1546
ALFREDA. SCHILTAND G. FREDERICK SMITH
Vol. 60
ACID DISSOCIATION CONSTANTS OF SUBSTITUTED l,lO-PHENANTHROLINES1 BYALFREDA. SCHILTAND G. FREDERICK SMITH Contribution from the Noyes Chemical Laboratory, University of Illinois, Urbana, Illinois Received June 11, 1966
The acid dissociation constants of forty substitution derivatives of 1,lO-phenanthroline were determined potentiometrically at 25". A linear relationship which permits extrapolation of data to find pK, values in water for derivatives insoluble in water was shown to be obtained between per cent. dioxane and relative p K , values in water-dioxane solutions. The Hammett substituent constants for methyl groups were shown to be additive for multiple methyl substituents.
Introduction The important applications of 1,lO-phenanthroline and its numerous substitution derivatives in the field of analytical chemistry have created considerable interest in their physical properties and those of their chelate complexes with various metal ions. Aside from its utilitarian value to the analytical chemist, this information should afford some insight into the related factors of structure and reactivity. At the present time relatively few physical constants of the 1,lo-phenanthrolines and of their complexes have been determined. The acid dissociation constant of 1,lO-phenanthroline as well as the instability constants of its Fe(I1) and Fe(111) complexes have been determined by Lee, Kolthoff and Leussing.2 Determinations of these same constants have also been carried out by Brandt and Gullstrom3 on some 5-substituted 1,lOphenanthrolines and on the unsubstituted compound. The basic strength of a chelate ligand, among other factors of its structure, determines to a large extent the stability of its metal complexes. Some exceptions have been noted to this general rule, but it should remain valid, excluding steric effects, when applied to a given metal ion and a series of closely related chelate molecules. In the case of the metal complexes with various substituted 1,lOphenanthrolines it is t o be expected that the relative order of stabilities of a given metal complex should be the same as that of the basic strengths of the ligands. This follows from the similarity of the 1,lo-phenanthroline derivatives. Any exceptions would most probably result from steric effects introduced by substitution on the phenanthroline ring. It is proposed that the empirical relationship, log k - log kO = pa, given by Hammett4 may find application in predicting approximate values of stability constants of the substituted 1,lO-phenanthroline metal complexes. Employing the acid dissociation constants of the substituted derivatives of 1,lO-phenanthroline to calculate +values and the experimental results of Brandt and Gullstrom on stability constants of Fe(I1) complexes of 5-substituted 1,lo-phenanthrolines to obtain p , it is believed that satisfactory values for the sta(1) Based upon a portion of the dissertation submitted by Alfred A. Schilt in partial fulfillment of the requiraments for the degree of Doctor of Philosophy in the Graduate School of the University of Illinois. (2) (a) T.S. Lee, I. M. Kolthoff and D. L. Leussing, J. A m . Chem. SOC.,TO, 2348 (1948); (b) TO, 3596 (1948). (3) W. W. Brandt and D. K. Gullstrom, ibid., 74, 3532 (1952). (4) L. P. Hammett, "Physical Organic Chemistry," McGraw-Hill Book Co., Ino., New York, N. Y., 1940,pp. 184-207.
bility constants of iron(I1) complexes of other substituted ligands may be obtained. A straight-line relationship has been demonstrated between the pKa of substituted 1,lO-phenanthrolines and the oxidation-reduction potentials of the ferrous-phenanthroline complexes with nitro, methyl, bromo, chloro or phenyl substituents in the 5-p0sition.~~6The ruthenium complexes appear to give similar correlation.6 In general such a relationship is to be expected in the event that the stability constants of both the oxidized and the reduced forms of a complex are linearly dependent on the basicity of the chelate ligand. The acid dissociation constants of the substituted 1,lO-phenanthrolines may, therefore, provide for the prediction of approximate oxidation-reduction potentials as well as of stability constants. Such an approach appears attractive in light of the experimental difficulties involved when the water insoluble phenanthroline complexes are to be evaluated. Experimental Methods and Results Materials.-The various substituted 1,lO- henanthrolines were prepared either in the laboratories of f'emple University' or by the G. Frederick Smith Chemical Company following these directions. They had previously yielded correct analyses for carbon and hydrogen, and were considered to be as pure as crystallization procedures afford. The best available commercial grade of 1,4-dioxane was purified by refluxing over a mixture of sodium hydroxide and silver oxide for 48 hours. The distillate obtained after this treatment was stored over sodium metal; prior to immediate use it was refluxed for at least 24 hours before distilling requisite amounts. The deionized water used was carbon dioxide-free as indicated by pH measurement. Measurement of Acid Dissociation Constants.-A semimicro balance was used to weigh out samples of the 1 , l O phenanthroline derivatives in exactly 0.125 milliequivalent amounts which were then transferred into 25-ml. volumetric glass stoppered flasks. Using a weight buret samples of standardized hydrochloric acid (equivalent to 0.5-0.6 of the free base added) were weighed into the contents of the flasks. Various amounts of the 1,4-dioxane (at 25') were delivered by calibrated pipets into the flasks, the contents of the flasks were diluted to the mark with carbon dioxidefree water and the flasks were then stoppered and placed in a thermostat to equilibrate to 25". All of these operations were performed without appreciable delay in order to minimize the slight tendency for the pH to decrease as. COZ is picked up by the relatively weakly buffered solutions. The solutions were measured a t 25" using a Beckman model G pH meter equipped with a glass-saturated calomel electrode ,pair. The meter was calibrated immediately before use with Leeds and Northrup Standard Buffer of (5) R. V. G. Ewens, Nature, 155,398 (1945). (6) F. P.Dwyer, J. Proc. Roy. SOC.N.S. Wales, 84, 83 (1950). (7) (a) F. H.Case, J. Am. Chem. Soc., TO, 3994 (1948); (b) Ti, 821 (1949); (o) 71, 1828 (1949): (d) J. O r a Chem., 16, 941 (1951): (e) 16, 1541 (1951); (f) F. H. Case and R. Sasin, ibid., 20, 1330 (1955); (9) F. H. Case and J. A. Brennan, ibid., 19, 919 (1954): (h) F. H. Case, 8. Catino and F. Scholniok, ibid., 19,31 (1054).
. 8
Nov., 1956
ACIDDISSOCIATION CONSTANTS OF SUBSTITUTED 1,10-PHENANTHROLINES 1547
pH 4.01 a t 25". The performance of the instrument was checked peoriodically using a standard buffer solution of pH 6.86 at 25 . After completion of a series of measurements on any one compound, the calibration was checked using the pH 4.01 buffer once more. It was observed that continued use of the dioxane-water solutions caused the attainment of constant readings to be sluggish. This behavior occurred twice during the entire investigation and was eliminated both times by flushing out the saturated potassium chloride salt bridges with fresh solution. Other than this, no difficulty was experienced in obtaining rapid and reproducible measurements. Relative acid dissociation constants in the dioxanewater solutions of the conjugate acids of t,he substituted 1,lo-phenanthrolines were calculated from the pH meter readings and the known concentrations of conjugate acid and free base (1,lO-phenanthrolines are monobasic). Because the pH meter was calibrated for aqueous solutions, rather than for each particular dioxane solution, the measurements provide only relative p€I values for the dioxane solutions. To obtain the acid dissociation constants in water the relative pK, values were plotted versus weight per cent. dioxane and the curve extrapolated to zero dioxane. Since activity coefficients were not employed and because of unknown liquid junction potentials the dissociation constants are not thermodynamic values. The ionic strengths
TABLE I ACIDDISSOCIATION CONSTANTS AT 25" 1,lO-Phenanthroline derivative
Unsubstituted 4-Bromo3-Chloro3-Ethyl4-Ethyl3-Methyl2-Phenyl3-Phenyl4-Phenyl5-Phenyl4-n-Propyl3,s-Dibromo4,7-Dichloro5,6-Dichloro4,7-Diethyl4,7-Dime thoxy5,6-Dimethoxy2,4-Dimethyl2,9-Dimethyl3,4-Dimethyl3,7-Dimethyl3,s-Dimethyl4,6-Dimethyl4,7-Diphenoxy4,6-Diphenyl4,7-D iphenyl3,4,6-Trimet hyl3,4,7-Trimethyl3,5,6-Trimethyl3,5,7-Trimet hyl3,5,8-Trimethyl2,4,7,9-Tetramethyl3,4,6,7-TetramethyI3,4,6,8-Tetramethgl3,4,7,&Tetramethyl3,5,6,8-Tetramethyl3,4-Cyclohexeno5,6-Cyclohexeno3,4-Cyclopenteno(3,4),( 7,s)-Dicyclohexeno-
Concn. limits of dioxane used (% by wt.)
0-40 12-40 12-40 12-40 12-40 12-40 20-40 20-40 20-40 20-40 12-40 52-68 40-56 28-40 12-40 40-60 0-40 12-40 12-40 12-40 12-40 20-40 12-40 38-60 40-72 40-60 12-40 12-40 12-40 12-40 12-40 20-40 12-40 12-40 12-40 12-40 12-40 12-40 12-40 40-52
Extrapolated PKS
4.86 4.03 3.99 4.98 5.44 5.00 4.90 4.82 4.90 4.72 5.45 3.90 3.03 3.47 5.60 6.45 4.42 5.96 6.17 5.62 5.57 5.23 5.71 5.34 4.69 4.84 5.93 5.99 5.34 5.90 5.27 6.50 6.45 6.07 6.31 5.54 5.66 5.30 5.78 6.23
-m
1.53 1.30 1.50 1.40 1.55 1.43 1.23 1.67 1.80 1.70 1.58 1.85 0.56 1.25 1.18 1.60 1.23 1.55 1.50 1.28 1.33 1.68 1.28 1.98 1.90 2.02 1.55 1.09 0.98 1.18 1.os
1.10 1.25 1.35 1.55 0.65 1.55 1.33 1.58 1.73
at which the measurements were made varied from 0.0025 to 0.0030. Judging from a comparison of the attendant error of the experimental procedure and the small error from neglecting activity coefficients at these low and relatively constant ionic strengths, it may be concluded that these values might well be considered thermodynamic constants. The results are presented in Table I. The lower limits of concentration of dioxane used depended upon solubility requirements; this information is included together with the maximum concentration of dioxane employed in the measurements as an indication of the reliability of the extrapolations. Negative values of the slopes of the linear curves obtained by plotting the experimental data are given in the table also. The slope, m, is here defined as the rate of change of the relative pKB values with change in weight fraction of dioxane. Use of extrapolated p K . values and the slopes permits comparison to be made of the relative basicities of the 1,lO-phenanthrolines as a function of dioxane content.
Discussion To obtain acid dissociation constants referred to water, the data were extrapolated to zero dioxane content. This was necessitated by the insolubility of most of the substituted 1,lO-phenanthroline derivatives in water. It was shown experimentally that this technique, although theoretically unjustified as yet, would yield reliable results. Two of the more water-soluble compounds (1,lO-phenanthroline and 5,6-dimethoxy-l ,10-phenanthroline) permitted study in aqueous solutions and were found to give a linear relationship, a t least within experimental error, between relative pKa values and per cent. dioxane in the range from 0 to 40% dioxane. The observed linearity is fortuitous and may possibly result from the use of relative instead of absolute acid dissociation constants. The acid dissociation constants of some anilinium salts as determined by Marshall and Grunwald8 in concentrations of dioxane in water of 0-20% do not vary linearly with respect to per cent. dioxane by weight. Similar non-linearity is given by data of Gutbezahl and Grunwaldg for anilinium salts in dilute ethanol solutions in water, The dissociation constants of some p-diketones in the system dioxane-water were found to deviate from linearity at lower dioxane concentrations by Van Uitert, et ~ ~ 1 . I~ n0 this case non-linearity was suggested to result from a shift to the enol form as the mixed solvent became more concentrated in water. Some of the aliphatic carboxylic acids have been found to give essentially linear dependence of pKa with respect to dioxane content in the dilute dioxane solutions.l 1 The order of relative basic strengths of the derivatives of 1,lO-phenanthroline is as one would predict. The various substituents produce different effects in electron density on the nitrogen atoms depending on both their nature and position, in keeping with their known effects on other ring systems. I n the 2- and 4-positions a phenyl substituent increases the electron density on the ring nitrogens slightly by its tautomeric effect, whereas in the 3and 5-positions its electrophilic effect predominates. The effects of methoxyl groups are similar (8) H. P. Marshall and R. Grunwald, J . Am. Chem. Soc., 76, 2000 (1954). (9) B. Gutbezahl and E. Grunwald, ibid., 76, 559 (1953). (10) L. G.Van Uitert, C . G. Haas, W. C. Fernelius and E. E. Douglas, ibid., 76, 455 (1953). (11) H. 6. Harned and R. B. Owen, "The Physical Chemistry Of Electrolytic Solutions," Rcinhold Publ. Corp., New York, N. Y., 1043, p. 581.
1548
R. B. ANDERSON,W. K. HALL,J. A. LECICY, AND K. C. STEIN
Vol. 60
TABLE I1 to those of the phenyl but more pronounced. Halogen substituents exhibit slight tautomeric and proADDITIVITYOF METHYLSUBSTITUENT CONSTANTS nounced inductive effects. The inductive effects --c of methyl substituents, which do not give the tautoDifference Calcd. Obsd. meric effect, are in accord with previous predictions .... 0.60 .. based on 7~ electron density calculations for 1,lO+ O . 03 .17 0.14 phenanthroline by Longuet-Higgins and Coulson. l 2 .... .56 .. Multiple substitutions of methyl groups on 1,lO. ... .25 .. phenanthroline are found to increase its basicity in 40.06 1.16 1.10 an additive fashion. To illustrate this, use is made .ll 1.20 1.31 of the substituent constant, u, which is defined by .03 0.73 0.76 the expression cr = pKi - QK,,where pK1 is the 4 .02 .73 .71 negative log of the acid dissociation constant in wa.03 * 34 .37 ter of 1,lO-phenanthroline and PKa is that of the .04 .81 .85 substituted derivative. The additivity of the four .09 .98 1.07 individual methyl substituent constants when f .16 1.29 1.13 multiple substitutions are considered is shown in .06 0.98 1.04 Table 11. The calculated values of these four conf .19 .67 0.48 stants were obtained by averaging their individual 4 .18 .59 .41 contributions in all of the methyl derivatives stud+ .I6 .84 .68 ied. The good agreement, for the most part, of the + .01 1.46 1.45 calculated and experimental values indicates that - .05 1.54 1.59 reasonably accurate values for the acid dissociation .06 1.15 1.21 constants of the methyl derivatives of 1,lO-phen+ .68 2.32 1.64 anthroline may be calculated in this way. Acknowledgment.-The authors wish to express Temple University for supplying the comp unds their gratitude to Professor Francis H. Case of reported herein either directly or indirectly through his many outstanding contributions in the synthesis (12) H. C. Longuet-Higgins and C. A. Coulson, J . Chem. SOC.,208, of substituted 1,lO-phenanthrolines. 971 (1949).
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SORPTION STUDIES ON AMERICAN COALS BY ROBERT B. ANDERSON,W. KEITHHALL,JAMES A. LECKYAND KARL C. STEIN Contribution of Central Experiment Station, Bruceton, Pennsylvania Received June 91, 1966
The interpretation of sorption studies on coals haa not led to a clear picture of their physical structure. The sorption of inert molecules such as nitrogen and argon at liquid nitrogen temperatures is low, corresponding to surface areas of 0.5 to 15 m.2/g. The sorption of polar molecules such as water, alcohols and amines a t about room temperature is high, tori-2sponding to surface areas of 50 to 400 m.Z/g. The present data, obtained from sorption of normal and isobutane at 0 , and nitrogen isotherms at temperatures of -195 to 0", indicate that the surface areas of coals have some intermediate value. Density measurements show that coals have a sizable pore volume. Sorption of normal butane was considerably greater than that for isobutane on most coals, indicating that at 0" a large fraction of the pores have openings of about 5 A. Adsorption of nitrogen at -195" was low indicating that at this temperature the pore openings are smaller, the order of 4 A. or less. The sorption of polar molecules is apparently complicated by swelling and imbibition, involving weak chemical bonds between the sorbate and polar groups in the coal.
Coal is an amazingly complex substance, and that have not yet been settled. In addition to the physical and chemical studies have not led to any British group, researchers in South Africa6-' and simple, useful structure of coal. The interpreta- the NetherlandsQ have contributed valuable extion of sorption studies is equally difficult. Work of perimental data and interpretation. The controthe British Coal Utilisation Research Association1-6 versy resulted from the interpretation of the obbeginning about 1940 has produced a large amount servations that (a) the sorption of inert molecules of pertinent data for sorption on coal; however, the such as argon and nitrogen at liquid nitrogen teminterpretation of these data has led to polemics peratures is relatively low corresponding to surface areas of 0.5 to 15 m.2/g., and (b) the sorption of (1) D. H. Bangham, R. E. Franklin, W. Hirst and F. A. P. Maggs, Fuel, 28, 231 (1949). polar molecules such as alcohols and amines at (2) D. € Bangham, I. F. A. P. Maggs, Proceedings of a Conference about room temperature is high, corresponding to on the Ultrafine Structure of Coals and Cokes, The British Coal Utilisasurface areas of 50 to 400 m."/g. The British group tion Association, London, 1944, pp. 118-130. (3) R. E. Franklin, Trans. Faraday SOC.,46, 274 (1949). (4) M. G r i 5 t h and W. Hirst, Proceedings of a Conference on the Ultrafine Structure of Coals and Cokes, T h e British Coal Utilisation Aasociatinn, London, 1944, pp. 80-94. (5) P. A. P. Maggs, rel. 4. 1944, pp. 9g-109.
(6) P. Le R. Malherbe, FueE, 80, 97 (1951). (7) P. Le R. Malherbe and P. C. Carman, ibid., 81, 210 (1952). (8) P. Zwietering, A. P. Oele, and D. W. van Krevelen, ibid., 80, 202 (1951). (9) P. Zwietering and D. W. van Krevelen, ibid., 88, 331 (1954).
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