NOTES
661
Table I : Wave Number (in 108 cm.-') of Observed Maxima of Charge-Transfer Compounds
---CTBl
Benzene Naphthalene Anthracene Phenanthrene Chrysene Pyrene Perylene a
C
==
29.8 21.0 16.7 22.0 19.2 17.2 14.5
cyclohexane.
-
TiBn--
Tic14
Solvent
CTBa
Au1-a
25.8 25.0b
4.8
..
8.3
..
27.0* 23.5 21.0
7.8 6.3 6.5
C"
.. ..
.. ,.
CTBI
27.0b 21.3 16.5 21.3 18.5 16.8 14.2
CTBa
Au1-1
601vent
C C
22.gb
19.2
6.4
c
5.0
C C C C
CTBi
23.8 16.0 12.2 16.2 14.5 12.5 8.0
vocll CTBI
CTBa
Au1-a
Aul-a
Solvent
C ..
C .. ..
18.8 15.2
20.4
63 7.2
C
12.4
C
Inaccurate value.
ization potential of the donor molecule, E A = electron affinity of the acceptor molecule, Ec = Coulomb energy of the ions D f and A-, and p = a constant. McConnell, et aL,4 give a linear relationship of the type h.v = aID b. In Figure 2 the hv values of maximum absorption of
+
3 2 ~
4 L 6,
P O
I
the longest wave length band observed in the complexes are plotted against the ionization potential of the aromatic hydrocarbons. The observed linear relationship strongly supports the charge-transfer character of the complexes. In some charge-transfer complexes a number of charge-transfer bands have been observed. It appears that the differences Au are independent of the acceptor molecules. So the reason for these additional bands will have to be found in the donor molecules. The Au values observed by us agree reasonably well wit.h a number of Au values found by Briegleb, et, al.,5 who measured the charge-transfer spectra of complexes of tetracyanoethylene and chloranil with aromatic hydrocarbons. In agreement with the interpretation of these authors we assign these additional absorption bands to transitions from the ground level of the complex to charge-transfer levels of which the D + ion is excited. Acknowledgments. The author wishes to thank Mr. J. P. C. van Heel and Mr. J. P.G. Rousseau for their experimental assistance and Dr. C. Bokhoven for his stimulating remarks. (4) H. McConnell, J. S. Ham, and J. R. Platt, J . Chem. Phys., 21,
u
66 (1953).
I
(5) G . Rriegleb, J. Caekalln, and G. Reuss, 2. physik. Chem. (Frankfurt), 30, 316 (1961).
6,
P
E
3 C 6,
>
4
t 6,
70
-
The Role of S i l v e r N i t r a t e Ion P a i r s in the Alkyl H a l i d e - S i l v e r N i t r a t e R e a c t i o n Ionization potential in
I
I
8.0
9.0
Figure 2. 0,complexes with TiC1,; 0 , complexes with TiBrd; A, complexes with VOC1,. 1, benzene; 2, naphthalene; 3,phenanthrene; 4, chrysene; 5, pyrene; 6, anthracene; 7, perylene.
Q.V.
by G. D. Parfitt, A. 1,. Smith, and A. G. Walton Chemistry Department, University o j Kottingham, S o t t i n g h a m , Enpland (Received July 2'7, 1964)
In connection with some light scattering studies' on the precipitation of silver iodide from homogeneous Volume 69,Slrmber 2
February 1966
I~OTES
662
solution, using the rcaction between silver nitrate arid ethyl iodide in ethanol, an expressioii for the over-all kinetics of this reaction was assunicld involving the degree of dissociation of the silver tiitrate. This expression is now discussed further i n the light of subsequent conductonietric studies of silver nitrate i n ethanol-water mixtures. The kinetics of the reactions between alkyl halides and silver nitrate i n a variety of solvents have been extensively studied since the papers of Donnan, Burke, and I’otts, published betweeii 1904 and 1910.2-4 A recent review of the literature is that of JIelendezand re^.^ Iiinetic data have usually been interpretqd in ternis of the stoichiometric coricentratioiis of the reactants, leading to fractional over-all apparent orders between 2 and 3. The contribution of the alkyl halide to the over-all order has invariably been found close to unity, but attenipls to separate the contributions of the Ag+ and NO3- iotis to the over-all order, again using stoichiometric concentratioris, have led to a coniplicated picture. Jlany of the apparent anomalies are, however, resolved if quantitative allowance is made for ionic strength effects, association of silver nitrate to ion pairs, and the effect of reaction products. The suggestion that silver nitrate ion pairs niay be involved in the reactioii was, in effect, made in 1909 by Do~itiati,~ who postulated that in ethanol “undissociated silver nitrate niolecules” niay be the reacting species. SenterV6however, found that the apparently similar reaction bet wee11 silver nitrate and salts of broiiiiriated aliphatic carboxylic acids i n aqueous solution was catalyzed by the solid silver bromide produced, arid the iori pair niechariislii was also strongly criticized by Hughes, Ingold, and JIasterman,’ who suggest ed that het erogeiieous catalysis was “diagnostic of the reaction betwxri silver ions and alkyl halides.” Although het erogcneous catalysis iiiay bc significant in aqueous atid soiiie other solutions, there is good evideiice that, e.g., in ethanol4 arid acetotiitrile,* it is negligible. The autocatalytic effects observed in the reaction can. nioreover, be explained without recourse to the postulate of hetcrogeneous catalysis. Hughes, et al.,’ have cointnented that to invoke the participation of “undissociated silver nitrate molecules” i n the reactioii is “a convenient theory since nobody claitlis to linow the concentration of these ti~olccules.” The sit uatioii has now changed, however, and sonic confiderice can be placed in iori pair coiicetitraf ions obtained froni coiiductance Ineasurenients analyzed by recetit ly refined conductance equatio~is.~ It is, moreover, oiilv the variation of ion-pair concentration with over-all concentration which is required Th? Joilrnal of Physical Chemistrg
to test the hypothesis and this quantity is not very sensitive to the value of the association constant taken. In the following discussion association constants for silver nitrate in ethanol-water niixtures which have been calculated froni conductance measurements in this laboratorylo will be used.
Kinetic Analysis Since the reaction rate is affected by both the silver and nitrate ion concentrations, the over-all kinetic process might tentatively be writtenR
NO3--
+
+ RI + Ag+ +KO3-RIAg
----t
products
transition state
with the *litrateiorl participatiIlg in the trarlsitioIlcolll11 plex essentially as forlllulated by Defining the rate of reactiorl as the rate of decrease of stoichiometric silver nitrate concelitration, arid ignoriIlgthe effect of reaction products -dc/dt
k’~~y‘f*’
(2)
where c is the stoichiometric concentration of silver nitrate and alkyl halide (taken equal) at any time t , y is the degree of dissociation of silver nitrate, and J’* is the mean ionic activity coefficient. Since eq. 1 can hardly represent a single kinetic step, it may be split into a postulated association equilibrium between two of the three species, and the ratedetermining step proper may be taken as the reaction between the result of this association and the remaining species. Since silver nitrate ion pairs are known to exist in solution (especially at low dielectric constant), it seems reasonable to interpret eq. 1 as (Agx03) RI -t transition state, where parentheses are used to indicate an ion pair. This leads at once to the kinetic expression
+
-dc/dt
=
h 2 ( 1 - y)
(3)
(1) M. J. Jaycock and G. D . Parfitt, Trans. Faraday Soc., 5 7 , 791 (1961). (2) K. A. Burke and F. G. Donnan, J . Chem. Soc., 8 5 , 555 (1904). (3) F. G. Donnan and K. A. Burke, 2. physik, Chem., 69, 148 (1909). (4) F. G. Donnan and H. E. Potts, J . Chem. Soc., 97, 1882 (1910). (5) E. Melendez-Andreu, Ann. Chim. (Paris), 7, 695 (1962). (6) G. Senter, J . Chem. Soc., 97, 346 (1910). (7) E. D. Hughes, C. K. Ingold, and S. Masterman, ibid., 1236 (1937). (8) G. S. Hammond, SI. I?. Hawthorne, J . H. Waters, and B. SI. Graybill, J . Am. Chem. Soc., 8 2 , 704 (1960). (9) R. >I. Fuoss, ibid., 81, 2659 (1959). (10) G. D. Parfitt and A. L. Smith, Trans. Faraday Soc., 5 9 , 257 (1963). (11) N . Kornblum, R. A. Smiley, 11, K. Blackwood. and D. C . Iffland, J . Am. Chem. Soc., 77, 6269 (1955).
NOTES
663
but, since K A = (1 - y)/~yy*~, where K A is the association constant for the silver nitrate equilibrium, eq. 3 reduces to -dc/dt
= kKACa7*fk2
which is equivalent to eq. 2. An alternative interpretation of eq. 1as RIAg'
+ NO3- -transition
state
will also lead t o expressions equivalent to eq. 2 and 3. I n hydroxylic solvents nitric acid appears as a product of the reaction yielding nitrate ions which will participate in the reaction. Then
KA
=
(1 - ?')/?'f,tZ(c?'
+
cry')
(4)
where the primed symbols refer to nitric acid. I n the simple case that the reaction is supposed to produce nitric acid in stoichiometric quantities (actually 70Oj, or more in ethanol-water mixtures3) and KA' is assumed equal to K A (in fact, KA' < K A in ethanol), then y = y' and c' = co - c, where the subscript 0 refers to initial quantities, yielding, from eq. 4, 1 - y = K A ~ o y y & ~Since, . under the assumed simple conditions, the ionic strength remains constant, the activity coefficient is also constant, leading to 1 - y = constant, which, combined with eq. 3, gives -dc/dt
= k"C2
reactants (at equal concentrations) according to k" = ( c o n ~ t a n t ) c ~ ~Thus, ~ ~ ~ .application of eq. 5 predicts that a plot of log (1 - y)ovs. logcowill havea slope of 0.53. Such a plot, using y-values calculated from conductance datalo over the range co < 20 mM, is shown in Figure 1. The result, is a curve of slope varying between 0.4 and 0.6, which is in good agreement with the experiments of Burke and Donnan. If the value of K A from conductance data (210 M-l) is in error, it is most probable that association has been overestimated.lZ Reductmionof K Ato 8 M-l, as foundio in the 70 wt. % ethanol-water mixture, only causes an increase in d log (1 - y)o/d log co to -0.7.
(5)
where k" = k ( l - y ) o so that a given reaction would follow a simple second-order equation in stoichiometric concentrations. A plot of c-' us. t is now linear and of slope k" whereas eq. 3 predicts that such a plot is concave to the time axis of slope, a t any time, of k(1 y). Even though the simple cqnditions assumed above do not fully apply, some straightening of the c-1 us. t plot will enable initial slopes (independent of products) and, hence, k values to be obtained more accurately. The addition of the nitrate of an indifferent cation will, on the assumed mechanism, increase the reaction rate by increasing the concentration of silver nitrate ion pairs. Equation 4 again applies, and, for nitrates such as KN031° and "1N03, the K A value for the added nitrate may be taken equal to that for silver nitrate without great error.
Comparison of Kinetic Analysis with Experimental Data Burke and Donnan2 found that, for the reaction between silver nitrate and both ethyl and butyl iodides in ethanolic solution, the (apparent) second-order rate constant, k", calculated from -clc/dt = k"c2, remained almost constant in the course of any one reaction but varied with the initial concentration of
0.3
I
I 0.2
0.4
0.6
Log co
(CO.
0.8 mM).
1.0
1.2
Figure 1. Plot of log (1 - y),, vs. concentration co (mM) for silver nitrate in ethanol; y is t h e degree of dissociation of silver nitrate.
Burke and Donnan also studied the effect of adding ammonium nitrate to the initial reaction mixture. I n the range of concentrations for which values are available, they found that the value of k" for a reaction initially 12.5 m M , with respect to ethyl iodide, silver nitrate, and ammonium nitrate, was greater than that found in the absence of ammonium nitrate by a factor of 1.2. If the ammonium nitrate is taken to have approximately the same value of KA as silver nitrate, this factor is expected, on the ion pair mechanism, to be the ratio of the (1 - y ) value at 25 mM t o that a t 12.5 mM, which, from the data quotedlO and plotted in Figure 1, is 0.46/0.38 = 1.2. Calcium nitrate caused smaller increases in rate, as expected in view of increased ionic strength and association effects introduced by the divalent cation. For reactions in solvents containing significant amounts of water where the nitric acid is appreciably ionized, the virtual constancy of k" for a given reac~~~
(12) R. L. Kay and (1963).
~
J. L. Dye, Proc. Natl. Acad. S c i . C' S., 49, 5
Volume 69,Number 8 February 1965
NOTES
664
tion is explained immediately by eq. 5: I n pure ethanol the limited dissociation of nitric acid, together with the side reaction producing ethyl nitrate rather than nitric acid to the extent of -30%, makes the virtual constancy of k“ more difficult to explain, though the ionization of the nitric acid is still sufficient to make the reaction difficult to follow conductometrically.13 I n the case of nonhydroxylic solvents, such as acetonitrile, where no nitric acid is produced, the value of IC!’ does, as expected, decrease during the course of reaction.8 Furthermore, when silver lactate replaces silver nitrat’e in ethanol-water mixtures, the value of k“ again decreases as the reaction proceeds14which is exactly as expected in view of the weak nature of lactic acid. The ion-pair mechanism predicts that the reaction rate will decrease as the dielectric constant is raised in moving from ethanol to water as solvent, by the ratio of the (1 - y ) values in the two solvents which is 4 3 0 ( y in water calculated from an extrapolated KA valuelo and stoichiometric activity coefficient^'^). As observed, the reaction with methyl or ethyl iodide3 is only decreased by a factor of -7; the rate does not change monotonously, but it is riot expected that the silver nitrate ion pair will remain equally reactive in conditions of competitive solvation obtaining in niixtures of polar solvents. The recent work of ;\Ielendez-Andreu5 using an acetone-water mixture provides good support for the proposed mechanism involving the silver nitrate ion pair. The reactions of alkyl bromides, from methyl to nhexyl, with silver nitrate and perchlorate were studied and found (by means of the Leffler ploti5) to be of the same mechanism for all the bromides in all the solvent mixtures. The reaction with silver nitrate is particularly rapid in acetone, and it is significant that silver nitrate is a particularly weak electrolyte in this solvent. Silver perchlorate, which is not as weak in acetone,16 reacts niuch more slowly, even though in water-rich mixtures, where the two electrolytes are both strong, the rates are very similar. It is evident that the rapid rate of the silver nitrate reaction in acetone cannot be due to a high reactivity of the silver ion in this solvent and also can hardly be due to the increased nucleophilic character of a desolvated nitrate ion since the transport number for this ion in silver nitrate in acetone (0.58) is little different from that in water (0.54) and less than that in ethanol (0.61). The considerable association of silver nitrate in acetone, which may be due to stabilization of the ion pair itself in this solvent , I 7 explains the high reaction rate, at once, if the 1011 pair is the reacting species. We conclude, therefore, that thc suggestion of The Journal of Physical Chemistry
Donnan as to the mechanism of the silver nitrate-alkyl halide reaction was essent,ially correct for many solvents despite later criticisms. (13) A. G. Walton, Thesis, University of Nottingham, 1960. (14) R. Parsons, “Handbook of Electrochemical Constants,” Butterworth and Co., Ltd., London, 1959. (15) J. E.Leffler, J . Org. C h a . , 20, 1202 (1955). (16) V. S. Griffiths, K. S. Lawrence, and AM.L. Pearce, J . Chem. Soc., 3998 (1958). (17) W. R.Gilkerson, J . Chem. Phy8., 2 5 , 1199 (1956).
The Reaction of Methyl Radicals with Methyl and Methylene Fluoride by G. 0. Pritchard, J. T . Bryant, and R. L. Thommarson Department of Chemistry, University of California, Santa Barbara, California 93018 (Received July 29, 1964)
Raal and Steacie‘ have observed a significant decrease in the activation energy for H atom abstraction by methyl radicals with increasing halogenation of methane
CHs
+ CH, - .X,
+CH4
+ CHa - ,Xn
(1)
For X = chlorine, and n = 1, 2, and 3, they obtained values of El of 9.4, 7.2, and 5.8 kcal. mole-’, respectively. For X = fluorine, and n = 1 and 2, El was found to decrease from 8.7 to 6.2 kcal. mole-’. A similar trend was obtained with methyl and methylene bromide. They conclude that there is a progressive decline in the magnitude of the C-H bond strength with increasing halogenation, leading to an enhanced “activity” of hydrogen atoms in substituted methanes. While this generalization is correct for chloromethanes,2 it is certainly not correct for fluoromethanes.2 We3 have found El using CF,H to be 10.2 f- 0.2 kcal. mole-’, and the most likely value314for D(CF3-H) is close to 105 kcal. mole-’. We have therefore redetermined El for CFHl and CF2H2. (1) F.A. Raal and E. W. R. Steacie, J . Chem. Phys., 20, 578 (1952). (2) C . T. Mortimer, “Reaction Heats and Bond Strengths,” I’ergamon Press, London, 1962,pp. 132-134. (3) G. 0.Pritchard and R. L. Thommarson, J . Phys. Chem., 68, 568 (1964). (4) E.Whittle, private communication: from unpublished bromina tion experiments. This is some 4 kcal. lower than the previous value, see 1’. Corbett, A. h f . Tarr. and E. Whittle, Trans. Faraday Soc., 59, 1609 (1963).
