D. P. STEVENSON, G. 11. COPPINCER AND J. W. FORBES
4350
included. The relative efficiencies of the gases in the three systems seem to follow the same order. It is not possible to extend the comparison to a detailed picture of the transfer of energy with the order of accuracy obtainable a t present.
[CONTRIBUTIOS FROM THE S H E L L
Solvent Effects on n
Vol. s3
Acknowledgments.-The author wishes to thank Professor W. Albert Noyes, Jr., for his advice and encouragement during the course of this work. He is also grateful to Professor D. J. IVilson for many illuminating discussions.
DEVELOPMENT COMPANY, EMERYVILLE, CALIFORNIA]
Transitions of the Bases, Water, Ammonia, Hydrogen Sulfide and Phosphine BY D. P. STEVENSON, G. &I. COPPINGER AND J. L\-. FORBES -+ C*
RECEIVED JUNE 10, 1961 I t is shown that the quartz and near vacuum ultraviolet absorption bands of water, ammonia, hydrogen sulfide and phosphine in water solution show a "blue shift" relative to the location of these absorption bands in the vapor or in aprotctropic solvent solution spectra. The magnitudes of the blue shifts are approximately equal to the heats of solution a t -25' of these hydrides or bases in water. It is suggested that an upper limit to the conceutration of such base molecules in water solution that are not engaged in hydrogen bond complexes with the water solvent can be estimated from the ratio of the absorptivity of the water solution to that of the vapor or aprototropic solvent solution a t a particular wave length. It is show1 that the effect of small concentrations of water on the spectrum of ammonia in otherwise dry diethyl ether are compatible with this suggestion.
It is generally accepted that the lowest energy (longest wave length) electronic excitation process of such substances as ammonia, water, phosphine, hydrogen sulfide, hydrogen chloride and the alkyl derivatives of these substances is associated with the promotion of a non-bonding electron from a 2p or 3p level of the most electronegative atom to an anti-bonding orbital, ;.e., the transition is of the type n -+ r*.lZ2Stated in other terms, the optical electron for the longest wave length ultraviolet absorption band of such substances is one of the unshared electron pair with which the base properties of these substances are associated. Rather naive considerations3 suggest that if an unshared electron pair of a base is involved in hydrogen bond formation there should be a significant blue shift of the associated electronic absorption band. That is, hydrogen bonding should markedly lower the energy level of the non-bonding electrons while having relatively small effect on the energy level of the anti-bonding orbital. To the extent that this conclusion is valid, one may further conclude that spectrophotometric observations of the n 3 U * absorption band of such bases in various solvents should provide a means of distinguishing the degree of specific solvation of bases in prototropic solvents from the non-specific solvation as Brealey and Kasha3 have suggested for the solvent effects on n --t T * transitions. Direct evidence of the general validity of the concept of a dramatic blue shift of n +. 5* absorption bands accompanying solution of a base in a prototropic solvent is provided by a comparison of the spectrophotometric behavior of water vapor with that of liquid water. Wilkinson and Johnston4 have described the absorption spectrum (1) R. S. Mulliken, J . Chem. Phys., 3, 50G (1935). (2) A. D. Walsh, J . Chem. SOL.,2260, 2296 (1953). (3) Based on t h e discussion of t h e effect of solvent polarity o n t h e location of ?z T * absorption hands. H. McConnell, J . Chein. P h y s . , 20, 700 (1982); G. J. Brealey a n d M. K a s h a , J . A i n . Cheiiz. Soc., '77, 4462 (1955). (4) R. G. Wilkinson a n d H. L. Johnston, J . Chem. Phrs., 16, 190 (1950).
-
of water vapo: for wave lengths in the range of 1450 to 1850 A. From their Fig. 3 one finds the absorptivity of water vapor to be 1G2 and 22 l./ mole cm. a t 1800 and 1850 8.,respectively. Ley and Arends: reported measurements on the absorption spectrum of liquid water in the range 18201983 8.; at 1524 8. and 1850 A.they report A = 0.11 and 0.04 I./mole cm., respectively. This decrease in the absorptivity of water that acco5panies its condensation corresponds to a -150 A. blue shift of the long wave length edge of the absorption band. Such a wave length shift in the vicinity of 1800 8. corresponds to -13 kcal./mole stabilization in the liquid phase of the orbital carrying the optical electron. Ley and Arends6 reported on the absorption spectrum of ammonia in hexane (1900 5 X 5 21008.) andammoniainwater (1820 5 X 5 20208.) where the absorptivities are in the range from 2000 to 50 l./mole cm. Their results for the hexane solution spectrum of NH3coincide with the envelope of the minima of the vibronic bands of the ammonia vapor spectrum reported by Tannenbaum, et al.,7 while the aqza ammonia absorption band is displaced about 120 A. to the blzte from that of the hexane solution. The present authors have measured the absorptivities of solutions of ammonia in water and in dry diethyl ether with the results shown as curves I and I1 of Fig. 1. In the region of overlap between our measurements of the aqua ammonia spectrum and those of Ley and the absorptivities agree to about 50y0 corresponding to a difference of 10 to 20 8. a t constant absorptivity. The ether solution spectrum coincides more or less with the smooth curve through the ammonia vapor band spectrum.' In the region of overlap between our ether solution observations and those of Ley and Arends on hexane solutions, we find perhaps a 10 ( 5 ) H. Ley a n d B. Arends, Z . Physik. Chem., B6,240 (1929). (6) H. Ley a n d B. Arends, i b i d . , B17, 177 (1932). (7) E. T a n n e n b a u m , E. M. Coffin a n d A. J. Harrison, J . Chcnl. Phys., 21, 311 (1933).
Nov. 5 , 1961
SOLVENT EFFECTSON TRANSITION OF BASES
A. red shift for the ether solution absorption band
4351
3,000
i
i
that may be due to experimental error. Ley and Arendsa reported measurements of the absorption spectrum of hydrogen sulfide in hexane and water solutions. We have measured the hydrogen sulfide vapor spectrum with the results shown as curve I11 of Fig. 1. The measurements of Ley and Arsnds in hexane solutio? in the vicinity of 2300 A. are shifted perhaps 15 A. to the blue from our vapor spectrum. We have also measured the absorption spectrum of hydrogen sulfide in aqueous 0.01 N and 0.1 N HzS04solution with the results shown as Curve I V of Fig. 1. The dilute acid was used as a solvent for the measurement of the aqueous hydrogen sulfide spectrum when we found very large apparent deviation from Beer’s law in measurements of water solutions. For example, we found the apM H:S in parent absorptivity of a 1.5 X water solution to be 135 l./mole cm. a t 2300 A. Il I\ 1 I l \\ 03 while a t this wave length the apparent absorptivity of a 1.3 X M solution in water was found to be I I 0 lL I 1 I \I 0.1, 390 l./mole cm. The difficulty is clearly due to ..1900 2000 2100 2200 2300 2400 2500 dissociation of the HzS; HS- $as a strong absorpWave length, a.u. tion band with Xmax = 2300 A. and A m a x E 5000 l./mole cm.8 In the acid solutions the dissociation Fig. 1.-Segments of the quartz region ultraviolet absorpis suppressed completely and no apparent deviation tion spectra of ammonia, hydrogen sulfide and phosphine: from Beer’s law was found a t either acid concentra- curve I, KH3 in dry diethyl ether; curve 11, NH3 in water; tion. curve 111, hydrogen sulfide vapor; curve IV, hydrogen sulThe water solytion absorption band of H2S is fide in dilute sulfuric acid; curve V, phosphine vapor; curve shifted about 50 A. to the blue relative to the vapor VI, phosphine in water. All measurements were made with band, corresponding to about 4 kcal./mole stabili- a Cary Model 14 ultraviolet spectrophotometer. The zation of the non-bonding electron (unshared pair) dotted portion of I V from Ley and Arends, ref. 8 of text. orbital. Curves V and VI of Fig. 1 show the results of attention on ammonia. Aqua ammonia usually is discussedlc in terms of the two equilibria measurements of the absorption spectra of phosphine in the vapor state and in water solution reNHa(free) HzO NHiOH (1) spectively. We have been able to find no prehX40H NH4+ OH(2) viously reported quantitative measurements of the electronic absorption spectrum of this substance. and over the years the conclusion of Moore and Further, we found no evidence of the diffuse bands Winmillll that the equilibrium constant for (l), K Cheesman and Emeleusgreported for the 2200-2300 TABLE I A. region. The absorption kand of PH3 in water BLUE SHIFTS vs. TROUTON’S COXSTANT A N D HEATS OF appears to show a small 30 A. blue shift from the SOLUTIOS~ position of the band in the vapor spectrum (corAS vap. ( b . p ) , - A H sol., “Blue shift,” responding to 2 kcal./mole). e.u. kcal./mole kcal./mole The reader may have noted that the energy CH4 17.51 ..... .. equivalents of the blue shifts of the absorption ”3 23.28 8.28 -10 bands of the four substances, H20, NH3, HeS and H20 26.04 4-10.51 -13 PH3,in water solution relative to vapor or non-polar 1.96 .. Ne 17.54 solvents parallel in excellent fashion the tendency of SiHa 18 ..... .. the four substances to associate in the liquid phase PH3 18.82 ? -2 as measured by the departure of the Trouton’s HzS 20.97 4-4 . 6 -4 constants of the substances from those of the cor17.85 2.P .. Ar responding iso-electronic rare gas or non-polar hya A S vap. and AH sol. (in HzO) from N.B.S. Circular dride. Similarly, the energy equivalent of the blue No. 500. * W. Latimer, “Oxidation Potentials,” 2nd shift is approximately equal to the heat of solution Ed., Prentice-Hall, New York, N. Y . , 1952, p. 367. of the “base” in water in those cases where the is approximately equal to 1, has been heats of solution are known. The data pertinent to - [NH40H1, .TNHz- 1_ these conclusions are summarized in Table I. generally accepted. It is not clear to the authors I n order t o show the potential utility of the data what structural concept should be associated with in Fig. 1 for the discussion of the specific solvation “free” ammonia molecules that have been assumed of bases in prototropic solvents, i.e., for the discus- to constitute about 50% of the ammonia in the sion of hydrogen bond formation, we will focus \
I
+
+
+ +
+
( 8 ) H. Ley and B. Arends, 2 . Physik. Chem., B16,311 (1932). (9) G. H. Cheesman and H. J. EmelCus, J . Chem. Soc., 2847 (1932).
(10) See for example, N. V. Sidgwick,“Chemical Elements and Their Compounds,” VOX.11, Oxford Press, New York, N. Y., 1950,pp. 88C--881. (11) T. S. Moore a n d T. F. Winmill, J . Chem. Soc., 107, 1G35 (1912).
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EPHRAIRI BEN-ZVJ
AND
THOMAS L. ALLEN
Vol. 53
aqueous solutions. However, it seeins reasonable vations were mnde of the effect on the aiiirnonin to assume that the implication of “free” is that such spectrum of the addition of water to diethyl ether molecules are not specifically solvated but are solutions of ammonia. There were measursd the simply subjected to the forces of the general dipolar optical densities a t 2125, 2150 and 2175 A. (1 0 environment of the medium, water. If “free” has cm. cell), of ether solutions of NH3 with various this meaning, then it seems reasonable to assume concentrations of water. As may be seen in Table that the spectral behavior of such “free” ammonia 11, the presence of 5 to 10 ml. of HZO per liter of molecules would be very like that of ammonia ether causes a large reduction in the effectivc. abmolecules in diethyl ether solution, where there are sorptivity of the ammonia solutions a t these wnvc the strong general dipolar forces associated with lengths. the C-0 bonds but no possibility of specific solvaTABLE I1 tion such as is implied in the formula NHbOH. l 2 Subject to the assumptions indicated in the pre- E F F E C T O F HsO O N THE SPECTRUM O F ~1~~I N DIE,TIITL ETHER vious paragraph, one can conclude that the absorpM 0.0139 0.0130 0.1::o tion spectrum of aqua ammonia, as given by Curve ? f (”I), __ Co (HzO), M 0.0000 ,0334 I1 of Fig. 1, is the sum of two quantities A ( a q . NHs)
= (1
-
W )
‘lNIi8
+
O.D. (1 cm.)
aAhIIpOIi
(3)
2125 2150 2175 K
1.460
0 . (io5
,686 332 . 1.A7
. ZY:,
.2fj
0.337 ,132 where (Y is the fraction of the dissolved ammonia in 0.454 + 0.012 0.406 i 0 . 0 0 5 the form of NHdOH, ie., specifically solvated, 3 . 4 i 0.1 2.6 * 0.1 ANH*OH is the absorptivity of NH40H a t the wave If one assumes the optical density of each of these length A, and is the absorptivity of “free” solutions is proportional to the concentration of ammonia molecules and is to be taken equal to that “free” ammonia in the solution, one can calculate of a diethyl ether solution a t A. Since OIAXH~OII is of necessity a positive quantity, we h a r e from (3) an apparent association constant for the reaction for the fraction of free ammonia molecules, 1- a , XHd H20 S H a O H ( i n ether) (:) the inequality from the expression
+
and the conclusion that less than 0.5y0of the animonia in water solutions is “free.” Similar reasoning applied to the data on tlie vapor’ and water solution spectra of water, phosphine and hydrogen sulfide leads to the conclusion that less than 0.2% H20, 30% PH3 and 4276 H,S are “free” in water solution a t about 23”. In order to ascertain that the decrease in absorptivities associated with change of solvent from ether to water in the case of ammonia is appropriately attributable to specific complex formation, obser(12) By t h e f o r m u l a , XHiOH, we imply an “ o u t e r complex” in t h e sense of Fig. 2 of t h e definilivc paper of R. S . hlulliken, J . P h y s . Chem., 66, 801 (1952).
since the total NHs concentration in the solution is small compared with the water concentratioii. As may be seen in the last line of Table 11, tlic spectral data give -3 l./mole for the (NHa--I-IGQ) association constant in the ether solution. Such a value for the association constant would imply that considerably less than lY0 of the ammonia in water solution is free, if the association constant were independent of solvent. Clearly the results of the measurements on the effect of water on the spectrum of the ether solution of ammonia are completely compatible with the interpretation given above of the difference between the spectra of the water and the ether solutions of ammonia.
[COSTRIBU TION PROX THE DEPARTXENT OF CHEMISTRY, v s I V E R S I T Y O F CALIFORSIA,
DAVIS,CALIFORNIA]
The Oxidation of Oxalate Ion by Peroxodisulfate. 11. The Kinetics and Mechanism of the Catalysis by Copper(11) BY EPHRAIM BEN-ZVI~ AND THOMAS L. ALLEN RECEIVED MAY20, 1961 The oxidation of oxalate ion by peroxodisulfate is strongly catalyzed by copper( 11),but the catalysis is subject to inhibition by molecular oxygen. An investigation has been made of the kinetics of the catalyzed reaction in the absence of oxygen. The rate law is first-order in peroxodisulfnte, zero-order in oxalate, and half-order in the catalyst. A free-radical chain mechanism, involving oxidation of copper to the terpositive state, is postulated for the reaction. It is shown that copper does not take part in the chain-initiating step. Reduction of copper t o the unipositive state prohably does not occur.
I n the first part of this series3 it was shown that copper(I1) is a very effective catalyst for tlic oxi-
dation of oxalate ion by the &Os= and a brief study of the kinetics of the copper-catalyzed re-
(1) T h i s work was assisted b y a research g r a n t from tlie National Science Foundation. ( 2 ) Abstracted in p a r t from t h e P h . D . Dissertation 01 Ephraim Ben-Zvi, University of California, Davis, 1900. 73, 3ZSQ (1951). ( 3 ) T. L. Allen. J. A m . Chem. SOL.,
(4) I n common usage 5 2 0 8 - is persulfate. Under t h e “ 1940 Rules” [W. 1’. Jorissen, H. Bassett, A . Damiens, F. Fichter a n d 1%.R i m y , ibid., 63,889 (1911)) a n d in Chemical Abstracts i t is designated peroxydisulfate. I n t h e ’‘ Definitive Rules for Nomenclature of Inorganic Chemistry,” i b i d . , 8’2, 5j23 ( l 9 0 0 ) l i t is peroxodisulfate.
