Nov., 1960
SOLUBILITIES OF CALCIUM SOAPSOF LINEARCARBOXYLIC ACIDS
1741
tion constant itself. Though these workers chose to attribute the whole effect to activity change, the second alternative cannot be ruled out. As is now well-known10-12optical rotation is related to the The symbols fo, f-, f- represent the activity co- existence of absorption peaks in the far ultraviolet. efficients of the undissociated acid, bitartrate and Alterations in rotation are to be related to changes tartrate ions respectively, f + that of the proton. in the wave length and relative intensity of these The symbol C is the sum of tartrate species as absorptions. Such changes infer alterations in the energy levels of the molecule, and therefore possibly before and in the binding of the acidic protons to the carboxB = ( H + ) (HT-) 2(HzT) = (HT-)initial 2(H~T)initia1 ylate oxygens in the case of tartrate. It is therefore possible that there may in fact be changes in the Such a set of determinations would be valuable in dissociation of the acid under the influence of salts, tracing the actual course of the activity coefficients and in a manner specific to the cations, as is true of of the several species with concentration and with the optical rotatory effects themselves. Extension salt cation. of this to other weak acids with electron-donor As pointed out by Kolthoff and Busch,6 the salt groups is obvious. effect on tartaric acid dissociation may be due to (10) E. U. Condon, Rev. M o d . Phys., 9 , 432 (1937). activity effects alone, but there also exists the (11) W. Kuhn and E. Braun. Z. physik. Chem., [B]8 , 445 (1930). possibility of some real alteration of the dissocia(12) W. Moffitt and A. Moscoaitr, J . Chem. Phys., SO, 648 (1959).
+
+
+
METAL COMPLEXING BY PHOSPHORUS COMPOUXDS. 11. SOLUBILITIES OF CALCIUM SOAPS OF LINEAR CARBOXYLIC ACIDS' BY R. R. IRAN AND C. F. CALLIS Mansanto Chemical Company, Research Department, Inorganic chemicals Division, St. Louis 66, Mo. Received
May $1, 1980
Solubility products of calcium soaps of linear carboxylic acids from Ceto CI8are reported from measurements of the competition for the calcium between the insoluble soap and the soluble calcium tripolyphosphate complex. Data are given for various temperatures and ionic strengths. For the saturated soaps, the negative logarithm of the thermodynamic solubility = -2.63 1.24 product at 25', p [ K B P-,] is related to the number of carbons in the soap chain by the equation p[K.,] (No. of Carbons). The unsaturated linear calcium soaps, oleate and linoleate, were significantly more soluble than calculated from the expression for the saturated soaps.
+
Introduction One of the techniques for measuring complexity constants is to use a nephelometric end-point for determining the competition for the metal ion between complexing and precipitating anions. I n a previous paper12 this procedure was utilized to measure the complexity constants of calcium with linear polyphosphates, using measured values of the solubility product of calcium oxalate. In the present study, the reverse is done; linear carboxylate anions are made to compete with the tripolyphosphate anion in tying up calcium, in order to measure the solubility product of these calcium soaps. Experimental Chemicals.-Sodium tripolyphosphate hexahydrate was used as the source of tripolyphosphate anions. It was prepared by four repeated fractional crystallinations of commercial sodium tripolyphosphate from aqueous solutions of ethyl alrohol. The final sample analyzed to better than !)9.7% Na6PrOto. Tetramethylammonium tripolyphosphate was prepared by ion exchanging the sodium salt with the hydrogen form of 100-200 mesh Dowex 50W-X2 and neutralizing the resulting acid immediately with tetramethylammonium hydroxide, as previously described.8 The (1) Presented before the Division of Inorganic Chemistry, 138th Meeting of the American Chemical Society, New York, September 1060. (2) R. R. Irani and C. F. Callis, THISJOURNAL, 64, 1398 (1960). (3) J. R. Van Warer, E. J. Griffith and J. P. McCullouph, J . A m . Chem. S O L ,11, 287 (1955).
stock solution was maintained a t 25' and a pH of 12 to avoid hydrolytic degradation. The C.P. grade hexanoic, octanoic, decanoic, lauric, m-yristic, palmitic, stearic and oleic acids were obtained from Eastman Kodak. The C.P. grade heptanoic, undecylic and linolcic acids were purchased from Matheson, Coleman and Bell. The nonanoic acid was technical grade from Matheson, Coleman and Bell. The tridecylic acid wm obtained from the K and K Labs., Jamaica, New York. The acids were ronverted to tetramethylammonium derivatives by neutradaation with Eastman Kodak reagent grade tetramethylammonium hydroxide. Other chemicals were C.P.grade. Procedure.-The nephelometric titrations were carried out at a pH of 12 using the same procedure previously descTibed2 except for temperature control and precipitating anions. I n the prcsent experiments, temperature was controlled to 3: 0.1' using a heater in combination with a heat-sensing Thermotrol unit, manufactured by Hallikainen Instruments, Berkeley, California. The concentrations of the calcium-precipitating anions, the linear carboxylates, were chosen below the critical mirelle concentration, and such that the competition for calcium with tripolyphosphate was favorable. In addition, no PZOS was detected in the precipitates formed upon addition of a slight excess of Ca++ to the solution containing P3Ol0-s and linear carboxylate anions. The ionic stlengths were adjusted to the desired values using tetramethylammonium bromide.
Results and Discussion The raw data showing the number of cc. of a calcium solution that must be added to a solution containing tripolyphosphate and linear carboxylate
R. R. IRANIAND C. F. CALLIS
1742
anions to reach a point of incipient precipitation, are given in Table I. TABLE I SUMMARY O F DATA FOR THE BOXYLATE
Temp., "C.
Tripolyphosphate. M (C1 x 103)
25
2.23
C~'+-P~~I~-~-LIXEAR CARSYSTEMS~
Linear carboxylate, M (Cd
Cc.of8.86 X l O - * M C a + + s o h . (t~) to nephelometric end-point at ionic strengths of 0.1 1.0
Heptanoate (7)
0.416 0.498
3.92b 4.15b
.. *.
Octanoate (8)
25
25
2,2:3
2.23
37
1.94 I .94
50
1.94
0.208 0.277 Nonanoate (9) 3.16 x IO-*
3.79 x 3.79 x 2.53 x 3.79 x 2.53 x 3.79 x
10-2
10-2
10-2 10-2
5.98" 5.65"
5.87d 5.2gd
2.33
25
2.23
37 50
25 37 50
25
2.910 1.94 1.94
2.33 1.92 1.94
2.23 2.90
25
2.33
1.16x 1.74x
5.2 4.8
..
..
..
4.6
4.79 4.00 4.82 4.43
..
6.20
5.75
Undecylate (11) 1.21 x 10-3 6.11 .. 2.42 x 10-3 .. 3.63x 10-3
2.16 x 3.23 x 6.44x 8.04 x 5.36 x 6.44x 8.04 x
.. .. ..
50
25
1.94 1.94
2.33
10-3
.. ..
.. 5.75 4.85 5.90 7.37
4.49 4.30 5.24 5.20 5.04
..
Laurate (12) 5.0 x 10-4
6.03
1 .o
10-3
4.96
10-3
3.73
10-3
2.58
10-3 10-3
5.11 4.51
5.50 3.56 .. .. ..
x 2.0 x 3.0 x 2.0 x 3.0 x
10-3
10-3 10-3 10-3 10-3
Tridwylate (13) 1.86x 1 0 - 4 5.29
2.33x 10-4 3.2i X 10+ 2.33x 10-4 3.27x hlyristate (14) 4.37x 10-5 8.74
37
10-3
x
1.31x 0.84x 1.67x 0.84x 1.67X
10-5 10-4
10-4 10-4 10-4
Oleate (unsat. 18)
25
2.33
3.53 X 7.06 x 10-5 1.06x 10-4
4.61 2.53 3.37 2.92 1.17 Linoleate (double unsat. 18) 25 2.33 3.57 X 0.69 .. 1.43X 2.69 3.74 1.58 2.14X 3.18 2.86 x 10-4 2.92 .. a Total volune of solution was alwars 250 r c . Ionic strength = 0.56. Ionic strength = 0.35. Ionic strength = 1.25.
For any ionic strength and temperature a t a pH of 12 and a nephelometric end-point2 - =K8P bCaPaOio-*
Decanoate (IO)
25
Vol. 64
1.86 4.40
5.89
..
5.36 3.40 2.27 3.98 2.45 4.91 3.68
lo-' Palmitate (16) 7.62X lod 1.57 1.15 x 10-5 0.86
..
..
.. .. .. .. 4.04 2.40
4.30 1.92 0.95 .. ..
..
.. 1.56 0.55
(1)
where Ksp is the solubility product of the calcium soap K.,
=
(Ca++)(CH3(CH2)nCOO-)2
(2)
and /3capsolO3- is the complexity constant, A is the
volume of the solution in cc., C2and C1 are the molar concentrations of the linear carboxylate and tripolyphosphate anions, respectively, and z and y are the molar concentration and cc. of calcium to the nephelomet,ric end-point,, respectively. Using the known2z4values of P c ~ P ~ o the , ~ ~ data in Table I were interpreted with equat'ion 1 to give t>he apparent solubility products given in Table 11. The uncertainties shown as f are statistical 95% confidence limits. TABLE I1 SOLUBILITY PRODUCTS OF CALCIUM SOAPS
..
..
Cz ( C A / y Z - 1)
Calciiiin map Heptanoate Octanoate Nonanoate
Temp., OC. 25 25 25 37 50 Decanoate 25 Undecanoate 25 37 50 Laurate 25 37 50 Trideoylate 25 hlyristate 25 37 50 Palmitate 25 Oleate 25 Linoleate 25
a Extrapolated. strengt,h = 0.35.
Neg. log solubility product a t ionic strength of 0.1 1.0 o.o= 6.12 f 0.05b . . . . . . . . . . 6.3 f0 . 2 6 . 1 0 i .03' 5 . 5 0 f 0.07d 6 . 4 O i .06 8.79 i .04 ......... 9.10 f .06 8.69 3z ,04 .......... 9.00 i .06 8.53 i .04 .......... 8 . 8 4 f .06 9 . 1 0 3~ . 1 8.47 f . 0 2 9.32 f . I 5 10.04 i . I 3 9.90 i . 0 3 10.90 f .18 10.00 rt .04 . . . . . . . . . . 1 0 . 3 1 & .OF 9 . 3 8 f .04 ...... 9.69f .08 11.93 i .02 11.27 i. .02 12.16 f .04 11.49 f .02 . . . . . . . . . . 1 1 . 8 0 4 ~.04 10.74 i .OS .......... 11.05 f . I O 13.14 i .10 12.08 i .03 13.30 i , l 5 14.46 f .03 13.88 rt .06 14.66 f . l 5 14.01 i .05 .......... 14.32 f .08 13.49 i .02 .......... 18.80 & .04 17.10 f .02 16.18 i .07 17.42 i 04 14.94 f .OG 13.54 i .1 15.42 & .OS 14.20 i .01 12.75 rt .09 14.69 f . 0 3 Ionic st,rength = 0.56. Ionic Ionic strength = 1.25.
*
The solubility products a t infinite dilution, [KBP] -, were estimated using the relation plia, = P [ K * , I ~ ad; (4) where p is the ionic strength. It was found that the ratio of the solubility product's a t ionic strengths of 0.1 and 1.0 mas fairly independent of the specific soap. This justified extrapolation for calcium
+
(4)
J. I. Watters arid 6. AI. Lanibort, J . Am. Chem. Soc.. 81, 3301
(1959).
Kov., 1960
SOLUBILITIES OF CALCIUM SOAPS OF LIXEXR CARBOXYLIC ACIDS
1743
soaps, e.g., nonanoate, for which no data at p = 1 were collected. I n all cases, the activity coefficient correction factors were much smaller than the actual values, so that no serious errors were introduced by using equation 4 rather than the extended Debye-Huckel r e l a t i ~ n . ~ The thermodynamic solubility products for the saturated soaps, Ca [CH3(CHz),COO]z, a t 25O, shown in Fig. 1, were found to fit the semiempirical equation proposed by Skau and Boucher.6 Specifically, it was found that p[K,,],
=
+
neg. log. of thermodynamic KED= -2.63 1.24 X ( S o . of Carbons) ( 5 )
The solubility data for the saturated linear soaps at 25O, fit equation 5 with a standard deviation of f 0.34 pK,, unit, immaterial of whether the chain length of the soap was even or odd. The solubility product values at 25' for calcium oleate and calcium linoleate, also shown in Fig. 1, lay significantly below those for the saturated calcium soaps, presumably due to unsaturation in the chain. The solubility products from this work are more precise than those obtained by Yoke7 on calcium palmitate, laurate and oleate using radio-tracer techniques. However, the two sets of data agree within experimental error. Solubility product measurements for calcium soaps of saturated linear carboxylic acids having chain lengths of 6 or lower, or 18 or higher, were not possible with the technique used in this work due to unfavorable competition with the tripolyphosphate anion; the former gave too weak, and the latter too strong a competition. Nevertheless, the estimated values a t 25' would lie very close to the line shown in Fig. 1. This implies that under experimental conditions in the presence of Ca++ and immaterial of the relative concentrations of all species, presence of P3010-5would pro( 5 ) H. S. Harned and B. B. Oven, "The Physical Chemistry of Electrolytlc Solutions." Reinhold Publ. Gorp., New Pork, N. Y., 1950, 2nd ed., p. 121. (6) E. L. Skau and R. E. Boucher, THISJ O U R U ~ L68, , 460 (1954). (7) J. T.Yoke, ?bid., 62, 753 (1958).
6
8
10
12
14
16
18
30. of carbons in linear carboxylate. Fig. 1.-Thermodynamic solubility products of calcium soaps at 25": 0 , Ca[CHa(CH2),C00I2;0, calcium oleate (one double bond); o, calcium linoleate (two double bonds).
hibit Ca [CH3(CH2)&OO]2 precipitation, but cannot solubilize Ca [CH3(CH&aCOO]2. The solubility product data a t various temperatures fit the equation d In K., dT
-
AH
RT2
(6)
where AH is the heat of solution of the precipitate in water a t infinite dilution. The values of AH for calcium undecanoate, laurate and myristate were calculated to be - 22.2 f 0.8, - 19.8 f 5 and -15.8 i 3.6 kcal./mole, respectively. For precipitation of calcium nonanoate, the value of AH was found to be - 4.7 f 1.2 kcal./mole, significantly lower than that determined for the other calcium soaps. Since the nonanoic acid was the only technical grade material, no valid arguments for the discrepancy can be presented.
