POTENTIOMETRIC TITRATION OF DIBASIC ACIDS
May, 1942
[CONTRIBUTION PROM THE
DEPARTMENT OF CHEMISTRY O F THE UNIVERSITY
1153
O F DELAWARE]
Potentiometric Titration of Dibasic Acids in Dioxane-Water Mixtures BY RICHARD H. GALEAND CECILc. LYNCH In a paper published in 1938 Lynch and LaMer' showed that the theory of acid-base titration in dioxane-water mixtures for weak monobasic acids leads to substantially the same relations obtained by Auerbach and Smolczyk2 and by Soderback3 for titration in aqueous solutions. In this article we have extended the theory to include the titration of weak dibasic acids in dioxane-water mixtures. Titration data on oxalic, malonic, succinic, and glutaric acids are presented in support of the theory. Theory of Acid-Base Titrations (Dibasic Acids) .-From Auerbach and Smolczyk2 we have the relation between the KI and K2 for a weak dibasic acid
univalent ions, association to neutral pairs occurs for dielectric strengths above 10; for univalent and divalent ions together in such media association to neutral pairs and neutral triplets only may be assumed. We shall consider the influence that such ion association has on acid-base titration calculations for dibasic acids. Consider a solution containing dioxane, water, lithium sulfate of relatively high concentration, a weak dibasic acid, being titrated with a strong base MOH, as sodium hydroxide. We are concerned with the following (a)
HzA J_ H+
ib)
(HA)-
(d)
MpA
+ (HA)-;
H'
+ A';
2Mf
+ A-;
K't = (mH+mt,:HA-)
Kc =
(mPM+
mA-
nz M ~ A
(1)
This is essentially the same as the relation obtained for dibasic acids by Soderback3 Electroneutrality demands
+
+
where B = ( m x / m ) ( m ~ + / c [I ) (x/vO)]if the dilution is such that the activity coefficients may be taken as unity; where x c
= volume of base used = initial concentration
dJ = initial volume of acid solution fi = activity coefficient of the univalent ions fi = activity coefficient of the divalent ions m = molarity of base w0 = initial weight of solvent m0 = initial molality of acid
(g) m y +
+
mH+
+
= mHA-
+
2mA-
f
mH80r-
+
2mSOr
If m0is the initial acid molality, then (h)
mHA
f
mMHA
+
mLiHA
f
mMiA
++ + mLi2A
mHA-
%A-
=
??to
Since (mx/w0) is the stoichiometric molality the base would have at any time if no reaction occurred in the solution, we obtain a relation (i) m M +
+
mMHA
+
2mMgA
f
mMH80r
f
2mM2804 =
(mx/wo)
In a medium where the ionic environment is kept where the m's refer to the molalities of the subsensibly constant, Eq. (2) reduces to Eq. (1) stances in subscripts (m is the concentration of where K , = Ky/fq and K2 = K:/f2. This con- the base in moles/1000 g. of solution; x is grams dition can be achieved by maintaining a rela- of base solution added a t any time; wo is the tively high concentration of inert neutral salt initial weight of solvent taken). M is the metal which furnishes both univalent and divalent ions ion of the base MOH. We shall make the following assumptions to to the solution. simplify the problem : (j) Bjerrum associations Fuoss and Kraus4 have shown from conoccur in the solvent to the extent of neutral pair ductivity studies that ion association in the Bjerand neutral triplet formation only. (k) Sufficient rum sense occurs in dioxane-water mixtures; for lithium sulfate has been added to maintain a (1) Lynch and LaMer, THISJOURNAL, 60, 1252 (1938). (2) Auerbach and Smolczyk, 2. physik. Chcm., 110, 66 (1924). constant ionic environment for both mono- and (3) Sderbpck, Arkiu. Kcmi. M i n . och. Gcol., l l A , 1 (1934). &-valent i ~ n s . (1) The weight of base solution (4) FUOM and IKWB, THXB JDEBNAG, 81, a397 (I%%?),
RICHARDH. GALEAND CECILC. LYNCH
1154
added to the end-point does not appreciably change the total solution weight. (m) For MHSO4 and LiHA the same extent of association; for hI2SO4 and LiZA the same extent of association. Such assumptions would appear reasonable for salts of same ionic type in the same ionic environment.
+
also also
~ L ~ H A 3m1,igA
mLl*
=
mEBo,-
=
~
U
+ 2msor-
H
+
2niA"
- f?tH+ + m W H A f
Equating through mA- in (A) and (B), and collecting all terms of (KIK2) and (KJ
+
S~ ~ W L~M ~ S(in') O ~
--
(m")
Combining (g), (i), (m'), and (m") tl?IlA-
Vol. 64
2mMd
+
m L % H Af
2rnLI2A =
(7nx:Ze'O)
Substituting for mMHA from (f) ; for m M , A from (d) ; for mLiHA from (e) ; for mLi,A from (c) ; for mM + from (g) ; and for f f t - ~from ~ (b)
Now, mH+ is negligible compared with 2mAand also as compared with (mx/wo); hence
r
1
2m0 (1
+ K')
= 0
where K = (mLi+/K6) and K' = (mZLi+/K3). In case of a great extent of ionization of salts, the values of K and K' approach zero, and the Eq. (1) is obtained. Multiplying both sides of the above equation by (1 K"),we obtain
+
or
13I
+
+
where KT = KI(l K ) and K,* = K2 [(l K')/(l K)]. Thus, we obtain an equation identical in form with Eq. (1) in which, however, the constants, K: and Ka, differ from the classical constants, defined in Eq. (A) and (R), by a constant factor. The solution of the problem of determining KT and K,* is through plotting of C; as abscissa and 4 as ordinate, where
+
When m3A- and m2A- are small compared with mA-, a condition true for weak acids a t low concentrations, only the first term on the left of the equation is significant; hence
's
(At
Combining (a), (b), (h), (m'), and (m"), where
we obtain the relation
VigIliA
+
7% M2A
+
WZLiHA
+
I:TLi:A
=
Wo
Similar substitution as used above leads to the expression m.4- =
r
1
+
+
for the equation KTKf KT C; 7 = 0, is that of a straight line; 17 = - KT C; -K:K,*, where ( - KF) = m (slope); ( - K:K,*) = b (intercept on 7 ) ; and ( - K : ) = a (intercept on 5). The plots obtained from titration of dibasic acids in dioxane-water mixtures should give support for this modified theory of titration. Figure 3 shows results obtained for the acids investigated. The Reference E. m. f.-In order to apply the general theory to a titration, it is necessary to know the relation between e. m. f. of the cell measured and the activity of the hydrogen ion, or far constant ionic environment, the molality
POTENTIOMETRIC TITRATION OF D'IBASIC ACIDS
May, 1942
of the hydrogen ion. If we impose the latter conditions during titration, the relation can be simply established in a manner similar to that described by Lynch and LaMer. Thus, for the cell described in the experimental part for obtaining Eovalues E298 =f Eo - 0.05915 log mH+ (4) and (E298
+ 0.05915 log
WZH&
= E"
- 0.05915 log (mHB0,) -??!E-
at W Z H ~ S O= ~ 0, we set log (mH+/mH,S04) = 0 and E = ( E N ~ O.OS915 log VZH~SO,). This limiting value may be obtained by extrapolation, and the Eq. (4) employed for calculation of the hydrogen ion molality from e. m. f. measurements.
+
1155
carbonate-free sodium hydroxide solution in aqueous solution. All were chemically pure within 0.1%. Lithium Sulfate.-Merck c. P. Lithium Sulfate was used without further purification. Aqueous solutions of this salt had a PH of 6.0. Solutions.-All solutions were prepared by direct weighing, and all titrations were made by weight. Thus, the concentrations are expressed in weight per cent. or by molality. I n order t o prepare the dioxane-water-lithium sulfate solutions, aqueous solutions of lithium sulfate were prepared and analyzed for sulfate. The aqueous lithium sulfate was then mixed with dioxane to the proper ratio. This was also the method of obtaining dioxane-watersulfuric acid solutions which were added to the dioxanewater-lithium sulfate solutions to give the sulfuric acid solutions used in determining the EO values for the cells. The other acid solutions were prepared by adding the acids directly to the dioxane-water-lithium sulfate solutions.
'I/I 0.04
0
0.02
0.04
0.06
v5KiG. Fig. 1.-Extrapolation plots for EO values of the cell: 1, dioxane/water ratio 1.002, m ~ f i o ,= 0.3002, EQ = -0.0344 v.; 2, dioxane/water ratio 1.856, mL1ao, = 0.05263, EO = -0.0347 V.
Experimental Potentiometric titrations were made at 25' in two dioxane-water mixtures in which the ratios were 1.002 and 1.856, respectively, using oxalic, malonic, succinic, and glutaric acids as examples of the series of weak dibasic organic acids. As in the case of lithium chloride, large additions of lithium sulfate produce immiscibility in the dioxane-water system, which sets an upper limit to the permissible concentration of lithium sulfate. A previous study by Lynch and LaMer' of the system sodium hydroxide-dioxane-water a t 25 ' showed that small amounts of sodium hydroxide produced immiscibility. Thus it was necessary to add approximately 1.6 molal caustic in aqueous solution in the form of the titrating base from a weighing pipet to keep the dilution factor under one per cent. Preparation of Materials Dioxane.-Technical 1,4-dioxane from the Eastman Kodak Company was purified by the method described by Eigenbergerb with slight modifications. The product was kept over metallic sodium, from which it was distilled from an all glass apparatus, as needed. Acids.-Reagent grade acids from the Eastman Kodak Company were recrystallized and analyzed for purity using ( 5 ) Eigenberger, J . pro&l. Chcn.. 180, 75 (1931).
0.5 1.o 1.5 2.0 Moles of NaOH per mole of acid. Fig. 2.-Titration curves for the acids: Dioxane-water Curve 1, Glutaric; Curve 2, Succinic ratio = 1.856 Curve 3, Malonic; Curve 4, Oxalic Molality of Li2SOI = 0.05263 Dioxane-water Curve 5, Glutaric; Curve 6, Succinic Curve 7, Malonic; Curve 8,Oxalic 0
1
= 0.3002
Apparatus.-The bolically as
titration cell is represented sym-
RICHARD H. GALEAND CECILC. LYNCH
i 156
VALUESOF K f
TABLE 1 K,* FOR
AND
THE
Acid
lMolalltr (inn) acid
IXoxanewater ratio
Dielectric constant
Molality Lips04
Oxalic Malonic Succinic Glutaric Oxalic Malonic Succinic Glutaric
0.01010 . 00i825 .01020 ,007056 00ii16 .0102ti ,007358 ,008334
1.002 I . 002 1.002 1.002 1.856 1.856 1.856
34.3 34.3 34.3 34.3 31.8 21.8 21.8 21.8
0.3002 ,3002 .3002 ,3002 .05263 ,05263 ,05263 05263
1.856
Vol. 64
ACIDSAT 25' K:/K~
K:
k': ,.
. . . ...
..... .
1.92 X 1 V 3 2.10 X 1W4 8.70 x 2.66 x 2.60 X 10-3 1.09 x 10-4 4.40 x 10-6
1.60 X 2.60 X 3.42 X :+.05 x 4.20 x i.54 X 1.13 X
I
.
...
,
120 8.1 2,5 8.7
10-j 10-4 10-6 lWfi
ti20 14..i 3.9
lo-"
reproducible to 1 0 . 1 mv. over a period of four days Measurements were reproducible to within *0.3 mv.
the solvent, one of the dioxane-water mixtures, extends throughout the cell, This cell avoids liquid junctions by making contact in the acid solution through a lithium sulfate-dioxane-water bridge connected by a ground glass joint. The bridge, by means of the glass joints, also avoids mixing and loss by diffusion during the titration. All measurements werc made with the cell in a water therniostat a t 25.0 * 0.1 '
Experimental Data and Results
Eo Of the were determined by ~~leasuring a series of cells of the type Hg,SO?, Hg Au, quinhydrone, (LizSoz'm3), H2S04(ml)
in which the solvent was dioxane-water. Linear extrapolation of the function (E298 0.05915 ____ --2 log mH2& against the d m H p S O to 4 zero concentration of the acid gave the Ea value. Fig--4 ure 1 gives the data and Eo values for the two dioxane-water ratios chosen. X E The titration data for the four acids are as shown in the curves of Fig. 2. It was observed that the cell e. in.f. drifted when the end-point --8 was overreached. Accordingly, only data to the neutralization point are included. - 4 - 2 o 2 I ii -in o 10 20 30 From the titration data, the and q values in E x 10: :x 106 the neighborhood of half neutralization were obOxalic acid Malonic acid tained for each acid in each concentration of dioxane. The plots of these functions are given in 4 x 108 5 x 10'. - 5 0 5 10 1.5 20 -15 - 5 0 5 in li Fig. 3, and the values of the constants as calculated from the data are collected in Table 1. Values in the range 2575% neutralization were -4 used in the extrapolations. The dielectric constants are from the data of Akerlof and Shorth x I t should be noted that the dioxane-water ratios and the molalities are initial values, and -io that during titration small changes occur in all these values. -2-10 1 L 3 4 -3-2-10 1 L s 4 The straight lines are evidence that the rela5 x 10s t x 105 tion (Eq. 3) is satisfied for titration of dibasic Glutaric acid. Succinic acid. acids in such media. Only the oxalic acid in I. ig. 3 -Graphical method for determining dissociation approximately 50y0 dioxane showed variance constants a t 25". from the straight line plot. This may be exE m. f. measurements were made with a Leeds and plained by the assumption that the oxalic acid Korthrup Type K potentiometer, and a Type P galvanome- first dissociation is that of a relatively strong acid. 0
+
' ie
ter, sensitive to 0.00072 microampere per scale unit. The Weston standard cell was checked frequently against other standard cells. The mercurous sulfate electrodei in the dioxane-water-lithium sulfate sohtion werc found
Summary 1. J, I
The Xuerbach and Smolczyk treatment of Akerioi .,nd i h o r t THISJ O L I R N ~ L68, . 1241 (1936)
1,1,~-BIFLUOROPROPANE 1157
May, 1942
acid-base titration for weak dibasic acids has been extended to include the influence of ion-association of salts in low dielectric media. An equation indicating the effect of ion-association on the acid dissociation was obtained. 2. Potentiometric titrations of the dibasic acids : oxalic, malonic, succinic, and glutaric, have been performed in dioxane-water mixtures containing 50 and 657, dioxane, with a quinhy-
[CONTRIBUTION FROM
THE
drone-mercurous sulfate electrode chain. Relatively high concentration of lithium sulfate was employed for keeping the ionic environment sensibly constant during the progress of titration, for necessary conductivity, and for the practical elimination of junction potentials. These data were used in support of the extended theory of acid-base titration presented. NEWARK, DELAWARE
RECEIVED DECEMBER I O , 1941
DEPARTMENT OF CHEMISTRY AT THEOHIO STATE UNIVERSITY]
The Preparation and Directed Chlorination of l,l,l-Trifluoropropane BY ALBERTL. HENNEAND ATHERTON M. WHALEY
Trifluoromethyl groups have been synthesized before by halogen exchange between an inorganic fluoride and an organic trihalide RCX3 (where R was H or CH3),192or by the interaction of antimony trifluoride and an allylic trihalide.3 Attempts to prepare CH3CH2CF3 by an extension of these methods indicated the need for modifications.
The experimental procedure was then modified and successfully yielded the desired trifluoride, C H ~ C H Z C F ~The . physical properties, together with those of the intermediate mono- and difluorides (CH&H&C12F and CH3CH2CClF.J are listed in the general table.
Preparation of CH&!HzCFs.-Ten moles of CH&H= CClz (prepared as indicated above) and twenty-five moles Preparation of CH&H&!Cl8.-Ten moles of commercial propylene dichloride was chlorinated a t the boiling point, of hydrogen fluoride were heated in a steel bomb at 75" in the dark, and with iron filings as catalyst4 to give six for five hours, and then a t 95' until a pressure of 20 atmosmoles of CHaCHClCHClz, b. p. 130-133'. Ten moles of pheres was obtained. This treatment generated hydrogen this was heated to reflux and vigorously stirred with fifteen chloride, which was tapped off intermittently through a moles of 20% aqueous sodium hydroxide,6 then distilled condenser held a t -78', until the pressure no longer rose to 20 atmospheres by further heating at 95'. The bomb (instead of decanted) to give nine moles of CHsCH= CC12, b. p. 75-77'. Anhydrous hydrogen chloride was was then cooled to -78', and its contents poured into anled into ten moles of this olefinic dichloride in the presence other bomb containing 1300 g. of antimony trifluoride of 2 t o 3% of aluminum chlorides to yield 4.5 moles of combined with 250 g. of chlorine. This bomb was then heated with a free flame and the product tapped off a t a CH3CH2CC13, b. p. 106.5-108.5'. pressure of 13 atmospheres. The tapped gases were First Attempted Synthesis of CH3CH2CFa.-CH8CH2washed through dilute sodium hydroxide, dried over calcc1s was subjected to the customary halogen exchange cium chloride and condensed with solid carbon dioxide. with antimony trifluoride2 in the presence of catalyst. The condensate amounted to 800 g., from which 3.6 moles The treatment gave only about 5-10% of CH3CH2CF3, of CH3CH2CF3,b. p. -13', and 3.6 moles of CH8CHzand 10% of a mixture of C H ~ C H Z C F Zand C ~ CH3CH2- CClF2 were separated by low temperature distillation. CFClz. When the halogen exchange was performed withMore trifluoride was obtained from this intermediate out catalyst, practically no fluoride was obtained. Much difluoride as follows. Ten moles of CH8CH2CClF2 were decomposition was observed, hydrogen chloride was freed, placed in a bomb containing 1300 g. of antimony trifluoride and the only recovered product was CH3CH=CC12, which combined with 400 g. of chlorine. Ten moles of hydrogen accounted for 50% of the original CH3CHZCCl3. fluoride was condensed into the bomb which was then Second Attempt.-The allylic trichloride CH,=CHCCl, heated a t 95' on a water-bath. The reaction products was prepared by dehydration with phosphorus pentoxide were tapped off at 12 atmospheres through a water reflux of CHsCHOHCC13, obtained by condensing chloral with a condenser, and gave on redistillation 8.5 moles of trifluomethylmagnesium halide.'J It failed to exchange halo- ride and 0.3 mole of recovered material. gens with antimony f l ~ o r i d e . ~ The strongly directing effect of fluorine clusters (1) Henne, THISJOURNAL, 69, 1200 (1937). (2) H e m e and Renoll, ibid., 68, 889 (1936). (3) Henne, Whaley and Stevenson, ibid., 63, 3478 (1941). (4) Levine and Cas, U. S. Patent 2,119,484,May 31, 1938. (5) Cass, U.b. Patent 2,134,102,Oct. 25, 1938. (6) Levine and Cass, U. S. Patent 2,179,218,Nov. 7, 1939. (7) Vitoria, Rec. frav. chim., 24,265 (1905). 18) Kharasch. el al.. THISJOCJRXAL, 63, 2.568 (1911)
upon c h l o r i n a t i ~ n was ~ ~ ' ~verified once more when C H ~ C H Z Cwas F ~ subjected to the action of chlorine in sunlight, in a glass container and in the presence of water. The distillation of the reac(9) Heme and Renoll, ibid., 69,2434 (1937). (10) Henne and Haeckl, ibid., 63,2692 (1941)
