T H E ELECTRICAL CONDUCTIVITY O F ORGANIC ACIDS I N WATER, ALCOHOLS, AXD ACETOKE, AND T H E E L E C T R O N C STRUCTURES OF T H E ACIDS* BY HERSCHEL HUNT WlTH H. T. BRISCOE
I n a previous paper1 the authors presented data dealing with the conductances of twenty-four organic acids in water and ethyl alcohol. The work reported in this paper represents an extension of that study to solutions of some of these acids in methyl, propyl, and butyl alcohol, and acetone. There is a pronounced lack of data dealing with the effect of properties and constitution of such solvents upon the ionization and conductance of solutions of weak electrolytes. This paper also includes the work done to date upon the study of ethyl alcohol solutions of certain dibasic acids, a few acids whose molecules contain double bonds, and several aromatic acids not included in the previous report. The dissociation constants of various benzoic acid derivatives have been calculated from the values of the limiting conductivities of HCl, KaC1, and the sodium salts of the acids at 3ooC.
Experimental The alcohols were treated with 20 cc. of sulphuric acid and 20 cc. of water per liter to free them from amines and were then distilled. They were freed from aldehydes by distilling from an alkaline silver nitrate solution. Highcalcium hydrated lime was dehydrated in an electric furnace at 600-700°C. Each alcohol was dehydrated by refluxing one liter of alcohol with 500 gms. of this lime for eight hours. After distillation, the product was again refluxed with zoo gms. of lime, and the alcohol was distilled. After fractional distillation the methyl alcohol was found to have a specific conductance of 3.5-2.69 X IO-^ reciprocal ohms. This conductance was lowered t o 2.45 X IO-’ reciprocal ohms by distilling the alcohol from metallic calcium. Conductance measurements were made at 3oOC. Ethyl alcohol prepared in the above manner had a specific conductance of 7.5 X IO-^ reciprocal ohms; normal propyl alcohol gave a specific conductance of 9.1; X IO-^ reciprocal ohms; and normal butyl alcohol gave a specific conductance of 9 . 1 2 X I O + reciprocal ohms. The specific gravity of the ethyl alcohol at 2jDC. varied from 0.;8508-0.;8510. Pure alcohol2 has a specific gravity of 0.78jo; at this temperature. The water content of our alcohol is therefore less than 0.005 S. * This paper is a part of a thesis submitted by the first-named author as partial fulfillment of the requirements for the degree of Doctor of Philosophy in Indiana University. Hunt and Briscoe: J. Phys. Chem., 33, 190 (1929). “International Critical Tables,” 3, I 1 7 .
HERSCHEL HUST WITH H. T. BRISCOE
I496
Acetone was prepared by long standing over calcium chloride. It was then fractionally distilled with a thermometer and later with a long fractionating column. The last distillation was performed with the use of a block tin condenser. Acetone, thus purified, had a specific conductance of 12.8 X IO-^ reciprocal ohms at 3oOC. The bridge assembly was the same as that used in the previous work, except for the alterations described below. The microphone hummer was replaced by a Type E Vreeland Oscillator. This instrument, as well known, produces a pure sine wave of constant frequency independent of fluctuations in the actuating direct current. It is also noiseless in operation and may be conveniently started and stopped. Although the frequency may be altered, all measurements made in this work were made with the oscillator set to give a wave of 1000 cycles per second. Variable air condensers ( 5 X 1oC4 to 1 2 X IO-' microfarads) were placed in series with the resistance arm of the bridge to overcome electrode effects and to aid in producing a definite and easily determined minimum in the telephones. Curtis coils (IO-110,ooo ohms) xere used for resistances above 1000ohms in order to aid in the elimination of errors due to inductance and capacity. These coils had an accuracy of 0.04% and wcre checked against a standard resistance. Resistances below 1000 ohms were measured by means of a small Leeds and Korthrup calibrated, four-dial box. The cells used in the previous work were employed for the solutions of greatest conductivity. For solvents and very feebly conducting solutions, a cell similar to that described by Danner and Hildebrand' was used. This cell had a constant of 0.0036685. All cell constants were frequently checked. The bath, bridge, and method of cleaning apparat'us, calibration, filling the cells, and preparing solutions have been described in the previous paper.
Data In the following tables the molecular conductances of the acids are given for the various solvents. The temperature was 3oOC. in each case. Blanks in the tables are due to the fact that some of the most dilute solutions were such feeble conduct,ors that a significant reading could not be made or because of the insolubility of the acid in the solvent a t these particular dilutions. TABLE I Nonochloroacetic Acid MeOH
EtOH
PrOH
BuOH
0.0592
0.00288
0.1111
o.ooj2o
32
0,1451 0.5358 2 ,0314
0.2205
0.01122
0.00157 o ,00267 0.00j30
128
8.012 j
0.4809
0.03535 0 . I03 j j
Dilution 2
8
512
1024 1
3j.6723 62.9393
I
,1951
I
,6363
Danner and Hildebrand: J. .4m. Chem. Sac , 44,
o.orzj1
Acetone 0.0091 O.OIj2
o ,0369 0.0776 0.2206
0.2509 2824 ( 1 9 2 2 ) .
ELECTRICAL CONDUCTIVITY O F ORGANIC ACIDS
I497
TABLE I1 Dichloroacetic Acid bIeOH
EtOH
PrOH
BuOH
Acetone
0,3375 0.6697 I ,4030
0.2531 0.4576
0.0182 0.0258
0.0096 0.0131
0.0232 0.0388
0.0455
0.0218
0.0963 0.2037 0.3118
0.0523
0,0777 0.1642
1024
3,3549 9.9886 18.2681
0.89j8 I . 7619 4.3369 7.6226
Dilution
JIeOH
EtOH
2
1.7340
2
,0474
8
3.3155 6.8765
3.6805
0.IOjI
6’5904
0,1865
11 ,7716
0,3562
512
14,0599 30.4619
20.6543
0.6984
1024
40.7705
27.38j8
0.9627
Dilution 2
8 32 128 512
0.1389 0.2474
0.4567 0.68j4
PrOH
BuOH
Acetone
0.0649
0,0249 0.0383 0.0673 0.1290 0.2639 0.3207
TABLE I11
32 128
0.0372 0.0575
0.0982 0.2797 0.7110
0.8368
TABLE Is’ Cyanacetic Acid SleOH
EtOH
PrOH
BuOH
Acetone
0.0295 0.036j
0.3356 0.3808
15.0968
0.0479 0.0639 0.1129 0.1323
0.4460
1024
o 20j3 0.3292 0.580; I ,0989 2.1651 3.0183
0.0549 0.0676 0.0807
512
0.3790 0,5915 1,0778 2.3813 7.6400
Dilution 2
8 32 128
0.1072
0,1855 0.2500
0,5347 0,6333 0.6863
TABLE s’ Glycollic Acid Dilution 2
8 32 I28 jr2 1024
MeOH 0.0766 0.2094 0.6431
EtOH
PrOH
BuOH
Acetone
0.0416 0.0716
o ,0062
0.1210
0.0116
2.5744 10.7402 20.1687
0.2523 0.606j 0.9072
0.0225
0.0035 0.0043 0.0068 0.0133
0.0349 0.0457 0,0543 o ,0856
0.0077
0.1473
0,1882
HERSCHEL HUNT WITH H. T. BRISCOE
I498
TABLE VI Monochloroacetic Acid in Mixtures of Alcohols ( j o c c each by volume) Dilution
MeOH-BuOH
MeOH-PrOH
o
0224
0
0
0409
0
4040
o 0832
I 5819 6 1683 11 9962
o 1726 0 3965 o 6176
0
0364
0
8
0 1171
4077
32
128 512 1024
0
1 5896 6 4601 1 2 1364
EtOH-PrOH
0336 I097
2
EtOH-BuOH o 1092 0
0369
o 0812 o 2240 o 7162 I
3861
TABLE VI1 Xlolecular Conductivity of Acids in Acetone Dilution
Acetic Acid
2
0 0020
8
0
32 128
0
0039 0084
0
0247
o 0926
0
512 1024
0
Oj8I
0 2 j Z j
o 6886
o 4721
1
Brornoacetie k i d o ooj9 0 0108 0 0340
Iodoacetic Acid 0
0804
0
1346
o 1848 3538 I392
Discussion of Results The order of conductivity values for the various solvents is water, methyl, ethyl, propyl, and butyl alcohol for each acid studied. The solution of an acid in acetone has a conductance between that of ethyl and propyl alcohol, except in the case of trichloroacetic acid, in which case the values fall between those for propyl and butyl a t most dilutions. The dielectric constants' of the various solvents at 2 0 O C . are: Methyl alcohol Ethyl Alcohol Propyl alcohol
31.2
25.8
Butyl alcohol Acetone
19.2 20.7
22.2
The dielectric constant of acetone is nearer the value for butyl alcohol than that of any of the other alcohols, but the acetone solutions are better conductors than those of both butyl and propyl alcohol in most cases. This cannot be explained on the basis of fluidities. The fluidity of acetone is almost as much greater than that of methyl alcohol as the fluidity of the latter exceeds that of ethyl alcohol.2 It is a t least interesting to note that the CH3 group attached to the hydroxyl forms molecules of a solvent much weaker in ionizing power than water. Separation of the CH, group from OH by CHX (CHS-CH2-OH) does not greatly alter the ionizing power of the medium. The addition of a Landolt-Bornstein: Tabellen, 2, 1036 (194 d). Ramsay and Shields: Z. physik. Chem., 12, 433 (1893).
ELECTRICAL COSDUCTIVITY O F ORGANIC ACIDS
I499
second CH, group (CH3-CH2-CH2-OH), however, again causes a very pronounced decrease in ionizing power. The third CH, group (CHa-CHz-CHzCH2-OH) causes very little change. For instance, in the case of dichloroacetic, water solutions show conductances of 20-166 times the conductances in methyl alcohol, dependent upon the dilution. Ethyl alcohol solutions are 14-24 times better conductors than propyl alcohol solutions of the same acid concentration. The conductances in methyl alcohol solutions, however, are only 1.5-2.4 times those in ethyl alcohol and the values in propyl alcohol are only 1.3-2 times the conductances in butyl alcohol. These relations are shown graphically in Fig. I . Regardless of the acid dissolved in the alcohols, the conductivities in the various alcohols stand in approximately the same proportion as mentioned here for dichloroacetic acid. The variations are probably very nearly within the range of experimental error. Most certainly the approximate variations are the same for different acids with the one exception of glycollic acid in the more dilute solutions in methyl and ethyl alcohol. The almost uniform variation in conductivity with change in solvent and the great variations between water and methyl alcohol, and between ethyl and propyl alcohol, with insignificant changes between methyl and ethyl, and propyl and butyl, have no direct relationship to the dielectric constants, fluidities, and association tendencies of the alcohols. I t seems that they must be explained by the effect of the carbon-hydrogen group upon the power of the oxygen atom of the OH group to combine with the hydrogen ion of the acid. I n other words, the presence of the methyl, ethyl, and other radicals affect the basicity or hydrogen accepting properties of the solvent. As might be expected, the acceptance of hydrogen by the oxygen atoms of the hydroxyl groups will also be controlled somewhat by the nature of the substituent atoms or radicals in the acid molecule. The substitution of chlorine in the molecule of acetic acid causes the hydrogen atom of the carboxyl to be more easily detached through the influence of the negative chlorine upon the electron-sharing properties of the oxygen atoms of the carboxyl group. To the extent that it is freed, however, it may be expected that hydrogen ion will be accepted by the oxygen atoms of a particular alcohol to approximately the same extent and that, the acceptance of H by the oxygen atoms of different alcohols will vary in a fairly constant ratio regardless of the acid which acts as solute. The absolute basicity or hydrogen accepting property of the solvent is determined by the electronegativity of the group that is attached to the hydroxyl radical of the alcohol. I t seems, therefore, that an equilibrium between the hydrogen attached to the oxygen atoms of the carboxyl group (and to various other atoms of its molecule which may act as donors‘ of free electron pairs) on the one hand, and the oxygen atom of the solvent molecules on the other, really determines the extent of ionization of the acid and the conductance of the P O ~ U tion. This equilibrium may be shifted in one direction or the other by substitutions in the molecule of solvent or acid. Sidgwick, in his “Electron Theory of Valency,” refers to an atom surrounded in whole or in part by unshared electron pairs as a “donor” atom.
HERSCHEL HUNT WITH H. T. BRISCOE
I500
The basicity or hydrogen-accepting property of acetone is believed to depend upon the presence of an enol modification. Such a modification has been assumed to account, a t least in a measure, for its association and rather high dielectric constant. We note that acetone has a fairly constant ability to accept hydrogen ion, falling just below ethyl alcohol in this regard.
FIGI hleOH Solutjons B. EtOH C. PrOH D. BuOH A.
::
FIG.2 A. Iodoacetic Acid
B. Glycollic Acid C. Chloroacetic Acid D. Acetic Acid
Its acceptance of hydrogen, as in the alcoholic solutions, is limited by the strength of the bonds by which hydrogen is attached to the donor atoms of its own molecule. These facts and ideas lend support to the suggestions of Latimer and Rodebush' and others, that the dielectric constants, molecular association, and ionizing power of the solvent are not directly dependent upon one another, but are to be considered as individually dependent upon the molecular structure of the solvent and solute. Latimer and Kodebush: J. Am. Chem. Soc., 42, 1419 (1920).
ELECTRICAL CONDUCTIVITY OF O R G h S I C ACIDS
I jOI
Bronsted' has defined the basicity constant of the solvent by means of the following equation: KA = KAcld K ~ a s in which KA is the dissociation constant of the acid, KAcrd its acidity constant, and K B ~is~ the . basicity constant of the solvent. Assuming that the acids have approximately the same effect upon the K B ~ of ~ .the medium, it would appear that the basicity constants of methyl and ethyl alcohol are nearly the same, while propyl and butyl have nearly thp same basic strength. The latter are much weaker than methyl and ethyl alcohol, which in turn, are weaker than water. We are considering conductivity data rather than dissociation constants, but this is permissible since it has been shown? that the limiting conductivities of all these acids are approximately the same. So long as the dissolved acid does not affect the basicity constant of the medium or affects the basicities of the different solvents in the same direction and to the same extent, the dissociations of an acid in different solvents will stand in the same ratio to one another as the basicity constants of the solvents stand to one another, regardless of the nature of the acid. This is true, provided that the solvents do not alter in any way the acidity constant of the acid; that is, the tendency to permit the dissociation of hydrogen ion from its carboxyl linkage must be the same in all solvents. If the acidity constant is altered to different extents or in different directions by the solvents, the dissociation ratios will vary accordingly. This effect is determined by the relative electron-sharing properties of the groups attached to the hydroxyl radical of the solvent on the one hand and those of the substituent group in the acid molecule, as well as the position of this substituent with reference to the carboxyl radical and its effect upon the oxygen to hydrogen bond, on the other. Fig. z shows typical curves for the conductances of acids in acetone solutions at different dilutions. Table VI shows that propyl and butyl alcohols affect the conductivity in methyl alcohol almost equally, while the effect of propyl on the conductivity in ethyl alcohol exceeds that of butyl. Experimental The molecular conductances and dissociation constants of several acids in ethyl alcohol at 30°C are given in the following tables. The limiting conductivities of the acids were computed by means of the following equation: L a c i d = & S a salt A,HCl - h,NaCl. The conductivities of the sodium salts at 3ooC were calculated from the work of Lloyd and Pardee.3 The conductivity of KaC1 a t infinite dilution was determined in our laboratory.
+
Bronsted: Chem. Reviews, 5 , 308 (1928).
* Lloyd and Pardee: Pub. Carnegie Inst. Kash., S o . 260, p
99. Lloyd and Pardee: Pub. Carnegie Inst. Wash., X o . 260, p. 99.
HERSCHEL HUNT WITH H. T. BRISCOE
I502
Dilution
hfo. Cond. 3 o T .
400
37 6 40.93 43.6
800 2000
‘
Dilution
Mol. Cond. 30°C
4000
45.6 53.675
00
The conductivity at infinite dilution was found by means of Kohlrausch’s formula: A, = Av (aCI’3).
+
There is a pronounced disagreement in the data dealing with the conductivity of HC1 a t infinite dilution. Goldschmidt’s‘ value is 89.4 a t 2 j°C. This is higher than that first obtained by him (74.3) and also considerably higher than that obtained by Partington and Lapworth* (66.5). Goldschmidt’s value for the limiting conductivity of HCl has been calculated by use of the Kohlrausch formula. Lloyd and Pardee question this method of calculation and recalculate the value at infinite dilution from Goldschmidt’s data, by using the equations of Xoyes3 and Randall.4 Lloyd and Pardee’s value for the conductivity of HC1 at infinite dilution is 82. Goldschmidt has determined the limiting conductivities of the sodium salts of a few organic acids. These and others have been determined by Lloyd and Pardee. Because of differences in the methods of calculating Ao,GoldSchmidt’s values are uniformly about I O reciprocal ohms higher than those of Lloyd and Pardee. Partington6has measured the limiting conductivity of HCl in ethyl alcohol a t oo, 18’, and zs0C, and has found a value for the temperature-coefficient. From these data we have calculated the value of A, for HC1 at 30’ as 7 2 . 4 2 . Using this value for the limiting conductivity of HCl, Lloyd and Pardee’s values for the sodium salts of the acids, and our own value for the limiting conductivity of KaC1, we have determined the limiting conductivities of the acids. Their dissociation constants have been calculated from the well known Ostwald equation I( = a*/(1 - a)\’. Goldschmidt has calculated the dissociation constants for a few of these same acids. We find good agreement between his values and ours when we use the higher values that he gives for the limiting conductivities of HC1 and the sodium salts. Since there are such great discrepancies in the observations of the limiting conductivity of HCl and the sodium salts of certain organic acids, the authors propose to check these values as soon as possible. I n the meantime we have used Partington’s value for HC1 as the only one available a t 3ooC, and Lloyd and Pardee’s data on sodium salts as these cover nearly all the cases which we have investigated. We do not claim absolute accuracy for the constants. Our dissociation constants for the various acids ’Goldschmidt: Z. physik. Chem., 89, 131 (1914). Partington and Lapworth: J. Chem. SOC.,99, 1419 (1911). Xoyes: J. Am. Chem. Soc., 30, 335 (1908). Randall: J. Am. Chem. Soc., 38, 788 (1916). Partington: J. Chem. Soc., 99, 1937 (1911).
ELECTRIChL COSDUCTIVITY O F ORGANIC ACIDS
I503
are only relative, but they enable us to compare the acids on the basis of the effect of substituent groups upon the dissociation of hydrogen from its carboxyl linkage. The dissociation constants obtained for a few acids by GoldSchmidt were calculated from molecular conductivity data in solutions in which the dilution ranged from about 10-20 liters. The constants which we are quoting herewith are obtained from molecular conductivity data a t dilutions of 8, 32, 128,512, and 1024liters. It will be noted that these values, in general, seem to decrease as the dilution increases to what would eventually prove a constant value in very dilute solutions. There is, however, no evidence of constant values over the wide range of dilution a t which we have calculated the dissociation constants.
TABLE VI11 Dilution
o-Sitrobenzoic
m-Sitrobenzoic
g,3;,0-
o-Toluic
0.0718 0.0942 0.0569 0,1239 0,0741 0,IjIs 32 0.1228 0.2952 0.1836 128 j12 0.2393 0.4928 0.3397 I024 0.7019 0,4751 0.3177 Sp. Cond. of alcohol 1.61 X IO-’ 2
8
0.0j94 0.09jr
m-Toluic
p-Toluic
0.0098
0,0772
0.0133
0.1231
0.0107
0.0234 0.0442 0.1273
0.2479 0.3431 0.4852 0.5564
0,0223 0.0383 0.1098
0.1502
0.1520
TABLE IX Dilution
o-Chlorobenzoic
z
0.0208
8
0.0317 0.0604 0.1006
m-Chlorobenzoic
p-Chlorobenzoic
o-Aminobenzoic
0,3485 0.011; 0.4388 0.0219 32 128 0,4919 0.0606 512 0.222j 0.5688 0.1475 0.6319 0.2141 0,4792 1024 Sp. Cond. of alcohol 145 X IO-^ T.4BLE
Dilution 2
o-Hydroxybenzoic
m-Hydroxybenzoic
0.0123
0.0131
0,0299 32 0.0488 128 0 .I079 512 0.2690 1024 0,4149 Sp. Cond. of alcohol 2 0 0 8
o.oj13 0.1036 0.1307 0.1989 0.4530 0.7629
m-Aminobenzoic
p- Aminobenzoic 0.02j7
0.0374 0~1070
0.0421
0 .I754
0.2207
0.0568 0.0957
0.2jIO
0.1270
x
p-Hydroxybenzoic
Benzoic
Phthalic
0.0018
0.0190
0,1352
0.0045
0.0591
o.oz8j
0,1408 0.I833 0,3975
0.0110
0 .I235
0.0376 0.0912 0.1519
0.2553
0.0673 0 .I709 0.2473 x IO-^
0.5222
0.5394 0.8211
HERSCHEL HUNT WITH H. T. BRISCOE
1504
TABLE XI Dilution
Succinic
2
8 32 128
0.0528
512
0.1153
1024
0.202'j
0.0123 0,0239
hIalonic
Oxalic
0.0386 0.0576 0.0905 0.1672 0.3229
0.2016 0,3939 0.6459 1.1630 2.9489 4.2390
Fumaric
Maleic
Adipic
0.2658 0.514j 1.018; 1.9814 3.7961
0.0253 0.0j1j
0.16jj 0.2516 0.3127
5.22oj
o.oojo 0.0138 0.0474
0.1249 0.18j8
Sp. Cond. of alcohol 196 x IO-^
TABLE XI1 Dilution
MesoTartaric
DextroTartaric
Crotonic 0.0044
2
8 32 128 512
0.3727 0.7344 0.9811 I . 1045
I024
1.1415
0.0104 o ,0189
0.0221
0.0448 0.0912 0.2308 0.3'58
0.0425
0,1089 0.1538
Dissociation Constants Note All values recorded below are t o be multiplied by IO-^, except values for the conductivity of the acids at infinite dilution.
TABLE XI11 Dilution
0-Sitrobenzoic
m-Sitrobenzoic
p-Sitrobenzoic
o-Toluic
m-Toluic
8 32
30.8 19.7 18.7 13 .o
9.6
0.66
4.5
0.51
0,43 0.47
3.1 3 .o 2.6
0.45
56.7 Si . 8 2 7 .7 13.8
13.1
30.8 13.3 7.3 6.3 6.1
9.1
0.69
60.3
59.8
61.2
57.67
128 512 1024
of acid
0.94 0.66
p-Toluic
0.34 0.71
Ao
57.6
57.3
TABLE XIV Dilution
8
32 I28 512
1024 A, of acid
o-Chlorobenzoic
3.1 3.1
m-Chlorobenzoic
417
.O
p-Chlorobenzoic
0.45
o-Aminobenzoic
39 . o 15.7 9.2
2.2
jI , 2
0.42 0.8
2.7
17.4
1.17
I1
6.2
10.5
I.24
16.5
165 . o
.7
p-hminobenzoic
5.4 1.7
0.78 0.jj
0.48
I joj
ELECTRICAL COSDCCTIVITY OF O R G A S I C -4CIDS
TABLE SS’ Dilution
8 32 I28
o-Hydrosybenzoic I .6 1.9
m-Hydroxybenzoic
13 .6 0.7
2 3
I .0
I.?I
I024
3 6 4.3
*io of acid
62.4
j12
I
.8
57.6
p-Hydroxybenzoic
Benzoic
73.2
0,075
19.8 8.4 9.9 , 8.5
0.107
0,459 0.64
jj.8
59.5
0.312
Discussion of Results The structure of benzene and its derivatives, from the viewpoint of the electronic theory of valence, is well summarized by the concepts advanced by Huggins’ and Crocker.z Huggins favors a structure of benzene in which the carbon atoms, with the bonds which tie them to hydrogen or other atoms or radicals, are in two parallel planes, each of which contains three carbon atoms. There is much evidence in favor of this view in the results that have been obtained by the analysis of X-ray measurements of benzene derivatives and by absorption spectra studies. Recently MacInnes3 has found evidence to support this view, which is essentially that of the Korner centroid structure for benzene, in his study of the ionization constants of the halogen and methyl substituted benzoic acids in water solutions. MacInnes finds the relative values for the dissociation constants of these acids, when the second substituent is in the ortho, meta, and para position with reference to the carboxyl group, agree with the spatial requirements of the Huggins structure. The hydroxyl acids agree only fairly ell with this relation, however, and the nitro-benzoic series furnishes complete exception. An attempt to apply MacInnes’ methods of calculation and graphical representation to the dissociation constants of these acids in alcoholic solutions does not lead to any conclusive results. Crocker’s theory of the electronic configuration of the benzene molecule has proved helpful in reconciling the views of many investigators who have made suggestions concerning structure and orientation in benzene derivatives. Berliner’ has recently summarized Crocker’s theory and has used it in interpreting his studies of association and vapor pressure measurements of the nitroanilines, mononitrotoluenes, and toluidines. Crocker distributes the electrons in the benzene molecule as follows. Each of the six carbon atoms in the benzene ring is attached to the two adjacent carbon atoms and to its hydrogen atom by a pair of electrons. This accounts for txenty-four of the thirty valence electrons possessed by the six ‘Huggins: Science, 5 5 , 674 (1922); J. .Im. Chem Soc., 44, 1607 (1922); 45, 264 (1923). * Crocker: J. Am. Chem. SOC.,44, 1618 (1922). AlacInnes: J. Am. Chem. Soc., 50, 2587 (1928). ‘Berliner: J. Phys. Chem , 32, 307 (1928).
I j06
HERSCHEL HUST TVITH H. T. BRISCOE
carbon and six hydrogen atoms. The remaining six electrons are placed between the carbons and in the plane of the ring, as shown:
If an ortho or para orienting group replaces hydrogen, theqe six electrons, which are referred to as the "aromatic" electrons, are attracted toward the I ) 3, and 5 positions and repelled at the 2 , 4 , and 6 positions If the substituent orientq in the nieta position the electrons are attracted towards the 2 , 4, and 6 positions and repelled at positions I , 3, and j
CH,
A The carboxyl group is meta orienting, and so in the molecule of benzoic acid we may expect to have the six aromatic electrons grouped in pairs at the 2 , 4 and 6 positions. Thus, we may say, that the carboxyl substitution tends t o make the carbon atom in positioil one more electro-positive and the carbon atom in position two more electro-negative by the shift that its introduction into the molecule produces in the arrangement of the aromatic electrons. The change is one in the degree of polarity between the carbon atoms of the ring. This effect mill be noticed in the rate and extent of substitution of various atoms or radicals for the hydrogen attached to different carbon atoms, and the polarity of the ring and carboxyl carbons will, in turn, be influenced by the character of the substituent which replaces hydrogen. If an atom of chlorine or any other substituent, which strongly attracts the paired electrons between itself and carbon, is substituted in the ortho position to the carboxyl, the carbon atom to which it is attached becomes less electro-negative in character because of the attraction of the chlorine nucleus and the consequent displacement of the pair of electrons away from the ortho carbon atom. Stieglitz' makes this carbon altogether positive. This view agrees with the shift in electrons in unsaturated hydrocarbons. Kharasch* and others have developed the conception that in vinyl chloride and similar structures the bonding electrons are shifted away from the carbon atom to which chlorine is attached, causing the adjacent carbon atom to assume a relatively electro-negative character. Vpon the substitution of chlorine in the ortho position, therefore, the carbon atom to which the carboxyl group is attached becomes more negative and the carboxyl carbon more positive. The opposite effect is produced when chlorine is attached 1
Stieglitz: J. h m . Chem. Soc., 44, 1293 ( 1 9 2 2 ) . Kharasch: Chem. Reviews, 5 , j 1 1 (1928).
ELECTRICAL COSDUCTIVITY O F ORGAKIC ACIDS
I507
to the carbon atom in the meta position with reference to the carboxyl group. These differences in polarity, although only slight, should cause a t least a noticeable difference in the dissociation of the hydrogen from the carboxyl. RIeta-chlorobenzoic acid should be more strongly dissociated than orthochlorobenzoic acid. If this were true, it would seem to bring into partial agreement the views of Stieglitz, Lewis,’ Kharasch, and others. Lewis’ contention that chlorine causes a shift of the electrons about the nuclei of all the intervening atoms in the same direction, regardless of the presence of unsaturation (“aromatic” electrons), is not consistent with this idea, however. So far as his theory is applied within the carboxyl group itself, there is agreement. I n view of the data supplied by the chemical reactions of halogen acids with the chlorine derivatives of unsaturated compounds, it is believed that we must come to the conclusion, contrary to Lewis, that chlorine substitution results in the production of an electro-negative condition for the carbon atom opposite that to which chlorine is attached. Meta-chlorobenzoic acid is not stronger than the ortho derivative in water solutions. Lewis calls attention to the fact that the dissociation constants for the chlorobenzoic acids are: ortho, 1 . 3 X IO-^; meta, 1.6 X IO-^. He uses this as an argument against the views of Stieglitz, who assumed that chlorine makes the carbon, to which it is attached, positive. Kharasch argues that the halogen benzoic acids should be only slightly more dissociated than benzoic acid, since, even in the ortho position, the halogen is on the beta carbon atom, and in the aliphatic acids the effect of a halogen on a beta carbon is very slight. He notes that the effects in the aromatic acids are of the same order as that produced when the halogen is substituted on the alpha carbon in the aliphatic acids. But here and in his statements concerning the toluic acids, he disregards the six aromatic electrons, the unsaturation of the benzene ring, and the fact that a substituent produces a change in the electropositive and electro-negative conditions of the various carbon atoms of the ring and of the carboxyl. Kharasch also finds the fact, that formic is a stronger acid than benzoic, an argument aginst Lewis’ statement that “methyl alcohol is a weaker acid than water, phenol is a stronger acid.” I n this case he disregards the fact that in benzoic acid, hydrogen is attached to the phenyl group through the carbon atom of the carboxyl and that in phenol the hydroxyl group is attached directly to the phenyl group. If the phenyl group is more electronegative than hydrogen, phenol should be a stronger acid than water, (CgHjO-Ht); and again if phenyl is more negative than hydrogen, the
:0:
.. ,.
.C:O: . .. group is more positive in benzoic acid than in formic, and benzoic acid should therefore be weaker. That the electro-negativity of the phenyl group does Lewis: “Talence and the Structure of .Itoms and Molecules,” 142-146.
HERSCHEL HUKT WITH H.T. BRIGCOE
1508
produce such an effect upon the carbon atom of the carboxyl, and even in a chain attached to it, is evidenced by the fact that the carboxyl group is meta orienting, while the CH2COOH group, in which the carboxyl is attached to the ring by a carbon atom twice removed, is ortho and para orienting. Lewis has pointed out this fact in support of Flurscheim’s’ alternation of “residual affinity” within the ring and in the attached carbon chain. The views which we express have much in common with Flurscheim’s hypothesis. Kharasch goes on to state that ”the only conclusion that one may draiv from these data (ionization constants) is that, ionization is a molecular effect, and that in our iniperfect knowledge of the effect of the solvent in causing ionization . . that it is not permissible to draw concliisions in regard to the electro-negativity of organic radicals from these premises.” I n thip statement we thoroughly agree with Kharasch and are herewith presenting dissociation data in ethyl alcohol solutions in order that these views may be tested in some solvent other than water. We believe that u-ater is not a suitable solvent in which to coinpare the dissociation of organic acids as regards the electro-negativity of the radicals, induced polarities, inner salt formation, or any other property which may depend in whole or in part upon the oresence of polarity, wholly or partially within the acid molecule. The electrical moment of the water molecule and its hydrogen-accepting properties (basicity) is so great that any phenomenon which depends upon the presence of minor degrees of polarity and hydrogen-accepting properties within the acid molecule or its substituent groups is almost completeiy overshadowed by the effects due to the stronger influence of the water molecule. Bronsted2 has shoivn from the data of Larsson3 and Xichaelis and RIizutani4 that et’hyl alcohol is about 600 times as weak a base as water. Briinsted has pointed out the significance of the basicity of the solvent in determining the true acidity constant of the acid from its dissociation constant. I n ethyl alcohol solutions, where the hydrogen of the carboxyl is permitted more completely to display its natural dissociation tendency in the presence of a solvent whose niolecules are not so highly polarized, and in which the oxygen atom of the hydroxyl group is not an extremely strong “donor,” the meta-chlorobenzoic acid is much stronger than the ortho derivative. Para is slightly weaker than the ortho-chlorobenzoic acid, as might be expected, because of the greater distance in this case between the carbon atoms to which the chlorine atom and the carboxyl group are attached. This also holds true for the ortho, para, and meta toluic acids. N e t a toluic acid is decidedly stronger in alcoholic solutions than either the ortho or para isomer, and of the two, ortho is slightly stronger than para. This is in agreement with the electronic configuration of the benzene molecule which has been discussed earlier in this paper. We are not comparing the electronegative character of methyl to hydrogen, but rather the effect of the same
.
f-liirscheim: J. prakt. Chem., (z),66, 321 ( 1 9 0 2 ) ; 71, 497 (1905). Bronsted: Chem. Reviews, 5 , 2 3 1 (1928). Larsson: Dissertation, Lund (1924). * 3Iichaelis and Mieutani: Z.physik. Chem., 116, 1 3 5 (192j).
ELECTRICAL CONDUCTIVITY O F ORGANIC A C I D S
I509
group, CHR,in two different positions with reference to the carboxyl. Crocker considers the methyl group as acting in much the same manner as the chlorine atom, because of the strong nuclear positive charges on the carbon atom and the three hydrogen nuclei attached to it. The results for the hydroxyl, amino, and nitro-benzoic acids are not so clear cut as we have found them to be in the series of chlorine and methyl substituted benzoic acids. Meta-hydroxy-benzoic acid is decidedly weaker than the para, but stronger than the ortho derivative. K e must remember that in both t>hehydroxyl and amino groups the oxygen and nitrogen atoms respectively have strong hydrogen accepting properties. It is the hydroxyl group that gives water and alcohol their basic properties; it is the amino group that so decidedly decreases the dissociation of the acid when XHz replaces a hydrogen in the methyl group of acetic acid. The presence of these atoms, possessing, as oxygen and nitrogen do, free pairs of electrons to which hydrogen ion may become attached in the same manner as it does in forming (H,O)+ and (C2H50H2)+, must certainly affect the dissociation of the acid in a marked manner. And it is only logical to assume that the hydrogenaccepting properties of these groups is influenced by their position on the ring. When the amino group occurs in the ortho position to the carboxyl, there is less tendency for it to combine with hydrogen through the free electron pair on the nitrogen atom than when the KH? is in the meta position. This is due to the fact that the ortho carbon is relatively electro-negative while the meta carbon is electro-positive. When the amino group is attached to the ortho carbon there is less tendency for the nitrogen atom to accept hydrogen; that is, it is less negative. When in the meta position, the nitrogen atom is more negative and hydrogen is accepted more freely. This effect causes the meta derivative to be weaker, as shown by dissociation data, than it would be if the NH2 group had no hydrogen-accepting properties, and operates in a direction opposite to the ortho and para orienting effect of the amino group upon the arrangement of the aromatic electrons and the polarity of the carbon atoms in the ring and in the carboxyl group. The same is true for the hydroxyl derivatives. This explanation agrees with the dissociation data in alcoholic solutions. I n each case, the meta derivative is not as strongly] and the ortho is more strongly, dissociated than one would be led to expect from comparison with the relative values of the two chloro and two methyl derivatives. There are no free pairs of electrons (unsaturated bonds) attached to the carbon atom of the methyl group; while in the case of the strongly negative chlorine atom, although it possesses three free pairs of electrons] there is no indication of its tendency to accept hydrogen. For instance, H:C1: is highly ionized, perhaps
1007~.
I t is of interest to note in this connection that RIacInnes' has found the chloro and methyl benzoic acids to be the only derivatives that regularly fall into accord with his graphical representation of their dissociation conhIacInnes: J. Am. Chem. Soc., 50, 2587 (1928).
HERSCHEL H U S T WITH H. T. BRISCOE
'I
I
\ \
I
:t \\\ I
FIG. 3 A. m-nitrobenzoic Acid B. o-nitrobenzoic Acid C. p-nitrobensoic Acid
I
FIG.4 A. m-chlorobenzoic Acid B. o-chlorobenzoic Acid C. p-chlorobenzoic Acid
stants and the relative distances that separate them from the carbon atom to which the carboxyl group is attached in accordance with Huggins' conception of the electronic configuration of the benzene molecule and its derivatives. Ortho-nitrobenxoic acid is slightly more dissociated in alcoholic solution than the meta derivative. I n the nitro group there is no single atom possessing a nuclear charge sufficiently effective to overcome the repulsive effect of the free electron pairs about the nitrogen and oxygen atoms. The bonding electron pair between carbon and nitrogen, therefore, may be repelled rather than attracted as in the case of chlorine. Meta-nitrobeneoic acid should be weaker, therefore, than the ortho isomer, for the same reason that
ELECTRICAL CONDUCTIVITY O F ORGANIC ACIDS
IjII
the meta form is stronger than the ortho in the chlorobenzoic acid series. Furthermore, the structure of the nitro group is probably
--N::O .. * . :0: as shown by Sidgwick.1 Here the nitrogen atom shares a complete pair of its own valence electrons with oxygen to form a coordinate bond. This results in a slight degree of polarity in which the nitrogen atom becomes slightly more electro-positive. The effect upon the carbon atom to which it is attached is to make that atom slightly more negative, and if this carbon is in the ortho position with reference to the carboxyl group, the acid should be more strongly dissociated than if the nitro group is in the meta position. The presence of unshared pairs of electrons about the atoms that compose the substituent group should have an effect upon the attraction of the substituent for the bonding pair of electrons between it and the carbon atom in the ring. This effect should influence polarity within the ring and also in attached chains. That such is the case may be seen from the fact that the percent of increase in conductivity (over that of ben. zoic acid) is almost directly proportional t o the number of unshared electrons about the various atoms that constitute the substituent. To obtain this result the conductances of the ortho, para, and meta membersof each series must be averaged. When considered individually, successive changes from a group possessing no free electrons (CH,) to one possessing ;wo (r\TH2), then four (OH), then six (Cl), FIG. 5 and finally ten (NO*) show an opposite .L p-hydroxybenzoic Acid B. o-hydroxybenzoic Acid effect when the substituent is placed on C. m-hydroxybenzoic Acid the ortho atom to that obtained when it is in the meta position. The effect in the para position is smaller than that produced when the substituent is in the ortho position with reference to the carboxyl group. Sidgwich. “The Electron Theory of Valence,” 6 j .
1512
HERSCHEL HUNT WITH H. T. BRISCOE
The following additional facts may be noted concerning the relative values of the conductivity of various acids in alcoholic solutions. The conductance of dibasic acids decreases as the number of CH2 groups separating the carboxyl groups increases. I n maleic and fumaric acids, the cis-acid is about ten times stronger than the trans-acid. These relations are the same in water solutions.
FIG.6 A. o-aminobenzoic Acid B. m-aminobenzoic Acid C. p-aminobenzoic Acid
FIG.7 A. m-toluic Acid B. o-toluic Acid C. p-toluic Acid
The effect of the double bond upon the dissociation of the hydrogen atom in the carboxyl is illustrated in crotonic acid. Crotonic acid is from four to five times weaker than butyric acid, as shown by conductivity data in alcoholic solutions. This leads to the conclusion that there is at least a slight degree of polarity between the doubly bonded carbon atoms, and since the carboxyl carbon atom in crotonic acid is evidently more electro-negative than that in butyric acid, we may assume that, in crotonic acid, the alpha carbon atom is electronegative and the beta carbon atom is relatively electropositive. This agrees with the fact that crotonic acid yields beta-bromobutyric acid upon treatment with HBr, while butyric acid reacts with bromine to form the alpha derivative.
ELECTRICAL CONDUCTIVITY O F ORGANIC ACIDS
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Figs, 3, 4, 5 , 6, and 7 show curves for the molecular conductivities against the cube root of the concentration of the different series of substituted benzoic acids.
Summary The conductances of various organic acids in methyl, ethyl, propyl, and butyl alcohol and acetone have been measured a t 3oOC. The conductances of substituted benzoic acids, and a few unsaturated and dibasic acids have been measured in ethyl alcohol. The results have been interpreted in terms of the electron theory of valence and current theories of molecular structure.
