August, 1963
KINETICSOF SILVER(II)IN PERCHLORATE MEDIA
freezingz7might then be taken as an indication of the enhancement of such migration in the solid phase. There is some recent evidence that energy transfer to solutes in aromatic media is indeed accelerated by crystallization. 28 Conversely, the absence of LET effectsz9and indifference to phase boundaries of product yields from aliphatic media3*are possibly then symptomatic of rapid energy migration even in their liquid phases. That the liquid aliphatics should be superior to their aromatic counterparts in this respect is contrary to one's intuitive expectations, grounded iii small quantum phenomena. It is further consistent, however, with other evidence of apparently facile energy transfer in aliphatic ~ o l u t i o n s . ~ ~ ~ ~ ~ Paradoxically, those excitations having high oscillator strengths and thus most likely to be induced (vide i n f r a ) are also just those, in most theories of energy transfer33 and collective excitation^,^^ which should (27) J. Y . Chang and M. Burton, 137th National Meeting, American Chemical Society, Cleveland, Ohio, 1960. (28) E. Collinson, J. J. Conlay, a n d F. S. Dainton, Nature, 194, 1074 (1962). (29) H. A. Dewhurst and R. H. Schuler, J . Am. Chem. Soc., 81, 3210 (1959); R. H. Schuler, J . Phys. Chem., 68, 925 (1959); R. H. Schuler and A. 0. Allen, J . Am. Chem. Soc., 77, 507 (1955). (30) H. Hamashima, M. P. Reddy, and M. Burton, J . Phys. Chem., 62, 246 (1958). (31) T. J. Hardwick, ibid., 66, 2132 (1962). (32) P. J . Dyne and J. Denhartog, Can. J . Chem., 40, 1616 (1962).
1617
couple most readily with the surrounding medium. A resolution will necessarily require a detailed understanding of intramolecular processes. Tentatively, one may ascribe the apparent lack of coupling in aromatic systems to their expectedly efficient predissociative and internal conversion mechanisms.35 A parenthetical remark may be made concerning reaction mechanisms in these mixtures. To account for LET and isotopic distribution effects, various aut h o r ~ " ,have ~ ~ envisaged hydrogen and acetylene production as arising through bimolecular reactions of the type
B*
+ B* +products
Such a mechanism probably is incompatible with the apparent first-order dependence of acetylene production on precursor generation (Fig. 2) and with the likely role of hydrogen atoms in the production of hydrogen gas, as discussed above. The concurrent LET effects on hydrogen and acetylene formation must arise then through other, more complicated, track reaction sequences or through track-associated thermal effects. (33) M. Burton, et al., Ed., "Comparative Effects of Radiation," John Wiley and Sons, Inc., New York, N. Y., 1960. (34) U. Fano, P h y s . Rev., 118, 451 (1960). (35) H. Sponer, Radiation Res. Suppl.. 1, 658 (1959). (36) J. &I. Scarborough and J. G. Burr, J . Chem. Phys., 87, 1890 (1962).
A KINETIC AND SPECTROPHOTOMETRIC EXAMINATION OF SILVER(I1) I N PERCHLORATE MEDIA BY J. B. KIRWIN,F. D. PEAT, P.J. PROLL, ASD L. H. SUTCLIFFE Donnan Chemical Laboratories, University of Liverpool, Liverpool, England Received January 9, 1965 The absorption spectra of solutions of Ag(I1) in perchloric and nitric acids have been determined in the region 350 to 760 mM. A change in the spectra with anion concentration provides evidence for complex formation. The kinetic data obtained for the decomposition of Ag(I1) in perchloric acid also indicate the participation of perchlorate complexes. A mechanism is proposed for the reduction of Ag(1I) t o Ag(1).
lmtroduction A s part of a detailed study of the silver(1)-catalyzed reaction between cobalt(II1) and chromium(II1) in perchloric acid media, it became necessary to study reactions of silver(I1). This paper reports the results obtained from a detailed spectrophotometric and kinetic decomposition study of divalent silver in perchloric acid along with a spectrophotometric examination o€ divalent silver in nitric acid. Previous work in this field includes the pioneer work of Xoyes, et U Z . , ' . ~ who studied the decomposition of silver(I1) in nitric acid media and also investigated the possibility of the occurrence of silver(I1)-nitrate complexes. They found, from e.m.f. measurements, that between one and two nitrate ions are bound to the silver ions,2 which suggests the occurrence of both AgN03+ and Ag(NO3)z species; this conclusion is based upon the assumption that silver(I1) in perchlorate media is uncomplexed which, as will be shown later in this paper, may be incorrect. (1) A. A. Noyes, C. D. Coryell, F. Stitt, a n d A. Kossiakoff, J . Am. Chem. Soe., 69, 1316 (1937).
(2) A. A. Noyes, D. De Vault, C. D. Coryell, a n d J. J. Deahi, t b d , 59, 1326 (1937).
Probably because of the difficulties in handling the very reactive silver(I1) solutions, no previous work of a quantitative nature has been reported on the absorption spectra of this valency state in any medium. High acid concentrations are essential to keep the rate of decomposition in control. Since silver(I1) is a d9system, it should, in addition to any charge transfer bands, contain an internal d + d transition the strength of which will depend upon the configuration of the ligands around the metal ion, and it should be possible to suggest from this band if any Jahn-Teller distortion, as found for other d 9 systems, is present,. Experimental Materials. Silver(1) Perchlorate.-The B. Newton Maine commercial product was used. The silver concentration was estimated by titration with thiocyanate in the usual manner. Perchloric Acid.-The British Drug Houses Analar reagent was used, the concentrations being estimated by titration with standard caustic soda or sodium bicarbonate solution in the usual manner. The Perchlorates of Sodium, Barium, Calcium, Lithium, Magnesium, and Lanthanum.-The above perchlorates were prepared by the action of perchloric acid on the metal carbonates (the carbonates of the first three elements being Analar).
J. B. KIRWIN,F. D. PEAT,P. J. PROLL, AKD L. H. SUTCLIFFE
1618
3
Vol. 67
acid media whereas the reaction of Ag(I1) with Ce(II1) is rapid and gave consistent results. Spectrophotometry.-The results were obtained from two types of instruments [a] a Unicam S.P. 500 spectrophotometer and [b] a Perkin-Elmer recording spectrophotometer, the latter being used when the rate of decomposition of the silver(I1) product was rapid. Kinetics.-All measurements were made on a Unicam S.P. 500 spectrophotometer at a wave length of 475 mp. The temperatures of solutions \\ere maintained a t 1 0 . 0 5 " using a cell block through which thermostated water flowed.
Results and Discussion (1) Spectrophotometry.-The absorption spectrum of silver (11) in various concentrations of perchloric acid is shown in Fig. 1. As can be seen from this diagram there is a maximum in the spectrum a t a wave length of 475 mp with an extinction coefficient of 140 j= 7 M - l crn.-l, which is clearly a n internal ligand field d d transition. There is also a shoulder a t 575 mp and the tail of a charge transfer band starting at about 380 mp and increasing sharply in intensity toward the ultraviolet region. The position of the maximum iii the visible region a t 475 mp shows no change with changing perchloric acid concentrations but there is a slight change in the shape of the band and aIso large changes in the regions of the shoulder (-575 mp) and the charge transfer band (-380 nip) which is indicative of more than one species being present. This suggests the possibility of silver(I1) complexing with C104-, since any absorption in the visible region must be due solely to Ag(II), as the coiicentration of Ag(II1) present cia the equilibrium
-
350
450 660 Wave length, mp.
650
750
Fig. 1.-The spectrum of Ag(I1) in various concentrations of perchloric acid: A, 6.0 M; B, 3.0 M; C, 1.5 J!f a t 25".
+
350
450
550
650
750
2Ag(II) Ag(U Ag(W is very small* and Ag(1) absorbs oiily in the ultraviolet region. Because of this equilibrium, which as will be shown later is established very rapidly, the applicability of the Beer-Lambert law cannot be checked in the normal manner, but it was found that a t 475 mp the 1 1 6 Ag(I1) was absorbance over the range to proportional to the oxidizing power. Thus if we assume the concentration of Ag(II1) is very small, the above law is applicable to the system. Since Ag(I1) is a d Qsystem it should, in addition to being similar to Cu(II), also be similar to Ti(III), hence it should be possible to calculate the position of the absorption maximum relative to that of Ti(II1) which is a t 490 mp, using the formula
T a r e length, mp.
Fig. 2.-The absorption spectrum of silver(I1) in nitric acid of concentrations: A, 6.0 M ; B, 3.0 114; C, 1.5,%I a t a temperature of 25.0". Nitric Acid and Silver Nitrate.-The Analar Hopkins and Williams products were used. Silver(I1) Perchlorate and Nitrate.-The silver( 11)compounds were prepared by ozonolysis of the silver(1) compound (0.15 to 1.2 M )in the corresponding acid media. Preliminary ozonolysis was always carried out to remove all traces of impurities which react with the silvcr(I1). A method was evolved for estimating the concentrations of silver(I1) by utilizing the oxidation of cerous solutions. The ceric produced was measured either by direct spectrophotometry a t 375, 400, 425, and 450 mp or by titration with ferrous sulfate. The ceric produced was always measured relative to the amount produced by cobalt(II1) perchlorate or nitrate under identical conditions, since the spectrum of ceric is very dependent upon the anions p r e ~ e n t . ~ The estimation of silver(I1) by oxidation of iron(I1) to iron(II1) as previously used,' was found to give very variable results in perchloric ( 3 ) G. (1955).
Harqreaves and L. H. Sutcliffe, Trans. Fapaday Soc., 61, 1105
AE
5epa4
= -
r6
where AE is the splitting of the levels; p is the dipole moment of surrounding water molecules; T is the distance between ligands and central metal ion; a4 is the average 4th power of the radius of a 4d electron; and e is the effective charge on the nucleus of the central metal ion. The predicted peak would occur at 460 nip compared with the value of 475 mp that was observed. By analogy with Cu(II), Jahn-Teller distortion effects are the probable reason for the two bands, ie., that a t 475 mp and the shoulder a t 575 mp, although the presence of different absorbing species cannot be ruled out. However, if it is assumed that the two bands originate from the Jahn-Teller effect alone, then the distortion produced must be the same order of magnitude as that in the case of Cu(11).
August, 1963
KINETICS OF SILVER(II) I N PMWIII,OR.ITIC MEDIA
Figure 2 shows the spectrum of silver(I1) in nitric acid of strengths comparable with those of the perchloric acid. As can be seen from this figure a d + d transition band again occurs, the position of maximum absorption being a t shorter wave lengths than for the perchlorate media, namely, in thc region 380-400 mp, there being a shift in the wave length with nitric acid concentration; this is indicative of complexing. The charge transfer band (not shown in thc figure), also begins at shorter wave lengths than for the perchlorate, namely at 330 mp. Rnothcr important fcature of the nitric acid spectra is the largc over-all increasc in extinction coefficients. Estimates of the oscillator strengths of these transiand for the tions are, for the perchlorate 3.8 X nitrate, using the spectrum obtained with the 6.0 M These values are very largc nitric acid, 5.5 X and hence they both suggest that the observed spectrum is a combination of two bands as is the case for copper(II).4 The absorption spectra in the ultraviolet region could not be obtained in a very satisfactory manner, owing to the unavoidable presence of Ag(1). lpigure 3 shows thc effect of small amounts of nitric acid 011 perchloric acid solutions of silver(II), illustrating the stronger complexing nature of the nitrate ion compared to the pcrchlorate ion, towards silver(J1). While the change of extinction coefficient a t 400 mp with nitrate ion concentration would be sufficiently great for quantitative measurements of complex formation to be made, the difficulty of handling the highly reactive Ag(I1) solutions plus the rapid establishment of equilibrium betivecn Ag(II), &(I), and Ag(I1I) militate against reliable measurements being made. (2) Kinetics.-The kinetics of the decomposition of hg(I1) in perchloric acid solution were followed from the disappearance of the silver(I1) optical absorption at the wave length of 475 mp. The results reported here were obtained by a sampling technique, the solution being kept in blackened flasks owing to the lightsensitive nature of the reaction. Light from the spectrophotometer was found to have no effect on the reaction. The disappearance of the silver(I1) was found to be sccond order under all the conditions mentioned in this section (a typical set of rate data is shown in Fig. 4); no trace of the fourth-order term proposed by Soyes, et al.,l for the reaction in nitric acid solution was observed. Oxidizable impurities, as has been mentioned previously, were removed by a preliminary ozonolysis. The ozone left dissolved in the solution after ozonolysis was removed by passing a stream of oxygen through the solution until no ozone odor could be detected. The products of the decomposition were found to be silver(I) and oxygen. The rates observed were found to bc reproducible to within *5%, and all t1ie;results presented here are averages of a t least three experiments, reactions being followed to not less than 75% complction. No trace of hydrogen peroxide was detected a t any time during the reaction. Figure 5 shows a plot of the rate of decomposition against the inverse of the silver(I1) concentration a t four temperatures in the range 0 to 30°, with sodium perchlorate being used to replace the silver(1) perchlorate in order to keep the ionic strength constant. The linearity of the plots in Fig. 5 suggests that there is (4) J. Bjerrurrl, C. J. 13slllausen, and C, K. Jp'rgenson, Acta Chem. Scand., 8, 1275 (1954).
350
430 330 Wuvc lcngtli,
1619
1130
7,jU
inp.
Fig. 3.--l'he effect of ttic addition of small riinounts of nitric, wid to Ag(I1) in 3.0 rll pcrc.hloric acid: A, 0 3 .If; B, 0.12 121; C , 0.06 .I1 a t a tcmpcrature of 26.0".
1 /
I-
ILdd
2100
JtiOJ
13JO
Time, sec.
Fig. 4.-The inverse of the absorbance of Ag(I1) (measurcd a t 476 mp) plotted against time for the conditions: Ay(1) conceritration = 0.288 M , ionic strength = 4.48 MI tempcraturc = 25.0" and perchloric acid concentrations 1.50 d l (A), 2.00 Jf (B), 3.00 M (C), and 4.20 M (U).
a term [Ag(II)I*/ [Ag(I)] in the rate expression. The over-all activation energy derived from the data shown in the figure is 11 2 kcal. mole-'. Figure 6 shows the plot of the observed second-ordcr rate constant against the concentration of perchloric acid for a constant concentration of Ag(1) at a temperature of 25'. It can be seen that there is a minimum in thc rate at a perchloric acid concentration of 3.0 ill. J'rom the shape of the curve, it is obvious that no one effect could produce this variation in the rate: the variation in the perchloric acid conccntration must have two effects, one having a retarding and one an accelerating influence. Since we are forced by the very nature of this reaction to work a t very high ionic strengths, in order to find the
*
J. B. KIRKIS, E'. D. PEAT,P.J. PROLL, AXD L. H. SUTCLIFFE
1620
Vol. 67
Square of perchlorate concn., SO. 10 20 30
I
0.28
40
I
I
I 0.1
I
I
I
0.2
0.3
0.4
Inverse square of acidity, M .
4.0 Inverse Ag(1) ooncn.,
2.0
6.0 *Vf-1.
Fig. 5.-The plot of the observed second-order rate constant against the inverse of the Ag(1) concentration a t an ionic strength of 7.17 III and a constant perchloric acid concentration of 6.0 M a t temperatures: A, 30.0"; B, 25.0"; C, 17.0"; D, 0.0".
** c
0.14
c
Fig. 7A.-Thc plot of the observed second-order rate constant against the function [H+]-2, a t a constant [ClO*-] of 4.48 M and [Ag(I)] of 0.288 ilf a t a temperature of 25.0". 7B.-The plot of the observed second-order rate constant against the function [C10k-]2 a t a constant [HC104] of 1.5 Af and a constant [Ag(I)] of 0.288 Jf a t a temperature of 25.0".
i
I
4.0
2.0
6.0
8.0
[CIOa-lz/ [H +I?.
Fig. 8.-The plot of the observed second-order rate constant against the function of [C104-]2/[H+]2a t 25.0" for a Ag(1) conare for a constant ionic strength of centration of 0.288MJ 4.48 M ; o are for a variable ionic strength of between 1.8 and 4.80 d f obtained from Fig. 5 and 6.
+
4.0 6.0 Perchloric acid concn., AX. 2.0
Fig. 6.-The plot of the observed second-order rate constant against the concentration of perchloric acid, a t a temperature of 25.0'. The concentration of Ag(1) was kept constant a t 0.288 M.
[a] second order with respect to silver(I1) [b] a n inverse dependence upon silver(1) [c] an inverse square dependence upon the acidity [d] a dependence upon perchlorate ion concentration of either power two or powers two plus one
variation of the rate with hydrogen ion concentration E+ must be replaced by Li+ as this is the nearest replacement one can obtain from the point of view of activity coefficients. Figure 7 h shows a plot of the By analogy with other silver(I1) systems studiedlJ rate against the inrerse square of the hydrogen ion [a] and [b] must obviously enter into the rate expresconcentration at a constant ionic strength of 4.48 M sion via an equilibrium of the type and a perchlorate ion concentration of 4.48 ;II. A ~ O K shown (in Fig. 7B) is the effect of increasing the per2Ag(II) J_ Ag(1) hg(II1) rapid chlorate ion concentration using lithium perchlorate, a t followed by the rate-determining decomposition of a constant concentration of perchloric acid of 1.5 M. silver(II1) Unfortunately, the ionic strength cannot be kept constant, since perchlorate affects the rate (see this figure). k Ag(II1) -+- products rate determining The results have been plotted as a dependence upon the square of the perchlorate ion concentration alSince the concentration of Ag(II1) is small and can though a t low concentrations of perchlorate ion a linear also be assumed to be stationary, the following rate law dependence is reasonable, and any ionic strength effect can be derived is included in this effect. The dependence of the observed second-order rate -d[Ag(II)] - kK[Ag(II)]' constant upon the inverse square of the hydrogen ion at [*%(I)1 concentration is convincing as may be seen from Fig. 7h. Considering observation [c1, the most probable reason for this inverse square dependence upon the The Mechanism of Decomposition acidity is the presence of an equilibrium of the type From the experimental evidence presented above it is obvious that any mechanism proposed to account for Ag3+ H,O A g o + 2Hf the decomposition of Ag(I1) must reproduce the follswing features ; ( 5 ) B. &I. C Q I ~and Q ~A. C. Wrthl, J . Am. Chsm. Soc., 80, 273 (195S)i
+
+
5
+
ADSORPTION OF WATERAND CARBON DIOXIDE BY LINDEMOLECULAR SIEVEX
August, 1963
which has been postulated already6 to account for the effect of p H on the rate of electron exchange between the valency states of silver. Considering [d], the effect of the perchlorate ion on the rate is most likely to be combination specific complexes of Ag(I1) and an ionic strength variation. Evidence for the former conclusion comes from slight changes observed in the spectrum of Ag(I1) on addition of the perchlorate ion. Attempts mere made to keep the ionic strength constant while varying the perchlorate ion concentration by using di- and trivalent perchlorates as well as monovalent ones. Barium, calcium, magnesium, lanthanum, sodium, and lithium perchlorates were employed. Unfortunately, the ionic strength could not be kept constant due to the large variation in the activity coefficients of the various cations a t the necessarily high values of ionic strength. The detailed mechanism probably is K
+ Ag3f rapid A g o + + 2H+ rapid
2Ag2+ I_ Ag+ hg3+
+ HzO
-=
k
A g o + -+ Ag+
KI
+ l/~Oz rate determining
hgain, assuming a stationary state for the concentration of silver(III), the rate lam is (6) J. A. hIcMillan, Chem. Rev., 62, 65 (1962).
1621
-d[ilg2+] - kKKl[Ag2+j2 -
dt
[Ag+l W+IZ
KO satisfactory mechanistic explanation for the perchlorate effect can be put forward. The experimental dependence of the rate on [C104-]2/[H+]2is shown in Fig. 8, hence the full rate law should include this function. Presumably the addition of nitrate ions would replace perchlorate by nitrate complexes hence the form of the rate law would remain unchanged. However, the reactivity of nitrate complexes might be very different and the stationary state assumption might no longer be strictly applicable to the data obtained by Noyes, et u1.l; this might account for the introduction of a fourth-order term in the rate expression by these workers. The suggestion could be tested by studying the influence of nitrate ions on the electron exchange reaction. The effect of varying the perchlorate ion concentration on this reaction would also be interesting. We have found that the addition of small amounts of nitrate ions to the reaction causes a decrease in the rate of reaction, while larger amounts of nitrate ions to the reaction causes an increase in the rate of reaction. These effects are only possible if a t least two complexes of silver(I1) exist, which supports the suggestion made previously by Koyes, et aL2 Acknowledgment.-P. J. P. wishes to thank D. S. I. R. for the award of a postdoctoral fellowship during the tenure of which the work was completed.
,4N INFRARED SPECTRUSCOPIC STUDY OF lTHE ADSORPTION OF WATER AND CARBOS DIOXIDE BY LISDE MOLECULAR SIEVE X1 BY L. BERTSCH~ AND H. W. HABGOOD Research Council of Alberta, Edmonton, Canada Received January 16, 1963 Infrared spectra a t 20" are reported for small amounts of water and carbon dioxide (up t o 1 molecule per cavity) adsorbed on Li-, Na-, and KX-zeolites. The water spectra show no evidence of structural surface hydroxyl but may be interpreted in terms of isolated water molecules adsorbed simultaneously by an ion-dipole interaction with the exchangeable cation and by hydrogen bonding of one of the hydrogens to an oxygen of the zeolite surface. The remaining hydroxyl is free and gives a sharp OH band around 3700 em.-'. Carbon dioxide is both physically adsorbed in a linear configuration at a cation and also chemisorbed in one or more bent configurations which are assumed to be carbonate ions formed by interaction with surface oxide ions adjacent to exchangeable cations. One such species common t o all three zeolites gives bands around 1700 and 1340 cm.-l. On NaX a second species with bands a t 1485 and 1425 em.-' is more strongly held a t low pressures but is readily converted to the first form in the presence of any additional physically adsorbed molecules. Additional bands are also present a t different frequencies for carbon dioxide adsorbed on LiX and KX but these do not show such behavior. Small amounts of preadsorbed water greatly accelerate the rate of carbon dioxide adsorption-presumably by catalyzing the chemisorption step.
Introduction In studies of the surface chemistry of ionic solids, the dehydrated crystalline zeolites are of interest both in their own right as solids possessing crystallographically well defined surfaces and also as examples of possible limiting structures approached by local regions of silicaalumina cracking catalysts. As part of a general study of the adsorptive and catalytic properties of these substances we have examined, using a pressed disk technique, the infrared spectra of the lithium, sodium, aad (1) Contribution No. 202 from the Research Council of Alberta. Presented to the 142nd National Meeting of the American Chemiral Society, Atlantio City, N. J., September, 1962. (2) Researoh Council of Alberta Postdoctoral Fellow, 1960-1962.
potassium forms of Linde Molecular Sieve X containing small amounts of adsorbed water and carbon dioxide. Zeolite X is the most open-structured of the various synthetic zeolites commercially available. According to the crystallographic study of Broussard and Shoemaker3the aluminosilicate framework encloses a cavity network in which cavities of approximately 12 A. diameter are interconnected by 9-A. windows in a diamond configuration. Apart from the negligibly small contribution by the external crystal surface, the adsorptive surface is made up of the internal surfaces of these (3) L. Broussard and D. P4 Shoemaker, J . Am. Chem. (1960),
8%, 1041
