direct titration of fluoride with cerous solution was tried using Ampho Magenta as indicator but the accuracy was only 100 mg. F at best. Subsequently, Batchelder and Meloche (B27) used methyl red as indicator in the direct titration of alkali fluorides with cerous nitrate. They claimed that methyl red adsorbs on (CeF3)riF~ micelles formed during the first portion of the titration and that, at and after the end point, the methyl red is desorbed and hydroxyl ions are adsorbed on (ĎâŃ3)çĎâ+++ micelles, thus changing the indicator color. Henne and coworkers (H75, S37) used a similar method with a mixed methyl red- bromcresol green indicator. The method was fairly satisfactory for large amounts of fluoride. Nichols and Olsen (N7) also used methyl red with bromcresol green as a visual indicator in the cerous titration, using a special titration flask of two round-bottom flasks connected together. They titrated a hot solution containing 50% alcohol, with pH adjusted to pale pink using cresol red, until a maximum color difference occurred between the two flasks. The contents of the flasks were intermixed be­ tween drops of titrant. However, their (N7) potentiometric method was more satisfactory. Potentiometric methods (A14, N7, T50) using cerous nitrate and a conductometric method (W55) using cerous picrate or trichloroacetate as titrants for fluoride are described elsewhere in this chapter. Use of this titrant is also discussed by Kolthoff and Stenger (K55). Frère (F57), and later Bégué (B32), titrated 03 ot 300 mg. fluoride, as alkali fluoride, ni a solution neutralized ot phenolphthalein, with yttrium nitrate using methyl red indicator. Alkali chlorides ro nitrates cause errors varying from —1 ot —6%; sulfates lead ot positive errors. Frère (F57) compensates for these ions yb titrating blanks containing the same amount fo sodium and/or potassium chloride, nitrate and/or sulfate. The volumetric titrations mentioned using cerium (III) ro yttrium nitrates apply only ot samples fo relatively high fluoride content and ni ANALYTICAL CHEMISTRY OP FLUORINE 121

general do not give very good accuracy. Thus, these methods are not recommended.

5. Indirect Titration of Lead Chlorofluoride and Calcium Fluoride Starck (SI 14) first attempted an indirect volumetric method for fluoride based on precipitation of lead chlorofluoride. He dissolved the precipitate and attempted to determine the chloride content by standard technics; his results were less accurate than by his gravimetric procedure outlined in a previous section. Hawley (H41) had better results by dis­ solving the precipitated lead chlorofluoride in nitric acid and titrating the chloride by the Volhard method. The precipitate was filtered at 15°, washed with water saturated with lead chlorofluoride, and finally washed with cold water. Hoffman and Lundell (H66, H67) added lead nitrate and sodium acetate to a heated solution neutralized to bromphenol blue containing some sodium chloride and 2 ml. 1:1 hydrochloric acid to precipitate lead chlorofluoride. After standing overnight, the supernate was decanted through Whatman No. 42 paper, and the precipitate washed once with water, four or five times with a saturated lead chlorofluoride solution, and once more with water. Finally, the lead chlorofluoride was dissolved and the chloride titrated by the Volhard technic. Others (C48, E27, M47, M67, P52, R27, S56, S102, S103) have used a similar procedure. Kapfenberger (K9) precipitated lead chlorofluoride from a solution neutralized to methyl orange with 0.5 ml. 1 Ν hydrochloric acid in excess by adding saturated lead chloride solution and then base to the yellow color of methyl orange (pH 4.5-4.7) with constant stirring. After standing overnight, the precipitate was filtered, and its chloride content determined by the Volhard method. Tananaev (T8) precipitated lead chlorofluoride using a known excess of sodium chloride solution and an excess of lead nitrate. After filtration, excess chloride in the filtrate was determined volumetrically using an adsorption-type fluorescein indicator. Geyer (G34) and Rinck (R37) used a diphenylamine blue indicator in a similar method. Rinck (R37) found it necessary to age the lead chlorofluoride precipitate several hours before filtration. The A.O.A.C. has adopted a method for fluoride in insecticides de­ pending upon formation of lead chlorofluoride as one of three suggested procedures (L32, L33); the method is a variation of the Hoffman and Lundell procedure (H66) described. The Aluminum Company of America (A 18) also uses in some cases a volumetric method based on this reaction. The adoption of this method by the A.O.A.C. was based in part on collaborative studies of this indirect volumetric determination of lead chlorofluoride as compared to other methods reported by Donovan (D62, 122 PHILIP .J ELVING, CHARLES A. HORTON AND HOBART H. WILLARD

D64, D65, D66) and Graham (G62). Satisfactory results were obtained (D65) for more than 01 mg. fluoride ni insecticides, but results were low if over 05 mg. boron was present (D62). Frère (F57), however, has found that the volumetric lead chlorofluoride method gives results about 1% low for cryolite analyses. A group fo workers (Si6) studied several modifications fo volumetric and gravimetric methods for fluoride involving the formation fo lead chlorofluoride, including the Hoffman-Lundell (H66) volumetric method and a gravimetric modification, the Kapfenberger (K9) and Hawley (H41) volumetric methods, and two new modified precipitation technics. They found that conditions must eb carefully controlled ot obtain a precipitate fo lead chlorofluoride fo the stoichiometric composition. When the chloride ot fluoride ratio was equal ot ro less than ,1 the lead ot chloride ratio ni the precipitate was greater than 1 yb 2 % ro more ; when the chloride ot fluoride ratio was over 1.5, the lead ot chloride ratio ni the precipitate was less than 1 yb about 2%. Their work indicated that equilibrium between lead chloride and fluoride for formation fo lead chlorofluoride si very slow ni aqueous solutions ta either 05 ro 70°. They preferred Kapfenberger's (K9) volumetric method but reached the con­ clusion that methods based no formation fo lead chlorofluoride can never be more precise than 5 parts ni 1000. A complicated indirect method (15), which si not recommended yb the present authors, uses a known amount fo chloride ni precipitation fo lead chlorofluoride. A standard mercurous nitrate solution si added ot the filtrate ot remove excess chloride sa mercurous chloride. After filtration, excess mercury si determined yb titration with sodium chloride. Amperometric ro polarographic methods (H32, S101) and a potentio­ metric method (C89) for determination fo fluoride based no formation fo lead chlorofluoride are discussed ni another section. Fluoride has also been determined yb precipitation fo lead bromo- fluoride and determination fo the bromide ni the precipitate (C52, E13) or the bromide left ni the filtrate when a known amount was added for the precipitation (V2). A modified Volhard procedure was used for the titration ni all cases. Precipitation as Calcium Fluoride. Garcia (G6, G7) first used na in­ direct volumetric method for fluoride depending no precipitation fo calcium fluoride. Aknown amount fo calcium and oxalate was added ot the sample ot precipitate calcium fluoride and oxalate ; the excess calcium in the filtrate was determined yb precipitating ti sa oxalate, followed yb titration with permanganate ni the usual way. Olivier (09), no the other hand, determined calcium volumetrically ni calcium fluorides ores after leaching the powdered sample with dilute acetic acid. Scott (S42) and ANALYTICAL CHEMISTRY OF FLUORINE 123

Uebel (Ul) used modifications fo both fo these methods, precipitating only the fluoride and titrating the excess calcium ni the filtrate ro that ni the precipitate. Carrière and coworkers (C22, C24), Shinkai (S55), and others (T18, T57) also precipitated calcium fluoride from na ammoniacal solution using a known excess fo calcium and then precipitated the excess as oxalate and titrated ti with permanganate. They obtained na accuracy of ± 1 % for precipitation ni platinum but poorer results for precipitation in porcelain vessels. Frost (F69) reported na accuracy fo ±0.2 ot 0.4% for a similar procedure. Mougnaud (Ml 18, Ml 19) criticized the volu­ metric methods using precipitation fo calcium fluoride from both acetic acid and ammoniacal solutions, claiming some solubility fo the fluoride in the first medium and adsorption fo calcium yb the precipitate ni the second medium. As discussed elsewhere ni the chapter, Piccard and Buffat (P41) titrated fluoride conductometrically, Stevens (S121) and Peterson (P31) nephelometrically, and Ryss and Bakina (R77) poten- tiometrically ni methods which depended no the insolubility fo calcium fluoride.

C. TITRIMETRIC METHODS: NEUTRALIZATION

1. Reactions Involving Fluorosilicic Acid and Potassium Fluorosilicate Numerous volumetric methods for fluoride are based no the formation or hydrolysis fo the fluorosilicate ion, sa indicated yb the following reaction : 4H+ + Si02 + 6F- ^ SiFe— + 2H20 (1) One method, proposed yb Siegel (S67), depends no formation fo fluoro­ silicate yb reaction fo a neutral alkali metal fluoride with hydrochloric acid and silica gel. nI this method, for every 6 moles fo fluoride, 4moles of acid are consumed ni formation fo 1 mole fo fluorosilicate. Completion of the reaction was detected with methyl red indicator. Fluorosilicates of the other metals give na acidic reaction due ot hydrolysis before all the fluoride has reacted. Ruiss and Bezmenova (R68) showed that SiegePs method (S67) gives low results owing ot slow reaction fo the silica gel with the fluoride and the unsuitable pH color range fo methyl red. They and others (C9) substituted sodium silicate sol for the gel, and a 1:1 mixture fo 0.1% dimethyl yellow ro methyl orange with methylene blue as the indicator. They claimed that these mixed indicators are neutral ot fluorosilicate ion, os that fluorosilicate need not eb removed sa ni the other methods subsequently discussed. Their precision no the determina­ tion fo about 0.2% F was better than ±0.5% fo the value. Neudorffer (N3) published a similar procedure. 124 PHILIP J. ELVING, CHARLES A. HORTON AND HOBART H. WILLARD

Decomposition of the fluorosilicate by a base is better suited for the alkalimetric determination of fluoride. This can proceed in two steps; or in both steps at once:

H2SiFe + 2KOH K2SiFe + 2H20 (2) K2SiFe + 4KOH -> 6KF + Si02 + 2H20 (3) H2SiFe + 6KOH -> 6KF + Si02 + 4H20 (4) Penfield (P27, P28) separated fluoride from interferences as silicon tetra­ fluoride, which was absorbed in water forming hydrofluorosilicic acid :

3SiF4 + 2H20 -> 2H2SiFe + Si02 (5) It is necessary, in the titration of the fluorosilicic acid with base by reac­ tion 2, to prevent hydrolysis of the potassium fluorosilicate formed. Penfield (P27, P28) added to the fluorosilicic acid a large excess of potas­ sium chloride, liberating 2 moles of hydrochloric acid, which were titrated in a cold 50% alcoholic solution with standard base using litmus or methyl red indicator. Van Kampen (K4) and others (A26, L2, R50) have used modifications of the same method. In a cold solution containing alcohol, the potassium fluorosilicate is rather insoluble and hydrolysis according to equation 3 is minimized. Schucht and Moller (S30) determined fluoride in fluorosilicic acid by addition of excess calcium chloride rather than potassium chloride, and titration, with base, using methyl orange indi­ cator. The total reaction is :

H2SiFe + 3CaCl2 + 6NaOH -> 3CaF2 + 6NaCl + H2Si03 + 3H20 (6) Thus, this procedure is a variation of reaction 4. The calcium fluoride and silica formed remained in colloidal solution. A similar procedure was used for analysis of alkali metal fluorosilicates using phenolphthalein indicator (reaction 3). Shuey (S60) and Babko (B2) stated that the method is satisfactory for pure fluorosilicates, but Anosov and Chirkov (A26) found this method and that of Treadwell (T49), mentioned below, less satisfactory than the Penfield (P27) method. Zarin and Dubnikov (Z3) recommended a variation of the Penfield method in which barium fluorosilicate was precipitated by addition of barium chloride and the liberated acid titrated (equation 2). Katz (K12) titrated fluorosilicic acid with base first in an aqueous solution and then in a 50% alcoholic medium, and claimed, incorrectly, that the acid requires equal amounts of titrant for both media. Dinwiddie (D52) showed that alkalimetric titration of fluorosilicic acid in 50% alcohol containing potassium chloride required 1 to 3% more than one- third of that required in an aqueous titration, apparently because of slight ANALYTICAL CHEMISTRY OF FLUORINE 125 hydrolysis of the potassium fluorosilicate formed in the alcoholic-KCl solution titration. Treadwell and Koch (T49) and Honig (H69) also titrated fluorosilicic acid in 50% alcohol, using either potassium or barium hydroxide as the titrant and phenolphthalein as indicator; the reaction proceeds according to equation 2. Liverside (L42) used a similar titration technic after treat­ ing the fluorosilicic acid with silicic acid and ammonium hydroxide at 100°. Offerman (02, Ul) titrated fluorosilicic acid according to reaction 4, using cochineal as indicator. Others (A26, H63, S43) used phenolphthalein as indicator under similar conditions. Reaction 4 is favored by titration in a hot solution. Schucht and Moller (S30) showed that the results by reaction 4 are always a bit high, increasingly so as the volume of solution titrated increases. Wagner and Ross (W4) also used a procedure based on the reaction of equation 4 and speeded the reaction by boiling the so­ lution; however, the titration was completed at room temperature. Patten (P17), as the result of a collaborative study, recommended the Wagner and Ross (W4) method. Later, similar studies (Bll, Mill, M112, Ml 13, Ml 15, S60) criticized this method mainly because of its use of the tedious separation of the fluoride as silicon tetrafluoride. Stolba (S123) was probably the first to suggest precipitation of potas­ sium fluorosilicate (equation 2) in the presence of alcohol, filtration, and titration of the precipitate with caustic (equation 3) as a method for the determination of fluoride. Travers (T41, T42) improved and checked this technic, as did Specht (S102) and others (B29, D62, D64, S88) later. Modifications of the Stolba-Travers method are still accepted for certain types of samples by the A.O.A.C. (L33, L34). The A.O.A.C. modified Travers method for insecticides calls for boiling the sample for a minute with precipitated silica and a slight excess of hydrochloric acid, cooling, and precipitating potassium fluorosilicate by the addition of potassium chloride and alcohol. The precipitate is dispersed in water and titrated with standard sodium hydroxide to the phenolphthalein end point. Donovan (D64) found more variability in results using a modified Travers method than for two other methods tested. Brinton and coworkers (B98) determined both fluorosilicic and hydro­ fluoric acids in mixtures by titrating both acids in the cold (reaction equation 2 for H2SiFe) and then at 80°, using phenolphthalein indicator. In the second titration the fluorosilicate formed in the first step is decom­ posed according to equation 3. For good results the authors caution that the base used must be free of silicates. Jacobson (J3) and Vasil'ev (V3) used a similar technic. The latter showed that the silica added to titrate hydrofluoric acid was incompletely precipitated in presence of potassium 126 PHILIP J. ELVING, CHARLES A. HORTON AND HOBART H. WILLARD chloride. Jacobson (J3) used methyl orange as the first step indicator and applied a correction factor to the second titration value. Hart (H23) discussed the determination of fluorosilicate in the presence of boric acid, and the analysis of alkali fluoride, bifluoride, and fluorosilicate mixtures in insecticides. He determined total acidity in both cases by titration with base. To determine the bifluoride he removed fluorosilicate by precipitation induced by the addition of alcohol and potassium chloride, and then titrated the fluoride with base at a low temperature. Kern and Jones (K16) recommended titration of fluorosilicic acid with base at room temperature (equation 2) using dimethylaminoazobenzene as indicator; they tried rosolic acid, phenol red, thymol blue, and phenol­ phthalein for titration of the alkali fluorosilicate formed (equation 3). Kobayaski (K43) titrated hydrofluoric acid in the presence of fluorosilicic acid at room temperature in the presence of 5 to 10% potassium nitrate using phenol red indicator. He claimed that results at room temperature were just as satisfactory as for titrations near 0°. Tananaev (T18) also used phenol red indicator for the titration of fluorosilicic acid. Others (B80, S78) added silica gel, potassium chloride, alcohol, and excess stand­ ard hydrochloric acid to neutralized fluorosilicic acid and then back- titrated with standard base using methyl orange indicator; bromcresol green has been used as the indicator in a similar procedure (S13). Geffcken and Hamann (G27) used a similar technic for hydrofluoric acid, without addition of the silica, with bromcresol purple as indicator, in the analysis of mixtures of hydrofluoric and fluorosilicic acids. For the fluorosilicate content, they titrated further at the boiling point to a naphtholphthalein end point. Thomsen (T35) has recently studied the alkalimetric titration of fluorosilicic acid. For titration of samples containing high amounts of silica, he recommended a titration with caustic at room temperature followed by one in the hot, using a mixture of bromthymol blue and phenolphthalein as the indicator. The second end point was used to cal­ culate the fluorosilicate content. For samples containing low amounts of silica or an excess of hydrofluoric acid, he found that in the absence of excess hydrofluoric acid one-third as much caustic was required to titrate to methyl orange as compared to titration using phenolphthalein. For titration of sodium fluoride containing impurities, he added colloidal silica and standard acid and titrated with base to the methyl orange end point to minimize hydrolysis of the fluorosilicate ions formed by interac­ tion of the fluoride and silica. The fluoride present is 1.5 times the equiva­ lents of titrant used. In connection with the titration of fluorosilicic acid, it is interesting ANALYTICAL CHEMISTRY OF FLUORINE 127 to note that in 1921 Hudleston and Bassett (H78) found that the complete neutralization of this acid in hydrofluoric acid mixtures by base requires an appreciable length of time. They recommended plotting the base required versus time and extrapolation back to zero time for calculation of the fluorosilicate concentration. Kubelka and Pristoupil (K68) deter­ mined the equilibrium constant for the reaction :

SiF6— + 2H20 -• Si042 + 4H+ + 6F~ K = ^ ·[H+l mk= 10-26·6 Ruiss and Bakina (R66) report an almost identical constant. The alkalimetric methods discussed in the preceding paragraphs are now used principally for the determination of fluorosilicate rather than as a primary method for determination of total fluoride. This is true in spite of the fact that fluoride is generally separated from interferences as fluorosilicic acid by some modification of the Willard-Winter (W52) technic. The reason is that volumetric methods using thorium are gen­ erally more specific, sensitive, and precise.

2. Miscellaneous Neutralization Titrations Free hydrofluoric acid, a moderately weak acid, may be titrated directly with caustic using phenolphthalein as indicator. This technic was used as early as 1893 by Meslans (M65) for the hydrofluoric acid derived from organic compounds. Others (H6, S101, W65, Z6, Z10) also have titrated hydrofluoric acid with sodium hydroxide and phenolphthalein indicator. Winteler (W65) was the first to show that methyl orange is unsatisfactory as indicator for alkalimetric titration of hydrofluoric acid. Ginsberg (G43) used this type of titration after evolution of fluoride as hydrogen fluoride. Watson and coworkers (W17) absorbed HF from the decomposition of fluorine-containing plastics in standard sodium carbonate and titrated this solution with base or acid as necessary using phenolphthalein indicator. They also titrated with thorium nitrate in the presence of Alizarin Red S, using the procedure of Matuszak and Brown (M52). The titration with base has also frequently been used in Atomic Energy Commission laboratories for samples from which fluoride has been separated by the pyrohydrolysis technic discussed in another section. Generally, however, this simple titration procedure is not applicable because of the presence of other stronger acids, of buffering salts, or of ions which complex fluoride. The Fisher Scientific Company (F39) has supplied electrodes of gold and calomel in plastic for the pH titration, in the presence of quinhydrone and alcohol, of hydrofluoric acid with 128 PHILIP J. ELVING, CHARLES A. HORTON AND HOBART H. WILLARD barium chloride. These electrodes could be used, with plastic beakers, for the ordinary neutralization of hydrofluoric acid. Ryss (or Ruiss) and coworkers (R75, R79) studied formation and hydrolysis of fluoroborate ion, BF4~, by titration with base or addition of excess base and back-titration with standard acid. Kern and Jones (K16) had earlier indicated that only approximate results are obtained in the titration of fluoroboric acid with sodium hydroxide. Ramsey (R9) titrated acetic, mono-, di-, and trifluoroacetic acids with base after separation by partition chromatography. Savchenko and Tananaev (Sll) have titrated K2TaF7 in dilute hydrofluoric acid, in the presence of calcium chloride, with sodium hydroxide using methyl red indicator. Each mole of the tantalum complex fluoride required 4.5 moles of base. Zschacke (Z10) titrated a mixture of hydrofluoric and sulfuric acids with base first in the cold and then in the hot, using phenolphthalein as indicator to give the total acidity ; fluoride was determined separately by titration with ferric solution using thiocyanate indicator, or by dis­ tilling the hydrofluoric acid at 50 to 60° over a 24- to 36-hour period. Ginsberg (G43), however, analyzed a similar mixture of acids by distilla­ tion of the hydrofluoric acid from the sulfuric at 110 to 160° in three steps requiring a total of 45 minutes. The separated acids were both titrated with caustic. Swinehart, Bumblis, and Flisik (S130) analyzed boron trifluoride for various impurities and BF3 itself. To determine the boron trifluoride they weighed a sample into water, added neutral calcium chloride, titrated in the cold with caustic, brought the solution to a boil, heated at 90° for 10 minutes, and again titrated, using phenolphthalein indicator both times. Booth and Martin (B79) discuss analysis of boron trifluoride extensively in their book.

D. ELECTROMETRIC METHODS The relative inertness of fluoride ion electrochemically has resulted in the proposal of almost no electrometric methods which involve the direct measurement of fluoride ion. Instead, methods have been developed which involve the effect of fluoride ion upon the observed behavior of other electroactive ions such as ferric and uranyl ions. Titrimetric methods for fluoride have utilized the various commonly available technics for equivalence point detection as well as the more recently developed high frequency oscillators. 1. Potentiometric Titration A number of titration procedures involving potentiometric end points have been proposed for the determination of fluoride ion. These have ANALYTICAL CHEMISTRY OF FLUORINE 129 included complex-formation reactions and precipitation reactions, with some of the latter being observed through measurement of the change in pH. The most widely studied reactions have been those involving formation of ferric fluoride complexe++ s++ and precipitation of cerous fluoride. The electrode system, Fe , F+ /Pt, has been used in the potentio­ metric titration of fluoride by the supposed formation of FeF6 (B105, G64, T50, T51). The relation between measured potential and fluoride ion concentration is empirical and does not fit with the proposed FeF6 complex. Study of the complex ions formed by ++iron wit+h fluoride (D56, D58) leads to the postulation of the species FeF , FeF2, and FeF3. The break in potential apparently does not coincide with the equivalence point; the latter is best determined graphically from the data (T2). The potential break at the end point is increased and the potential read­ ings are stabilized by (a) saturating the sample solution with sodium or potassium chloride; (b) adding ferrous ion to the ferric chloride titrant solution; (c) making the solution 50% in ethyl alcohol; and (d) keeping the sample volume small, preferably 5 or 10 ml. (T2, T3, T50). The accuracy seems to be only fair. Iodide ion interferes; bromide, nitrate, and sulfate do not. Treadwell and Kohl (T50) suggested making the solution, which was 48% alcohol and saturated with sodium chloride, barely acid to phenol­ phthalein and titrating with 0.1 TV" ferric chloride solution containing about 1 % ferrous chloride. The solution is stirred with a stream of carbon dioxide. Slightly past the equivalence point, there is a sudden change in potential which can be simply measured on a millivoltmeter connected to a platinum wire electrode and a silver chloride reference electrode. Other cations, e.g., chromium, vanadium, uranium, titanium, columbium, and zirconium, which form fluoride complexes, were less satisfactory as titrants than iron. Later, the same authors found (T51) that, since complex FeF6 is decomposed by the addition of aluminum ion, it is possible to titrate fluoride ion accurately by adding ferrous-ferric solution to the concentrated solution containing alcohol and sodium chloride, and then titrating with aluminum chloride solution. Treadwell and Zurcher (T53) found that solutions containing silicic acid can be rendered silica- free and suitable for potentiometric titration with ferric ion by boiling the sample solution with freshly precipitated cadmium hydroxide. The ferric fluoride complex can also be utilized as the basis for the potentiometric titration of fluoride ion with lead ion (C89). The sample solution is 50 % ethyl alcohol and contains sodium chloride and an acetate buffer (pH 3.5). About 90% of the expected volume of 0.1 Ν lead nitrate solution is added; then, some ferrous-ferric ion solution is added, and titration is continued. The dissociation of the ferric fluoride complex 130 PHILIP J. ELVING, CHARLES A. HORTON AND HOBART H. WILLARD at the equivalence point du++e to precipitatio+ n of the fluoride as PbCIF causes a change in the Fe+ , Fe+ redox potential. The measurement of low fluoride ion concentrations of the order of 0.2 to 60 mg. per liter, e.g.,.the amounts encountered in water or phos­ phate rock analysis, ha+++s been performed using a concentration cell involving the Fe++, Fe /Pt electrode (L51). The detailed procedure must be carefully followed to obtain the possible 1% accuracy. The method has been refined (B122), using gold electrodes and measuring potentia3 l changes rather than potentials themselves. A sensitivity of 10~2 μg. of fluoride has been attained in 10 to 20 ěŔ. of solution containing 10~ μg.; the best sensitivity is in the concentration rang6 e belo5 w 0.01 M. Fluoride ion concentration in the range of 10~ to 10~ M can b+e+ determined by a potentiometric measurement in the presence of Fe+ ion; the e.m.f. measured varies almost linearly with fluoride ion concen­ tration (B105). The method is based on the apparent fact that, in 4th4e presenc+e of fluoride, ferric ion forms the series of complex species, FeF * ", FeF2, up to FeF β ; in low fluoride ion concentration, only equilibrium processes involving formation of the first two ions need be considered. Cerous solution can be used to titrate fluoride potentiometrically using a ferri-ferrocyanide-platinum electrode (A14, P53) or a glass elec­ trode (N7) as indicating electrode system. Fluoride (4 to 120 mg.) in 100 ml. of 50% alcohol solution (kept at 70° ± 10°), containing a minute amount of K3Fe(CN)6 and some KCeFe(CN)e suspension, is titrated with Ce(N03)3 solution using platinum and calomel electrodes. The average error was about 0.3% in the determination of 0.1 to 50 mg. fluorine'. Acidic solution and large amounts of salts, particularly sulfate, obscure the potential break. The fluoride in the solution obtained from a sodium peroxide bomb fusion can be titrated in hot 50% alcohol solution with 0.01 Ν Ce(N03)3, using the glass electrode as an indicator for the maximum potential (pH) change per unit volume of titrant; the error is 1% (N7). In addition to the foregoing pH titration for fluoride precipitation, fluoride has been titrated with calcium nitrate in solution saturated with sodium fluorosilicate, using a quinhydrone electrode system (R77). Similarly, after adjustment to pH of about 7, an alkali fluoride solution can be titrated with 0.01 Ν thorium nitrate of about pH 4, using a glass electrode; the end point is the inflection point on the pH-volume curve (B122). Hydrofluoric acid can be determined or assayed potentiometrically by an indirect titration (F39). The sample, 0.6 to 1.0 g., is added to a Monel beaker containing 20 ml. of 30% barium chloride solution and 50 ml. of methyl alcohol; the beaker is set in a cracked ice-salt mixture. After the ANALYTICAL CHEMISTRY OF FLUORINE 131 addition of 0.5 g. of quinhydrone, the hydrochloric acid, formed by precipitation of barium fluoride, is titrated with 1 Ν sodium hydroxide solution, employing a gold electrode and a Bakélite calomel electrode. The procedure is applicable to samples containing from a few to 52% HF; the sensitivity of the method is 0.1% as is also th+e accuracy. The reaction between alkali fluorides and U ; to form the slightly soluble MUF6 can be followed potentiometrically in the presence of a small amount of uranyl salt (F43, F44) ; there is little change in potential until completion of the reaction. A sulfanilat+4 e buffer is used to keep the pH low enough to prevent hydrolysis of U salt but high enough to pre­ vent formation of slightly ionized HF. Up to 0.2 g. fluoride can be titrated in 100 to 200 ml. of solution. Air is first removed from the test solution with a carbon dioxide stream. Calcium, aluminum, iron, and phosphate interfere; the fluorine in fluorosilicates can be titrated. The study (H10) of fluoride solutions at a tantalum electrode indicates that the system F~, TaF— Ta may have analytical possibilities. The shifts in the potentials of iron, titanium, and vanadium redox systems on the addition of fluoride ion have been studied (S134, S135, S136, Table I of Section V-E), as has the cell system, sodium amalgam-sodium fluoride solution-lead fluoride solution-lead amalgam (19).

2. Conductometric Titration Fluoride ion can be titrated to a conductometrically determinable end point with aluminum chloride to form A1F6 , especially if a large excess of sodium chloride is provided to cause the precipitation of Na3AlF6 (H20, J5, T50, T51). The most favorable conditions seem to be in acetate- buffered 30% ethyl alcohol solution where 2.5 to 20 mg. per 50 ml. can be titrated with 0.02 Ν aluminum chloride solution with an average error of 1% (H20, J5). Chloride, nitrate, sulfate, and silicate do not interfere. By means of test tube cells and microburets, as little as 13 μg. of fluoride can be satisfactorily titrated. Willard and Williams (W55) were able to titrate fluoride ion with cerous nitrate, an increase in conductivity being obtained due to the difference in mobility between fluoride and nitrate ions. The sharpness of the angle obtained on plotting the conductivity data before and after the equivalence point was increased by substituting a slower moving ion, such as acetate or picrate, for the nitrate. The procedure using cerous acetate is as follows: Add bromthymol blue indicator to the sample fluoride solution and adjust the acidity with 0.1 Ν sodium hydroxide or acetic acid solution to the neutral or slightly acidic shade of the indicator. Dilute to such a volume as to require 2 ml. of 0.1 Ν or 4 ml. of 0.01 Ν reagent per 100 ml. of sample solution. With the sample immersed in a 132 PHILIP J. ELVING, CHARLES A. HORTON AND HOBART H. WILLARD constant temperature bath, take five conductivity readings on either side of the equivalence point; allow 1 minute for equilibration after each addition of reagent. The maximum error was 0.05 mg. on 15 mg. and 0.003 on 0.04 mg. of fluorine in 20 ml. of solution. Nitrate, chloride, perchlorate, and silicic acid cause negligible error. Carbonate and sulfate cause high results; the former is readily removed by prior careful acidification with acetic acid and boiling. Cerous picrate is satisfactory as titrant for amounts of fluoride less than 5 mg. Fluoride can be titrated conductometrically in pure sodium fluoride solution by precipitation as BiF3, using 0.5 M BiOC104 solution as titrant (J5). The solution should contain from 30% alcohol for large amounts of fluoride to 80 or 85% for amounts under 0.05 mg. The application is limited, since the titration cannot be done in acidic solution and since many anions interfere. The conductometric titration of fluoride with calcium chloride solution has been reported, but no experimental details were given (P41). Fluoro­ silicic acid in 67% alcohol solution can be titrated satisfactorily with sodium or potassium hydroxide solution, using conductometric determina­ tion of the end point (K8). The existence of the mixed lead halides (PbFCl, PbFBr, and PbFI) has been confirmed conductometrically (D33). Solutions of strong acids and of alkaline earth metal salts can be titrated conductometrically with a sodium fluoride solution (S28) ; beryllium salts have been titrated with ammonium fluoride (P61).

3. Amperometric Titration Small amounts of fluoride (5 iig. to 1 mg.) can be titrated in 0.1 Ν potassium chloride solution (10% ethyl alcohol by volume) with dilute (0.01 to 0.1 N) thorium nitrate solution (B122, L16, L54). There is no current flow, except for the residual current, until the equivalence point. The cause of the current flow is apparently due to the reduction of nitrate in the presence of thorium or lanthanum [lanthanum nitrate can also be used as titrant (L16)]; the wave height is proportional to the thorium concentration. The sample solution should be about 0.002 to 0.006 Ν in fluoride and should be freed of oxygen. The precision is about 1 %. The results are lower than in the gravimetric determination of fluoride. Aluminum, magnesium, and phosphate interfere. Radimer (Rl) found the amperometric thorium titration to be time-consuming as well as subject to interference by silicate and fluorosilicate. Willard and co­ workers (W50) found the method unsatisfactory below 100 μg. of fluoride. Fluoride ion can also be titrated amperometrically with standard uranyl acetate solution, forming a uranium-fluoride complex; the first ANALYTICAL CHEMISTRY OF FLUORINE 133 excess of uranyl solution produces a wave due to the apparent catalytic reduction of nitrate in the presence of uranium (B122, F40, K53). The sensitivity of the method is 1 /xg. per milliliter; the average deviation is 7%. The method has been claimed to be best adapted for the rapid analysis of relatively pure fluoride solutions (F40). Lead nitrate solution can be used to precipitate fluoride in amounts of 0.5 to 10 mg. or more per 100 ml. as PbCIF in undegassed 0.1 Ν potassium chloride solution adjusted to pH 5.5 to 6.5 (P35). Current readings due to lead can be taken 3 to 5 minutes after each addition of titrant. The method thus complements the thorium titration, which is applicable to much lower concentrations of fluoride. The accuracy of the lead titration is 0.5% or better. The similar titration of 1 to 50 millimolar fluoride solutions in 50% alcohol solution has been described, but an empirically determined end point must be used (H32). Aluminum can be titrated with fluoride ion in 50% ethyl alcohol solution saturated with sodium chloride and containing a small amount of ferric iron (R38). The diffusion current due to the iron is measured. The end point is the disappearance of this current; a correction is made for the fluoride which has combined with the ferric iron. The method may have some possibility as a direct amperometric titration for fluoride or ferric ions. An amperometric method described by Castor and Saylor (C37) is based on the decrease in diffusion current of an aluminum-dye complex in the presence of fluoride ion as a result of the formation of A1F6 , with a corresponding increase in the diffusion current of the free dye. Superchrome Garnet Y (Colour Index 168) (sodium salt of 5-sulfo-2- hydroxybenzeneazoresorcinol) is used at pH 4.6 where there is 0.2-volt' separation between the two waves. The fluoride solution is titrated with a solution of the aluminum-dye complex and the current is read at —0.37 volt versus S.C.E., where only the free dye is reduced. Satisfactory results were obtained on test solutions containing 2 to 8 mg. F per 100 ml. ; in this volume range, up to 20 mg. sulfate or silicate and up to 10 mg. phosphate do not interfere. Cobalt, nickel, iron, thorium, titanium, vanadium, and zirconium interfere seriously; these ions can be removed by electrolysis at a mercury cathode or by a Willard-Winter distillation of the fluoride.

4- High-frequency Oscillator Titration Encouraging results even in the presence of large amounts of salts were obtained by Harley and Revinson (H19), who used a high-frequency oscillator circuit to determine the equivalence point in the titration of micro amounts of fluoride using thorium as the titrant. Subsequently, they reported (R21) an improved procedure in which successful results 134 PHILIP J. ELVING, CHARLES A. HORTON AND HOBART H. WILLARD were obtained using lanthanum nitrate as the titrant although thorium nitrate is almost as good. A pH range of 4 to 6 was found to be best. Chloride, borate, silicate, and sulfate do not interfere; interference due to small amounts of phosphate may be eliminated by titrating in 0.02 to 0.03 Ν hydrochloric acid solution. Many polyvalent cations interfere, from which the fluoride can be separated by distillation.

5. Polarography Fluoride ion can be determined by the difference produced by its addition upon the lead wave in a solution containing a known amount of lead nitrate and chloride ion (H32). The measured diminution in lead concentration, after correction for dilution and the solubility of lead chlorofluoride, is a measure of the added fluoride. Since the solubility correction is a complex approximation and the difference in two not overly precise measurements must be used, the method does not seem to have anything to recommend it; the amperometric procedure certainly seems preferable.

E. PHOTOMETRIC METHODS

1. Colorimetric Methods Involving Bleaching Chemistry of Fluoride Complexes. Most of the colorimetric methods for fluoride depend on competitive complex formation or precipitation of a metal ion by fluoride as compared to a chromophoric complex of the metal with some other inorganic ion or organic compound. Thus, the decrease in color of the solution is a function of the tendency of the fluoride present to combine with the metal. The metal ions usually used in such methods are: zirconium, thorium, aluminum, titanium (IV), and iron (III). Most authors assume that only the insoluble fluoride or the complex containing sufficient fluoride to satisfy the maximum coordina­ tion number is formed. These forms are: ZrFe , ZrF4, ThF4, A1F3, A1F6 , TiF6—, and FeF6 , with the insoluble compounds under­ lined. Several studies have shown, however, that the situation is not as simple as this. The existence of other complexes with fluoride or with the chromophor and, in some cases, other equilibria often account for the fact that most of the colorimetric methods for fluoride do not follow Beer's Law. For aluminum, Brosset (B103), Brosset and Orring (B106), and Kleiner (K33, K34) have shown that all possible complex ions between A1F++ and A1F6 can exist in solution. Lacroix (L6>, by following the pH of aluminum solutions containing fluoride during titration with base, indicated the presence of A1F++, A1F3, and AlFe in solution. The ANALYTICAL CHEMISTRY OF FLUORINE 135 work of Caglioti (C3) indicated that dilute solutions of aluminum fluoride contain A1F4~ and A12F9 ions as well as A1F3 and AlFe . The equilibrium constants for these complex ions are discussed later. Other workers (C86) have shown the existence of basic aluminum fluorides of 4-1- 44 formula Al2+X (OH)6F3a;, where χ varies from 0.4 to 14.0. Kleiner (K33, K34) established that AlF is ten times more stable than FeF " ", and that fluoride represses the hydrolysis of aluminum to Al(OH)3. Phase studies of ternary systems of A1F3, water, and HF, NaF, or NH4F showed formation of the solids: A1F3-3H20, A1F3-3HF-3H20, A1F3-3HF-6H20, 3NaFAlF3, HNaF4AlF3 (T10, T14, T15) and the complexes: A1F4", NH4A1F4, 2NH4FA1F3H20, and (NH4)3AlFe (N25). In Nikolaev's (N19) study of the solubility of CaF2 in hot aluminum sul­ fate solutions he isolated A12S04F412H20, Al2OF4, and more complex salts. Carter (C25) reported that aluminum fluoride is soluble to the extent of 0.559 g. per 100 ml. of saturated solution (pH 5.2), and synthetic Na3AlFe 0.041 g. per 100 ml. Nikolaev and coworkers (N20) gave the solubility of 2A1F3-5H20 as 0.127 g. per 100 g. of saturated solution at 0°, and 2.42 g. at 100°; Novoselova (N25) gave the solubility as 0.55% 9in4 8 5 water at 25°. Lacroix (L6) gave the solubility product of A1F3 as 10~ - , and of Na3AlF6 as 10~~ . Because of this relatively large solubility of A1F3, Bushey (B123) depends on precipitation of K3A1F6, which is much less soluble, in his acidimétrie method for aluminum. In the case of chromophoric complexes of aluminum with organic compounds which have been used to indicate the extent to which the aluminum is tied up by fluoride, some work has indicated that more than one complex is formed. For the lake itself formed by aluminum with alizarin, Babko (Bl) showed that varied complexes are formed, depend­ ing on the relative concentration of the two components. Parker and Goddard (Pll) found that both pH and the presence of calcium ion can affect the complex formed between aluminum and Alizarin Red S. Babko and Rychkova (B6) have shown that salicylic acid combines with alumi­ num in ratios of 1:1, 2:1, and also 3:1. The monosalicylate complex is less stable than that of copper and more stable than the ferric complex. Doubtless a similar situation exists with Eriochromcyanin R, morin, and other compounds used to measure the aluminum uncomplexed by fluoride. Brosset and Gustaver (B105), Babko and Kleiner (B4, B5), and Dodgen and Rollefson (D56, D58) found that ferric ion also forms all possible complex ions between FeF++ and FeFe in solution. Weyl and ftudow (W29) also demonstrated the formation of ferric fluoride com­ plexes through the shift toward shorter wave lengths of the absorption bands of thiocyanate, thiosalicylate, salicylate, and acetylacetone com­ plexes of iron (III). Syrokomskii and Avilov (S134) showed that fluoride 136 PHILIP J. ELVING, CHARLES A. HORTON AND HOBART H. WILLARD as well as other iron complexers like phosphate, citrate, oxalate, and tartrate change the redox potential of originally equimolar ferric-ferrous solutions. The relative changes are listed in Table V. Brosset and Gus- taver (B105) used this change in potential in a method for6 determination of fluoride. The change in potential for 0 to 12 X 10~ M fluoride is plotted in Fig. 4. Thomas and Gantz (T29) showed that fluoride is a relatively weak complexing agent for iron (III). They ranked the stability

TABLE V Change in Potential of Various Redox Pairs on Addition of Various Complexing Agents (2 to 24 MIL.)

Fe (III, II) +++ 0.02 M Ti (IV, III) VO++, v in 1 Ν H2S04, in 1 Ν H2S04, in H2S04,

Acids: Hydrofluoric 0.422 0.124 Tartaric 0.560 0.348 Citric 0.206 Oxalic 0.462 Phosphoric 0.516 0.621 Acetic 0.356 Formic 0.247 Salts (plus acid) : NH4F (1 N H2S04) 0.362 0.279 NH4F(1#HC1) 0.504 NH4F(2iVHCl) 0.524 (NH4)2S04 0.286 0.351 Reference S134 S135 S135 in the decreasing order: cyanide, citrate, oxalate, tartrate, acetate, phos­ phate, fluoride, thiocyanate, tetraborate, sulfate, chloride, bromide, nitrate, and perchlorate. Phase studies of ternary systems of FeF3, water, and HF, NaF, or KF (T12, T13, T15) showed formation of the solids: FeF3-3H20, FeF3-3HF- 3H20, 2KFFeF3H20, llKF-4FeF312H20, and 2FeF3-5NaF. The last compound is suitable for the quantitative separation of iron (III) from other soluble fluorides. Other work (N26) showed that iron is precipitated as hydroxide from a FeF3 solution, beginning at pH 3.5 and completed at pH 6.7, while if NH4F is also present, precipitation begins at pH 5.5 and is complete at pH 7.2. Correspondingly, beryllium begins to precipi­ tate from a BeF2 solution at pH 6, while a (NH4)2BeF4 solution starts to ANALYTICAL CHEMISTRY OF FLUORINE 137

20

15

CONDITIONS: ( KN0 ) - 0· 53 3 3 3 5 ( Fe|,) = 1.007 χ I0" ( Fe+'2+ ) = 0.642 χ ΙΟ" -3 ( H A) = 3.705 χ 10-3 ( Crt) = 8.01 χ ΙΟ ( F, ) = Variable

0 2 4 6 6 θ 10 12 ( F " ) χ ΙΟ

FIG. 4. Change in iron redox potential with fluoride ion. 138 PHILIP J. ELVING, CHARLES A. HORTON AND HOBART H. WILLARD precipitate at pH 7.5 to 8.0. This work, and the relative equilibrium con­ stants (K35, P62), indicate that BeF4~ is a more stable complex than the iron (III) fluoride complexes. Foley and Anderson (F46, F47) showed that iron can form more than one complex with sulfosalicylic acid ; others (H56) studied complexes with salicylaldehyde, ethylacetoacetate, and m-cresol. These studies indicate again the many equilibria possible in colorimetric methods using iron. For beryllium, Kleiner (K35)4 and Purkayastha (P62) found that all ions between BeF+ and BeF6~ exist. Ruiss and Bakina (R67) have determined the hydrolysis constant for dissociation of the BF4~ ion in water. Ryss and coworkers (R75, R78, R79) found that the hydrolysis of fluoroborate is not a simple process but that hydroxyfluoroboric acids, such as HBF3OH, form as intermediate complexes. To date, no colori­ metric methods for fluoride using beryllium or fluoroborate ions have been published. For thorium, Dodgen and Rollefson (D58) showed that ThF+++ ThF2++ and ThF3+ can exist in solution and that ThF4-4H20 is the form which precipitates. Day and Stoughton (D27, D28), by use of TTA (theonyltrifluoroacetone), cited additional evidence for the complexes ++ listed as well as for Th(N03)F++ and Th(N03)F2+. Kraus and Holmberg (K63) 6by a pH study showed that thorium hydrolyzes to ThO and Th20+ in dilute acids 1 M in NaC104. Other workers (T26, T27), in a study of the coagulating effect of thorium4 nitrate solutions on colloidale silver halides, indicated that the free Th+ ion exists only below 53 X 10~ N, while a Th complex with +3 charge occurs4 below 3 X 10~ Ν and a Th complex with +2 charge around 4 X 10~ N. The existence of such hydrolysis products and soluble fluoride complexes explain why titrations of fluoride with thorium nitrate often are not stoichiometric. Phase studies of binary systems of ThF4 and alkali fluorides (D45) indicated complexes of thorium such as: K2ThF7, KThF6, KF-3ThF4, Rb3ThF7, RbThF6, and RbF-3ThF. There was no evidence for salts of the type (M+)2ThF6 as there is in the case of silicon, titanium, iron (III), and zirconium. Wadhwani (Wl, W2) recently showed by both conductometric and indicator types of titration that the reactions between thorium ion and an alkali metal fluoride, and between thorium ion and fluorosilicic acid, are not identical. Thorium fluoride, ThF4-4H20, is formed from the simple fluoride; insoluble H2ThF6 is formed from the fluorosilicic acid. He also noted differences in the composition of the insoluble fluorides formed by addition of thorium ion to distillates from Willard-Winter separation of fluoride under varying conditions. These were probably due, at least in part, to his use of sulfuric acid in the procedure. He noted (W2) a com- ANALYTICAL CHEMISTRY OP FLUORINE 139 petition between the thorium and the fluoride ions remaining ni solution because fo the solubility product relationships for the Alizarin Red S indicator. The effect fo the hydrogen ion concentration no the indicator's dissociation also influenced the extent fo reaction fo the indicator with thorium ion. Tolerable limits ot prevent interference ni the thorium titration were given for nitrate, perchlorate, sulfite, sulfate, phosphate, arsenite, arsenate, and the other halide ions. For zirconium, Connick and McVey (C81) showed yb a study fo the two-phase equilibrium fo zirconium ion ni acidic solution and its TTA +++ H chelate ni benzene that even ni 2 M HC104 this element exists principally as Zr(OH) , while ta lower acidities Zr(OH)2" " predominates. Ion + + exchange studies (D50) indicated similar ions for solutions fo ZrOCl2. Fluoride formed complexes such sa ZrF++, ZrF2+, and ZrF3+ sa the fluoride concentration was increased (C81). tI si also well known that zirconium forms other fluorozirconate ions such sa ZrFe (F66). Table V shows the change ni redox potential fo the ferrous-ferric, titanium (III)-(IV), and vanadium (III)-(IV) equimolar solutions no addition fo various tïomplexers, including fluoride. The bleaching fo the yellow color fo peroxytitanic acid yb fluoride si employed ni a colorimetric method subsequently discussed. No colorimetric methods depending no the weak complexes formed with vanadium have been published. ä A study with radioactiv3 e cerium (M54) has shown that, below 10~ M Ce+++ and 7 X 10~ M HF, cerous fluoride does not precipitate an++d 20% of the cerium (III) in solution is associated with fluoride as CeF . The equilibrium constant and a calculated solubility product for CeF3 are given in Table VI. Although cerium (III) has not been used in colori­ metric methods for fluoride, it is widely used in the volumetric methods for fluoride (see Section V-B-4). The equilibrium constants for the fluoride complexes of several metals and hydrogen ion are given in Table VI. Fernelius (F31) discusses some of the problems related to such complexes. Comparison of the values for the monofluoro complexes gives the following decreasing order of stabil­ ity: Th, Al, Be, Zr, Fe(III), and H+. Similar considerations apply for the difluoro and trifluoro complexes. This fact alone should indicate a prefer­ ence for colorimetric methods using thorium or aluminum rather than zirconium or iron, providing other factors are neglected. These factors include the number of interfering ions, the sensitivity of the colorimetric reagents, and the influence of other equilibria, principally hydrolysis of the metal ion and association of the fluoride as the weak acid, hydro­ fluoric acid. The constants also explain the reason why empirical plots of absorbance versus fluoride concentration are necessary for most of the colorimetric methods. 140 PHILIP

TABLE VI Equilibrium Constants for Fluoride Complexes of Several Metal Ions J.ELVING,CHARLESAHORTONANDHOBART

K33 Β103 D56 D56 Reference Β104 D23 R56 K34 Β106 Β105 B 4 D58 D58 C81 K35 M54 3 3 3 3 3 + + +2 Metal ion H+ H+ H+ A1+ A1+ Fe+ Fe+ Fe+ Th < Zr < Be Ce

\ogki -0 -0.18182 23.14 3.148 8 3.623 .6200 6.36.322 6.16.13 3 5.1552 .15 25.30 15 .301 55.27 .276b 77.6D. 6533 5τ.oU. 500U 5.88o.8S6 b 4.0 4 3 .952 5. ,747 4. 322 4.943 log &2 0.591 0.672 1 .519 5.31 5.02 3 .929 .398 log k z 3.82 3.85 3 .222 2 .70 4. 440 2 .826 3.561 log k A 2.14 2.74 2 .00 1.991 log k b -0.43 1.63 0 .357 log k 6 0.4 log K 0.409 3.820 4 .139 19.77 16.38 t 8 .771 logK..p. 8.77 1 6.835 Condition 1 2 2 3 4 5 3 6 6 7 3

Notes on conditions: Form of the constants+: + 1: 0.53 M NH4NO3, 25° Jbi = (A£+*F-)/(M *)(F-) + 2: Infinite dilution, 25° k2 = (A/ *(F-)2)/(M+*F-)(F-) 3: 0.1 Μ HN0 kt = (Jlf «(F-).)/(Af+«(F-).)(F-) 3 +2 +2 4: 0.53 M KN03 or NH4NO3 Kt = ki - k2 - - • k n H.WILLARD 5: 0.53 M KN03, 0.01 M HCi KB.V. = (ThF2)(HF)V(H ) 6: 0.4 M NaC104, 0.1 M HCIO4 7: 2 M HCIO4 ANALYTICAL CHEMISTRY OF FLUORINE 141

Colorimetric Methods Using Titanium. Steiger (SI 16) and Merwin (M63) developed a colorimetric method depending on the bleaching by fluoride of the yellow color of peroxytitanic acid which is widely used in various modified forms today. The reactions involved (M103) probably are: TiO++ + H202 + H20 = TiO — + 4H+ (1) TiO++ + 6F- + 2H+ = TiFe— + H20 (2) Equilibrium 2 lies far to the right, thus sending equilibrium 1 to the left, and thereby decreasing the concentration and color of the peroxytitanic ion. Babko and Volkova (B7) showed that the yellow 4complex between titanium and peroxide formed below pH 3 is Ti(H202)+ , with maximum absorption of ligh4 t at 465 đŔě++. Above pH 3 a colorless 1:1 complex forms, 4 either Ti(H02) ++ or Ti(02) . The dissociation constant of the colored complex is 0.9 X 10~ . Probably, the presence of two interrelated equilib­ ria account for the fact that the bleaching is only proportional to the fluoride concentration over a limited range (B38, SI 16, M103, W42). Dahle and Wichmann (W42) showed how important control of pH is for good results; the optimum bleaching occurs at pH 1.5. A recent study by Kleiner (K36) showed that the stability constant of6 the monofluoro complex of titanium in 0.1 Ν nitric acid is 2.6 × 10 for the reaction TiO++ + F~ = TiOF+. Various studies have shown that it is better to measure the color photometrically rather than visually. Parrish and coworkers (P14) measured at 440 đŔě, where bleaching was reproducible to 0.003% F~ and accurate below 2% to ±0.01%; Ovenston and Parker (016) used a violet Ilford No. 801 filter; Bendig and Hirschmuller (B38) used a 546-đŔě green mercury light in preference to a sodium light; and Monnier, Vaucher, and Wenger (M103) favored either the Ilford 801 or a Hilger H505 filter with a tungsten lamp. The last workers found that the optical density was an inverse linear function of fluoride concentration between 0.6 and 3.2 mg. per liter within ±0.01 mg. per liter. Alkali sulfates cause additional bleaching, while aluminum and ferric ions inhibit bleaching, since they form more stable complexes with fluoride. Phosphate causes some bleaching by forming a titanium complex itself (SI 16, M63). However, at a lower pH of 0.9 sulfate does not inter­ fere and small amounts of phosphate interfere to a negligible extent (M104). Sulfate and phosphate interfere more seriously at pH 1.5 or higher. Hill and Reynolds (H64) were able to analyze monofluorophos- phates using this colorimetric method in a more acidic medium. Monnier and coworkers (M104) and Dahle (D6) preferred to eliminate the effect of phosphate by measuring the color before and after addition of alumi­ num. The aluminum ion eliminated the effect of fluoride on the titanium color. Nitrite does not affect the method (K57). 142 PHILIP J. ELVING, CHARLES A. HORTON AND HOBART H. WILLARD

The work of Fenton (F30, M107) has suggested the only other colori­ metric method for fluoride using titanium4 . He found that dihydroxy- maleic acid gives a red color with Ti+ which is more sensitive to titanium ion than the peroxide yellow color. All published papers which have used the colorimetric methods with titanium or the other colorimetric methods subsequently discussed are listed in Table VII. Colorimetric Methods Using Zirconium. The most widely used colori­ metric method for fluoride in water is one using zirconium with Alizarin Red S (sodium salt of l,2-dihydroxyanthraquinone-3-sulfonic acid). Besides forming a colored lake (or complex) with this compound, zir­ conium ion changes the color of other 1,2-, 1,4-, 2,3-, 2,6-, or 2,7-dihy- droxy-; 1,2,4-trihydroxy-; and 1,4,5,8-tetrahydroxyanthraquinones (C50, W24). Many of these types of compounds have been investigated as indicators for colorimetric methods for fluoride (see Table VII). An Alizarin Red S indicator was first used by DeBoer (B69) for the volumetric determination of zirconium using an alkali fluoride as the titrant. As a result, the colorimetric reagent consisting of zirconium and Alizarin Red S is usually known as DeBoer's reagent. Such a reagent is also often used in solution or on test papers to detect fluoride (see Sec­ tion IV). Recent work (A36, B114) has indicated that lakes prepared in H2S04 are more stable than those prepared in HC1 or mixtures of these two acids. All these acids have been used in preparing the lake reagent and in the final solutions. The acidity in the lake reagent and in the final samples for color comparison have run from 1 to 12 Ν and 0.01 to 12 N, respec­ tively. The molar ratio of Alizarin Red S to zirconium has varied con­ siderably from the lowest ratio of 0.0028 (B70) to the highest ratio of 0.94 (F9). A ratio of 0.15 to 0.16 has been used most (C34, D30, D74, Lll, S100, etc.). Most of the reagent lakes are aged for 1 hour to several days before use. The color comparisons using zirconium-Alizarin Red S have been made both visually and photometrically. A recent work (A36) measures the lake absorption at 520 đŔě and the free dye absorption at 420 đŔě to obtain the fluoride concentration. Some other workers (B114) measure the color only at 525 đŔě in a 50-mm. cell. A linear curve was obtained for the 0- to 100^g. range. Although several workers have had success with photometric methods, Richter (R33) claims that the photometric pro­ cedure is not very successful. In general, the color change with fluoride concentration does not follow Beer's law (W67). Besides sulfate, which interferes somewhat, traces of phosphate, aluminum, ferric (W6), arsenious, arsenate (S66), thiosulfate, oxalate ANALYTICAL CHEMISTRY OF FLUORINE 143

TABLE VII References for Colorimetric Fluoride Methods Complex With or As Metal Ion Used References Titanium Peroxytitanic acid A3, A12, B38, B74, B122, D2, D4, D6, D7, D8, D12, D26, D73, F29, F60, F73, H2, H64, K14, K58, L18, M63, M103, M104, 016, P14, P36, RI, R7, S34, S47, S60, S65, S101, S116, S140, W15, W31, W36, W42, W43, Y7 Dihydroxymaleic acid F30, M107 Thymol M107 Zirconium Alizarin Red S (1,2-dihydroxy- A10, All, A16, A18, A27, anthraquinone-3-sulfonic acid, A36, B23, B45, B68, B69, sodium salt) B70, B77, B86, C30, C31, C34, C35, C36, C42, C50, C51, C60, C68, C90, D29, D74, E22, Fl, F3, F9, F10, F45, G4, G68, H3, H21, H42, J19, K62, L10, LU, L45, M44, M83, M95, M97, M123, R4, R5, R6, R7, R73, S6, S7, S38, S39, S40, S65, S66, S89, S90, S91, S100, S109, ΤΙ, T23, T34, T39, T55, T56, W6, W7, W67, W68, W69 Alizarin Red S, indirect R33 Purpurin All, B69, F73, G50, G51, J4, K54, K61, M35, S66 Purpurin & Solway Green G.S. R71 Quinizarin B69, R71 Quinalizarin R71, S96 2-Hydroxyanthraquinone C50 Rufigallic acid B69 Alizarin Blue S B69 Hystazarin B69 Purpoxanthin B69 Anthrapurpurin S66 Flavopurpurin S66 Anthragallol S66 1,2,4,5,8-Alizari ncyanine S66 Cerulin S66 2-Hydroxy-5-methylazobenzene- 4-sulfonic acid K78 144 PHILIP J. ELVING, CHARLES A. HORTON AND HOBART H. WILtARD

TABLE VII (Continued) Complex With or As Metal Ion Used References p-Dimethylaminoazophenyl- arsonic acid B122, H22, R2, R3 Hemotoxylin G4, J8, J9 Iron (III) Thiocyanate A32, B4, B5, C53, D2, F49, G64, 14, K70, M7, 05, S46, S90, S97 Salicylate A32, Β122, G65, K60, R34, R35, VI1 Sulfosalicylate L7, M102 Acetylacetonate A32, W45 8-Hydroxyquinolate A32 "Ferron" F6, U5, W2, Y9 Ferrocyanide K70 Fluoromethemoglobin F4 Bromide E28 Chloride F50 Cresotinic acid B122 Aluminum 8-Hydroxyquinolate C69 Hemotoxylin 06, 07, 08 Eriochromcyanin R G37 (vol.), T37 " Aluminon" C71, C72, F16, 014, R32 Thorium Alizarin Red S B39, C71, K29, M3, M62, S5, T21, T23, Y9 Chrome Azurol S R22 "Thoron," (1-arsenophenylazo)- 2-naphthol-3,6-disulfonic acid H71 Molybdenum (Blue) Complexes Silicate B43, Cl, M56, P13 Phosphate B43 Arsenate B43 Benzidine R15

(B69), and sulfide (S6) ions interfere seriously in this method. Metaphos- phate interferes more seriously than orthophosphate (T23). Chloride, bicarbonate (L10), manganous, ferrous, silicate, alkali metal (L10), calcium, magnesium (S6), and borate (H3) ions interfere very slightly or not at all. In the analysis of sea water (T34), however, NaCl and MgS04 have been added to the standards to compensate for the slight effects caused by these compounds in the water. Millner and Kunos (M83) removed phosphate by a preliminary treatment with silver, followed by chloride to remove excess silver, before using the zirconium-Alizarin Red S method. Because of the stability of the zirconium fluoride formed in this ANALYTICAL CHEMISTRY OF FLUORINE 145 colorimetric method, it can be used for detection or estimation of fluoride in such compounds as CaF2, K2SiF6, K2BeF4, PbF2, and PbCIF, which are often intractable by other methods. Scott (S38, S39) has found it possible to distinguish a difference of ±0.05 p.p.m. F~ in water at the 1 p.p.m. level with this method. This gives a representative idea of the precision of this colorimetric method. Stone (S124) and Koone (K56) reported that alizarin is a more sensi­ tive reagent for fluoride with zirconium than Alizarin Red S. The former also noted the interference of oxidizing agents like chlorate, bromate, and iodate due to destruction of the dye. Other workers, such as Jaki (J4) and Kolthoff and Stansby (K54, F73) favor purpurin lakes with zirconium. They claim that this lake has greater stability. This reagent can be used to determine less than 3 Mg. F~ (K54) with a probable precision of ±0.2 Mg. and a precision of ±0.6 Mg. at the 3- and 200-Mg. levels, respectively (J4). Gad and Naumann (G4) and others prefer hemotoxylin black as the chromophoric dye and can determine 0.1 mg. F~ per liter in 50-ml. samples of water. Harrold and Hurlburt (H22) developed a field test for gaseous fluorides based on the change in color of the zirconium complex of p-dimethylaminoazophenylarsonic acid on paper. Radimer, Smiley, and Lafferty (B122, R2, R3) improved this latter method by extracting the free red dye in acetone and measuring the color of the extract. This procedure avoids the interference of the color of the brown metal-dye complex. At 2 p.p.m. F~ in air, their precision was <0.2 p.p.m. The method is applicable over the range of 1 to 500 p.p.m. by volume. Am­ monia interfered in the method, but N20, Cl2, C2H2, H2S, Br2, and C02 did not. It can be used for F2 which hydrolyzes to HF in the reagent which is moistened with acid. Numerous other dyes which have been used with zirconium are listed in Table VII. Colorimetric Methods Using Ferric Ion. Smitt (S97) was the first to adapt GreefTs (G64) volumetric method using ferric chloride titrant and thiocyanate as indicator for use as a colorimetric method. Foster (F49, F50) first studied the ferric thiocyanate method extensively. He found that the acidity (pH 4.5), temperature, and concentrations used must be carefully controlled. The range of his method was 25 Mg. to 45 mg. F; excess sulfate, chloride, and nitrate as well as the expected ions, such as phosphate, interfered. Early comparative studies (D2, S90) indicated that results using this method compare favorably with the peroxytitanic colorimetric or Willard-Winter (W52) volumetric methods, providing the salt content of the samples was not too high. Chemodanova (C53) modi­ fied the method and used a photophotometer for the automatic deter­ mination of 1 to 15 Mg- of HF per liter of air. He improved the color 146 PHILIP J. ELVING, CHARLES A. HORTON AND HOBART H. WILLARD

stability of the ferric thiocyanate by the addition of K2S208. Hydrogen chloride or nitrogen oxides only bleached the solution very slightly, and S02 could be tolerated up to ten times the HF concentration, but fluoro­ silicate, sulfate, and phosphate interfered seriously. Okuno (05) analyzed water using pH 2.8 rather than pH 4.5. Babko and Kleiner (B5, B4) studied the equilibria involved in the method and used it to establish stability constants of the ferric fluoride complexes. The most recent work on the ferric thiocyanate method was reported by Ingols, Shaw, Eberhardt, and Hildebrand (14). They used pH 1.9 to 2.0 and decreased the effect of interferences in water by using a blank to which ZrOCl2 was added in addition to the regular reagents. The zirco­ nium restored the color which had been bleached due to fluoride but did not appreciably affect the bleaching. due to the interferences. They measured the color at 490 đŔě. However, in spite of considerable work, even the reactions involved in the thiocyanate method for iron have not been completely elucidated, and the method, as Babko stated (B3), "is characterized by the low stability of the ferric-fluoride complex and a weak color," so that it cannot be recommended very highly. About the same time, Armstrong (A32) developed a method based on bleaching of the ferric acetylacetonate complex. He found that this re­ quired less critical control of the experimental conditions than did thio­ cyanate, salicylic acid, or 8-hydroxyquinoline. More than 50 mg. NaCl, 100 mg. Na2S04, or 400 mg. NaN03 affected the color, making the method somewhat less sensitive to these compounds than the thiocyanate method. Wilcox (W45) studied the method photometrically; as expected, bleaching did not follow Beer's law. Somewhat later, Greenspan and Stein (B122, G65, R34, VI1) used ferric salicylate for the determination of 1 to 140 Mg. F per ml. Acidity was controlled at pH 3.1 with hydrochloric acid, and absorbency was measured at 530 đŔě. Alkali metal, Ni, Cu, Ag, Hg, Mn(II), Cd, Ca, and Sr metal ions and halide, nitrate, and acetate anions had practically no effect. Interferences included Co, Bi, As, Sb, and Al metal ions, and tartrate, citrate, phthalate, borate, silicate, oxalate, sulfate, and phos­ phate anions. Accuracy was ± 1 Mg- for 0 to 35 Mg- F per milliliter for 5-ml. aliquots. Determinations are possible in the presence of ethanol, glycol, dioxane, pyridine, and small amounts of uranyl ion. Kortum- Seiler (K60) developed a similar method but used a pH of 2.7 and meas­ ured the absorbency at 550 đŔě. Rickard, Ball, and Harris (R34) used the method on organic compounds decomposed at 1100° in a stream of moist oxygen. Their overall precision was ±3.9% of the fluorine present for a single determination in the 0.2- to 1.5-mg. range. Monnier, Rusconi, and Wenger (M102), and Lacroix and Labalade (L7) used 5-sulfosalicylic acid as the chromophore with ferric ion for ANALYTICAL CHEMISTRY OP FLUORINE 147 colorimetric determination of fluoride. The former workers used pH 3.0; the latter used pH 2.9, controlled by monochloroacetate buffer, and measured the bleaching at 520 m/z. Other methods using ferric ion are listed in Table VII. Kullberg (K70) developed one of the few methods for fluoride based on an increase in color with F concentration. He used a solution of ferric ferrocyanide and could detect 0.03 Mg. F in a drop of water. Fabre and Bazille (F4) determined fluoride in aqueous solution by forming fluoromethemoglobin from methemoglobin and measuring its absorbency at 610 to 620 đŔě. The method is not very sensitive. Colorimetric Methods Using Aluminum. Chariot (C49) in his excellent discussion of chromophoric organic reagents for aluminum calls attention to the adverse effect of fluoride on many of the systems. In 1938 Oshero- vich (014) developed the first colorimetric method using aluminum. He used the well known "Aluminon" or aurintricarboxylic acid as the chromophoric chelating agent and an ammonium chloride-ammonium acetate-ammonium carbonate medium. This method was investigated further by Clifford (C71, C72), using a 10-cm. cell and measuring the absorbency at 524 τημ. Interferences included phosphate, nitrate, nitrite, sulfite, free chlorine, and reducing agents, but sulfate did not interfere. Okuno (06, 07, 08) used hemotoxylin with aluminum in a visual colori­ metric method for Japanese waters. Presence of fluoride changes the color from the purple of the aluminum-hemotoxylin complex to brown, reddish brown, yellowish brown and, finally, yellow, the color of the free hemo­ toxylin. By use of three different reagent concentrations, each added in a stepwise fashion, he was able to detect differences of 2 Mg., 5 to 10 Mg-, and 20 Mg- F for the ranges of 1 to 10, 0 to 100, and 160 to 400 Mg- Okuno's order of increasing sensitivity for various colorimetric methods used visually was: peroxytitanic acid, ferric thiocyanate, zirconium-Alizarin Red S, and aluminum-hemotoxylin. Richter (R32) and later Thrun (T37) used Eriochromcyanin R (sodium salt of o'-sulfohydroxydimethyl- fuchsondicarboxylic acid). The former used pH 3.8; the latter, pH 5.5, and visual comparison of colors. Most accurate results were obtained on waters after a Willard-Winter separation for the 0.1- to 6-p;p.m. range. Gilbert and Saylor (G37) have used the same compound as the indicator in a volumetric method using aluminum chloride as the titrant. Clewett (C69).has attempted to use 8-hydroxyquinoline in a colorimetric method; however, the fluorometric method subsequently discussed which uses this reagent is much more sensitive for the determination of traces of fluoride.

2. Miscellaneous Colorimetric Methods Methods Using Thorium. Clifford (C71) in 1940 made a photometric study of the color change of the thorium-Alizarin Red S complex in con- 148 PHILIP J. ELVING, CHARLES A. HORTON AND HOBART H. WILLARD nection with a study of the titration of fluoride using thorium as the titrant. The end point color change was found to be sluggish, indicating some equilibrium between thorium and the dye as well as with fluoride. Van der Merwe (M62) was first to use color comparison of thorium- Alizarin Red S lakes at pH 2.6 with fluoride present. As the fluoride content increased, the color changed from the red of the thorium-dye complex (or lake) to rose, pink, pinkish buff, and finally yellow. The yellow color is due to the absorption by the free Alizarin Red S. Mc- Clendon and Foster (M3) made a spectrophotometric comparison of the absorption of the thorium complex and of the dye in basic and acidic media. The similarity of the spectrum in alkali to that of the thorium complex in acidic medium suggested to them that thorium either changes the ionization constant of the dye or the ionic strength of the solution. If the lake is a true complex, the structure is: OH

\Th+/4

and the bond of thorium to the 1-oxygen atom is somewhat ionic, as it would be in the rose or pink sodium salt in alkaline solution. Talvitie (T4) and, later, Salsbury, Cole, Overholser, Armstrong, and Yoe (S5, Y9) also used the thorium-Alizarin Red S system in visual colorimetric methods. They used a formic acid-sodium formate buffer to maintain a pH of 3.5 and allowed 30 minutes for equilibrium. The method is applicable for over 0.1 p.p.m. F and has a precision of about 5 % of the amount present. The fluoride equivalent to 1 mg. of several common ions is: sulfate, +0.02 mg.; Ca, Mg, chloride, —0.01 mg.; nitrate -0.005 mg.; Na, K, 0.000 mg. Phosphate or aluminum ions naturally interfere seri­ ously. The temperature was regulated within one degree by Talvitie (T4). Revinson and Harley (R22) prefer Chrome Azurol S to Alizarin Red S. They use an o-toluidine-perchloric acid buffer of pH 4.0 in a medium 4 M in NaC104 and measure the absorbency at 605 đŔě. Gum arabic is used to stabilize the thorium lake. The range of their procedure is 0 to 80 Mg. F. Horton, Thomason, and Miller (H71) measure the color decrease of the thorium-"thoron" ((l-arsenophenylazo)-2-naphthol-3,6-disulfonic acid) complex. The fluoride was separated by a modified Willard-Winter technic using perchloric acid, the distillate acidified to about pH 2, iodide ANALYTICAL CHEMISTRY OF FLUORINE 149 added to reduce chlorine, hydrazine hydrate to reduce iodine, then thorium solution (120 Mg.) and thoron reagent. The optical density was measured at 545 πΐμ. The range was 0 to 50 Mg. F with an overall accuracy of 4%, although a calibration curve is necessary, since the bleaching did not follow Beer's law. Excessive acid in the distillate gave high results and the usual extraneous ions interfered. Flagg and McCarty (F41) found by a spectrophotometric study that alizarin proper forms a one to one complex or lake with thorium at pH 4 to 4.2 in an alcoholic medium. Methods Using Molybdenum. The first method (M58) suggested de­ pended on the reaction of the silicon in the fluorosilicic acid distilled. The latter was reacted with ammonium molybdate, and the complex reduced to the blue complex by the usual technic and the color measured photo­ metrically. Cade (CI) developed a similar method for determination of H2SiFe in HF. He could determine 0.005 to 0.3% in HF with an error of ±0.003%. Most of the HF was volatilized from a hydrochloric solution and the remainder complexed with boric acid before forming the "molyb­ denum blue." An empirical calibration curve was necessary. Parri (P13) noted that fluoride changes the yellow color of acetone solutions of molybdosilicate or molybdophosphate. He used such a solu­ tion for titration of fluoride solutions as well as suggesting their colori­ metric possibilities. Bergamini (B43) studied photometrically the masking by fluoride of the blue reduced molybdenum complexes with silicon, arsenic, and phosphorus. The absorption of these complexes is maximum at about 470 đŔě, but fluoride changes most the absorption at 700 raju. The same bleaching occurred for 1.9 mg. F using silicon, 3.75 mg. F using arsenic, and 8.55 mg. using phosphorus. Thus, the effect of fluoride is greatest on the molybdosilicate blue complex, and it was recommended for use in the determination of fluoride. Redox Methods. Elemental fluorine liberates chlorine, bromine, or iodine from the respective solid alkali halides. The amount of the halogen formed may be determined colorimetrically. These methods are discussed in another section. 3. Fluorometric Methods Surprisingly little work has been done on the detection and estimation of fluoride using fluorescence methods. Goto (G58) described three fluorescence methods for the qualitative detection of fluorine. One method depended on the bleaching by fluoride of the strong yellow-green fluores­ cence of the zirconium-morin complex in H CI solution. The others were based on the restoration of the violet fluorescence of salicylic acid in acetate solution, which had been bleached by ferric or titanium ions, due 150 PHILIP J. ELVING, CHARLES A. HORTON AND HOBART H. WILLARD

to formation of fluoride complexes of these two metal ions. Feigl and Heisig (F16) developed a qualitative test based on the quenching of the fluorescence of paper treated with aluminum 8-hydroxyquinolate. Okac (03) used for large amounts of fluoride the quenching of the fluorescence of the aluminum-morin complex in a visual volumetric method with aluminum chloride as the titrant. Recently, Willard and Horton (H72, W49) published a quantitative method for microgram quantities of fluoride based on the same system. Spence, Straetz, Krause, and Watters (B122, S105) developed a very similar method. The results and conditions used for these two independently developed methods are compared in Table VIII. Both are based on the measurement of the decrease in fluorescence of the aluminum-morin complex in a semi- alcoholic solution. Spence and coworkers also investigated the use of maclurin (pentahydroxybenzophenone) instead of morin in this method. Willard and Horton (H72, W49) also developed a quantitative method for traces of fluorine based on the fluorescence of uncomplexed aluminum extracted in chloroform as the 8-hydroxyquinolate from a solution buffered at pH 4.7. They also investigated other systems which proved to be unsatisfactory. Willard and Horton (H72, W47, W48) also developed a photofluoro- metric titration method for larger amounts of fluoride using thorium as titrant and morin or quercetin as indicator in 50% alcohol. They used a monochloroacetate buffer whose pH is 3.0 in water and 4.9 in 50% alcohol. The range of their method was from 0.04 mg. to over 40 mg. F. The precision for a single determination at the 95% confidence level was 1.3% of the value at 1.0 mg., 0.6% at 15 mg., and 0.5% at 30 mg. fluoride. The method is suitable for determination of fluoride in organic compounds decomposed by alkali, or it can be used following a Willard-Winter (W52) separation. 4- Nephelometric Methods Relatively little use has been made of the possibilities of determining fluoride ion nephelometrically. However, the well-known tendency of calcium fluoride to form as a fine precipitate or colloidal suspension has been used as the basis for such a method. Apparently originally suggested by Peterson (P31), the method was first extensively studied by Stevens (S121), who precipitated colloidal calcium fluoride in 25% ethyl alcohol solution. Gelatin is added as a protective colloid to stabilize the suspen­ sion. The accuracy is ±1%; arsenate, phosphate, silica, and sulfate interfere and must be removed. For samples requiring a carbonate fusion, the lower limit is 0.3% F; smaller amounts may be determined in water-soluble material. ANALYTICAL CHEMISTRY OF FLUORINE 151

TABLE VIII Comparison of Two Fluorometric Methods for Traces of Fluoride

Workers Willard, Horton Spence et al. References H72, W49 B122, S105 Instrument Klett, later model Klett, early model Primary filter Corning 5850 Corning 5583 Secondary filter Lumetron B530 or others Corning 3385 Exciting wave length, ταμ 436 Maximum fluorescence, m μ 500 Cell type and volume Test tube, 10 ml. Curvette, 1 or 10 ml. Total final volume 100 ml. 1 or 10 ml. Alcohol concentration 50 % 60 % Equilibrium time allowed 20 min. 20 min. Range of method 54 Mg. Al, 0.1-25 Mg. F 10 ml., 0.02-0.50 gM. F 27 Mg. Al, 0.03-18 Mg. F 1 ml., 0.005-0.06 Mg. F Precision 54 Mg. Al, ±0.2 to ±0.6 10 ml., ±0.02 Mg. F Mg. F 27 Mg. Al, ±0.02 to 1 ml., ±0.01 gM. F ±1.0 Mg. F Type of curve Nonlinear Nonlinear Interferences Similar for both procedures

A rather lengthy and complicated procedure for fluoride is based on comparison of the colloidal lead sulfide derived from lead fluoride forma­ tion with that obtained from a standard lead solution (G18). The Pb:F ratio has to be determined empirically with known quantities of fluorine and varies with the amount of fluorine. The procedure is applicable to the determination of 0.05 to 2 mg. fluorine. Procedures for applying the method to fluorine in water, carbonates and silicates, and vegetable and animal matter are given by Snell (S98). Giammarino (G35) described a nephelometric method based on the formation of colloidal lanthanum fluoride for samples containing 0.03 to 0.05% fluoride. Chepelevetskii (C57) suggested the titrimetric determination of fluoride by titration to maximum turbidity, using a photoelectric photom­ eter. Proposed titrants were thorium in the presence of alizarin and lead in the presence of excess sodium chloride and alcohol. The variation using thorium is probably really photometric detection of the end point of the usual thorium titration with alizarin or Alizarin Red S indicator as pro­ posed by Nichols and Kindt (N6), although it is impossible to determine definitely from a translation of the original paper (C57). Preliminary unpublished work at the University of Michigan (W50) has indicated that it may be possible to titrate fluoride in a buffered aqueous solution nephelometrically with thorium fluoride without the 152 PHILIP J. ELVING, CHARLES A. HORTON AND HOBART H. WILLARD use of an indicator. The turbidity of such a solution increases almost to the stoichiometric end point and then decreases with an excess of thorium. The thorium fluoride formed during the titration increases turbidity if it is kept suspended by stirring, but it is coagulated, so that it scatters less light, by the excess +4-charged heavy-metal thorium ion. Such coagula­ tion of a sol by a multicharged ion is common in colloid chemistry studies. Unfortunately, insufficient work has been done on the method and the effect of extraneous ions to warrant publication. The investigators obtained a V-shaped curve which required extrapolation for the end point; the effect was decreased in the presence of alcohol or other sol coagulants.

5. Emission Spectroscopy The detection and determination of fluorine by emission spectroscopy has until recently not been too satisfactory owing to the high ionization potential of fluorine and the presence of the principal spectral lines in the far vacuum ultraviolet region. The raies ultimes or principal lines of fluorine are all at wave lengths below 1000 A. However, other sensitive lines of fluorine may be excited in the region usually employed in spectro- chemical analysis (2000 to 10,000 A). Some of these lines are excited in spark sources ; others in the hollow cathode discharge. The most sensitive fluorine line in the photographic region is that at 6856 A due to singly ionized fluorine. Analytically useful spectrographic methods for fluorine have been based on the molecular band spectra produced by alkaline earth metal fluorides, particularly that of calcium, when excited in the D.C. arc. The demands of atomic energy research have resulted in the development of spectrographic methods for the determination of as little as 1 p.p.m. fluorine present in 0.04 μg. quantity. The nature of the fluorine spectra and some of the previous attempts at emission spectrographic methods for fluorine are summarized by McNally, Harrison, and Rowe (M29). The latter found a practical sensi­ tivity of 300 p.p.m. for the determination of fluorine in the vacuum ultra­ violet region using a hot-spark source. Mitchell (M89) has reviewed the spectrographic determination of fluorine in soils, plants, and related materials. For convenience, the work on the determination of fluorine by emission spectroscopy is discussed according to the type of excitation employed. D.C. Arc. In 1920 Mott (M117) noted that fluorine compounds can be quickly identified in the D.C. arc by adding calcium carbonate and observing the bright-green bands characteristic of calcium fluoride. Since then, many others have used this basic technic. Datta (D22) was apparently the first to investigate systematically the ANALYTICAL CHEMISTRY OF FLUORINE 153 band spectra of the alkaline earth metal fluorides originating in the elec­ trical arc. There are no fluorine lines or bands in the normal arc discharge. However, if the fluorine is combined with beryllium, magnesium, calcium, strontium, or barium, diatomic molecules are formed in the arc or spark discharge which show spectral bands. Ahrens (A5, A6, A7, A8, A9, S44) has carried on extensive studies with the D.C. arc procedure, particularly in reference to the analysis of rocks, minerals, and soils. In his later work he has used the calcium fluoride bands extensively. To obtain optimum sensitivity, calcium should be present in the arc column in excess. For quantitative work, Ahrens used a CaO band as an internal standard and determined the intensity ratio, CaF/CaO (S44). In one sample of granite, with 12 replicates, the mean reported was 0.047% F with a standard deviation of 11% of the amount found (A8). The CaF2 band with a head at 5291 A (due to CaF emitter) is the most generally used band, since it is intense enough to detect 2.5 Mg. F; the most sensitive band is the CaF band with a head at 6064 A (S44). Although the latter is eight times more sensitive than the head of CaF (5291), it is subject to excessive interference from CaO bands. The sample (finely powdered if a solid, or a dried residue if a solution) is mixed with calcium oxide or a calcium salt, if there is no calcium in the sample, and with charcoal, if a phosphate rock. The mixture is placed in the cup of a graphite electrode and, e.g., given a 30-second exposure at 0.5 ampere for a 5-mm. arc gap (A6). For example, solutions containing 0.001 to 0.005 mg. F as NaF and 5 mg. of calcium acetate are evaporated in the electrode cup; in water analysis, a residue containing 0.05 to 1.5% F is added to an electrode impregnated with calcium chloride (B38). At one time Ahrens (A6) recommended the SrF emitter band head because of its greater sensitivity, but owing to interference by cyanogen bands he now prefers the calcium band head. The BaF band with head at 4992 A has a bad background. Fluorine in minerals, bones, teeth, and similar refractory material may be concentrated by the addition of sulfuric acid and distillation of HF; the latter is absorbed in calcium hydroxide solution, and the dried calcium fluoride is added to the electrode (P3). Owing to the rapidity with which fluorine is volatilized, it is advisable to use fast photographic plates. Ryde and Yates (R74) determined fluorine in glasses by adding calcium oxide to the powdered sample, exciting in a D.C. arc, and observ­ ing the CaF2 band head at 5291 A. Papish, Hoag, and Snee (P3) excited the ash of bones and teeth in a D.C. arc with graphite electrodes and a current of 8 amperes; they reported a sensitivity of 10 Mg- F with this procedure. Petrey (P33) analyzed water using a weighed portion of the residue after evaporation in a graphite electrode impregnated with a 154 PHILIP J. ELVING, CHARLES A. HORTON AND HOBART H. WILLARD

calcium solution. When excited at 15 amperes, a sensitivity of 0.05 p.p.m. of fluoride on the original water basis was reported. Others who have used the D.C. arc and the CaF2 band at 5291 A in various materials include Churchill (C62), Braidech and Emery (B89), Boissevain and Drea (B75), Glock, Lowater, and Murray (G45), and Lowater and Murray (L52). Data on fluoride band spectra for analytical and other purposes are also given elsewhere (F75, F76, M108, P24, P58, Y2). A.C. Spark. Johnson and Norman (J13), using a spark in air with no inductance with the powdered sample in a cupped graphite electrode, found that the fluorine lines at 6856 and 6902 A gave a sensitivity limit of 200 Mg. fluoride. De la Roche (R41) used a spark discharge between aluminum electrodes in a closed system with F2, BF3, or SiF4 in nitrogen. Sventitskii (S132) reported detection of fluoride in aqueous solutions with a rotating electrode arrangement. With a condensed low-voltage spark Pfeilsticker (pp. 107-108 of K40, P37, P38) was able to detect fluorine lines at 4025.5 and 3847.1/50.0 A from NH4F on a copper elec­ trode. CaF2 in powdered samples of the electrolytes of alumina-cryolite baths has been determined with a Feussner spark using the line pair Ca 3158.87-Na 3302.32/98 A and a rotating electrode; 60 to 80 samples can be analyzed in 8 hours (E24). By means of high-temperature arc excitation, fluorine could be deter­ mined in powdered ores and minerals or dried solution residues deposited on copper electrodes, using the F lines 6856.0, 6239.6, and 6348.5 A (B82). Hollow Cathode Discharge Tube. McNally, Harrison, and Rowe (M29) used a stainless steel hollow cathode source in a tube containing helium at 10 mm. mercury pressure and a procedure which involved 15 minutes or less per sample. With a low dispersion glass prism spectrograph and 800 volts across the sample tube at 0.2 ampere, the limiting sensitivity with the 6856-A line was 0.01 Mg. F (10-mg. sample containing 1 p.p.m. F as NaF in a metal oxide matrix). The greatest concentrational sensitivity was 0.1 p.p.m. F. The reproducibility on eight 20-mg. samples (30 p.p.m. F) was ±15%, while the average deviation of the line intensity ratio, F/He, was 6%. As little as 1 p.p.m. of fluorine in uranium metal could be detected (B122). Gillieson and Newcombe (G40) also used the hollow cathode to deter­ mine organically bound fluorine; they reported a sensitivity of 5 Mg- F. Electrodeless Discharge. Gatterer (G12) found that the fluorine spec­ trum itself can be excited by placing a tube of glass with a high melting point, containing 10 to 20 mg. of sample and evacuated to 1 M> in the high- frequency magnetic field of a coil. Intense spectral emission is obtained ANALYTICAL CHEMISTRY OF FLUORINE 155

in the region of 4500 to 7000 A. The original powdered sample can be used, and as little as 0.01% fluorine can be determined with 10% accuracy; the sensitivity is 0.001%. Large amounts of other constituents do not interfere or decrease the sensitivity. Indirect Determination. Paul and Karreth (P18, P19, P20) have described a method for fluorine based on the silicon spectrum. The sample is heated with sulfuric acid and quartz sand to evolve SiF4, which is absorbed in a few drops of potassium hydroxide solution in a silica-free lead-boron glass cup which is subsequently melted in a platinum crucible, or is directly absorbed in a melt of 10 parts PbO and 90 parts B203 at about 560°. The spectrum of the resulting sample, at the intensity ratio of Si 2881 to Pb 2873 A, permits the determination of as little as 40 Mg. of Si corresponding to 0.11 mg. F. Ahrens (A7) studied the detection of fluo­ rine in a siliceous matrix. The latter method was refined by Spence (S104), who determined micro amounts of fluorine, using a platinum distilling apparatus to recover the SiF4 by absorption in a quaternary ammonium salt solution. With copper or platinum electrodes, the Si line at 2881.6 A is used for up to 0.3 Mg- F. With copper electrodes the Si lines at 2516.1 and 2538.5 A can also be used. As little as 0.01 Mg- Si corresponding to 0.03 Mg- F can be detected; differences of 0.02 Mg. Si can be readily detected. Samples con­ taining as little as 1 Mg. F have been analyzed. The chief limitation is the silicon blank. F. MISCELLANEOUS METHODS 1. Enzymatic Activity Inhibition Amberg and Loevenhart (A19) developed a method for detection of fluoride based on the inhibiting effect fluoride has on the action of en­ zymes. Lipase enzyme normally hydrolyzes neutral ethyl butyrate to butyric acid. The action is inhibited by fluoride, so that the difference in volume of standard alkali required in the presence or absence of fluoride is a measure of the fluoride content. Scott (S41) made the method quan­ titative by control of the incubation period and other variables. Phos­ phate did not interfere, which is unusual for methods for the determination of fluoride. Stetter (SI 19) developed a similar method using potato phos­ phatase enzyme. At pH 3.8 his method had a sensitivity of 0.02 Mg- F per miltiliter. Up to a total of 1 Mg. fluoride could be estimated to ±0.01 Mg. Titanium, aluminum, beryllium, and molybdate ions interfered in the method. Schlamowitz (see VI1, p. 183) used a similar method with hog liver esterase stabilized with glycerol as the enzyme and ethyl butyr­ ate as the substrate. The enzyme and ester were incubated with the sample at 25° for 30 minutes; then, further action of the enzyme was 156 PHILIP J. ELVING, CHARLES A. HORTON AND HOBART H. WILLARD stopped by addition of alcohol before titration of the acid formed with standard caustic. The method's range was 0 to 0.5 p.p.m., with an accuracy of ±5% for buffer-free fluoride samples. Others (A20, C85, M36, W19) have made biochemical investigations of the inhibition of enzyme activity by fluorides. For example, Contardi and Ravazzoni (C82) detected soluble fluoride or hydrogen fluoride on leaves by their effect on the enzymatic action of the phosphates of the short cuticle. Insoluble fluorides did not affect the enzyme. 2. Etching and Wettability of Glass Measurement of the extent of etching of glass by hydrofluoric acid, evolved by heating a sample with sulfuric acid, was a method used fre­ quently (K44, L31, 015, W72, Z2) in early analytical methods for fluoride. Often the etching was compared to that produced by standards given a similar treatment (C16, R55, S107), or the depth of an etch groove measured (F60, K19, O10). Woodman and Talbot (W79) dis­ tinguishe7 d between various concentrations of fluoride down to 1 part in 10 by heating the sample with acid at certain defined increasing tem­ peratures. Brauns (B94) measured the change in weight of a Pyrex flask caused by fluoride etching in the presence of acid. Although this method is very simple, it is certainly not accurate. Dubnikov and Tikhomirov (D72) measured the change in wettability of a glass capillary due to attack by fluoride in a microscopic method with a sensitivity of 0.3 μg. F. A warm dichromate-sulfuric acid mixture was circulated in the capillary by shaking. Before adding fluoride the solution wet the wall uniformly. On addition of one drop of a solution containing fluoride, the time required until the glass capillary was not wetted was noted by means of a microscope. The authors (D72) claim better results with quartz capillaries and poorer results with high boron glasses. In earlier work, however, Gautier and Clausmann (G21) claimed that hydrofluoric acid attacks quartz to one-tenth the extent of the attack of glass. Canneri and Cozzi (CIO) measured the contact angle of an air bubble at the end of a boric acid-free capillary tube immersed in a mixture of 100 ml. concentrated sulfuric acid and 2.5 g. potassium dichromate. Presence of fluoride changed the contact angle. A plot of contact angle versus the negative logarithm of the fluoride concentration, pF, was linear for the range pF 2.5 to 4.2. The method was most satisfactory between 20 to 200 μg. of sodium fluoride per milliliter of sample. Boric and silicic acids were the only reagents of those tested which decreased the accuracy and sensitivity of the method. Mayhofer and Wasitzky (M57) and Théophile (T28) also have de- ANALYTICAL CHEMISTRY OP FLUORINE 157 veloped quantitative methods for fluoride based on the etching of glass. Etching has been used widely in qualitative tests (Section V-A-l).

3. Catalytic Activity Hemmeler (H47) studied the catalytic effect of various ions on the autooxidation of sodium sulfite by measuring the time required to attain a standard color using hydroquinone as a chromophoric agent. The color effect of various halides of sodium differed; the autooxidation was retarded by fluoride and iodide and enhanced by bromide and chloride.

VI. Fluorine Compounds

A. SPECTROPHOTOMETRY TECHNICS 1. Infrared Absorption and Raman Scattering Infrared absorption and Raman scattering have been used mainly for qualitative identification or molecular structure studies of both inorganic and organic fluorides. Very little quantitative work has been reported, although some government and industrial laboratories have used infrared absorption for quantitative determination of Freon-type organic com­ pounds or gaseous inorganic fluorides. Williams and McDonald (F55, W58, W59) have developed methods for determining fluorocarbons in air by means of a double-beam infrared analyzer. Smith and coworkers (S81) report the quantitative determination of CC1F3, CHF3, and other im­ purities in C2F4. Benning and coworkers (B41) described a method for determining water in Freons, which do not contain hydrogen atoms, using the water band at 2.76 ě. Gaunt (G13) used an infrared technic to determine im­ purities in boron trifluoride. For qualitative purposes, methane derivatives containing one carbon1- fluorine bond have a C—F stretching vibration at 1045 to 1072 cm.- ; those with two fluorine- 1atoms have two principal bands at 1075 to 1085 and 1140 to 1265 cm. or, if the molecule1 s contain no hydrogen atoms, at 1075 to 1085 and 1140 to 1150 cmr Trifluor1 o derivatives show two bands at 937 to 1102 and 1117 to 1210 cm." Norma-l1 higher alkforanes have-1 two strong infrared bands at 1120 to 11801 cm. and 1340 to 1350 cm. , and a Raman band at 1340 to 1380 cmr Saturated cyclic fluoro­-1 carbons have a typical infrared band between 1140 and 1350 cm. Unsaturated aliphati-1 c fluorocarbons have bands at 1300 to 1340 and 1725 to 1800 cm. , the fluorine-1 atoms enhancing the double-bond stretch­ ing vibration about 100 cm. No definite trends have been established for the aromatic-type compounds (S81). A good discussion of the interpretation of the infrared and Raman 158 PHILIP J. ELVING, CHARLES A. HORTON AND HOBABT H. WILLARD

spectra of fluorine-containing organic compounds has been given in a report by Smith and coworkers (S81). Glockler has discussed some of the simpler molecules in Volume I of the present treatise (G46). The Raman or infrared spectra of several inorganic fluorine com­ pounds have been reported and would be useful in the identification and perhaps in the quantitative determination of these molecules. The com­ pounds studied include HF (B124, E32, R12, S18, S50, W5); SiF4 (B9); BF3 (B9, G13, L36, S106) ; BFrdimethyl etherate (D75) ; KHF2 (C83, K17, P45); KPF6 (L26); alkali fluoroborates (C84); GeF4 (W78); NF3 (B9); NOF (J17, M32, W77); PF3 (Y14); POF3 (D35); POF2X (D35); POFX2 (D35, D38); POFXX' (D36, D38); PFXX' (D37); PSF3 (D39); PSF2X (D39) ; AsF3 (H74) ; MoF6 (B120, G14, T20) ; WFe (B120, G14, T20) ; UFe (B51, B120); U02F2 (B24); F20 (B21, H62, J15, S127); SF6 (E7, E8, E30, E31, L8, SI, Y15); S2F10 (B9, E6); SOF2 (Y13); SeF„ (Si, Y15); T.eFe (SI, Y15); C1F (N9); C1F3 (J16); BrF5 (B119); IF6 and IF7 (L49, L50); and Na2SiF6 (M74). For the phosphorus compounds mentioned, X stands for some halogen atom other than fluorine. Corresponding infrared and Raman studies of fluorine-containing carbon compounds, references for which are not included in the bibli­ ography of Glockler's discussion in Volume I (G46), are: CF20 (N8, W76); CF3OF (L9); CF4 (B9, BIO, L36, PI, P47, R13, W78); CHF3 (A20, B46, P47); CH2F2 (P47, S122); CH3F (B40, M51, Y4, Y5); CC12F2 (P47) ; CC13F (D40, P47, T31) ; CBr3F (D34, D40) ; CBr2ClF (D34, D40) ; CBrCl2F (D34, D40); CH2BrF (D40); CHBr2F (D40, G47); CHC12F (A20, D40, P47, T31); BrCHF2 (P46); Br2CF2 (P46); BrClCF2 (P46); BrCF3 (P46) ; ICF3 (P46) ; CHCIBrF (A20, D40) ; CHC1F2 (P47) ; CBr2F2 (A20); CC12F2 (T31); CC1F3 (T31); C2F6 (B20, B48, N13, P2, R13); C2F6C1 (B20, N12); C2F6I (H33); CH3CF3 (S81, T32) ; CF3CFC12 (N12); CC13CF3 (S81); CH3CHF2 (S81); CF3CHO (S48); CC1F2CC1F2 (S81); CBrF2CHF2 (L5); CBrF2CHCl2 (L5) ; CBrF2CHClF (L5); CC1F2CHC1F (P7); CBrF2CHBrF (P7); CH2=CHF (T40); CH2=CC1F (T40); CH2=CF2 (E9, S81, S83, T40); CC12==CF2 (H20, H31, N10, S81, T40); CC1F=CF2 (S81); CHF=CF2 (P7); CF2=CF2 (Nil, S82, T40); C3F8 (E10, S81); n-C3F7I (H33); 1,1-diiodoethforane (H33); 02N—(CH2)2 —CF3 (S48); CF3—C=CH (H50); CF2=CF—CH3 (S81); (n-C3F7)2S2 (H34); (n-C3F7)2S3 (H34);

Aromatic compounds studied include fluorobenzene (C88, Kl, K49, K51, M124, S81, W73); p-difluorobenzene (S81); m-difluorobenzene (K50); 1,2,4-trifluorobenzene (S81); 1,2,4,5-tétrafluorobenzene (S81); p-amino- ro bromo-, hydroxy-, ro iodofluorobenzene (K52) ; p-nitrofluoro- benzene (R17); p-methoxyfluorobenzene (R18); benzotrifluoride (P26, S81, T33); m-(or p-)fluorobenzotrifluoride (S81, T33) ; 2,5-difluorobenzo- trifluoride (S81, T33); m-chlorobenzotrifluoride (P5); p-fluorotoluene (H60, K49, S81, T33) ; m-fluorotoluene (T33) ; o-fluorotoluene (H59, K49, S81, T33); 2,4-difluorotoluene (K50, T33); and varied derivatives fo methforyl phenyl ketone (P5). Other compounds studied yb Raman scattering ro infrared absorption include: 2-methoxy(or ethoxy, n-propoxy, n-butoxy, n-pentoxy)-l,1,2,2- tetrafluoroethane (P8) ; 2-methoxy-2,2-difluoro-l-chloro-l-fluoroethane and corresponding ethoxy, n-propoxy, and n-butoxy compounds (S81); 2-methoxy-l,2,2-trifluoroethane (P7); 2-methoxy(or ethoxy, n-propoxy, n-butoxy, n-pentoxy)-2,2,difluoro-l,l-dichloroethane (P9); n-dibutforyl oxide (S81); tri(n-butforyl) nitride (S81); "Teflon" (CF2—CF2)n (S81); and "Fluorothene" (CF2—CClF)n (K26, S81). Other recent studies include various partially fluorinated normal alkanes (H39, L3, S82, S84), alkenes (H25, H26), and alkdienes (H26, S82); chlorofluoro alkenes (H30, P6); chloro-, bromo-, and iodofluoro alkanes (H25, H26, H36, H38, H35, H37, P10, S82) ; 1,1,1-trifluoroacetone (H27); 3,3,3,4,4-pentafluorobutyne (H27); and cis- ro £rans-l,2-dichloro- cyclobutforane (L3). Studies have been continued no alkyl and fluoroalkyl disulfides (B91) and alkyl fluorophosphates (Dl) and extended ot tri- fluoroacetylaniline derivatives (B84) and fluorine-substituted betaines (P61). 2. Ultraviolet Absorption A relatively moderate amount fo work has been done no the ultraviolet absorption fo fluorine-containing compounds. Typical inorganic studies are those fo the ultraviolet absorption spectra fo HF (B67, S2); F202 (B99); C1F3 (S26); MoFe (T20); and WFe (T20). Contrary ot other reports, Safary, Romand, and Vodar (S2) found on absorption bands ta 2050 ro 2800 A for pure HF. The probable impurity causing these bands is S02. Ultraviolet absorption spectra have been reported for the following organic compounds sa well sa others: fluorochloromethanes (L4); fluoro- bromomethanes (D22a) ; fluoroethylenes (L4) ; fluorochloroethylenes (L4); fluoroacetylacetones and metal derivatives (H28); fluorobenzene (R6); fluorotoluenes (M2, S87); p-fluoronitrobenzene (S87); p-fluoro- aniline (S87); various substituted benzotrifluorides (M2, S87, SI 10); 160 PHILIP J. ELVING, CHARLES A. HORTON AND HOBART H. WILLARD various sulfides and disulfides (B19) ; and monofluoropyridines and mono- fluoroquinolines (M79). In the simpler compounds the replacement of chlorine by hydrogen or fluorine shifts the absorption to shorter wave lengths (L4). B. X-RAY DIFFRACTION PATTERNS In view of the unique applicability of X-ray diffraction measurements for the identification and determination of individual crystalline com­ pounds in mixtures, it is unfortunate that more data on the results of such measurements as applied to fluorine-containing compounds are not available. The Atomic Energy Commission and its predecessor have made avail­ able a large number of reports covering inorganic compounds containing fluorine. These reports usually cover the X-ray powder pattern or crystal structure study of the compounds. These reports are not listed in the bibliography to this chapter but are identified in Table IX by MDDC, AECD, or other identifying numbers. Two reports have been issued on fluorocarbons (U. S. Atomic Energy Commission No. NYO-715 and NYO-930). A few studies of the X-ray diffraction patterns of fluorine-containing compounds have been published in normal channels, e.g., cyclohexforane (C61) and a variety of inorganic fluorides (F62, F63, F64). Brosset (B103) studied the complex aluminum fluorides by electrochemical and X-ray crystallographic technics and concluded that ther++ e is a stepwise combina­ tion of fluoride and aluminum ions from A1F to A1F6 . The available X-ray diffraction data on a number of fluorine-contain­ ing compounds are given by Wyckoff (W81) in his comprehensive collec­ tion of such data. Similar data is available in a form specially prepared for

TABLE IX Available X-Ray Diffraction Data for Fluorine-containing Compounds

Class of compound Atomic Energy Commission Reports

Uranium and uranyl fluorides MDDC—67, 364, 427, 1151 to 1154, 1283, 1588, 1674, 1675 AECD —2093, 2140, 2523 BR —589 Alkali metal fluorouranates MDDC—67, 1152 AECD —1798, 2089, 2093, 2162, 2163, 2628 Other metal (La, Ce, Th, Pu) fluorides MDDC—67, 1151 to 1154, 1286, 1396 AECD —2628, 3167, 3168 Alkali metal fluorometallates (La, Ce, MDDC—67, 1396, 1675 Th, Pu) AECD —2093, 2162, 2163 ANALYTICAL CHEMISTRY OF FLUORINE 161 analytical application in the card file of X-ray diffraction data for several thousand compounds available from the American Society for Testing Materials.

C. MISCELLANEOUS METHODS BASED ON MEASUREMENT OF PHYSICAL PROPERTIES A number of specialized technics of an instrumental nature based on measurement of fundamental physical properties are of significance for the future analytical chemistry of fluorine-containing compounds. In some cases, limited use of these technics has already been made. The succeeding paragraphs indicate such applications as well as the fact that fluorine compounds have been measured or examined by such technics; the latter fact suggests the possible fruitful utilization of the specific technic for analytical purposes. The discussions in Volume I by Glocker (G46) on the theoretical aspects of fluorine chemistry and by Brice (B96) on the physical properties of fluorocarbons contain considerable material which supplements and enlarges the subsequent discussion. The mass spectrometer cracking patterns for a number of fluorine- containing compounds have been described. One study (M93) covered the mass spectra of a number of lower aliphatic cyclic and straight-chain fluorocarbons; the spectra were compared to those of the corresponding hydrocarbons. Other cracking patterns are still buried in unpublished data at Atomic Energy Commission installations. The parent ion peaks of the fluorine-containing compounds differ from those of the hydrocarbons in that they are all small. The mass spec­ trometry method is feasible for the analysis of mixtures of fluorocarbons of nine or fewer carbon atoms. Pyrolysis of complex compounds to simpler products (R44) might be used to obtain species which could be examined in the mass spectrometer. Inorganic compounds reported include SiF4 (R69), VOF3 (B13), and UF3 (G42). The mass spectrometric determina­ tion of oxygen in fluorocarbon derivatives is discussed in Section VI-D-3. The mass spectra patterns of some of the volatile inorganic fluorides and interhalogens are as yet available only in the files of industrial laboratories. Allen and Sutton (A15) in a compilation of electron diffraction data published in 1950 give the bond distances and references for a large number of fluorine-containing inorganic and organic compounds of various types. The compounds investigated included fluorine derivatives of most of the Ci and C2 hydrocarbons, acetic acid and its derivatives, etc. Sutton and coworkers (C40), later reported the data on GeF4 to complete the data for the fluorides of the first three rows of Groups IV to VII of the Periodic Table. Other diffraction data are available also (B25, B107, S49, S58). 162 PHILIP J. ELVING, CHARLES A. HORTON AND HOBART H. WILLARD

In a review of microwave spectroscopy, Gordy (G57) lists PF3, AsF, 32 34 16 18 CH3F, and CHF3 as having been studied. The microwave spectra of PS F3, PS F3, P0 F3, and P0 F3 have been reported (H40), as have those of trifluorosilane derivatives (S51), CF3C1 (C79), CH2F2 (L37), CH2C1F (M121), and trifluoromethyl cyanide and halide (S52). Micro­ wave methods may sometime be applicable to analytical problems. A microwave refractometer could be used for analysis of binary gas mix­ tures containing hydrogen fluoride. Gutowsky and Hoffman (G75, G76) have examined the nuclear magnetic shielding for hydrogen fluoride, a large number of inorganic fluorides, and saturated solutions of fluoride ion and of a number of complex fluoride anions. Other data on magnetic measurements for com­ plex fluorides are summarized by Nyholm and Sharpe (N27). Fujiwara (F71) studied mixtures of paraffin and Na2SiF6 in an appa­ ratus for observing nuclear magnetic resonance for qualitativ1 e and19 quantitative analysis. The energy absorption intensities by the H and F nuclei are in good agreement with the relative amounts of paraffin and Na2SiF6. The energy absorption line may be readily observed with as little as 0.1 mg. F. Neutron diffraction has also been used to locate the position of hydrogen in potassium bifluoride (P32). The use of neutron scattering in determining the residual hydrogen content of fluorocarbons is discussed in Section VI-D-3. In the same sec­ tion there is also discussed the application of measurement of the follow­ ing physical properties to the same determination: refractive index, specific inductive capacity, and fluorescence. Other properties which might have application in detection of fluorine compounds are dipole moments and dielectric constants (B108, M106) and quadrupole spectra (L43). The latter would be of value mainly in estimation of minute impurities in pure compounds. Exchange reactions involving radioactive fluorine, such as have been used in kinetic studies (R45), may sometime be of use in analytical prob­ lems. Traces of fluoride may some day be determined also by activation analysis using pile or cyclotron activation sources. An unusual method for estimating fluoride depends on viscosity as a measure of the hydrolysis of ethyl orthosilicate in acidic solution. Hy­ drolysis, proportional to the fluoride present, is followed by measuring the increase in viscosity of the test solution in a certain fixed period of time. Aluminum and zirconyl ions interfere seriously, while ferric, chromic, and cobaltous ions interfere progressively less in the order listed (E26). Fabre and coworkers (Fl, F2, F3, F4, F5) used electrodialysis with a three-compartment cell to separate fluoride from suspensions of biologi­ cal-type samples. ANALYTICAL CHEMISTRY OF FLUORINE 163

Recent use of interhalogens in chemistry (G66a) and the variation of electrical conductivity in solutions of BrF3, C1F3, and IF5 (W79a) sug­ gest probable future use of such solvents in analysis. For example, bromine trifluoride acts as an ionizing solvent forming BrF2+ and BrF4~ ions which could be used in isolation of acids such as (BrF2)2-SnFe and BrF2-SbFe and bases such as KBrF4 and Ba(BF4)2 (W79a). Conduc­ tometric or potentiometric titration based on formation of such salts should be possible. Correspondingly, liquid hydrogen fluoride is an ioniz­ ing solvent, and preliminary polarographic studies have already been made in this medium (C70).

D. ANALYSIS OF FLUOROCARBONS Owing to the stability of the carbon-fluorine bond, especially when more than one fluorine atom is on the same carbon atom, the analysis of fluorocarbons has given the analytical chemist a most difficult problem. In a sense there has been a contest between the synthesizer and the analyst. The synthesizer has tried to make carbon-fluorine compounds of ever-increasing stability to temperature and other energy conditions, and of ever-increasing inertness to chemical reactivity. At the same time, the synthesizer has expected and demanded that the analytical chemist obtain more and more precise and accurate analytical data on these com­ pounds which are so resistant to many of the usual methods of functional and elemental analysis. Although there is still much to be done in refining existing methods and in developing new methods, a group of satisfactory methods has been developed in recent years for the elemental analysis of fluorocarbons with the beginnings of methods for the determination of them as compounds. In addition, methods have been developed for analyzing halocarbons in general and for determining foreign elements, impurities, and water. Many of the good analytical methods are apparently still unpublished, reposing in the files of various companies and government installations. In the subsequent discussion of methods for determining elements other than fluorine in fluorocarbons, an attempt has been made to survey the general problem of analyzing organic compounds in the presence of fluorine; in one or two cases cogent analytical problems concerning inorganic compounds have been reviewed. 1. Decomposition to Obtain Fluoride Ion The principal problems in the analysis of fluorocarbons have probably been the decomposition of the compound to obtain fluoride ion and the measurement of the fluoride so obtained. The magnitude of the problem from the viewpoint of obtaining highly precise and accurate results is 164 PHILIP J. ELVING, CHARLES A. HORTON AND HOBART H. WILLARD apparent from the fact that a fluorocarbon, CnF2n+2, contains 77% or more fluorine. The most satisfactory methods for the decomposition of fluorocarbons seem to be those based on (a) heating with sodium or potassium metal in a glass tube or metal bomb (B18, E20, K22, S72); (b) prolonged treat­ ment with sodium in liquid ammonia (M76); (c) combustion in moist oxygen (M84) or steam (R34); and (d) combustion in the presence of silica (C66, R34, T25). These methods and others are discussed in more detail in the section on the decomposition of fluorocarbons and organic compounds; detailed procedures for some methods are given by Rodden (R42). 2. Determination of Constituents Other Than Fluorine Halogens. Some of the procedures which have been described for the decomposition of fluorocarbons to form fluoride ion will convert any other halogen present to the halide ion; the latter can then be determined by standard methods, e.g., Volhard titration. Fusion with alkali metal or treatment with alkali metal in liquid ammonia allows fluorine and other halogens to be determined in aliquot portions of the sample solution obtained (B97, E20, K22, M76). In one combustion procedure, other halogens can be absorbed in silver maintained at 295° ; the gain in weight represents halogen present other than fluorine (T25). Carbon and Hydrogen. Carbon and hydrogen in fluorocarbons are usually determined separately as a result of the difference in stability of carbon-fluorine and carbon-hydrogen bonds, and of the ease of formation of HF where the recovery or removal of the fluorine is more essential than the measurement of the hydrogen. In combustion methods for the decom­ position of fluorocarbons using conventional combustion tube technics, carbon is converted to carbon dioxide which can often be absorbed by Ascarite as in the conventional carbon and hydrogen determination after removal of the fluorine product (M84, T25). Carbon can also be deter­ mined from the carbon residue obtained on heating the compound with sodium contained in a boat in an evacuated system (S72). Until quite recently there was little need for methods involving the simultaneous determination of carbon and hydrogen in samples of high fluorine content. For samples such as fluoroacetyl sugars, benzoyl fluoride, and aromatic fluorides, it was usually sufficient to add lead chromate or potassium chromate to the cupric oxide of the conventional tube packing and, perhaps, to mix powdered cupric oxide with the samples (B31, B93, G73, L12, M109, N17, S21). Occasionally, a copper combustion tube was used, as was a long layer, e.g., 10 cm., of lead dioxide to trap volatile fluorine compounds. Granular sodium or potassium fluoride is used in ANALYTICAL CHEMISTRY OF FLUORINE 165 some cases to remove SiF4 (B35, K72, N17, Y6). For example, the com­ bustion tube may be packed with a mixture of cupric oxide and lead chromate, followed by a layer of lead dioxide, while the external absorb­ ents used to remove the combustion products are concentrated sulfuric acid for water, granular potassium fluoride for SiF4, and soda-asbestos for carbon dioxide (N17). To prevent hydrolysis of SiF4, water vapor must not be condensed before it is absorbed (K72). In compounds containing high percentages of fluorine, the stability of the carbon-fluorine bond makes completeness of combustion in the carbon and hydrogen determination difficult. A further difficulty is introduced by the formation of SiF4 which vitiates both carbon and hydrogen values. The normal Pregl combustion tube packing is stated to give satisfac­ tory results for the micro determination of carbon and hydrogen in com­ pounds containing up to 25% fluorine; compounds containing more fluorine require temperatures of 750° and higher for combustion (Jll). At 750° the lead chromate in the tube fuses and attacks the silica com­ bustion tube; a 2-cm. plug of lead chromate outside the furnace and immediately preceding the silver wire avoids this difficulty. The tube is good for 30 to 50 combustions. Although good results for hydrogen are obtained in the range of 0.5 to 1% H, compounds containing large amounts of fluorine give blanks of the same order of magnitude. Some compounds such as carbon tetrafluoride can be burned only in an oxy- hydrogen flame. Chambers (C44) has suggested the following arrangement for micro carbon and hydrogen determination. The tube filling consists of a layer of cupric oxide between two layers of silver wool, all heated to 800°; SiF4 is removed by lead dioxide kept at 190° (2 g. of lead dioxide will hold 60 mg. of fluorine). Polytetrafluoroethylenes have been satisfactorily analyzed. In a recent thorough study of the micro determination of carbon and hydrogen in organic compounds containing fluorine, Belcher and Goulden (B35) suggested the following tube packing, starting from the entrance end: 14 cm. of platinum contact for oxidation, and 2.5 cm. of silver wool for the removal of sulfur and halogen compounds, including the fluorine, from any hydrogen fluoride formed, kept at 750° ± 20°; then, 2.5 cm. of silver wool, 6 cm. of granular sodium fluoride to absorb SiF4, and 2 cm. of silver wool, all at 270° ± 20°. The silica combustion tube lasts for at least 200 to 300 combustions. If nitrogen is present, an external absorbent for nitrogen oxides is introduced between the water and carbon dioxide absorbents. The standard deviations on a variety of samples were ±0.1 % H and ±0.2% C. Preliminary work indicates that the method serves for fluorocarbons if water vapor is added to the combustion zone. 166 PHILIP J. ELVING, CHARLES A. HORTON AND HOBART H. WILLARD

Hydrogen. It has been recently stated (M28, p. 257) that no com­ pletely satisfactory method has been found for the determination of hydrogen in fluorocarbons, particularly in the low concentration ranges where it is most important. The primary importance of the hydrogen determination has been due to two facts: the preparation of many fluoro­ carbons by fluorination of the corresponding hydrocarbons and the necessity of keeping the hydrogen content to a minimum. Furthermore, since the amount of hydrogen left in the molecule is a good measure of the degree of fluorination, the determination of small quantities of hydrogen in the presence of a large quantity of fluorine is an important problem. Hydrogen in low concentration can be rapidly and satisfactorily determined in fluorine-containing halohydrocarbons by pyrolysis to form hydrohalic acids (M75). The compound is swept in a stream of nitrogen through a platinum tube, packed with platinum gauze, and heated to 1300°. The exit gas is passed through freshly boiled redistilled water, which absorbs acids and free halogen; the hot, freshly boiled solution is titrated with standard base to a phenolphthalein end point. If less than 1 atom of H per atom of halogen is present, all the hydrogen will be con­ verted to hydrogen halide; fluorine is converted only to HF, but chlorine forms some free chlorine for whose contribution to the acidity a correction must be made on the basis of an iodometric titration. If the ratio of hydrogen to halogen is less than 1, chlorine must be added to the nitrogen. Hydrogen in the range of 0.27 to 1.21 % in three fluorine-containing com­ pounds was determined with an average error of 1 part in 25. For hydro­ gen content of less than 0.1%, the method is not quite satisfactory as a control procedure due to the high blank and difficulty in use (M28). Hydrogen in fluorocarbons containing small percentages of hydrogen can be determined by decomposing the sample by heating it with mag­ nesium in an evacuated quartz tube at 700°, passing the hydrogen pro­ duced over copper oxide heated to 300°, and absorbing the water formed in phosphorus pentoxide (P22). Provision is made in the analysis train for an automatic gas pump which serves to recycle unreacted hydrogen and sample. The hydrogen in three samples, 0.4 to 1.3%, was determined to about 1 part in 20. Since the cross section of hydrogen for the capture of slow neutrons is about ten times that for either carbon or fluorine, the decrease in trans­ mission of slow neutrons by a sample of fluorocarbon is a measure of the hydrogen-fluorine ratio (B115, B117, M28, R8). The obvious limitation on the use of this method is the need for an almost monoenergetic source of neutrons, such as would be available from bombardment of beryllium by deuterons produced by a cyclotron or from a nuclear pile. The precision is ±0.5% of the hydrogen content; the uncertainty due to calibration is ANALYTICAL CHEMISTRY OF FLUORINE 167

±0.02% above 0.1% H and ±0.5% below 0.1% H. It has been used on samples containing 2 H atoms per 1000 F atoms. Any element with high neutron capture or scattering cross section will interfere, e.g., chlorine, as will a change in carbon-fluorine ratio. The specific refraction, Τ = = n2 " 1 I ~ Μ η + 2 ' d where η is the refractive index and d the density of the sample, can be used as the basis for an empirical method for determining hydrogen- fluorine ratio with a precision of ±5% (B18, B115, B117, G69, G70, M28). The method is limited to F/H values exceeding 3, since the con­ stancy of the value of the atomic refraction of fluorine is uncertain. If the parent compound is known, this ratio permits calculating the percentage composition of a saturated fluorocarbon or hydrofluorocarbon and the molecular weight. The average difference between the direct analytical and the refractometric results is about 1 %. Fluorocarbon lubricant oils, e.g., O21F44, which contain residual hydro­ gen show greater fluorescence intensity than the completely fluorinated compounds (B115, B117, M28). Standard technic for measuring solution fluorescence is used. The results obtained must be interpreted very cautiously, since the effect due to impurities and structural differences may exceed that due to the presence of hydrogen. The accuracy of the method is limited by the knowledge of the hydrogen content of the calibration standards and the use of standards and samples with similar history. The hydrogen content of fluorocarbons can also be determined by measurement of the capacitance of the sample, since there is a direct proportion between specific inductive capacity and hydrogen content for fluorocarbons of similar structure (R42). The attainable precision is ±0.01% by weight, while the accuracy depends on how accurately the hydrogen content of the reference standard is known. In the range of 0.1 to 5.0% H, the accuracy may be as good as ±0.02% (absolute?). Oxygen. The usual methods for the direct determination of oxygen in organic compounds are not usually applicable to fluorocarbons and their derivatives, owing to the stability of the latter to thermal decomposition and hydrogénation. Kirshenbaum, Streng, and Grosse (K27) have described an isotopic dilution method in which known weights of the compound and oxygen 18 are equilibrated by heating at 800° in a plati­ num tube, followed by the determination by mass spectrometric measure­ ment of the oxygen isotope ratio in any stable oxygen compound formed. Results on substituted fluorocarbons containing 12 to 17% oxygen had an average error of 0.11 %. 168 PHILIP J. ELVING, CHARLES A. HORTON AND HOBART H. WILLARD

Nitrogen. Satisfactory results for nitrogen are obtained with the usual micro Dumas method; high values are occasionally obtained with com­ pounds containing over 75% F, probably due to formation of CF4 and its subsequent measurement as N2 (Jll). When fluorocarbon compounds of the type R3N and R2NF are fused with metallic sodium in a vacuum system, the nitrogen is quantitatively converted to NaCN (S71, S72). The cyanide formed is determined by Liebig's method, i.e., titration with standard silver nitrate solution to the first faint permanent turbidity. In the usual procedure for the decomposition of fluorine-containing com­ pounds by alkali metal fusion, the cyanide formed is destroyed (E20). Water. Water present in fluorocarbons has been determined by titra­ tion with the Karl Fischer reagent (G2, M28), by absorption by phos­ phorus pentoxide (Gl), and by infrared spectrophotometry (B41). In the first method, potentiometric titration (A17) is used with an excess of Karl Fischer reagent solution being added to the sample (diluted with a Freon if necessary) and a standard water-in-methanol solution being used in back titration. Measured quantities of water added to samples could be recovered within ±5 p.p.m.; this precision could be obtained only by skilled, experienced operators (G2). An accuracy of ±4 p.p.m. was ob­ tained in the range of 4 to 400 p.p.m. (0.04%) water (R42). In the determination of water in dimethforylcyclohexforanes using phosphorus pentoxide as absorbent, measured quantities of water were recovered except for an average 3 p.p.m., which remained in the sample (Gl). The P206 method, while more time-consuming, was believed to be more reliable and to give results nearer the absolute value than the Karl Fischer method. In a series of experiments, the average of the average deviations and the greatest single average deviation was 0.7 and 1.4 p.p.m. for P205 and 2.4 and 5.0 p.p.m. for the Karl Fischer method. Concentrations of 0 to 10 p.p.m. of water in Freons and other gases and solids can be measured in a few minutes with an accuracy of 1 p.p.m. by determining the absorption of the sample at the 2.67-ě water bond (B41). Samples such as C2F3C13, which absorb at this wave length, interfere. 3. Determination of Fluorocarbons As Compounds The identification, characterization, and determination of specific fluorocarbons is a matter that presents considerable analytical difficulty. Infrared absorption spectroscopy has proved to be the most useful in dealing with mixtures of fluorocarbons, especially of those containing isomeric species which boil at about the same temperature and, of course, have the same elemental composition. Libraries of the spectra of various fluorocarbons have been prepared by various industrial concerns and ANALYTICAL CHEMISTRY OF FLUORINE 169 governmental installations; such information is slowly being made available. Examination of the infrared spectra of about seventy fluorohaloear- bons has shown that the frequency range- 1of the fundamental C—F vibra­ tion runs between 1080 and 1115 cm. (9.3 to 9.0 ě); in multifluoro, multiearbon atom molecules, the range is somewhat greater (E4). Infrared and Raman spectra of various fluorine-containing ethylenes and other hydrocarbons have been examined from the viewpoint of the assignment of frequencies. Further discussion of absorption spectro­ photometry technics will be found in the section on photometric methods for determining fluorine. Some mass spectra (cracking patterns) of fluorine-containing com­ pounds have been accumulated in industrial laboratories, but, unfor­ tunately, little of note from the analytical viewpoint has yet been published. Traces of fluorocarbons in air can be detected by aspirating the air through a platinum tube, packed with platinum and heated to 950°, and then over a moist piece of thorium-alizarin or zirconium-alizarin im­ pregnated paper (K29). On reaction with water, the pyrolysis fragments form HF, which causes the paper to change color. The sensitivity of the technic is 1 p.p.m. of C8Fi6 by weight in air in less than 1 minute, or 100 p.p.m. in 10 seconds. The use of one of the Freons, izns?/ra-difluorotetrachloroethane, as solvent in the cryoscopic determination of the molecular weights of fluoro­ carbons has been suggested; the molecular freezing point depression constant was 37.6 ± 0.4 (B47, L41). Few data were, unfortunately, pre­ sented. A modified Victor Meyer method has also been used in determin­ ing the molecular weight of fluorocarbons and gave results correct to within 1 or 2% (B18). The Bratton and Lochte modification of the Victor Meyer vaporimetric method gives good results for the microscale molecu­ lar weight determination on liquid organic compounds containing fluorine; the Rast micro cryoscopic procedure, using s//m-tetrachlorodi- fluoroethane as a substitute for camphor, is suitable for solids (J11 ). Benzotrifluoride has also been suggested as a cryoscopic solvent for organic compounds containing fluorine (H13) ; the freezing point constant is 4.90 ± 0.05. The preferred method of determining the molecular weights of fluoro­ carbons and fluorocarbon derivatives is by vapor density measurement. The Victor Meyer method is much less satisfactory than the use of gas density balances. Methods based on the latter are particularly useful and precise as a result of the high density of fluorocarbon vapors. Suitable vapor density balances have been described by Simons and coworkers for 170 PHILIP J. ELVING, CHARLES A. HORTON AND HOBART H. WILLARD the determination of the molecular weight of compounds boiling below 100° (S70) and for that of compounds boiling above 100° (S75); the precision of measurement is ± 1 % or better. Fowler, Burford, and coworkers (F53, F54) have described the follow­ ing analytical procedures used to characterize pure fluorocarbon mate­ rial : (a) molecular weight determination by Victor Meyer vapor density method (accuracy of ± 2 %) ; (b) molecular weight determination of high- boiling compounds by the freezing-point depression method using unsym- difluorotetrachloroethane (B47) ; (c) refractometric method as an indirect check on the molecular weight determination by comparing the experi­ mentally determined specific refractivity with that calculated on the basis of an assumed formula; (d) determination of fluorine by decomposi­ tion with metallic potassium (E20) and modified thorium titration of the fluoride ion.

E. ASSAY AND ANALYSIS OF HYDROFLUORIC ACID AND HYDROGEN FLUORIDE The manufacturing Chemists Association has recommended a set of methods for the analysis of anhydrous hydrofluoric acid, i.e., liquid hydrogen fluoride (S53). The precautions and care necessary in the sampling procedures are discussed in considerable detail. The design and use of sampling cylinders and sample weighing tubes are explained. In addition to the description of sampling technics, the following analytical procedures are given: (a) sulfur dioxide, by reaction with an excess of a standard iodine solution (iodate-iodide solution) and back-titration of the excess iodine with thiosulfate; (b) total acidity, by reaction with stand­ ard sodium hydroxide; (c) hydrofluorosilicic [fluorosilicic] acid, by neu­ tralization to phenolphthalein with standard sodium hydroxide, first in the cold and then continuing at about 100°; the volume used in the hot titration is equivalent to the silica ; (d) sulfuric acid, in which sulfuric and fluorosulfonic acids are calculated as sulfuric acid on the basis of the alkali consumption of the residue obtained on evaporation on the steam bath; (e) water computed as the difference from 100% by subtracting the assay and impurities. The procedures involve many critical details which must be taken into account in order to obtain trustworthy results. These pro­ cedures are based in large part on a study of the sampling and analysis of anhydrous hydrogen fluoride published by Swinehart and Flisik (S131). Cook and Findlater (C78) have also described a procedure for the analysis of anhydrous hydrogen fluoride which is claimed to be suitable for rapid routine work and to have experimental hazards reduced to a minimum. The experimental simplification is largely due to the use of glacial acetic acid as a diluent for the anhydrous hydrogen fluoride ANALYTICAL CHEMISTRY OP FLUORINE 171 samples. Sulfur dioxide is determined by reaction with excess standard iodine solution and back-titration with thiosulfate with a correction for the H2S present; hydrogen sulfide by treatment with cadmium acetate solution and measurement of the precipitated sulfide by the usual iodine- thiosulfate method; and total sulfur by treatment with bromine water and gravimetric measurement as barium sulfate. Total sulfur less the sulfur due to S02 and H2S equals the sulfur due to sulfuric and fluoro­ sulfonic acids. Silica is determined on a moist residue by dissolution in boric acid solution, addition of ammonium molybdate, and photometric measurement. Water is determined by Karl Fischer reagent in a cyclo- hexene-pyridine solution of the sample. The details, basis, and time requirements of the different determinations are given. The specifications of the American Chemical Society for reagent chemicals include a set for hydrofluoric acid (48 to 51 % HF) (W34, pages 172-74). Hydrogen fluoride and fluorosilicic acids are determined on the basis of titrations at 0° and 80° with sodium hydroxide. Tests for maxi­ mum permissible amounts of chloride (silver chloride turbidity) phos­ phate (molybdenum blue color), sulfate (barium sulfate turbidity), heavy metals (sulfide precipitation), and iron (thiocyanate color) are made on individual residues. Sulfite limit is ascertained by an iodine test. Tests for reagent grade sodium fluoride are also given (W34, pages 332-33). Similar procedures for the assay and testing of hydrofluoric acid are given in the U. S. Pharmacopeia (P40). Laszlo (L25) has recently described a method for the simultaneous determination of hydrofluoric and fluorosilicic acids which is based on the precipitation of K3A1F6 (derived from the HF) and K2SiF6 and determina­ tion of the aluminum. McKee and Hamilton (M26) have described a pro­ cedure for the control of hydrofluoric-nitric acid stainless steel pickling baths, which includes the determination of total acidity, iron, nitrate, and fluoride. The latter in the range of 1 to 3 g. per 100 ml. is determined by a Willard-Winter distillation and thorium titration.

VII. Determination of Fluorine in Specific Materials Many of the analytical procedures and methods discussed in the pre­ ceding sections were developed for the determination of fluorine in specific types of samples. The specific handling of a number of such sample cate­ gories has already been discussed, as for example, in the determination of fluorine in fluorocarbons or in different gaseous materials. In addition, frequent reference has been made in the discussion of general analytical methods for fluorine to their suitability for certain classes of materials. Frequently, specialized pretreatment and handling is necessary before 172 PHILIP J. ELVING, CHARLES A. HORTON AND HOBART H. WILLARD fluorine can be determined in many types of samples by one of the general methods. No attempt will be made in the subsequent discussion to indicate the particular variations in sampling, handling, preliminary treatment, and analytical measurement necessary for analyzing any specific type of sample. Instead, references will be given to procedures which have been developed for or applied to various types of samples. Accordingly, anyone desiring information on the general approach to the determination of fluorine in a particular type of sample can consult the references cited, either for that type of sample or for one similar to it.

A. BIOLOGICAL SAMPLES The use of fluorine in plant sprays and insecticides in general, the role of fluorine in dental caries, fluorine intoxication, the various aspects of the industrial hygiene of fluorine, and the many other ways in which fluorine is related to animal and plant welfare have resulted in numerous studies on the determination of fluorine in biological samples. In many or most of these, the concern is with inorganic fluoride ion which is usually separated from the rest of the sample by a Willard and Winter distilla­ tion. Often, the sample is first ashed to destroy organic material; calcium or magnesium oxide, peroxide, or other compounds may be added previous to ashing to fix the fluorine (e.g., see Sections III-A and III-B). Proce­ dures generally applicable to biological samples are given in references Al, A33, C93, Fl, F2, F40, G52, G66, H70, K25, M34, M35, M58, S32, and S37. As might be expected, the determination of fluorine in toxicological and forensic analysis has received considerable attention. Fluorides and silicofluorides are important industrial dust hazards (F8). Jacobs (J2) has a good discussion on the determination of hydrofluoric acid and of fluorides in toxicological work. In the volume on the pharmacology and toxicology of uranium com­ pounds resulting from the Manhattan Project (VI1), there is a helpful discussion by Flagg (F40) of methods for the determination of minute amounts of fluoride ion in a variety of substances such as air, tissues, and biological fluids. Other references on the detection and determination of fluorine in toxicological work include B16, B87, B100, Bill, G33, G52, G54, G66, LI, L29, L53, M40, P19, and S87. /. Plants Typical studies on the determination of fluorine in plants, vegetables, foliage, and similar samples are given in references C35, C36, C65, D20, ANALYTICAL CHEMISTRY OF FLUORINE 173

D21, E5, F26, G25, G59, H55, 16, 17, J2, K67, M13, M14, M17, M34, M35, M56, M58, M89, 015, RIO, R15, R36, S62, S63, S98, S119, W31, W42, W66, W67, W68, W69, W70, and Z5. Many of the foregoing pro­ cedures are also applicable to the analysis of soils (see also L34, M7, M8, M14, M15, M16, M18, M22, 015, R15). Kruger (K67) has discussed the detection of fluoride in plants injured by fluorine. The detection and determination of fluorine on plants and vegetables due to sprays is covered in references E5, G67, L34, and W31. Mitchell (M89) reviewed the spectrographic determination of fluorine in plants, soils, and related materials (see other references on emission spectroscopy in Section V-E). The determination of fluorine in insecticides and sprays is discussed in references B32, B76, C28, D62, D63, D64, D65, D66, E5, E17, G67, H23, K67, L33, L34, RIO, S60, S61, S64, S88, and W31. A number of procedures (B72, C29, C95, E19, II, 12, K62, L2, W16) cover the deter­ mination of fluorine in wood preservatives and in treated wood.

2. Animals A number of analytical methods have been described for the detection and determination of fluorine in body fluids and tissues (e.g., C93, F40, G32, G33, G52, L20, M34, M35) and, more specifically, in blood (G17, K61, M62, S85, S86, W80). Other typical procedures are given in refer­ ences A33, C14, C15, C17, C18, C93, G17, G19, G20, G16, G32, K61, K66, L20, M34, M35, M55, M56, M57, M58, M62, M105, RIO, S32, S37, S85, S86, S119, W19, and W66. Specialized methods have been developed for the determination of fluorine in bones and in teeth. In general, these methods can be supple­ mented by procedures intended for the determination of fluorine in minerals and phosphates. Procedures specifically applicable to bones and teeth include those in references B45, B86, C33, C68, C94, D26, Fl, F5, F26, F27, G48, L47, M4, M8, M40, M103, O10, 09, 012, P3, S47, S59, S92, S109, T54, and VI1.

B. FERTILIZERS, PHOSPHATES, AND PHOSPHATE ROCKS The question as to whether fluorine-containing fertilizers can cause plants to have a fluorine content toxic to grazing livestock or to other animals and human beings eating such crops has stimulated search for reliable methods for the determination of small amounts of fluorine in phosphatic material. Typical methods for fertilizers and phosphatic samples (including inorganic and organic fluorophosphates) are given and discussed in references A5, B88, C28, D7, F42, G30, H21, H64, 16, 17, Jl, K24, M12, M15, M19, M24, M48, M83, M103, 014, P43, R16, R20, R23, R24, R25, R26, R27, R29, R31, R50, S77, T7, U2, V9, W4, and W32. 174 PHILIP J. ELVING, CHARLES A. HORTON AND HOBART H. WILLARD

Isakov (16) has covered the determination of fluorine in the gases and solutions obtained in the production of superphosphate. Ahrens (A5) has described an emission spectrographic method for fluorine in phosphate rock which is applicable to as low a concentration as 0.02 % F. Mason and Ashcraft (M48) have described the determination of as little as 0.0003 to 0.09 % F in phosphoric acid and trisodium phosphate. Hill and Reynolds (H64) have described the colorimetric determination of fluorine in monofluorophosphate. C. FOODS AND BEVERAGES The literature on the general determination of fluorine in foods, food products, and foodstuffs is quite large, as is indicated by the following specific references: A16, A19, A23, C39, C55, C64, C71, C72, C73, C75, C76, C93, D2, D3, D4, D5, D7, D9, D10, D46, E17, F24, F27, G9, G10, G17, G30, G54, G67, H73, K25, K45, L19, L32, L33, L34, L45, M3, M5, M58, M105, R55, R59, S22, S47, Vil, W12, W18, W35, W36, W38, W39, and W40. References for the toxicological assay of fluorine in food can be found in Section VII-A. Representative procedures for related types of products include those for butters and margarines (D46); table salt (C54, C55, F29); baking powder (Bll, D8, L33, L34, MHO, Mill, M112, M113, M114, M115, P17); meat (B71, C76, M96, N21); fish (L32); gelatin (M95); preserves and jellies (D46); alcoholic beverages (A16, B12, C66, C76, C80, D46, F26, F 27, F42, H55, K19, L27, M61, R19, SI 19, W18, W79); spray residues on food products (H73, L34, see Section VII-A-1); and feeding stuffs (D5, D9, E17, M34).

D. ROCKS, MINERALS, AND ORES Many of the methods described in previous sections of the present chapter for the detection, separation, and determination of fluorine were originally developed for the examination and analysis of various types of minerals and rocks. Typical procedures for various types of rocks and minerals which are also usually applicable to ores and soils are covered in references All, A12, A13, A14, A16, B82, C96, D21, D29, D30, F6, F12, F13, F14, F20, F45, F73, G17, G22, G28, G30, G31, G48, G68, K58, L24, L28, L33, L47, M33, M39, MHO, 08, O10, 014, P51, P52, R23, R24, R27, R28, R29, R31, R39, S3, S59, S77, S98, S102, S109, T28, T41, W45, and Y2. Emission spectroscopic methods have also been described (A5, A7, A8, B82, P3, S44). The specific determination of fluorine in soils and liming materials are discussed in references L33, M7, M8, M9, ĚÇ, M12, M14, M16, M15, M18, M20, M21, M22, M23, M89, R15, S28, and S44 (see also Section VII-B). ANALYTICAL CHEMISTRY OF FLUORINE 175

The specific determination of fluorine in silicates and fluorosilicates has also been thoroughly investigated (A7, All, B109, D71, HI, H2, H3, K3, M18, M83, N22, T8, T28; see also Section VII-G). Another area which has been well studied is the determination of fluorine in phos­ phorites and apatites (P43, R24, R47, R48, R49, R50, S78, T7, U2; see also Section VII-B). Methods have been described for various special types of ores and minerals, e.g., bauxite (P4); cryolite (C28, D70, F57, R37, S108, T18, T44, Z9) ; fluorite and fluospar (A22, B92, D71, L40, L46, M42, P50, P57, R37, T57); lithium minerals (G74); zinc ores (K20, R40, S51); blendes and sulfide ores (F60, K20, O10, 012, T28); lime (K46); and coal (C90, C91).

E. WATER The literature on the determination of fluoride ion in water is volumi­ nous with colorimetric methods, as might be expected, predominating. The general references include the following: A29, B23, B56, B58, B74, B77, B83, C13, C16, C31, C42, C51, C56, C68, D2, D7, D29, D73, D74, E22, E23, F12, F27, F48, F50, F51, G4, G17, G23, H2, H42, 14, 18, J9, J8, K61, K62, L10, LU, L33, L34, M23, M31, M34, M39, M91, M92, M97, P33, P48, R25, S5, S6, S8, S38, S39, S45, S90, S92, S93, S95, S96, S98, S119, S120, Tl, T4, T23, T39, T56, Vll, W6, W7, Wll, and Y12. Specific references to potable water supplies are C62, D15, F25, F26, J19, L20, M123, S7, S40, S65, S89, S91, S109, T37, and T55. The A.P.H.A. (A20) uses a colorimetric procedure (zirconium-alizarin method) with or without a prior Willard-Winter distillation. A number of methods are concerned with the determination of fluorine in mineral waters (B14, C12, C18, C30, C31, C34, G36, 08, P48). Okano and others have investigated the determination of fluorine in hot springs (B72, 04, 05, 06, 07, S31). Wilcox (W46) studied water for irrigation use. Koehler (K46) critically investigated the determination of fluorine in lime intended for use in water treatment. Goldman and coworkers (G53) and Yoe and coworkers (S5) surveyed the detection of organic fluorine compounds in water, summarizing the results of a considerable number of investigations carried out by the Chemical Warfare Service. The determination of fluorine in sea water is discussed in references A27, A37, M90, and T34. A rapid (5-minute) test for fluoride in the range of 0.8 to 1.4 p.p.m. for use in the installation and adjustment of chemical-feed apparatus for the fluoridation of water supplies has been described by Rubin (R58). The LaMotte Company has available a kit for the colorimetric determination of fluoride in water in the range of 0.2 to 2.0 p.p.m., with an accuracy of 176 PHILIP J. ELVING, CHARLES A. HORTON AND HOBART H. WILLARD

0.1 p.p.m. A method has been described for fluoride in tannery soak waters (D59). F. AIR The detection and determination of fluorine and fluorine compounds such as hydrogen fluoride in air has received considerable attention, although much of the resulting data and methods have remained unpub­ lished in industrial and governmental files. Typical of the investigations which have been published are those of Yoe and coworkers (S5) on the determination of physiologically active monofluoro organic compounds in air, and of Yaffe (Yl) on atmospheric concentrations of fluorides (dusts, fumes, and gases) in aluminum-reduction plants. The various subsections of Section II on the analysis of gaseous sam­ ples contain references for and descriptions of various devices and methods for the detection and determination of gaseous fluorine com­ pounds. Attention should be called to the halide meters (D24, H49, H79, M70) described in Sections II-D and IV for the continuous analysis of air for halogen-containing compounds. Fluorocarbons in air can be determined by infrared absorption (F55, W58, W59). Flagg (F40) describes a sam­ pling device suitable for air containing hydrogen fluoride. In Section IV various devices are described for the detection of fluorine and its compounds in air, e.g., a device consisting of a zirconium- salt-alizarin impregnated absorbent for detecting hydrofluoric acid vapor in air (M38). Thermal conductivity methods for hydrogen fluoride and fluorine in nitrogen or air-like mixtures have been described (K28, SI 12). Luzina (L55) has described the determination of fluorides in aerosols by electrical precipitation and absorption in a liquid absorbent. Other references which would be helpful in the analysis of air include C53, C77, D54, F40, H79, J18, K29, L55, M43, NI, N2, RI, R6, R71, S23, S45, S113, T52, T53, T58, VU, W52, W74, and Y10. Reference should be made to the paragraphs on toxicological analysis in Section VII-A-1. G. MISCELLANEOUS A considerable number of specialized methods have been described for the detection and determination of fluorine which do not fit into the preceding categories; a representative selection are indicated in the following paragraphs: In glasses and enamels in the range of hundredths of 1 % to over 12% (B22, D26, G28, H66, L47, P14, R74, S98, T38) ; sulfuric acid and oleum (S107); gases and solutions obtained in superphosphate production (16). In chromium-plating baths (M47, 013, W60); nickel-plating baths (C67, H15); hydrofluoric acid-nitric acid stainless steel pickling baths ANALYTICAL CHEMISTRY OF FLUORINE 177

(M26) ; heat-treating salts (E24, P34) ; basic slag (S98, W15) ; welding flux (C52) ; steels (S68). In various stages of aluminum production (E24, P4, V6) and of magnesium production (P50, P51, V6). In aluminum compounds (R7) ; beryllium compounds (C58, L39, R33, V2); borates and fluoroborates (A23, A24, B79, D62, K15, K20, L34, M3, M94, R33, R67, R75, R76, R79, S102, T42, T45, W1C); boron trifluoride (B79, M77, S130, W9); lime (K46); silicic acid "soots" (R5); fluoro­ silicates (A26, B2, B76, B98, B109, Cl, D52, J10, K20, K43, L33, L34, L34, M18, M53, MHO, M116, S60, T9, T35, V3, V5); sodium chloride (F29); zirconium metal and compounds (C59, D25). In rodenticides (E17) and insecticides (see Section VII-A-1). In tannery soak waters (D59). In pharmaceuticals (Al, P40, R4, S87). In textiles (E19, H14, J16, S99).

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