in solutions 0.1JI in HClO,. The reduction with tributylphosphine qeerns to be efficient in solutions containing no electrolyte r r with acetic acid and ammonium acetate or sodium hydroside present. Refluxing did not appear to aid the redurtion. Elemental sulfur was added prior to titration to remove any excess phosphine which would cauze some difficulty in the location of the amperornetric end point (3). hlso, the end point was sharper if the solution was made 0.1Jf in HClOl prior to titration and after the reduction had occurred. The presence of elemental sulfur with the thiol causes an apparently complex
reaction to occur during addition of silver nitrate, but quantitative results are still possible (4). The solutions turn yellow or orange during titration and a dark precipitate, presumably silver sulfide, is produced. ACKNOWLEDGMENT
The authors thank 11 & T Chemicals for supplying the tributylphosphine. LITERATURE CITED
(1) Henderson, W. A., Jr., Buckler, S. A., J . Am. Chem. Soc. 8 2 , 5794 (1960).
( 2 ) Humohrev. It. E.. Hawkins. .T. %I --, ANAL.'CHE;;. 36, 1612 (1964). ( 3 ) Humphrey, It. E., Webb, R. M., I
Sam Hbusthn State College, unpublished work, 1964. ( 4 ) Karchmer, J. H., ANAL. CHEM.2 9 , 425 (1957). ( 5 ) Moore, C. G., Trego, B. R., Tetrahedron 1 8 , 205 (1962). (6) Schonberg, A , , Barakat, M. Z., J . Chem. SOC.1949, p. 892.
R A Y E. HUMPHREY JAMESL. POTTER
Department of Chemistry Sam Houston State College Huntsville, Texas SUPPORT of the Robert A. Welch Foundation of Houston, Texas, is gratefully acknowledged.
Investigation of 1,3-Diaminopropanol-2 Interactions with Copper and Nickel Ions SIR: During an investigation of enzyme-chelate interactions, we were concerned with the chelating properties of 1,3-diaminopropanol-2 (AOH) with Cu+2 and Ni+2. Isolation of a 2-1 AOH-CU+~complex has been reported (3) and Gonick ( 7 ) found that the C u f Z chelate reached a maximum li value of 1.5 on the basis of the following equilibrium : Cu+2
+ ;iOH
~
Cu.40Hf2
Later, Duertsch, Fernelius, and Block ( I , 2) showed that if the ethanolic hydrogen is released during titration with Cu +2, analysis of the interaction yielded steady values for the stability constant and a readily explainable maximum value of 1.0. The stability constant for the interaction Cu+2
+ AOH e C u h O + + H+
was reported as
K,
=
(CuXO+) (H+)
/ (Cut*) (XOH) (1)
because the ionization constant of the ethanolic hydrogen cannot be easily ascertained. The experimental method used by both Gonick and Bertsch consisted of titrating the metal ion in the presence of mineral acid with a solution of the amine. We found that the nature of the interaction is readily apparent if solutions of various amine to metal ratios are titrated with base as described by Chaberek and Martell ( 4 ) . When the differences in the two methods are taken into consideration, our results are in close agreement with the results obtained by Bertsch, Fernelius, and Block ( 1 , 2 ) . In the N i i 2 systems, there is evidence that the ethanolic group participates in chelation but that the
ethanolic hydrogen acts as a very weak acid. To clarify the interactions further, spectrophotometric methods were used to determine chelate ligand to metal ratios and the k , of the ethanolic group. Although the latter experiments were not entirely successful, we were able to obtain an approximate k , value which agreed well with a theoretical calculation based on inductive effects in organic compounds according to the method of Hine ( 8 ) and Dewar (6). EXPERIMENTAL
Analytical reagent grade hTi(NO& and Cu(NO& were used and the solutions standardized according to the chelatometric titrations described by Wilson and Wilson (11). T h e dihydrochloride salt of AOH was prepared by dissolving the amine in excess concentrated HC1, analytical reagent grade; flash-evaporating to dryness; and recrystallizing several times from methanol-ethanol-water mixtures, m.p. 163-4' C. All titrations were made a t 30 i 0.1' C. and at an ionic strength of 0.162 with NaCl as the supporting electrolyte to correspond with our enzyme studies. Ligand concentration was kept constant a t 0.02JP in all studies and the metal concentration was varied appropriately. I n all titrations, 0.1933M NaOH was the base employed. The experimental setup and procedure were as previously described (4). A Beckman Zeromatic pH meter was used to follow the potentiometric titrations of solutions containing various ligand to metal ratios. When necessary, a sufficient interim of time was allowed to elapse between additions of base until steady pH values were attained. Beckman Models DK-2 and DU spectrophotometers were used to analyze solutions of various ligand to
metal ratios according to the method of Job (9). The DK-2 wa5 also used in the experiments to determine the k , of the ethanolic group. This latter experiment consisted of preparing a series of AOH-SaOH solutions of known OH- concentration prepared with Fischer reagent ION NaOH. T h e ionic strength was kept constant with NaCI. The p H of these solutions was calculated using known activity values for XaOH solutions of known molarity. The solutions were scanned from 200 to 350 mp. T h e peaks for the dissociated and undissociated forms were not separable. The poor resolution observed may be attributed to the high OH- concentrations used in these studies. RESULTS A N D DISCUSSION
Titrations of Cu +2-AOH(HCI)2 mixtures and Ni +2-;\OH(HC1)2 mixtures, as well as free amine di HCI are shown in Figures 1 and 2. Ligand-Cu+2 ratios of 1-1 and greater all yield exactly three equivalents of H + per metal ion. This can only be explained on the basis of formation of a 1-1 terdentate chelate with simultaneous release of the ethanolic proton. That a 2-1 chelate does not form is evident from a comparison of curves .I ! B , and C in Figure 1 ; curve C is a summation of curves .I and B . Since a 2-1 chelate does not form, the possibility that the third proton comes from Cu+2-bound water is excluded. This system was also analyzed according to the method of .Job, utilizing the peak at 623 mp exhibited by the C u f 2 chelate. Unfortunately, mixtures containing a mole fraction of amine lethan 0.5 precipitated a t the pH of the experiment. Therefore no absorption VOL. 37, NO. 1 , JANUARY 1965
165
14 I
m
Figure 1. Titration of C U + ~in the presence of 1,3-diaminopropanol-2 A = Titration of the diamine with NaOH; B = equimolar amounts of C U + ~ and diamine; C = 100% excess of diomine; m i s moles of base per mole of metal; for curve A, abscissa taken as moles of base per mole of ligand
maximum could be observed at a 1-1 ratio, but solutions containing larger ratios of amine (at constant total concentration) showed decreasing absorption which is corroborative evidence of only 1-1 chelate formation. The K + 2 chelate titration curves, shown in Figure 2, indicate formation of a 2-1 chelate which probably involves the ethanolic group. In the 2-1 hOH(HC1)2-Ni+2 mixture, the break a t four equivalents of base per metal corresponds to titration of the protons coordinated to the amine nitrogen atoms. If the ethanolic group is involved in metal bonding, there should be no evidence of formation of a 3-1 chelate. Titrations of 3-1 and 5-1 amine-Nif2 mixtures do not indicate formation of a 3-1 complex. If the curves in Figure 2 are compared at p H 10.5 (the inflection point in the titration curve of free amine), it can be seen that approximately one additional mole of base per metal ion is used beyond that expected if only a nitrogen coordinated complex is formed; the value of m is 2, 5.1, 6.9, and 10.8 for curves .1, B , C, and D, respectively, a t pH 10.5. This suggests that, at this pH, one additional proton from metal-bound water or the ethanolic group in the chelate is being titrated. .in attempt to determine the stoichiometry according to the method of Job was unsuccessful; the spectrum of the chelate mixtures changed when the ratio of amine to Ni+* was varied from 1.5-1 to 2-1, and 1-1 mixtures could not be prepared because of precipitation. In any case, this analysis gave no evidence of a 3-1 species a t pH 9.00. .4t pH’s > 10 no precipitate aplleared and the solutions lost their color. Ionization 166
0
ANALYTICAL CHEMISTRY
of the extra proton could bc responsible for this effect. Evaluation of stability constants as usually defined would require the evaluation of the k , of the ethanolic group. R e had reason to believe that we could determine this value spectrophotometrically. Ideally, we should observe two distinct peaks corresponding to the undissociated and dissociated species. Unfortunately, the necessity of using systems of very high pH introduced solvent effects which prevented clearcut resolution of the peaks. Instead one peak appeared which continually shifted toward higher wavelengths with a corresponding lowering of absorption as the pH was increased. Then a t pH 14.06 the absorption began to increase while the peak continued to shift as before. This increase in absorption was attributed to [ -01- formation and the pH of the solution when this was first observed was considered to be the pK, of the ethanolic hydrogen. Admittedly, the method was only an approximation and some corroborative evidence was needed. h theoretical calculation of the pK, of the ethanolic hydrogen was made using the method of inductive effects outlined by Hine (8).
For a series of substituted acetic acids ( 5 ) , u * , the measure of electron donation or acceptance, was +3.651. When this value was used in the expression derived for substituted alcohols and corrections were made for carbon chain lengths and the presence of two positively charged nitrogens (10), the pK, was calculated to be 14.35. The estimated pK, value of 14.06 found in the spectrophotometric method appears therefore to be the correct order of magnitude. Our calculation of the stability constant for the (CuhO)+ chelate ( K 1 ) agrees well with the log K , value of 3.57 rt 0.17 obtained by Bertsch, Fernelius, and Block (8)if the K , is defined as in Equation 1 because
K,
KS k,
= -
where
K1
=
(CUAO) (CU +2) (a0-) +
(3)
the value of K 1 is dependent on the k.. Using our estimated k , we found an average value of 2.5 X lo1*for K1. The stability constants for the 1-1 and 2-1 Ni-AOH chelates reported by Bertsch, Fernelius, and Block ( 1 ) of 5.42 for log K1 and 4.16 for log K z compare well with our values of 5.47 and 4.14 for log K I and K 2 ,respectively, a t 30 f 0.1” C.
O
P
4
6
8 m
1
0
1
P
Figure 2. Titration of N i f 2 in the presence of 1,3-diominopropanol-2 A = Titration of the diamine with N a O H ; B = solution contains diomine ta Nii2 ratio of 2 to 1 ; C = 3 to 1 ratio; D = 5 to 1 ratio; m is moles of base per mole of metal; for curve A, abscissa taken as moles of base per mole of ligand
ACKNOWLEDGMENT
The authors thank Scott MacKenzie, Jr., for his assistance with the theoretical organic calculations and William E. Ohnesorge for his aid in interpreting the spectral data. LITERATURE CITED
( 1 ) Bertsch, C. R., Fernelius, W. C., Block, B. P., J . Phys. Chem. 62, 4 4 4 4 0 flQ.581. ,-~--,. ( 2 ) Zbid., pp. 503-4. ( 3 ) Breckenridge, J. G., Hodgins, J. W. R.. Can. J . Research 17B. 331-5
(1939). (4)Chaberek, S., Jr., Martell, A. E., J . Am. Chem. SOC.74, 5052-56 (1952). (5) Cohn, E. J., Edsall, J. T., “Pyoteins,
Amino Acids, and Peptides, ACS Monograph, Itheinhold, New York, 1962. (6) Dewar, AI. J. S., J . Am. Chem. SOC. 84, 35-9 (1962). ( 7 ) Gonick, E., Doctoral Dissertation, The Pennsylvania State College (1951). (8) Hine, J., “Physical Organic Chemistry,” 2nd ed., McGraw-Hill, Kew York, 1962. (9) Job, P., Ann. Chim. 9, 10, 113 (1928). (10) Taft, It. W., “Steric Effects in Organic Chemistry,” Wiley, New York, 1956. (11) Wilson, C. L., Wilson, U. W., I ‘Comprehensive Analytical Chernistry, p. 351, Elsevier, New York, 1960. ERNEST MARIO SANFORD M. BOLT ON^
Pharmacy Department University of Rhode Island Kingston, R. I. THISwork supported in full by Public Health Service research grant NB04580-02 from the Institute of Neurological Disease and Blindness.
Present address, Department of Biostatistics, Columbia University, New York, N. Y .
