DISTRIBUTION STUDIES OP FORMIC, ACETIC, PROPIONIC,
AND BUTYRIC ACID BETWEEN BENZENE AND WATER
BY
DONALD H. ZIPPER
A THESIS
Submitted to the Faculty of The Creighton University in Partial Fulfillment of the Requirements for the Degree of Master of Science in the Department of Chemistry
/
OMAHA, 1943 TABLE OP CONTENTS
INTRODUCTION
EXPERIMENTAL
RESULTS .,
DISCUSSION .
SUMM ARY . . ..
BIBLIOGRAPHY INTRODUCTION
Henrys law, published first in 1803, relating the solubility of a gas in a given solvent to the pressure of the gas, was a limited statement of the distribution law. In 1872, Berthelot and Jungfleisch found the same principle applied to the partition of iodine between the two immiscible solvents carbondisulfide and water. Nemst, in 1891 formulated the law in general terms, thus: At a given temperature a substance present in two phases in eauilibrium is distributed between the two phases in a definite ratio of concentrations, provided it is not different chemically in the two phases. The law may be expressed mathematically, thus: C2 - K, (1) °1 ~ *
In which Cj and Cg are the concentrations of the sub stances in the two phases, and is the distribution constant." When passage of the solute from one solvent to the other is attended by partial ionization or dissociation or polymerization, the distribution ratio is no
tucker and Meldrum, Physical Chemistry p. 255. 2
longer a constant, If and Cg denote total concen trations, for it will be remembered that the distribu tion ratio is constant only with respect to a single molecular species. Some slightly ionized organic acids exist almost wholly as single molecules in water and almost wholly as double molecules in benzene. The various equilibria involved are shown in the following scheme, in which the dotted line represents the benzene-water interface and the solute is benzoic acid, written HBz for C6H5C00H:
Benzene layer , Water layer i £(HBz)2i==? HBz — ] = ? HBz + Bz” I K1 K2 K3 In the benzene layer (HBz) - KXV ((HBz)2j where the parentheses indicate molar concentrations. Since a very large portion of the solute in benzene is double molecules, it is nearly true that (HBz) — 0^% where is the total concentration In the benzene layer, expressed as moles of single molecules per liter.
In the water layer Kg is the Ionization constant of benzoic acid, and ifoc is the fraction ionized, (HBz)
In the water layer is Cw ( 1 - Q£ ). Equation (1) applied to the distribution eauilibrium for single molecules Is thus 3
C Kg — w ( 1 ~ oC ) _ const* (2) V Sine© K-^ is unkown (though constant), we may combine it 2 with Kg in a single constant called the distribution ratio*
Recently, Wall end Rouse1"' studied the association of benzoic acid in benzene using an isothermal distilla tion method similar to that used by Lassettre and Dickln- 4 son , which involved the distribution of a volatile solvent between two solutions of (practically) non volatile solutes. Using this method Wall and Rouse were able to develop an equation which most nearly represented the dissociation of the dimer of benzoic acid. The equa tion which best fitted the data is
log K = 3.790 -( 1977/T ) g Later, Wall, using the equilibrium constant calculated * from the above equation, calculated the degree of dis sociation of the dimer of benzoic acid,|ft , using the expression ? K - 4(i C ~ T ~ " <> 7 in which C is the concentration of the dimer in moles' per liter. Wall modified equation (2), taking into 2 ’ Millard, Physical Chemistry for Colleges p. 360. 3F.T. Wall and P.E. Rouse, Jr., J. Am. Chem. Soc. 63, 3002 (1941). ' 4E.N. Lassettre and R.G. Dickinson, ibid.. 61, 54
F.T. Wall, ibid., 472 (1942). 4
consideration the dissociation of the acid dimer in benzene,@ , and obtained the expression
K - °w( 1 - PC ) (3)
V 1 - 0 )
Wall used the data for the distribution of benzoic acid
between water and benzene reported by Creighton to calcu
late • These values substituted in equation (4) gave
remarkably consistent values of the constant K.
It is the purpose of this paper to determine ana
lytically the apparent distribution of formic, acetic, propionic and butyric acid between benzene and water, and to determine as nearly ss possible the true distribu tion constant of the of the acids by use of corrections foroC , the degree of dissociation of the acid monomer in water, and @ , the degree of dissociation of the acid dimer In benzene.
0C may be determined quite easily by use of the dissociation constant of the respective acids and the common equilibrium expression for the dissociation of a 1 weak acid. o K - OC C 1 - X where C is the concentration In normal moles per liter and oc is the degree of dissociation of the acid in water.
As far as is known there has been no work done on the determination of the dissociation constants of the 5
polymers of either formic acetic, propionic or butyric
acid in benzene or solvents similar to benzene. However, the dissociation of some of these acids in the vapor state in air has been determined in terms of the partial pressures of the monomers. 6 Coolidge determined the dissociation constant, Kp, of the dimer, in the vapor state, of formic acid at temp eratures ranging from 10°C to 156°C, and derived an equa tion which best represented the data, especially at low temperatures and pressures. The equation is
log Kp _ 10.755 - 3090/T .
MacDougall7 determined the dissociation constant the acetic acid dimer in the vapor state at temp eratures ranging from 25°C to 40°C and gave the equa tion as best representing his work
log K = 11.789 _ (3590/T) . 8 ^ MacDougall also worked on the association of propionic acid vapor and determined the reciprocal dissociation constant l/Kp of propionic acid dimer and trimer in the / vapor state at temperatures ranging from 50°G to 65°C. He gave no general equation of Kp for the temperature
6 ------A.S. Coolidge, J. Am. Chem. Soc«^ 50, 2166 (1928)
F.H. MacDougall, ibid.. 58, 2585 (1936). 8 F.H. MacDougall, ibid.. 63. 3420 (1941). 6
range but calculated the quantities of heat absorbed in
the formation of one mole of the dimer and one mole of
the trimer, by applying the Clausius-Clapeyron equation
to the Kp*s experimentally determined. There has been no work reported on the properties
of butyric acid in the vapor state.
By using the Kp of the above acids and making the
assumption that the dissociation of the polymers of the
acids would be approximately the same in benzene as in
air because air and benzene have dielectric constants
of somewhat the same magnitude, it is possible to change the Kp to Kc by use of the eauation derived from the
ideal gas laws Kp - Kc (RT)AR (4) and use concentrations instead of pressures. With the
Kc thus obtained the degree of dissociation of the
dimer (J may be determined from the equilibrium expression
Kc - 2c(*‘ 1 -(> (5) derived from the general reaction
An ;f= ? nAl in which c equals Concentration of the monomer as analy zed. Ajj represents the polymer and A^ the unassociated acid* 7
EXPERIMENTAL
Materials Used. Merck’s A.C.S. reagent grade formic acid and Mallinckrodtfe analytical reagent grade acetic acid were used without purification.
The propionic acid was a technical grade from Welsch Scientific Company. Avolume of about 500ml. was distilled using a 12-inch reflux column. The first and last third portions were discarded. The portion collect ed boiled at 139° _ 141°C (uncorr.).
The butyric acid was a technical grade which was purified by distillation using the seme column as above.
Of the 1800ml. used in the distillation, the first 500ml. were discarded, and 400ml. were then collected. The purity is indicated by the refractive index, which was 1.3917 at 25°C.
Apparatus and Technique. The only apparatus used that is worthy of mention is the micro-burettes and the flask in which the mixture was allowed to come to equilibrium.
The micro-burettes could be read to an estimated 0.001ml.
The flask used was a 50ml. round-bottom distilling flask from which the side-arm was removed and the opening sealed. The reason for using this flask was that it gave a wide surface of contact between the two immisci ble solvents with a small total volume of solutions.
The technique employed in the analysis of the 8
solutions of the equilibrium mixture is as follows.
The acid being studied was added to equal volumes of benzene and water in the flask described above. To ob tain results which would show a somewhat regular increase of acid concentration, the amounts of solvents were kept constant and the acid added was increased by eaual amounts for each run. The flask was then stoppered and thoroughly mixed by vigorous shaking, and placed in a constant temperature bath at 25°C + 0.04°. The flasks were held in the water with burette clamps, which were fastened by means of rods, to the stirring device of the bath. The vibration of the electric motor caused a slight but sharp vibration of the flasks, thus provid ing ideal conditions for equilibrium to be attained.
A check run on the same solution showed that 5 hra* time to reach equilibrium agreed very well with one kept in the bath 22 hrs. During these analyses it was never convenient or necessary to analyze thè mixtures at intervals which would allow less than 5 hrs. in the bath. / Therefore, a determination of minimum time was not carri- 9 ed out. Brown and Bury in similar analyses found one hour to be sufficient for strong solutions and even less time for dilute solutions.
The solutions, at equilibrium, were then pipetted — ------§------— ------Brown and Bury, J. Chem. Soc.. 125. 2430 (1923). 9
off, placed In weighing bottles, weighed, and the amounts of acid determined by titrating with standard sodium hydroxide solutions, using phenolphthalein indicator^ Ihe standard solutions were of different strengths. The strength used depending on the concen tration of the acid solution being titrated. A sodium hydroxide solution was standardized against constant boiling hydrochloric acid and also against potassium acid phthalate. The two methods agreed very well and all subsequent standardizations, and checks on the solutions, were based on potassium acid phthalate as a standard, because of the ease of handling the latter.
Results. The results of the analysis are given in
Table 1. The distribution ratios of formic, acetic, propionic, and butyric acid in benzene and water plotted against molal concentration of the acid in water, are shown in Figures 1, 2, 3, and 4 respectively.
The molarities of the solutions were calculated.
The densities of the water solutions of formic and acetic acid were obtained from the Handbook of Chemistry.^ The density of water solutions of propionic acid was 11 obtained from the International Critical Tables. ~~10Lange. Handbook 'of "Chemistry. 2nd.' id. p. 1061. 1040. Int. Crlt. Tables, vol. 3, p. 112. 10
TABLE 1.
Molality Molality C water Density Density Molarity Molarity in in C benzene of of in in Benzene Water Benzene Water Benzene Water Soln. Soln.
FORMIC ACID 0.00742 2.352 316.9 0,8740 1.0240 0.00674 .01578 4.547 2.173 288.1 .8740 1.0424 .01378 3.919 .02369 6.266 264.5 .8 740 1.0547 .02068 .03927 8.987 5.129 228.8 .8 740 1.0712 .03428 .04745 6.805 11.35 239.3 • 8 740 1.0831 ,04138 .06777 8.078 14.86 219.3 • 8 740 1.0980 .05905 .08273 16.73 9.690 .8740 1.1049 .07205 .08233 17.23 202.2 10.45 209.3 .8 740 1.1067 .07170 10.63
ACETIC ACID
0.00404 0.2674 66.29 0,8737 1.0002 0.00352 0.2646 .00705 .3960 56.20 .873 7 i.0010 .00615 .3871 .01028 .4927 49.74 .8737 1.0017 .00898 • 4794 • 01341 .5971 44.52 .8738 1.0022 .01171 • 5777 .02374 .8465 35.65 .8738 1.0066 .02072 .8109 .04909 1.345 27.40 .8 740 1.0090 .04278 1.256 .06883 1.667 24.22 .8741 1.0113 .1370 .05992 1,531 2.462 17.97 .8745 1.0168 .1188 • 2021 2,176 3,302 16.34 .8749 1.0217 .1747 • 3306 2.818 4,500 13.61 .8757 1.0279 .2839 ,4120 3.641 5.373 13.04 .8762 1.0316 .3523 4.191 • 6 215 7.335 11,80 .8 774 1.0391 .5256 5.294 • 9292 10.22 10.99 .8793 1.0470 .7743 6.624 11
TABLE 1. (cont.)
Molality Molality C water Density Density Molarity Molarity in in C benzene of of in in Benzene Water Benzene Water Benzene Water Soln. Soln.
PROPIONIC ACID
0.03370 0.1691 5.018 0.8843 0.9982 0.02971 0.1667 .05320 .2102 3.592 ,8845 .9985 .04689 .2067 ,1151 .3386 2.940 .8847 .9993 ¡1011 .3301 .2132 .4991 2.341 .8854 1.0004 .1858 .4923 .3723 .6996 1.879 .8865 1.0015 .3212 .6663 .5396 .8983 1.665 .8878 1.0028 .4607 .8445 .9042 1.266 1.355 .8903 1.0048 .7556 1.156 1.891 1.993 1.054 .8963 1.0088 1.487 1.732 3.020 2.771 0.9173 .9018 1.0128 2.226 2.334 4.317 3.702 0.8576 .9077 1.0170 3.073 2.955
BUTYRIC ACID •
0,1283 0.1646 1,282 ,1598 .1939 1.213 .2493 .2483 0.9962 .3442 .3014 .8756 .7893 .5054 ,6403 1.6301 .7893 .4842 2.8571 1.121 .3925 3.0726 1.153 .3752 4.9900 1.580 ,3166 320
300
280
260 iter mze:
240
220
200
FIGURE 1 FIGURE 2 S 14
Propionic acid molality in water solution
FIGURE », 15
Butyric acid molality in water solution
FIGURE V The densities of the benzene solutions of formic acid were considered the same as that of pure benzene because of the very dilute solutions. Densities of the benzene solutions of acetic12 and 13 propionic acid were also obtained from the Inter national Critical Tables.
12 Ibid.. p. 156. 13 Ibid.. p. 164 17
DISCUSSION
The ratios of the distribution of the acids be tween benzene and water in Table 1 show definitely that some change of the acid molecule has taken place in one of the solvents, because the ratios are far from being constant. If no change had taken place In either of the solvents or if the change was the same in both solvents, the equation r
where Cw is the concentration of acid in water and is the concentration of acid in benzene, would hold over a range of concentrations.
There Is reason to believe that the acid forms a dimer in the benzene. This being the case we would » have the following equilibrium in which HA represents the acid monomer in water and (HA)g represents the acid dimer in benzene.
2HA (HA)2
The equilibrium constant K would be
(HA)2 with concentrations expressed in moles per liter of the single molecules. Substituting Cw and for con centrations of HA and (HA)g respectively, we have the constant which in this case would be the distribution 18
constant K, a K, c ~ W Using this expression for the distribution constant, a much better constant is obtained. The results of which are shown in Table 2* Prom these results it is apparent that the dimer is formed in the benzene.
In the above it was assumed that no dissocia tion of the acid monomer in water or of the acid dimer
in benzene took place* When these dissociations are
taken into consideration, the following eaullibrium would exist. water benzene 2A~ + 2H+ ^=? 2 HA 2HA (HA) „ k 3 Prom this equilibrium the equilibrium expression
for the dissociation in the water and dissociation in benzene may be derived.
For any mono-basic acid in water the dissociation constant is given by the following expression.
K, _ oC2M where M is the molar concentration and oC is the degree
of dissociation. Prom this cC may be calculated for any concentration If the dissociation constant K2 is known. For the dissociation of the acid dimer in benzene the equilibrium expression is derived somewhat differently. 19
If we represent the dissociation of the dimer in benzene by the equation ( HA ) 2 — ___^ 2 HA then p K„ - (HA)15 5_ i®a Let (3 = degree of dissociation of the dimer.
M = concentration of the monomer as analyzed.
Then M — w " " dimer. 2 “ " HA
M M ft - " (h a )2 ------T* -
Substituting in these values we have
K, - (HA)2 2 - 2Mft2 (6) _P >i= f-f From this (3 may be calculated for any concentration if the dissociation constant Kg is known. If a trimer is formed as is the case for propionic acid the same method may be used except that the concen tration of the trimer must be represented by M • 3 If we assume that the acid in the benzene is all in the form of double molecules and all single molecules in the water the equilibrium may be represented by this equation, w^ter 'benzene 2HA ± ^ (HA)2 cw where and Cw are the concentrations in moles of 20 single molecules per liter and the dotted line re presents the water-benzene interface* If the dissociation of acid monomer in the water is to he considered in the equilibrium, Cw would be Cw (l — OC) inhere oC is the degree of dissociation of the monomer* The equation is then water 1 benzene 2HA -*--- '--- y. (HA)„ 0 „(1 -OC) ' C„ 2 Prom this the dissociation constant is (7) c ' ^ T ^ ) When the dissociation of the acid dimer in the benzene is also considered the Cb is C^(l — (J ) , where (3 is the degree of dissociation of the dimer. The equation is then water , benzene 2 H A ± = ! (HA) 2 Cw (l -OC) - Cb (l ~ / ( ) w ¡K if 8 The degree of dissociation of the acid, oc , in water was calculated from the ionization constant of the acids by use of the general equilibrium expression where M is molar concentration, 21 K, oÇ2-M (9) 1 -OC The constants used for formic, acetic, and propionic acid were 1.4X10~ , 1.75X10” and 1.4X10"Respectively. The oC for butyric acid was not calculated because no values for the dimer-monomer equilibrium constant could be found in the literature, and, hence, ($ could not be evaluated. Unless both In calculating (Î we must have a dissociation constant which is in terms of concentrations. As stated earlier the constants reported in the literature are in terms of pressures. We must change the Kp to Kq. This may be done by use of equation (4) which is based on the ideal » gas laws, Kp = K0 (RT)in . For formic acid we substitute into the general equation of Kp for the dissociation of the dimer, as 14 given by Coolidge, the temperature at which the Kp is desired. From the equation of Coolidge for the varia tion of Kp for formic acid with temperature, log Kp = 10.755 - 3090/T the value of Kp at 25°ç is 2.451 mm. 14 Coolidge, 0£. cit. . 50 p. 2166. -4 Using equation (4) a value of 1.32X10 is obtained for Kc with the concentration in moles per liter. Using this value for Kc for in equation (6),@ , the degree of dissociation of the dimer of formic acid in benzene was calculated. The constant Kp chosen for acetic acid is a direct measurement of Kp at 25°C. This directly determined value of Kp was used instead of one calculated from the equation derived from the data at several tempera tures because the experiments at 25°C were reported as being most accurate. The Kp at this temperature is 0.547 mm. Prom this value the Kc was determined in the manner described for formic acid. The Kc for acetic acid in benzene is 2.94X10 , when concentrations are in moles per liter. The @ is then determined in the manner described for formic acid using equation (6). The constant K for propionic acid was not given at 25°C nor was any general equation given relating Kp and temperature. The AH for the dissociation of the, di mer was calculated from the Kp given at temperatures ranging from 50° to 65°C. Using the average AH for these temperatures and the value of Kp at 59.98°C as 15 ^ r\ given by MacDougall, the Kp at 25°C was calculated by 15 " ------MacDougall, 0 £. clt.. 65 p. 3420. 23 use of the integrated Clausius-Clapeyron equation. The Kp at 59.98°C is given a 6.897 mm. The Kp at 25°C, calculated, using 18,430 cal. as the average A H for the range reported, is 2.630 mm. Changing from Kp to Kc by means of equation (4), the K is found to be 1.414X10 . Using this constant and equation (6) the degree of dissociation of the dimer, Q> , was calculated. The Kp and AH for the dissociation of the trimer of propionic acid were also given by MacDougallî5 The value of the Kp at 59.98°C is 9.80 mm. The average AH for the temperature range reported is 23,850 cal. Using these values in the Clausius-Clapeyron equation the Kp for the dissociation of the trimer at 25°C is 14.30 mm. The corresponding K is 4.135X10”8. c The degree of dissociation of the trimer to the mono mer was not determined because a trial calculation showed that its effect would be inappreciable. The dis tribution constant K^ for propionic acid considering the dissociation of the trimer as well as that of the dimer would be /cbd -Ç- y) K- (1 0 ) C j l -oc) where is the degree of dissociation of the trimer into the monomer, ( HPr ) 5 ± =---- ^ 5HPr 24 A trial calculation of the value of showed, that its effect would be inappreciable, so it was not used in the calculation of the constant given in Table 2. The values of oC and Q> are shown in Table 2. The values of the various distribution constants which were derived above, using the values ofoc and Q of the corresponding concentrations, are also shown in Table 2. The constant obtained when the acid in the benzene is considered wholly as double molecules is certainly a much improved constant over the constant where the acid is thought to be single molecules in both solvents* At first glance a great variation still seems to exist, but it must be considered that a wide range of concentrations has been used in determining the constant. The use ofoc » alone tends to make the constant vary still more, espec ially in the more dilute solutions, where the effect of oC is greater. When the effect of oc and Q are both con sidered there is an improvement in the constant. This improvement is not so very great but it is definitely- a trend in the right direction, especially in the more dil ute solutions where one would expect a greater effect due to dissociation. The reason for the constant to appear better in the more concentrated solutions is believed to be, that in concentrated solutions the effect of dissoc iation of the dimer and the monomer is negligible and 25 TABLE 2. Molarity- Molarity- v^T VET" V cb'(r-pj in in oC Benzene Water Cw Cw (l — ci) ^ ^ FORMIC ACID 0.C0674 2.173 0,010 0.094 0.0378 0.0382 0.0363 ,01378 3.919 .0074 .067 .0300 .0302 .0292 .02068 5.129 .0064 .055 .0280 .0282 .0274 ,03428 6,805 .0056 .043 .0272 .0274 .0268 .04138 8.078 .0051 .039 .0252 ,0253 .0248 .05905 9.690 .0047 .033 .C251 .0252 •.0247 .07205 10.45 .0045 .029 .0256 .0258 .0254 .07170 10.63 .0045 .029 .0251 .0252 .0250 ACETIC ACID 0,00352 0.2646 (J .0081 0.063 0.224 0.226 0.219 .00615 .3871 .0068 .048 .202 .204 .199 .00898 .4794 .0060 .042 .198 .199 .194 .01171 .5777 .0055 .035 .188 .189 .186 .02072 .8109 .0044 ,C 26 .178 .178 .176 .04278 1.256 .0037 .Ci9 .165 .165 .163 .05992 1.531 .0034 .016 .160 .160 .159 .1188 2.176 .0028 .011 .159 .159 .158 .1747 2.818 .0025 .009 .149 .149 .147 .147 .2839 3.641 .0022 .007 .147 .147 .3523 4.191 .0020 .006 .142 .142 .142 .5256 5.294 ,00i9 .005 .137 .137 .137 .133 .7743 6.624 .0016 .004 .133 .133 26 TABLE 2, (cont.) Molarity Molarity in in Benzene Water "“V PROPIONIC ACID 0.02971 0.1667 0.0092 0.014 1.038 1.041 1.030 .04689 .2067 .0082 .012 1.046 1.053 1.042 .1011 .3301 .0065 .0088 .960 .9 70 .964 ,1858 .4923 .0056 .0062 .875 .880 .878 ,3212 .666 5 .0046 .0047 .850 .855 .852 .4607 .8445 .0041 .0039 .805 .808 .806 ,7556 1.156 .0035 .0031 .752 .755 .752 1.48 7 1.732 .0028 .0022 .70 5 .708 .705 2,226 2.334 .0025 .0018 .670 .672 .670 3,073 2.955 . .0022 .0015 .5 70 .571 .570 BUTYRIC ACID Molality Molality Ratio in in * using Benzene Water Molality 0.1284 0.1646 2.18 .1598 .1939 2.06 ,2493 .2484 2.00 .3442 .3014 1.95 .7893 .5054 1.75 1.630 .7893 1.64 2.857 1.121 1.51 3.073 1.153 1.52 4.990 1.580 1.42 > o £ 4 0 7 27 therefore the acid is more nearly composed of wholly double molecules in the benzene and of wholly single molecules in the water* If» however, the true values of oc and @ for all concentrations were known, the con stant should hold for the weaker concentrations too. The reason OC is not the true degree of dissociation is that the ionization constant does not hold over a range of concentrations as widely as those used. In dilute solutions, however, the oc should be quite exact. The Q> is the largest error in determining the constant. First of all, the assumption was made that the dimer would dissociate the same in benzene as in air. This assumption is probably not far from right but neverthe less the dielectric constant of benzene is approximately twice that of air, so it is likely that the degree of dissociation of the dimer in air is somewhat less than in benzene* The error in this assumption, therefore, would be in the direction of a value for Q> which Is too small. However, this error cannot be considered very significant since the value of Q> is not very sensitive to the value of K used* From the constant expression it is seen that Q> being larger, especially in the dilute "St solutions, would improve the constant* An additional error in calculating the value of $ is in the conversion of Kp to Kc. Equation (4) as siamés 28 thst ideal gases are used. The values of Kp used for all three of the acids had been determined at tempera tures considerably below the boiling point of the acids. Under these conditions, the vapors of the acids deviate significantly from ideal behavior* Finally, it must be mentioned, the distribution law itself is of the nature of a limiting law, and, hence, consistent values of the constant can only be expected, at best, in very dilute solutions. It is significant that the corrections here applied for the ionization and dimerization of the solutes do appreciably improve the value of the distribution constant in dilute solutions. It is evident that these corrections alone are insuffic ient to extend the application of the distribution law » to fairly concentrated solutions. The eouilibrium in appreciably strong concentrations obviously involves additional factors which at this time cannot be determined. 29 SUMMARY 1. The distribution of formic, acetic, propionic and butyric acid between benzene and water, 25°C, has been determined over a wide range of concentrations. 2# A correction for the dissociation or ioniza tion of the acid monomer in the water was made. 3. An attempt was made to correct for the poly merization of the acid in benzene in calculating a distribution constant of these acids# > 3 0 BIBLIOGRAPY Books. Gucker, Prank Thomson, Jr. end Meldrum, William Buell. Physical Chemistry. New York: American Book Co., 1942. International Critical Tables. 1st. ed., Vol. III. New York: McGraw-Hill Book Co. 1928. Lange, Norbert Adolph. Handbook of Chemistry. 2nd. ed. Sandusky, Ohio: Handbook P u b l i s h e r s I n c . 1937. Millard, E. B. Physical Chemistry For Colleges. 5th. ed. New York: McGraw-Hill Book Co., 1941."" Articles Brown, Frederick S. and Bury, Charles R. »The Distribu- tion of Normal Fatty Acids between Water and Ben zene. Journal of the Chemical Society. 123. 2430 v ) • Coolidge, Albert S. »The Vapor Density and some other Properties of Formic Acid». Journal of the Ameri can Chemical Society. 50, 2166 (1928). ------*” Lassettre, Edwin N. and Dickinson, Roscoe G. »A Compara- tive Method of Measuring Vapor Pressure Lowering with Application to Solutions of Phenol in Benzene». Journal of the American Chemical Society. 61, 54 MaoDougall, F.H, »The Molecular State of the Vapors-of Acetic Acid at Low Pressures at 25, 30, 35, and 5 Q ) 0f the Am9I*i-can Chemical Society. MacDougall, F.H. »A Study of the Vapors of Propionic Acid at 45, 50, 55, 60, and 65°». Journal of the American Chemical Society. 63. 3420 (1941). 3 1 Wall, F.T. and Rouse, P.E. Jr, "Association of Benzoic Acid in Benzene". Journal of the American Chem- cal Society. 65, 30012, (1941). Wall, Frederick T. "Distribution of Benzoic Acid between Water and Benzene". Journal of the American Chemical Society. 64, 472 (1942). /