454
JOHN H. ELLIOTT AND MARTIN KILPATRICK
T H E EFFECT OF SUBSTITUENTS ON T H E ACID STREKGTH O F BENZOIC ACID. 11
IN METHYL ALCOHOL JOHN H. ELLIOTT A N D MARTIN K I L P A T R I C K
Department of Chemistry and Chemical Engineering, University of Pennsylvania, Philadelphia, Pennsylvania Received August 20, 1940
The purpose of this work is to investigate, by an experimentally consistent method, the effect of substituents on acid strength. The acids chosen for study were the monosubstituted benzoic acids. These acids have the advantages that much theoretical work has been done on the subject of benzene derivatives, and that the geometrical relationships between the groups in the benzoic acids are relatively simple. Ideally, a thorough investigation of relative acid strengths and the effect of substituents on acid strengths would involve the determination of the thermodynamic quantities free energy, entropy, and heat of reaction in a number of solvents. In aqueous solution and in dioxane-water and water-methyl alcohol mixtures, a part of this task has been completed by Harned and his coworkers (18) for the aliphatic, carboxylic acids by measurement of cells without liquid junctions. Very little has been done with this method in anhydrous solvents; the cells employed by Harned and his coworkers are not applicable to the substituted benzoic acids, since these substances are reduced by hydrogen, and the use of the quinhydrone electrode leads to difficulties, due to the interaction of the quinhydrone and the silver-silver halide electrode. There remains the use of concentration cells (17, 30) and other cells involving liquid junctions (1, 15). The conductance method for determining the dissociation constants of substituted benzoic acids has been employed in ethyl alcohol by GoldSchmidt, and a few values are available for the solvent methyl alcohol (13, 14). The kinetic method has been applied in the solvent ethyl alcohol ( 5 ) . The colorimetric method has many advantages as a general method and has been used with success by Kolthoff and Guss (23) and by Minnick and Kilpatrick (27). This method has the advantage that the results can be obtained a t zero ion concentration by a simple linear extrapolation when the carboxylic acid and the indicator are of the same charge type. However, the concentration cell employed by Hammett and Dietz (17) and Wooten and Hammett (30) has the advantage of economy in the use This paper was abstracted from the dissertation presented by John Habersham Elliott to the Faculty of the Graduate School of the University of Pennsylvania in partial fulfilment of the requirements for the degree of Doctor of Philosophy, .4pril, 1940. It was read a t the Xinety-ninth hleeting of the American Chemical Society, which was held in Cincinnati, Ohio, April, 1940.
E F F E C T O F SUBSTITUENTS ON ACID STRENQTH
455
of the solvent and yields fairly precise values of the relative acid strength in a direct measurement. The difficulty of the high internal resistance of the cell can be overcome by the use of an amplifier. The disadvantage of the method is that the results are obtained in the presence of an added salt, such as lithium chloride. EXPERIMENTAL METHOD
The type of cell measured was H B 0.005Ml Licl Au LiB 0.005 M I LiCl 0.045MI
HX 0.005M LiX 0.005 M Au LiCl 0.045M
HB signifies benzoic acid, H X the substituted benzoic acid, and LiB and LiX the corresponding lithium salts. It is necessary to examine closely the assumptions involved in the use of this type of cell to determine the relative acid strength KA.B~. The cell reaction is C~H402H2= CaHaOz
+ 2H+ + 2e
(1)
and the potentials of the half-cells are
or
and
At constant temperature the second and third terms on the right are equal if the electrolyte concentration is the same on each side, the electrolyte being chiefly lithium chloride, and for the same reason [ j H + ] I = [fH+]II (2). Neglecting the liquid junction we have
456
JOHN H. ELLIOTT AND MARTIN KILPATRICK
and putting in the dissociation constant
K
=
CH+C--B CA
at 25OC. we obtain
If I is benzoic acid and
["I ["I p
CB I
C B 11
where KAzBo is the equilibrium constant for the reaction
A,
+ B,*& + B;
A, being the substituted benzoic acid and B, the corresponding base' while A0 and Bo represent benzoic acid and benzoate, respectively. Since this cell involves liquid junctions, it is necessary to test the results against another experimental method and to examine the cell in the presence of various concentrations of lithium chloride. The cell itself is shown schematically in figure 1. Diffusion of the salt-bridge solution contained in A into that portion of the cell containing the buffers B is prevented by the coarsely ground plugs C. These plugs are firmly seated a t the beginning of each series of experiments. The electrodes, D, consist of a number of turns of gold wire, E, which are connected to a metal cap F. The cell caps, G, which hold the electrodes, are fitted to the body of the cell by means of the ground-glass joints H, and the electrodes are held in this cap by the ground joints I. Nitrogen, saturated with the solvent vapor by means of the bubbler, J, is passed through the two buffers. The nitrogen then passes from the cell through the escape vent L, and leaves the cell by way of the bubbler K. As can be seen from figure 1, two electrodes are immersed in each buffer, affording four possible electrode combinations. This makes the detection of a poisoned electrode a simple matter. Very little difficulty was experienced with poisoned electrodes during the course of this work. The method used in cleaning the electrodes is as follows (25): The electrodes are rinsed in tap water and then placed in warm cleaning solution (40°C.) for 15 min. After being rinsed with water containing a little ammonia, they are thoroughly rinsed with distilled water, wiped with an absorbent tissue, and stored in the solvent to be used. Before use they are carefully wiped and immediately placed in the buffer solution to be measured.
EFFECT OF SUBSTITUENTS ON ACID STRENGTH
457
Since the cella studied have a high resistance (ca. lo7 ohms), a thermionic amplifier (4) was used, together with a potentiometer and suitable galvanometer. The cell waa thermostated in an oil bath controlled electrically at 25OC. to fO.l°C.
FIG.1 PURIFICATION OF MATERIALS
The acids were purified by recrystallization if their melting points were lower than the values reported in the literature or if titration indicated that they were impure. The acids were recrystallized from water or a suitable acetone-water mixture, dried in an Abderhalden drier, and stored over phosphorus pentoxide in separate desiccators. The p-hydroxybenzoic acid was recrystallized from an acetone-toluene mixture. The 0- and m-fluorobenzoic acids were prepared by the method of Dippy and Williams (9). Table 1 summarizes the data on the acids. The lithium chloride was prepared by treating the best commercial product with concentrated hydrochloric acid, evaporating to dryness, grinding, and drying to constant weight in an Abderhalden dryer, the final
458
JOHN H. ELLIOTT AND MARTIN KILPATRICK
temperature being 110°C. This method is not as satisfactory as that used by Krieger and Kilpatrick (24), but is convenient when large quantities are required and the product is sufficiently pure for these experiments TABLE 1 Acid data ACID
1 Benzoic.. . . . . . . i
99.8
j
"C.
121.5
!
'C.
121.7
"C. ~
121
~
(10)
Ortho-substituted benzoic acids ~
Substituent:
NOz. . . . . . . . . . I,. . . . . . . . . . . . Br . . . . . . . . . . c1, . . . . . . . . . . . F............. CHs . . . . . . . . . . OCHs. . . . . . . .
OH . . . . . . . . . . .
148.9 162.7-163.5 150.4-150.6 141.9 127.5 104.8 101.0-101.5 158.0-158.5
99.8 100.2 100.1 99.9
*
100.1 99.9 100.0
147.5 162 148 140.7 122 102.4 98 159
148 162.5 150 142 127.5 107.5 101.5
Meta-substituted benzoic acids
Br.,.
, , , ,
..,.
,I
C1. . . . . . . . . . . , ' F . .. . . . . . . . . CH1. , . , . . , , . . OH . . . . . . . . . . , /
A-02. . . . . . . . .
I. . . . . . . . . . . . . Br . . . . . . . . . . . c1, . . . . . . . . . . . F. . . . . . . . . . CHs . . . . . . . . . . OCHs. . . . . . . OH . . . . . . . . . . .
99.8 99.8 100.3 100.2 99.8
99.8 100.1 99.9 99.8 100.4 100.2 99.9 99.9
1' j
156.8-156.9 154.8-155.0 1 123.5-124.5 112.8 1 202.7-203.5
152 154.9 124 110.5 201.3
240.3-240.6 272.5-273.0 256.0 242.1 188.5 181.3 184.0 216.3
242.4 266 251 211.5 182 176.8 184.2 213
1
I
'
155 158 125 112.5
1
211
,
1
i
(7) (11) (7) (7)
1
'
~
1
I
254.5 241 182 181
184
I
* The small sample did not permit titration. in buffer solutions. The alkalinity of the product was never greater than 0.05 per cent as lithium hydroxide and generally less than 0.01 per cent. The quinhydrone from Leeds and Northrup was used without further purification.
EFFECT OF SUBSTITUENTS ON ACID STRENGTH
459
The methyl alcohol was dried by the method of Lund and Bjerrum (26). After reaction with magnesium, the alcohol was distilled and the middle fraction refluxed with anhydrous copper sulfate to remove traces of amines (19). The alcohol was then redistilled, using a Snyder column, and was stored in all-glass containers protected from moisture. The boiling range was never greater than 0.1"C. and determinations of the density indicated a purity of 99.94 per cent. In order to prepare the buffer solutions, a standard solution of lithium methylate was prepared by reacting metallic lithium with methyl alcohol and storing the solution in an all-glass system protected from moisture, carbon dioxide, and oxygen. The buffer solutions were prepared by weighing out the acid, dissolving it in the solvent, adding the required volume of methylate and lithium chloride solution, and making up to 25 ml. in a calibrated flask. TEST OF THE EXPERIMENTAL METHOD
To carry out a measurement the solution for the salt bridge was pipetted into the cell, the plugs placed in position, and the bridge solution allowed to come to thermal equilibrium in the thermostat. The plugs were then firmly seated, the excess solution removed and, after rinsing with the buffer to be measured, approximately 5 ml. of the buffer was introduced. Solid quinhydrone was added to each buffer, the electrodes were fitted in place, and nitrogen was passed through the buffers for 10 min. The E.M.F. of the cell was then measured, using all possible combinations of electrodes. It was found that there was no change of E.Y.F. with time after 15 min. and that small equal quantities of water in each buffer had no effect on the observed E.Y.F.. As a test of the experimental method the following cell was set up 0.005 A l HB I LiCl 0.025 ' i h HB IAu A' 10.005 M LIB 1.00 ill 10.005 M LiB
!
~
Quinhydrone
Quinhydrone
and lithium chlcride was then added to each arm. As there is no appreA ciable correction for the change in the C - ratio due to dissociation the CB ratios are 1 and 5 and the theoretical E.M.F. of this cell is 41.35 millivolts a t 25OC. Table 2 gives the measured E.M.F. as the LiCl/LiB ratio is inLiCl creased. It is seen that the E.M.F. is constant when > 5. LiB As a further check, o-nitrobenzoic acid buffers in stoichiometric CA CB ratios of 1 and 5 were measured in aqueous solution. Here the dissocia~
460
JOHN
a.
ELLIOTT AND MARTIN KILPATRICK
tion of the acid changes the ratios to 0.285 and 1.022 a t 0.045 M lithium chloride. The theoretical E.M.F. is 32.8 millivolts; the observed E.M.F. was 32.9 millivolts. To compare this method with the conductance method, and to see if the ratio of the dissociation constant of the substituted acid to that of benzoic acid changed with electrolyte concentration, a few measurements in aqueous solution were carried out with the cell given on page 455. The results are shown in table 3. In order to compute the correction in acidbase ratio due to the dissociation of the acid, it is necessary to know the dissociation constants of the acids in a solution of ionic strength 0.05, mostly lithium chloride. This was done by measuring the buffer solution against aqueous hydrochloric acid (3). The value of K,for benzoic acid was found TABLE 2 Eflect of lithium chloridc
LiCl/LiBratio . . . . . . . . . . . . . . .
1
1.6
inmilliv millivolts . . . . . . . . . .1 41.4
1 3.2
1
41.5
'
4.8 41.5
~
6.4 16.0 1 40.0 I 41.4 1 41.5 1 41.4 1 41.7 8.0
~
I _ _ -
Comparison
OJ
1
O - N O ~. . . . . . . . . . . . . . . . . . . 0-Cl., . . . . . . . . . . . . . . . . . . . . ~
0-F . . . . . . . . . . . . . . . . . . . . .
i
\
m-NO*. . . . . . . . . . . . . . . o-CH~. ................./
TABLE 3 the E.M.F. and conductance methods i n water
2.009 1.245 0.892' 0.889 0.721 0.306
-
-1. !
-
2.029 1.279
(8)
0.936
(7)
0.715 0.295
(7) (8)
(28)
EFFECT OF SUBSTITUENTS ON ACID STRENGTH
461
urement, and a further comparison will be made after presentation of the experimental results. From these comparisons the conclusion is drawn that the maintenance of constant ionic strength (predominantly solvent salt) in both arms of the cell minimizes the liquid-junction potentials and justifies their neglect. An analysis of the errors involved in the meaaurements yields a precision measure of 0.007 log unit. This is based upon an estimated variation of 6E = 0.3 millivolt.
where
EXPERIMENTAL RESULTS
Table 4 summarizes a typical experiment. Table 5 gives the results for all the experiments at 25°C. in methyl alcohol. I n column 2, the values for the equilibrium constant for the reaction given by equation 8 are listed. The ionic strength is 0.05, and 90 per cent of the electrolyte is lithium chloride. In the third column are given the corresponding values at infinite dilution from the colorimetric data of Kilpatrick and Mears (21). Comparison of columns 2 and 3 shows that the agreement is better than 0.02 log unit in most cases. This means that the two methods give essentially the same values at two different ionic strengths and that the ratio of the dissociation constants of the substituted benzoic acids to benzoic acid is independent of ionic strength up to p = 0.05. Further evidence that the same statement is correct for higher ionic strengths is furnished by the results reported in another paper. This statement applies to the isoelectric equilibrium of equation 8. If we relate the acid strengths to the solvated proton and compare the equilibrium constants for the reaction
A'
+ S e H * S ++ B-
(10)
the large effect of ionic strength in this solvent is immediately apparent. Column 4 gives the negative logarithm of the equilibrium constant for the above reaction computed as in a previous paper (27). To refer the values relative to benzoic acid to the solvated proton requires the knowledge of the dissociation constant of one of the acids in a solution 0.045 M in lithium chloride and 0.005 111 . in the anion of the acid. This value was determined by measurement of the cell
Quinhydrone
Quinhydrone
462
JOHN H. ELLIOTT AND MARTIN KILPATRICK
The addition of water makes a marked difference in the E.M.F. of this cell, and the reproducibility of the results is not as close as in the measurements of buffered solutions. The results with benzoic acid calculated on the basis of complete dissociation of hydrogen chloride in the alcoholic solution give pK, = 8.682 and a corresponding experiment with salicylic TABLE 4 m-Bromobenzoic acid i n methyl alcohol at 26.0"C. Reaction: m-bromobenzoic acid benzoate- = m-bromobenzoatebenzoic acid 0 005MHB 0 005 M HX Cell: ilu 0 005M LiB LiCl 1.000M I 0 005M LiX o 045 il.~L i c l 0 045M LiCl [ Quinhydrone Quinhydrone 1
+
+
1
!
~
X ncidused:
M.p., 156.8-156.9°C.
Purity by titration, 99.8% RUN 1
~
LiCi used: HC1 trastment Purity
