T H E HYDROFLUORIDES OF ORGANIC BASES AND A STUDY O F HYDROFLUORIC ACID BY J . F. T. BERLINER* AND RAYMOND M. HA”**
The literature of organic chemistry abounds in the description and methods of preparation of organic bases and their derivatives. The hydrochlorides, hydroiodides, hydrobromides, sulfates, oxalates and picrates of a majority of the commoner bases have been prepared, analyzed and their physical properties recorded. During the investigational work on the quantitative preparation of jiodo-2-amino-toluene and some of its more common derivatives’ the authors prepared the hydrofluoric acid addition compound of the above base. Upon estimation of the nitrogen content of the above salt (Kjeldahl-GunningArnold method) 4 44 per cent of nitrogen mas found instead of the 5 . 5 2 per cent calculated for a compound CF,H~.(CHJ).J.?;H*.HF.This material was prepared three times and recrystallized repeatedly on the supposition that some impurity was affecting the results, but upon reanalysis of these highly purified materials practically the same nitrogen content was found as in the first preparation. At an earlier date one of the authors had prepared and analyzed some otoluidine hydrofluoride and found that although recrystallized several times, the nitrogen content remained practically constant between the limits 7.40 to 7.47 per cent instead of yielding the i i . 0 2 per cent nitrogen calculated for the compound CsH4.(CH3).NH?.HF.At the time this discrepancy was thought to be due to some extraneous cause and no further thought was given to it. This compound and its analysis were recalled to mind when the disparity in the nitrogen content of the iodo-o-toluidine was noted. Upon inspection of the analytical results obtained by the nitrogen analysis of the o-toluidine and iodo-o-toluidine, it seemed evident that there might be two or three factors acting separately or together, which could be held responsible for the great difference between the amount of nitrogen actually found and the theoretical amount calculated from the general i : x basehalogen acid addition product. Possible factors were that the derivative could contain water of crystallization, that a polymer of hydrofluoric acid combined with the base, or that several unpolymerized molecules of the acid combined separately. Should the first factor be true, the nitrogen content of o-toluidine hydrofluorlde would correspond to a compound containing about 33.1j per cent of water and therefore equivalent to the formula (CsH, * J. F. T. Berliner. Associate Chemist, U. S. Bureau of Mines, Kash.,, D. C.; formerly with U. S Bureau of Chemistry. * * Raymond M. H a m , Industrial Fellow, Mellon Institute, Pittsburgh, Pa.; formerly with U. S. Bureau of Chemistry. I H a n n and Berliner: J. .Im. Chem. SOC., 47, I709 (1925).
THE HYDROFLUORIDES OF ORGANIC BASES
1I43
CH3.XH?.HF)?.7H?O, According to the other lat’ter suppositions, the salt would contain 4 2 . 7 1 per cent of hydrofluoric acid or equivalent to the formula CP,H~.CH3.KHz.4HF. It was noted that upon heating o-toluidine hydrofluoride it readily sublimed and upon analyzing the sublimed product it was found to be of the identical composition as before sublimation. A portion of this o-toluidine derivative was allowed to stand in a vacuum desiccator over PzOj for eight days and another portion placed in an electric d q i n g oven at 106°C. for 14 hours. These samples, upon analysis, revealed a nitrogen content (7.3 j per cent and 7 . 4 2 per cent, respectively) corresponding almost exactly with the previous analysis,- 7.40% N. These results seemed to indicate conclusively that the material did not contain any water of crystallization and that it was probable that four molecules of hydrofluoric acid combined with o-toluidine and iodo-o-toluidine. A very careful survey of the available literature revealed the fact that little work had been done upon the reactions of hydrofluoric acids and organic bases. Beilstein? reports aniline hydrofluoride as C6HjXH?.HF,and cites a paper by Beamer and Clark3 published in 1879. These authors during the course of study of common salts of aniline report as follows on aniline hydrofluoride. “Splendid pearl white scales mere formed by the mixing of aniline and hydrofluoric acid. The salt is easily soluble in boiling alcohol, however, separating so rapidly on cooling from solution that i t stops up the mouth of the funnel through which it is being filtered.” S o analysis is given or formula presented and it appears as if Beilstein had himself introduced the above formula. Wallach4, a few years later, in discussing the preparation of organic fluorine compounds, indicates the preparation of crystalline aniline hydrofluoride (see p. 260) but gives no analysis or further description. No other work appeared until 1903, when Hantzch and Vock5 reported on some diazonium hydrofluorides which mere water soluble, quite unstable, and were capable of undergoing the general diazonium reactions. They gave meager analytical data and assumed that the compounds formed were of the general formula R.K2.F.HF, viz., C 6 H 2 , B r ~ . N ~ , F . H F . 2 Hand ~0 C,jH4.NO?.iYz.HF.H?0. A few years later Weinland and LewkowitzGpublished a paper on hydrofluorides of some anilides and substituted anilines, an extension of Weinland’s earlier investigation’ upon some double salts of hydrofluoric acid. Somewhat later Weinland with Reischles reported on some hydrofluorides of weak organic bases. A more manifest discussion of the work of Keinland and his
* Beilstein:
“Handbuch der anorganischen Chemie”, 3rd Ed., 2, 310 Beamer and Clarke: Ber., 12, 1067 (1879). ‘Wallach: .Inn.,235, z j j (1886). Hantzsch and Vock: Ber., 36, 2059 (1903). Weinland and Lenkonita: Z. anorg. Chem., 45, 39 (19051. .Inn., 315, 357 (1901); 328, 1 4 j , 149 (1903). Keinland and Reischle: Ber., 41, 3671 (1908).
’
1144
J. F. T. BERLINER AND RAYMOND
M. HANN
collaborators is possible after presentation of our own analytical and experimental results. Since no further record of hydrofluoric acid derivatives could be found and in view of the unusual results noted with the previously mentioned bases, a study of the compounds of hydrofluoric acid and various types of organic bases was instituted. Derivatives of over twenty-five bases were prepared, carefully purified and analyzed. The analyses of these compounds in every instance indicated the derivative to be of the type B.4HF where “B” represents an aliphatic or aromatic primary, secondary, or tertiary base. Similarly the diamines reacted with hydrofluoric acid to form derivatives of the type R.(KH2)2.(4HF)2. Compounds of this general composition were prepared in dilute and concentrated, heated and cooled aqueous, ethereal, and alcoholic solutions of hydrofluoric acid. The results were concordant throughout and indicate that in some manner organic bases add on four molecules of hydrofluoric acid.
Experimental Preparation and Properties. The general procedure employed in the preparation of the hydrofluoric acid derivatives was to add 0.4 to 0 . 7 moles (IS to 2 5 cc.) of pure 48 per cent hydrofluoric acid directly to the base or to an ether or acetone solution containing 0.1 mole of the purified base. The reaction was exothermic, and upon cooling the hydrofluoride separated in nearly quantitative yields. The acid salts are readily recrystallized from hot dilute aqueous solutions of hydrofluoric acid, separating as colorless crystalline compounds. As a rule three recrystallizations were made although analyses indicated that a single recrystallization gave a pure product. With the exception of the dimethyl and diethyl amines, pyridine, quinoline and quinaldine, all the hydrofluorides crystallized well. Even monomethyl aniline and monoethyl aniline whose mineral acid salts are not crystallizable, readily form beautiful crystals of the hydrofluoric acid derivatives. A rather remarkable property noted was that nearly all the hydrofluorides sublimed unchanged upon heating. Hydrofluorides of metanilic and sulfanilic acids and of o-toluidine decomposed. The only hydrofluorides examined which gave a characteristic melting point were m-nitraniline (m.p. 2 0 7 - 9 O C . cor.), p-nitraniline (m.p. 173-4’C. cor.), anthranilic acid (m.p. 217-18°C. cor.) and ethyl aniline (m.p. 170-1%. cor.). I.
Analytical (a) Nitrogen All the hydrofluorides prepared were analyzed for nitrogen by methods of the Association of Official Agricultural chemist^.^ These results are presented in Table I and it will be noted with what close agreement our results correspond with those required by the general formula B .4HF. 2.
“Methods of Analysis”, A. 0. A. C. and Edition
(1925).
THE HYDROFLUORIDES O F ORGANIC BASES
TABLE I Pr'itrogen Analyses on Hydrofluorides Compound
Aniline o-Toluidine m" pJ' m-Nitraniline p'J m-Xylidine Cymidine +Cumidine o-Tolidine Di-phenylamine Tri-phenyl Guanidine p-Phenetidine o-Anisidine p-Xitro-o-Anisidine a-Xaphthylamine
p-
"
Iodo-o-Toluidine Anthranilic Acid Methyl Aniline Ethyl Aniline hletanilic Acid Sulphanilic Acid
Analysis Weight of Cubic centimeters sample taken of N/xo Acid for analysis consumed
,1065
. I330 ,1030 , I847 ,2152
,1274 ,1197 ,1052
I633 ,1074 '
'1336 ,1006 ,1087 ,1442 '1339 ,1029 . I304 ,1205
,1160 ,2145 ,1256 ' I093 ,1329
6. I 7 .o 5.5 9.8 19.6 11.4 6.0 4.6
7.7 5.8 5.4 7.6
5.0 7.2
10.6 4.6 5.8 3.8 5.3 11.4 6.2 4.3 5.3
% Ir'
found
8.02 7.36 7 ' 48 7.43 12.76 12'53 7.02
6.12 6.60 7.56 5.66 10.58 6.44 6.99 11.09 6.26 6.23 4.41 6.40 7.44 6.91 5.51 5.58
%X
theory
B .4HF 8.09 7.48 7.48 7.48 12.84 12.84 6.96 6.11 6.51 7.52
5.62 10.44 6.44 6.89 11.28 6.28 6.28 4.47 6.45 7.48 6.96 5.54 5.54
(b) Carbon While the determination of nitrogen content presented a rapid and convenient method of analysis, a more complete analysis was considered essential. Attempts to determine carbon and hydrogen by combustion in the usual manner failed to lead to concordant results. The hydrofluoric acid was apparently not oxidized under the conditions; but passed on into the absorption train and was absorbed, yielding high results for carbon and low results for hydrogen. h second contributing factor t o discordant results was traced to reaction of the unoxidized hydrofluoric acid with the combustion tube. About this time JTilde and LochtelO published their work upon a rapid method for the determination of carbon in organic compounds. I n the described procedure the compound is burned in the presence of compressed oxygen in a calorimeter containing standard barium hydroxide solution. The carbonate formed is determined by differential titration using phenolphthalein and methyl orange. Several hydrofluorides were successfully lo
Wilde and Lochte: J. Am. Chem. SOC., 47, 440 (1925).
1146
J. F. T. BERLINER AND RAYMOSD hl. HANN
analyzed by a modification of this procedure and incorporation of Lindner's" process for the determination of carbonate.12 Following combustion in presence of an excess of Ba(OH)2,the solution was titrated in the cold until colorless to phenolphthalein. The solution was now transferred to a 350 cc. Erlenmeyer flask with a measured excess of .zKhydrochloric acid, a reflux condenser attached and the solution boiled gently for I j minutes. The excess of hydrochloric acid was now titrated back with Ba(OH)* and the volume of acid consumed by the carbonate obtained by difference. This procedure was necessary because barium fluoride reacts acid in water solution, preventing direct titration. The summary of the analytical results on carbon is given in Table 11.
TABLE I1 Carbon Analyses on Hydrofluorides P
Weight sample
Hydrofluoride
Aniline
o-Toluidine
Cc. SI5 HC1 Keight C consumed
a b
,1715 .1j22
j9.26 j2.68
C
,1155
40.20
,1212
45.31 63.64
a b Ethyl aniline a b p-Phenetidine a b $-Cymidine a b
,1704
.136~ ,1311 ,1693 ,1263 ,1164 ,1394
Found
/c
c Calculated B.4HF
41.46 41,54 41.77
41.59
44.87
44.89
53.98 51.92
.05438 ,07637 ,06479 ,06232
47.64 47.54
47.73
62.98
,07559
.oj6oj ,06190
44.65 44.38 j3.18 j3.21
44.21
46.70
51.58 61.80
,07110
,06322 ,04824
,07417
44.82
52.36
(c) Fluorine No particular trouble was experienced in the determination of the fluorine content of these conipounds, since the attached hydrofluoric acid may be removed by solution in strong alkali. The sample (.. gm.) was treated with 3 cc. of 4 s HOH, water added to a volume of 2 0 cc., the solution adjusted to methyl red neutrality with dilute HCI, j cc. of iV Na2COs added and then heated to boiling. Ten cc. of 20 per cent CaCL solution was now added and the solution allowed to digest on the steam bath for 1,!4 hour. This was cooled to s4'C., 3 to j cc. of concentrated acetic acid added and after allowing it to cool to about ~ j - z o ' c . , filtered through a weighed Gooch crucible, washed sparingly with cold dilute acetic acid, dried and weighed. The filtrate was measured and the loss for solubility of CaFl in filtrate corrected for (.0016 gms. per so0 cc.). The results are tabulated in Table 111. Lindner: Ber., 5 5 , 2025 (1922). We are indebted to Dr H. L. Loehte and X r , G. Decherd, of the Cnlverslty of Texas for devising the method and performing the analyses on carbon. 11
12
THE HYDROFLUORIDES O F ORGASIC BASES
TABLE I11 Fluorine Analyses on Hydrofluorides 7cF Hydro fluoride
a b a C ymidine b a p-Phenetidine b a lIethy1 aniline b a ni-Sitraniline b o-Tolidine a b 0-Saphthylamine a b Aniline
Weight Sample T e i g h t CaFz ,2000
Found Calculated B . 4HF
,1804 . I801 .1362 ,1362
43.90 43 ' 83 33.15 33.15
43.91
,1436 ,1432 .1670 ,1666
34,95 34.85 40.64 40.55
3j.00
,1431 , I429 ,1677 .16;5 , I397 ' 1396
34.83 34.78 40.81 40. j 6 35.00 33.98
34.85
33 16 '
40.62
40.84 34 06
From these data it is apparent that the nitrogen, carbon and fluorine contents of a variety of hydrofluorides agree remarkably well with a general formula B . 4HF. It is noteworthy that o-tolidine, a diamine, adds eight moles of H F per mole of base. (d) Titration studies Compounds of strong acids with weak bases are normally titratable with a strong base. I t was observed that the addition compounds under discussion dissolved in warm water to yield acid solutions, presumably due to a partial hydrolysis of the compound. That some hydrolysis actually takes place is quite apparent in the instance of the nitraniline hydrofluorides which on contact with water are immediately transformed froin colorless to yellow crystals, yielding yellow solutions. rln attempt was made to titrate solutions containing known amounts of these hydrofluoride salts with various strengths of standard sodium hydroxide solutions and using indicators with transformations in several ranges of hydrogen ion concentration. All attempts led to the general result that when an indicator having a range above a pH of 6 was employed only three-fourths of the combined acid (on the assumption of 4 H F ) could be titrated. A great number of trials were made varying the several conditions such as temperature, concentration of salt and alkali, solvent, addition of various neutral salts, and the use of indicators corresponding to several ranges of acidity. The results of all these experiments were analogous-only threejozoths of the hydrofluoric acid could be accounted for. K i t h indicators whose transformations corresponded to ranges higher than a pH of 0, the amount of acid that appeared titratable increased, but the results were extremely erratic. Determinations varied from 7 j per cent
I 148
J. F. T . BERLINER AND RAYMOND M. HAXN
to 85 per cent and in one instance as high as 92 per cent of the calculated amount of hydrofluoric acid present in the compound. I n these higher ranges the end points of the titrations were never sharp and in most cases extended Over unusually wide ranges. It was thought that this extraordinary behavior was due to some peculiar property of hydrofluoric acid and a search was made for reports of previous titrations of this acid. The only study that was found that dealt specifically with this subject was by Haga and OsakaI3 who titrated various strengths of
FIG.I Typical electrometric titration curve for the organic base hydrofluorides Dashed vertical lines represent moles of alkali per mole of hydrofluorides.
hydrofluoric acid in aqueous solution with alkali using a number of indicators. I n the use of litmus as an indicator the results obtained caused the authors to suggest that the molecule of hydrofluoric acid might be HzF2,H3F3 or H ~ F I . The analysis of the neutralized solution indicated a composition of K3HF4. However, their results show that in general all the hydrofluoric acid present can be readily and completely titrated by the indicator method and that a number of indicators may be employed, (phenolphthalein was considered the most satisfactory). From our previous results and the above contribution it seemed apparent the peculiar behavior noted was due primarily to the types of compound with which we were dealing. These data indicate that one of the hydrogen fluoride molecules of the (HF)*complex is in combination in a manner differing from that of the other three. Since coordinate results could not be obtained by the indicator method of titration, a series of electrometric titrations were made with the hope of being able to interpret the above data and gain an insight into the mechanism of the complete neutralization of these compounds. l3
Haga and Osaka: J. Chem. SOC.,67,251-255 (1895).
THE HYDROFLUORIDES OF ORGANIC BASES
1149
These titrations were made using a normal calomel cell and the Hildebrand type of hydrogen electrode. The volume of the solution was approximately IOO cc. a t the start and the concentration of the salt.about .OI N Several of the salts were very carefully titrated electrometrically, and the results of the electromotive force measurements converted into their equivalent pH values and plotted. It was of interest to note that all the compounds thus titrated (9) yielded analogous types of neutralization curves. h typical curve is represented in Fig. I . This curve, representing the titration of 0.1966 grams (0.001 136 moles) of aniline hydrofluoride with 0. I X NaOH, has several interesting and noteworthy features. The acid nature of the solution, undoubtedly due t o hydrolysis, is evidenced by the high hydrogen ion concentration (low pH) of the initial solution. As the alkali is added neutralization takes place smoothly until an equivalent of three moles of the alkali have been added or threefourths of the available hydrofluoric acid accounted for. As neutralization proceeds from this point, a very peculiar condition appears. Instead of the curve rising rapidly to a value of pH of about 10.0 to 11.0 (within the range of 0.1% excess alkali) as is the case in the course of a normal type of neutralization, there is a sharp slowing up in the rate of the decrease of the hydrogen ion concentration in a range represented by about pH 7. j-8.0. From this point the addition of alkali is in a nearly linear relation to the increase in the value of pH. The equivalent of approximately one more mole of alkali, making a total of four moles, must be added to bring excess alkali in a the curve to the range attained by the addition of 0.17~ normal type of neutralization under the above conditions. Further addition of alkali causes the curve to act in the normal manner. From an inspection of this curve which is typical for all the compounds titrated electrometrically it will be immediately apparent why difficulties were encountered in titrations by the indicator method. It may be noted at once that for indicators whose point of equilibrium in the color transformation range represented by a pH value in the proximity of 8 will give an apparent end point at 7 ~ of7 the~ total available acid while those indicators having transformations a t lower hydrogen ion concentrations will show a proportionately higher acid content. The very wide range of neutralization between the alkali equivalent of three and four moles shows that the end point of any titration that lies within this range will be extremely difficult to distinguish with an accuracy closer than I O to 2 0 7 ~of the total alkali necessary to cause complete neutralization. I t is, therefore, quite apparent that the base hydrofluorides cannot be completely titrated satisfactorily by means of the indicator method. However, by the use of indicators of the proper range it would be possible to titrate exactly 7 jYc of the total acid present in the compounds. I n Table IV some of the numerical results for the electrometric titrations of some of the compounds are given.
J . F. T. BERLISER AZiD R A Y U O S D M. H A S S
I150
TABLEIV Typical Electrometric Titration Results Aniline hydrofluoride 0.1966g. in roo cc.
cc N / I O SaOH 0
5 IO
15 20 25
30 32 34 35 35.5 36 36,s 37 37.5 38 39 40 41 42 43 44
Cymidine hydrofluoride 0.1689g. in roo
p-Phenetidine hydrofluoride 0.1 I j6 g. in roo cc
CC
Ph
2,30 3.07 3.52 3.89 4.I 7 4.42 4. i 4
4.90 5.30 5 93 7 . 2 2
8.16 8.51
cc S/ro SaOH 0
3.19
4
I3
3.60 3 ' 92 4 . I9 439 4.68 4.92
11
5.22
I4 j
5,42
j.27
22.5
5.73 7.8j 8 . 70 9.07 9.31 9.58
21
23 23.5 24
24.5
9.00 9.17 9.46 9.73
26 26.5
IO.0 1
27.5
10.31 10.59
25.5
27
45 46 47
10.95
48
11.24
50
11.36
28 28. j 29 29.5 30 31 32 36
55
11.54
40
60
11.68
50
11.07
2.63
2
22
16 20
Ph
0
I2
8 I2
25
11.16
cc S/IO SaOH
2.73 3.47 3.84 4. I7 4.42 4.76 4.93
4
8.78
10.80
Ph
6 8 IO
15
IS
5
9.80
16 16.j '7
IO.00
Ii.5
10.23 10.42 10.58 IO. 70
I8 18.5 I9 19.5
10.8; 10.95 11.03
20
5.66 598 6.76 8.54 9 . IO 9.43 9.71 9.97 10.24
11.10
21.5
10.49 10.71 10.86 10.98 11.08
11.20
22
11.1j
23
11.27
24
11.36 11.67 11.96
11.30 11.55 11.66 11.84
20.5 21
30 45
A consideration of the reasons for this remarkable behavior is beyond the present scope of this paper. It is incumbent, however, t o prove that the effect is due to a peculiar property of these addition compounds only, and not to some extraneous factor. The factor that might be considered to be the most probable cause of any such deviation from the normal is that of the hydrofluoric acid, another consideration is that of the mechanism of neutralization of organic base salts such as the hydrochlorides. I n view of these possibilities, a series of electrometric titrations were made on aqueous solutions of hydrofluoric acid of various strengths (0.001, 0.01
THE HYDROFLUORIDES OF O R G A S I C BASES
1151
and 0.13)and on aqueous solutions of aniline hydrochloride. Both of these materials give results strictly analogous to the normal strong acid-base neutralization. I n Fig. z are represented the superimposed neutralization curves of aniline hydrofluoride and hydrofluoric acid (o.rN). The curve for aniline hydrochloride has been omitted since it so closely parallels that of the hydrofluoric acid.
MOLFS Of ALKALI
FIG.2 Comparison of electrometric titration curves of hydrofluoric acid (A) and aniline hydrofluoride (B).
Consequently, it may be stated that without doubt the peculiar characteristics noted in the mechanism of neutralization of the hydrofluorides of the organic bases are an intrinsic and individual property of these compounds, and that the separate compound constituents of these salts behave in the normal manner. I n consideration of the above contentions and experimental results, it is postulated that in these organic base hydrofluorides there are present four molecules of hydrofluoric acid, three of which are bound to the molecule in a manner distinctly differing from the fourth. The constitutional relations of three molecules of the acid to the base are identical. I n solution, doubtlessly, these three molecules are hydrolyzed from the base simultaneously and t o the same extent, and the neutralization in relation to these proceeds in a normal manner. However, at the very point of the neutralization of these three molecules of hydrofluoric acid there is left in solution the base hydrofluoride with the one molecule of hydrofluoric acid that is bound in an entirely different manner to the base than the other three molecules of acid. From this point on, the course of the neutralization is distinctly not normal, and it is not within the domains of this discussion to attempt an explanation of the actual mechanism of the neutralization of this last molecule of hydrofluoric acid.
1152
J. F. T. BERLINER AND RAYMOND M. HANN
That the first three molecules of hydrofluoric acid attached t o the organic base are not individually neutralized or rather hydrolyzed from the base is evidenced by the smooth, unbroken portion of the electrometric titration curve up to the utilization of about 7 5 n of the available acid. A consideration of the possible structural constitution of these compounds based on the above results will be found in a subsequent portion of this discussion. Discussion of the Work of Weinland and Lewkowitz6 and of Weinland and Reischle * I n entering upon a discussion of the work of these authors, it is our intention to merely point out possible sources of agreement between their data and our own. These authors investigated the action of hydrofluoric acid on aromatic anilides and amines and described in all twenty-four compounds, of which about half are of the amine type. I t is with this latter group that we are a t present concerned. They described the following compounds: Aniline monofluorhydrate (C6HjXH2,HF.I .gH20), aniline difluorhydrate (CsHjNH2.zHF),aniline trifluI .gH20),p-amidophenol trifluorhydrate (CsH*(OH) orhydrate (C6Hs?\TH2.3HF, K H z . ~ H F . I . ~ H,~ O) p-amidophenetol difluorhydrate ( C ~ H ~ ( O C ~ H S ) N H ~ . Z H F ) , p-amido-benzoic acid difluorhydrate (C,H,(COOH) NH2.zHF),p-bromaniline tri-fluorhydrate (CsH4BrXH2.3HF.Hz0),and 0-, m-, and p-nitraniline trifluorhydrates (C6Ha(K02)?JH2~3HF.H20), triphenyl amine monohydrofluoride (C&)39.HF) and diphenyl amine, di- and trihydrofluorides ( C S H ~ ) ~ NH.2HF and (C6Hj)2NH.3HF). Since such diversity of combination was in sharp contrast to the uniformity which we had experienced, an examination of the analytical and experimental data of Weinland and Lewkowitz was undertaken. This survey revealed the fact that in a number of instances, the formulae assigned to the compounds were based upon titration values with tenth-normal alkali. It has been demonstrated that such titration allows the estimation of but about threefourths the combined hydrofluoric acid. Formulae based on the assumption of total titration would therefore be incorrect. Another source of difference between our compounds and some of those of Weinland and his collaborators is that in all cases our derivatives were prepared in the presence of an excess of acid while this was not the case with several of the preparations described by the above authors. For one of the compounds of the amine type, complete analytical results on carbon, hydrogen, and fluorine are given. N o indication is supplied regarding the procedure employed, and from a consideration of the difficulties enumerated in a previous section regarding the estimation of carbon and hydrogen in these compounds, we would prefer not to discuss this compound. Four of the remaining compounds were not prepared in an excess of acid, namely aniline mono- and dihydrofluorides, triphenylamine monohydrofluoride and diphenylamine dihydrofluoride. I n the instance of two other
THE HYDROFLUORIDES OF ORGASIC BASES
1153
derivatives no indication is given of the amount of acid used, but evidently there was not an excess of acid (p-amido-phenetol dihydrofluoride and pamido benzoic acid dihydrofluoride) . Eliminating these hydrofluoride derivatives from further consideration, but six compounds remain to be discussed. These will be separately considered. I , Aniline hydrofluoride.--The empirical formula for this compound is given as C6H51YH2.3HF.+H20,based on four titrations with N / I O KOH which averaged, show a hydrofluoric content of 36.6%. Assuming that but three-fourths of the acid was titrated and the empirical formula was that of the cornpound described in this paper, viz. C6H5SH?(HF)r, the result theoretically mould be 34.770 which is reasonably close to the above result. There is no analytical evidence given which could possibly allow the assumption of the constitution of the compound t o be as presented. 2, p-Bromaniline hydrofluoride. --The empirical formula proposed for this compound is CsH4,Br,SH2.3HF,H?0,and as in the previous case, the only analytical data presented is t’hat of titration with 1 / 1 0 KOH. It is immediately noted that the molecular weight corresponding to the above would be almost identical with that corresponding to CBHr.Br.T\rTH2.AHF. Therefore, assuming that but three-fourths of the total acid was titrated, the latter formula may be considered the more correct. 3. (a, b and c ) 0-, m-, and p-nitroaniline hydrofluorides.-The empirical formulae presented for these compounds are the same, and correspond to C s H a . S 0 2 . ~ ~ ? . 3 H F , H 2This 0 . is likewise based on acidimetric titrations. I n the case of the ortho isomer, a nitrogen analysis is included which agrees very well with our result (IT. &. L. 1 2 . 8 % ~ our value 12.76%). As in the instance of the p-bromaniline hydrofluoride, the molecular weight of and these compounds is almost identical with that of C~HI,NOYIKH~(KF)~, since the nitrogen analysis corresponds with this formula as well as for that given and the formula proposed by Keinland! et al., is based on the incorrect assumption that all the hydrofluoric acid was titrated, it is contended that the compounds he was describing are of the type C6H4X0,~KH2(HF),. 4. Diphenyl amine hydrofluoride.--The empirical formula proposed for this compound is (CsH5)2.KH.3HF based on titration data and a nitrogen determination. The nitrogen result lies practically midway between the calculated value for the above compound and that for (CeHs)z,NH.sHF. Since the titration values represent but three-fourths of the total acid, the proposed composition is considered erroneous. The compound has the composition represented by the tetrahydrofluoride. From this brief summation it appears evident that the analytical data presented by Feinland and his coworkers to establish 6he empirical composition of the hydrofluorides of the various amines are in agreement with our results and support the contention that when an excess of hydrofluoric acid acts on any amine, a compound of the general type R.?r”p,(HF)4is formed.
1154
J. F. T. BERLISER AND RAYMOSD M. HANN
It can be shown by a recalculation of the data that the analytical results on the aniline hydrofluoride are also in concordance with this precept. However, since we have not prepared any of this type of compound, we do not desire to enter into a discussion of this phase at the present.
The Abnormality of Hydrofluoric Acid The present conflicting evidence of the nature of hydrofluoric acid and the limited knowledge regarding its reactions make any discussion of the constitution of the compounds of the type considered in this paper, rather uncertain. Much work has been done on anhydrous hydrofluoric acid. both in the gaseous and liquid states. It has been shown through vapor pressure measurements that the vapor has a molecular weight of 37.32 at 30 j"C. indicative of a molecular formula (HF)214.Simons and HildebrandI5 have shown the molecular weight of the vapor to be 39.74 at 305.1'K (40 at 30j"K) and from data on the vapor pressure of highly purified anhydrous hydrofluoric acid they calculated that at any temperature it may consist of an equilibrium mixture of the isomers HF and (HF)G. They present evidence to indicate that the polymerization is exothermic and may be represented as 6 HF
=
(HF)s
+ 40 cals.
Iron Kartenberg and Fitzneri6 have recently correlated and made a study of the thermochemical data on fluorine using the above assumption of Simons and Hildebrand in order to correct for the degree of polymerization of hydrofluoric acid in the calculation of the heat of solution and other derived values. There has been quite some discussion upon the molecular complexity of hydrofluoric acid, as evidenced by its abnormal boiling point, by Forcrand" and by Berthoud1gi9. The former is of the opinion that the abnormal boiling point of hydrofluoric acid as well as that of water and of ammonia are not due to molecular association but to the high heat of formation of these compounds due to a di-symmetry of their molecules. Berthoud, however, does not accept this view and shows that the abnormally high points of ebullition are readily accounted for by the polarity of these substances and that molecular asymmetry is not sufficient to cause such high boiling points. Recently Kolossowsky20has pointed out an error in Simons' calculations of the heat of vaporization, thereby showing that it is in agreement with previous determinations. From the extrapolated vapor density data of Thorpe and Hambly and applying Trouton's rule according to the principle of DeHeen, he arrives at the remarkable conclusion that hydrofluoric acid boils as a normal, nonMallet: Am. Chem. J., 3, 189 (1881). Simons and Hildebrand: J. Am. Chem. SOC., 46, 2187 (1924). 16 Von Wartenberg and Fitsner: 2. anorg. allgem. Chem., 151, 313 (1926). 17Forcrand: J. Chim. phys., 15, 517(1917). 18Berthoud: J. Chim. phys., 15, 3 (1917). LgBerthoud: J. Chim. phys., 16, 245 (1918). *O Kolossowsky: Bull., (6),51-52, 422-28 (1927). 14
15
THE HYDROFLL-ORIDES O F ORGANIC BASES
IIjj
associated liquid, i.e. its degree of association at its boiling point and in the saturated vapor at the same temperature is the same. This association is shown to correspond to the empirical formula H3F3. HengleinZ1in a study of the physical properties and molecular structure of the halogens and their compounds finds that all of these substances, with the exception of hydrofluoric acid, are in very fine agreement with the rule proposed by BiltzZ1,Le., that the physical properties (melting point, optical properties, thermal relations, etc.) of the halogens and their compounds are in linear relationship to the molecular volume. However, hydrofluoric acid was so abnormally different that it TTas completely excluded. While much study has been devoted t o the anhydrous acid, relatively slight consideration has been given to the molecular complexity of hydrofluoric acid in aqueous and other solutions. Some work has been done on the conductivity and activity of hydrofluoric acid in aqueous solution. Wynne- Jones and HuddlestonZ3found that for concentrations up to 0.4 S the activities of the hydrogen ions were in concordance with the calculated values while above this concentration a marked abnormal deviation occurred. Kreman and DecolleZ4from conductivity data rendered an opinion that hydrofluoric acid is dibasic and that fluorine is divalent, while WaldenZ5also measured the conductivity relations and on the contrary postulated its monobasicity. By studying the change of conductivity of hydrofluoric acid solutions during neutralization Pellini and Pegoraroz6 found that a minimum in the specific conductivity occurred at 0.5 moles of alkali and contended that this proved the basicity of the acid. However, BruniZ7has shown that the results of Pellini and Pegoraro were merely fortuitous, for their results would have been different if any other concentration of acid had been investigated. Bruni shows that the minimum of the specific conductivity curve varies with the concentration. However, there seems to be more or less complete agreement among investigators that in aqueous solution the acid exists partly as the HFZ ion. Pickz8assumed that HF?existed in solution while there was no appreciable amount of H ~ F present z and the complex ion was formed by the dissociation of mono-basic hydrofluoric acid, and the reaction of the fluorine ion with undissociated mono-basic hyfluoric acid. He calculated the equilibrium constant of this reaction (HF + F = HA) t o be 5.5 and also pointed out that the formation of the complex ion was independent of the acid concen-
Henglein: Z. anorg. allgem. Chem., 118, 165 ( 1 9 2 1 ) . 2z Biltr: Z. anorg. allgem. Chem., 115, 241 ( 1 9 2 1 ) . 23 Wynne-Jones and Huddleston: J. Chem. SOC.,125, 1031 ( 1 9 2 4 ) . z4 Kreman and DecoUe: Monatsheft, 28, 9 1 7 (1907). 25Walden: Z. physik Chem., 2, 58 (1888). 26 Pellini and Pegoraro: Atti Accad. Lincei, 16 11, 273-279 (1907); also Z. Elektrochemie, 21
13, 621-22 (1907). 27
zB
Bruni: Z. physik. Chem., 69, 69-74 (1909). Pick: Ternst Festschrift, 360-73 ( 1 9 1 2 ) .
11j6
J. E'. T. BEIiLISER A S D RAYMOND 11. HAS?:
tration. Davies and Huddleston29 confirmed the existence of this complex i m between the concentrations of 0.16s and 2 . 3 s and calculated the equi~ ~ aslibrium constant t o bc, 4.7. Further Anthony and H u d d l e ~ t o nhave certained by freezing point methods that no appreciable amount of H2F2 exists in solution X very good discussion of the general physical properties of aqueous solutions of hydrofluoric acid will he found in the new edition of Gmelin's "Hantlbuch der anorganischen C'hemie" ( S o . 5 . 8th Edition, 1926). Pick's contribution contains an excellent discussion and summary of the knowledge of the molecular composition and ionization of aqueous solutions of hydrofluoric acid. In the instance of thc compounds with organic bases rThich we have described, it appears practically inipossible to explain the observed phenomena by application of the above briefly outlined investigations. At present, attention to these data merely emphasizes the peculiar and abnormal behavior of hydrofluoric acid. In proceeding to a survey of the general reactions of this acid, it may be noted that usually when a monatomic fluoride is allowed to crystallize from a n excess of hydrofluoric acid, an acid salt is obtained. I n a surprisingly large number of instances the acid compounds may be considered as salts of the hypothetical quadribasic hydrofluoric molecule H4Fa. Many double salts may also be refrrred to as derivatives of a parent substance H4F4 and listed below are empirical formulae of numerous types of compounds which may be so considered. SaF.3HF31;E(F.3HF3l;RbF.3HF32;LiF,3HF33;nZgF2.zSaFR1;~\lgFz,2KF36; ZnF2.zKF3"."'; ReF2,z?;H4F3J; BeF2.2SaF33;and Poulenc40 has prepared a series of double salts of nickel and cobalt with alkali fluorides of general type ?;iF2.zNH4F; CoF2,zT\'aF, as well as 3 series containing MnF2 in place of the alkali fluoride such as CoF2,MnF2. Weber4' has described a similar series containing ferrous fluoride, i.e., FeF2 zM'F. A rather detailed compilation of double salts of hydrofluoric acid has been made by Bartexkoa2and also by Wellsd3. In their papers are listed many Davies and Huddleston: J. Chem. Soc., 125, 260 (1924). 3°.Anthonv and Huddleston: J. Chem. Soc., 127, 1 1 2 2 ( 1 9 2 j ) . " d e Forcrand: 2. anorg. Chem., 28. 384 (1903); Guntz: Bull., ( 3 ) 13. 114 (1895); Moissan: Compt. rend., 106. 547 ( 1 8 8 1 ) ;.Ann. Chim. Phys., (61 3, j (1884). 32 Eggelin and hleyers: Z. anorg. Chem., 46, 174 (1905). Chahrie: Compt. rend., 132, 680 (1901). B4 S e t t o : Z. angea. Chem., 4 , 4j ( 1 8 9 0 ~ . 3 j Duboin: Compt. r e n d , 56, 848 (1863). 36 JVagnrr and von Helmholtz: Ber., 19, 896 ( 1 8 8 6 ) . 3 7 Berzelius: Pogg. :inn.. 1, 2 2 . 26 (1824!. 3 8 de Marignac: Arch. Sei. G e n h e , (4) 3 0 , 4j (18731. 3 9 Gav 1,ussac and Thknard: a'Recherchesphysico-chimiques" (181I j; Berzelius: Pogg. .Inn., 8,"187 (1826:; von .i\r.dejeff: 5 6 , I O I (1842). 40Poulenc: Ann. Chim. Phys.. ' 7 1 2 . 47 (1894'. 41Weher: J. prakt. Chem.. S O , 2 1 2 (1860?. brr Iloppelfinoride" , 1909). 4 3 l'--l!s: Am, Cheni. .J., 2 6 . 389 !1901:. 20
THE HYDROFLUORIDES OF ORGANIC BASES
IIji
derivatives which, empirically a t least, may be considered as derived from a hypothetical acid H4F4. The alkali metals, silver and cuprous copper form salts KM’aF4 where M’ represents any univalent basic substance. Divalent metals yield compounds such as Mg M“F4 while trivalent’ metals produce double salts, an example of which is SbAf’F4. Marked variation in composition may occur even in these major divisions t o form salts such as KzM’IF~, KIhi’F,,,lfgh!I’2F4, etc. Tetravalent and quinquevalent salts may be considered as derived from H4F4 since they form compounds ZrM’,Fs and TaZvl’3F8. Berzelius3’ describes ZnF2.zA1F8(ZnA12F~) and Brannerd4prepared PbF4.3KF.HF(HPb”” KaFs). Bartzeko mentions oxide salts which he formulates as (ASO).M.FI, (YO) and (UO) ,M2.FI. Tellurium salts of this type have been investigated by Metznerlj and also by Prideaux and Millot46. Little work has been accomplished upon metallic organo fluorides. K r a ~ s has, e ~ ~however, prepared a number of alkyl tin double fluorides similar to (C2H5)2 SnFz. zKF. Of possible interest in connection with salts of this type are the general reactions: and
MO ,110
+ 4KF.HF = hfF2.zKF + z K F + z H F + H20 + 4NH4F = ;LIFZ.ZNH~F + z K H ~+ Hz0
advanced by Ruffd8in an extensive treatise on the chemistry of fluorine. I t has been mentioned that in every instance the compounds described in this paper were recrystallized in the presence of an exwss of hydrofluoric acid. I t is quite possible that an organic hydrofluoride of lower acid content than four moles might be obtained in the presence of an excess of the amine. However, the present evidence would strongly indicate that an excess of acid being present insures an acid-base combination containing four moles of acid for every amino group present. There seems t o be a striking analogy between these compounds and inorganic compounds resembling KF.3HF, which are prepared in a similar manner. That double salts should show a fairly strong tendency to form what may be conceived to be derivatives of the hgpothetical acid H,F, also seems significant. There are of course many double salts of the metals whose empirical formulae do not fall into line, but there is in general a decided analogy t o be found between a majority of double fluoride salts and the organic compounds we are discussing. The usual stability of the base hydrofluorides is paralleled by the behavior of inorganic fluorides, several of which are more stable as acid fluorides than as normalfluorides. Theearliest observations of this curious phenomenon were Branner: Z. anorg. Chem., 7 ; I f1894). hIetzner: Ann. Chim. Phys., ( 7 1 15, 203 (1898). 46 Prideaux and hlillot: J. Chem. SOC., 129, 520 (1926). 4i Krause: Ber., 51, 1447 (1918). 4 8 Ruff: “Die Chemie des Fluors”. 17 (1920); Mellor: “ A Comprehensivo Treatise on Inorganic and Theoretical Chemiatrv”, 2 , 129-141, 512-521 (1922); Abegg and Auerbach: “Handbuch der anorganischen Chemie”, 4 I1 (1913); Gmelin’s “Handbuch der anorganischen Chemie”, 8th Ed., S o . 5 (19261. 14
41
I158
J. F. T. BERLINER AND RAYMOND M. H A S S
made by Wiegleb49. BerzeliusJoand RoseS1made corresponding observations on the same salt, ammonium fluoride, and found that in hot aqueous solution ammonia is evolved and ammonium acid fluoride results. Even in moist air solid ammonium fluoride was transformed upon standing into the acid salt. There has been some discussion relative to the valence of fluorine and the structure of its compounds, but no satisfactory or even adaptable hypothesis has as yet been advanced. assumed that in double fluorides the halogen atoms are united t o the central atom or group so that the group acts as an acid anhydride such as (F.K.F)H or (F. (F,B:FIK F) in which the halogen atoms act as intermediate or connecting links between two elementary atoms or groups. Pauling and Hendricks” in their work on the structure of azides and cyanides by means of X-rays, discuss the possible configuration of hydrofluoric acid and its derivatives and remark that the acid fluoride ion (HF,)- may, because of the small charge on the proton as compared with the other kernels, possess a unique electronic arrangement. I t is generally believed that in the acid fluoride ion the proton is sharing only two electron pairs with other nuclei. The structure of the ion may be represented thus :
..
..
(: F : H : F :)-
..
..
The Electronic Structure of Organic Base Hydrofluorides The above section is a very brief r6sumC of the state of knowledge in regard to the structure of hydrofluoric acid derivat,ives. I n all this mass of data there seems to be no direct fact or hypothesis that may be utilized to assist in an explanation of the phenomena of the hydrofluoric acid salts of the organic bases. Yet throughout the presented data there may be discerned a relationship between the inorganic and the organic derivatives. 1here exists a strong tendency for hydrofluoric acid or inorganic fluorides to become attached to other inorganic fluorides and to form quite stable derivatives. Quite often these compounds contain four atoms of fluorine to the molecule. I n the acid fluorides the tendency is to form tri-hydrofluoric acid salts, viz. KF.3HF, when an excess of acid is employed. The latter property is directly analogous to the mode of formation of the organic compounds described in this paper. I n every case these were prepared in the presence of a n excess of hydrofluoric acid. Another analogy is the fact that, as aforementioned, in the organic hydrofluorides one of the acid groups appears to be Kiegleb: Creel’s “Die neuesten Entdeckungen in der Chemie”, 1, 1 3 ( 1 7 8 1 ) . Berzelius: “Lehrbuch del Chemie”, 3 . 282 (18561. “ R o s e : Pogg. Ann., 108, 19 (1859). 52 Kerner: “Seuere hnschauungen auf dem Gebiete der anorpnischen Chcmie,” 68 (19051. j3 Pauling and Hendricke: J. Am. Chem. Soc., 47, 2904 (1925);48, 643 (19261. p9
THE HYDROFLUORIDES OF ORGANIC BASES
1159
attached to the base in a manner differing in some respect from the accompanying other three. There may exist an interpretation of these results in the application of principles of molecular structure that have been developed by the studies of Lewis, Thompson and others. Fluorine, in common with other halogens, may be considered capable of possessing several valences, and as a matter of fact, its most common valence is a t present unknown. We may a t present assume the existence of the same valences as exhibited by the other halogens. Electronically, fluorine consists of the helium nucleus surrounded by seven orbital electrons. The tendency of such a configuration is to capture an additional electron in order to complete the octet shell. For this reason one would anticipate a valence of one. Many of the reactions of fluorine and its compounds are quite complex and difficult to explain on the basis of a rigidly univalent element. Even hydrogen is under certain conditions conceived to be divalent and in fact this circumstance may act,ually exist in hydrofluoric acid. This idea was suggested by Huggins&'and was also advanced by Latimer and Kodebushjj. It is interesting to note in this connection that liquids which are highly associated, and possess the high di-electric constant and ionizing power which apparently accompany this association, are substances containing the type of structure postulated for hydrofluoric acid. This idea of the bivalent hydrogen offers a rather clear conception of the association of such molecules as water, hydrogen peroxide, hydrocyanic acid, ammonia and hydrofluoric acid. Pauling and Hendricksj3, as well as several other investigators, have 9 .
..
considered the electronic structure of the hydrofluoric acid ion to be (.F:H :k':)..
..
Ivhere the hydrogen is considered to share four electrons with the fluorine atoms. It may be argued from this that fluorine has the power to cause hydrogen to share this number of electrons when they are in combination. In other words, the hydrogen is rendered strongly negative. S o w if one considers the molecule of hydrofluoric acid as normally represented, .. H:F: and assumes that through some agency it has been rendered relatively more positive than in the free condition, there would be a tendency for other molecules of hydrofluoric acid, through the negatively charged hydrogen atom, to attach themselves to this molecule in the manner represented beloly :
:F: H :H+:F:H:F: H :F: 54
55
Huggins: Science, 40, 679 (1922). Latimer and Rodebush: J. Am. Chem. SOC.,42, 1 4 1 9 (1920'.
1160
J. F. T. BERLINER A S D RAYMOND M. HA”
This type of structure is possible since even the octet does not necessarily have to be considered completely saturated, for though it cannot receive any more electrons, its own electrons may be shared in such a manner as to fill the shells in other atoms, thus completing their octets. As ThomsonS6says, “Any of the electrons in an octet might act in this way to fill gaps in the layers of electrons around an atom. Thus an octet might act as a nucleus from which chains and side chains of atoms ramified in every direction. As far as geometrical considerations are concerned, there is nothing to limit the number of certain atoms which could be linked together in this way.” I n this structure, it will be noted that there are four molecules of hydrofluoric acid and that one molecule differs from the other three. I n the organic hydrofluoridesthe first molecule to attach itself to the amino group is rendered relatively more positive than the other acid molecules, which is the condition under which this structure is assumed to be possible of formation. Similarly in the potassium fluoride, the fluorine atom attached to the potassium is rendered relatively positive and thereby allows the formation of a compound of the structure KF.3HF. :F: H :K:F ..: H : F..: H :F:
analogously
: .. F: H RNH? :H:F:H:F: .. .. H :F:
I n this conception of the structure of the organic compounds the reason why one of the acid molecules is not readily titrated and the other three are so easily determinable is apparent. The three hydrofluoric acid molecules attached through their hydrogen atoms are very readily hydrolyzed while the other acid group is strongly attached t o the amino group. This conception of fluoride derivatives would also allow such compounds as KF.HF, KF.2HF and the corresponding organic compounds, (Ref. to 6 and 8), for it would not be essential or necessary for all the electrons of the fluorine atom to be shared. This complete sharing would be the tendency in the presence of an excess of hydrofluoric acid. Probably other conceptions of the structure of these compounds could be formulated, but of the several which we have considered, the one just given presents a satisfactory explanation for ( I ) formation of compounds of the type B(HF)? or better B.HF 3HF and K F 3HF, ( 2 ) the possibility of formation of mono-, di- or tri-acid salts such as KF.HF, KF 2HF, etc., (3) the titration phenomena described in which the facts are presented t o show that but three-fourths of the total hydrofluoric acid present in the molecule can be titrated directly by indicator method and (4) the improbability that more than four molecules of hydrofluoric acid may combine with a basic group (or more than three molecules with the type KF). b6
Thomson: Phil. Nag., 41, j44 (1921)
THE HYDROFLUORIDES O F ORGANIC BASES
1161
The singular properties of fluorine, as compared to other halogens, are too well known to require comment, However, under certain circumstances there are exhibited similar types of reactions. That compounds containing more than one halogen acid molecule to each molecule of amine can exist and that such combination is a circumstance of environment has been shown by several investigators. Working with dry acids and dry bases Kaufler and K u ~ +produced hydroscopic salts which evolved acid in the air, from dimethyl amine, dimethyl aniline, dimethyl-otoluidine and other bases. I n a second paper the same authorss8 described the preparation of other salts of the general formula B.2HX from numerous aliphatic and aromatic amines and pointed out that the ability to form such s d t s is decreased by the introduction of negative groups in the molecules. This is in direct accord with the structure presented in this paper. Korcznskis9 has measured the quantity of pure dry hydrogen chloride absorbed by certain amines a t temperatures varying from room temperature to 7 j°C. and found that salts were formed containing two and three molecules of hydrogen chloride to each one of base. I n a second paper60 he extended this work and summarizes his conclusions ascribing the constitution of the bodies to the coordination value of nitrogen. From these facts it may be assumed that under the proper conditions the other halogen acids may form similar types of compounds as the hydrofluoric acid. Hydrochloric acid, being the halogen acid which, in electronic structure is most comparable to hydrofluoric acid, would naturally form acid salts far more readily than hydrobromic or hydriodic acids. No doubt, judging from the difficulty in preparing acid hydrochloride salts, it would be rather improbably that even relat,ively stable acid bromides or iodides could be prepared. summary Evidence has been presented to prove that whenever an organic amine is treated with an excess of hydrofluoric acid the resulting compound has a base to acid ratio of 1:4; i.e. an empirical composition of B.(HF)+ It has been demonstrated by electrometric titration studies that of the four molecules of acid attached to the base, three are bound and react in an identical manner, while the fourth differs distinctly from these in the nature of its reactions and union to the base. The work of previous investigators in this field has been discussed and their results, though apparently quite different, are shown to actually be in complete accord with those herein reported. An attempt has been made to present an explanation of this phenomena and to correlate this unusual behavior with the inorganic compounds of Kaufler and Kunz: Ber., 42, 3 8 j (1909). Kaufler and Kunz: Ber., 42, 2482 (1909) 18 Korcznski: Ber., 41, 4379 (1908). 6 o Korcznski: Ber , 43, 1820 (1910). 57
5B
1162
J. F. T. BERLINER AND RAYMOND M. HANX
hydrofluoric acid as well as with some of its general physical and chemical properties. A structure is proposed for this type of compound which appears to satisfactorily explain the formation of compounds of the type B.HF.3HF, the possibilities for the formation of compounds containing less but no more hydrofluoric acid and the fact that but three-fourths of the acid present is directly titratable acidimetrically.
Acknowledgment The authors wish to express their deep appreciation to Professor William Mansfield Clark for the many helpful suggestions regarding the electrometric studies of the compounds described and to Dr. William Blum for generously allowing us to use the electrometric apparatus and the facilities of the Electrochemical Division of the Cnited States Bureau of Standards.
