Pure B F3 was claimed to be prepared by the reaction of KBF4 and BaCl2 at 600° by Vinnik et alSlm\ The ionization potential was given as 10.2 volts by Kaufman*3 8 9); the bond strength for B F3 was given as 6.86 ± 105 dynes per cm, and for BF4~ as 5.278 x 105 as calculated by Goubeau and Bues*2 7 7). The infrared spectra of B F3 was given by Van-derryn*7 4 4). The heat capacity and vapor pressure were measured by Oliver and Grisard*5 1 9). The heat of fusion was found to be 2874.6 ± 3 cal per mole at the triple point 281.93°K. Vapor pressures from 38 to 155°K were given by logio Pmm = 7.74853 - [1685.8/( + 220.57)]. The heat of vaporization at the boiling point 125.75°K is 10.2 kcal per mole.
The entropy of the liquid and ideal gas at 298.16°K are 42.57 ± 0.10 and 71.57 cal per deg per mole, respectively. The dissociation energy of BF was calculated to be 185 kcal per mole by Barrow*42).
Diboron tetrafluoride, B2F4, has been made by Finch and Schlesinger*224) by the reaction B2C14 + S b F3 -8 0 to - *s B2F4. It had a melting point of —5.6°. The vapor pressure of the solid is given by logio Pmm = 10.84 (1856/T) and the liquid, logio Pmm = 9.009 -(1466/T). These give —34° for the boiling point and about 6700 cal per mole for the heat of vaporization. It apparently formed a dietherate stable to —23° which then decomposed to a monoetherate which was stable to 0° but which gave ethyl fluoride on standing at this temperature.
24 J. H. SIMONS
B2F4 with Me3N gave a solid of the formula [B2F4(Me3N)2]4. The crystal and molecular structure of B2F4 were determined by Trefonas and Lips
comb*7 3 7) by means of X-ray diffraction.
Boron trifluoride monodeuteriate D[BF3 • OD] melting at 11.0° with a density, in the temperature range 5 to 20°C, of
= 1.8482 - (1.11/10000
and boron trifluoride dideuteriate D30 ( B F3O D ) melting at 11.3° and with a density, in the temperature range 8 to 20°, of
rf4* = 1.7232 - (1.46/1000/)
were reported by Greenwood*2 7 9). H[BF3OH] and H30 [ B F3O H ] melted at 6°.
Boron trifluoride is a reactive chemical. Sprague, Garrett, and Sisler*7 0 6) caused it to react with the oxides of nitrogen and studied the physical properties of the products. N203 and B F3 gas gave N203 • 2BF3; N204 + B F3 gas gave N204 • 2BF3; N204 + B F3 + 03 gas gave N206 • 2BF3; N205 + B F3 in CC14 gave N205 • B F3; N204 + B F3 in S 02 gave N204 • 2BF3; and N204 + B F3 in H N 03 gave NOBF4. Bachman and Hokama*35) reported N203 • B F3 by combining the gases in nitro-ethane. Bachman et #/.*34) reported N204 • BF4 by a similar method.
Boron trifluoride resulted from the fluoridation of BN with either F2 or AgF2 according to Glemser and Haeseler*2 7 0). B F3 reacts with 02N F or with N2C>5 and H F to form nitrylborofluoride, 02N B F4, which is a nitrating agent as reported by Olah and Kuhn*5 1 6). Glemser and Liide-man*2 7 1) found that B F3 gas + solid S4N4 at room temperature forms 4S4N4 • BF3, carmine red, which sublimes at 95° in B F3 atmosphere.
B F3 diluted with N2 and S4N4F4 gives S4N4F4 • B F3 which is green.
Monoammonia-boron trifluoride, B F3N H3, was found by McDowell and Keenan*4 7 0) to react in liquid ammonia with the alkali metals. With po
tassium and cesium the mole ratio of the reactants is one and the products are presumably the alkali fluorides and B F2N H2. With sodium, 2.5 moles of the metal and with lithium 3 moles of the metal react with one mole of B F3N H3. Parry, Kodama, and Schultz*5 3 0) found that B F3N H3 slowly crystallized from an aqueous solution as BF4NH4. B F3 was shown to form addition compounds with SF4, SOF4, and IF5 with the formulas S F4 • BF3, SOF4 • BF3, and I F5 • B F3 by Seel and Detmer*6 5 1). Schmeisser and Jenkner*6 2 1) found it to react with ethylene oxide, 2BF3 + 2 C2H40 ^ C4H s 02 • 2BF3 which decomposes at 130° to B F3 and C4H s 02, dioxane.
Paterson and Onyszchuk*5 3 0 t >) studied the reaction of boron trifluoride and hydrazine. At 25° in vacuo B F3N2H4, which melts at 87°, is formed. In tetra-hydrofuran, 2BF3 • N2H4, which melts at 260°, is formed. The thermal de
composition of B F3N2H4 gives N2, N H3, N H4B F4, and BN. Parsons et a/#( 5 3 0 a ) prepared C F3B F2 in two ways. From the reaction product of
K B( n C 4 H g ) 2 and CF3I a yellow ether solution was obtained in addition to KI, which on treatment with BF3 yielded BuBF2 and CF3BF2. In the other method, CF3SCI reacted with B2H6 at 60° for 15-18 hr to form CF3BF2. This compound formed adducts with (CH3)3N and (C2H5)3N.
Salts of methforylfluoroboric acid were prepared by Chambers et tf/.*120a).
First ( C H3)3S n C F3 combined with B F3 to form (CH3)3Sn+(CF3BF3)-.
From this salt were obtained KBF3CF3, which is water soluble, and Ba(BF3CF3)2 and NH4BF3CF3. The acid H B F3C F3 was also obtained, The crystal structure of H3C H2N • B F3 and H3CCN • B F3 were studied by Geller and Hoard*2 5 9). Borontrifluoride thiocyanate has been shown to decompose at —40° by Seel and Miiller*6 5 5) and BF3 has been found by Goubeau and Bergmann*2 7 6) to react with sodium hydride at —70° in ethyl ether to form NaHBF3. This compound reacts with water to form Na[BF30H] and H2 and at 200°, the compound reacts with NaH to form Na[H2BF2] and NaF. The hydroxytrifluoroborates such as KBF3OH, NaBF3OH, and K2B3F4O3OH were prepared by Ryss and Slulskaya*608>.
The tetrafluoroborates are usually hydrated. Some of the anhydrous salts are either unstable or do not form. L i B F 4 was prepared by Shapiro and Weiss*666> at 35° in ether solution 3Li2C03 + 8BF3 -> 6LiBF + 3CO2 + B2O3. The barium salt Ba(BF4)2 was reported made at 130°
from BaF2 and B F3 • 2 H20 by Pawlenko*532>. El'Kenbard*2 0 0> studied the dehydration of a number of hydrates. He found that both Sr(BF4)24H20 and L i B F 4 • 3H2O could be dehydrated to the anhydrous salt with P2O5.
Mg(BF4)2 • 7 H20 and Cd(BF4)2 • 6H20 formed dihydrates by the same method, and Cu(BF4)2 • 6H2O formed an unstable tetrahydrate. Ni(BF4)2 • 7H2O and Co(BF4)2 • 6H2O formed no lower hydrates and the anhydrous salts were unstable. A toluene solution of silver fluoroborate was made by reacting anhydrous silver fluoride with BF3 in toluene, and the copper salt was similarly prepared in toluene solution using a mixture of copper and cupric fluoride by Warf*7 6 0). The anhydrous salts were not obtained upon evaporation of the hydrocarbon. Pure colorless A g B F 4 was prepared by Heyns and Paulsen*3 2 8) by passing BF3 through a suspension of AgF in C e H 6 until nearly all the AgF is dissolved. Upon concentration of the filtered solution (CeHe)2 • A g B F 4 was precipitated. At 50° and 1
mm-C e H 6 was lost and A g B F 4 obtained. Olah and Quinn*5 1 8 a) used nitro-methane as the solvent for this reaction.
The infrared spectra of ammonium, sodium, and potassium fluoro-borates were studied by Cote and Thompson*1 4 4). Ray and co
workers*1 6 2'5 6 4'5 6 6) prepared oxyfluoroborates such as BaBFsO, PbBFsO, and SrBFsO but were unable to make M2BF3O compounds. They did obtain KHBF3O • 3BF3 and ( N H4)2B F30 • 3BF3. They also obtained C u B F30 • 5 H20 , N i B F30 • 7 H20 , C0BF3O • 7 H20 , ZnBF3Q • 7 H20 ,
26 J. H . SIMONS
and C d B F30 • f H2O. In addition they made double salts of the formulas.
(NH4)2S04MBOF3 • 6 H 20 and (NH4)2BeF4MBOF3 • 6 H 20 where M = Cu, Co, Ni, Zn, or Cd. Chackraburtty*1 2 0) studied the crystal struc
ture of BaBOF3.
Long and Dollimore*4 3 7) believe that dihydroxydifluoroboric acid, H3BC>2F2, is not capable of independent existence. In the vapor it is probably H2O and BF2 • OH. Ryss*6 0 8) reported the trimeric difluoro-orthoborates, N a3B 30 3F 6 and K 3B 30 3F 6.
Aluminum fluoride has been so extensively studied in the past that most new information consists of physical measurements, particularly of increased precision. Staritzky and Asprey*7 0 9) discussed the crystallo-graphic data of A1F3. The structure of A1F3 • 3 H 20 was determined by Freeman*2 4 4^. Aluminum monofluoride A1F was prepared from the metal and A1F3 at temperatures in excess of 800°C*63°).
Its band spectra was determined by Rowlinson and Barrow*599^, the emission spectra by Naude and Hugo*4 9 8) and the thermodynamic and spectroscopic properties discussed by Barrow, Johns, and Smith*4 6).
The dissociation energy of A1F was calculated to be 156 kcal per mole by Barrow*42). The infrared spectra of synthetic cryolite Na3AlF6 was investigated by Lattre*4 2 7). The vapor pressure of A1F3 was determined by Evseev et a/.*2 1 4) between 980 and 1123°K, logp = 14.44 - 16,967/T and the heat of sublimation is 77.63 kcal per mole at 1051.5° and 80.28 at 0°K. Witt and Barrow*7 9 6) also determined the vapor pressure and found l o g Pm m= 1 2 . 7 1 4 - 15,122/T from 955 to 1400°K, giving a heat of sublimation of 72.72 ± 0.46 kcal per mole. They also determined the reaction A1F3 ( C) + 2A1 ^ 3 A l F( g) in the pressure range 10~4 to 10~2 mm Hg. They found the heat of formation of A1F at 298°K = - 60.97 ± 0.41 cal per mole and D0 of A1F = 156.1 + 1.8 kcal per mole.
The heats of the six successive reations of F~ with A l3+ to form A1F63~
have been measured by Latimer and Jolly*4 2 6).
Porter and Zeller*5 5 4 a) determined the heats of dissociation of Al2F6(g) and LiF • A l F3 ( g) by mass spectrometric methods. They found for the former AH0 at 1000°K to be 48.0 ± 4.0 kcal per mole of dimer and AH° at 1000°K for the latter to be 73 ± 4.0 with the products as gases.
King*4 0 0) measured the low temperature heat capacities of cryolite and aluminum fluoride and found the entropies at 298.15°K for Na3AlF6
F " + A l3+ -> A1F++ A S0 = 32 \ F - + A1++ - * A1F2+ = 26
F " + A F2+ -> AlF3(aq) = 18 18 I cal per deg per mole
to be 57.0 ± 0.4 cal per deg per mole, and for A1F3, 15.85 ± 0.08. The entropy of formations of cryolite from sodium fluoride and aluminum fluoride is 4.3 ± 0.5 at 298.15°K.
(NH4)sAlF6 was prepared by Haendler, Johnson, and Crocket*2 8 9) by reaction of a methanol solution of AlBr3 and NH4F. The high temperature heat contents of cryolite, anhydrous aluminum fluoride, and sodium fluoride were measured by O'Brien and Kelley*5 1 2). The heat content (in cal per mole) of Na3AlFe(a) is given by the equation # r — # 2 9 8- 1 5 =
45.95T+ 14.73 x 1 0~3r2 + 2.78 x 105T-i - 15,942 from 298° to 845°K, where it is transformed to the /3 crystal form with a heat absorption of 2160 calper mole. For Na3AlF608) HT - # 2 9 8. 1 5 = 52.15T + 7.93 x 1 0 -3T2 -13,840 from 845° to the m.p., 1300°K. The heat and entropy of fusion are 27,640 cal per mole and 21.26 cal per deg per mole respectively. For liquid Na3AlF6, HT - #2 . i 59 = 93.40T - 26,480 from 1300 to 1400°K. 8 For AlF3(a) HT - # 2 9 8. 1 5 = 17.27 T + 5.48 x 10"3r2 + 2.30 x lO^T"1
- 6480, from 298 to 727°K where transformation to AF3(j8) occurs with the absorption of 150 cal per mole. For AlF3(j8), # r — # 2 9 8 . 1 5 = 20.93 T + 1.50 x 1 0~3r2 - 6500, from 727 to 1400°K. The heat and entropy of fusion are 8030 cal per mole and 6.25 cal per deg per mole, respectively.
HT - # 2 9 8 . 1 5 for liquid N a F is 16.40T + 170 from 1285 to
1800°K. Coughlin*1 4 5) determined the heat of formation of Na3AlF6 ( C) from the elements at 298.15°K = — 784.8 kcal per mole. For its reaction 3 N a F( c) + AlF3 (c) -> N a3A l F6 ( c )a t 298.15°K. A # = - 20.3 kcal per mole,
AF = - 21.6, and A F1 480° K = - 29.1.
The compound NaAlF4 has been obtained by Howard*3 4 5) from the quenched vapors from molten cryolite. Upon reheating it disproportion-ates to chiolite and aluminum fluoride. Consider the reaction 5NaAlF4 -> 5NaF • 3A1F3 + 2A1F3. The melting point of NaAlF4 was given by Mashovets et #/.*4 5 9) as 1020°. T h e electrical conductivities of molten Li3AlF6, Na3AlF6, and K3A 1 F 6 were measured by Yim and Feinleib*8 1 0). X-ray evidence for both mono-and tri-hydrates of aluminum fluoride was given by Fisher, Bock, and Meisel*2 3 2). A series of hexa-fluoroaluminates were described by Mitra*4 8 0) of the form M'M"A1F6*
6 H20 . They are made by dissolving the carbonates or hydroxides of the three metals in 10% hydrofluoric acid, heating to 80° in a platinum vessel, and then crystallizing in vacuo. The salts are all crystalline and water soluble. They lose water above 80°. The following salts were reported:
N a C u A l F e • 4 H20 N H4F e A l F6 • 6 H 2 O K C u A l F e • 6 H 2 O N H 4 C 0 A I F 6 • 6 H 2 O N H4C u A l F6 • 6 H 2 O N a N i A l F e • 4 HaO
R b C u A l F e • 6 H 2 O N H4Z n A l F6 • 6 H 2 O
28 J. H . SIMONS
K Z n A l F e • 6 H 2 O K N i A l F e • 6 H 2 O
K C d A l F e • 6 H 2 O N H4N i A l F6 • 6 H20 N H4C d A l F6 • 6 H20 T I N i A l F e • 6 H20
Scandium (III) is complexed by fluoride in aqueous solutions as shown by Kurry, Paul, Hepler, and Connick*4 2 2).
Sc+++ + H F - ScF++ + H + ScF++ + H F = S c F2+ + H +
S c F2+ + H F = S c F3 ( aq ) + H + S c F3 (a q ) + H F = S c F4~ + H +
Yttrium fluoride of low oxygen content was prepared by Smutz, Burnet, Walker, Tischer, and Olsen*7 0 1) by heating a yttrium compound such as the oxide in H F vapor at 600°. Its preparation from Y2O3 and N H4F • HF at 400° by Walker and Olson*751) and from Y203 or YC13 and F2 by Tischer and Burnet*7 3 4) gave a product containing more oxygen. Carlson, Schmidt, and Spedding*1 0 9) prepared the metal from the fluoride by the reaction with calcium or magnesium at 1000°.
Crystallographic data for YF3 was given by Staritzky and Asprey*7 0 8) and its crystal structure by X-rays by Zalkin and Templeton*8 1 8). Hund*3 4 9) prepared YOF from the combination of Y2O3 and YF3 in vacuo at 900°
and found two crystalline forms. Zachariasen*816) studied the crystal structure of YOF. Hund*3 4 8) prepared NaYF4 in two crystalline forms and studied their crystal structure. Bode and Voss*75) reported the crystal structure of K3SCF6, K3YF6, Rb3YF6, and CS3YF6. The density of L a F3 in the range 1750 to 2450°K is given by Kirschenbaum et «/.*401b) as 5.793 - 6.82 x 10"4r. Roberts*575*) prepared L a ( C F3C 02)3 from L a203 and C F3C 02H .
The crystal structure of a K L a U F 4 was studied by Zachariasen*811) who also studied L a F 3 and LaOF.
Actinium trifluoride, A c t F 3 and oxyfluoride ActOF were prepared by Zachariasen*8 1 3'8 1 6) and X-ray crystal information obtained. T h e former compound was obtained by reaction with H F and the latter by hydrolysis of the trifluoride. The trifluoride was used by Stites, Salutsky, and Stone*7 1 8) to prepare the metal by reduction with lithium vapor at 1000°.