with satisfactory reproducibility of the oxygen value, gave assurance that the temperature was sufficiently high to attain a complete reduction reaction of these oxides. Optimum temperatures for the determination of oxygen in silica and uranium osides were 2050” and 2400” C., respectively. Heating above 2150” C. was avoided with silica, as silicon carbide, which forms inside the capsule, sublimes a t 2200” C. If allowed to sublime, this would coat the inside crucible walls, thereby inhibiting the active carbon surface. Table I1 lists a number of oxygen determinations successfully obtained by this inert gas diffusion method. Some of the inorganic samples were chemically analyzed. The oxygen found by difference compared favorably with the oxygen results of the direct inert gas diffusion method.
The reaction crucible together with the flared lip thimble when cooled can be lifted out of the reaction chamber and the used capsules removed with tweezers. By reinserting the thimble and the empty crucible into the reaction chamber and degassing for 15 minutes at 2600” C., the assembly can be reused. Blanks should again be well Rithin limits after this degassing time. This operation of emptying and reusing the crucible can be carried out as many times as desired, as long as a film of metal or carbide does not form on the inside of the reaction chamber. This method offers a variety of possibilities on the future application for the determination of oxygen in volatile samples. Precision and accuracy combined with speed and ease of operation make this a highly reliable and economical method for the direct determination of oxygen in a wide variety of materials.
DISCUSSION
ACKNOWLEDGMENT
Tables I and I1 indicate various inorganic samples that have been successfully analyzed by this technique. Four to five determinations can be carried out in any one reaction crucible.
The authors acknowledge the assistance of William E. Chambers, and the x-ray diffraction interpretation of Donald E. Mentzer. The met chemical analyses were performed by Delores
Leonard and Martha F. Sweeney to confirm total add-up of compounds. The assistance of the National Carbon Co. Research Laboratory in carrying out this investigation is deeply appreciated. REFERENCES
(1) Banks, C. V., O’Laughlin, J. W., Kamin, G. J., ANAL.CHEM.32, 1613-
16 (1960). (2) Bennet, E. L., Laboratory Equipment Corp., St. Joseph, Mich., personal communication. (3) Crumpler, T. B., Yoe, J. H., “Chemical Computations and Errors,” pp. 131-4, M7iley, New York, 1940. (4) Elbling, P., Goward, G. W., ANAL. CHEM.32,1610-13 (1960). (5) Kallman, S., Collier, F., Ibid., 32, 1616-19 (1960). (6) Laboratory Equipment Corp., Instruction Manual for Operation of LECO Oxygen Analyzer N o . 534300, 1958. (7) McKinley, T. D., private communication, “Procedure for Preparation of Absolute Oxygen Standard.” (8) Singer, L., IND.ENQ.CHEM.,ANAL. ED. 12. 127 (1940). (9) Smiliy, W. G.; ANAL. CHEM. 27, 1098-102 (1955). RECEIVEDfor review June 12, 1961. Accepted September 1, 1961. Division of Analytical Chemistry, 140th Meeting, ACS, Chicago, Ill., September 1961.
Behavior of Substituted Aromatic Acids in Selected Nonaqueous Solvents ROY R. MlRON California Research Corp., Richmond, Calif.
DAVID M. HERCULES Departmenf of Chemistry, Juniata College, Huntington, Pa.
b The effect of nonaqueous solvents on the acidic behavior of substituted benzoic acids and phenols has been correlated with structural properties of the acids in a variety of nonaqueous solvents, Specifically, a linear relationship between AHNP and pK, has been demonstrated. A correlation has also been established between Hammett’s u value and AHNP in the same solvent series. Four of the acids studied deviated from “normal” behavior and the causes for these deviations are discussed.
S
1936 there has been considerable effort directed toward improving acid-base titrations in nonaqueous solvents, and extensive reviews have been written (11, 12, IS). Streuli and Miron (17) have correlated the behavior of substituted acids in pyridine with their behavior in water INCE
1770
ANALYTICAL CHEMISTRY
by plotting the difference in halfneutralization potential from a standard acid (AHNP) against pK, in water. Recently this work has been extended by Streuli to titration of bases in nitromethane (16). Hall (3) has conducted a n extensive investigation of the titration behavior of a large number of acids in a variety of solvents. The present investigation was undertaken to determine if the relationships established by Streuli and Miron in pyridine held for the titration of substituted benzoic acids in solvents of widely differing dielectric strength and basicity. This was accomplished by studying the relationship between AHNP and pK, for the substituted benzoic acids in a variety of solvents. Also, it was felt desirable to correlate AHNP with some property of the acid, more readily available to investigators than pK, values. Because tables of Hammett’s u values for substituent
groups on the benzoic acid nucleus are readily available, and are related to pK,, we felt that a correlation between AHNP and u would be desirable. The investigation has demonstrated that linear correlations between AHNP and pK. do exist in a variety of solvents, and that linear relationships between AHNP and Hammett’s u exist in the same solvents. Several of the submstituted benzoic acids-namely, aminobenzoic, p-aminobenzoic, pmethylbenzoic, and p-nitrobenzoic acids -show consistent deviations in one or the other of these plots. In most cases, these deviations can be explained in terms of equilibria or resonance considerations. EXPERIMENTAL
Reagents and Solutions. All of the acids studied were either Eastman Kodak White Label grade or were obtained from university stock.
Melting points of all compounds agreed with those recorded in the literature. All solvents were dried and distilled prior to use, taking the middle fraction from the distillation. The solvents, their sources, and drying agents used are: pyridine (J. T. Baker, KOH); acetonitrile (Matheson, Coleman & Bell, AlzOa); 4-methyl-2-pentanone (Matheson, Coleman & Bell, MgS04); 2nitropropane (Commercial Solvents Corp., MgS04); o-nitrotoluene (Eastman Kodak White Label, hlgS04); nitrobenzene (university stock, MgS04) ; N,N'dimethylformamide (Merck reagent grade, MgS04) ; chlorobenzene (Matheson, Coleman & Bell, Al203); bromobenzene (Matheson, Coleman 8: Bell, A1203). Tetrabutylammonium hydroxide was used as the titrant, and was prepared from tetrabutylammonium iodide by the method of Cundiff and Markunas (2). Excess titrant was stored under nitrogen to ensure its long-term stability. The reservoir for the titrant was fitted with a drying tube containing Ascarite and magnesium perchlorate to prevent absorption of carbon dioxide and moisture from the atmosphere. Apparatus. Electrode potentials were measured using a d.c. electrometer designed by Serfass (9, I 4 ) , covering a range of 0 to 1500 mv. A glasscalomel electrode system was used, and the calomel electrode was modified by replacing the saturated aqueous KCl solution with a saturated solution of KC1 in methanol (2). The average deviation of electrometer output for a single reading was ==I 1.2 over the entire 1500-mv. range. The output of the electrometer was connected to a Brown recorder having a chart speed of 1 inch per minute. The syringe drive system was based on a design by Lingane (8),modified to improve operation by utilizing a combination of gears, a rollerless chain, and 1/160-hp. synchronous motor. Delivery rates of 0.48, 0.95, and 1.90 ml. per minute were available. A switching arrangement n-as provided to allow the recorder chart-drive motor and the syringe-drive motor to be activated simultaneously. A 50-ml. hypodermic syringe was used for the titrations, having the tip replaced by a standard-taper 12/5 female ball joint, joined to the syringe barrel through a graded seal. The plunger of the syringe was filled with lead shot to increase drive stability. Procedures. The titration apparatus was calibrated incrementally, using standard solutions of acid and base. The buret was capable of delivering 1.90 ml. per minute, with a standard deviation of = t O . O l ml. per minute. For titrations of acids using the apparatus described, 1.00 meq. of the acid was dissolved in 100 ml. of the solvent. In the few cases of limited solubility encountered, 100 ml. of the solvent were saturated with the acid. KO particular precaution mas taken to exclude carbon dioxide from the titration vessel, since no titration blanks
ACETONITRILE
120
80 P"PlC,Ni
A 13
40 .6C
0
4c
.2c
- 40
4
p - 80
- 2c -4c
-120
- 6C 1
2
3 pK,
4
5
6 -80
(H20)
-40 0 40 AHNP (UILLIYO_TS)
80
I
0
Typical behavior of substituted benzoic acids in nonaqueous
Figure 1 . solvents 1. 2. 3.
4. 5. 6. 7. 8.
9. p-Bromobenzoic 10. p-lodobenzoic 1 1. p-Aminobenzoic
rn-Chlorobenzoic rn-Bromobenzoic m-lodobenzoic m-Aminobenzoic rn-Methylbenzoic rn-Nitrobenzoic rn-Methoxybenzoic p-Chlorobenzoic
12. p-Methylbenzoic 13. p-Nitrobenzoic 14. p-Methoxybenzoic 15. p-Ethoxybenzoic 16. p-lsopropylbenzoic
17. Benzoic
were encountered in any of the solvents studied. Temperature was kept constant at 22' C. The titration curves were recorded automatically and the half-neutralization potential was determined graphically. RESULTS FOR SUBSTITUTED BENZOIC ACIDS
The first phase of the investigation mas to study the relationship between AHNP and pK, for selected benzoic acids in a series of solvents. A linear relationship between these parameters was found in all of the solvents studied. Figure 1,A, shows a plot of AHNP us. pK, for the substituted benzoic acids in acetonitrile, which is typical of plots obtained for all solvents. Table I presents equations defining the relationships between AHNP and pK, for the substituted benzoic acids in all solvents studied. These equations are derived from least squares plots of titration data. Also presented in Table I are the slopes of the least squares lines, in units of millivolts per pK,. The slopes measure the sensitivity of AHNP to changes in pK, for the acids dissolved in the various solvents. Because AHKP-pK, plots have relatively constant slopes from one solvent to the other, one may conclude that the relative behavior of meta- and para-substi-
tuted benzoic acids differs little from solvent to solvent. However, the slopes of the curves obtained in this investigation indicate that the solvents studied are better titration media for benzoic acids than water, and that one may resolve the titration curves for mixtures of two benzoic acids having pK,'s differing by 0.94 pK, unit. Also shown in Table I are the standard deviations from linearity for AHNPpK, plots. From these data it might appear that the correlation between AHNP and pK, is not too precise. However, the deviations presented in Table I are for all acids studied in a particular solvent, and they are heavily weighted by certain acids showing large deviations in most solvents. Therefore, for most substituted benzoic acids, the relationships given in Table I are precise and should be useful to the analytical chemist in predicting the behavior of substituted benzoic acids in these solvents. Because of the consistent relationship between AHNP and pK,, and because of the well known relationship between pK, and Hammett's u value, it was anticipated that there should be a relationship between u and AHNP in any given solvent. Plots of AHNP against u should be of greater utility to the analytical chemist than plots of VOL 33, NO. 12, NOVEMBER 1 9 6 1
1771
Table 1.
Relation of AHNP with pK, and Sigma for Solvents Investigated
Relation with pK. (PKo = )
Solvent Pyridine Acetonitrile CMethyl-2-pentanone %Nitropropane o-Nitrotoluene Nitrobenzene N,N'-Dimethylformamide Chlorobenzene Bromobenzene
0.00755 AHNP 0.00637 AHNP 0.00688 AHNP 0.00719 AHNP 0.00687 AHNP 0.00659 AHNP 0,00695 AHNP 0.00704 AHNP 0.00644 AHNP Table II.
++ 4.21 + 4.17 4.24 + ++ 4.12 4.09 4.17 ++ 4.18 4.10 + 4.16
Deviation from Linearity, Slope, S, pK. Mv./pK, Unit 133 0.11 157 0.09 145 0.14 139 0.17 146 0.16 152 0.15 144 0.11 142 0.15 155 0.11
Deviation from
Linearity, Slope, Mv./u -147 -178 -160 -180 -176 -196 -160 -156 -166
Relation with Sigma, u -0.00678 AHNP 0.053 -0.00560 AHNP 0.073 -0.00624 AHXP 0.027 -0.00555 AHNP 0.157 -0.00568 AHNP 0.215 -0,0051 AHNP + 0 . 9 0 -0.00623 AHNP 0.103 -0.00639 AHNP 0.136 -0.00602 AHNP 0.142
+ +++ +
++ +
s,
rJ
Unit 0.064 0.095 0.086 0.18
0.19 0.16 0.089 0.16 0.14
Deviations from Normal Behavior for Substituted Benzoic Acids
Std. Dev. for Other Acids ApKa AU
PABA PMBA PNBA ApKo Au ApK. Au ApK. A0 0.203 (0.10) -0.101 (0.08) (0.002) (-0.14) Pyridine 0.148 (0.11) -0.145 (0.00) (0.072) (-0.09) Acetonitrile 0.197 (0.17) -0.159 (0.06) (0,020) (-0.13) CMethyl-Zpentanone 2-Nitronronane 0 . -10 0.060 -0.31 10.0121 (-0.181 0.341 -0.39 0,432 (0.22) -0.250 . . --- - r -- r - - L-?Jitrotoluene 0.07 0.075 Insol.-- Ins'ol.---' ( - o . i g j 0.291 -0.38 0.464 '0.31' -0.268 0.362 0.29 -0.331 0.291 -0.33 0.09 0,041 (-0.14) -0,147 (-0.16) Nitrobenzene N,N'-Dimethylformamidea 0.06 0.054 -0.31 (-0.036) (0.03) (0.072) (-0.09) 0.186 (0.14) -0.121 0.11 0.117 Insol. Insol. (-0.02) (0.113) (0.00) (0.064) (0.29) (-0.291) Chlorobenzene 0.09 0.098 Insol. Insol. (-0.10) (0.176) (-0.14) (0.246) (0.14) (-0.071) Bromobenzene An apparent discrepancy exists for the dielectric constant of DMF at 25" C. Leader and Gormley ( 7 )report 36.71 as compared t o 26.6 routinely measured by Merck & Co., Inc. Because Merck's reagent grade DMF was used in these studies, the value of 26.6 has been adopted. Solvent
MABA ApKa AIS (0.044) 0.06 0.023 -0.32 -0.204 0.08 0.052 (-0.16) (0,025) 0.07 0.057 -0.40
AHNP vs. pK., because tables of u values are more readily available than those of pK, values. Such plots will allow extension of AHNP correlations to acids other than benzoic acids but still allow correlation against a common parameter. Figure 1,B, shows a typical plot of AHNP vs. u for the substituted benzoic acids dissolved in pyridine. Table I also presents least-squares equations for plots of AHNP vs. u. Again, the slopes of these lines are relatively constant, indicating the consistency of behavior of the benzoic acids from solvent to solvent. The standard deviations from linearity shown are excessively large because of consistently large deviations in the behavior of three acids. The causes for these deviations and their effects are discussed in the following section. Deviations
of Individual Acids.
In this investigation four of the acids consistently deviated from the "normal" behavior of the other acids. m-Aminobenzoic acid showed large deviations from the linear relationship between AHNP and pK, in several solvents. Rather large deviations from the correlation between AHNP and u were also encountered for p-aminobenzoic acid, p-methylbenzoic acid, and p-nitrobenzoic acid. Table I1 shows the deviations of the four acids listed for plots of A " P US. pK, and AHNP us. u, as well as standard deviations from linearity of the 1772
*
ANALYTICAL CHEMISTRY
other acids but excluding these four. We have generally assumed that deviations for a single acid, larger than three times the standard deviation calculated for all other acids, can be counted as significant. (Deviations for acids not considered significant by this criterion are enclosed in parentheses in Table 11.) The standard deviations from linearity for plots of AHNP us. pK, and AHNP vs. u shown in Table I1 are a better measure of the correlation between these parameters than are those shown in Table I. M-AMINOBENZOIC ACID (MA%). This acid shows a consistent negative deviation in plots of AHNP us. pK,, particularly in solvents of low dielectric constant. Such behavior indicates that MABA is behaving as a stronger acid relative to the other acids titrated in the same solvent. The fact that the para derivative shows no significant deviations from normality in any solvent indicates that the deviations encountered in the meta isomer are inherent with meta arrangement rather than with the amino group itself. MABA exists largely in the zwitter ionic form in aqueous solution (1). MABA is insoluble in several low dielectric solvents in which p-aminobenzoic acid is soluble (see Table 11). Hence, it is probable that MABA exists in the zwitter ionic form when dissolved in the solvents used in this investigation. Therefore, the ionization of MABA would be represented by
COO* I
COO0 I
+ R.H+ @NH2 (where R equals solvent molecule), while the ionization of other acids would be represented by
ooH +
R
e
NHz
V
NHz
I n MABA we encounter little additional charge separation in the unprotonated acid compared with the protonated acid, because the major amount of charge separation has occurred in the formation of the zwitter ion. However, for other acids (like the para derivative) a complete separation of charge must occur when the molecule ionizes. Therefore, it is reasonable to assume that the dielectric properties of the solvent should not affect the ionization of MABA as st,rongly as they would affect the ion-
ization of other acids. Furthermore, as the dielectric constant of the solvent becomes lower, MABA should tend to show stronger acidity relative to the other acids because ionization would be easier, since charge separation in the process of ionization would not be essential. Such an explanation is consistent with the data of Table 11, where MABA shows increased acidity in solvents of low dielectric constant but relatively normal behavior in solvents of high dielectric constant. ~-AMIXOBEKZOIC ACID(PABA). This acid does not show the deviations in AHSP-pK, plots shown by the meta isomer but does show rather large deviations from normal behavior for plots of AHKP os. u in three solvents. The deviations indicate that PABA is behaving as a stronger acid, relative to the other acids, in the solvents 2-nitropropaneJ 2-nitrotolueneJ and nitrobenzene. Large deviations for this acid are encountered only in these three solvents and in no others. The explanation for this behavior of PABA probably lies in a change of resonance interaction between the amino and carboxyl groups in nitro solvents. K’ormally when amino and carboxyl groups are para t o each other, they interact to produce a more negative charge on the carboxyl oxygen because of contributions from structures such as I to the ground state of the molecule.
NH2 I
NOz I
I
COOH I1 This effect causes a decrease in acidity of PABA relative to benzoic acid because of the more negative character of the carboxyl group. However, if the solvent were to interact with PABA in a fashion that would decrease this resonance interaction, the acid would appear to be stronger, relative to other acids. Aromatic nitro compounds complex with aromatic amines (10); a complex between aniline and nitrobenzene can be demonstrated by the deepening of the yellow color of the latter on mixing. The complexes formed between the aromatic’ nitro compounds and amines are A electron complexes (chargetransfer complexes) formed because of differences in formal charges of the A electron systems involved. Since the amino group of PABA has a tendency to donate an electron to the ring, this will result in a formal negative charge on the ring and i t is likely that a weak complex will be formed with
the “positively” charged ring of an aromatic nitro compound. In this case structures like I1 would tend to keep the negative charge on PABA localized within the ring and would reduce the contribution of structures like I to the ground state energy of the molecule. This would result in a relative increase in acidity. MABA shows no deviation in AHKP -u plots for solvents in which PAB.4 shows large deviations. This behavior is consistent with the fact that there is much less resonance interaction between the amino group and the carboxyl group in MABA than in PABA (18).
~-METHYLBENZOIC ACID (PMBA), As can be seen from Table 11, PMBA showed large deviations from normal behavior in plots of AHNP us. u, tending toward greater relative acidity. Significant deviations were observed in all solvents studied except p-chloroand p-bromobenzene. (In these low dielectric solvents deviations in general were large and although the deviations of PMBA are larger than for other acids, they cannot be considered significant.) Deviations for PMBA in nitro solvents were larger than in other solvents. Such behavior indicates that there are probably two factors operating to cause deviations for PMBA, a factor operating in all solvents and a factor specific to the nitro solvents. The large deviations encountered in the nitro solvents for PMBA can probably be accounted for by considerations similar to those for PABA in the same solvents. Normally structures like I11 and IV contribute to the ground state of PMBA, causing a decrease in acidity relative to benzoic acid.
0 .d b H I11
alous behavior when the rates of solvolysis of meta- and para-substituted benxotosylates were correlated with u in acetone-water mixtures. Also, Kloosterziel and Bacher (6) found the behavior of p-toluidine anomalous when the ionization constants of meta- and para-substituted anilines were compared with their catalytic constants for a prototropic rearrangement of a sulfone in n-hexane. Indeed, these authors concluded that u for the paramethyl group is not constant but varies as a function of solvent. The data of Table I1 indicate this to be the case and can be considered as evidence to support the postulate of Kloosterziel and Bacher regarding the lack of constancy of u for the p-methyl group. ~-NITROBENZOIC ACID (PNBA). The data of Table I1 indicate rather constant deviations from normality for PNBA in all solvents studied except those of low dielectric constant. The nature of these deviations indicates that PNBA behaves as a relatively weaker acid. Normally the resonance interaction between the nitro group and the carboxyl group would tend to make PNBA a stronger acid than benzoic acid. However, we are a t a loss to explain what type of interaction could occur in all of the solvents studied that would interfere with this resonance interaction to cause the relative decrease in the strength of PBNA. Therefore, we can merely propose that the u value for the p-nitro group is a function of the solvent, as is the u value for the pmethyl group. The only other justification for such a proposal lies in the fact that Hammett originally found it necessary to propose two values for the u of the p-nitro group, one for use in anilines and phenols and one for use in other compounds. RESULTS FOR SUBSTITUTED PHENOLS
I 0;-OH
IV
On A complex formation with nitro compounds, the contributions of such structures to the ground state energy would be less and the relative acidity of PMBA would increase. The fact that PMBA shows large deviations from normal behavior in all solvents is not readily explained. In general, it appears that pmethylsubstituted derivatives are anomalous when correlations between equlibrium properties or rate properties are attempted with u. Kochi and Hammond (6) have observed that p-methyl and p-methoxy derivatives showed anom-
Because of the general applicability of AHNP-pK. and AHNP-u relationships in a variety of solvents, these investigations were extended to substituted phenols, to see if similar relationships could be established. Titration curves were obtained for the phenols using 4-methyl-%pentanone and N , N ’dimethylformamide as solvents. These solvents were chosen because the substituted benzoic acids showed the least deviations from normality in them. Since completion of this work, Streuli (16) has published an extensive study dealing with the titration of phenols in pyridine and has demonstrated the validity of the AHNP-pK6 relationship in this solvent. Figure 2,A, shows the relationship between AHNP and pK. for the substituted phenols dissolved in 4-methyl2-pentanone. Figure 2,B, shows a plot of AHNP us. u for the substituted phenols in the same solvent. The VOL. 33, NO. 12, NOVEMBER 1961
1773
Figure 2.
Typical behavior of substituted phenols in nonaqueous solvents 1. 2. 3. 4. 5. 6.
-40
Table 111.
Solvent 4-Methyl-2-pentanone N,N’-Dimethylformamide
ANALYTICAL CHEMISTRY
80 120 160 AHNP (MILLIVOLTS)
40
240
200
280
Tabulation of pK. and Sigma Equations for Substituted Phenols
Relation with pKa (PKa = ) 0.0082 AHNP 7.42 0,0264 AHNP 5.24
curves obtained for the phenols titrated in N , N ’dimethylformamide were similar in appearance to those shown. Table I11 tabulates the equations for the AHNP-pK, and AHNP-u relationships, the slopes of these lines, and the deviations from linearity. None of the substituted phenols studied showed any significant deviation from normal behavior in these solvents. The slope for the phenols (mv./pK,) is dehitely smaller than for the substituted benzoic acids titrated in the same solvent. In &methyl-2-pentanone the phenols show a slope of 122 mv./ pK,, while the benzoic acids show a slope of 145 mv./pK.. In D M F the slope for the phenols is 38 mv./pK,, while for the benzoic acids it is 144 mv./pK,. This indicates that these solvents will show poorer resolving power in the titration of phenol mixtures than in the titration of mixtures of substituted benzoic acids. Initially it was hoped to correlate all AHNP data for benzoic acids and phenols by plotting them against the pu product for the Hammett equation (4). Such a relationship was not realized for the solvents studied. Although the phenols show linear relationships between AHNP and u, these relationships are distinct and cannot be correlated 1774
0
Phenol m-Chlorophenol m-Nitrophenol p-Chlorophenol p-Nitrophenol p-Bromophenol
++
Deviation from Slope, Linearity, Mv./pKo pKo Unit 122 38
0.15 0.07
Relation with Sigma ( u =)
with the comparable relationships for the benzoic acids. In plots of p u the phenols gave slopes of -125 and -42 mv./pu in 4methyl-2-pentanone and DMF, respectively, while the benzoic acids gave slopes of -160 mv./po in both of the solvents. Obviously the data for each solvent would give rise to discontinuous curves when plotted on the same graph. Therefore we may conclude that although linear relationships exist between AHNP-pK, and AHNP-u within the phenol family for a given solvent, such relationships are distinct and separate from those of the substituted benzoic acids in the same solvent. ACKNOWLEDGMENT
The authors thank Richard M. King and Fred A. Achey for their assistance in building the apparatus; and Velmer B. Fish, Robert S. Sprague, and E. J. Serfass for their many valuable discussions. One of us (R.R.M.) thanks the United States Steel Foundation of New York for some financial support. LITERATURE CITED
(1) Cohn, E. J., McMeekin, T. L.,
Edsall, J. T., Blanchard, M. H., J . Am. Chem. SOC.56, 784 (1934).
++ 2.153 1,184
0.004 AHNP 0.0120 AHXP
Slope, Mv./u 250 83
Deviation from Linearltp, c Unlt 0.06 0.06
(2) Cundiff, R. H., Markunas, P. C., ANAL.CHEM.28,792 (1956). (3) Hall, H. K., J . Phys. Chem. 60, 63 (1956). (4) Hammett, L. P., “Physical Organic \ - - - - I
Chemistry,” McGraw-Hill, NEWYork,
--
1440
(5) Kloosterziel, H., Bacher, H. J., J . Am. Chem. SOC.74, 5806 (1952). (6) Kochi, J. K., Hammond, G. S., Zbid., 75,3445 (1953). (7) Leader, G. R., Gormley, J. F., Ibid., 73,5731 (1951). ( 8 ) Lingane, J. J., ANAL. CHEM.20, 285 (1948). (9) Miron, R. R., Ph. D. thesis, Lehigh University, 1959. (10) Muller, R. E., WynneJones, W. F. K., J . Chem. Soc. 1959,2375. (11) Reilley, C. N., ANAL. CHEM.28, 671R (1956). (12) Riddick, J. A,, Zbid., 30, 793R (1958). (13) Zbid., 32, 172R (1960). (14) Serfass, E. J., Lehigh University,
unpublished work.
(15) Streuli, C. A., ANAL. CHEM.31, 1652 (1959). (16) Ibid., 32, 407 (1960). (17) Streuli, C. A., Miron, R. R., Zbid., 30,1978 (1958). (18) Wheland, G. W., “Resonance in
Organic Chemistrv.” “ , Wilev. New York. 1955.
RECEIVEDfor review April 6, 1961. Accepted Se tember 11, 1961. Division of Analyticaf Chemistry, 139th Meeting, ACS, New York, N. Y., Se tember 1960. Work performed at Lehigt University, Bethlehem, Pa.
