GO4
SAMUEL KAUFMAN AND C. R. SINOLETERRY
Vol. 50
THE REACTION BETWEEN TERTIARY AMINES AND ORGANIC ACIDS IN NON-POLAR SOLVENTS' BY SAMUEL KAUFMAN AND C. R. SINGLETERRY Naval Research Laboratmy, Washington, D. C. Received Februarv 86, 1061
+
Cryoscopic data indicate that tertiary aliphatic amines can react with carboxylic acids according to the equilibria N AS NAe and NA2 ++A* Na, (N = monomeric amine; A = monomeric acid). It, is probable that the dimeric acid molec u b reacts directly with such amines. This mode of combination is consistent with compositions of solid phases isolated from non-aqueous solvents as well as with the cryoscopic data presented. A possible bonding mechanism and structure are suggested. The amine-carboxylic acid complexes are appreciably dissociated into free acid and free amine in dilute benzene solutions. The weakly basic aryl tertiary amines such as dimethylaniline do not react with carboxylic acids in benzene to n detectable extent. The heterocyclic conjugated pyridine reacts with these acids to only a slight extent. There is no evidence in the data presented to indicate micelle formation by the high molecular weight complexes formed between tertiary amines and carboxylic acids.
Introduction Liltle information is available concerning the stability and the state of dispersion of the soluble compounds formed between amines and carboxylic acids in non-aqueous solvents. The amine picrates2-6 have been found to be stable 1-1 compounds, which in the case of primary and secondary amines are strongly associated to form dimers. With the exception of the tribenzylammonium salt the picrates do not dissociate appreciably into free acid and free amine. Certain primary amines have been reported to form 1-1 compounds with carboxylic acidsam7in non-aqueous solvents. Studies in the absence of a solvent have been considered to give evidence of the existence of compounds containing from one to three molecules of acid per molecule of amine.+" The present investigation was intended to establish the presence or absence of colloidal aggregates of amine-acid compounds in benzene solutions. Preliminary cryoscopic observations indicated that micelle formation did not occur to a detectable extent, and suggested further that the amine-acid compounds were dissociated into free amine and free acid to a degree that permitted detailed study of the reactions. It was expected that the amine and acid might combine in equivalent proportions to form salts that could be recrystallized for cryoscopic study, but all of the crystalline compounds isolated contained more than one equivalent of acid per mole of amine. The variety of possible compositions deduced from studies of binary systems and from (1) The opinions or assertions aontained in thi oomrnunication are those of the authors and are not to be construed as official or reflecting the views of the Navy Department. This is a oondensation of Naval Research Laboratory Report 3743 (October 9, 1950). (2) F. M . Batson and C. A. Kraus, J . Am. Chsm. Soc., 16, 2017 (1934). (3) R. M. Fuoss and C. A. Kraua, {bid., 67, 1 (1935). (4) C. A. Kraua and G . S. Hoopsr, Proc. Nall. Acad. Sci., 19, 939 (1933). (5) A. A. Maryott, J . Reeearch Natl. Bur. Standards. 41, 1, 7 (1948). (6) E. B. R. Prideaux and R. N. Coleman, J . Cfism. Soc., 462 (1937). (7) B. A. Hunter, Iowa State Coll. J . Sci.. 16,223 (1941). (8) W. W. Lucasse, R. P. Koob and .I. G. Miller, THISJOURNAL, 48, 85 (1944). (9) P. Matavuli, Bull. soc. ofiim. roy. Yougoelm, 10, 25, 35, 51 (1939). (10) E. A. O'Connor, J . Chem. Soc., 196, 1422 (1924): 119, 401 (1921). (11) W. 0. Pool, H. J. Harwood and A. W. Ralnton, J . Am. Chsm. Sor., 67, 775 (1945).
analysis of the solid amine-acid compounds rcported here do not prove the existence of the indicated compounds in dilute solutions in non-polar solvents. The combining ratios observed might be determined as a consequence of the relative solubilities of the possible compositions and the geometry of the crystal lattice. The equilibrium between acids and amines was therefore examined by adding amine and successive increments of acid to a quantity of benzene in which they would remain comi pletely dissolved. Cryoscopic determinations were made with solutions of cetyldimethylamine, triisoamylamine, pyridine, dimethylaniline, and acetic, myristic and picric acids, respectively. Further observations were made with additions of myristic acid to solutions of cetyldimethylamine, triisoamylamine, pyridine and dimethylaniline, and with additions of picric and acetic acids to solutions of triisoamylamine. Because of the pronounced association reported for primary and secondary amines and their saltslSthe scope of this investigation was limited to tertiary amines in order to simplify the interpretation of the data. The four amines studied were chosen to provide a wide range of basicity, structural type, and symmetry of substitution about the nitrogen atom. Materials.-The myristic acid (Eastman Kodak No. 1116) was prepared for use by drying in vacuo to 60' (m.p. 52.353.0 ). Ralston'a considered 54.4' the best value for this acid. The neutral e uivalent was 232.8 (theoretical 228.36). The CetyldimethyTamine was prepared by the method of Ralston, et U Z . , ' ~ from hexadecylamine having a melting point of 45.646.6' (literature value 46.77').14 The cetyldimethylamine was fractionally distilled i n vacuo, and stored in sealed ampoules. Before use, i t was dried by heating in vacuo to 80". Its neutral equivalent was 278.7 (theoretical 269.5). Spot tests16 indicated that primary and/or secondary amines were present to the extent of less than 0.1%.
The triisoamylamine (Eastman Kodak No. 1880) was dried over solid potassium hydroxide, and fractionally distilled in vacuo. Refractive index (ne%) 1.4330; neutral equivalent, 228.4 (theoretical 227.4). The dimethylaniline (Eastman Kodak No. 97)was dried and fractionally distilled in vucuo. Its refractive index (nBD)was 1.5589 (I.C.T. value, 72% 1.5587). The pyridine (Baker C.P. grade) was fractionally distilled (12) A. W. Ralston, "Fatty Acids and their Derivatives," John Wiley and Sons, Ine., New York, N. Y., 1948, pp. 32-33. (13) A. W. Ralston, D. N. Eggenberger, H. J. Harwood and P. L. DuBrow, J . Am. Chem. Soc., 69,2095 (1947). (14) A. W. Ralston, C. W. Hoerr, W. 0. Pool and H. J. Harwood, J . Org. Cfiem., 9, 102 (1944). (15) F. Feigl, "Qualitative Andy& by Spot Tests," 3rd edition, Elsevier Publishing Co., Inc., New York, N. Y., 1946, PD. 36881.
&AMINESAND ORGANICACIDSIN NON-POLAR SOLVENTS
May, 1952
with appropriate safeguards against the absorption of moisture. Its refractive index (n%) was 1.5101 (I.C.T. value, n% 1.509): The picric acid (Eastman Kodak No. 210) was recrystallized from aqueous alcohol and from benzene, and desolvated in vacuo a t 62" (m.p. 122.3-122.9'). The acetic acid (Baker C.P. A.C.S.) was purified in a dry atmosphere by four partial freezings. The fraction reserved for cryoscopic work froze at 10.4". Thiophene-free benzene (A.C.S.)was percolated through activated silica gel to remove polar material and moisture, and was !tored over sodium until used. Its freezing point was 5.37 . The naphthalene (Eastman Kodak No. 168) was purified by double resublimation in vacuo and was heated to about 50" in vacuo just before use. Special precautions were taken to protect chemicals and equipment from moisture. Experimental Procedure.-A Beckmann cryoscopic apparatus wm used for the freezing point measurements. The mean jacket temptrature was held at 3.3" - AT The contents of the inner tube within a tolerance of 0.1 were agitated by an externally driven glass stirring screw. Reciprocal vertical stirring was found to pump out solvent vapor a t an appreciable rate. The losses with rotary stirring were about 0.05g. per hour and corrections to observed data were made on this basis. The stirrer shaft entered the inner freezing tube through a glass T-tube. Dried nitrogen was forced past the shaft; omission of the nitrogen current allowed the entry of moisture, pFogressively depressing the freezing point. Before use, the inner tube was flushed with dried nitrogen for one hour or more to remove moisture. Omission of the flushing produced a variable and marked depression of the freezing point of the solvent. The amount of supercooling was controlled a t 0.2 f 0.05" by seeding with minute dry crystals of benzene. Correctionsle were made for withdrawal of solvent by freezing. Within 1% equal quantities of benzene were used for all determinations. Observations were made in a constanttemperature room at 25" and 20 to 50% relative humidity. The temperature .of the solution was read at half-minute iiitervals preceding freezing and after freezing began, until the temperature showed no detectable change during a minimum interval of four minutes. This constant temperature was taken as the observed freezing point. When all of the data for a given solute were combined in a single plot of AT, the freezing point depression us. concentration, the average deviation of the experimental points from thc best smooth curve was f0.002".
.
GO5
0.7
06
0.:
0.4 0
c
4
0.3
6.2
0 CETYLDIMETHYLAMINE TRIISOIMYLAMINE A DIMETHYLANILINE
0.1
A PYRIDINE 0 NAPHTHALENE
0 '0
0.05
0.10
0.15
0.20
CONCENTRATION, EOUIVALENTS/ KG BENZENE,
Fig. 1.-Depressions of the freezing point of benzeiic by amines and by naphthalene.
molal freezing point constant, Kf, of benzene fouiicl is 5.17'. The equation employed by the Bureau of Standards for the cryoscopic determination of the Although attempts to prepare solid compounds purity of benzene" is consistent with a value of were not pursued far enough to produce materials of Ir'f = 5.13 a t infinite dilution. The constant obhigh purity, the results are reported briefly because tained in the present measurements is thought to of their relation to those obtained cryoscopically. differ from the accepted value because of small sys(a) Cetyldimethylamine with molten stearic acid tematic errors arising from the materials and techgave a solid which was recrystallized three times niques employed; it was used for the treatment of from acetone, m.p. 64.2-64.5'; N, 1.31 (calcd. for other data obtained under the same conditions ill triacid complex 1.24%). (b) Cetyldimethylaminc order to minimize the effect of such errors. titrated with stearic acid in 95% ethanol to phenol The cryoscopic effect of the amines (Fig. 1) is red end-point gave a solid which was recrystallized very near that of the naphthalene, except for the from ethanol and from acetone, m.p. 77-79', N, case of pyridine, which is known to be abnormal in 1.62 (calcd. for diacid complex, 1.67%). (e) its cryoscopic behavior,l8 probably because of,the Cetyldimethylamine with equivalent proportions formation of solid solutions. Picric acid (Fig. 2) of lauric acid in ethyl acetate solutioii gave a solid shows nearly ideal behavior. The dotted line, which was recrystallized from acetone and from drawn for comparison purposes, represents oneethyl acetate, m.p. 56.8-57.5'; N, 2.13 (calcd. for half the freezing point depressions expected for diacid complex 2.09y0). (d) Cetyldimethylamine and corresponds to the slope for a with an equivalent amount of oleic acid gave a mush naphthalene, totally dimerized acid. Inspection of the graph of crystals which were centrifugally filtered and suggests that appreciable quantities of the carboxywashed with benzene, m.p. 36.2-37.3'; N, 1.80 lic acids are in the monomeric form, but that the (calcd. for diacid complex, 1.68). Figure 1 presents the cryoscopic data for naph- dimeric form preponderates.lS There is little dif(17) A. 1-1. G!usguw, A . ,I. Streiff rtiicl E'. D. Roaaiiii, J . Research Naif. t,li:deiie aiitl the anlines. The il.vei:~ge esperimeii(d Bur. Sta!rdardr, 56, 355 (1945).
Results
(1G) IC Arridt, "llsndbucli der Pliy~ilil~ii~c1i-(:heiiii-iclieii'l'aaliiiilL." 2nd edition, Fcrdinnnrl ICnkc, Stuttanrt, 1023, p. 110.
(18)C . R. Burg and 11. 0.Jeiikiiia, J . Chem. Sue., (is8 (1934). N. Brocklenby, Can. J . /le.roorch, 14B,222 ( I X c j ) .
(19) 11.
SAMUEL KAUFMAN AND C.
606
R. SINGLETERRY
Vol. 56
The brackets signify concentrations expressed’ as moles per kilogram of solvent. The values of K A show ~ a scatter which is not greater than is to be expected from the experimental uncertainty of the freezing point measurement. If the concentrations are expressed as mole fractions instead of molalities, J C A ~ for myristic acid is 6.0 X and for acetic acid is 1.25 X loa4, where k~~ is the dissociation constant on the mole fraction basis. The corresponding JCA~ derived from dielectric measurements20 at 30’ for stearic acid is 1.7 X low4,and that for acetic acid is 2.4 X lod4. If it is assumed that the heats of dimerization in benzene solution of both stearic and acetic acids are equal to 8.6 kcal. per mole of dimer, the value reported for benzoic acidlZ1then the constants at 5.4’ for these acids may be computed from the Van’t Hoff equation. At this temperature, approximately that of the cryoscopic observations, the computations lead to = 4.8 X 10-6 for stearic acid and JCA, = 0.68 X for acetic acid. These estimated constants are in reasonable agreement with those obtained from the cryoscopic measurements.
0.10
0.N
OX
0.4[
u I-
a 0.3C
TABLEI1 CRYOSCOPIC EFFECTSRESULTINQ FROM THE ADDITION OF ACIDSTO AMINESIN BENZENE
0.21
Increment symbol
0 ACETIC ACID MYRISTIC ACID A PICRIC ACID 0 NAPHTHALENE
0.10
E 0
1
1
I
0.05
0.10
0.15
0.20
CONCENTRATION. FQUIVALENTS / KG. BENZENE.
Fig. 2.-Depressions of the freezing oint of benzene by acids and by naphthaine.
ference between the behaviors of myristic and acetic acids in this respect. Concentrations recorded in the tables and figures were computed from measured neutral equivalent weights for cetyldimethylamine, triisoamylamine and myristic acid, and from formula weights for the remaining compounds.
Solute added” e uiv./ ATtotal ( %g./ (“C.) solvent)
+
Cetyldimethylamine myristic acid Amine 1. 0.205 0.0393 Acid l a .217 .0337 .243 .0723 Acid 2 a .293 .1127 Acid 3a Acid $. .363 .1514
+
Triisoamylamine myristic acid Amine l b 0.182 0.0350 Acid Ib .I99 ,0338 .227 .0665 Acid 2 b Acid 3b .281 .lo62 Acid 4b .347 .1395
+
0.0404 ,0826 .1240 .1649
(KAz)
9.2 12.0 7.2 4.7 Av. 8
(equiv./kg. &vent)
0.0153 .0363 .0703 .lo37
(KAz)
1.8 2.0 1.4 1.2 Av. 1 . 6
(Table I) of the The dissociation constants, dimeric acids were calculated from Kt and the cryoscopic data. The following reaction was assumed Az
2A
where A represents the monomeric acid. Thus KA, = [Ala/[A21
(1)
symbol ment
Solute added“ ATtotd (“C.) ’% esolvent) .F(
+
Pyridine myristic acid Amineld 0.180 0.0393 Acid l d .270 .0437 Acid 2d .370 .084’3 Acid 3d .475 .1243 Acid 48 .572 .1629
+
Triisoamylamine acetic acid Amine le 0.200 0.0380 .210 .0372 Acid le Acid 2, .229 .0759 Acid 3. .337 .1546 Acid 4e .488 .2286 Acid 5, .658 .3026
+
Triisoamylamine picric acid Amine 1r 0.200 0.0384 Acid l r ,198 .0170 Acid 2r ,178 ,0334 Acid 3r ,192 ,0434 Acid 4r .296 ,0655 Acid 5r .438 ,0959 a Values for “Amine” increments are those for solutions of amine before additions of acid. Values for “Acid” increments represent total acid added.
Dimethylaniline myristic acid TABLE I Aminelo 0.212 0.0403 Acid lo ,330 .0425 DISSOCIATION OF ACIDDIMERIN BENZENE FROM CRYOSCOPIC .440 :0866 Acid 2, OBSERVATIONS Myristic acid Acetic acid .558 .1288 Acid 3, Dissn. Dissn. .671 ,1717 Acid 4, constant constant Concn., x 104 Concn x 103 (equiv./kg. solvent)
Incre-
The primary data of Table I1 were obtained by the addition of successive increments of various acids to benzene solutions containing fixed amounts of different amines. Plots of these data are not directly comparable since the initial concentrations of the pure amine solutions were not precisely equal. (20) A. A. Maryott,.M E. Hobbs and P. M. Gross, J. Am. Chem. Soc., 71, 1671 (1949).
(21) L. Pauling, “The Nature of the Chemical Bond,” 2nd edition, Cornel1 University Press, Ithaca, N. Y., 1940, Chap. IX.
1
&AMINESAND ORGANICACIDSIN NON-POLAR SOLVENTS
May, 1952
607
For graphical purposes, the primary data have and [N,] is the concentration of amine in the solubeen reduced to direct comparability with respect tion to which the acid was added. to equivalence points, relative amounts of acid Discussion of Results present, and slopes. The derived results are The interpretation offered for these experimental plotted (Figs. 3 and 4) with the relative cryoscopic effect, ATtotal/ATamineas the ordinate, and the results is based on the assumption that, in these diacid-amine ratio, equivalents acid/equivalents lute solutions, the individual kinetic units present amine, as the abscissa. The depression ATtotal is have substantially ideal effects on the activity of that of the observed solution, while ATamine 1 is benzene. The validity of the assumption is supthat of the amine solution before acid additions. ported by the cryoscopic effects found for three of It is true that the relative cryoscopic effect for a the four amines (Fig. l), and by the satisfactory given acid-amine ratio will vary somewhat with the values calculated for the dimerization constants of level of absolute concentration. However, in. the myristic and acetic acids. The cryoscopic data of ~ ~distinctly ~ ~ ~ ~ polar comrange studied, if the initial amine concentration Kraus and ~ o - w o r k e r sfor differs by lo%, the maximum change in the relative pounds such as triisoamylammonium picrate4 have been successfully explained on the basis of a similar cryoscopic effect is of the order of only 1%. The dotted lines represent the hypothetical cases assumption. Accepting the validity of Raoult’s law in the syswherein the reaction of an amine with an acid attains irreversible completion a t 0, 1, 2 or 3 equiva- tems studied, it is clear that the compounds resultlents of acid per equivalent of amine, following ing from the reaction between amines and carboxwhich additions of excess carboxylic acid exert an ylic acids in these solutions are in equilibrium with independent cryoscopic effect. The data for my- substantial amounts of free acid, even in the presristic acid were used for plotting the diagonal ence of excess amine. Dimethylaniline gives no dotted lines in both figures. The difference be- indication of reaction with myristic acid; the points tween the cryoscopic behaviors of myristic and ace- for the freezing point depression after the addition tic acids is so slight as to be imperceptible in these of acid lie quite precisely on the dotted line 0 (Fig. 3) which corresponds to the independent cryoscopic graphs. Because pyridine does not conform to the cryo- effect of the acid. This behavior is in accord with scopic behavior of the other amines studied, the that observed by O’ConnorlO for the same amine pyridine-myristic acid data in Fig. 3 were treated with acetic acid. There is evidently some interacdifferently from .the rest to permit convenient vis- tion between pyridine and myristic acid, but the compound formed must be highly unstable. Even ual comparison. The ordinate in this case is in the case of the two aliphatic amines the mini1 ( A T a o i d / [ N ~ ] X Kt) mum amount of uncombined acid that must be where AT,,id is the total additional depression assumed to explain the data is considerable. found after additions of acid to the amine solution, The data for triisoamylamine plus picric acid (Fig. 4) indicate the formation of a stable amine picrate whose composition is one equivalent of acid to one of amine. No significant dissociation can
+
3’0r-
t
t 0
t
CETYLDIMETHYLAMINE TRIISOAMYLAMINE A OlMETHYLANlLlNE 0 0
A
0 ACETIC ACID
A
MYRISTIC ACID PICRIC ACID
0
PYRIDINE
0
1.0
2.0
3.0
4.0
E O U I V A L E N T S ACID / E Q U I V A L E N T 1.0
2.0
3.0
4.0
E O U I V A L E N T S ACID /EPUIVALENT
5.0 AMINE.
Fig. 3.-Relative cryoscopic effects resulting from additions of myristic acid to solutions of various amines.
5.0
6.0
’
AMINE.
Fig. 4.-Relative cryoscopic effects produced by additions of various acids to solutions of triisoamylamine. (22) D.A. Rothrook, Jr., and C. A. Kraus, J . Am. Chem. Soc., 69, 1699 (1937).
,
SAMUEL KAUFMAN AND C. R. SINGLETERRY
608
be inferred from the data in this case. The abnormally small depressions produced by the compositions approximating stoichiometric proportions of the amine and acid indicate an appreciable interaction between the amine picrate molecules (N
2NA NzA2 monomeric amine)
=
(2)
Vol. 56
fact that the cryoscopic effects are substantially less than can be accounted for by equilibria .(1) and (4b) indicates some further interaction between the solutes present. By analogy to the behavior of triisoamylammonium picrate, the equilibrium 2NA2
NzA4
(6)
which is not reflected in the dielectric measurements was postulated, which should conform to the exof Maryott,b but which is supported by the data of pression Batson and Kraus.2 The intersection of the exI ~ N ~=A [,N A I /[NAZI trapolated branches of a plot of AT vs. total picric As an ahernative, the equilibrium acid added to triisoamylamine falls within 1% of the abscissa corresponding to a 1:1 ratio of acid to NA,f-NAz+Az (7) amine. The curve beyond the stoichiometric point was p,roposed, for which the expression is very nearly that to be expected from the indeK N A r = [A21 [NAZI /["Ad pendent action of the excess picric acid. It has a slope corresponding to K I = 4.7 as compared with would hold. For the case in which an interaction represented Kr = 5.1 for a benzene solution of picric acid alone. The difference exceeds the experimental error, and by equilibrium (6) or (7) leads to a complex having may indicate a slight tendency of the excess picric a stability a t least an order of magnitude less than acid to associate with the amine picrate. Strong that of NA2, it is possible to determine the approxi~ reference t o these inacids such as hydrogen chloride may be expected mate value of K N Awithout to react as completely with aliphatic tertiary amines teractions. This is done by computing apparent values of KNA*from the experimental data with as does picric acid.23~~~ A satisfactory explanation of the data for the re- the assumption that only equilibria (1) and (4b) action of triisoamylamine or cetyldimethylamine apply, and extrapolating the resulting plot of KNA, with myristic acid requires that one equivalent of a vs. acid-amine ratio to zero acid-amine ratio. tertiary aliphatic amine combine with two equiva- The extrapolation is improved in linearity but not lents of myristic acid monomer (or directly with altered in final value if, for a second approximaone unit of the acid dimer) to form a molecular com- tion, a reasonable value of K N ~orAK~N A ,is introplex that is considerably dissociated in dilute solu- duced into the computations for K N A ~ . By use of the value K N A=~ 3 X obtained tions. An attempt was made to compute equilibrium constants for some possible reactions between by assuming equilibria (1) and (4b) only, an atA the ~ data cetyldimethylamine and myristic acid by a method tempt was made to fit a value of K N ~to that took account of the relative concentrations of for the system triisoamylamine plus myristic acid. was ~ consistent with a the monomer and dimer forms of the acid required No assignable value of K N ? A by the constants of Table I. No one of the equilib- constant value of K N P ~ .On the other hand the similarly computed values of KNA'corresponding rium expressions corresponding to the reactions to increasing total acid content was respectively N A ~ N + A (3) 0.71, 0.110, 0.0714 and 0.0855. The first of these N A z q N +2A (4a) was discarded as unreliable because of the large effect of experimental errors a t this acid-amine NAz - N + A z (4b ) ratio, and K N A=~0.0855 was taken as the probable NApf;-N +3A (5) value. was satisfactorily constant, although i n the range To test the hypothesis that equilibria (l), (4b) between 0 and 1.5 formula weights of acid per for- and (7) are the controlling equilibria in this system, mula weight amine, the expression ATtotal was computed for each of the experimental compositions, using the determined values of KNA? [Nl [AzI/[NAzl , the calculated values of corresponding to equilibrium (4b) varied less rap- ATsmine I, Kr and K A ~and KNA? and K N A ~ For . comparison a similar test was idly with changing acid to amine ratio than did those for equilibria (3) and (5). Equilibrium (4a) made, assuming equilibria (l), (4b) and (6). A A ~ was chosen, which would could .be used as an alternative for (4b) because value of K ~ J =~ 0.071 produce a fit most favorable to the two uppermost [A,] and [A] are mutually dependent upon K A and ~ points in Fig. 5 , since these two' points are most equilibrium (1). ~ . result of If the only equilibria to be considered were (1) sensitive to the variation of K N ~ AThe and (4b), A T t o t a l should approach the dotted line 2 the computations demonstrates that equilibrium (Fig. 3) asymptotically as the excess of the acid be- (7) is consistent, while equilibrium (6) is inconsistcomes sufficient to suppress the dissociation of the ent, with the experimental data for the triisoamylcomplex. The curves for both cetyldimethylamine amine-myristic acid system. Examination of Fig. 4 shows that equilibrium (6) and triisoamylamine with myristic acid, however, is not valid for the triisoamylamine-acetic acid cross this line. The same is true for the system triisoamylamine plus acetic acid (Fig. 4). The system. The validity of this equilibrium requires that the C I I I ' V ~approach the dotted line 3 awppto( 2 3 ) >r. R I . I h v i s . .I. h t t . Chern. SVC.,71, 3544 ( l H 4 H ) ; (Ir) h1. h l . tically, wliereus it very deliiiitely c,rosses ittiis liric. Davis aiid IG. A. A I c U w i d d , J . IZeseaxh NULL Bur. Slondurdy, 42, 696 Computations similar to those described above, (I!NU). assuming equilibria (I), (4b) and (7) were made for (2.1) J. A . hlocrle niid C . Curtail, J . A m . Ckeni. Sor., 71, 852 (1949).
.
I
--f
.
May, 1952
l - h X N E S A N D ORGANIC
ACIDSIN NON-I~OLAI~ SOLVENTS
the system triisoamylamine-acetic acid. Results for this system were K N A=~ 1.02 X and KNA, = 2.38 x low2; these values were insensitive to further approximation cycles. A comparison (Table 111) of the computed and measured freezing
60‘3
0’
TABLE I11 TE~T OF EQUILIBRIA IN SOLUTIONS OF TRIISOAMYLAMINE PLUSACETICACID hlnine content = 0.0380 equiv./kg. benzene; KA, = 1.6 X ~ 1.02 ~ x 10-3; Z C N A ~ = 2.38 x 10-2 10-3; K N = Total acid added,
(equiv./kg.benzene) 0.0372 .0759 .1546 .2286 .3026
exptl.
ATtotal, ‘C., coinpiited NA2 Az a NA4
0.210 ,229 ,337
0.210 ,229 ,335
ATcataIr ‘C.9
+
,488
.489
,658
,665 0
EXPERIMENTAL
I
point depressioiis shows that equally good agreement is obtained, except for the mixture of highest acid-amine ratio. It is concluded that equilibria (4b) and (7) are valid for dilute benzene solutions of tertiary aliphatic amines plus fatty acids. 0 15 0 0 05 0 IO 0 15 0 29 The instability of the compounds formed beACID A D D E D . E P U I V A L E N T S / K G B E N Z E N E . tween amines and carboxylic acids in benzene solation, and also the surprising amount of acid bound Fig. 5.-Test of alternative equilibria in solutions of Criisoamylsmine plus myristic acid. per equivalent of amine, indicate that the forces between the acid and amine are different in type from those operating in the amine picrates or the amine bered ring formed by the hydrogen bridges of the hydrohalides. The reaction appears to be the for- two carboxyl groups with such an orientation that mation of a molecular complex between acid dimer the unshared electron pair of the nitrogen was :uid the amiiic rather than the formation of a true equally available for dipolar interaction with either of the hydrogens of the dimeric acid, as represeritcd ion pair. The successful description of the experimental in I. 0-H -----0 rcsults i n terms of one or of a series of equilibrium / \\ coiistaiits does not, of course, furnish direct inforIt- -C mation concerning the type of bonds involved, nor \O-----I~-O Nn: as to whether the monomeric or dimeric form of the acid is the chemical unit actually reacting with the I amine. It might have been postulated that the observed phenomena were the result of the stepwise Two six-membered chelate rings would result from reaction of the amine with four units of acid mono- the formation of I, and might lead to moderate mer. Thus the choice of the acid dimer as the stability of a type of hydrogen bridging that would reactingunit may appear arbitrary and justified only otherwise be too weak to detect. The bonding by a single hydrogen of more thaii 1)y the resulting coiivenience i l l the analysis of the (lata. However, the assumpt,ionsemployed for calcii- t,wo electroiiegative atoms at8one t,ime i s iiii~isiial, laidng the coiicentrations of the various species in- R I ( Iiough Pading?’ cites lhe “bifurcaled Iiydrogen \wived i n Lhe equilibria imposc :t further restrictioti. b011d” N-K; *o iiilerred froin lhe cryslid slruc? K N Awere ~ computed from The cotistaiits K N Aand a ooticeiitratioiis of N, NA2 and NA, that were dr- t u l e of gly~ine.~5Tlie simultaiieous direction of rived from the freezing point data on the assump- two hydrogen bridges toward ,z single oxygen atom tion that neither NA nor NABwere present in sig- has been postulated for the dimeric form of s:tlicylic nificant amounts. The fact that the values derived acid?’ as well as in 1,8-dihydroxyanthraquinone for these constants by ignoring the possible pres- and some similar compounds. Oxygen differs from ence of NA or NA, lead to computed depressions of nitrogen, it is true, in having two pairs of electrons the freezing point in close agreement with experi- not involved in covnleiit bonding, but hydrogen ment supports the idea that the acid dimer is the bonding is usually considered to be essentially an actual reacting unit. electrostatic or dipole effect and not electron sharing. It is difficult to construct a model for the associa- The bonding of a second dimer by the NA2 is diftion of a single amine molecule with four separate ficult to explain. The first dimer unit may, howacid monomers that implies an energy of bonding ever, be sufficiently polarized by the presence of the sufficient to compete effectively for monomeric nitrogen atom so that the loose association of a secunits with the dimerization equilibrium (1). On ond unit becomes possible. the other hand, a direct bonding of the dimeric acid The strength of bonding between the amine and iinit to ,z tertiary amine might result if the nitrogen ( 2 5 ) G. AILrccht and II. B. Ccrcy, J. .1m. Chcrn. Soc., 61, 1087 of tlic amine approached one side of the ciglit-mcm- (1930)
7-lL
#
610
HITOSIHAGIHARA
acid is not great; for triisoamylamine with myristic acid the degree of dissociation of NA2 estimated for a concentration of 0.01 formula weight per kilogram of benzene is 46% as compared, for example, with 13% dissociation of the hydrogen-bonded dimer of myristic acid to monomer a t the same concentration. (A smaller degree of dissociation in paraffinic than in aromatic solvents is mggested by the work of Pohl, Hobbs and Grossz6who found that (26) H.A. Pohl. M. E. Hobbs and P. M. Gross, J . Chem. Phys., 9, 408 (1941).
Vol. 56
dimeric acetic acid was more readily dissociated in an aromatic than in a paraffinic solvent.) A bond of the type indicated in I implies a smaller heat of reaction and a smaller dipole moment than would be expected if a tertiary amine reacted with a fatty acid to produce an ionic compound. Acknowledgment.-Acknowledgment is made to Kenneth L. Temple, formerly of the Naval Research Laboratory, for the preparation of two compounds of cetyldimethylamine with stearic (b) and lauric (c) acids.
SURFACE OXIDATION O F GALENA I N RELATION TO I T S FLOTATION AS REVEALED BY ELECTRON DIFFRACTION BY HITOSIHAGIHARA~ Taihei Mining and Metallurgical Laboratory,2 Omiya City, Japan Received February $5, 1861
The initial oxidation of galena surfaces was studied in air, in an enclosed atmosphere in alena owder, in a vacuum furnace, in water, and during dry and wet grinding. The electron diffraction examination of t8e oxilieed faces revealed that the lowest oxidation product is in all cases lead sulfate (PbSOd). Under dry conditions the sulfate crystallites were oriented in .two ways on the cleavage face and no pseudomorphism was observed. In air the next higher oxidation product was the basic sulfate, PbzS06. Considerations are given to the orientation relations of these oxidation products and their growth. In no instance were crystalline carbonate, hydroxide or the lower sulfoxides, PbS,O, (with n / m less than 4) observed.
Introduction Electron diffraction investigation of the initial surface oxidation on galena is of interest from the theoretical as well as the practical viewpoint. The former interest is concerned with the mechanism of crystal growth of oxides on solid surfaces. With single-crystal substrates, particularly single crystal cleavage faces where the arrangement of atoms a t the boundary surface is known, the above problem will be given the most detailed considerations. Interesting electron diffraction studies have already been reported with such sulfide minerals as zincblende, ZnS, 3-6 molybdenite, M o S ~ and ,~ stibnite, Sb&8 Galenag is well suited for this purpose: its crystal structure is simple (NaC1type cubic),lo its cleavage face is simple (cube face), and yet the detailed study of its initial oxidation is still lacking. The practical viewpoint concerns itself with the separation of sulfide minerals by flotation. It is generally accepted that sulfide minerals undergo some degree of surface oxidation in the processes of crushing and grinding before entering into the (1) Kobayasi Institute of Physical Research, Kokubunsi, Tokvo, Japan. (2) Former Mitsubishi Mining and Metallurgical Laboratory. (3) T. Yamaguti, Proc. P/iys.-Malh. SOC.Japan, 17, 443 (1936). (4) G. Aminoff and B. Broom& KQZ. Svenska Velenskapsakad. Handl., 16, 3 (1938). (5) R. Uyeda, S. Takagi and H. Hagihara, PTOC. Phys.-Math. Soc. Japan, 23, 1049 (1941). (6) D. M. Evans and H. Wilman, Proc. Phys. SOC.(London), 8 6 8 , 298 (1950). ( 7 ) R. Uyeda, PTOC.Phys.-Math. SOC.Japan, 20, 656 (1938). (8) 8. Miyake, Sci. Papers Inst. Phys. Chem. Reaearch ( T o k y o ) , 84, 565 (1938). (9) H. Hagiliarrt, Proc. P/ws.-Malh. Soc. Japan, 2 4 , 762 (1942). (10) R.W.G. Wyckoff, “The Structure of Crystals,” The Chemical Catalog Co., Inc., (Reinhold Puhl. Corp.), New York, N. Y., 1931. TI& monograph contaiw tlic structural data and references thereupon.
flotation cell. The reaction of the aqueous solutions of the collecting reagent with such oxidized surfaces has been of importance in the “chemical theory” of flotation advocated by Taggart, and accordingly a great deal of research has been devoted t o its understanding.l1-l6 It is now widely accepted that the surface of galena exposed to air either through dry or wet grinding is a complex mixture of sulfoxides, hydroxides and carbonates. l7 This conclusion is based upon the chemical analyses of ions liberated into distilled water or into the aqueous solutions of the collecting reagent from galena. It is obvious that such deductions cannot be free from unavoidable ambiguities in interpreting which compounds actually exist on the mineral surfaces. The rate of oxidation is also controversial since reported studies range from “rapid”11- 18,19 to ‘Lslow.”20*z1 It is, therefore, desirable that these conclusions be checked by an independent research technique such as is afforded by electron diffraction examination. The present paper deals mainly with the initial (11) A. F. Taggart, T. C. Taylor and A. F. Knoll, Trans. A m . Inst. Mining Met. Engrs., 87, 217 (1930). (12) A. F. Taggart, T. C. Taylor and C. R. Ince, ibid., 87, 282 (1930). (13) A. F. Taggart, G. R. M. del Guidice and 0. A. Ziehl, ibid., iia, 348 (1935). (14) A. Knoll and D. L. Baker, Am. Inst. Mining Met. Engrs., Tech. Pub. No. 1313 (1941). (15) A. F. Taggart and M. D. Hassialis, Trans. Am. Inst. Mini7~u Met. EngTs., 169, 259 (1946). (16) T. C. Taylor a n d A. F. Knoll, ibid., 112, 382 (1935). (17) I. W. Wark, “Principles of Flotation,” Australian Institute of Mining and Metallurgy, Inc., Melbourne, 1938, p. 184. (18) P. A. Lintern a n d N. K. Adam, Trans. Faraday Soc., 81, 564 (1936). (19) H.H.Herd and W. Ure, T H I s JOURNAL, 48, 93 (1941). (20) H.B. Bull, B. S. Ellefson and N. W. Taylor. THISJOIJRXAL, S8, 401 (1934). (21) P. Siodler, h’olloid-Z., 68, 89 (1934).
