426
D. L. LEUSSING AND JERROLD JAYXE
water mixture. Both hydrogen azide and water can react with the lead nuclei which are necessary for the catalytic acceleration of the reaction. This observation is confirmed by the present data. In the case of the untreated lead azide where the initial gas evolution is lorn ( a = 0.007) the autocatalytic reaction cannot be stopped completely. However, the induction time becomes longer than in the case where the hydrogen azide-water mixture has been frozen out. I n addition, the reaction rate is lower and the sample does not explode. In all cases where the initial gas evolution attributable to the hydrazine hydrate treatment and the ensuing formation of hydrated basic lead azide was high (a = 0.06-0.08), a complete poisoning of the autocatalytic reaction was observed if the gasses were not frozen out. This poisoning effect occurs also when the liquid nitrogen is removed at higher conversions, permitting the hydrogen azide-water mixture to exert its influence on the lead azide. We then observe a reduction of the maximum reaction rate and a shift of the curve toward longer times (curve 7, Fig. 5 ; curve 11, Fig. 6). The decomposition of all samples where hydrazine hydrate was involved was less complete ( a , around 0.85) than the decomposition of the untreated lead azide ( a , around 0.95). This is attributed to the presence of the basic lead azide which, due to its lower azide content, is less likely to decompose, even at 300'. When the hydrogen azide-water mixture is removed from the silver-coated samples, the silver shows its catalytic activity. The induction periods become smaller than those of the untreated material. In the initial stages of the reaction the curves are practically identical regardless of whether the amount of silver is 1.0 or 10.0 atomyo. This suggests that it is not the amount of the external catalyst that is important, but the contact area between the silver grains and the lead azide, which should be not much different for 1 and 10 atom% of Ag for the following reason: After the first silver
Vol. 66
azide has been formed on the lead azide surface, it is more likely that further silver azide is deposited on its own lattice than on new sites on the lead azide surface. This arrangement will not be changed when the silver azide is reduced to metallic silver. The good reproducibility of the curves for the silver-coated samples in the initial stage of decomposition as well as the sudden rate increase (Fig. 8), which cannot be observed in either the untreated or the hydrazine-treated material, are not inconsistent. Upon completion of the outgassing process there exists a relatively large metal-lead azide interface in the silver-coated samples. In the other cases this interface will have t o be formed by the creation and growth of lead nuclei. Obviously the area of the preformed interface with the silver-coated samples is appreciable and permits good reproducibility of the curves and a high rate of reaction in the initial stage. It must be pointed out that the maximum rate is smaller with the silver-coated samples than with the untreated lead azide, in contrast to the respective behavior in the initial stage as pointed out above. It is assumed that two different propagation mechanisms are operating. I n the early stages the reaction front originating from the metallic nuclei is confined t o the individual lead azide grains. Later the growing metal specks come into contact with still undecomposed portions of adjacent lead azide grains, and the reaction propagates via a bridge mechanism. Due to the treatment with hydrazine hydrate the silver-coated samples are covered with a surface layer of basic lead azide. Furthermore, the silver grains on the surface separate the individual lead azide crystals, which results in the reduction of the bulk density, as has been found. Both factors reduce the probability of the bridge mechanism occurring. This is not the case for the untreated lead azide, and the bridge mechanism can contribute to the reaction rate so as to result in an explosion.
MONONUCLEAR AKD POLYNUCLEAR COMPLEX FORMATION BETWEEN IROX(I1) ASD ~,~-DI~!vIERCAPTO-~-PROPAKOL BY D. L. LEUSSING~ AND JERROLD JAYNE Department of Chemistrg, University of Wisconsin,Madison, Wis. Receiued August 849 1861
Iron( 11) and 2,3-dimercapto-l-propanolions react t o form intensely colored dark red polynuclear complexes conforming to the series, DMP(FeDMP),', and a mononuclear complex, Fe(DMP)2', which has a less intense red color. The formation constant for Fe(DMP)2' was found t o be about 6 X l O + l 5 in 0.10 M potassium chloride at 30".
Recent investigations have shown that mercaptide containing ligands form polynuclear complexes with some of the divalent metal ions of the first transition series. These polynuclear complexes can be characterized using the "core plus links" postulate of S i l l h 2 The bidentate monothiols p(1) To whom inquiries should be addressed a t the National Bureau of Standards, Washington, D. C. (2) L G. SillCn, Acta Chem. Scand., 8 , 299, 318 (1954).
mercaptoethylamine3 and mer~aptoacetate~ react with nickel(I1) to form complexes belonging to the series Ni [Ni(RS)2],2-2n while 2,3-dimercapto-1propanol (DMP) forms complexes of the type DMP(MDMP),- where M is zinc(II)6 or nickel (3) D. C. Jicha and D. H. Busch, 135th National Meeting of the American Chemical Society, Boston, Mass., April, 1959. (4) D. L. Leussing, R. E. Laramy, and G. S. Alberts, J. Am. Chem. Soc., 82, 4826 (1960).
Ma,rch, 1962
427
COMPLEXES O F I R O N .4XD 2,3-DIMERCAPT0-1-PROPA4KOL
I .o (II).6 Mercaptide ions act as bridging groups in these polynuclear complexes and a feature has been 0 postulated to be d,d-n bonding where electrons are donated from the metal ion d orbitals to the vacant sulfur d orbitals.6,' I n this respect it is significant that manganese(I1) ions with only a half-filled d shell weire found not to form polynuclear species under comparable conditions where the nickel(I1) and zinc(I1) complexes obtain. We have undertaken the present study of the iron(I1)-DMP system to characterize further the properties of the mercaptide complexes with the divalent metal ions of the first transition series; in particular to determine with which of the divalent metal ions in the series Mn to Zn the onset of pol ynucl~earcomplexes occurs under the conditions used in lhese studies. The results which we obX tained for this highly air-sensitive system were not Fig. 1.-Complex formation in the region E > 1.00: y = as precise as those which were obtained for the 2 - Ti; z = -log (DMP') + log Fer; A, Fet = 1.98- 2.05 X l o p 3M ; 0,Fet = 5.16 X IOM3M ; c]Fet = 5.80 x zinc(II)-DAIP system. However, conclusions of a 10 - 3 M. semi-quantitative nature could be drawn. calculated is the average for the solution and solid phases Experimental Ferrous chloride stock solutions, approximately 0.1 M , were prepared freshly for each series of experiments by dissolving C.P. iron powder in a slight excess of hydrochloric acid under a nitrogen atmosphere. After completion of the reaction (24 hr. with continual agitation or several days with frequent shaking) aliquots of the solutions were analyzed for iron and chloride. The difference between the chloride as ferrous chloride and the total chloride was taken as the excess of hydrochloric acid concentration. To study the complexes, the techniques and procedures used were essentially the same as those used earlier.6 The titration-pH method was used but each point on the titration curve was the result of a single batch-wise experiment because of the necessity of equilibrating many of the solutions with a solid phase, the high sensitivity of the system to oxygen, and the general sluggishness of D M P systems. Two levels of iron concentration were used, approximately 2.5 or 5 mM. The D M P concentration in each experiment was maintained a t a level that was two or three times that of the iron. The required volumes of ferrous chloride, DMP, and standard potassium hydroxide solutions were mixed under air-free conditions with sufficient water and potassium chloride to bring the final volume to 100.0 ml. and be 0.100 M in potassium chloride. After equilibrating a t least 36 hr. in a water bath at 30" the pH of each solution was determined. Values of 3 and (DMP') were calculated using the equations
W
=
DMP+,
- [l+Iil.+ UH 2DMPt - OHt
+ [OH-]
- [H+l
K2a.
2+z
1
(1)
(2) where the subscript t designates the initial analytical concent,ration of the substance indicated, the quantities in parentheses refer to equilibrium concentration, and U H refers to the hydrogen ion activity as calculated from the measured pH. The value of OHt was corrected for the excess acid in the ferrous chloride solutions. The values of K I , and K z , were determined to be 2.03 X 10-9 and 1.91 X respectively, in 0.100 M potassium chloride a t 30.0'. It should be noted that the assumption was made in the derivation of eq. 1 and 2 that both thiol protons of a complexed D M P molecule are replaced. Also the value of w so ( 5 ) D. L. Leussing and T. N . Tischer, J . A m . Chem. S O ~ .89, , 65 (196:l). (6) D. L. ILeussing, i b i d . , 80, 4180 (1958). (7) J. Chatt and F. A. Hart, J . Chem. Soc., 2807 (1960).
when both are present. The results for < 1 are given in Table I and those for @. > 1 are given in Fig. 1 as a plot of y, which equals 2 W, us. x, which is -log(DMP') log[FeIt.
+
-
TABLE I THE APPARENTVALUESFOR THE SOLUBILITY PRODUCT, K,, = (Fe++)(DMP-) K,,x l o t 1 3 Fet DMPtv M M 103 x 108
x
2.01 2.01 4.76 5.80 4.76 2.01 5.16 4.76 5.16 4.76 2.01 2.01 2.04 4.76 1.98
5.77 3.46 15.00 14.58 15.00 5.77 15.00 15.00 15.00 15.00 5.77 3.46 3.59 15.00 3.60
OHt M
x
108
0.31 0.31 1.77 2.00 3.67 1.81 5.16 5.57 7.15 6.99 3.30 3.30 3.34 8.42 3.62
PH
5.73 5.89 5.52 5.62 5.70 6.23 5.78 6.01 5.99 6.10 6.59 6.82 6.76 6.24 6.97
-log
(DMP-)
it
10.20 0.08 10.12 .OS 10.22 .17 10.04 .20 9.89 .38 9.26 .45 9.76 .50 9.30 .58 9.37 .69 9 . 1 5 .73 8.62 .SO .81 8.52 8 . 6 1 .81 8.90 .88 8.22 .91
Fe: DMP-
Fe: DMP'
1.2 1.4 2.4 4.4 2.6 6.0 4.9 10 6.8 8.9 8.3 12 9.2 6.9 10
1.1 1.4 2.4 4.5 4.1 6.9 5.2 12 9.4 13 15 19 16 16 28
1:1
1:1.2
Results and Discussion Qualitatively, the following behavior was observed. At th.e lowest value of ri (0.08) a small amount of precipitate was obtained and the solution was tinged faintly red. As A increased both the amount of precipitate and the color intensity of the supernatant liquid also increased. The solutions were dark red at ri values around 0.8 and from visual observations no changes in color intensity occurred from this point until ri attained a value greater than 1.4. Above this value the color intensity decreased with increasing f i until at equals two the color had changed from a dark red to a light red. No furt'her color change occurred as the solutions were made more basic except in the most alkaline solut'ions (pH ll), where a slight yellowish tinge was assumed by t'he red solutions. The precipitate appeared to be absent at a equal to 1.2 and greater. As in the case with zinc(II), it is necessary to know this transition point in order to calculate t,he sta-
D. L.LEUSSING AND JERROLU JAYNE
428
bility constants, but because of the deep color it was not possible to determine visually the point a t which the turbidity disappeared. This point was taken as the minimum value of a where centrifuging failed to bring down solid material. The FZ values in Table I and as implied in Fig. 1 are seen to depend on both the (DMP-) concentration and the level of iron. This behavior is expected where the solution species are in equilibrium with a solid phase but when the solid phase is absent polynuclear complexes are indicated. Polynuclear complexes also are suggested by the observations described above, since if only mononuclear complexes were formed a progressive color decrease should be noted with increasing values of beginning at 1.OO. The data in the region a > 1 will be discussed first since their treatment is the most straightforward and the results are similar to those observed with the zinc(I1)-DMP system. I n Fig. 1 the points which would fall along separate curves in the usual plot of a us. -log (DMP-) are seen to fall aproximately along the same curve regardless of the iron level. This is in agreement with the behavior expected for the formation of a "core plus links" series DMP-(FeDMP),-. Employing t h e curve fitting methodsZv5a curve was fitted to the points in Fig. 1 using the relationship qn = qoqn * with qo equal to 2.0, yielding a value of q equal to 1.3 X The solid line in Fig. 1 represents the curve calculated using these values of q o and q. This is the same relationship that was found to describe the behavior of the zinc(I1) polynuclear complexes but in this latter case values of PO and q equal to 1.5 and 1.4 X 10-6 were obtained. Since the shape of the curve in Fig. 1 depends on PO, the nearly identical q values show that for a given degree of complexing the distribution of the various species is about the same in the iron(I1) system as with zinc(I1). Thus, as was shown with zinc(II), complexes containing a fairly large number of iron(I1) ions (up to about eight) are appreciably stable in the solution phase. The absolute stabilities of the iron(I1) complexes are, of course, lower than those of zinc(I1). As 5 amroaches infinity, the ratio of Fe:DMP in the polynuclear complexes approaches 1: 1. This yields a neutral species which most likely comprises the solid phase as was demonstrated with nickel(II)6.9 and zinc(II).6 On the basis of this assumption values of the solubility product, (Fe++) (DMP-), were calculated for the two extremes; (a) ferrous and DMP= ions combine in a 1:1 ratio and with regard to the unprecipitated iron the concentrations of polynuclear species are negligible compared to that of free ferrous ions; (b) ferrous and DRIIP- ions combine in a 1:1.2 ratio'" and the complexes are in equilibrium with a minute but negligible amount of solid FeDMP. The apparent values of Ksp are given in Table I. It is seen that (8) The quantity pn is the equilibrium constant for the reaction l)Fe(DMP)%"% D;MP*(FeDMP),' n DMP". (9) P. Zuman and R. Zumonova, Tetrahedron, 1, 289 (1957). (IO) In the presence of solid FeDMP, the equilibria (n 1) FeDMP.-DMP*(FeDMP)nF e + + hold. Therefore, the ratios of the ooncentrations of the various divalent species are constant as long aa solid is present and the average composition of the complexes in solution also remains constant. (n
+
+
+
+
Vol. 66
for neither case constant values are obtained. I n the lowest pH range the "constants" become smaller with decreasing pH. We have not been able to account quantitatively for the drifting "constants" but a reasonable explanation lies in the formation of protonated complexes, which are not taken into account in eq. 1 and 2. This complication does not present a major difficulty because the results indicate that protonated species are absent in the region a > 1. These species then must disappear as approaches 1 from the low side. Therefore, eq. 1 and 2 should give increasingly more valid results as a increases and the true value of K,, should be approached in the limit and bracketed by the apparent values calculated for (a) and (b). A system in which the over-all value of a equaled 0.73 was analyzed and it was found that about onehalf the iron had precipitated from solution, presumably as the 1:l neutral complex. This result indicates that a closer description of the system is given by case (a). Further evidence favoring this conclusion lies in the fact that the K,, values calculated for case (a) appear to lie randomly scattered about a common value in the higher a region, while those values calculated for case (b) continue to increase with increasing 12. Accordingly, we estimate the value of KsPto be about 10 X 10-13. The value of Qz, the constant for the equilibrium Fe++ SDMP- e Fe(DMP)2-, was calculated to be 6 X 10+l6 using the equation QZ = usol/qKsp. The value of usol was obtained from the data for those solutions where the solid phase just redissolves (a equals 1.2). The method of calculation is described in ref. 5. A rough estimate of 5 X 10+lefor Qz can be made using the results 3 X 10+loand 2 X for the formation of Mn(DMP)2- and Zn(DMP)2-.11 Although the agreement between the two results for &z is only fair, at least some assurance is given that the experimental value of QZ is of the correct order of magnitude and therefore the assumptions made in its calculation are valid to a first approximation. The constant for the reaction 2Fe++ 3DMPFe2(DMP)r is calculated to be about That for the formation of Zn2(DMP)3- is about 4X The ratio of the one-fourth root of the constant for Fez(DMP)3' to the one-fourth root of that for Zn2(DMP)3= is about 1:10+3. For the monomeric bis-DMP complexes of iron(I1) and zinc(I1) the ratios of the square roots of the formation constants is also about 1:1Of3. From this comparison it appears, at least within the limitations of comparing constants for reactions in which different numbers of particles are involved, that the low result for Fez(DMP)3= is due to nothing more than those factors which normally operate to give a lower stability of iron(I1) complexes relative to those of zinc(I1). The relatively strong absorption of the polynuclear species relative to the mononuclear suggests that .rr-bonding also occurs in the iron(I1)D M P system. If so, the parallel decrease in the stabilities of the polynuclear and mononuclear complexes in going from zinc(I1) to iron(I1) can be
+
+
(11) D. L. Leussing, Talanta, 4, 264 (1960).
March, 1962
ACIDIONIZATION CONSTANTS IN H20 AND D 2 0
explained as a manifestation of the “synergistic effect”I2 existing between c and T bonds. (12) L.E, Orgel, “An Introduction t o Transition Metal Chemistry,” John Wiley and Sons, New York, N. Y., 1960.
429
Acknowledgment.-We wish to thank the National Science Foundation and the Wisconsin Research Foundation for grants which supported this and related work.
R,ELATIVE HYDROGEN BONDING OF DEUTERIUM. 11. ACID IONIZATION CONSTANTS IN H20 AND DzO’ BY A. 0. MCDOUGALL AND F. A. LONG Department of Chemistry, Cornell University, Ithaca, N . Y . Received August $0, 1961
Studies have been made of the ionization constants of a variety of weak acids in the solvents HzO and DzO by glass electrode or b y conductometric procedures. Emphasis has been on acids which in either the acid or conjugate base form would be expected to form intramolecular hydrogen bonds. These data, when combined with data of other workers, offer support to the proposal that for weak acids in general ( ~ K D A ~ K H Aincreases ) with KEA. However, the data also support the suggeEtion that the reference line of such a correlation is different for acids o f different types. Several of the present results are consistent with earlier studies on maleic acid in that intramolecular hydrogen bonding, relative to bonding to the solvent, appears t o be weaker for deuterium than for hydrogen. However, there are enough exceptions to this rule to suggest that a true picture will not be obtained without more explicit consideration of the properties of particular solutes as well as of the solvent differences in the two cases.
Reasons for interest in the relative abilities of hydrogen and deuterium to form hydrogen bonds have been summarized in Part I of this series.2 Part I also reported measurements of the ionization constanta of acids, one of which involved an intramolecular hydrogen bond; it was found that the ratio of these constants in water and deuterium oxide, for any particular acid, varied according to whether an intramolecular hydrogen bond was concerned in the ionization. The comparison in such a case is of the competition between internal hydrogen bonding and hydrogen bonding with the solvent, and it is not a direct measure of the strength of the intramolecular hydrogen bond invol~ed. The present work extends this comparison to other ,acids, some of which are hydrogen-bonded, and considers the results on the basis of a general correlation for the dissociation of acids in water and deuterium oxide. Such a relationship first was postulated by Rule and LaMer,3who found for the small. number of acids at that time investigated that log (KHAIKDA) was proportional to log KHA, KHAand KDAbeing the ionization constants of H- and D- acids in water and deuterium oxide, respectively. This relation has been support’ed by later moinkers4 but HEgfeldt and Bigeleisen5 have suggested recently that the type of acid (e.g., phenol, carboxylic acid, etc.) may change the constant of proportionality, i e . , the slope of the line, and this proposal is supported by the present work. Most of the conventional methods for determination of ionization constants have been applied to studies in deuterium oxide; Dahlgren and Long,2 for example, used e.m.f. measurements with a quinhydrone electrode. A much simpler method is suggesLed by the work of Glasoe and Long6 and (1) Work supported by a grant from the Atomic Energy Commission. (2) G. Dahlgren, Jr., and F. A. Long, J. Am. Chem. Soc., 82, 1303 (1960). (3) C. K.Rule and V. K. LaMer, ibid., 60, 1974 (1938). (4) P. Ballinger a n d F. A. Long, ibid., 82, 795 (1960). (5) E. Hiigfeldt and J. Bigeleisen, i b i d . , 84, 15 (1960). (6:) P. K.IGlasoe and F. A. Long, J . Phys. Chem., 64,188 (1960).
others’vs which has established that satisfactory measurements with deuterium oxide solutions can be made with a glass electrode. The determination of hydrogen ion concentration with a glass electrode is by no means as accurate as that with a hydrogen or quinhydrone electrode but it was hoped that satisfactory pK differences could be obtained readily for a series of acids. Studies now have been made for a number of carboxylic acids and phenols, several of which involve intramolecular hydrogen-bonding. In addition some data have been obtained for certain other acids (not hydrogen-bonded species) by conductance measurements, which method of course is considerably more accurate than the glass electrode procedure. Experimental Materials.-Reagent grade inorganic chemicals were used throughout. Deuterium oxide was supplied by the Liquid Carbonic Company; it contained at least 99.5% DzO. Salicylic acid (m.p. 159-160”), o-nitrophenol (m. 46”), p-nitrophenol (m.p. 112-1 14”), 2,4-dinitrophenof (m.p. lll’), 2,6-dinitrophenol (m.p. 62-63’), and y-resorcylic acid (2,6-dihydroxybenzoic acid) (m.p. 167”), were purified by recrystallization from water. Glycolic, oxalic, iodic, and chloroacetic acids were used without further purification except that the oxalic acid first was dehydrated over sulfuric acid. Reagent grade phosphoric acid (containing not less than 85% H,PO,) was employed. Sodium hydroxide solutions were British Drug Houses “Concentrated Volumetric Solutions” diluted appropriately. A stock solution of sodium deuteroxide in deuterium oxide was made up by allowing metallic sodium in toluene to react with boiled-out deuterium oxide in a separatory funnel. This stock solution was diluted according to requirements. All solutions were made up with boiled-out distilled water or deuterium oxide, and were estimated either by titration against standard base or in the case of certain of the phenols by a bromination procedure or by spectrophotometry. pH Measurements.-The potential was measured with either a Beckman Model G pH meter or (in some of the later experiments) a Cambridge Instrument Company Research Model Electron-ray pH meter. The electrode assembly, consisting of a Beckman No. 30167 glass electrode and a
.
(7) R. Lumry, E.L. Smith, and R. R. Glante, J. Am. Chem. Soc., 18, 4330 (1951). (8) K. Mikkelson and 8. 0. Nielsen, J. Phys. Chem., 64,632 (1960).
