Table 11. Nitrogen by Modified Dohrmann Dohrmann No. of End point, Analysis ppm Kjeldahl “C Sample type Processing 3 0.28 i 0.02 0.25 product 445 Processing 15 11.5 i 0 . 4 12 product 520 Processing 3 0.19 i. 00 0.21 product 545 Processing 4 0.08 i. 0.02 540 1 1715 1740 (1 wt S) Vacuum bottom 1 1730 1740 >540 Lubricating .1 352 340 oil blend .1 19 16 Butene gas Olefin-amine 4 1160 =t18 1080 copolymer MW >lO,oOO
erally fall between +lox of Kjeldahl. With no sampling problems, the reproducibility and accuracy of the KjeldahlIndophenol method for most petroleum samples above 10 ppm nitrogen are about =tlOz. Two types of petroleum material did not give good recovery. Crude oils gave nitrogen results u p t o 35z lower than that obtained with the Kjeldahl method. Asphaltenes, which were analyzed by dispersing in benzene, gave very poor recovery. When the sample was injected, a fine spray of black particles was noticed in the outlet. Evidently the solvent flashed from the dispersed asphaltene particles, and they passed essentially unreacted through the catalyst section. High level samples (>150 ppm) were determined after dilution in spectroquality benzene. Isooctane was originally tried, but a number of samples precipitated after a short period of time. When some heavier type samples, including carbazole standards, were dissolved in benzene or isooctane and left standing out in light, the nitrogen level as determined by the Dohrmann microcoulometer decreased with time. Evidently, light catalyzes a photochemical reaction is which nitrogen is converted to a species which might absorb on the glass walls of the container as they are formed.
cluded vacuum bottom cuts, atmospheric residua, solventdeasphalted oils, lubricating oil blends, hydrofined and partially hydrofined stocks and products, etc. The results gen-
RECEIVED for review September 30, 1970. Accepted January 8, 1971. Presented at the Division of Petroleum Chemistry, 155th National Meeting, American Chemical Society, San Francisco, Calif., April 1968.
z
Solvent Effects on Protonation Constants Ammonia, Acetate, Polyamine, and Polyaminocarboxylate Ligands in Methanol-Water Mixtures D. B. Rorabacher, W. J. MacKellar, F. R. Shu, and Sister M. Bonavita Depcrrfment r$ Chemistry, W a y n e Sttrte Unicersity, Detroit, Miclz. 48202 The protonation constants for ammonia, acetate, and eight multidentate polyamine and polyaminocarboxylate ligands have been determined as a function of solvent composition in methanol-water mixtures containing 0-99% methanol (by weight). The data have been obtained by means of potentiometric titrations using conventional glass and calomel electrodes to measure pH*, the nonaqueous equivalent to pH as referenced to the standard state in the identical solvent composition. I n agreement with previous observations, the protonation constant for ammonia passes through a minimum in the region of 6 5 7 0 % methanol (wt/wt) while acetate exhibits a continually increasing value with increasing methanol content. The amine nitrogen and carboxylic oxygen donor atoms i n the multidentate ligands exhibit corresponding behavior with slight deviations apparent for the nitrogen atoms as a result of increasing substitution and charge effects. The general behavioral patterns are interpreted in terms of electrostatic effects, base solvation, and proton solvation as a function of the solvent composition.
THEPOTENTIAL ADVANTAGES of utilizing nonaqueous solvents for improving analytical methods involving complexation have been noted previously ( I ) . However, the lack of reliable data for protonation and complex stability constants in other (1) A. Ringboni, “Complexation in Analytical Chemistry,” Interscience, New York, N. Y . , 1963, p 14.
than aqueous media has largely hindered the exploitation of nonaqueous approaches except on an occasional empirical basis. I n connection with our current interest in solvent effects (2, 3), we have undertaken the measurement of protonation equilibria for a number of analytically important ligands, principally of the polyamine and polyaminocarboxylate families, in methanol-water solvents ranging from pure water to 99% methanol (by wt). I n so doing we have taken advantage of recent advances in the understanding of these solvent mixtures, including the establishment of thermodynamically valid pH scales. Thus the protonation constants reported are of equal validity to those normally utilized in aqueous solutions. The choice of methanol as the nonaqueous solvent component for this study was based on both experimental and theoretical considerations. Alcohols closely resemble water in the nature of their protolytic behavior and are readily amenable to study without specialized techniques. Yet these (2) D. B. Rorabacher and F. R. Shu, unpublished data; W. J. MacKellar and D. B. Rorabacher, J . Amer. C/zem. Soc., in press. (3) D. B. Rorabacher, F. R. Shu, W. J. MacKellar, and R. W. Taylor, Abstract 127, Proceedings of the XIIIth International Conference on Coordination Chemistry, Cracow-Zakopane, Poland, September 1970. ANALYTICAL CHEMISTRY, VOL. 43, NO, 4, APRIL 1971
561
solvents also provide definite contrasts to water in their physical properties. I n addition, sufficient physical data are available for methanol and methanol-water solvents to permit acidity measurements to be made which are of equal thermodynamic validity to those commonly made in pure water. The multidentate ligands studied include three polyamines: ethylenediamine (en), triethylenetetramine (trien), and tetraethylenepentamine (tetren); four aminocarboxylates: glycinate ion (gly), ethylenediamine-N,N,N‘,N’-tetraacetate ion (EDTA), trans-1 ,2-diaminocycloh~xane-N,N,N’ ,“-tetraacetate ion (CDTA), and ethyleneglycol-bis-(fl-aminoethylether)-N,N,N’,N’-tetraacetate ion (EGTA); and a related poly(amino alcohol): N,N,Nr,N’-tetrakis-(2-hydroxyethyl)ethylenediamine (TKED). In addition, the protonation constants of ammonia and acetate ion were determined for comparison with values reported by previous workers. In this way the validity of the current method is tested and information is obtained pertaining t o the manner in which solvent composition affects base protonation. The results show distinct behavioral trends for the two types of donor atom undergoing protonation in accord with previous observations. For carboxylic oxygens, continually increasing protonation constants are observed with increasing methanol content amounting to a total change of about 4-5 log K units over the entire solvent range. By contrast, the amine nitrogen donor atoms exhibit a minimum value in the region of 70-80Z methanol. These phenomenological observations are shown to be consistent with electrostatic and solvation effects. The term “nonaqueous solvent,” as used throughout this paper, refers to any solvent other than pure water. EXPERIMENTAL
cipitation and recrystallization of the perchlorate salt as described previously (5). Tetren was precipitated as the acid sulfate salt (6) and recrystallized twice. Because of the low solubility of this salt in the methanolic solvents, the colorless neutral ligand was reconstituted by neutralization with concentrated aqueous NaOH solution followed successively by azeotropic distillation with benzene to remove the water, distillation of the benzene itself, and a vacuum distillation of the tetren (135 “C, 0.6 mm Hg). Trien was obtained as the reagent grade sulfate salt and neutralized with concentrated NaOH, the neutral trien separating out as a discrete layer upon standing overnight in the presence of saturated Na?S04. The trien layer was removed and vacuum distilled (104 “C, 1 mm Hg). Ethylenediamine was purified by fractional distillation. The sample of glycine (Sigma Chemical c o . ) was of unspecified purity but no attempt was made at further purification. The concentrations of the ligand solutions were determined either directly from the potentiometric titrations used to determine the protonation constants or by quantitative dilution of previously standardized solutions. Sodium perchlorate (G. F. Smith Chemical Co.) and tetraethylammonium perchlorate (Eastman Organic Chemicals) were recrystallized before use. Absolute methanol assayed at 99.98% (0.02% HrO) as obtained from the J . T. Baker Chemical Co. was used for preparing all solutions by diluting with a calculated amount of conductivity water to obtain the desired solvent composition, reported in all cases as per cent by weight of methanol. For solutions containing 95 methanol (wtjwt) or greater, the water content was routinely checked using the Karl Fischer titrimetric method. Apparatus. All titrations were made using an expanded scale p H meter (either a Sargent Model DR or a Corning Model 12) incorporating the electrode system:
i
Agar 1 Ag;AgCl, HCI:,Glass:;Solnin C H 3 0 H - H r 0 Bridge’ KCl (satd in HrO),HgCl?, Hg
(1)
Titrants. All protonation constants were determined by means of potentiometric acid-base titrations of the ligand. For basic species (TKED and the polyamines) perchloric acid was used as the titrant. For acid species (NH4+, CH,COOH, EDTA, CDTA, and EGTA) either sodium hydroxide or tetraethylammonium hydroxide was the titrant. I n the singular case of glycine, both an acid and a base titrant were used to titrate separate samples. Perchloric acid solutions were prepared by diluting 73.6 % HCIOl (G. F. Smith Chemical Co.) with methanol and adjusting the solvent composition with water followed by standardization against standard sodium methoxide solution. The water contained in the 73.6% HClO4 restricted the maximum methanol content of the resultant solution to about 99.5% methanol (by wt) for 0.10M HC10, and n o attempt was made to achieve more anhydrous solutions [although such methods have been described elsewhere (41. Because of the large carbonate contamination found in commercial sodium methoxide, solutions were prepared by reacting sodium metal with methanol and diluting with water to obtain the desired solvent composition. (Depending on the solvent composition, methoxide may convert partially or completely to hydroxide, according to the reaction: CH30H 2 0 -+ C H 3 0 H OH-.) Standardization was by titration against primary standard grade potassium acid phthalate. Tetraethylammonium hydroxide was obtained as a 10% solution in water (Eastman Organic Chemicals) and was utilized without further purification, being used primarily for comparative purposes. Reagents. Acetic acid, ammonium perchlorate, EDTA, CDTA, and EGTA were obtained in reagent grade and used without further purification. T K E D was purified by pre-
As indicated, the reference electrode (Leads & Northrup 491 3-D10 aqueous saturated calomel electrode) was impregnated at the liquid junction with a layer of agar KCl to prevent flow of the internal aqueous KC1 solution into the external nonaqueous solvent. Of the glass electrodes tested, the Beckman 41263 pH 0-1 1 electrode was found to have the fastest response time in the methanolic solvents and was used preferentially. For titrations requiring measurements at high pH*, any one of the following p H 0-14 glass electrodes was used: Beckman 19002, Corning 476034, or Sargent S-30050-15. To minimize dehydration of the glass membrane, glass electrodes were soaked in water when not in use and occasional checks were made to ensure that linear response was maintained over the entire pH* range of interest. Once an electrode showed evidence of nonlinear response, it was immediately discarded since none of the recommended methods for re-constituting “poisoned” electrodes was found to be effective. Prior to commencing a titration, electrodes were “conditioned” by soaking in a buffer of the desired solvent cornposition until a stable response was obtained (generally achieved within a few minutes to an hour-longer times being required for solvents with lower water content). Conditions. Temperature was controlled at 25.0 + 0.1 “C by inserting the titration vessel into a snugly fitting hollow brass jacket through which water was circulated from a Sargent “Thermonitor” Water Bath. To prevent C 0 2 absorption, the vessel was covered and the surface of the solution was continuously swept with nitrogen which had first been bubbled successively through two towers containing solvent
(4) C. D. Ritchie and P. D. Heffley, J. Amer. CJTem. Soc., 87, 5402
( 5 ) D. B. Rorabacher, T. S. Turaii, J. A . Defever, and W. G. Nickels, Imrg. Chem., 8, 1498 ( I 969). (6) C. N. Reilley and A . Vavoulis, ASAL. CHEW.31, 243 (1959).
+
+
(1965). 562
ANALYTICAL CHEMISTRY, VOL. 43, NO. 4, APRIL 1971
of identical composition to the titration solution in order to facilitate saturation of the gas with the solvent vapor and thereby minimize evaporation. Ionic strength was controlled at 0.10M using sodium perchlorate or tetraethylammonium perchlorate. For all ligands other than the polyaminocarboxylates, identical protonation constant values were obtained with either salt. For the titrations of EDTA, EGTA, and CDTA the presence of sodium perchlorate caused a significant decrease in the apparent value of the first protonation constant as a result of sodium complex formation. The extent of such complexation was quantitatively investigated only for EDTA in the aqueous solvent along with additional studies using LiC104 and K N 0 3 . In the nonaqueous solvents the de-ision was made t o utilize the sodium salt based o n the practical utility of these measurements t o other studies. Qualitatively, it appears that the error thus generated in the first protonation constant increases with increasing methanol while successive protonation constants are essentially unaffected. MEASUREMENT OF ACIDITY I n principle, any consistent reproducible scale for measuring acidity may be utilized for the determination of relative protonation constants. However, for a meaningful interpretation of such values, it is desirable to utilize an acidity scale which has thermodynamic significance. A survey of previous studies in which protonation constants have been measured in methanol-water mixtures (7) (nearly all of which were limited to the determination of a single protonation constant per base) reveals several novel techniques which have been used to circumvent real or imagined difficulties in measuring nonaqueous pH. Historically, attempts at measuring p H in nonaqueous solutions have met with the inevitable controversy surrounding the difficulties of establishing p H scales in such solutions which are referenced to the standard state of hydrogen ion activity in water. Many suggestions have been made relative t o establishing a “universal scale of acidity” which would presumably include the effect of the medium upon electrode response. A discussion of the relative merits of this approach is given in a recent review by Popovych (8). For the present purpose there is no advantage to be gained by referring all acidity values to an aqueous standard as long as the values are internally consistent and, in fact, equilibrium and kinetic measurements are more significantly related t o the hydrogen ion activity for each solvent composition. Therefore, the operational pH* scales established by deLigny and coworkers for methanol-water mixtures ( 9 ) , and later elaborated by Bates and coworkers (IO),in which the hydrogen ion activity is referred to the standard state in the same solvent composition, were chosen for utilization in this study. deLigny has defined the function paH* as the nonaqueous equivalent of pari paE*
=
-logaH*
=
-logmHyH*
(2)
where ari*, m ~and , yH*represent the molal activity, molality, and activity coefficient, respectively, for the solvated proton in (7) R. G. Bates and R. A. Robinson in “Chemical Physics of Ionic Solutions,” B. E. Conway and R. G. Barradas, Ed., John Wiley and Sons, New York, N. Y., 1966, p 211 ff; and references
therein.
(8) 0. Popovych, Crit. Rec. Ami/. Chrm., 1, 73 (1970). (9) C . L. deligny, P. F. M. Luykx, M. Rehbach, and A. A. Weinecke, R e d . Truc. Chim. Puys-Bns, 79, 699 (1960). (10) R. G. Bates, M. Paabo, and R. A. Robinson, J . Phys. Chem., 67,1833 (1963).
the specified solvent mixture, the starred quantities being referred to the nonaqueous standard state. Since paH* cannot be measured accurately for most solutions, deLigny and coworkers have rigorously determined the pari* values of standard buffer solutions in a manner analogous to the method applied by the National Bureau of Standards for aqueous solutions ( I I ) . Operationally the pH* value of a n unknown solution, designated as pH,*, is then determined in the normal manner by comparing the measured potential value, E,*, with that obtained for a standard buffer solution, Es*, of known pH,* in an identical solvent composition, Liz., (for 25 “C), pH,* = pH,*
- E,* + E,*0.05916
(3)
~-
Bates, Paabo, and Robinson (10) have shown that the potential across thej unction between alcohol-water solutions and a saturated calomel electrode, although large, is as reproducible as that at the junction between aqueous solutions and the salt bridge so that the measured pH,* values, obtained with the electrode system shown in Equation 1, closely approximate palr* under optimum conditions of measurement in the same way that aqueous p H approximates paIi. This approximation has been assumed to be valid throughout this study. In standardizing the p H meter for a specific solvent composition, the oxalate buffer (0.01M oxalic acid 0.01M ammonium hydrogen oxalate) and succinate buffer (0.01M succinic acid 0.01M lithium hydrogen succinate), prepared in the same solvent composition, were set to the appropriate pH,* values as interpolated from deligny’s established values for these systems (12). Deviations up to 0.04 unit from the interpolated pH,* values were tolerated. Although other buffer systems have since been established for methanol-water mixtures (IO), the pH,* values have not been rigorously defined for solvents exceeding 50z methanol. Thus, in this work the oxalate and succinate buffers were utilized in all solvent compositions for consistency. As an additional check these nonaqueous buffers were generally compared with meter readings obtained using standard aqueous buffers and the deviations thus obtained were compared with deligny’s published “6” values (13)
+
+
pH*
=
pH - 6
(4)
where 6, as identified by Bates, Paabo, and Robinson (IO), represents the difference between the liquid junction potential, E, (in pH units), and the medium effect, log,nyrr(8),Liz.,
The titration values of pH,*, hereinafter called pH*, were then read directly from the meter. RESULTS
Resolution of Protonation Constants. The following symbols are used to represent the designated quantities and constants: [HILI, [Hj-lL], [H+] = molar concentrations of the conjugate acid, conjugate base, and solvated proton, respectively. (11) R. G. Bates, "Determination of pH,” John Wiley and Sons, New York, N. Y., 1964. (12) C. L. deligny, P. F. M. Luykx, M. Rehbach. and A. A. Weinecke, Rrcl. Trcic. Chim. Pciys-Bus, 79, 713 (1960). (13) W. J . Gelsema, C. L. deligny, A . G. RemljnSe, and H. A . BhJlewI, ibid., 85, 647 (1966). ANALYTICAL CHEMISTRY, VOL. 43, NO. 4, APRIL 1971
563
~~
~
~~
~
~~
Table I. Physical Data for Methanol-Water Solvents Utilized in This Work (25 “C) PH,*~oxalate pH,*e succinate w t % C H ~ O H Ti* ( g = 0 . 1 ~ ) 5 pKs** d2o ‘0‘ Dd buffer buffer 0 0.796 14.00 1 .oooo 78.5 2.15 4.12 25 0.747 14.06 0.9608 67.0 2.27 4.57 50 0.678 14.10 0.9169 55.3 2.47 5.07 0.8849 48.5 2.66 5.41 65 0.637 14.17 0.8483 41.6 3.13 6.01 80 0.590 14.42 14.84 0.8218 37.0 3.73 6.73 90 0.550 95 0.514 15.41 0.8078 34.7 4.23 7.26 0.7964 32.9 5.20 8.23 99 0.464 16.32 32.4 5.79 8.75 16.70 0,7934 100 0.427 Extrapolated from data in Ref. 14 which extend only to 0.10 molal. b Interpolated from data in Ref. 16. c Cited in “Handbook of Chemistry,” N. Lange, Ed., 10th ed., McGraw-Hill Book Co., New York, N. Y., 1961, p 1182 ff, d Interpolated from the combined data of G. Akerlof, J. Amer. Chem. SOC., 54,4125 (1932); P. S.Albright and L. J. Gosting, ibid., 68, 1061 (1946). e Interpolated from data in Ref. 12. Q
aHjL*,
aHjW1L*,
aR*
=
molal activity of the corresponding
species. KHjC* = concentration protonation constant (units: molar-’) as defined by the equation
KHjm* = mixed-mode protonation constant (units :molal-’), convenient for application to measured pH* values, as defined by the equation
data are readily amenable t o treatment in this form and the resulting values can be directly applied to subsequent pH* measurements. As reported in this paper, these constants were resolved by the Bjerrum half-n-bar method (17)modified to allow the use of all experimental titration points in the calculation. For this purpose the Bjerrum RH function, representing the average number of protons bound per ligand molecule
(9)
where d represents the specific gravity of the mixed solvent (Table I). KHjo* = activity protonation constant (units :molal-’), defined by the relationship
y** = mean activity coefficient extrapolated t o 0.10M
ionic strength in each solvent composition from Oiwa’s experimental values (14) determined using HCI solutions (Table I). [Oiwa’s values, which extend t o a maximum ionic strength of 0.1 rn, were chosen for application in this work since they represent the most consistent set of experimental values available for the entire solvent range from pure water to pure methanol; furthermore, a critical comparison with other published activity coefficient values (7, 15) indicates that Oiwa’s values are as reliable as any others available over extended solvent ranges. J K,, K,* = activity solvent autoprotolysis constants (units : molal2) in pure water and methanol-water solvents, respectively, as interpolated from experimental data reported by Koskikallio (16) (Table I). Starred symbols, in general, refer to nonaqueous equivalents of the correlated aqueous quantities, referenced to the corresponding nonaqueous standard state. All protonation constants were initially computed as mixedmode constants, defined by Equation 7, since the experimental (14) I. T. Oiwa, J. Phys. Chem., 60, 754 (1956). (15) “Electrochemical Data,” B. E. Conway, Ed., Elsevier Publishing Co., Amsterdam, 1952; and references therein. (16) J. Koskikallio, Suom. Kemistilehti B, 30, 111 (1957). 564
ANALYTICAL CHEMISTRY, VOL. 43, NO. 4, APRIL 1971
can be related to terms which are experimentally accessible. F o r titrations involving the addition of acid to the unprotonated ligand, this relationship can be expressed as
where CH represents the total molar concentration of acid added and CLrepresents the total concentration of ligand in solution. In titrations where a protonated form of the ligand was titrated with standard base, Equation 10 can be applied by redefining C, as equal t o CA - CoH where C, represents the total initial molar concentration of acid and CoH represents the total molar concentration of base added. Alternatively, one may utilize the Q function defined as the degree of neutralization of the original protonated ligand, n
c[HJl
L L
j=O
where n represents the number of titratable protons on the ligand species present at the start of the titration (N. B., Q = n - fia; C A = nCL). Operationally, the values of [H+] and [OH-] were calculated from the relationships
(17) F. J. C . Rossotti and H. Rossotti, “The Determination of Stability Constants,” McGraw-Hill Book Co., New York, N. Y., 1961.
Table 11. Protonation Constants for Acetate Ion in Methanol-Water Solvents at 25 "C, p = 0.1M (Et4NC10r) log KHO* log K H ~ * Wt % This This Ref. 18 Ref. I 9 Ref. 20 Ref. 21 CH30H work work 4.768 4.756 0 4.63 4.73 4.91 4.90 10 5.011 16.55 5.08 5.08 20 25 5.05 5.18 5 334 34.59 5.45 40 50 5.55 5.72 5,808 53.99 5.90 60 65 5.84 6.04 6.500 76.03 6.64 80 6.34 6.57 7.31 90 7.10 7.36 7.858 93.78 8.00 95 7.46~ 7,750 99 8.69 9.02 -99.8 9.07 9.34 9.72 9.52 100 a Values in 95 methanol appear to be low. I
[OH-]
K,*d
= ___
as*y**
(13)
where d i s introduced to correct the experimentally determined uIi* values (aR* = 10-PH*) from molal to molar units, and -y+* is used to approximate yH*. For species involving a single protonation or neutralization step (NH4+, CHsCOOH, neutral glycine), computer calculation (using all titration points) of the mixed-mode protonation constant, KITm*,was facilitated by combining Equations 7 and 9 and converting t o the form
(14) which is equivalent to plotting the logarithmic form of Equation 7 (the so-called Henderson-Hasselbalch equation)
(7') The Ksm* values for ammonia and acetate were converted to the corresponding activity constants, KHo*, by using the extrapolated values of the mean activity coefficient, -y+* (Table I), to approximate the values of ~ N H ~ +and * -~CH~COO-* while assuming that ~ B E I ~and * ~ c ~ ~ ~ were c o o unity. ~ * The least squares values of the experimental mixed-mode protonation constants for acetate ion and ammonia and the corrected activity constants are listed in Tables I1 and 111, respectively. For comparative purposes, values obtained by other workers for these ligands are also listed. It should be noted that, in attempting to determine the protonation constants for ammonia, a general lack of reproducibility was encountered. Similar difficulties with ammonia solutions have been reported by several other workers. (18) T. Shedlovsky and R. L. Kay, J . Phys. Clzetn., 60, 151 (1956). (19) A . L. Bacarella, E. Grunwald, H. P. Marshall, and E. L. Purlee, J . Org. Cliern.. 20, 747 (1955). (20) H. S. Harned and N. D. Embree,J. Amer. Chem. Soc., 57, 1669 (1935). (21) L. J. Minnick and M. Kilpatrick. J . Phys. Cliem., 43, 259 (1939).
Table 111. Protonation Constants for Ammonia in Methanol-Water Solvents at 25 "C, p = 0.1M (NH4CI04) log KRo* log KH')t* Paabo, Bates:' Wt This This and CH30H work work Robinson (22) 0 9.31 9.21 9,2456 9.146 10 9.014 20 25 -9.16 -9.03 33.4 8.893 50 8.87 8.70 8,687 65 8.86 8.66 70 8.571 8.65 80 [8. 93Ic 90 9.32 9.06 95 9 ,67d 9.38d 99 I O . 37 10.01. -99,8 -10.72 -10.36 a Values shown were calculated from experimental mixed-mode constants for p = 0.10M in Ref. 22 using extrapolated y=* values (Table I). R. G. Bates and G. D. Pinching, J. Res. Not. Bur. Sid., 42, 419 (1949). e Experimental K t p * value shown was determined at p = 0.25M; KIT"*value was calculated using y= value for p = 0.25. Values in 95 methanol appear to be low.
z
Everett and Wynne-Jones, in studying ammonia and several substituted amines in 6 0 x methanol, noted that errors were greatest with the former compound and suggested this might be due to the absorption of CO? (23). Bates, Paabo, and Robinson suggested that observed deviations in methanolwater solvents might be due to incomplete dissociation of the ammonium salts (10). Other workers in experiencing similar problems in aqueous solution have surmised that significant amounts of ammonia are lost by evaporation or that the amines interact with the electrodes. Although no attempt was made in this study to pinpoint (22) M. Paabo, R. G. Bates, and R. A . Robinson, J . Plrys. Chem., 70, 247 (1966). (23) D. H. Everett and W. F. K. Wynne-Jones, Trom. Fnradcry Soc., 48, 531 (1952). ANALYTICAL CHEMISTRY, VOL. 43, NO. 4, APRIL 1971
565
Table IV.
Mixed-Mode Protonation Constants for Polyamines and a Poly(amino alcohol) in Methanol-Water Solvents at 25 "C, I-( = 0.1M (All values in molal-1) Wt % Methanol Ligand log K H , ~ " * 0 % 25 % 50 % 65 % 80 % 90 95 % 99 % 9.97 9.30 9.11 en KI 9.99 9.16 9.72 9.80 10.73 6.42 6.32 7.01 6.39 6.77 Kz 7.31 6.89 7.64 trien K1 10.09 9.86 9.36 9,3@ 9,73h ... 10.37 11.25 Kz 9.31 8.95 8.54 ... 8.63 8.87b 9.55 10.36 K; 6.75 6.44 6.10 5.99a 6 . 2jh ... 6.80 7.56 K4 3.39 -2.81 2. 276 ... -2.80 2.13 -2 I05Q 3.33 tetren KI 10.36 10.03 9.83 9.77 9.80 10.23 10.67 11.36 Kz 9.65 9.31 9.07 9.49 9.93 9.13 9.06 10.60 8.16 8.11 8.62 9.77 K3 8.50 8.28 8.25 9.01 K4 4.70 4.33 4.00 3.20 3.76 3.89 4.14 4.78 K6 2.40 2.01 1.86 (2,24) 1.96 2.08 2.26 (2.13) TKED KI 8.38 8.52 8.24 8.02 7.77 8.14 8.41 9.27 3.44 3.14 (2.75) 2.94 Kz 4.37 3.90 3.48 4.50 Values designated are for 70% CH,OH. Values designated are for 85 % CH,OH.
z
~
-99 8 % 11.10 8.18 ... ...
...
... ... ...
... I .
... ... ...
Table V. Mixed-Mode Protonation Constants for Aminocarboxylates in Methanol-Water Solvents at 25 "C, p = 0.10M(NaC104) (All values in molal-l) Ligand log KIT,%* 0% 25 % 50% 65% 80% 90% 95 99 -99.8% glY KI 9.65 9.65 9.26 9.15 9.30 9.83 10.09 10.97 11,35 KP (2.19) ... (2.93) 3.31 3.89 4.46 4.76 5.71 6.49 EDTA Kl 9.4W 9.27a 8.69 8.43a 8,30a ... 9.1w 9.9lQ ... K2 6.08 6.36 6.32 6.45 6.77 ... 8.28 8.88 K B 2.69 3.43 3.80 4.23 5.03 ... 6.78 7.53 ... K4 (1.57) (2.14) (2.70) (3.35) 3.85 , . . 5.70 6.60 ... EGTA KI 9.3P 9.24a 8.47& 8.3ga 8.1Y ... 8 . 6aa 9 , 25a ... Kz 8.71 8.61 7.79 7.74 7.60 ... 7.87 8.48 ... K3 2.77 2.92 3.02 (3.67) 4.19 ... 5.35 6.32 ... K4 ... 2.57 (2.15) ... 3.30 4.62 5.62 ,.. CDTA KI . . ~ 9.45arb 8.82azc ... 8.75a 7,52" 9.9@ ... Kz ... 6.17b 5.92c ... 6.42 ... 6.58 7.71 ... K3 ... 3.96b 4.03c ... 5.04 ... 5.17 6.67 ... K4 ... 2.93b 3.0Y 3.81 ... 4.50 5.53 ... Values reported were determined in the presence of 0.10M Na+, thus yielding low values for K H m* due to Na+ complexation. Values designated are for 20% methanol. Values designated are for 40% methanol.
z
z
I
I
.
.
.
.
. . I
.
(1
the source of error, a similar lack of reproducibility was observed for titrimetric data obtained using dilute solutions of ammonia such that apparent deviations as large as 0.3 unit occurred in some solvent compositions among the log K H ~ * values obtained by repetitive titrations. This problem was finally minimized by utilizing solutions of greater concentration (0.25M NH4C104in 80% CHBOH; 0.10M NHICIOI in all other solvent compositions) for the titrations. For species involving two or more stepwise proton additions, the several protonation constants were calculated using a modification of the Bjerrum half-n-bar method essentially identical to that applied by Jonassen and coworkers for calculating stepwise acid dissociation constants for protonated polyamines (24). As redefined for mixed-mode protonation constants, the applicable equations for a base capable of adding n protons under the experimental conditions are
(24) H. B. Jonassen, F. W. Frey, and A. Schaafsma, J . Pliys. Cliem., 61, 504 (1957). 566
ANALYTICAL CHEMISTRY, VOL. 43, NO. 4, APRIL 1971
I
.
where Equation 15b is applicable for the case 1
