Acid-Base Equilibria in tert-Butyl Alcohol L. W. MARPLE and G. J. SCHEPPERS Chemistry Department, Iowa State University, Ames, Iowa
b The acid-base equilibria in tertbutyl alcohol i s reviewed. Reiterative calculations of acid dissociation constants are improved. New values of dissociation constants show better internal agreement between measurement techniques than previously reported. Results of surveys of conductance titrations and EDTA chelation reactions are presented. Reactions of Mg, Ca, Cd, and Cu with the ditetrabutylammonium salt of EDTA give two types of precipitates in tertbutyl alcohol. They are of the form MH2Y and (MCIOJHMY and their precipitation depends upon the quantity of metal in solution, and the ratio of reactants. These species are more acidic than the EDTA salt, and can be determined b y potentiometric titration using tetrabutylammonium hydroxide titrant.
F
a PRACTICAL STANDPOINT, tertbutyl alcohol is an excellent solvent for the titration of acidic compounds. The solubility of most organic acids is sufficient to do excellent titrimetric work, even when the molecular weight of the acid is high. Most monofunctional acids give flat potentiometric titration curves indicating negligible acid-anion association or hydrogen-bonding in the solvent. Very weak acids, such as phenols, can be easily titrated in this solvent. Hindered phenols, such as 2,6di-tert-butyl phenol, and other weakly acidic compounds such as m-hydroxybenzoic acid, can be titrated potentiometrically or conductometrically with accuracy. These characteristics prompted further investigation of equilibria in this solvent. This paper reports the results of a survey of conductance titrations, and chelation reactions in pure tert-butyl alcohol. Pertinent to this study is the previous work of Marple and Fritz (6). These workers investigated the utility of the glass electrode for the measurement of hydrogen ion activity, and determined dissociation constants of acids and their corresponding salts in pure tert-butyl alcohol. The standard potential for a glass electrode was determined from the e.m.f. a t unit concentration perchloric acid in the presence of tetrabutylammonium perchlorate (evaluated by extrapolation of an e.m.f. us. concn. HC104 plot), and dissociation constants ROJI
obtained from conductance measurements. Dissociation constants were determined by the method of Fuoss and Kraus (S), and were in the range expected when compared to those in isopropyl alcohol solvent, ( 7 ) . Once the standard potential was established for an electrode, the Nernst equation was used to calculate hydrogen ion activities. Overall dissociation constants, for 1: 1 acids, of the type defined by Kolthoff and Bruckenstein (6) were found to be equivalent to the product of two conditional constants when a large excess of tetrabutylammonium perchlorate was present in solution. The measurement of conditional constants by e.m.f. and absorption spectrophotometry has been discussed in detail (6). Dissociation constants for benzoic acid and 2,4dinitrophenol have been recalculated by a more extensive reiteration procedure, which took into account the variation of activity coefficients with ionic strength. These constants showed better internal agreement than those previously reported ( 6 ) ,and are given in Table I.
Table 1. Over-All Dissociation Constants of Acids in terf-Butyl Alcohol Solvent
Solution components Log KHX 2,CDinitrophenol HDNP BuaNClOi -10.77 -10.53 HDNP BuaNUNP HDNP BuaNC10a -10.43 BuaXDNP Benzoic acid
+ ++
+
-14 98
EXPERIMENTAL
Conductance titrations in tert-butyl alcohol were performed using an Industrial Instrument conductivity bridge RC 16132 and bright platinum electrodes about 1 om.* in area. The cell constant (0.0715) was determined from the conductance of O.0200N KC1. All measurements were made a t 27' C. All acids used were reagent grade, with the exception of p-methoxyphenol,
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MI BASE Figure 1. Conductance titrations of phenols with 0.1 64M tetrabutylammonium hydroxide in tert-butyl alcohol solvent A. 8.
p-Methoxyphenol, 0.35 mmole p-Methoxyphenol, 0.70 mmole C. o-tert-Butylphenol, (displaced 1 .OO ml.) D. o-Phenylphenol, (displaced 1.00 ml.) E. o-Ethylphenol, (displaced 2.25 ml.) F. Thymol [displaced 3.25 ml.) G. Phenol (displaced 3.50 ml.)
VOL. 38, NO. 4, APRIL 1966
e
553
’I-
;
251
4
4 I
0
.02
.04
.06 C
~
.00
.IO
.I2
‘4
Figure 2. Equivalent conductance of (A) HC104and (B) Bu4NCI04in tert-butyl alcohol as a function of solute concentration
which was sublimed from material of unknown origin. No attempt was made to dry the acids before use, since the amount of water added with the titrant was large compared to that added with the acid. Tetrabutylammonium hydroxide (0.16J1) was prepared by the silver oxide method as described previously (1). The aqueous base solution was not purified by ion exchange, however, and slightly more isopropyl alcohol (27%) was needed to keep the final base mixture homogeneous. Nost stock solutions of the acids were prepared by dissolving 7.0 mmoles acid in 100 ml. tert-butyl alcohol. Ten milliliters of stock was put in the conductance cell, diluted with 25 ml. pure tert-butyl alcohol, and titrated with base. Where solubility of the acid was low, the stock solution was prepared a t half concentration and twice as much taken for titration. Potentiometric, titrations of (ethylenedinitri1o)tetraacetic acid di-tetrabutylammonium salt, ( (BuJ)~H~Y)-
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3
metal mixtures with tetrabutylammonium hydroxide were carried out in a cell described previously, ( I ) . Here also, tetrabutylammonium perchlorate was used to maintain a hinh ion strength, (0.01). ( B U ~ S ) ~ Hwas ~ Y prepared by neutralization of a n aaueous slurrv of (ethylenedinitrilo) tetraacetic acid “with tetrabutylammonium hydroxide to p H 4.8. Most of the water was removed at reduced pressure. Ether was added, and then removed a t reduced pressure. The resulting solid was dissolved in acetone, and the solution was filtered. Acetone was removed from solution until the salt recrystallized. The slurry was then heated to dissolve the salt, and ether added to reprecipitate it. This material was isolated and dried in a vacuum desiccator. The salt is
0 W
0
z a
c
0 3
0
z
0 U I
u
LL
0 W
n.
v)
0
Figure 4. Conductance titration curves of aliphatic dicarboxylic acids with 0.1 6 4 M tetrabutylammonium hydroxide in tert-butyl alcohol solvent
554
ANALYTICAL CHEMISTRY
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5 6 7 8 9 IO II MI. B A S E Figure 3. Conductance titration of sulfuric acid with 0.164M tetrabutylammonium hydroxide in tert-butyl alcohol solvent
n 2
A. B. C.
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4
Oxalic acid Malonic acid Succinic acid
0
deliquescent. The purity of the ditetrabutylammonium salt, 99.5-99.9%, m.p. 152-154’ C., was determined by titration with standard calcium solution at pH 10 using Calmagite indicator. Calcium, magnesium, copper, and cadmium perchlorates (G. F. Smith Chemical Co.) were used without further purification. RESULTS AND DISCUSSION
Conductance titration curves for phenol and o-substituted phenols are shown in Figure 1. The curves are linear up to 1 eq./mole owing to a small change in equivalent conductance with concentration for the salts formed in the titration (vide infra). Salts of o,o-disubstituted phenols are dissociated to a larger extent than o-substituted phenols (as would be expected from steric considerations) and end points were not observed for the disubstituted phenols. The use of a smaller titrant cation, such as (CH3)4N+,would lead to better conductance titrations for these phenols, assuming association increased with decreasing ionic radius, as in the case of carboxylic acids (4). Polyhydric phenols, such as 1,2-dihydroxybenzene, cannot be titrated successfully in tert-butyl alcohol, owing to air oxidation during the titration. Use of a nitrogen atmosphere slowed the oxidation and allowed estimation of the end-point a t 1 eq./mole. Very sharp conductance breaks were observed in the titrations of “01, HC104,HCl, and H2SOa(at both equil-alence points). A “V” shaped curve was observed for perchloric acid, which was expected on the basis of the relatively high dissociation constant of HC104, and the change in equivalent conductance with concentration, Figure 2. Since the shape of the X eq. us. concentration curves is typical of most salts, the titration curves of monobasic
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6 MI. BASE
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Figure 5. Conductance titration curves of aliphatic dicarboxylic acids with 0.1 64M tetrabutylarnrnonium hydroxide in tert-butyl alcohol solvent A. 6.
C.
Maleic acid Tartaric acid Fumaric acid
acids a t initial concentration of 0.02M (C1’2 = 0.14) were expected to be linear over approximately 90% titration. I n the case of sulfuric acid, Figure 3, interpretation of the conductance curve is somewhat uncertain, owing to the lack of information on the conductance of unsymmetrical electrolytes. The low conductance at the start of the titration is consistent with the dissociation constant from potentiometric measurements. In general, the specific conductances of ditetrabutylammonium salts were found to be less than monotetrabutylamnionium salts. The nonlinear behavior in the region 1 eq./mole to 2 eq./mole for sulfuric acid is the result of a t least two effects. First, a common ion effect on the dissociation of the two salts formed. If the dissociation of one of the salts is extensive, then the tetrabutylamnionium ion formed will have a significant effect on the dissociation of the other salt. Second, a difference in equivalent conductance-concentration relationships for the two salts will lead to nonlinear behavior. These two effects can operate so as to yield conductance curves that are concave upward, or concave downward. Only in the case of very high dissociation of the monotetrabutylammonium salt, via intramolecular hydrogen bonding, can one reasonably expect the conductance curve to be concave downward. Conductance titrations of a number of aliphatic dicarboxylic acids are shown in Figures 4 and 5. Here also, the conductances of the ditetrabutylammonium salts, are less than that of the monotetrabutylammonium salts, and very sharp
compounds precludes the accurate measurement of end points a t 1 eq./mole. Conductance titrations thus appear to be less satisfactory than potentiometric titrations in these cases where intramolecular association is possible. Most aromatic carboxylic acids like aliphatic acids, give no indication of intermolecular bonding, except in the case of p-hydroxybenzoic acid. I n the titration of small amounts of p-hydroxybenzoic acid, a solid appeared and disappeared during the reaction of the carboxylic acid group. This clearly indicated an association complex, and attempts were made to identify it. I n the first method, one half the amount of base required to neutralize the carboxylic acid group was added and the resulting insoluble salt isolated. The filtrate was titrated with base, and the amount of w-hvdroxvbenzoic acid removed from soiutibn in both acid and salt forms calculated. Interestingly, the ratio of carboxylate ion to carboxylic acid was not 1: 1 (as would be expected for carboxylic acid-carboxylate association) but 2: 1. This suggested a trimer of the form
HOO-COOHO-COO or A
B end points a t 2 eq./mole are observed. However, malonic, succinic, and maleic acids show very high conductance a t 1 eq./mole owing to intramolecular association of the half-acids. The curvature on both sides of the maximum for these
The second method was based on the graphical extrapolation of a conductance curve, as shown in Figure 6, curves A, B, and C. While the method was less precise than the former, it yielded ratios very close to 2: 1.
Figure 6. Conductance titration curves of p-hydroxybenzoic acid with 0.083M tetrabutylarnrnoniurn hydroxide in ferf-butyl alcohol solvent VOL. 38, NO. 4, APRIL 1966
555
4 Figure 7. Conductance titration curves of substituted benzoic acids with 0.1 64M tetrabutylammonium hydroxide in tert-butyl alcohol solvent A. 6.
Benzoic acid in-Hydroxybenzoic acid C. 2,4-Dihydroxybenzoic acid
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6 MI. B A S E
Characterization of the trimeric precipitate by infrared spectrum showed no absorption peak a t 930 cm.-1, but a small absorption peak a t 1680 cm.-l This indicated that the carboxylic acid group was not free to dimerize to give a hexameric unit, but incorporated into the trimeric precipitate in a dimer-like structure. Potentiometric titration showed a distinct decrease in solution acidity after precipitation of the association complex, which indicated that the hydrogen in the precipitate was titratable but was considerably weaker than hydrogen of the carboxylic acid group. All these data are consistent with structure B.
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Conductance titrations of other aryl carboxylic acids are shown in Figures 7 and 8. The monotetrabutylammonium salts of terephthalic and trimesic acids precipitated from solution during titration, but dissolved upon conversion to the ditetrabutylammonium salts. The incomplete dissolution of the monotetrabutylammonium salt of trimesic acid a t 2 eq./mole and the inflection a t 2.35 eq./mole where dissolution is complete, suggested the formation of a
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0
MI. B A S E Figure 8. Conductance titration curves of substituted benzoic acids with 0.1 64M tetrabutylammonium hydroxide in fert-butyl alcohol solvent A. 6. C.
D.
556
o-Phthalic acid i-Phthalic acid Terephthalic acid Trimeric acid
ANALYTICAL CHEMISTRY
trimeric precipitate. This could not be confirmed by independent experiments, however. I n general, conductance titrations strongly indicate an absence of intermolecular acid-anion association in tert-butyl alcohol. The assumptions made in the determination of dissociation constants by potentiometry (6), and the interpretation of potentiometric titration curves (2) would thus appear t o be valid. Although there is a great deal known of complex formation of (ethylenedinitrilo) tetraacetic acid (EDTA) in aqueous solutions, relatively little is known about complex formation in nonaqueous solvents. Owing to the general enhancement of the complex formation of a metal in a mixed solvent system, it was of interest to examine reactions in tert-butyl alcohol. Preliminary studies using magnesium indicated very strong complex formation because strongly acidic species were produced from the reaction of magnesium ion with EDTA. As it turned out, the reactions were such
-800
I MOLES OF BASE PER M O L E
OF
2
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2.5
EDTA
Figure 9. Potentiometric titration of Mg(C104)Z(Bu4N)2H2Y mixtures with tetra butylammonium hydroxide in tert-butyl alcohol solvent A. 6.
C.
9 : 8 mole ratio Mg: (BuaNhHnY 1 : 2 mole ratio 1 : 4 mole ratio
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Figure 10. Potentiometric titration of Cu(Cl04)2(Bu4N)2H2Ymixtures with tetrabutylamrnonium hydroxide in fert-butyl alcohol solvent A. 6.
that differentiating titrations of acidic species formed from the reaction of metals with EDTA were possible. Figure 9 shows potentiometric titrations of various mixtures of magnesium and (Bu4N),HzY with tetrabutylammonium hydroxide. The titration of ( B U ~ N ) ~ H , Y mixture, Curve C, Figure 9, shows two well defined breaks. The first corresponds to two equivalents of base per mole magnesium: the difference between the first and second breaks corresponds to one equivalent of base per mole (Bu4N)2H2Y in excess of magnesium. The complex formed under these circumstances is undoubtedly ;LIgH2Y,and the reaction with base hIgH2Y 2Bu4h’OH e (Bu4N)&IgY 2H20. The salt ;\lgH,Y has limited solubility for when the total amount of magnesium was increased, (Curve B, Figure 9) a salt precipitated. The ratio Mg:Y was 1 : l by chemical analysis and the infrared spectrum was almost identical to known MgH2Y. When magnesium is in excess of ( I ~ U ~ T ’ ; ) ~ HCurve ~ Y , A, Figure 9, the solutions are initially very acidic compared to solutions of either reactant. This indicated that perchloric acid was produced in the reaction. The amount of base required to reach the first small break corresponded to 2 eq. of base/mole of magnesium in excess of ( B u ~ N ) ~ H z Y . The amount of base added from the second to the third break also corresponded to 2 eq. of base/mole of excess
+
+
-800
2 : 1 mole ratio Cu: (Bu4N)zHzY 1 :4mole ratio
I 2 MOLES OF BASE PER MOLE OF EDTA
Figure 1 1. Potentiometric titration of Ca(ClO& and Cd(C104)2-(B~4N)2H2Ymixtures witht etrabutylarnrnonium hydroxide in tert-butyl alcohol solvent A. 6. C.
magnesium. The most logical explanation for these titration data was that a second precipitate formed, in the presence of excess magnesium, with the composition (MgC104)HMgY. The chemical analysis of the precipitate that formed when the ratio of Mg:(Bu4N2)H2Y was 4:1 was Mgl.9(C104)1.~Y,with hydrogen undetermined, The infrared spectrum showed that the precipitate was distinctly different from the precipitate isolated in the presence of excess (Bu4N)2H2Y, that it contained perchlorate ion, and that it was not a mixture of Mg(C104)z and MgHzY. The formation of (MgC104)HMgYfrom MgHzY by the reaction MgHzY
+ h’Ig(C104)z s (MgC104)HMgY
+ HClO4
would also account for the acidity of solutions containing excess magnesium. The reaction is probably not complete, since the first break of the titration curve is ill-defined. The reactions that take place upon titration up to the first break are then HClOi
+ B u ~ N O H2 BuaNClOa
+ HzO
4 : 5 mole ratio Cd: ( B u I N ~ H z Y 1 :4mole ratio Ca: (6uaN)zHzY Titration of (Bu4N)zHzY only
and
+
(MgC104)HMgY BuaNOH e (MgC104) (Bu4PT)MgY
+ HzO
and from the first to the second break MgHzY
+ 2 B u ~ N O H2 (Bu4W)tMgY
+ 2H20
and from the second to the third break
+
(MgClOa) (Bu4N)AlgY 2Bu4NOH s (Bu4N)MgY B~4NC104
+ + Alg(0H)z.
Titrations of C U ( I I ) - ( B U ~ ~ Y ) ~ H ~ Y mixtures, Figure 10, indicated marked similiarities to titrations of magnesium mixtures. Two distinct precipitates were again obtained, depending upon the reagent in excess. Infrared spectra and chemical analyses showed the copper precipitates to be analogous to magnesium precipitates. It thus appeared that the reactions of copper(II), and subsequently calcium(l.1) and cadmium(I1) (Figure 11) are identical to those of magnesium, under the same conditions, with only slight differences in the extent of complex formation. VOL 38, NO. 4, APRil 1 9 6 6
557
The determination of magnesium by titration with the tetrabutylammonium salt of EDTA is very difficult, owing to the precipitation of compounds of different R.Ig:Y ratios. However, the reverse titration of ( B u ~ N ) ~ Hwith ~Y magnesium is successful, using Eriochrome Black T indicator and acetic acid-tetrabutylammonium acetate buffer. Eriochrome Black T shows the same color changes in tert-butyl alcohol as in water. (Alost common metal-
lochromic indicators are insoluble in tert-butyl alcohol.) Four replicate titrations of approximately 0.10 mmole Mg in both tert-butyl alcohol and water gave identical averages, with the standard deviation in alcohol slightly, but not significantly, less than in water, LITERATURE CITED
(1) Fritz, J. S., Marple, L. W., ANAL. 34, 796 (1962). (2) Ibid., p. 921.
cHEM.
(3) FUOSS, R. M., Kraus, C. A., J. Am. Chem. h‘OC. 55, 476 (1933). (4) Harlow, G. A., ANAL.CHEM.34, 1482 (1962). (5) Kolthoff, I. M., Bruckenstein, s., J. Am. Chem. SOC.78, l(1956). (6) w.9 Fritz, J. S.J ANAL. CHEM.35, 1223 (1963). (7) Smiley, H. M., Sears, P. G., Trans. Kentuckv Acad. Sci. 18.40 (19571. RECEIVED for review September 30, 1965. Accepted February 4, 1966. Division of Analytical Chemistry, 149th Meeting, ACS, Detroit, Mich., April 1965.
Estimation of Medium Effects for Single Ions and Their Role in the Interpretation of Nonaqueous pH OREST POPOVYCH Analytical Research Division, Esso Research & Engineering Co., linden, N. 1.) and Department o f Chemistry, Brooklyn College of the City University o f New York, Brooklyn, N. Y. A new method is proposed for the estimation of medium effects for single ions in nonaqueous solvents. It is based on the assumption that the medium effect of an electrolyte composed of large ions of equal size and solvation (a “reference electrolyie”), can b e apportioned equally between the anion and the cation. Triisoamylbutylammonium teiraphenylboride (TAB BPhd) was chosen as the “reference electrolyte” and its medium effects as well as those of several other electrolytes in methanol and ASTM solvents were determined from the solubility products and Eo’s in water and Ihe two solvents. ASTM solvent consists of 50.0% toluene, 49.5% isopropanol, 0.5% water, by volume, and is used in ASTM titrations of acids and bases ( I ) . Medium i calcueffects for single ions m ~ were lated on the assumption the log mYTAB = log m Y B P h 4 ‘/Z log my+z TAB BPhp. Values for log ,,,YE were used to interpret pH data and to estimate liquid-junction POtentials in the two media.
and n r ~ i Jthe medium effect for the solute, which becomes unity only a t infinite dilution in water. The medium effect is a measure of the difference between the free energy of the solute in its aqueous (,Gio) and nonaqueous (,Gio) standard states: aGio - ,,,Gi0 = RT In ,,,Ti (2) The effect of medium on p H and other acidity functions has been treated exhaustively in the monographs by Bates ( 2 ) and Izmaylov (17). For purposes of discussion, however, we will summarize those relationships which involve medium effects for single ions in the interpretation of electrometric p H measurements in nonaqueous solvents (subsequently referred to as LLsolvents’J and designated by subscript s). I n any amphiprotic solvent we can derive the value of pa=* = - log a ~ H )from the e.m.f. of hydrogen(VZH silver chloride cells without transference, provided we know the standard potential of the cell ,Eo in the same solvent : E - ,Eo P&H* = ___ k log mci log a ~ (3) ~ l In Equation 3, k = (RT In 1 0 ) / F and the acitivity coefficient of the chloride ion a ~ c lcan be estimated from a suitable form of the Debye-Huckel law. The above relationship has been used to establish p H reference standards in water and in some solvents ( 2 ) . For solutions of pure acid HA a t a n analytical molality mH& the same result can be obtained from a conductometric determination of its degree of dissociation a: paH* = - log HA a a ~ ~ ) . The paH* scale is different for each solvent, so that equal numerical values +
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ix PARTIALLY and totally nonaqueous solutions, the activity of a solute i can be referred either to infinite dilution in the given solvent, or in water as its standard state. The corresponding solute activities on the nonaqueous (a,*) and aqueous (a,) scales are given by a,* =
m,
(la)
and (1b) a, = mi B y 2m ~ = $ ar*mY, where m is molality, the familiar activity coefficient which becomes unity a t zero molality in the given solvent, 558
ANALYTICAL CHEMISTRY
+
of paH* in different media do not represent equal acidities. For example, we do not know off hand whether a solution of paH”2 in an alcohol is more or less acidic than a solution of paH 1 in water. I n order to express p H measurements in different solvents on a single scale, it is most convenient to refer activities in all solvents to their aqueous standard states. According to Equation l b , paa* (nonaqueous standard state) and paH (aqueous standard state) are related through the medium effect for the hydrogen ion,yH: paa = pas*
- log
(4)
Similarly, if the aqueous standard potential were used in Equation 3, paH could be obtained directly from cell e.m.f. E measured in the solvent, provided the medium effect for the chloride ion ,,,YCI were known (3) :
E - ,Eo paa = ____ + l o g log
*YCl
mci
+
+ log
mYCl
(5)
hIost practical p H measurements are made in e.m.f. cells with transference, using electrodes standardized against aqueous p H buffers. The resulting “operational pH” readings in any solvent deviate from paH by the difference in the liquid-junction potentials in the aqueous buffer and the unknown, E 3: pH
- pa= = Ej/k
(6)
Consequently, the relationship between p H and paH* (the latter is useful as a n approximation of mH in dilute solutions) involves both the medium effect and the liquid-junction term (3,30): pH
- pa=* = E 3 / k - log
m~~
(7)
