tlip final inflection. This is further dcmonstratrtl in the titration of a mixture of citric. maleic, benzoic, and butyric acids shown in Figure 5 . The dissociation constants and concentrations of the acidic groups present are listed in Table 11. The break is sharp, and the total acidity calculated from the midpoint agrees n i t h the theory. S o organic acids having dissociation cmstants in the range of 10-7 to lo-’ )wre stud!-. Acidic materials having dissociation constants
of about 10-lo (phenol) are not titrated in this system. ACKNOWLEDGMENT
The authors n-is11 to thank A. G. Hersog for his trchnical assistance in carrying out this study. LITERATURE CITED
(1) Am. Soc. Testing Materials, Standards on Petroleum Products and Lubricants D 939-54 (1954).
(2) Jenkins, J. W., J . Am. Oil Chemtsts’ SOC.33. 225 11956). (3) lIizukami, b., Ieki, T., J . Pharnz. SOC. J a p a n 76, 467 (1956). ( 4 ) Samuelson, O., “Ion Exchangers in Analytical Chemistry,” p. 129, Wile\, Sew York, 1053 ( 5 ) Van Etten, C. H., Wiele, >I. B., AKAL CHEW25,1109 (1953). ( 6 ) Keisenberger, E., Makrochenzze Lei Mzkrochzm. -1cta 30, 241 (1942). RECEIVED for review February 26, 1958. Accepted June 24, 1958. Second Ilelaware Vallej- Regional Meeting, ACS, Philadelphia, Pa., February 5 , 1958.
Titration of Weak Acids in Nonaqueous Solvents Potentiometric Studies in Inert Solvents G. A. HARLOW and 0. B. BRUSS Shell Development Co., Emeryville, Calif.
b Phenolic and carboxylic acids have been potentiometrically titrated in inert solvents such as benzene, toluene, and gasoline. Nonaqueous solutions of quaternary ammonium hydroxides are used as the titrant and a glass-calomel electrode pair is used with a vibrating reed electrometer. Normal titration curves are obtained when dilute titrants are employed, but 1 N) an with concentrated titrants additional inflection frequently occurs near the mid-point. These mid-point inflections are believed to be due to acid-anion complexes.
(>
T
potentiometric titration of acids in inert solvents such as benzene, toluene, and gasoline is of interest from both the practical and theoretical points of view. This type of nonaqueous titration has received very little study because of experimental difficulties, one of LT hich is the shortage of suitable titrants. Conventional titrants such as alkali metal hydroxides and alcoholates combine with acids to form salts which are insoluble ininert solvents. These titrants also desensitize the glass electrode, making it almost useless for acidity measurements in solutions of 1-ery lo^ hydrogen ion content. Organic amines, on the other hand, are too weak for the titration of weak acids. The development of nonaqueous auaternari ammonium hydroxides has alleviated the titrant problem (3-6). S o t only are quaternary ammonium salts much more soluble in inert solvents than those of the alkali metals, but also the quaternary ammonium ions show no HE
tendency to inhibit the hydrogen ion response of the glass electrode. The second major difficulty is the accurate measurement of electrode potentials in inert solvents. Conventional titrometers and p H meters do not operate satisfactorily ith these solutions of extremely high resistance. This investigation had as one of its principal aims the development of apparatus which mould be satisfactory for this purpose. Published information on potentiometric titrations of acids and bases in inert solvents such as benzene and toluene is meager. The work of La Mer and D o m e s ( 7 ) , carried out 26 years ago, appears to be the major contribution. They titrated trichloroacetic acid n i t h diethylamine in benzene solution. d quinhydrone electrode was employed, and the potentials were measured with a ballistic galvanometer. The titration curves obtained n-ere normal in appearance. INSTRUMENTAL CONSIDERATIONS
The high resistance of inert solvents leads to two types of difficulties in potentiometric measurements. One, the electrostatic pickup of extraneous potentials, can be greatly reduced or eliminated by careful shielding and grounding of the titration apparatus. The other, the voltage drop across the electrodes and the cell solution, is a function of the current drawn by the voltmeter or p H meter. Titrometers and p H meters generally have input currents of to 10-12 ampere. The best of these instruments
are satisfactory for potentiometric titrations in nonaqueous solvents such a> acids, amines, and alcohols, where the total resistance is 10’0 ohms or less. They are inadequate, however, for inert solvents where resistances a thousandfold greater are encountered. Although the solvent resistance may be decreased by adding polar compounds, this n-ould defeat the purpose of many investigations. RIany of the instruments usually described as electrometers have input currents much smaller than conventional titrometers and p H meters (< ampere). An excellent review of the various types of electrometers has recently been published by Palevsky, Swank, and Grenchik (Q), n h o point out that the dynamic capacitor-type instrument is wperior in several respects. A number of excellent instruments of this type are being inanufactured, mainly for use in radiochemical work. I n this study, the Carey (Applied Physics Corp.) Model 31 vibrating reed electrometer 11-as employed. REAGENTS AND APPARATUS
Titrants. Tetra-n-butylammonium hydrovide in isopropyl alcohol solution was used throughout. (For t h e purpose of simplicity the titrant is referred t o as hydroxide, although it is actually a n equilibrium mixture of hydroxide and isopropylate.) T h e 0.2,V titrant was prepared by the ion exchange procedure (4). The stronger titrants were prepared by two different methods. One consisted of concentrating the 0.2N titrant in a rotating flask evaporator a t room temperature VOL. 30, NO. 1 1, NOVEMBER 1958
1833
and the other was a modification of the silver oxide procedure ( 3 ) . Calculated amounts of tetra-n-butylammonium iodide, silver oxide, and isopropyl alcohol were sealed in a bottle and the mixture was shaken on a mechanical shaker for 4 hours. The titrant was then decanted and filtered under an atmosphere of nitrogen, and stored in a refrigerator. Solvents. Dioxane was prepared by I cfluxing the commercially available solvent over sodium for 2 hours and then distilling. The other titration solvents were the C.P. commercial products used without further purification. Blank titrations were carried out to make certain that acidic impurities nere not present. V a t e r content of the toluene and benzene, as determined hy the Karl Fischer method, was about 0.02%. Apparatus. All titrations w r e carlied out in a shielded titration assembly. Voltages n ere measui ed with a n Applied Physics Corp. vibrating reed electrometer. Beckman glass c>lectrodeswere used. A diagram of the titration apparatus is shown in Figure 1 . The titration beaker is housed in a brass shield connected to a suitable ground. The shield can is fitted with a Bakelite cover which has holes drilled in it for two electrodes, a nitrogen tube, and a buret. An adapter provides a means of connecting the glass electrode lead to the input of the vibrating reed electrometer. The lead from the calomel electrode requires no special adapter. The earlier titrations were carried out with a standard sleeve-type calomel clectrode as the reference electrode. It mas subsequently found that the small amount of potassium ion introduced from this electrode caused large fluctuations in the voltage readings. The electrode was modified, therefore, by substituting a 1 . 0 s solution of tetran-butylammonium chloride for the saturated potassium chloride bridge solution. The use of this electrode resulted in more reproducible voltage readings. Titration Procedure. T h e titrations were carried out in 100 ml. of solvent. The volume of titrant was kept as small as possible in order to introduce very little isopropyl alcohol. Most of the titration solutions contained less than 0.1% isopropyl alcohol when t h e end point vias reached. I n many cases the initial reading TI as very slon to come to a constant value, or nould give a value beyond the potential range of the instrument and nould then be ignored. I n every case, the readings mere stable after the introduction of the first 0.02-ml. increment of titrant. In general, inert solutions equilibrate less rapidly than polar solvents and it is frequently necessary to wait several minutes betneeen additions of titrant to allow the potential to become stable. Sometimes it was necessary to turn off the magnetic stirrer and remove the tip of the buret from the solution before
1834
ANALYTICAL CHEMISTRY
Vibratlng Reed Electrometer
Magnetic Stirrer
Figure 1,
Apparatus for potentiometric titrations in inert solvents
Figure 2. Titration of phenol in various solvents with 0.2N titrant 0
-
[Il
w u m e or
a steady nieter reading could be obtained.
[
'
,
,
;
TI~PXI~
Figure 3. Titration of various phenols in toluene with 1 - 3 4 titrant
EXPERIMENTAL RESULTS
The titration of phenol in various inert solvents is shou-n in Figure 2. The titrant used here vias 0 . 2 5 tetran-butylammonium hydroxide in isopropyl alcohol. Because the average titration was 1 to 2 ml. in volume and the volume of the solvent n a s 100 ml., the cell solution a t the end point contained 1 to 2% isopropyl alcohol. The reference electrode in all of the titrations shown in Figure 2 was the Beckman sleeve-type calon~elelectrode containing saturated potassium chloride as the bridge solution. The titration curves for phenol. o-chlorophenol, nz-nitrophenol, and onitrophenol are shon n in Figure 3. The titrant in this and subsequent titrations was 1.5-Y tetra-12-butylammonium hydroxide in isopropyl alcohol. Because of the concentrated titrant and the large percentage of its volume occupied by the tetra-n-butylammoniuni ion. only about 0.1% isopropyl alcohol was present a t the end point. The reference electrode in these and subsequent titrations was a sleeve-type calomel electrode (Beckman). n hich contained 1 . O S
tetra-n-but,ylanimonium chloride as thc bridge solution. I t is interesting to contrast thc shape of the titrat,ion curve for phenol in toluene obtained n-ith the 1,LY titrant, with the curve in Figure 2 obtained with the 0 . 2 s titrant. The significance of the additional inflection a t the mid-point of the titrat,ion curve in Figure 3 is discussed in the section on intcrpret'ation. The influence of the structure of alkyl phenols on the shape of the titration curves is illustrated in Figure 4, which also s h o m the disappearance of thc niidpoint. inflection, as alkyl groups are added t,o the carbons adjacent tc thc hydroxyl group. The titration behavior of acetica. formic. and trichloroacetic acids is shown in the curves in Figure 5 . the outstanding fpaturcs of n-liich arc their wide span in potential. and the ljrcscrice of mid-point inflections. The titration curves for thrcr dibasic phenols are shown in Figure 6. Curves -4 and B indicate t,hat in these compounds only one of the tn-o acitlitics
800 I
0-H
1
0-
Figure 4. Effect of structure on titration curves for phenols in toluene
500
-C200 2 . 4. b - t n - 1 Butyl P h e n o l
2 -1-Bulyl6-hlelhy:
Phenol 0. I
rnL
Phenol
m
0
I
I
I V0ILIT.C
1
I
I
I
I
j
>I 1,rrant
+ ROH
-4ooC 0 I rnl
/
-500
I
1
,
I
volumeor T > W * W
Figure 5. Titration of carboxylic acids in toluene
is h i n g titratrtl. I t is :ipp:ircnt that tlie oiw titratabk. Iiytirogen is stronger t h n ~ o u l t lbe t,xpected for a simple plrcnol. L' shows tn.o inflertions, int1ic:ttiny the presc~nc~ of tn.o acidities of \vitld>- ctiftermt s t r c q t h . INTERPRETATION OF TITRATION CURVES
Tlir nplic~araiic.c~ of t\vo inflectioiis in tlie titr:itioii w r v for ~ a nionoliasic~acid swiiis haffiing, hiit there is sufficaient \\-ark reported in thc litrrature to offer a logird c>xplanationfor tlii.; phenonit:non. The very reccirt n-ork of van tlvr Hvij(lc: (12) Iias / ) c ~ i i csliecinlly 11c4pful. Con(1ii v t om etr i (S) . rj.osro p ic ( 6 ) :tilt1 i n f r a r d ( 2 . 1 3 ) studies point to thv clxistmce of rartjoxylic :wid-carbox>-l:itc'ion coniplexcs in nonpolar solvents. Su(.li roiiiplexes appear to b ( ~niore st:rljlc th:tn t h r \v(>ll known canrboxj-lic acid tlinicrr. Thc. structure proposed for the cmhoxylic* ac,id~-carbosyl:tte ion c ~ ) n i l ~l)y l (Slaryott ~ (8)is: 4(
//0
R-C
(4
(-)
\OH----()//
~
0
C ' --R
Pimilnr \trurtureq have litrn wggeqtecl
ioiume
of 'T~llrant
Figure 6 . Titration of dibasic phenols in toluene
by Kaufiixiii and Singleterry ( 6 ) , ant1 by Terger and Barron- ( I S ) . That phenols will also forni acid-anion coniplexes has been rerently reported by van der H(,ijde ( 1 2 ) . He explained sonic of the anomalous cwves obtained in tlie potentiometric titration of phenols in pyridine on this basis. Collateral evidence Iiased on conductometric stuclies, and leading to the same conclusion, is disc*ussed in tlic :tccompanying paper ( 2 )' The potentiometric titration curves for both phenols and carboxylic acids in inert solrents give strong evidence for the existence of acid-anion complexes. The existence of two inflections in these curvcs indicates that two acids of different strcngths must be present. Because one of these inflections always occurs near tht. mid-point of the titration, it s e e m likely that a one-to-one coniplex is involved. The type of coniplex which could most readily account for thc rxperimental data is one in which a molecule of acid is hvdrogen bonded to an anion of the acid. i s suming that such a complex exists, the titration of plie~ioli n an inert solvent with isoprogylntc ion could be represented by :
(3)
Reaction 1 is the titration equation for phenol, with RO- representing the titrant anion (in this case the isopropylate ion). As the phenolate ion is fornicd. it is tied up by Reaction 2 into a complex which tends to stabilize it. This complex is apparently murh more stable than a Fimple phenol t i i m c ~ and increases the apparent acid strength of the phenol. Kheri the mid-point of the titration is reached, there nil1 1~ very little free phenol available, arid the hydrogen ion concentration ivill drcrease to the point where the hydrogen of t h e coniplex will begin to titrate. Here it i R the acid form which is stabilizd by the complex, and its acidity will be xeaker than that of the free phenol. dlthough the titrant cation is not shon-ii in Equations 1 to 3, it probably cxists in the solution largeljas an ion pair n-it,li the phenolate ion or the acid-anion complex ( 2 ) . Because the a(-id-aniori coniplex is held together by a hydrogeii bond, it n-ould bc expected to become most apparent under conditions favorablr for strong hydrogen bonding. I t should not o c w r in solvents which themselves form strong hydrogen bonds hecauw of thc competition of the solvent molecules. That this is the case ran be seen from a comparison of the titration c u r w s for phenol in toluene shon-n in Figures 2 and 3. The curves were obt'ainetl under similar conditions. wcept for the amount of isoprcpyl alcohol introduced by the titrant. Isopropyl alcohol. being a polar compound capablc of hydrogen bonding. apparently prevents the formation of the phenolphenolate complex Lvhen present in concent'rations of t h r order of 1%. Further evidence for the existence of a hydrogen-bonded complex in these solrents comes from a study of the tit'ration behavior of nitrophenole. oKitrophenol tends to form strong intramolerular hydrogen bonds, while mnitrophenol molecules tend t o associate with one another ( I O ) . This difference in behavior ariees from the geometric VOL. 30, NO. 1 1 , NOVEMBER 1958
1835
orientation of the hydroxyl and nitro groups. 0
0
Similar differences should exist in the tendency of these compounds to form acid-anion complexes and these can be noted from the titration curves of Figure 3. Only a single inflection is obtained with o-nitrophenol, while V I nitrophenol gives an additional inflection a t the titration mid-point. The substitution of a chlorine atom in the ortho position in phenol should not greatly decrease its tendency to form the complex, because the internal hydrogen bond between oxygen and chlorine is relatively weak. The t n o inflections in the titration curve for o-chlorophenol (Figure 3) bear out this view. The tendency of phenol t o form hydrogen bonds is known to be reduced by alkyl substituents on the carbon atoms adjacent to the hydroxyl group. The o-alkyl phenols should thus form m-eaker acid-anion complexes. The titration curves for phenol and three alkvl phenols (Figure 4) shom that the mid-
point inflection disappears as the degree of shielding is increased. The sulfhydryl group has less tendency to form hydrogen bonds than the hydroxyl group. The thiophenol-thiophenolate complex should thus be leqs stable than the phenol-phenolate complex. The single inflection in the titration curve for thiophenol (not shoiin) supports this vien-, The explanation given for the shape of the titration curves for phenols also applies to rarboxylic acids. As s h o m in Figure 5, t n o inflections are obtained for three common carboxylic acids, acetic, formic, and trichloroacetic, in the solvent toluene. The range of potentials spanned by these titrations is noten-orthy-over 800 mi-. for acetic acid, 1100 mv. for formic acid, and 1200 mv. for trichloroacetic acid. Dibasic phenols exhibit similar titration behavior in toluene as in other nonaqueous solvents ( 1 1 ) . The titration curves for the three dibasic phenols shown in Figure 6 can be explained n ithout the assumption of an acid-anion complex. K h e n the two hydroxyl groups are favorably oriented as in the structure shown, internal hydrogen honding causes one of the hydrogens to OH
/
0-H \
become more acidic, while the other decreases in acid strength to the point
where it cannot be titrated a t all. The titration curves for t n o compounds of this general structure are shown in curves A and B of Figure 6. T h e n the hydroxyl groups of a dibasic phenol are oriented so as to preclude the popsibility of internal hydrogen bond formation, the titration curve shows inflections for both hydrogens (Figure 6, C). LITERATURE CITED
(1) Barrow, G. M.>J . d m . C'henl. S i c . 78, 5802 (1956). ( 2 ) Bruss, D. B., Harlow, G. .I.,. ~ A L CHEV 30, 1836 (1958).' (1958). CHEX (3) Cundiff, R. H., LIarliunas, P. C C.,, IIbzrl., bzrl,
28,792 (1956).
C. M.,11-JId, (4) Harlow, G. A., Xohle, C G. E. --I.,IIbid., b i d , 28, 787 (1956). ( 5 ) Harlon, G. A, 1 Wyld. 1-Id. G. E. .I.,Zbid 30. 69 11958). 11058). (6) Kaufman, S.,Singleterry, C. R.,J . Phus. Chem. 56, 604 (1952). (T) La Mer, V. K., Don-nes, H. C., -1. -4m. Cheni. Soc. 53, 888 (19311. (8) Marvott, A. A,, J . Research AVat/.Bur. Standards 38, 527 (1947). (9) Palevsky, H., Sn-ank, R. IC., Grenchik, R., ReL. Scz. Instr. 18, 298 (1947). (10) Pauling, L., "The Sature of the Chemical Bond," 2nd ed., Cornel1 Cniversity Press. Ithaca, S. T., 1940. (11) Sprengling, G. R., J . A m . Chet,) Scc. 76. 1190 11954). (1%)van 'der Heijde, H. G., A n d . Chzni. Acta 16, 392 (1957). (13) Yerger, A , , Barron-, G. R-., J . A m . Chem. Soc. 77, 44T5 (1955). ~
RECEIVED for review February 25, 1958. Accepted June 23, 1958. Division of A4nalytical Chemistry, 133rd Meeting, AiCS,San Francisco, Calif., April 1958.
Titration of Weak Acids in Nonaqueous Solvents Conductometric Studies DOUGLAS 8. BRUSS and GERALD A. HARLOW Shell Development Co., Emeryville, Calif.
b Conductometric titrations of phenols were investigated in media of low dielectric constant using both alternating and direct current methods. Phenols which are not sterically hindered exhibit conductance mid-point maxima which indicate association of exceptional strength in benzene, xylene, toluene, carbon tetrachloride, pyridine, acetone, and methyl isobutyl ketone. ' Sterically hindered phenols do not exhibit maxima in the conductance curve. Formation of an ion pair between the titrant cation and the hydrogen-bonded acid-anion complex is postulated to account for the observed effects. 1836
ANALYTICAL CHEMISTRY
I
acid-base behavior in nonaqueous solution, potentiometric ( 5 )and conductometric titrations of weak acids were studied. Conductometric methods, particularly those measurements a t low frequencies, have found little application in the determination of weak acidity when compared n-ith potentiometric methods. Several investigators, hon ever, h a r e studied the conductometric behavior of \Teak and strong acids in nonaqueous solution. Lippmaa (12),using a solvent mixture of dimethylformamide, diethylamine, and triethylamine, titrated several phenols a t high frequency using potassium methoxide as the titrant. Similar N IKvE8TIG.4rING
titrations were performed by Lane (11) in ethylenediamine using sodium methoxide as the titrant. Phenols and phenolic mixtures were titrated by Karrnian and Johansson (9) a t high frequencies. They obtained resolution of certain mixtures using potassium methoxide as the titrant in benzenemethanol solution. Other authors have investigated acids of stronger nature in nonaqueous media. Higuchi and Rehm ( 7 ) titrated mixtures of sulfuric and hydrochloric acid in glacial acetic acid using acetate ion as the base. Masui (14) and Ishidate and Nasui (8) applied high frequency methods to the titration of dicarboxylic
