on Analytical Chemistry, Los $ngeles, June 16, 1956. (2) Brummet, B. D., Hollweg, R. M., ANAL.CHEM.28, 448 (1956). (3) Gerhardt, P. B., Hartmann, E. R., Thid..I -29. 1223 - I ---- (lR.571.
\ - - - . / .
(4) Godd;, R. F., Hume, D. N., Zbid., 22, 1314 (1950); 26, 1740 (1954). (5) Irving, H. &I., Williams, R. J. P.,
Nature 162, 746 (1948); Analyst 77, 813 (1952); J . Chern. SOC.1953, 3192. (6) Pflaum, R. T., Popov, 4 . I., Goodspeed, X., . ~ ? ; A L . CHEX 27, 253 (\ -19.551. ---,-
(7) Pilipenko, A . T., Zhur. Anal. Khim. 8 , 286 (1953). (8) Riddick, J. A . , A l v . 4 ~ . CHEX 24, 41
(1952); 26, 77 (1954); 28, 679 (1956). (9) Vallee, B. L., Zbid.,26, 914 (1954).
RECEIVEDfor review June 10, 1957. ilccented Februarv ~" 3. - , 1958. Division of Refig-ing, 22nd Meeting, American Petroleum Institute, Philadelphia, Pa., May ~
1957.
Improved Conductometric Titration of Weak Bases W. H. McCURDY, Jr., and JOHN GALT' h i c k Chemical laboratory, Princeton University, Princeton, N. 1. )A new solvent system composed of a 1 to 1 mole ratio of 1,4-dioxane and formic acid has been developed for conductometric titration of weak bases. Routine titration of 10-mg. samples is readily attained if the base is sufficiently strong to react with the solvent to some extent. Equilibria studies have shown the factors responsible for the improved conductometric end points observed. Direct titration of mixtures of weak bases can b e realized if a difference of 1 ~ K B H +unit exists. Even smaller differences in basic strength can be resolved in favorable circumstances. Compounds that yield N-formyl esters with this solvent system a t room temperature may b e successfully titrated a t 0' C. in most cases.
A
number of nonaqueous titrations of weak bases using high-frequency conductometric apparatus have been reported. Glacial acetic acid was generally employed (8, 11, 12, 21), although a mixed solvent of benzene, methanol, and acetic acid was tried (14). The results check satisfactorily with potentiometric and acid-base indicator determinations in spite of the fact that the conductometric end points are generally located at the intersection of two ascending lines. The slope of the titration line before the end point beconies more nearly identical with the slope of the excess acid line as the basic strength of the sample decreases and the end point becomes increasingly difficult to locate accurately. Many workers have shown the improvement in conduetometric end point which occurs when a strong acid is titrated with a weak base (5, 6, 9), but such a procedure eliminates the possibility of determining mixtures of weak bases. Considerable improvement CONSIDERABLE
1 Present address, Harvard Medical School, Cambridge, Mass.
940
ANALYTICAL CHEMISTRY
in potentiometric end points may be realized by changing from glacial acetic acid to nitromethane-formic acid (20), trifluoroacetic acid (S), 4methyl-2-butanone ( 2 ) )or other special solrents. This study was initiated to discover solvents or solvent mixtures which would alter the slopes of the titration and acid lines in such a manner as to increase the sharpness of nonaqueous conductometric end points. INVESTIGATION
OF
TITRATION SOLVENTS
Preliminary u-ork was carried out titrating salts of weak acids with perchloric acid in a variety of simple and mixed solvents. All of these salts and most of the solvents n-ere reagent grade and were not purified further. Solutions of perchloric acid were standardized potentiometrically by titration of samples of potassium acid phthalate in the appropriate solvent. Conductance measurements were performed at 10 megacycles per second, using a General Radio Tn-in-T impedance bridge (1'7) and a t 1000 cycles per second with a Serfass conductivity bridge Model RCN-15. A concentric electrode titration cell with very thin glass walls was employed with the highfrequency instrument and a dipping type cell having a cell constant of 0.100 cni.-l was used for low frequency conductivity measurements. Similar results were obtained n ith either instrument, although the sensitivity of the high-frequency measurements \vas somewhat superior, because the high-frequency cell electrodes were 25 sq. em. compared to 1.0 sq. cm. for the dipping cell and because the high-frequency response is enhanced by decreasing the solvent dielectric constant (1'7). The cell constant for the high-frequency cell was 0.025 cm.-l in glacial acetic acid. Temperature control was not carefully maintained in the early experiments but volume corrections were applied to all titration curves. Effect of Solvent. Five-milliliter portions of 0 . W potassium acetate dissolved in t h e appropriate solvent
nere titrated with 0.1F perchloric acid in the same solvent. Dilution of samples t o 160 nil. with solvent n n s held constant in all experiments. With solvents of lorn dielectric constant, 5 to 20% by volume Tvater !vas added to increase the solubility of the sample and titration products. Table I, A, lists seven solvents in order of decreasing utility in this titration. The curves \\-ere plotted in a standard manner in order that slopes of the titration lines before the end point (slope B.E.P.) and excess acid lines after the end point might be compared. The ratio of slope of acid line to slope of titration line and the end point angle were calculated in each case. Ethylene glycol and glycerol were not considered promising because of the large solrent viscosity. Results with nitromethane are not included because no end point was observed in this solvent. Effect of M i x e d Solvents. A number of mixed solvents Lvere prepared by adding 10% by volume of acidic reagent to ethylene glycol nionomethyl ether and 1,4-dioxane-20% ivater. ,4 few typical results are given in Table I, B, for the titration of 0.5 nieq. of potassiuni acetate when 10% by volume of acetic, formic. or monochloroacetic acid was present. I n all cases the starting volume was 160 nil. Acetic acid and nionochloroacetic acid solvent mixtures show only a slight effect on the slope ratio compared to the formic acid mixtures. Experiments with other acidic reagents such as dichloroacetic and anhydrous hydrochloric acid solvent mixtures did not yield any significant iniprovenient in 4ope ratio. A series of formic acid solrent mixtures was studied and many different bases titrated. d fen- examples are presented in Table I1 for solrent niixtures containing 10, 20, and 30% (by volume) 99% formic acid in purified and 14% (by volume) 1,4-dioxane 99% formic acid in glacial acetic acid. The increase in sharpness of conducto-
(e),
metric end points produced by formic acid is more readily seen in Figure 1 for the high-frequency conductometric titration of potassium formate. This effect does not occur with potentiometric titrations of weak bases-for example, the titration of 0.5 meq. of potassium acetate in 1,4-dioxane-l0'% water (Figure 2) is not improved by addition of formic acid. The increased acidity of the formic acid solvent considerably decreases the magnitude of the end point potential break. Similar results have been found previously with potentiometric titrations in anhydrous formic acid (19). The conductometric or potentiometric titer volume is increased slightly as larger amounts of formic acid are added to 1,4-dioxane. Titration Mechanism. The improvement in conductometric titration of weak bases occurs primarily because the slope of the titration line before the end point is decreased (Figure 1). This is the result of a n equilibrium reaction between the weak base and the formic acid solvent. The slope of the acid line after the end point is increased somewhat, because of the increase in dielectric constant which takes place when formic acid is added to the solvent. A brief study of solvent acidity and attending equilibria with several weak hases mas undertaken to explain the phenomena. SOLVENT ACIDITY. Thirty-milliliter samples of purified 1,4-diosane were titrated with 99% formic acid and with a 0.5F perchloric acid solution in 1,4dioxane. The conductometric titrations indicate t h a t 1,4-dioxane is a weak base. I n the case of the formic acid titration (Figure 3, A ) the conductance remains very small until the concentration of acid approaches 1 mole of acid per mole of 1,4-dioxane. Conductance measurements of formic acid solutions in tetrahydrofuran were also very IOK. This effect was not observed upon titration of 30 ml. of glacial acetic acid with 99% formic acid (Figure 3,B). Although part of the curvature in these titrations is caused by the increase in dielectric constant which occurs during the titration, the evidence is in favor of the following equilibria between 1,4-dioxane (1,4-DO) and formic acid (HFm): 1,4-D0
+ HFm
K1'
h',
1,4-DO.HFm F* 1,4-DO.H+ Fm- (1)
+
From the conductance measurements i t may be assumed that IC1 >> K1'. The small solvent blank found in the perchloric acid titration of 1,4-dioxane (Figure 3,C) represents the same type of equilibria given by Reaction 1, except the concentration of undissociated complex is small-Le., K2' >> K z .
Table I. Effect of Solvent (Titration of 0.5 meq. of potassium acetate with 0.1F perchloric acid) Slope B.E.P., Slope End Point Micromhos/Ml. Titrant Ratio Angle Solvent 7.9 139' A. Acetic acid (957,) 0.93 5.4 147' 1.08 1,PDioxane (goy0) 3.6 151" 6.90 Ethyl alcohol (957,) 1.8 166' 8.54 Ethylene glycol monoethyl ether 17.9 168" 0.08 Glycerol 9.8 1.24 129' Ethylene glycol 1.9 14.5 169" Ethylene glycol monomethyl ether B. 1,4-Dioxane 4.1 150" 7.42 +20% HsO 4.1 147" 6.02 10% CH3COOH 132" 7.1 + 10% HCOOH 3.23 143O 3.9 3.58 10% ClCHzCOOH Ethylene glycol monomethyl ether 159" 8.82 2.6 10% CHICOOH 129" 2.91 10%HCOOH 8.1 151° 2.7 6.12 10% ClCHzCOOH
+ + ++ +
0 00
10
?. " '
..^ -
4
%
-
* I
H
1
Fm
z
d
10 20 ML. 0.1000F H CLO,
Figure 1. Effect of formic acid on high-frequency titration of potassium formate
Table II.
Effect of W e a k Base
(Titration of 0.2 meq. of base with 0.1F perchloric acid)
ML.
G1y cin e 1,PDioxane +20% Hz0 +3070 HCOOH Glacial acetic acid +14%HCOOH 8-Quinolinol 1,4-Dioxane 20 % HzO
+
+30% HCOOH Glacial acetic acid 1470 HCOOH
+
-0.025 - 1.67 -6.69
174' 140' 82'
+P1.85
174' 97" 157" 110'
-3.68 +1.92 +0.83
No end point
observed
-6.34 +3.50 -3.60
81" 156' 84'
HCLOe
Figure 2. Effect of formic acid on potentiometric titration of potassium acetate
Solvent. 0 907, 1,4-dioxane-l0% water 807,. 1,4-dio~ane-107~water-lO% formic acid
1,4-I)0
Potassium formate 1,4-Dioxane + l o % HCOOH +20% HCOOH +307oHCOOH
0.1000 F
K2 + H+ClOa- e 1,4-DO.HC104 Kz e1,4-DO.H++ Clod- (2)
This solvent blank may be eliminated (Figure 3,D and E ) for perchloric acid titrations of solvents containing 1 and 1.1 moles of formic acid per mole of 1,Pdioxane. I n these cases an acidic solvent of moderately low dielectric constant has been prepared from a basic one. SOLVENTEQUILIBRIA. A group Of experiments was performed to study reaction equilibria between weak bases and various formic acid solvent mixtures. Bases mere selected with absorption spectra in the near ultraviolet which are not subject to interfering absorption from the solvents under consideration. The purpose of the first step was t o compare the dissociation constant of a VOL. 30, NO. 5, MAY 1958
941
4.01
!
18)
6 1
?
101
I? o
z 4
0 "1
"
'
0 4
'
:
Figure 4.
_.,, .___- ./., 05 MOLES FORMIC
OS
'
1I O
HCLO,
Acid dissociation constant
of 8-quinolinol in 1,4-dioxane
/*
,
"
0'
MOLES
5 w
'
*
I
p
.
,
IO
ACID/MOLE
15 SOLVENT
I
32 3 4 % H f m
30
100IC)
+I
J
eo
-
m
100-
-
$? 6 0 1
w
1 , ----r , e , . , ',
10
0
MOLES
Figure 3.
H CL 0 4 / M O L E
I 40-
.-I----._.---*
,
30 x
20
1
-
J
0
IO-^
20
-
1,4-DIOXANE
Basicity of 1,4-dioxane
/--e
1
02
0
Conductometric titration curves A . 1,GDioxane with 99% formic acid B . Glacial acetic acid with 99% formic acid C. 1,eDioxane with 0.5F perchloric acid D. 1,4-Dioxane-32'% formic acid with 0.58' perchloric acid E . lj4-Dioxane-34q formic acid with 0.5F perchloric acid
MOLES
1
1
/*
,
1
06
04 FORMIC
1
1
08
ACID/MOLE
1
1
1
IO
,
12
I.4-DIOXANE
Figure 5. Acid-base equilibria in 1,4-dioxaneformic acid solvent mixtures
Basic compound.
0 8-Quinolinol 8 m-Nitroaniline 0 Anthranilic acid
ciation constants in aqueous and nonaqueous solvents have been discovered with other indicator bases (16). Equilibrium constant determinations for the system 1,4-dioxane-formic acidweak base were performed by adding perchloric acid] us. total concentration known amounts of 8.238F formic acid of perchloric acid in Figure 4. BH+ and in 1,4-dioxane to 1.003 X 10-3F 8B refer t o the protonated and unprotoquinolinol in 1,4-dioxane. Also several nated forms of the base. Least squares measurements were made on 1,4treatment of the data yields an interdioxane-formic acid solutions containcept value of 4.96 + 0.04 for $BH+ ing 9.988 X 10m4Fm-nitroaniline and of 8-quinolinol in 1,4dioxane. This 5.026 X 10-4F anthranilic acid. The is in close agreement with the values data are presented in Figure 5 as a graph of mole per cent BH+ us. the 4.91 and 5.01 determined in aqueous formic acid-l14-dioxane mole ratio. buffers (15). Nearly identical dissoThe S-shaped curve observed with 8quinolinol is caused by competition between 1,4-dioxane and the weak base for formic acid. Assuming that Table 111, Equilibrium Constants in 1,4-Dioxane-Forrnic Acid formic acid is entirely complexed when Excess 1,Cdioxane is in excess of the acid HFm, 1,4DO, concentration, the following equilibrium [1,4-D0.HFm1 Moles/ Moles/ Absorb[BHfl Liter Liter [1,4-D01 K B H + F ~ - reaction may be formulated: [BI Base ance BH+ Fm1,4-DO K B= H+F~0,1627 2.51 0.0649 9.882 1.608 0.080 8-Qni1iolinol 0.4283 3.33 7.503 0 , iiii 3.217 0.149 B + 1,4-DO.HFm (3) 0.9410 1.55 5.126 4.826 0.6083 0.497 1.410 1.01 3.922 1.433 5.648 0,774 Equilibrium constants at 25' C. 21.79 1.11 0.369 8.043 1.250 19.53 1.314 at 370 mp Full acid form based on Reaction 3 are given in Table 1.033 9.01 0.1150 9.895 5.028 m-Xitroaniline 1.156 111. The concentration of excess 1,1Full basic form 1.289 at 370 mk dioxane listed in Table I11 was cal0.3444 43.5 9,079 6.753 0 06809 Anthranilic acid 1.498 culated from [moles of 1,4-DO added Full basic form 1,600 at 310 mp - (moles of HFm added - moles BHT
conjugate acid in l,&dioxane with that in water. The moderately strong base, 8-quinolinol, was selected because accurate pKBH+ values have been determined in water (15). This compound is also of interest because the conductometric titration curves are strongly modified by formic acid solvents (Table 11). Solutions of 1.003 X 10-3F 8-quinoliiiol in 1,Cdioxane containing known sinounts of 0.054798 perchloric acid w y e prepared. Measurement of the absorption a t 370 niM of these solutions and the full acid form of the base allowed calculation of pKBm. 811 spec-
trophotometric measurements were performed in 0.996-cm. quartz cells with a Warren Spectracord. The results of these experiments are shown as a graph of log - log [unreacted
[TI
+
942
a
ANALYTICAL CHEMISTRY
formed)]. The decrease in K B H + F ~ values for 8-quinolinol with increasing formic acid may be attributed to increased dissociation of the formic acid in solvent mixtures of higher dielectric constant than 1,4-dioxane. Although the accuracy of these measurements is not great, the magnitude of the equilibrium constant is roughly proportional to KBH+ for the three bases studied. It was of great interest to the present work to examine the extent of reaction of 8-quinolinol, rn-nitroaniline, and anthranilic acid with several acidic sokent systems. Data from the dotted vertical lines in Figure 5 as \\-ell as some values for tIvo acetic acid systems are summarized in Table IV. These results indicate that bases having ~ K B Hvalues + less than 2.0 do not react mith the solvents listed in Table IV to any appreciable extent. On the other hand, bases stronger than PKBH+ = 5.0 react completely with formic acid solvent mixtures. Differentiation of basic strength should be possible between these limits. The slope of the conductometric titration line before the end point (slope B.E.P.) will vary with the degree of solvolysis if one or more products of this reaction have large equivalent conductances. Conductometric titrations of bases in glacial acetic acid do not yield very large differences in slope R.E.P. because the equivalent conductance of acetate is rather small. The fact that formate ion has a n abnormally large equivalent conductance
Table IV.
Solvent 14-T)inunne -i - -
+30% HFm +32y0 HFm +34% HFm
Glacial acetic acid HFm p H ~ ain + water a
Solvolysis
8-Quinolinol Slope % [BH+] B.E.P.a 94.0 97.5 99.5 69.0 99.2 4.91
of W e a k Bases
m-Nitroaniline Slope % [BH+] B.E.P.a 25.0 30.5 37.0 20.4 45.4 2.45
-6:3 -6.5 $3.5 -3.9 , . .
Anthranilic Acid Slope 70 [BH+] B.E.P.O 12.0 16.8 22.5 4.1 27.7 2.08
... ...
17.7 +0.98 f9.i
..
... ...
711.2 t3.8 ~ 2 1 . 1 ...
Micromhos per ml. of titrant.
B
in formic acid w s proven by the early work of Schlessinger and Bunting (18). The effect of formic acid on the conductometric titration of potassium formate (Figure 1) suggests that a complex between the anion and excess acid may also be involved ( I S ) . The conductance of the formic acid solvent complex is small. Therefore, the slope of the titration line may be controlled by the change in conductance which occurs when the formate ion in equilibrium Kith the basic sample is reacted with perchloric acid. Two limiting cases may be considered : PICBE+
BH+C104-
+ 1,-1-D@
5)
The equilibrium constant for Reaction
4 may be estimated from
which is 3.7 X lo4 for 8-quinolinol in 1 , 4 - d i o ~ a n e - 3 0 ~formic ~ acid. The large value of this constant indicates complete reaction with perchloric acid. The Ktitration equilibrium constant for Reaction 5 is simply ~ / K B H + . Bases of intermediate strength involve the titration of both BH+ F m - and B with perchloric acid. The three bases selected for spectrophotometric study yield the conductometric curves shown in Figure 6 upon titration with perchloric acid in l,4-dioxane-32yo formic acid. The slope B.E.P. obtained by conductometric titration of equivalent amounts of these bases in several acidic solvent mixtures is also summarized in Table
> 5.
+ 1,4-DO.H+C10*-e BH+C104- + 1,4-DO.HFm (4)
BH'Fm-
pKBa,
+ 1,4-DO.H+ClOd- e
