64
J. Phys. Chem. 1981, 85, 64-68
Thermodynamics of Protonation of Tetramines with Different Degrees of N-Methylation Rotando Barbucci, * Vlncenro Barone, Istituto Chimico, Sezione di Chimica Industriale ed Inorganica, Universitti di Napoll, I-80 134 Napoli, Italy
Mauro Mlcheloni, Istituto dl Chimica Generaale, Universitci di Firenze, I-50123 Firenze, Italy
and Lulsa Rusconl Istituto di Chimica Industriale del Politecnico, I-20133 Milano, Italy (Received: March 10, 1980; In Final Form: October 1, 1980)
In an attempt to give further insight into the effect of increasing methylation on the protonation of polyamines we studied the tetramines hexamethyltriethylenetetramine (trienMe6) and N,N,N”’,N”’-tetramethyltriethylenetetramine (trienMe4). The results are discussed with special reference to the related compound triethylenetetramine (trien). Different methylation of both terminal and inner nitrogens apparently affects the thermodynamic functions. The AGO, AH”,and AS” values suggest a unique protonation mechanism starting with the terminal nitrogen atoms. The stepwise entropy changes are discussed on the basis of the releasing of water molecules. Introduction The thermodynamics of protonation of polyamines has been the object of several studies both in aqueous solut i o ~ and, ~ ~ -more ~ recently, in the gas phase.6-8 Many attempts have been made to correlate the thermodynamic functions with the structural features of the p~lyamines.~p~$ For instance, the effect of the replacement of carbon atoms by heteroatoms (e.g., by passing from monoamines to diamines, triamines, and so o ~ ) ~ J Oand - ~ the ~ effect of the length of the aliphatic chain connecting the aminic nitrogen~’~-~‘ have been extensively studied. The behavior of several polyamines of different structure in aqueous solution has been related especially with more or less effective solute-solvent i n t e r a c t i o n ~ . ~ J ~ - ~ ~
However, few complete thermodynamic studies have been performed on the effect of N-methylation on the protonation of diamines1°J4and, to our present knowledge, no such studies have been performed in the case of open chain polyamines. These studies would eventually lead to a better understanding of the thermodynamic behavior of the macrocyclic amine ligand^.^^-^^ The aim of this paper is to report the results we obtained on the thermodynamics of protonation of two Nmethylated tetramines: trienMe6 and trienMe4.
trienMe, (1) “Stability Constants of Metal Complexes”, Chem. SOC.Spec. Publ., No. 17 (1964); No.25 (1972). (2) P. Paoletti, R. Barbucci, and A. Vacca, J. Chem. SOC.,Dalton Trans., 2010 (1972). (3) J. J. Christensen, R. M. Izatt, D. P. Wrathall, and L. D. Hansen, J. Chem. SOC. A, 1212 (1969). (4) E. J. King in “Acid-Base Equilibria”, R. A. Robinson Ed., McGraw-Hill, New York, 1965 (topic 15 of the “International Encyclopedia of Physical Chemistry and Chemical Physics”). (5) G. Schwarzenbach, Pure Appl. Chem., 307 (1970). (6) J. I. Brauman and L. K. Blair, J.Am. Chem. SOC.,90, 6561 (1968); 91, 2126 (1969); 92, 5986 (1970); 93, 3911, 3914 (1971). (7) D. H. Aue, H. M. Webb, and M. T. Bowers, J. Am. Chem. Soc., 94, 4726 (1972); 95, 2699 (1973); 98, 318 (1976). (8) R. Yamdagni and P. Kebarle, J.Am. Chem. Soc., 95,3504 (1973). (9) R. Barbucci and V. Barone, J. Solution Chem., 8, 427 (1979). (10) P. Paoletti, R. Barbucci, A. Vacca, and A. Dei, J. Chem. SOC.A, 310 (1971). (11) A. Anichini, R. Barbucci, L. Fabbrizzi, and A. Mastroianni, J. Chem. Soc., Dalton Trans., 2224 (1977). (12) R. Barbucci and M. Budini, J. Chem. Soc., Dalton Trans., 1321 (1976). (13) R. Barbucci, P. Poaoletti, and A. Vacca, J. Chem. SOC. A, 2202 (1970). (14) R. Barbucci and A. Vacca, J. Chem. SOC.,Dalton Trans., 2363 (1974). (15) P. Paoletti, F. Nuzzi, and A. Vacca, J.Chem. SOC.A, 1385 (1966). (16) M. Ciampolini, P. Paoletti, and L. Sacconi, J. Chem. SOC.,2994 ~
(1961) ~____,_
(17) L. Fabbrizzi, R. Barbucci, and P. Paoletti, J.Chem. SOC.,Dalton Trans., 1529 (1972). (18) A. F. Trotman-Dickenson, J. Chem. SOC.,1293 (1949). 87, 4481, 4485, 4491 (1965). (19) F. E. Condon, J. Am. Chem. SOC., (20) A. G. Evans and S. D. Hamann, Trans. Faraday SOC.,47, 34 (1951). (21) R. G. Pearson and D. C. Vogelsang, J.Am. Chem. SOC.,80,1038 (1958). (22) H. K. Hall, J. Am. Chem. Soc., 79, 5441 (1957). 0022-3654/81/2085-0064$01 .OO/O
H CH3, CH3’
I
H
I
/c H3 NCH~CH~NCHZCH~NCH~CHZN \CH3 trienMe,
Experimental Section Ma t e ria 1sa N ,N ,N ’,N ’’,N ’”,N ’”-Hexamet h y 1t r i ethylenetetramine (trienMe6)was prepared by methylation of trien with a mixture of formic acid and formaldehyde. After all carbon dioxide had been eliminated, the mixture was heated under reflux for 8 h and then acidified with hydrochloric acid. The solution was evaporated to dryness, and the residue treated with a concentrated solution of sodium hydroxide. The crude amine separated as a dense oil and was purified by distillation (bp 110-112 “C (0.2 mmHg)). The trienMee was used as the hydrochloride. This was prepared by adding hydrochloric acid to an alcoholic solution of the amine. The product precipitated out. It was recrystallized from aqueous alcohol and dried to constant weight at 60 “C and 0.1 mmHg. Anal. Calcd for Cl2HNN4Cl4:C, 38.30; H, 9.11; N, 14.89; C1, 37.69. Found: C, 38.2; H, 9.2; N, 15.0; C1, 37.6. (23) F. P. Hinz and D. W. Margerum, Inorg. Chem., 13, 2941 (1974). (24) M. Kodama and E. Kimura, J. Chem. SOC.,Dalton Trans., 116 (1976). (25) M. Micheloni, A. Sabatini, and P. Paoletti, J. Chem. SOC.,Perkin Trans. 2, 828 (1978).
0 1981 American Chemical Society
Thermodynamics of Protonation of Tetramines
N,N,N”t,N’/‘-Tetramethyltriethylenetetramine (trienMe4) was prepared by heating a mixture of N,N-dimethylethylenediamine (100 g) and 1,2-dichloroethane(19 g) at 70 “C for 48 h. The reaction mixture was then treated with a concentrated solution of sodium hydroxide. The organic layer was dried over potassium hydroxide pellets and distilled at 0.2 mmHg. The excess N,N-dimethylethylenediamine passed over first (28-30 “C), followed by the product (69-70 “C): yield, 6.5 g. Anal. Calcd for C1J&N4C14:C, 34.49; H, 8.68; N, 16.09; C1, 40.73. Found: C, 34.6; H, 8.7; N, 15.9; C1, 40.8. Other Reagents. COz-free NaOH solutions were prepared, stored, and standardized as described elsewhere.26 Stock solutions of 0.1 M NaCl were prepared from sodium chloride (C. Erba, ACS grade) without further purification and used as the ionic medium for both potentiometric and calorimetric measurements. Emf Measurements. Potentiometric titrations were performed with a digital potentiometer, an ORION 9101-00 glass electrode, a silver/ silver chloride electrode, and a salt bridge containing 0.1 M NaCl solution. The output voltages (mV) were automatically recorded with a Printina Gay alphanumeric printer. The titration vessel was thermostatted at 25.0 f 0.1 “C. A stream of nitrogen, presaturated with water vapor by bubbling through a 0.1 M NaCl solution, was passed over the surface of the solution. For the titrations, the NaOH solution was dispensed from a Metrohm Multidosimat piston buret graduated in hundreths of a milliliter. Eo calibrations were performed before and after each titration. The experimental values of the emf in the strongly acid region (pH I2.5) have been corrected for the liquid junction potential following the procedure suggested by Bates.27 The corrections on emf values have been determined by titration with a standardized strong acid. The concentration of hydrogen ion was calculated from the emf values (in mV) by means of the formula [H+] = exp(E - E0)/25.693
No attempt was made to correct the data to zero ionic strength or to apply activity coefficient corrections since under the solution conditions used here the polyamine does not contribute very much to the ionic strength and the ionic strengths were quite high. The protonation constants were derived from 117 data points for the trienMe6 and from 119 data points for trienMe4. Some experimental details are reported in Table I. (Additional emf measurements are available as supplementary material. See paragraph at end of text regarding supplementary material.) Distribution diagrams and variation of pH vs. the millimoles of OH- added are plotted in Figure 1 for two representative emf titrations. The program MINIQUAD 76A used to calculate the basicity constants has been previously described.26~28 Calorimetric Measurements. A LKB Model 8700 calorimeter was used. The output voltages of the unbalanced Wheatstone bridge were automatically recorded every 20 s with a Solartron Model A200 digital voltmeter and a Printina Gay alphanumeric printer. Electrical calibrations were performed before and after the reaction. The enthalpies of protonation were obtained by using the continuous titration calorimetric method.29 The titrant (26) M. Micheloni, A. Sabatini, and A. Vacca, Inorg. Chim. Acta, 25, 41 (1977). (27) R. G. Bates, “Determination of pH”, Wiley-Interscience, 2nd ed, New York, 1973, p 261. (28) A. Sabatini, A. Vacca, and P. Gans, Talanta, 21, 53 (1974). Dalton Trans., 519 (29) A. Sabatini and A. Vacca, J. Chem. SOC., (1980).
The Journal of Physical Chemistry, Vol. 85, No. 1, 198 1 65
TABLE I : E x p e r i m e n t a l Details of the EMF M e a s u r e m e n t s ~~
nL,(I
nH ,a
curve
mmol
mmol
1
0.1945 0.3082 0.3943 0.3116
data pH range
points
2.5-10.2 2.3-9.5 1.9-8.5 3.0-10.0
32 28 21 36
2.3-10.0 2.7-8.5 2.4-9.8
48 31 40
trienMe,
2 3 4
0.7780 1.8180 3.0600 1.2464 trienMe,
1
2 3
0.3811 0.5511 0.5608
1.9631 2.4098 2.6297
T o t a l a m o u n t of a m i n e ( n L )or of acid (nH).
(NaOH) was introduced into the calorimeter, containing a solution of the amine in an excess of HC1, at constant rate by a Braun syringe modified by us. The same experiments, but with only HC1 present, have a value for the enthalpy of the reaction H+(aq) + OH-(aq) HzO of -13.44 kcal mol-l, in good agreement with the accepted v a l ~ e .The ~ ~experimental ~ ~ details are reported in Table 11. (A temperature-time plot obtained from a typical calorimetric measurement is available as supplementary material.) The stepwise protonation enthalpies have been obtained by handling the calorimetric data as previously described.29 For trienMe6 the same AH” values were obtained by using an LKB 100700 flow microcalorimeter following a procedure described elsewhere.32
-
Results and Discussion Basicity Constants. The log K values of the stepwise protonation of trienMe4 (I) and trienMe6 (11) are reported in Table 111. For comparison purposes we have also included in Table I11 the log K values of unsubstituted trie11,3~of N,N,N’-trimethylethylenediamine(enMe3) and of N,N,N’,”-tetramethylethylenediamine (enMe4).10By increasing the degree of methylation, the log K values decrease. This is demonstrated by the data reported in Table 111; they show that all the constants for I and I1 are lower than those of the unsubstituted trien; furthermore, the constants for the less methylated compound I are higher than those for 11. This trend has already been observed in the case of monoamines and diamines3J4J6and is probably related to the effect of the substituents on sol~ation.~~~ The difference between the first two constants in the tetramines reported in Table I11 may be accounted for by the statistical effect (log 4) connected with the protonation of two equivalent atoms. This indicates that the positive charge already present on the molecule does not affect the basicity of the nitrogen atom which is protonated in the second step. A similar behavior has been observed only when the two nitrogen atoms to be protonated are far apart and/or when they are separated by a strongly electronegative atom.9,11,12,14,16,34 The log K1 values of I and I1 are similar to each other and to that of enMe4, but are much lower than the log K1 of the unsubstituted trien. This suggests (30)J. D. Hole, R. M. Izatt, and J. J. Christensen, J.Phys. Chem., 67, 2605 (1963). ’ (31) J. J. Christensen, L. D. Hansen, and R. M. Izatt, “Handbook of Proton Ionization Heats and Related Thermodynamic Quantities”, Wiley-Interscience, New York, 1976. (32) R. Barbucci, P. Ferruti, C. Improta, M. La Torraca, L. Oliva, and M. C. Tanzi, Polymer, 20, 1298 (1979). (33) H. S. Greyf and L. C. Van Poucke, Thermochim. Acta, 4, 485 (1972). Dalton (34) R. Barbucci, L. Fabbrizzi, and P. Paoletti, J. Chem. SOC., Trans., 2403 (1974).
06
The Journal of Physical Chemistry, Vol. 85, No. 1, 1981
A mmol OH-
!
,
'.I i
I
i, I.
\
'
'
I
/
I
/.6
I
B Figure 1. Dlstrlbutlon curves (broken lines) and variation of pH vs. millimoles of OH- added (full lines) for two representatlveemf tltrations: (A) curve 1 of trienMe,; (B) curve 1 of trlenMee.
that the first atom to be protonated is in both cases a terminal tertiary nitrogen. Furthermore, the log K 2of I1 is almost equal to (though slightly lower than) that of I. It is interesting to note that the difference log Kl - log K2 (0.76) is equal for both I and 11. These facts lead to conclusion that the first two protons are taken up by the terminal tertiary nitrogens in both I and 11, while the small differences in the first two basicity constants are only due
Barbuccl et al.
to the different inductive effect of the nitrogen atoms present along the aliphatic chain (which are methylated in the case of 11). The third and fourth constants refer to the successive protonation of inner nitrogen atoms. As a consequence, they feel the repulsion of one or two already charged nitrogen atoms, which lowers their basicity. The fact that K,(I) is higher than K,(II) and K4(I) higher than K4(II) is reasonable considering that the last two protonations involve amine nitrogens with a different degree of substitution. Protonation Enthalpies. From the data reported in Table I11 it is apparent that all the stepwise protonation enthalpies of I are intermediate between those of the unsubstituted trien and 11. This agrees with the general finding that the protonation enthalpies of amines in aqueous solution decrease with methylati~n.'~~J~ As far as the first protonation enthalpy is concerned, lAHol(I)( lAHol(II)I < IAHo1(trien)l. Furthermore the AHo, values of our methylated tetramines are very similar to that of enMel. This is a further proof that the first nitrogen to be protonated is the same in I and I1 and, in particular, it is a terminal trisubstituted nitrogen. This result is in agreement with a 13C NMR study already performed on several triamines, which unambiguously show that in all cases the protonation of the inner nitrogen starts only after the protonation of terminal nitrogens is complete.36 As previously observed for triamines and trien,gJes36 lAHo21is higher than IAHoll,the difference being particularly high for I. The above NMR results35rule out the possibility of tautomeric equilibria between the different basic nitrogens on the molecule, which was previously invoked to explain the trend of stepwise protonation enthalpies.16 Furthermore, the presence of inner nitrogens could at most equalize the first and second enthalpy, by shielding the terminal nitrogens from each other. Therefore, the trend 1AHo21> IAHollis more reasonably explained by a particularly favorable disposition of water molecules around a diprotonated p~lyamine.~ In fact, after the second proton is transferred from the solvated hydroxonium ion to the singly charged base, the presence at the extremes of the molecule of two positive charges, and of negative charges in the middle, gives place to a new and more stable orientation of water molecules around the diprotonated species. The lAHo31value in I1 is very low, apparently because it is affected by a neighboring charged nitrogen atom. The same trend is found for I, but the IAHo31is higher because the nitrogen which undergoes protonation is disubstituted rather than trisubstituted. As for AHo4values, they are always very low due to the repulsive effect of two charged groups near the nitrogen which is being protonated. Once again, the protonation enthalpy of I1 is the lowest, due to the fact that protonation occurs on a trisubstituted nitrogen. Furthermore, the AHo., values are also affected by the degree of substitution of the already charged nitrogen atoms (Table 111). Recently, some papers have appeared showing that, in the case of monoamines, there is a linear relationship between proton affinities and protonation enthalpies in aqueous solution, and that the different enthalpy changes for different amines mainly depend on the net charge on , ~ have computed the the nitrogen to be p r ~ t o n a t e d . ~We net charges of the nitrogens which undergo protonation (35) M.Delfini, A.L. Segre, F. Conti, R. Barbucci, V. Barone, and P. Ferruti, J. Chem. SOC.,Perkin Trans. 2,900 (1980). (36) P. Paoletti, M.Ciampolini,and A. Vacca, J.Phys. Chem., 67,1065 (1963).
The Journal of Physical Chemistry, Vol. 85, No. 1, 1981 67
Thermodynamics of Protonation of Tetramines TABLE I1 : Experimental Details of the Calorimetric Titrations
a
experiment
nL,a mmol
nH,a mmol
initial vo1, mL
titrant (NaOH), M
injection rate, mL min-'
injection time, min
1 2
0.23603 0.151 92
1.2613 5 0.94454
trienhle, 81.6 82.9
0.1578 0.1578
1.04483 1.044 83
10 10
1 2
0.29 526 0.12872
1.28034 0.6 0824
trienMe, 79.5 80.0
0.1476 0.1476
1.04489 1.04489
5 10
Total amount of amine ( n L )or of acid ( n ~(initial). )
TABLE 111: Thermodynamic Functions for the Stepwise Protonation of trienhle, and trienMe, in 0.1 M NaCl at 25 " C -AGO ,a -AH'," ASD,U amine (L) reaction log K kcal/mol kcal/mol cal mol'' K-'
-
L + H+ LH+ LH+ + H+ LHz2+ LHz2+t H+ + LH,3+ LH,,+ t H+ = LHA4+ L+"H+-LH+ ' LH+ + H+ e LHm2+ LHZZ+ t H + - LH,3+ LH,3+ t H + - LHd4+ L + H+ LH+ LH+ + H+ = LHZ2+ LHz2+t H + = LHS3+ LH,,+ t H+ = LH,4+ L + H+ + LH+ LH+ + H+ + LH," L + H + + LH+ LH+ + H+ = LHz2+
trienMe,
trienMe,
trienb
enMe,C enMe,d
9.11(2) 8.35(1) 5.26( 3) 2.24( 3) 9.19( 2) 8.43(3) 5.62(4) 2.73(6 ) 9.95 9.31 6.86 3.66 9.281 6.130 9.979 6.827
12.43(1) 11.39(1 ) 7.18(2) 3.06( 2) 12.55(1) 11.50( 2) 7.67( 2) 3.72(4) 13.59 12.71 9.37 5.00 12.674 8.371 13.627 9.323
a The values in parentheses are standard deviations on the last significant figure. ference 33.
(QN) by a simple quantum-chemicalmethod including only inductive effects (ref 9 and references therein). This method has proven to be quite reliable in determining charges, bond energies, and protonation enthalpies (ref 9 and references therein) of saturated molecules, where inductive effects play a major role. In Figure 2 is reported a plot of Q N vs. the experimental stepwise enthapies of protonation for trien, I and I1 in the assumption that terminal nitrogen atoms are always the first to be protonated. The good correlation coefficient obtained shows that for polyamines also the protonation enthalpies depend mainly on the net charge on nitrogen, which are, in turn, related to inductive effects. We note that only the above protonation mechanism leads to a linear relationship between Q N and AHo. Protonation Entropies. On the whole, also the protonation entropy values of I are intermediate between those of trien and 11. It is interesting to note that, while in the first two protonation steps the entropy values of I closely approach the corresponding values of 11, in the last two steps the entropy values of I are more similar to those of unsubstituted trien. Once again we find a hint that the two first protonation steps always involve terminal nitrogen atoms, while the inner nitrogens are protonated in the last two steps. An amine is solvated in water mostly via its nitrogen atoms.37 These act as centers which tend to order the neighboring H20 molecules by forming hydrogen bridges. This kind of ordering effect is lost when a proton is transferred to the base: RNHJaq) + H+(aq) * RNH3+(aq) The onium group is more largely solvated than the amino -
-~~
(37) M. T. Emerson, E. Grunwald, M. L. Kaplan, and R. A. Kromhout, J. Phys. Chern., 20, 6307 (1960).
7.28( 10) 7.51(10) 6.60( 8) 3.81( 8) 7.65(6) 8.45( 5) 7.78( 5) 6.54( 3) 11.0 11.3 9.5 6.8 7.40 6.64 9.72 7.64 Reference 35.
17.3( 5) 13.0(5) 1.9(4) - 2.5( 4) 16.4( 2) 10.2(2) - 0.4( 2) -9.5(2) 7.8 3.7 - 2.0 - 8.1 17.6 5.8 13.0 5.6 Reference 10.
Re-
-AH 11
V
Figure 2. Stepwise enthalpies of protonation (kcal/mol) vs. net charges of the nitrogens which undergo protonation: (0)trien, (H) trienMe,; (A) trienMe,.
group7 (Le., its orienting effect on the solvent is greater), but much less solvated than the H+ ion.38 Therefore, a positive entropy change results from the release of water molecules, or, in other words, from the decreased orientation of the solvent. It has been shown that ASovalues increase in the following order: primary amines < secondary amines < tertiary amines. In particular, ASo has been found to increase by about 5 cal mol-l mol-' K-l for each substitution of a hydrogen by an alkyl (e.g., methyl) group.20 This trend has been related with the lower solute-solvent interactions of hydrophobic alkyl groups in respect to positive amine nitrogen atoms.20 In the present case, for the first two protonation steps, -
(38) R. P. Bell, "The Proton in Chemistry", Methen, London, 1959, pp 23-24.
68
J. Phys. Chem. 1981, 85,68-75
since we are comparing a monosubstituted nitrogen atom (trien) with a trisubstituted nitrogen (I and 11)a difference in ASo of about 10 cal mor1 K-l was expected. Our results (Table 111)are in perfect agreement. With the approach of a second proton, the order around another H30+ion is lost; however, this loss is now partly compensated by the new order which is created by the two charges on the diprotonated molecule. In the case of diamines when the distance between the two nitrogen atoms exceeds four methylene groups, ASoz becomes equal to AS01,9J4 taking into account the correction due to the statistical effect. On the other hand, in the case of t r i a m i n e ~ l ~ l l or ~ ~tetramines ’ ~ ’ ~ ~ ~ ~(see Table 111) ASo2 remains much lower than ASol also when the distance between the two terminal nitrogen atoms (first to be p r o t ~ n a t e d exceeds )~~ four methylene groups. This is due to the ability of inner amino groups to coordinate further water molecules even in the unprotonated form. As a consequence ASoz is much lower than ASol due to an increased order of water molecules around the differently charged (+ or -) amino group^.^ Obviously, the ASo values are also affected by the substitution degree of
neighboring nitrogens. Therefore, both ASo, and ASoz of I1 are lower than the corresponding entropy values of I (Table 111). The further reduction in entropy values on going to the third and fourth protonation may be explained in a similar way. The last two entropy values of I are more similar to those of trien than to those of I1 because these last protonations always refer to inner nitrogens which are disubstituted in the case of I and trien, but are trisubstituted in the case of 11.
Acknowledgment. The financial support of the Italian Research Council (CNR) is gratefully acknowledged. Supplementary Material Available: Individual data points for emf measurements including milliliters of OHadded, emf, total concentration of amine, total concentration of acid, and residual on the total concentrations of amine and acid (i.e., the difference between the analytical values and those computed by the program MINIQUAD 76A) (7 pages). Ordering information is available on any current masthead page.
Oxidation of Amino-Containing Disulfides by B r p and OH. A Pulse-Radiolysis Study A. John Elllot, Roderlck J. McEachern, and David A. Armstrong” Department of Chemistry, University of Calgary, Calgaty, Alberta, Canada, T2N 1N4 (Received; June 12, 1980)
The rate constant for reaction of B r p with dithiodipropionic acid (-4.2 X lo8 mol-l dm3s-l) was independent of pH in the range 6.611.0 and was -4.5 times smaller than those for the neutral dimethyl and diethyl disulfides. Bra. only reacted with the disulfides of cysteamine, cysteine, and penicillamine with an appreciable rate (>lo8 mol-l dm3s-’) when one or both of the amino groups were unprotonated. Homocystine was less sensitive to the degree of protonation. While the reaction of Brz-. with dithiodipropionic acid yielded a transient (A, N 450 nm) which possessed the characteristicsof a disulfide cation, the amino-containingdisulfides cited above produced transients which adsorbed with a ,A, near 380 nm. The latter transients decayed by second-order kinetics over the pH range studied, and they gave no evidence of reaction with OH-. These species have been tentatively identified as perthiyl (RSS.) radicals. Hydroxyl radicals produced composite spectra consisting of the 380-nm species and other transients which absorbed below 350 nm. Oxidized glutathione gave only weak absorptions on reaction with BrZ; and OH. The yields of sulfydryl molecules have been measured for cystine and dithiodipropionic acid. The mechanisms of radiolysis are discussed in the light of present results and earlier studies.
Introduction The 1960s and early 1970s saw a great deal of research on the radiation chemistry of aqueous solutions of sulfur-containing amino acids and related comp~unds.l-~ This arose because of the “in vivo” biological radiationprotection potential of these compounds. Although most research centered on the sulfydryl molecules, disulfides were also in~estigated.~~~ In deaerated solutions containing disulfides, the principal products were the corresponding thiol, sulfinic acid, trisulfide, and products arising from C-S ~leavage.~ Furthermore, the hydroxyl radical ap(1) J. E. Packer in “The Chemistry and the Thiol Group”, Part 2, S. Patai, Ed., Wiley, London, 1974, p 481 and references within. (2) G. C. Goyal and D. A. Armstrong, Can. J. Chem., 53,1475 (1975); J.Phys. Chem., 80, 1848 (1976). (3) T. C. Owen, A. C. Wilbraham, J. A. G. Roach, and D. R. Ellis, Radiat. Res., 50, 234 (1972). (4) J. W. Purdie, J. Am. Chem. Soc., 89, 226 (1967); Can. J . Chem., 47, 1029, 1037 (1969); 49, 725 (1971). 0022-3654/81/2085-0068$01 .OO/O
peared to be implicated to some extent in the formation of all of these product^.^ Until relatively recently, pulse-radiolysis studies of disulfides were concentrated on the one-electron reduction by e,; and COz-. because the disulfide anion absorption, A, = 400-450 nm,5,6was often observed when disulfidecontaining proteins were reacted with e,; and COz-..7-10 However, utilizing pulse radiolysis with both optical and (5) M. Hoffman and E. Hayon,
J. Am. Chem. Soc., 94, 7950 (1972). (6) J. W. Purdie, H. A. Gillis, and N. V. Klassen, Can. J. Chem., 51, 3132 (1973). (7) J. R. Clement, D. A. Armstrong, N. V. Klassen, and H. A. Gillis, Can. J. Chem., 50, 2833 (1972). (8) R. H. Bisby, R. B. Cundall, J. L. Redpath, and G. E. Adams, J. Chem. SOC.,Faraday Trans. 1, 72, 51 (1976). (9) M. Faraggi, M. H. Klapper, and L. M. Dorfman, J. Phys. Chem., 82, 508 (1978). (10) A. J. Elliot, F. Wilkinson, and D. A. Armstrong, Int. J. Radiat. B i d . Relat. Stud. Phys., Chem. Med., 38, 1 (1980).
@ 1981 American Chemical Society
