T H E COXDCCTIVITY OF SOLUTIOXS OF SOME ALIPHATIC ORGANIC ACIDS I N WATER AND ETHYL ALCOHOL* BY HERSCHEL HLWT WITH H. T. BRISCOE
Many investigations have been made to determine the nature of solutions, and many of these have deait with conductivity measurements. But much remains to be done, before our knowledge, particularly of non-aqueous solutions, is complete. This paper deals with the conductances of solutions of various fatty acids and their substitution products in water and in ethyl alcohol. The effect of the substituent upon the conductivity of the acids in both alcoholic and aqueous solutions and the relative effects of the substituent upon the conductivities of the acids in solutions of the two solvents constitute the chief interests of this paper. Both the character and the position of the substituent have been considered. I n so far as conductivity data will serve the purpose, the data have been used to measure the relative ease with which H may be dissociated from the remainder of the acid molecule in each of the two solutions. The effect of substitution upon the strength of the bond has been observed. It is planned to continue and extend these studies to include conductivity determinations of solutions of aliphatic, aromatic, and unsaturated acids, both mono- and dicarboxylic, in various solvents. The dielectric constants, viscosities, and other physical constants of such solutions will be determined. Such data as we hope to collect may be expected to throw some light upon the relation of the character of solute and solvent molecules to the properties of their solution mixtures. Much work has already been done on the conductances of aqueous solutions of organic acids. One may mention the work of Ostwaldl who studied the electrical conductivity of aqueous solutions of some 240 organic acids. Jones2 has studied the conductivity of solutions of many organic acids as affected by temperature and dilution. Very little wdrk has been done on solutions of organic acids in other solvents. Wightman, Wiesel, and Jones3 made conductivity measurements of nine organic acids in ethyl alcohol. Lloyd and Wiese14extended this work to include twenty-nine other acids, most of which were of aromatic character. The study of the conductivity of alcoholic solutions of acetic and other fatty. acids and their substitution products has been neglected and avoided because of the extremely low conductivity of such solutions and the difficulty of at-
* This paper is a part of a thesis presented by the first named author as a partial fulfillment of the requirements for the degree of Doctor of Philosophy in Indiana University. Ostwald: 2. physlk. Chem., 2, 561 (1888); 3, 170,241, 369 (1889). *H. C. Jones: Am. Chem. J., 44, 159 (1910). 3Wightman, Xiesel, and Jones: J. Am. Chem. Sac., 36, 2243 (1914); Carnegie Inst. Wash. Pub., S o . 210, Chap. 111, (1915) ' Lloyd and Wesel: Carnegie Inst. Wash. Pub., S o . 230, Chap. VI1 (191j).
COKDUCTIVITY O F SOME ALIPHATIC ORGASIC ACIDS
191
taining a sufficiently sharp minima to make the reading on the bridge possible. Using cells with very low constants and observing the best methods and all the precautions of modern methods of conductivity measurements, we have been able to obtain easily read minima in all the solutions reported herewith.
Experimental Reagents Eastman's and Kahlbaum's purest acids were further purified and thoroughly dried by standard and well-known laboratory methods. Purity was tested by melting points and boiling points. The water used in the investigation had a specific conductivity of 1.20.91 X I& and was prepared by distillation ( I ) from alkaline permanganate, ( 2 ) from barium hydroxide, and (3) from Kessler's reagent in a block tin still which was fitted with a block tin condenser. S o t all the steam of this final distillation was condensed. The distillate was caught in a quartz container. Ethyl alcohol was treated with sulphuric acid, distilled, and then refluxed for several hours with lime. The lime was prepared by heating hydrated lime for six hours in an electric muffle furnace a t a temperature of 600°-7000c. Five hundred grams of this lime were used for each liter of alcohol. After distillation from the lime the alcohol was fractionally distilled using a long fractionating head and a block tin condenser. I n all this work cork stoppers were carefully wrapped with pure tin foil. The alcohol prepared in this way was found to have a specific conductance of 7 5 X IO-O. In all cases the measurements of conductivity were made on solutions for which the solvents had been prepared on the same or the preceding day. Specific conductivity determinations for solvents were made, of course, directly preceding the preparation of the solutions. Solutions were made up by weight and the various concentrations were prepared by dilution of the mother solutions. Care was taken to make all transfers of liquid away from contaminating air and to prevent evaporation during preparation. I n each case at least two mother solutions mere prepared. Subsequent dilutions were made from each of them. Data were not accepted unless two such independent determinations agreed within less than 0.2 7,. The solutions were made in carefully calibrated Jena glass vessels a t zs0C. Hence, a correction for the expansion of the solution from 25-30' had to be made. The precautions of Morgan and Lammert' were observed in cleaning and filling the cells. A11 glassware used was first carefully treated and steamed to remove soluble impurities. The Kohlrausch bridge assembly was employed in the determinations. A Leeds and Northrup Kohlrausch slide wire bridge with extension coils, tunable telephones, a microphone hummer and carefully calibrated resistances were used. A thermostat filled with water was kept constant to within o . o I O C . Beckmann and 0.1' thermometers were used in measuring temperatures in all the work. These were all calibrated by the U. S. Bureau of hlorgan and Lammert. J. Am Chem. SOC., 45, 1693,(1923).
HERSCHEL H U S T WIT11 H. T. BRISCOE
192
Standards. Xll parts of the bridge assembly and all wiring xvere protected by properly grounded shields. The cells used were of the Washburn type. Their constants were determined by the methods eniployed by Kraus and Parker.' The values of these constants varied from 0.0300to 0.8752. ('ells with the larger constants were used only for water solutions. All parts of t,he apparatus, especially the cells, were frequently checked to observe variations. Readings for each solut,ion \yere made at three different known resistances, time allowancr being made for temperature adjustments. The assembly and procedure in iiiaking up solut,ions were tested by repeating conductivity measurements of solutions of organic acids in water. Such determinations checked the results of Jones, Ostwald, and others very closely. The specific conductances of solutions were corrected for the specific conductance of the solvent. TABLE I In Alcohol 30' Dilution
Icetic
0.00394; 8 0.015226 32 0.0;;696 128 0.218384 jIZ O.8jjOjO 1024 1.329120 2
Propionic
Butyric
o.00~060 0.002967 0.013j30 0.013476 0.01j4oj 0.013380 0.1j6912 0.1444jr 0.585798 0,553142 1.28160; 1.253376
Sp. Cond. of alcohol 7j -214 X
10-3
Iso-Butyric
S-Valeric
Iso-Valeric
0.002876 0 . 0 0 2 8 6 0 0.002498 0.012689 0 . 0 11744 0.009340 0.042663 0.041I jl 0.0390j6 0.134j18 0.12120; 0 , I 11400 O.jq6366 0 . j22 7 j 2 0 . jooor9 1.209628 I . 115629 I .000012
reciprocal ohms.
TABLE I1 In Water 30' Dilution 2
8 32 128 512 1024
2048
dcetic
2,3918 4.9456 9.9256 19.6480 38.1952 53.0125 ,,.2301
*r
Propionic
1,9676 4.2792 8.5664 I;.OIIP
34.0582 45.0150 62.0338
Butyric
1.9498 4.4296 8.9664 1;.8302 34.5497 46.2029 68.56;o
Iso-Butyric
1.8678 4.2572
S-Valeric 2.7044 4.0088
8.2048 8.6400 17.1008 16.3840 34.1jj5 32.6004 46.1414 47.2678 63.7337 63.4880
Iso-Valeric
4,4880 9.1936 18.3296 36.j 8 i 5 5 I .85 j 3 74.526;
SP. Cond. of water 0.94-1. j 6 X IO-^ reciprocal ohms.
TABLE 111 In Alcohol 30° Dilution
S-Caproic
IsoCaproic
MonochlorAcetic
DichlorAcetic
Trichlor.Icetic 2.04740
BromoAcetic 0 . 0 8 180j
0.003086 0.009441 0,059246 o.zj3086 0.008192 0.01j804 0.111056 0.457648 3.68048 0 , I 56424 0.035664 o.oq112j 0.220464 0.895856 6.j9040 0.302099 32 128 0 . 1 5 2 1 2 8 0.141911 0.48094j 1.7619i2 11 ,77165 0.668147 j12 0,635136 0.631603 1.195091 4.336896 20.65433 1.342223 1024 1.136128 0.682199 1.636352 7.622686 27.38585 I ,614028 2
8
Sp. Cond. of E t O H
161 ---zooX IO-^
reciprocal ohms.
* & a m and Parker: J. Am. Chem. SOC.,44, 2422 (1922).
COSDUCTIVITY O F SOME ALIPHATIC O R G A T I C I C I D S
193
TABLE IT' In T a t e r 2 j o * Dilu-
S-Caproic
t ion
IsoCaproic
llonochlor.Icetic
Dichlor.Icetic
-
-
-
-
-
-
-
._
-
32
7.45
-
128
13.89 29.00 40.31
-
2
8
BromoAcetic
TrichlorAcetic
-
68.j
253.1
323.0
127.7
317.5
341.0
122
20.5.8 249 2 I024 * Ostwald. 2. physik. Chem , 3, 170. (1889:.
.3j2.2
353.7 3j6.0
199.20
jI2
72.4
360,1
30
2 4 1 20
Ostwald's values are in Siemens' units.
TABLE T' I n Alcohol 30' Dhticn
CyanAcetic
2
0.2073.;
8
0.32930 0 . j80j4 1.0989j 2.16512 3.01822
32
128 ,j~z
Thioacetic
Pyruvic
0.0801oi
0.16019 0,31040 0 .j948j 1.19232 2.64090 3.14064
Iodo.Icetic
Glycollic
0.065808 0.117980 0 . 2 13873 0,483901
0.041606 0.0i1616 0,121032
0.2j2320 0.606464 0.90j161
1024 Sp. Cond. of EtOH 1 2 7 - 1 6 1 X
1.045504 I . 581466
0.147 508 0 . 2 8 7j08
0.581939 I . 16joj6 ~600jij
AminoIcetir
-
reciprocal ohms
IO-
VI In Kater 'TABLE
Dilu-* tion
8 32 128
j12
CyanAcetic
IOj.3 176.1 260.9 297.3
Thioacetic* Pyruvic'
Iodo-* Acetic
Glycollic*
30.104
-
-
24.79
50.60 94.31 164.jo
47.50 88.00
116.70 2 0 7 . 0 0 1024 * 30' valucs in reciprocal ohms. **Ostmald: Z . physik. Chem., 3, 1 7 0 (1889).
52.05
i9,84 139.10 176.80
59.025
105.644 169.454 236.707 257.720
Amino-" Icetic
0.276;R 0.33164 0.43891 o 66816 2.14016
2.45453
2j'
TABLE VI1 In hlcohol 30' Dilu- a-Bromoprop- b-Bromopropionic ionic tion 2 0.037452 0.02248i
8 128
o.ojj060 0.036360 0.1ji3zo 0.062;84 0.328611 0.11j760
512
0.727219
32
1024
0.303129
0.8j~j84 0.6j660i
Sp. Cond. of EtOII 1 3 j - 2 2 6 X
a-Bromobutyric
a-Brom-isobutyric
0.038j6 o.o;103 0.138j3
0.02;90
0.36122
0,j8080 1.38404
IO-?.
a-BromValeric
0.02j60
o . o j 6 0 ~ 0.05871
a-Bromo-iao Valeric
o.oj6j3 0.089jj
0.10804 0.139jz 0.14861 0.26j41 0,33139 o.zi960 0,j339~ 0.74393 0 . 5 9 2 4 1 1.~0934
o.jj62c)
0 . 8 ~ 0 8
HERSCHEL HUNT WITH H. T. BRISCOE
I94
T.4BLE VI11 In Water 30'
Dilution
a-Bromoprop- b-Bromopropionic** ionic*
-
2
-
110.4 185.3
8 32 128 512
1024
19.48 37.36 ;0.66 95.25
225.0
a-Bromo-* a-Bromo-is0 a-Bromo- a-Bromo-is0 butyric Butyric Valeric Valeric
109,o 180.0
-
-
-
36.84 83.81
-
40.34 79.51 138.80 225.84
157.79 262.63 330.70
218.0
2;1.08
*P. Walden: Z. physik. Chem., 10, 638 (1892). Walden's values are given in Siemens' units at 2 5 " .
Discussion of Results Variations in the conductivity of the same substance in different solvents are usually explained by differences in dielectric constants, viscosities, association factors, and molecular complexities of the solvents. The results of this investigation show that the molecular conductivity in aqueous solution is 40-1000 times as great as the condu;tivity of the same acid in ethyl alcohol. When one considers the conductance of strong electrolytes in alcoholic solutions, these great variations in conductivity values are difficult to explain on the basis of the above-named properties of the solvents. TABLE IX Organic Radicals in the Order of their Effect upon Conductivity Temperature 30' Concentration K/s I 2 Alcoholic Solutions
CH2BrCH2 - 2. CH&H(CH,)CH? - 3. CH3CHzCH&H? 4. CH3CBr(CH3) - 5. CH,CH(CH,) - 6. CHsCHzCHz - 7. CH3CHz - 8. CHaCH(CH3)CHBr - 9. CHzOH-I O . CHz(CH3)CHzCHzCHz - I I . CH3CH2CHzCzHZCHz - I 2 . CH3CHBr - 13. CH3CH2CH2CHBr- 14. CH3CH2CHBr- 15. CH3-16. CHzI - 17. CHzCl-18. CHzBr - 19. CHZ(CS) 1 0 . CHaC'O-::i . CHCl? - 2 2 . CCla-I.
~
~
Aqueous Solutions
CHz(NH2) - CH3CHzCHzCHzCHz - 3. CH3CHzCHzCHZ - 4. CH3CHz-5 . CHaCH(CH3) - 6. CH,CHZCHz - 7 . CH3CH(CHz)CHg - 8. CH3-9. CH2BrCH2 - I O . CHZ(0H) - 11. C H J - 12. CH3CH2CHBr- 13. CH3CHBr- 14. CHzBr - ~ j CHsCl-. 16. CH3CH(CH3)CHBr- I;. CHICO-IS. CHz(CN) - 19. CH3CBr(CH3)- 20. CHClz --? I . cc13-I.
2.
CONDUCTIVITY O F SOME APILHATIC ORGAXIC BCIDS
I95
The effect of increase in dilution is to increase the molecular conductivity. The increase in conductivity in most cases is almost proportional to the volume. This relation is shown graphically in Fig. I For the purposes of interpolation and extrapolation the values of the molecular conductances a t different concentrations are plotted against the cube root of the concentration. It is not possible to estimate the molecular conductivity a t infinite dilution from these graphs becawe sufficiently high dilutions of these weak acids could not be run in alcohol to determine accurately where the curve cut the Y-axis. We expect later to determine the percentage ionization of these acids in alcohol. I n Table I S \ye have arranged the acids in the order of their conductances in the two solvents. This was done in order that we may emphasize the im7 -
FIQ.I
FIQ.2
portant effect which constitution of the molecule of the solute has upon its conductance in the same and in different solvents. Acetic acid, for instance, stands fifteenth in the alcoholic series and eighth in the water series. Although beta-bromopropionic acid is the poorest conductor in alcoholic solutions it conducts better than eight other acids in the water solutions. I n many cases decided shifts in the order are to be noted. The greatest differences are
196
HERSCHEL H U S T WITH I f . T. BRISCOE
found in solutions of those acids containing the groups CH3CHBrCH?-; (CH3)XBr -;(CH3),CHCHBr -; and CBHHWith an increase in CH, groups one would look for a simultaneous decrease in conductivity. This is true in alcoholic solutions until we reach caproic acid, which has a conductance greater than either propionic, butyric, or valeric, and iso-caproic has a conductance greater than normal caproic in concentrated solutions but falls below caproic in a X i 5 1 2 solution. Even this regularity is not found in aqueous solutions. I
1
FIG.3
FIG.4
In alcoholic solutions the iso-acids have a lower conductance than the corresponding normal acids, but in water ieo-butyric has a lower molecular conductance than butyric acid, while is0 valeric acid has a greater conductivity than valeric acid. t-nfortunately. we have only these two pairs of isomers to compare in aqueous solutions. Xn increase in the number of chlorine atoms in thc molecule of acetic acid causes an increase in the rnolecular Conductivity. I t will be not,iced t,hat tht, C'Sradical and the OH radical both cause a great increase in conductance. The substitution of oxygen for the two hydrogens on the alpha-carbon ntum of propionic acid increases its conductance many times. The substitution of S H ? for a hydrogen atoni in acetic acid lowers the conductance almost 945;. This can probably be explained on the assumption that, part of thp hydrogen ions iyhich are formed by the ionization of the carboxyl radical unite with the basic S H ? radical. If such is the case, only after equilibrium has bccn reached in this system will JW h a w the acid liberating hydrogen ions.
COSDUCTIVITY O F SOME ALIPHATIC O R G A S I C ACIDS
I97
Table I S may be looked upon as a series of radicals or groups in the order of their electronegativities. The order, of course, can be looked upon as correct, only in so far as conductivity data alone can be taken as a measure of the strengths of the acids and the ease with which electronic bonds between hydrogen and oxygen in the carboxyl group may be broken. The Stark-Lewis hvDothesis as" _ sumes that the degree of ionization of an acid is dependent upon the strength of the bonds between the oxygen and the hydrogen in the carboxyl group. Furthermore, we would expect from this
FIG.6
hypotheeis that the more electronegative the substituent the more highly ionized the acid Other factors being equal the acid containing the more negative substituent group should have the greater conductance, provided conductance depends only upon the dissociation of the acid by breaking the oxygen-hydrogen bond in the carboxJ.1 group For example, chloracetic acid should be stronger than iodoacetic acid and therefore have a greater conductance \\-e would expect the two acids to have a structure somewhat as follows, since n e know that chlorine is more negative than iodine.
:o:
H :: .. H : C : C : 0 - : --H
..
:c1:
..
:o:
H .. 1: .. H:C:C':O-
..
: I :
..
: -H
HERSCHEL HUNT WITH H . T. BRISCOE
198
I n the acetic acid molecule the electron pair between oxygen and hydrogen lies somewhat nearer the oxygen atom than the hydrogen. When chlorine replaces one hydrogen in the methyl group, the electron pair between chlorine and carbon is pulled away from the carbon and toward the chlorine atom. This causes a rearrangement of the other electron pairs about the carbon atom, which in turn, causes a shift in the position of the electrons about the carbon and oxygen atoms of the carboxyl group. This shift is toward the part of the molecule into which chlorine is introduced. I n the case of the oxygen-hydrogen linkage of the carboxyl group, the electron pair is drawn
c
FIG.8 FIG.7
A. S-butyric acid I t was found that iso-butyric acid fell just a short distance below Ti-butyric acid, having the same curvature.
closer to the oxygen and farther away from hydrogen. This amounts to a decrease in the strength of this bond, and an increase in the extent to which the acid may be expected to dissociate. Since chlorine is more electro-negative than iodine, it should have a more pronounced effect when introduced as a substituent. On these assumptions we would expect chloracetic acid t o have a greater conductance than bromoacetic acid and it in turn to have a greater conductance than iodoacetic acid. This is true in aqueous solutions, but in alcoholic solutions we find the conductance decreases from bromoacetic to chloracetic to iodoacetic acid. Table IX reveals other instances in which conductivity data does not support the theory. It is understood that other factors, such as molecular association, may account for some of these discrepancies.
CONDCCTIVITY O F SOME ALIPHATIC ORGANIC ACIDS
I99
It is interesting to compare the order of organic radicals in Table JX with Kharasch’s table1 of electronegativities. It is found that the alcoholic series agrees with Kharasch’s work but that the water series reveals some differences in position. Conclusions A11 the acids studied have a much lower molecular conductance in alcohol than in water. The rate of increase of molecular conductance for the successive dilu2. tions is about twice as great in alcohol as in water. 3. The homologous fatty acids fall in a series with a decreasing molecular conductance as the number of carbon atoms increase. Isovaleric acid in water solutions and the caproic acids in alcohol are exceptions. 4. The data do not consistently confirm Lewis’ theory of the effect of substitution of groups possessing different electrical characteristics upon the strength of the acid. This conclusion is tnie, of course, only in so far as conductivity data can be taken as evidence of the stability of the electronic bond between hydrogen and oxygen in the carboxyl group. 5 . The acids do not have the same relative conductivities in alcohol and in water solutions. 6. The substituent radicals have been arranged in the order of their effect upon the conductivity of the acid molecule in both water and alcohol. 7. The log-volume conductance curves show agreement with the mass action law a t high dilutions. 8. Simple and efficient means for preparing conductivity alcohol have been given. I.
Khsraech: J. Chem. Ed., 5, 404 (1928).
