Dee., 1963 creasing values of K z . This effect is most likely due to the increasing surface concentration as a result of the decrease in the polar character with chain length. It can be secn from Table I that K z decreases with increasing carbon chain length. For low solute concentratioiis, many adsorbed surface films in aqueous systems can be considered to be “gaseous.”* l1 At higher concentrations, deviations from the ideal behavior occur. Similarly, deviations from the ideal behavior infused saltscan be seen in Fig. 1. At low surface pressure the plots of the surface pressure times area us. surface pressure for the varioys solutes approach the value kT = 607 ergs (area in A.2 units). With increasing concentration, the curves deviate from the ideal value IcT. The larger the size of the solute niolecule, the higher the deviations. The relatively small molecules of sodium acetate and sodium propionate approach closest to the ideal state. Curves similar to those plotted in Fig. 1 have been found for aqueous solutions of dibasic aliphatic esters with 10 and 11 carbon atoms.* The plots in Fig. 2 have the typical features of aqueous systems with adsorbed monolayers as specified by Adarn.ll Extrapolation of the curves in Fig. 2 to R = 0 gave area values which were in reasonable agreement with values given in the literature for the cross sections of adsorbed molecules in closely paoked films with chains arranged perpendicular to the surface of the solvent (for the hydrocarbon chain 20-21 A.2,89for the perfluorocarbon chain 29 8 . 2 10).
The curves shown in Fig. 1 and 2 are composed from several overlapping curves each of which represents the plots of the individual members of a series. As a result, the surface concentration-surface pressure relationship for any given members of a series is identical. This implies that the area occupied by the adsorbed molecules of different chain length of a homologous series is the same and that the ions are therefore arranged more or less perpendicular in the surface. If the same number of ions lay flat iii the surface, they mould occupy different areas and thereby alter the surface concentration-surface pressure relationship of the different inembers within a homologous series. The short carbon chain length and the close vicinity of the polar part of the molecule to the carbon chain are likely to prevent a lateral adhesion between hydrocarbon chains even a t high concentrations. It is known that lateral forces between hydrocarbon chains of adsorbed alkylates become appreciable only a t a chain length of n = 12, and the intermolecular forces between adja,cent perfluorocarbon chains are even weaker. The nearly quantitative agreement of the area occupied by the adsorbed ions a t high concentratioiis in the fused eutectic, determined by using the simplified form of the Gibbs adsorption isotherm and by using other methods in aqueous solutions,s-lo shows that this equation is applicable to fused salt systems. This appears to be reasonable, since ionized surface-active substances in fused salts can be compared to ionized surface-active substances in water, which contains an inactive electrolyte and for which the simplified Gibbs adsorption is0 t herm is known to be applicable. i l l ) N. K. Adam, “Physics and Cliemistry of Surfaces, ’ Oxford Umv. Press, London, 1941, p . 40.
NOTES
2853
T H E INSTABILITY CONSTANT AND HYDROGEX ION DISSOCIATION OF CALCIUM HYDROGEN PHOSPHATE BY ISIDOR GREENWALD Department of Baochemasti y, New York C’nauers.tty College of Medzczne, N E WYork 16, NEW York Recezned June 15, 196$
I n 1940, Greenwald, Redish, and Kibrickl reported that solutions of mixtures of ;“\Ta2HP04and XaH2P04 that contained CaClz or MgCl2 were more acid than were similar mixtures that contained NaC1, or KC1 of similar ionic strength. The differences were consistent with the view that undissociated CaHP04 and MgH P 0 4 existed in these solutions. Calculations of the instability constants [M2+][HP042-]/ [NIHP04] gave values that increased with greater ionic strength. I n the previous year, Shima2 had reported the effects of KC1, NaCI, MgC12,and MgSO4 upon the pH of halfneutralized 12.5 mM phosphoric acid. Our calculatioiis from his data yielded values for the instability constant similar to those we had obtained from our own experiments. Some years later (1953), Gosseliii and C ~ g h l a nem,~ ploying the equilibrium between Ca46,in the presence and absence of phosphate, and a cation-exchange resin, also came to the conclusion that the complex CaHP04 existed in solution. I n that same year, Davies and Hoyle4 claimed that the major component of the complex was CaH2P04+ and that the amount of CaHP04 was less than onethird as great. However, their determinations of pH mere made colorimetrically aiid the pH given for their phosphate buffer mixture does not agree with the values calculated from the P K ‘ ~of phosphoric acid. If one assumes pK’, calculated from the pH of their 19 : 1 (KH2PO4 :Sa2HP04) mixtures, one arrives a t values for the negative logarithms of the instability constant of CaHP04 of 1.81, 1.81, and 1.83, which are quite similar to those we had reported. If, on the other hand, one calculates the true pH of 19 : 1 mixtures, one finds the values given by Davies and Hoyle (at different ionic strengths) to be 0.16, 0.20, and 0.20 too low. If the reported pH of the solutions containing CaClz is raised by these amounts, the values for the negative logarithms of the instability coiistaiits become 2.00, 1.82, and 1.75, again quite consistent with our findings. More recentJy, Bjerrum6 reported the results of experiments in which he mixed solutions of phosphoric acid and solutions of calcium hydroxide. He was unable to decide between the formation of a complex in which P: C a = 1 and another in which it would be 2. Probably, this was due to his insistence that CaHP04 and Ca(H2P04)(HP04)- must be strong acids, with pK = 2. Our assumption that the complex was CaHP04 was applied to Bjerrum’s data. It was found that the (1) I. Greenwald, J. Redish, and A. C. Kibrick, J . Bzol. Chem., 1S6, 65 (1940). (2) K. Shlma, J . Bzochem. (Tokyo), 29, 121 (1939). (3) R. E. Gosselin and E. R. Coghlan, Arch. Bsochem. Bzophts., 26, 301 (1953). (4) C. W. Davies and B. E. Hoyle, J . Chem. Soc.. 4134 (1953). (6) N. Bierrum, KgZ. Danske Vzdenskab. Selskab, Mat.-Fys. Med., 31, 7 (1958); translated b y M. H. Rand, United Kingdom Atomic Energy Research Establishment, trans. 841. (6) N. BJerrum and A. Cnmack, zbzd., 9 , 1 (1929)
2854
Vol. 67
NOTES
values for the negative logarithm of the instability constant increased with increasing pH. Apparently, CaHP04is not quite so weak an acid as me had assumed in 1940.7 Calculations of the instability constant Fvere made, assuming that pK‘ for the hydrogen ion dissociation of CaHP04 was 8.0, 8.5, and 9.0. The most consistent values were obtained with pK’ = 8.5. These are given in Table I. With two exceptions, they agree quite well. One that does not is no. 3, in which the paH (Bjerrum’s nomenclature; usually written pH) is given as 5,.930. The difference in titration was only 0.008 meq. A change of paH of 0.01 would increase this by 23% and the value of pK by 0.1. The other is experiment 9, in which the reported paH of 7.001 is clearly inconsistent with the value reported for the very similar experiment 8. With so little change in the proportion of Ca to P, the paH could not be very different. Accordingly, this author also has made the calculation assuming that p,H in experiments 8 and 9 were the same, viz., 6.94. TABLE I CALCULATION OF -log [Ca2+][HP042-]/[CaHPO,] ASSUMING THAT -log [H+l[CaP04-l/[CaHPO~]= 8.5 IGal, [PI, and paH from Bjerrum” ICa 1 [PI pK instability No. x 103 x 10s P ~ H constant 0.956 7.807 2.65 1 c. 910 2 96 1 115 7.171 2 .916 3 .914 1.723 5,930 2.026 2.60 0 716 7 916 4 .688 7.916 2.44 5 .687 .718 6 ,693 .721 7 968 2.56 7.169 2.79 .$65 7 .688 6 940 2.61 .960 8 .689 9 .688 ,965 7 OOlC ( I .go) 2.52 965 6.94d 9 .688 2.36 1 241 6.183 10 ,688 2 50 1.246 6 147 11 ,688 2 68 0 473 8 002 12 .458 2.68 ,476 7 968 13 .458 2.82 .623 6 992 14 .455 6 883 2.72 .670 15 ,456 2.73 826 6 171 16 .458 5 948 2 49 .862 17 .458 a See ref. 5. See text. Obvious error, see text. More probable p,H +. +
+
Most of our experiments of 1940 were a t a pH below that (ca. 7.0) a t which CaHP04 liberates an appreciable proportion of hydrogen ion. Accordingly, this author has recalculated the data from the three experiments a t 7.14, 7.31, and 7.11. The pK of the instability constant should be lowered by 0.16, 0.21, and 0.09, respectively. There is an error in the first entry under CaCI, in Table VI. It, should be 1.08, not 1.80. (7) We had stated t h a t our experiments yielded consistent values for [Mz+][HPOa2-] / [MHP04]only if the hydrogen dissociation constant was assumed to be 10-10 os less This conclusion was based upon expeiinlents with MgClz. Later, C~eenwald(J. Rzol. CAem , 161, 697 (1945)) a h o n 4 that CaHP04 appeared t o hberate H’and replace it w i t h C a H C 0 3 + a t pl-I as low as 7, whereas MgHPOi did not, even a t pH 8.0. _ _ l l l l
DONOR PROPERTIES OF SOME SUBSTITUTED ACETANILIDES B Y RUSSELL s. DRAGO AiXD ROBERT L.
CA4RLSoVt
Wzllzam A . Noves LoborQtory U n a w r s l t y of Ill%flo?s,Urbana, Iihnoas Recetaed June 10, 1065
In earlier articles we reported2-6 the effect substit-
uents on the carbonyl carbon of amides, of general formula RC(O)K(CHJ,, have on the donor properties of the carbonyl oxygen. Both aliphatic and aromatic substituents were employed and the donor properties were evaluated toward the acids iodine and phenol in the solvent carbon tetrachloride. The equilibrium constants and enthalpies for the formation of iodine adducts with a series of para-substituted benzamides were found to correlate with the donor property of the amide predicted from the Hanimett substituent constant, a. A plot of log K/Ko, where KO is the equilibrium constant for the S,X-dimethylbenzainide adduct, vs. u gave a straight line with a p-value of -0.55, determined from the slope. The small negative p-value indicates that iodine withdraws only a
K
log-
KO
= ap
slight amount of electron density from the donor, and an electron-releasing substituent increases the formation constant. It was of interest to determine the effect of substitution a t the nitrogen on the donor properties of the carbonyl oxygen. The a-electron density in the 0-C-K system is delocalized, so appreciable changes in the a-electron distribution are expected with changes in the properties of the group attached to nitrogen. This change in a-electron distribution affords a means whereby inductive and resonance effects a t nitrogen can be transmitted to oxygen and can affect oxygen donor properties. It is of interest to compare the magnitude of the influence of effects brought about by substitution a t both the carbon and nitrogen on the donor properties of oxygen. The equilibrium constants for the formation of iodine adducts with N-methylacetanilide and some K-methylacetanilides substituted on the 2’-position mere measured to determine the effects of nitrogen substitution, and these results are compared with the results obtained on iodine adducts of some substituted ben~aiiiides.~A linear free energy correlation is expected. Experimental Preparation and Purification of Materials.-The purification of iodine and CCla have been described previ~usly.~Eastman White Label N-methylacetanilide, N-methyl-o-acetotoluidide, and E-ethylacetanilide were purified by recrystallizing twice from an ether-petroleum ether mixture. The 2’-chloro-Xmethylacetanilide was prepared from methyl iodide and 2’chloroacetanilide by a reported procedure.? Excellent analyses for carbon, hydrogen, and nitrogen were obtained on all materials. Procedures and Calculations.-The experimental procedures for carrying out measurements and performing calculations have been previously described.3 All reported errors are a t the 90% confidencelevel.
Results and Discussion The experimental data, calculated equilibrium constants for the formation of 1: 1 adducts, and differences in molar absorptivities of the complex, EC, and iodine, €1,are contained in Table I. (1) Abstracted in part from the Ph.D. thesis of R. L. Carlson, Cniversity of Illinois, Urbana, Illinois, 1962. ( 2 ) R. S. Dsago, D. A. Wenr, and R. L. Carlson, J . Am. Chem. Soc., 84, 1106 (1962). (3) R. L. Carlson and R. S. Drago, zhzd., 84, 2320 (1962). (4) R. L. Carlson and R. 9. Drago, zbzd., 86, 505 (19G3). ( 5 ) RI. D. Joesten and R. 9. Drago, tbzd., 84 2037 (1962). (6) M. D. Joestenand R. S. Drago, zbia.. 84, 2096 (19G2). !7) E. Thielenape and A. Fulde, Be?.. 68B,751 (1935).
