Anal. Chem. 1999, 71, 3140-3144
Noncomplexing Tertiary Amines as “Better” Buffers Covering the Range of pH 3-11. Temperature Dependence of Their Acid Dissociation Constants Ashoka Kandegedara and David B. Rorabacher*
Department of Chemistry, Wayne State University, Detroit, Michigan 48202
Most of a broad series of buffers developed by Good (often called “Good’s” or “Good” buffers) have been shown to cause metal ion interference as a result of complexation. A series of tertiary amines, devoid of hydroxy groups or other weak donors on the r, β, or γ carbons, have been developed as “Better” pH buffers which, as a result of steric hindrance, are incapable of forming even weak complexes with metal ions. As a result, they avert interferences of the type often encountered in metal ion studies which require buffer use. The compounds studied are alkyl and alkylsulfonate derivatives of morpholine (3), piperazine (3), ethylenediamine (3), and methylenediamine (1). With the exception of the morpholine derivatives, these compounds have two protonatable sites and, therefore, can be used to buffer two separate pH regions. A series of 10 buffer compounds span the entire range of pH 3-11. The acid dissociation constants for these compounds have now been determined in aqueous solution at 15, 25, 35 and 45 °C, µ ) 0.10 M. From these data, apparent enthalpy and entropy values have been calculated to permit application at other temperatures as well. These buffers are recommended for use in all situations in which metal ions are present. All of the reported compounds are now commercially available. Studies involving metal ions in solution which require pH control are inevitably subject to the possibility of buffer interference as a result of complexation. Recently, we reported the development of a series of sterically hindered tertiary amine buffers1 which can be used to cover the entire pH range from 3 to 11 while avoiding metal ion complexation. In our original paper, the acid dissociation constants were reported only for 25 °C. Since data applicable to other temperatures would be useful to many investigators, we now wish to report the acid dissociation constants for all nine of the original buffers, plus that for one additional compound, as determined at 15, 25, 35, and 45 °C and an ionic strength of 0.10 M. The apparent ∆H° and ∆S° values associated with these acid dissociation constants have been determined so * Corresponding author. E-mail:
[email protected]. Fax: (313) 577-1377. Tel.: (313) 577-2605. (1) Yu, Q.; Kandegedara, A.; Xu, Y.; Rorabacher, D. B. Anal. Biochem. 1997, 253, 50-56.
3140 Analytical Chemistry, Vol. 71, No. 15, August 1, 1999
that pKa values appropriate to other temperatures may be calculated as well. Background. The problem of buffer interference in metal ion studies arises from the fact that all Brønsted bases are, by definition, also Lewis bases. Thus, the possibility of buffer complexation is an insidious problem whenever metal ion phenomena are to be studied at controlled pH values. Several approaches, with varying degrees of success, have been reported for circumventing this problem. Good and co-workers were among the first to tackle this problem in a methodical manner.2-4 In generating a series of more than 20 buffers for use in biological studies, these workers listed the prevention of metal ion complexation as one of their desired goals. Good’s buffers (often called “Good” buffers) have, in fact, been widely adopted and at least one major chemical supplier devotes a special section of its catalog to them.5 Yet, there have been increasing reports of buffer interferences when using most of Good’s buffers in the presence of metal ions, especially copper.6-9 In an early study,10 we demonstrated that tertiary amine nitrogens, in which all substituents were ethyl or larger, were incapable of forming initial coordinate bonds to metal ions in aqueous solution (see Figure 1A). However, correlated mechanistic studies11-14 have demonstrated that metal ion complexation can occur whenever these sterically hindered tertiary amines (2) Good, N. E.; Winget, G. D.; Winter, W.; Connolly, T. N.; Izawa, S.; Singh, R. M. M. Biochemistry 1966, 5, 467-477. (3) Good, N. E.; Izawa, S. Methods Enzymol. 1972, 24, 53-68. (4) Ferguson, W. J.; Braunschweiger, K. I.; Braunschweiger, W. R.; Smith, J. R.; McCormick, J. J.; Wasmann, C. C.; Jarvis, N. P.; Bell, D. H.; Good, N. E. Anal. Biochem. 1980, 104, 300-310. (5) Sigma Chemical Company Catalogue; Sigma Chemical Company: St. Louis, MO, 1999; pp 1910-1917. (6) Nakon, R.; Krishnamoorthy, C. R. Science (Washington, D.C.) 1983, 221, 749-750. (7) Gregory, J. D.; Sajdera, S. W. Science (Washington, D.C.) 1970, 169, 9798. (8) Lleu, P. L.; Rebel, G. Anal. Biochem. 1991, 192, 215-218. (9) Kaushal, V.; Barnes, L. D. Anal. Biochem. 1986, 157, 291-294. (10) Turan, T. S.; Rorabacher, D. B. Inorg. Chem. 1972, 11, 288-295. (11) Rorabacher, D. B.; Turan, T. S.; Defever, J. A.; Nickels, W. G. Inorg. Chem. 1969, 8, 1498-1506. (12) Rorabacher, D. B.; Moss, D. B. Inorg. Chem. 1970, 9, 1314-1318. (13) Moss, D. B.; Lin, C.-T.; Rorabacher, D. B. J. Am. Chem. Soc. 1973, 95, 5179-5185. (14) Cooper, T. H.; Mayer, M. J.; Leung, K.-H.; Ochrymowycz, L. A.; Rorabacher, D. B. Inorg. Chem. 1992, 31, 3796-3804. 10.1021/ac9902594 CCC: $18.00
© 1999 American Chemical Society Published on Web 06/25/1999
Figure 1. (A) Schematic diagram illustrating the steric inability of an aquated metal ion (solid atom) to bond directly to a tertiary amine nitrogen (shaded atom) when all substituents attached to the nitrogen are ethyl groups (or larger). (B) Schematic diagram showing an aquated metal ion forming an initial coordinate bond to a weak donor atom attached to an R, β, or γ carbon of a tertiary amine, thereby facilitating subsequent ring closure to the nitrogen atom. In these figures, the buffer compound shown is HEPES, a common buffer of the Good series which has been shown to complex copper(II) ion (ref 1). Hydrogen atoms are omitted for clarity. The diagonally striped atom is the sulfur of the sulfonate group and the horizontally striped atoms are oxygens (alcoholic and sulfonate oxygens on HEPES and water oxygens attached to the metal ion).
Figure 2. Noncomplexing tertiary amine buffer compounds discussed in this work.
(which dominate in Good’s series of buffers) contain weakly coordinating groups, such as -OH, on the R, β, or γ carbons. Such weakly coordinating donor atoms permit the metal ion to gain a “toehold” on the buffer which then permits access to the tertiary amine nitrogen to complete a chelate ring. This latter mechanism is demonstrated in Figure 1B for a hexaaquametal ion reacting with HEPES [N-(2-hydroxyethyl)-N′-(2-sulfonatoethyl)piperazine], one of the most commonly used buffers in the Good series. In fact, of Good’s 20 buffer compounds, only three are completely noncomplexingsMES, MOPS, and PIPES1,8,9 (see Figure 2)sand the use of PIPES is restricted by its limited solubility. Direct spectral evidence of metal complexation by Good’s other buffers can, in fact, be readily observed in solutions with moderate concentrations of copper(II) ion.1 In 1973, Deutsch and Cheung15 proposed the use of 2,6-di-tertbutylpyridine and its substituted derivatives because they were compounds which could react with hydrogen ion but were too sterically hindered to complex with metal ions. Unfortunately, most of these compounds have very limited solubility in water and, thus, are of limited use for aqueous studies. A decade later, (15) Deutsch, E.; Cheung, N. K. V. J. Org. Chem. 1973, 38, 1123-1126.
Elias and co-workers16 took a similar approach in developing a series of 2,6-dimethylpyridine compounds as purported noncomplexing buffers. Their compounds were much more soluble but also much less sterically hindered, which presumably accounts for scattered reports of metal ion interference when using these buffers.17-19 All buffers based on pyridine also have the disadvantage that they absorb strongly in the ultraviolet region of the spectrum, making them unsuitable if UV monitoring is required. To circumvent the various problems associated with previous buffer series, we recently developed a series of noncomplexing buffers1 which rely on the use of sterically hindered tertiary alkylamines, most of which are closely related to Good’s MES, MOPS, and PIPES buffers. In our “Better” buffer compounds, we have avoided the use of possible weak donor atoms on the R, β, or γ carbons, with the exception of sulfonate groups which appear to be too weakly coordinating to promote interference. By incorporating more than one amine nitrogen into a single compound and/or varying the position of substituents with positive or negative inductive effects, a series of buffer compounds has been generated which covers the entire range from pH 3 to 11. Studies at pH values above and below this region generally do not require buffering for reasonably dilute solutions as a result of the high concentrations of H+ or OH- already present in solution. The 10 compounds for which acid dissociation constant data are now reported as a function of temperature are illustrated in Figure 2. EXPERIMENTAL SECTION Reagents. The buffer compounds MOPS, MES, PIPES, TEEN, and TEMN were obtained directly from Aldrich Chemical Co. TEMN was purified by vacuum distillation prior to use. The compounds DEPP, PIPPS, and PIPBS were synthesized by adaptation of literature methods as previously described.1 The newly added analogue, MOBS, was synthesized by the same approach used previously for PIPBS20 except that morpholine, (16) Bips, U.; Elias, H.; Hauro ¨der, M.; Kleinhans, G.; Pfeifer, S.; Wannowius, K. J. Inorg. Chem. 1983, 22, 3862-3865. (17) Raycheba, J. M. T.; Margerum, D. W. Inorg. Chem. 1980, 19, 837-843. (18) Elias, H.; Rass, U.; Wannowius, K. J. Inorg. Chim. Acta 1984, 86, L37L38. (19) Aoki, K.; Inaba, M.; Teratani, S.; Yamazaki, H.; Miyashita, Y. Inorg. Chem. 1994, 33, 3018-3020. (20) Jermyn, M. A. Aust. J. Chem. 1967, 20, 183-184.
Analytical Chemistry, Vol. 71, No. 15, August 1, 1999
3141
Table 1. Acitivity Coefficients for Hydronium and Hydroxide Ions and Autoprotolysis Constants for Water As Utilized in This Work
log γH log γOH- a pKw0 b pKwc c +
a
15 °C
25 °C
35 °C
45 °C
-0.0823 -0.1167 14.346 14.147
-0.0835 -0.1186 13.996 13.794
-0.0848 -0.1206 13.680 13.474
-0.0863 -0.1228 13.396 13.187
a
Activity coefficient values calculated from the extended DebyeHu ¨ ckel equation using optimal values of the “A” and “B” parameters (ref 26) with ion size parameters of 9 × 10-8 and 3.5 × 10-8 cm for the hydronium ion and hydroxide ion, respectively (ref 25). b Activity autoprotolysis constants utilized for water (ref 27). c Concentration autoprotolysis constants as employed in the PKAS program.
rather than piperazine, was utilized as the starting material. In the current work, samples of DEPP were titrated shortly after synthesis to minimize the possibility of degradation. DESPEN, which was first synthesized in our laboratory,1 was produced with improved purity by utilizing a slightly modified procedure. A mixture of 2.0 g (0.017 mol) of N,N′-diethylethylenediamine (Aldrich), 5.0 g (0.045 mol) of fresh 1,3-propane sulfone (Aldrich), 5 mL of triethylamine, and 12 mL of distilled water was refluxed for 20-25 min, chilled, and placed in a separatory funnel. Extraction with diethyl ether (3 × 5 mL) removed the unreacted materials. The remaining aqueous layer was adjusted to pH 1.5 with dilute HCl and evaporated to leave a brownish-yellow oily product. Approximately 25 mL of hot methanol was added to produce a pale yellow powder which was filtered on a fritted glass funnel and washed repeatedly with hot methanol to yield a pure white product which was free from the small impurity previously detected by NMR.1 All solutions were prepared using distilled deionized water. For the pH titrations, sodium hydroxide solutions were prepared daily from saturated NaOH to minimize carbonate contamination and were standardized against primary standard potassium hydrogen phthalate. pH Measurements. A measured quantity of each buffer was dissolved in a known volume of distilled deionized water placed in a 200-mL tall-form beaker, which was inserted into a water jacket. A measured quantity of HClO4 solution was added to ensure that the buffer compound was initially fully protonated. Ionic strength was controlled at 0.10 M using NaNO3. Temperature control was achieved by use of a circulating external water bath. The beaker contained a magnetic stirrer and was covered by a rubber stopper into which holes had been drilled to permit insertion of the electrodes, buret tip, thermometer, and a nitrogen gas inlet. Nitrogen gas was flowed over the surface of the solution during the titration to minimize carbon dioxide absorption. All pH measurements were made with an Orion Model 901 Ionalyzer equipped with glass and Ag/AgCl electrodes. The meter was standardized daily using (i) 0.05 molal potassium hydrogen phthalate and (ii) 0.025 molal KH2PO4 + 0.025 molal Na2HPO4 standard buffers (purchased as prepackaged powders from Fisher Scientific). The defined pH values for these two buffers at 15, 25, 35, and 45 °C are, respectively, (i) 3.999, 4.008, 4.024, 4.047 and 3142 Analytical Chemistry, Vol. 71, No. 15, August 1, 1999
(ii) 6.900, 6.865, 6.844, 6.834.21 Since these are the two most commonly used standard buffers specified for establishing the internationally adopted operational pH scale, their use permits the hydronium ion activity to be closely approximated as 10-pH. Further details of the experimental approach have been reported previously.22 RESULTS AND DISCUSSION All acid dissociation constants are reported as mixed-mode constants in which the hydronium ion is expressed in terms of activity (aH) and the conjugate acid and base species are expressed in terms of molar concentration (charges on the latter two species being omitted here for purposes of generalization):
HB h H+ + B
Ka1m )
H2B h H+ + HB
Ka2m )
aH [B] [HB] aH [HB] [H2B]
(1)
(2)
(For buffer solutions containing millimolar concentrations of buffer, the use of Kam values permits the direct application of the Henderson-Hasselbalch equation,
[B] pH ) pKam + log [HB] in the region of pH 5-9; and only minor corrections are required for pH ranges 4-5 and 9-10). For each experimental titration, a measured excess of a standard solution of HClO4 was added initially to ensure that the buffer compound was fully protonated. Approximately 50 incremental additions of the standard NaOH solution were made during the course of each titration with the pH values being recorded. The individual protonation constants were resolved using Martell and Motekaitis’ PKAS software23 based on Bjerrum’s nj H approach,24 where nj H is defined as the average number of ionizable hydrogen ions attached to the base species, B. For a diprotic species, nj H may be expressed as:22
CH - COH + [OH-] - [H+] nj H ) ) CB [B] + [HB] + [H2B] (3) [HB] + 2[H2B]
In eq 3, CH and COH represent the total concentrations of HClO4 and NaOH, respectively, added to the solution of base (buffer), CB represents the total concentration of the buffer compound in all forms, and [OH-] and [H+] represent the molar concentrations of hydroxide and hydronium ions, respectively, at each point in the titration as calculated by applying the respective activity (21) Bates, R. G. Determination of pH: Theory and Practice, 2nd ed.; Wiley: New York, 1973; p 73. (22) Westerby, B. C.; Juntunen, K. L.; Leggett, G. H.; Pett, V. B.; Koenigbauer, M. J.; Purgett, M. D.; Taschner, M. J.; Ochrymowycz, L. A.; Rorabacher, D. B. Inorg. Chem. 1991, 30, 2109-2120. (23) Martell, A. E.; Motekaitis, R. J. Determination and Use of Stability Constants, 2nd ed.; VCH Publishers: New York, 1992. (24) Bjerrum, J. Metal Ammine Formation in Aqueous Solution: Theory of the Reversible Step Reactions; P. Haase & Son: Copenhagen, 1957.
Table 2. Mixed-Mode Acid Dissociation Constants for Noncomplexing Tertiary Amine Buffer Compounds in Aqueous Solution As a Function of Temperature (µ ) 0.10 M)
TEMN, pKa2m b TEEN, pKa1m pKa2m DEPP, pKa1m pKa2m DESPEN, pKa1m pKa2m PIPES, pKa1m h pKa2m PIPPS, pKa1m pKa2m PIPBS, pKa1m pKa2m MES, pKam MOPS, pKam MOBS, pKam
15 °Ca
25 °Ca
35 °Ca
45 °Ca
∆H°, kJ mol-1
∆S°, J mol-1 K-1
11.23 (1)c 6.69 (1)c 10.10 (2)c 4.62 (4)e,f 8.68 (5)e,f 5.72 (2)c 9.17 (1)c 2.7 (5)e 6.81 (9)e 3.85 (0)c 7.99 (1)c 4.37 (2)d 8.78 (3)d 6.17 (1)d 7.23 (1)c 7.60 (2)c
11.01 (4)d 6.58 (7)d 9.88 (6)d 4.48 (1)c,g 8.58 (2)c,g 5.62 (1)c,g 9.06 (0)c,g 2.67 (12)d,f 6.78 (1)d,f 3.79 (2)e 7.97 (1)e 4.29 (4)c 8.55 (2)c 6.06 (10)d,f 7.09 (6)c 7.48 (2)d
10.68 (4)c 6.33 (3)e 9.72 (10)e 4.41 (2)c 8.47 (5)c 5.44 (3)e 8.88 (3)e 2.20 (8)e 6.70(3)e 3.74 (3)e 7.86 (1)e 4.17 (3)c 8.38 (1)c 5.97 (3)c 6.98 (0)c 7.37 (0)c
10.52 (14)c 6.16 (2)c 9.48 (1)c 4.31 (3)d,f 8.43 (4)d,f 5.35 (2)e 8.77 (3)e 2.60 (3)c 6.61 (1)c 3.65 (7)c 7.83 (2)c 4.13 (1)e 8.29 (5)e 5.91 (2)d 6.82 (4)c 7.22 (4)c
44 ( 4 33 ( 3 35 ( 4 18 ( 2 16 ( 2 23 ( 1 24 ( 1
-65 ( 12 -14 ( 9 -73 ( 13 -26 ( 5 -112 ( 6 -31 ( 4 -93 ( 4
12 ( 3 11 ( 2 11 ( 1 15 ( 1 29 ( 2 16 ( 7 24 ( 2 22 ( 2
-91 ( 9 -35 ( 6 -115 ( 4 -33 ( 4 -67 ( 7 -64 ( 2 -56 ( 6 -70 ( 4
a Mean pK m values are shown for each temperature with the standard deviation listed in parentheses relative to the last digit(s) given (e.g., a 11.01 (4) represents 11.01 ( 0.04 and 2.67 (12) represents 2.67 ( 0.12). b For TEMN, pKa1m < 1. c On the basis of two replicate determinations. d On the basis of three replicate determinations. e On the basis of four replicate determinations. f An additional replicate determination was discarded due to the fact that one (or both) pKa values exhibited internal standard deviations which exceeded (0.06 (in the lone case of PIPBS at 25 °C, two such replicate determinations were discarded); under these circumstances, the “goodness of fit” was considered to be unsatisfactory for that specific determination. g The 25 °C values originally reported for DESPEN and DEPP (ref 1) were found to be significantly in error due to the presence of a homologous impurity which has presumably been eliminated in the current study. h The pKa1m values for PIPES are inexact due to a combination of the limited solubility of this compound and the low pKa value; thus, meaningful ∆H° and ∆S° values cannot be generated for pKa1m.
coefficients for these ions to the measured pH value. All of the terms on the right of eq 3 are determined experimentally with the values of CH, COH, and CB being corrected for dilution at each step. For conversion of activity to molar concentration (and vice versa) for the aquated hydrogen and hydroxide ions, activity coefficients were calculated from the extended Debye-Hu¨ckel equation using ion size parameters of 9 × 10-8 and 3.5 × 10-8 cm for these two ions, respectively,25 and the values of the DebyeHu¨ckel “A” and “B” parameters were used as listed by Bates for each temperature.26 The resultant activity coefficient values (γΗ+ and γΟΗ-) used and the activity autoprotolysis constant (Kw0) for water27 at each temperature are given in Table 1. The values of nj H can be used to calculate the apparent Ka1m and Ka2m values for each titration point according to the following expressions:22
Ka1m )
Ka2m )
(2 - nj H)(aH)2 (nj H - 1) aH + nj HKa2m
(2 - nj H)(aH)2 + Ka1m(1 - nj H)(aH) nj H Ka1m
(4)
(5)
The contribution of Ka2m to eq 4 was ignored when obtaining an initial estimate of Ka1m using the titrimetric data for the range of 1.9 > nj H > 1.1; and the resultant Ka1m estimate was then inserted into eq 5 to obtain an initial estimate of Ka2m from eq 5 using data for the range 0.9 > nj H > 0.1. Refinements using the entire set of data were made by iteration. (NOTE: In practice, the commercial PKAS program calculates the concentration values, Ka1c and Ka2c, (25) Ref 21: Table 3-3, p 49. (26) Ref 21: Table 4, pp 449-450. (27) Ref 21: Table 2, p 448.
based on [H+]; these values were then corrected back to obtain the mixed-mode constants, Ka1m and Ka2m.) Upon final refinement of the two acid-dissociation constants, the PKAS program calculates the theoretical pH at each point in the titration and compares these values to the experimental data to provide a standard deviation for both pKa1m and pKa2m. In the current work, any titration which resulted in an internal standard deviation g 0.06 for either pKa value was arbitrarily discarded from the final data set (regardless of the level of agreement with other data) since the goodness-of-fit was considered to be unsatisfactory. (Of the 113 titrations included in this work, six were rejected on this basis, and one experimental pKa1m value for PIPBS at 45 °C was rejected using Dixon’s Q test at the 95% confidence level.)28 At least two titrations were carried out for each buffer compound at each temperature. When the results for a single temperature, or the trend in Ka values for the range of four temperatures, showed inconsistencies, additional titrations were carried out. The checks provided by the multiple temperature measurements yielded data which are considered to be superior to those reported earlier for 25 °C alone.1 The mean pKam values for all 10 buffer compounds at all four temperatures are provided in Table 2. The individual experimental pKam values were plotted against the reciprocal temperature for each acid dissociation process according to the expression
pKam )
∆S° ∆H° 2.303RT 2.303R
(6)
to provide the apparent enthalpy and entropy values. {NOTE: Such reciprocal temperature plots of pKam values for buffer (28) Rorabacher, D. B. Anal. Chem. 1991, 63, 139-146.
Analytical Chemistry, Vol. 71, No. 15, August 1, 1999
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compounds are often not completely linear as exemplified by the temperature dependence of the standard buffers used in this work (see Experimental Section). This presumably results from the fact that at least three processes contribute to each pKa value: (i) rupture of the nitrogen-hydrogen bond, (ii) formation of the water-hydrogen bond (to form the hydronium ion), and (iii) rearrangement of the hydrogen-bonded structure of the surrounding solvent matrix.} Plots of eq 6 for each set of pKam data appeared to be linear within the experimental accuracy, yielding the apparent ∆H° and ∆S° values listed in Table 2. The internal consistency of the data is indicated by the fact that pKam values calculated using these ∆H° and ∆S° values at each specific temperature are, in all cases, within (0.05 of the experimental mean pKam values, and all but five (pKa2m for TEMN at 35 °C, pKa1m for TEEN at both 15 and 25 °C, and pKa2m for PIPBS at both 25 and 35 °C) are within (0.03 of the experimental means. For all pKam values, the standard deviations are given in parentheses in Table 2 in terms of the last digit(s) listed (see footnote to Table 2). Comments. Of the newly determined pKam values, those for DESPEN and DEPP are significantly different than the values which we previously reported for 25 °C.1 The earlier samples were found to contain homologous impurities which were removed by improved purification procedures prior to the current measurements. The pKam values reported for PIPBS and MES, as determined in this study, also differ significantly from the values previously reported.1 The originally reported data for PIPPS were very imprecise. The current pKam values, though changed only slightly, are vastly superior in their precision. For TEMN, pKa1m < 1, too small to be determined titrimetrically. Fully protonated PIPES is of limited solubility (approximately 1 mM), thereby limiting the solution concentrations to levels which were too dilute to permit a precise titrimetric evaluation of the relatively small pKa1m value. As a result, the data for this last constant are too scattered to permit an evaluation of apparent ∆H° and ∆S° values. CONCLUSIONS The diagram in Figure 3 illustrates the effective buffering range of each tertiary amine compound for ratios of [HnB]/[Hn-1B] ranging from 1:5 to 5:1, based on a total buffer concentration of 1 mM. For pH values below 5 and above 9, corrections were made for self-dissociation. From this diagram it is readily apparent that the use of these 10 compounds provides effective buffering over the entire range of pH 3-11. None of these compounds shows evidence of complex formation with copper(II) ion even when the
3144 Analytical Chemistry, Vol. 71, No. 15, August 1, 1999
Figure 3. Schematic diagram illustrating the useful pH ranges for the noncomplexing tertiary amine buffer compounds discussed in this work at 25 °C. The ranges are based on solutions of millimolar concentration in which the conjugate acid-base ratio varies from 1:5 to 5:1.
latter is present at high concentration levels. In view of their freedom from metal ion complexation, we suggest that these buffers be termed Better buffers (to distinguish them from Good buffers). Although additional buffer compounds could be generated using the same fundamental principles involving steric hindrance, the current series of compounds should effectively solve the problem of buffer interference in aqueous studies on metal ions. If increased solubility is required, derivatives of the current set of compounds can be generated as long as donor atoms are excluded from the R, β and γ positions. The limited solubility of PIPES is not a matter of serious consequence in its upper pH buffering range since MOPS and TEEN can be used to cover this region adequately. All 10 compounds reported in this paper have recently been made commercially available by GFS Chemicals.29 ACKNOWLEDGMENT This work was supported by the National Science Foundation under Grant CHE-9528831. Received for review March 5, 1999. Accepted April 22, 1999. AC9902594 (29) GFS Chemicals, Inc., P.O. Box 245, Powell, OH 43065.
