the smoothing time of the electrometer could be increased. An increase in time constant to the order of 30 t o 40 seconds would give a n ultimate sensitivity of about ‘/s p.p.m. for oxygen. EXPERIMENTAL GLC APPLICATION
Recently a small-volume (3-ml.) cell was designed and constructed for experimental application as a GLC detector. Figure 17 shows a section of a recorder chart demonstrating the high sensitivity of the d e t x t o r to carbon tetrachloride. The large peak represents 1 part in 10’0 of carbon tetrachloride in a 1-p1. sample of nheptane. A similar sample of nheptane nominally free of carbon tetrachloride showed a peak about 20% of the magnitude of the one in Figure 17. This peak may represent a trace of carbon tetrachloride impurity or some other impurity with a boiling point close to that of carbon tetrachloride.
SUMMARY A N D CONCLUSIONS
A three-electrode ionization chamber has been constructed, in which it is possible to measure the current due to negative ions. The device can be used to give a quantitative measure of the concentration of trace amounts of electrophores in gas mixtures whose primary constituents remain fixed. The cell may be used to detect electronegative gases or vapors in gas chromatography, because the pure carrier gas guarantees constant background composition. The sensitivity is of the same order as that achieved by operating an ordinary ionization chamber in the voltage region where recombination diminishes the current slightly. A sensitivity of about p.p.m. should be possible for 02, in nitrogen, using a 30- to 40-second time constant in the electrometer. The sensitivity to water vapor is about one third that for oxygen a t the 0.1% level of concentration. The cell has been used to measure
oxygen in butane and butadiene with a sensitivity comparable to that achieved for nitrogen. ACKNOWLEDGMENT
This work was initiated a t the suggestion of the late D. J. Pompeo, who saw analytical possibilities in the fact that the performance of fast ionization chambers was impaired by the presence of traces of oxygen in the gas filling. J. R. Bailey assembled the experimental equipment and took many of the data, and John Nelson fabricated the cell. LITERATURE CITED
(1) Loeb, L. B:, “Basic Process of Gase-
ous Electronics,” p. 22, University of California Press, Berkeley, Calif., 1955. (2) Lovelock, J. E., ANAL.CHEM.33, 162 (1961).
RECEIVEDfor review April 17, 1962. Accepted July 30, 1962. Pittsburgh Conference on Analytical Chemistry and Applied Spectroscopy, March 1962.
Titration Characteristics of Quaternary Ammonium Titrants Influence of Cation Structure G. A. HARLOW Shell Development Co., Emeryville, Calif. ,The structure of a quaternary ammonium titrant may have a profound effect on the apparent strength of a negatively charged acid as determined b y nonaqueous potentiometric titration. The effect is sufficiently great to cause some dicarboxylic acids to titrate as monobasic acids with one titrant and as dibasic acids with another. The strengths of uncharged and positively charged acids are, on the other hand, independent of titrant structure. Potentiometrictitration data are presented for nine quaternary ammonium bases of widely differing cation structure. In some cases these data are supplemented by conductometric titration curves carried out simultaneously in the same titration cell. Both the potentiometric and conductometric curves can b e interpreted on the basis of competition between hydrogen bonding and ion pair formation.
S
THE INTRODUCTION of nonaqueous tetra-n-butyl-ammonium hydroxide in 1956(3,6) many other quaternary ammonium titrants have been INCE
1482
ANALYTICAL CHEMISTRY
investigated. These titrants have been prepared in a variety of solvents by three different methods of preparation. They differ greatly as t o the structure of the quaternary ammonium cation as well as t o the nature of the anion (hydroxide, methylate, etc.). T o the analyst interested in utilizing the great advantages of these titrants, little information is available in the literature which would enable him to choose the best one for his purpose. A4sa matter of fact, there is little in the literature to indicate that they are not all equal and interchangeable. This report is intended to show the profound influence of the structure of the quaternary ammonium cation on the titratability and the apparent strength of certain acids. A later report will deal with the relative stabilities of these titrants. The influence of the titrant solvent on the titration curves obtained Kith a single quaternary ammonium base has been reported previcusly ( 7 ) . In view of this effect and the multiplicity of titrant solvents used in previous studies, no useful comparison of titrant cation structures can be made on the basis of
published curves. It is necessary, as has been done here, to prepare the different titrants in the same concentration and in the same solvent and then t o utilize them in the titration of the same acid under identical conditions. In choosing the quaternary ammonium titrants for this study, an attempt mas made to utilize previously reported titrants where possible and yet provide a wide variety of structures. In all, 10 different titrants were prepared, all in isopropanol solution. Five of these have been described previously in the literature. They are tetrabutylammonium (3, 6 ) , tetramethvlammonium (15), tributvlmethylammonium (6), tetraethylammonium ( l a ) , and trimethylbenzylammonium (17) hydroxides. A sixth titrant. hexadecyltrimethplammonium hydroxide, has been mentioned in conjunction with a method of preparation, (g), but no titration data have been previously reported. The four remaining titrants, tetrapropylammonium, tetraheptylammonium, hexadecyldimethylbenzylammonium, and trimethyl-
TO POTENTlOhlETER A N D ELECTROMETER
BEAKEK
n
PUMP
Figure 1. Apparatus for simultaneous potentiometric and conductometric titrations
tentiometer, using a vibrating reed electrometer (Applied Physics Corp.) as a null detector. With this apparatus the midpoint potentials of negatively charged acids were repeatable to 5 to 10 mv. in duplicate titrations. Conductometric measurements were carried out in the same cell simultaneously with the potentiometric measurements. A pair of bright platinum electrodes approximately 1 cm. square and 5 em. apart were utilized. For the sake of clarity these electrodes are not shown in Figure 1. Conductivity was measured with a Leeds & Northrup a x . bridge. The conductometric data obtained cannot be considered accurate. A small but significant amount of reference bridge solution may have diffused into the titration cell even though the glass frits were the finest obtainable (Corning, ultrafine). Although semiquantitative, these data can be extremely useful in supplementing and interpreting potentio-
Conductometric electrodes not shown
phenylammonium hydroxides, have not been previously reported. The last mentioned compound proved to be so unstable that titration data obtained with it will not be reported. EXPERIMENTAL
Apparatus and Procedure. Two different titration procedures were used. T h e early potentiometric titrations were carried out with a conventional titration procedure (no temperature control) utilizing the ShellPrecision Dual A.C. Titrometer. The accurately weighed sample of acid was dissolved in 100 ml. of solvent in a tall-form 250-ml. titration beaker and the titrant was delivered from a IO-ml. buret. Beckman glass and calomel electrodes were used and stirring was accomplished magnetically. During the titration the beaker was blanketed with a slow stream of nitrogen. The midpoint potentials of negatively charged acids obtained under these conditions were generally reproducible to about 10 to 15 mv. in duplicate titrations. Simultaneous potentiometric and conductometric titrations were carried out with an entirely different apparatus shown in Figure 1. The temperature of the titration beaker was maintained constant a t 25' f 0.2"C. by circulating water from a constant temperature bath through its jacket. A Beckman blue-tip glass electrode was used as the sensing electrode and a sleeve-type calomel electrode (Beckman) was used as refFrence. The reference electrode was in a separate compartment connected to the titration cell with a salt bridge. Both the remote compartment and the bridge contained a solution of tetrabutylammonium iodide in isopropyl alcohol. The reference electrode was cleaned and the iodide solution was replaced by a freshly prepared solution each day. The potentials were measured to 0.2 mv. with a Leeds & Korthrup po-
I
-200
0
I
2 3 VOLKhlE O F T I T R A V T , ml
i
5
Figure 2. Titration of maleic acid in water with tetrabutylammonium and hexadecyltrimethylammonium hydroxides
metric results and will be presented here for that purpose. Reagents. T h e quaternary ammonium titrants used in this study were prepared as approximately 0.1N solutions from their respective halides by the nonaqueous ion exchange method which has been previously described (6). Isopropyl alcohol was used as the titrant solvent in all cases. The water content of the finished titrants was less than o.570. Their stability will be reported elsewhere. As in the past these solutions will be referred to as hydroxides for the sake of simplicity, but they are actually equilibrium mixtures of hydroxide and isopropylate. The acids used as samples were Eastman reagent grade or equivalent and were not further purified. The solvents were the best grade obtainable commercially. Most of the titrations were carried out in isopropyl alcohol (Shell Chemical Co.) which contained less than O.O8YOwater and an insignificant amount of acidic impurities.
POTENTIOMETRIC TITRATIONS
The structure of the titrant cation has little or no effect on potentiometric titration curves of simple acids in aqueous media [for effect on polymeric acids of high molecular weight and high charge density see reference (4).] This is true not only for uncharged and positively charged acids but, as can be seen from Figure 2 (maleic acid), for negatively charged acids as well. The two titrants used here were quaternary ammonium bases of widely differing structures. In nonaqueous solvents of intermediate dielectric constant, such as isopropyl alcohol ( D = 18), uncharged and positively charged acids continue to yield titration curves which are independent of the titrant structure, I n Figure 3, for example, identical curves are obtained with the same two quaternary ammonium titrants for a mixture of perchloric acid, acetic acid, and phenol. I n this solvent perchloric acid reacts with the solvent to yield the solvated proton, a positively charged acid. Acetic acid and phenol are both uncharged acids. The situation is entirely different when negatively charged acids are titrated in nonaqueous media with titrants of dissimilar structures. The titration curves obtained when maleic acid (Figure 4) is titrated with the same two titrants illustrates this point. Although the first acidity of maleic acid titrates in the same manner Kith both titrants, the negatively charged bimaleate ion appears much stronger when titrated with the hexadecylA trimethylammonium hydroxide. comparison of the curves in Figures 2 and 4 shows the tremendous effect of the change in solvent on the apparent
-4Cn
r 0
L
a
[L
L O
I
,
7
8
-i-n4?
Figure 3. Potentiometric titration of mixture of perchloric acid, acetic acid, and phenol with tetrabutyl and hexadecyltrirnethylammoniurn hydroxides in isopropyl alcohol VOL. 34, NO. 11, OCTOBER 1962
1483
50C
40:
30C
POC
E IC: I i
Li
c
-IC.
-2c3
...
-J.,
1
-4;;
I
I
2 3 4 VOLUME O F T i T R A S T , ml
6
5
Figure 4. Potentiometric titration of maleic acid in isopropyl alcohol with tetrabutylammonium and hexadecyltrimethylammonium hydroxides
strength of the two acidities and on the sensitivity of the negatively charged acid to titrant structure. Particularly striking is the fact that maleic acid titrates as a monobasic acid with tetrabutylammonium hydroxide. In the initial phase of this investigation, sulfuric acid was titrated in isopropyl alcohol with eight different quaternary ammonium hydroxides under normal titration conditions. It is well known that in water, sulfuric acid gives only one inflection with both acidities titrating as strong acids. If titrated in isopropyl alcohol with a quaternary arnmoniur& titrant, two good inflections, well resolved from one another, are obtained. Under these conditions sulfuric acid can be considered as a mixture of equal amounts of a positively charged acid, resulting from the first acidity, and a negatively charged acid, the bisulfate ion. If two
different quaternary ammonium hydroxides of widely differing structure are used for the titration, the first acidity will appear t o be of identical strength with both titrants but the bisulfate ion will not. This is shown in Figure 5 for the two titrants tetrabutylammonium and hexadecyltrimethylammonium hydroxide. The other titrants tested show similar results with the first acidity being identical in each case but with the apparent strength of the bisulfate ion varying considerably. The midpoint potentials obtained for the bisulfate ion with the various titrants are listed in Table I. These midpoint potentials will be used throughout this study as measures of apparent acidity. Under the convention used the stronger the acid, the more negative will be its midpoint potential. From a comparison of the approximate midpoint potentials with the titrant structures there is no obvious correlation between apparent acidity and molecular weight. There is, however, a good general correlation between acid strength and the shielding efficiency of the various alkyl group combinations. By shielding efficiency is meant the effective radius or the minimum distance to which another molecule or ion can approach the nitrogen atom. Thus a t one extreme is the tetrabutylammonium ion which has great shielding efficiency because the nitrogen is protected by a n alkyl chain of four carbons in each of four directions. At the other extreme is the hexadecyltrimethylammonium ion which is well shielded in only one direction in spite of its great total bulk. In three directions it is shielded only by methyl groups. Although the single hexadecyl group does not contribute greatly t o shielding efficiency it does perform an important
Table 1. Midpoint Potentials of Bisulfate Ion in Isopropyl Alcohol
Approximate
midpoint
potential, mv. Titrant Tetramethylammoniilm
(precipi-
Hexadecyltrimethylammonium Phenyltrimethylammonium Hexadecyldimethylbenzyl-
ammonium Tetraethylammonium Meth ylt~ributylammonium Tetra ropylammonium Tetragut y lammonium
1484
ANALYTICAL CHEMISTRY
tated) 110
110
0
130 170 180 220 240
I
4
5
i O L L Y E OF 7 T R A l i T m l
Figure 5. Potentiometric titration of sulfuric acid in isopropyl alcohol wiih tetrabutylammonium and hexadecyltrimethyl ammonium hydroxides
-400
I
I
I
i 4 XOLUYE OF I I T R I \ T . ml
!
J
5
6
Figure 6. Potentiometric titration of sulfuric acid in isopropyl alcohol with piperidine, morpholine, and diphenylguanidine
function in the titrant by increasing the solubility of the salts formed when negatively charged acids are titrated. I t is for this reason that hexadecyltrimethylammonium hydroxide, instead of tetramethylammonium hydroxide, is used throughout this study as an example of a lightly shielded titrant. Tetrabutylammonium hydroxide will serve as a good example of a titrant with high shielding efficiency. The influence of titrant structure on apparent strength of the bisulfate ion is apparent in other solvents as well. The same trend was shown when sulfuric acid was titrated in pyridine and in acetone (curves not shown). It is interesting to compare the titration characteristics of some other organic bases with those of the quaternary ammonium hydroxides. The curves for the titration of sulfuric acid in isopropyl alcohol with diphenylguanidine, piperidine, and morphohe are shown in Figure 6. Note that the midpoint potential of the bisulfate ion varies from one titrant to another but that in every case it is 200 mv. or more negative (stronger) than with the quaternary ammonium titrants. A comparison of the structures s h o w that the shielding efficiency in these three amine titrants is very low even when compared with cetyltrimethylammonium hydroxide. The greater the distance between two carboxyl groups of a dicarboxylic acid the smaller will be the influence of one group on the acidity of the other and the lower the K I K z ratio ( 1 ) . A similar relationship would be expected between the distance separating the negative charge from the carboxyl group in the negatively charged acid and the sensitivity of the acid to differences in titrant structure, The experimental data indicate that this can be considered only as a very crude rule of thumb, even for a series of straight chain aliphatic acids such as oxalic,
I I
rl
1
-300
0
acid s h o w a difference of about 65 mv. in the two midpoint potentials while the unsaturated maleic acid shows twice this value. This difference must be attributed at least in part to the ability of the double bond to orient the carboxyl groups in a planar configuration. In maleic acid the groups are fixed in the cis position whereas in succinic acid they are free t o rotate. Phthalic acid serves as another example of a structure in which the carboxyl groups are favorably oriented. It is not surprising therefore, that it yields very similar titration curves under the same conditions. These curves will be shown and discussed in a later section. 1
i I i O L U U E O l TITRA\T
1
I
1
1
rnl
Figure 7. Potentiometric titration of phthalic acid in isopropyl alcohol with tetrabutylammonium and hexadecyltrimethylammonium hydroxides
malonic, succinic, and adipic. When titrated with the two titrants of contrasting structure, tetrabutylammonium and hexadecyltrimethylammonium hydroxides, the first plateaus of each pair of titration curves coincides perfectly, but the second plateaus show varying degrees of separation. The general trend is the expected one: the influence of titrant structure falls off rapidly with the number of carbons separating the carboxyl groups in the acid molecule. The change from oxalic t o malonic, however, constitutes a very significant exception. Despite the added carbon in the bimalonate ion, the separation between the plateaus is equal (within experimental limits) t o the corresponding separation for the bioxalate ion. It is apparent that some factor besides simple coulombic attraction is involved. Hydrogen bonding could be such a factor ( I I , I 4 ) . The titration of the same series of dicarboxylic acids in pyridine with several different titrants showed a similar decreasing effect as the number of carbon a t o m between the carboxyl groups was increased. Again, the difference was large in the case of oxalic acid, considerably less for succinic acid, and very small for adipic acid. The influence of titrant structure on the apparent strength of the bimaleate ion was shown in Figure 4. A comparison of these curves with those for succinic acid (not shon-n) demonstrates how the presence of a double bond increases the sensitivity of a n acid to titrant structure. Although the carboxyl groups in both compounds are separated by tn-o carbons, the saturated
20
I
SIMULTANEOUS POTENTIOMETRIC AND CONDUCTOMETRIC TITRATIONS
It has been previously shown by Higuchi and Rehm (IO), as well as others (Is, 16) that the shape of nonaqueous conductometric titration curves for dibasic acids is influenced by the radius of the titrant cation. It has also been shown by Bruss and HarloTv ( 2 , 5 ) that conductometric data can be combinedwith potentiometricdatato achieve a more complete understanding of the chemistry involved in nonaqueous titrations than is possible with either It seemed adtechnique alone. vantageous, therefore, t o obtain both types of data simultaneously and to utilize both in explaining the influence of the structure of quaternary am-
VOLUME OF TITRANT, ml
Figure 8. Conductometric titration of phthalic acid in isopropyl alcohol with tetrabutylammonium and hexadecyltrimethylammonium hydroxides
hydrogen bonding does, however, increase the acidic strength of the acid ( A ) so that the initial potential of the solution is more negative than would be expected from a monobasic carboxylic acid of similar structure. The addition of titrant to the solution converts the neutral acid to the negatively charged
0
0 I1
v'1
OR
0 .I
monium cations on the apparent strength of negatively charged acids. This approach will be applied first to interpret the behavior of phthalic acid. The potentiometric curves obtained in the titration of phthalic acid in isopropyl alcohol with the two titrants of dissimilar structure are shown in Figure 7 . The corresponding conductometric curves, obtained simultaneously, are shonn in Figure 8. Both figures have been divided into three areas which represent the titration of the neutral acid (I), the titration of the biphthalate ion (11), and the addition of excess titrant(II1). Let us first consider the reaction taking place in region I. Before the addition of titrant, phthalic acid exists in the solution largely as undissociated molecules \vhich contribute little to the conductivity of the solution. Internal
R
0
biphthalate ion ( B ) as shown above. The conversion is acompanied by an increase in conductivity as the undissociated acid is replaced by the more highly dissociated acid salt. Ion pairing of the anion with either quaternary ammonium cation is minimized by the chelate structure which spreads the negative charge over a relatively wide area. Thus there is little opportunity in this region for differences in ion pairing tendency of the two cations t o become apparent, and the conductometric curves are similar in shape. In the present study we are primarily concerned with region I1 of the titration curves as this is where the titrant cation begins to exert its influence. As titrant is added beyond the first equivalence point a profound change takes place in the structure of the anion. This is shown by the following reaction. VOL. 34, NO. 1 1 , OCTOBER 1962
1485
n
c
I
0
+a II
\-)
C-0
-1
t iOR NR4
C-0
6
I
KR, +
I
-
I
ROH
I
/
luRi
(-1
/
/
/
D B
C
It is impossible to specify the sequence in which the ion pair (C) is formed and the proton is captured. A concerted attack by an ion pair of titrant is possible. I n any event the loss of the proton takes place a t a much lower basicity (lower potential) in the presence of that cation which has the stronger tendency to form ion pairs (see Figure 7 ) . The difference in the ion-pairing ability of the two cations is reflected also in the conductometric curves (Figure 8) which diverge sharply in this region. The sharp drop in conductivity obtained with the lightlyshielded titrant is due to the formation of two new sites for ion pairing (the free carboxylate groups) with the addition of each isopropylate (or hydroxide) ion. The ion pairs in turn stabilize the open anion ( D ) . No such sharp decrease in conductivity occurs in the corresponding reaction with the heavily-shielded titrant cation. The cation in this case is incapable of forming an ion pair of comparable stability and consequently a higher basicity must be reached before the chelated proton can be removed and the open anion structure obtained. In the present case the acidity of traces of water in the solvent as well as the acidity of the isopropyl alcohol itself places a limit on the basicity possible in this solution. Thus no inflection is obtained in the potenti -netric curve a t the second equivalence point. Region I11 of the two figures shows the results of adding an excess of titrant. I n the case of the conductometric curves, a sharp reversal 1 of slope occurs with hexadecyltrimethylammonium titrant due to the relatively great dissociating ability of the quaternary ammonium base. The change in slope with the tetrabutylammonium titrant is much more gradual as reaction with the acid continues beyond the calculated equivalence point.
'
If the above explanation is correct, the potentiometric and conductometric titration curves for phthalic acid with other quaternary ammonium titrants should be predictable on the basis of the ability of the cation to form ion pairs. Titrants with cations having a small effective radius should give results similar to those obtained with hexadecyltrimethylammonium titrant while those with large effective radius should resemble tetrabutylammonium titrant in their characteristics. With one exception, the experimental results are as predicted. The midpoint potentials of the biphthalate ion as obtained with six different quaternary ammonium titrants are shown in Table 11. These data were obtained from titrations carried out under the carefully controlled conditions described earlier. Consequently they are considered more precise than the values for the bisulfate ion shown in Table I. A comparison of the midpoint potentials of the biphthalate ion with the structure of the titrant cation shows again a strong correlation between acid strength and effective radius. It is apparent from a study of the structure of the first three titrants and their midpoint potentials that the smallest alkyl group present is most influential in determining effective radius and hence apparent acidity rather than the average of the four alkyl groups. Thus the value for hexadecyltrimethylammonium ion with an average of over
500
400
1
c I
Table 11. Midpoint Potentials of Biphthalate Ion in Isopropyl Alcohol
(25'
C.)
Titrant Tetramethylammonium
Midpoint potential, mv .
Hexadecyltrimethylammonium
Tetraethylammonium Tetra ropylammonium Tetraguty lammonium Tetraheptylammonium 1486
0
ANALYTICAL CHEMISTRY
322 333 394 461 477 500
I
-200 O
,
1 I 2 I 4 V O L U Y E OF TITRAVT. m l
I
I
5
b
Figure 9. Potentiometric titration of phthalic acid in 90 to 10 isopropyl alcohol-water mixture with potassium, hexadecyltrimethylammonium, and tetrabutylammonium hydroxides
0
1
2
3
4
5
6
VOLUhlE OF TITRANT, m:
Figure 10. Conductometric titration of pkthalic acid in 90 to 10 isopropyl alcohol-water mixlure with potassium, hexadecyltrimethylammonium, and tetrabutylammonium hydroxides
four carbons per alkyl is much closer to tetramethylammonium than to tetrabutylammonium (or even tetraethylammonium). It is also apparent from the data in Table I1 that the effect of additional methylene groups rapidly decreases as the chain length of the alkyl groups increases. This is shown by the small decrease in apparent acidity in going from the tetrabutylammonium to the tetraheptylammonium ion. Conductometric titration curves, obtained simultaneously with the potentiometric curves from which the data in Table I1 were obtained, follow the expected pattern with the sharpness of the angles of inflection bearing an inverse relationship to the effective radius of the cation. The behavior of tetraethylammonium is, however, somewhat unexpected. The conductometric curve reflects a greater degree of ion association than would be expected from the cation structure. This behavior cannot be explained a t the present time. I t is of interest to compare the potentiometric and conductometric titration characteristics of an inorganic titrant such as potassium hydroxide which has a small and symmetrical cation with those of the quaternary ammonium titrants. Such a study cannot be carried out in pure isopropyl alcohol because of the limited solubility of the potassium salts of dibasic acids. This difficulty can be overcome, however, by adding about 10% by volume of water to the alcohol. The potentiometric and conductometric titration of
phthalic acid in such a solvent with potassium, tetrabutylammonium, and hexadecyltrimethylammonium hydroxides is shown in Figures 9 and 10. I n considering the effect of the added water on the potentiometric titration curves for the two quaternary ammonium titrants, two striking differences are immediately apparent. First, in contrast to the titration curve in pure isopropyl alcohol (Figure 7 ) ,the titration curve obtained with tetrabutylammonium hydroxide now shows two definite inflections. Thus the effect of the water has been to increase the relative acidity of the biphthalate ion to a sufficient degree (as compared to the solvent) to make it titratable. Second, the difference in the apparent acidity of the biphthalate ion, as titrated with the two quaternary ammonium titrants, has been drastically reduced by the presence of the water. Whereas the midpoint potentials differed by about 150 mv. in Figure 7 , they differ by only about 50 mv. in Figure 9. The midpoint potential in the titration curve obtained with the isopropyl alcoholic potassium hydroxide for the
biphthalate ion is 100 mv. more negative than t h a t for hexadecyltrimethylammonium hydroxide and about 150 mv. more negative than that for tetrabutylammonium hydroxide. It is interesting to speculate as to the great difference in apparent acidities that might be obtained if potassium hydroxide could be compared to tetrabutylammonium hydroxide in anhydrous isopropyl alcohol. The conductometric titration curves obtained in the mixed solvent with the two quaternary ammonium titrants are remarkably different in shape. The tetrabutylammonium curve shows no sharp reversal of slope as does that of hexadecyltrimethylammonium. Apparently water has largely eliminated ion association in the case of the heavily shielded cation. As would be expected, potassium hydroxide shows the greatest tendency to form ion pairs and this behavior is reflected in the shape of its conductometric curve.
(2) B ~ s ED. , B., Harlow, G. A., ANAL CHEM.30.1836 (19581. -, (3) Cundiff,' R. H., Markunas, P. C., Ibid., 28,792 (1956). (4) Gregor, H. P., Frederick, M., J . Polymer Sci. 23,451 (1957). (5) Harlow, G. A., Bruss, D. B., ANAL. CHEM.30,1833 (1958). (6) Harlow, G. A., Noble, C. M., W Id, G. E. A., ANAL.CEEM.28,787 (19567. (7) Harlow, G. A., Wyld, G. E. A., Ibid., 30,73 (1958). ( 8 ) Ibid 34, 172 (1962). (9) Heij';le, H. B. van der, Dahmen, E. A. M. F., Anal. Chim. Acta. 16, 378 (1967). \ - - - . I -
(10) Higuchi, T., Rehm, C. R., ANAL. CHEM.27,408 (1955). (11) Jones, I., Soper, F. G., J. Ch?m.SOC. 1936, 133. (12) Kreshkov, A. P., Bykova, L. N., Mkhitaryan, N. A., J. Anal. Chem. (U.S.S.R.) 14, 530 (1959). (13) Masui, M., J. Pharm. SOC.Japan 75,1519 (1955). (14) McDaniel, D. H., Brown, H. C., Science 118, 370 (1953). (15) Meum, N. van, Dahmen, E. A M. F., Anal. Chim. Acta. 19,64 (1958). (16) Ibid., 21, 10 (1959). (17) Patchornik, A., Rogozinski, S. E., ANAL.CHEM.31,985 (1959).
.
LITERATURE CITED
(1) Bjerrum, N., Z. Physik Chem. 106, 219 (1923).
RECEIVEDfor review May 31, 1962. Accepted August 10, 1962.
Factors Affecting Stability of Nonaqueous Quaternary Am mo nium Tit rants G . A. HARLOW Shell Developmenf Co., heryville, Calif.
b The influence of cation structure, solvent composition, water content, and temperature on the stability of nonaqueous quaternary ammonium titrants has been studied. A comparison of the decomposition rates of 12 titrants of widely differing structure has shown tremendous differences in stability. For example, tetraethylammonium hydroxide in isopropyl alcohol has a half life of about 0.2 days at 50' C., while tetramethylammonium hydroxide has a half life of approximately 26 days under the same conditions. The presence of water has a profound stabilizing action on the titrants but at the sacrifice of basic strength. Inert and basic solvents increase the rate of decomposition. Temperature is very important; a titrant stored in a refrigerator at 15' C. is about 16,000 times more stable than when stored at room temperatures. In contradiction to published data, the decomposition approaches first order kinetics. .This is explained b y assuming an ion pair as an intermediate.
-
A
of quaternary ammonium bases have been proposed for use as nonaqueous titrants for the determination of acidity. Although it is known t h a t these titrants tend to decompose and that such decomposition seriously impairs their usefulness (l7),no information is available on their relative stabilities. An attempt will be made here to provide data on the influence of cation structure, titrant solvent composition, and storage temperature on the rate of decomposition. It is hoped t h a t this information, and the recent report describing the titration characteristics of these compounds (6), will enable the analytical chemist to use existing titrants more efficiently and to devise new titrants with superior qualities. Titrant stability is especially critical in the determination of very weak acids because of the extreme basicity which must be created to titrate these compounds. Because the major route of decomposition of the quaternary ammonium bases involves an attack on the acidic hydrogens of the cation VARIETY
by the anion this high basicity also reduces stability. No attempt will be made here t o review the extensive literature on the decomposition of quaternary ammonium hydroxides in aqueous solution. The subject has been thoroughly discussed by Ingold (IS) and many others. It will be helpful, however, t o summarize briefly one of the theories which has been advanced to explain the relation between structure and stability. Quaternary ammonium hydroxides decompose by two different routes as shown below. NRiOH + NRa ROH (1) NR4OH + NRj H& CHR-9
+
+
+ &0
(2)
Reaction 1 is generally slow compared to Reaction 2 and will be the principal reaction only with titrants such as tetramethylammonium hydroxide. The value of this particular compound as a titrant is greatly diminished by the relative insolubility of its salts in nonaqueous solvents. VOL. 34, NO. 11, OCTOBER 1962
1487
