The Joimal of Physical Chemistry, Vol. 83,
Comrniunications to the Editor
cyanogen with hydrogen, the nonlinear concentration profiles for formation of hydrogen cyanide were again identical from a two apparatus determination. In the first, hydrogen cyanide profiles were measured in reflected shock wave experiments by infrared emission. These were compared to results from reflected shock waves where the sampling was through an end plate nozzle, and the detection was TOF mass spectrometry. The formation rates as a €unction of temperature were again identical within experimental error. Both of these examples demonstrate that time errors and composition errors are negligible from end wall sampling. A l ~ otemperature errors are negligible in the cited work, however it must be emphasized that these two examples pertain to reactions in which the overall enthalpy change is small. Lastly, we wish to emphasize that TOF mass spectrometry in reflected shock tube experiments offers one of the most general detectors developed to date. This is particularly important in complex chemical systems where the temporal behavior of not only reactanh but products and intermediates (including radicals in favorable cases) can be measured. Such information can lead to a clearer mechanistic understanding and is to be preferred over measurements of a single quantity in similarly complex systems. (1)G. P. Glass, G. B. Kistiakowsky, J. V. Michael, and H. Niki, J . Chem. Phys., 42, 608 (1965). (2) J. M. Brupbacher and R. D. Kern, J. Phys. Chem., 77,1329 (1973).
C. T. BOWMAN. In our paper we noted that time-resolved mass spectrometer measurements can provide qualitative information on intermediate and product species. We also indicated that there have been successful applications of time-resolved mass spectrometry to shock tube determinations of elementary rate coefficients. However, gas dynamic effects can introduce significant errors in mass spectrometric measurements behind reflected shock waves. Careful experimental design is required
No. 6, 1979 763
to minimize the magnitudes of these errors and additional diagnostic techniques should be applied to corroborate tht-= mass spectrometer results. WINGTSANG. The accuracy goal (f25%) which is held up as ultimately achievable in shock tube studies is now routinely met if not surpassed in comparative rate single pulse shock tube studies. The lesson I draw from this is the necessity of working under conditions where a particular reaction is truly being studied (no mechanistic artifacts), and the reaction temperature is measured directly. S. H. BAUER.It may be useful to mention that the ab’iguity of shock temperatures, which for two decades conventional ki-
neticists pointed to as the uniquely undesirable feature in the use of shock heating for quantitative investigations, may soon be resolved. The advent of high precision spectroscopic technqiues based on lasers permits the measurement of state populations on a microsecond time scale. Some improvement in precision is still required. However old style diagnostics, such as spectral line-reversal and measurement of rotational state populations via absorption or Raman spectra are enjoying a resurrection; laser-induced fluorescence is a commonplace diagnostic in many laboratories. C. T. BOWMAN. Measurements of post-shock temperature are useful, especially in interpreting results from single-pulse shock tube experiments and from conventional shock tube experiments on reactions with large enthalpy changes. Some of the new laser-based diagnostics discussed in the paper can provide temperature measurements, but are not yet routinely applied for this purpose. However, the accuracy of these measurements likely will be of the same order as calculated shock temperatures in experiments where the shock speed has been accurately measured and where boundary layer and shock attenuation have been minimized.
COMMUNICATIONS TO THE EDITOR Ionic B Coe!fficients of Aqueous Solutions Containing Tetraalkylarnmonium Ions
Sir: In a recent publication’ we discussed the differences in the structure-making effects of inorganic and tetraalkyl(ary1) onium ions on water structure. It was shown that i t is moire accurate to describe the physical nature of the influence of alkyl onium ions on the water structure b y the term “reinforcement of the water structure”, since “structurizing” in their solutions is determined by the formation of clathrate structures through the filling of the intermolecular cavities of water by organic chains of alkyl onium ions. Due t o this filling, the number of degrees of freedom of translational and rotational movements of the water molecules surrounding the organic portions of the ions decreases. This nature of ion-solvent interaction manifests itself in the B coefficients of the Jones-Dole equation too. Stokes a n d Mills2 have shown that viscosity (7) of aqueous electrolyte solutions can be described by the equation + Velectr + 7Einst + qorient+
(1) where qo is t h e viscosity of solvent; qelectris the positive viscosity increment conditioned by electrostatic interaction; qEinst is the positive viscosity hydrodynamic increment 7
70
Vstr
0022-3654/79/2083-0763$0 1.OO/O
arising from the shape and size of hydrated ions (this increment is closely related to the Einstein effect3); qorIent is the positive viscosity increment due to the orient iI t’l0n of polar solvent molecules b y the ions; Tstr is the viscosity increment related t o the destruction of solvent structure in t h e region of t h e ionic cosphere. This increment for inorganic ions is negative and for the large ones in aqueous solutions at comparatively low temperatures is the reason for “negative viscosity”. T h e comparison of eq 1 with the Jones-Dole equation4
shows, after eliminating the electrostatic contribution which is proportional t o c1I2, that
BcrO
=
7Einst
+ Vorient + 7str
(3)
However, in our view, in the case of aqueous solutions containing tetraalkylammonium cations (from Pr,N+) and other large tetraalkyl(ary1) onium ions, one more item (qrelnf) should be added to eq 1 7
=
70
+ Velectr + 7Einst + Vorient + 7str + Vreinf
and consequently
0 1979 American Chemical Society
(4)
764
The Journal of Physical Chemistry, Vol. 83, No. 6, 1979
Communications to the Editor
This contribution, vrelnf,which is positive and large, is conditioned by the effect of "reinforcement of the water structure"' by large tetraalkylammonium ions. Not dwelling in this communication on the methods of separation of molal coefficient B of aqueous solutions into ionic contributions, the ionic B coefficients can be assumed to be due to four major components: Bion = (6) BionEinst + Bionorient + Bionstr + Bionreinf
Bn-R4N+
p 0°C
, I mol"
1.50
T h e values of B'OnorIent and BIOnstrfor aqueous solutions containing large tetraalkylammonium ions are negligible, the former due to the small surface charge of these ions and the latter due t o the fact t h a t the hydrophobic hydration of these ions, as opposed t o inorganic ions, is realized without destruction of t h e intrinsic water ~ t r u c t u r e . ' , ~Then, for the above-mentioned solutions BRIX'
= BR4N'
Einst
+~~4~+reinf
(7)
According to the Einstein theory3
B = 2.59
I
(8)
where Y is the effective incompressible molal volume occupied in liquid by hard spherical obstacles, a t the surface of which the liquid remains stationary. Robinson and Stokes6 showed that eq 8 can also be used for electrolyte solutions. Assuming t h a t the V values for large hydrated R4N+ ions will be equal t o their volumes in solutions, one can write
I
I
100
50
I50
rh3, H
Figure 1. Variation of 6 R 4 N t coefficients in aqueous solutions with the hydrated radii of these ions at different temperatures.
= 2 . 5 ( 4 / 3 ) ~ r ~ ~=~6.307 + ~ NX 1021rR4N+3L mol-' (9)
BR4h+Elnst
where N is the Avogadro number and rR4N+ (in cm) is the radius of tetraalkylammonium ions. T h e question is, what radii must be used when calculating the BRfltEulst value. T o our mind, the hydration radii of these ions are more suitable since they characterize the real ion dimensions in aqueous When temperature increases, the hydrophobic hydration of large tetraalkylammonium ions decreases, because of the destruction of the ice-like water structure and the disappearance of intermolecular cages for placing the organic chains of these ions. As this takes place, the BR4Ntrelnf value must decrease and approach zero when hydrophobic hydration disappears. It is precisely this fact t h a t explains the negative values of aB/aT for aqueous solutions containing large tetraalkylammonium ionsn8 At the temperatures when hydrophobic hydration disappears, the value V in eq 8 will be equal to the volumes of unhydrated ions. Thus, if the speculations are correct, when the temperature increases the Bionvalues of large tetraalkylammonium ions in aqueous solutions must tend according to eq 7 and 8 t o the theoretical values of BR4N+E1nst for unhydrated ions. T h e BR4N+E1nst values calculated by e 9 using the radii of the unhydrated ions rpr4N+ = 3.33 and r g , , p = 3.85 A:,7 are equal t o 0.234 and 0.360 ~ m - ~ / m o respectively, l, and are presented in Figure 1 by dot-and-dash lines. Examination of Figure 1, constructed by using the BRlN data of Kay e t a1.8 and our valueslO of the radii of hydrated tetraalkylammonium ions (rh)calculated from conductometric data,ll corroborates our speculations. The extrapolation of the temperature dependence of the values BPrdN+and BBUdNf (dotted lines in Figure 1) up to the intersection with the straight line of the Einstein equation (dot-and-dash line in Figure 1)permitted us to estimate the dimensions of these ions when t h e hydrophobic hydration is absent: rPrrN+ = 3.50 8, and rBU4N+=
%I1,'
Figure 2. Temperature dependence of aqueous solutions.
and
values for
3.59 A. These values are close to the real dimensions of these ions found earlier7,l0from conductometric data (3.33 and 3.85 8,, respectively). There is a fairly good agreement between these two sets of values calculated by different methods, if one takes into consideration all assumptions made. BY,extrapolating the temperature dependencei of the BRINvalues to the temperatures at which the BRlh values will be equal to the BR4N+Einst values of bare (unhydrated) ions, we can estimate the temperature of the disappearance of hydrophobic hydration. This procedure (Figure 2) shows that, for Pr4N+ and Bu4N+ ions, the hydrophobic hydration completely disappears a t temperatures between 100 and 150 OC.16 This fact is corroborated by the IR data of Philip and Jolicoeur.12 Desnoyers and PerronI3 showed that the item De2 could be added to the Jones-Dole equation, where coefficient D depends on higher terms of solute-solvent and solutesolute interactions. The use of experimental results of Kay e t ale8permitted us to calculate by a least-squares method the B and D values for aqueous Pr4NBr and Pr4NI solutions. The B values of Pr4NBr solutions were 0.90 z!= 0.02
The Journal of Physical Chemistry, Vol. 83,
Communicatioris to the Editor
(10 "C), 0.765 f 0.009(25 "C), and 0.62 f 0.01 (45 "C), and for Pr4NI 0.824 f 0.008 (10 "C). Since the D values of an aqueous solution of inorganic ions are small,13Kaminsky's14 anion coefficients B were used for the separation of the molal B values into ionic contributions. Finally, the following values of BPra' were obtained: 0.95 (10 "C), 0.80 (25 " C ) , andl 0.64 (45 "C). These values are in good agreement with the data of Desnoyers and Perron13 a t 25 "C. Using the obtained values of B P r d N * instead of t h e experimental values8 for the construction of a figure such as Figure 1, 'we found, by the extrapolation of the temperature dependence of the values BPrdN+ up t o the intersection wi1,h the straight line of the Einstein equation, t h a t the dimension of bare Pr4N+ ions are 3.42 This value is only greater by 2.7 % than the real ionic dimension. Unfortunately, a similar calculation for Bu4N+ ions was impossible due t o the lack of experimental data for solutions of moderate concentrations.
No. 6, 1979 765
A Na-M o Na-Y A H-Y
-1I
a.
Acknowledgment. This research was supported by the Centre for Absorption in Science, the Ministry for Immigrant Absorption, State of Israel. References and Notes (1) E. S. Krumgalz, Faraday Discuss Chem. SOC.,No. 84, 247, 336 (1978). (2) R. H. Stokes and R. Mills in "The International Encyclopedia of Physical Chemistry and Chemical Physics", Vol. 3, "Viscosity of Electrolytes and Relateij Properties", Pergarnon Press, New York, 1965. (3) A. Einstein, Ann. Phys., 19, 289 (1906). (4) G. Jones and M. Dole, J . Am. Chem. Soc., 51, 2950 (1929). (5) E. S. Krumgalz, J . Gen Chem. USSR, 44, 1585 (1974). (6) R. A. Robinson and R. H. Stokes, "Electrolyte Solutions", Butterworths, London, 1966. (7) E. S. Krumaalz. Russ. J . Phvs. Chem.. 45. 1448 119711. i8j R. L.Kay, T-Vtuccio, C.Zawiyski, and D.F. Evans, J: Pbys. Cbem., 70, 2336 (1966). E. S. Krurngalz, J . Solution Chem., in press. B. S. Krumgalz, J . Struct. Chem., 13, 727 (1972). D. F. Evans and R. L. Kay, J . Phys. Chem., 70, 366, 2325 (1966). P. R. Phillip and C. Jolicoeur, J . Phys. Chem., 77, 3071 (1973). J. E. Desnoyers and G. Perron, J. Solution Chem., 1, 199 (1972). M. Kaminsky, Z. Phys. Chem. (FrankfurfamMain), 8, 173 (1956); Discuss. Faraday Soc., 24, 171 (1957). Based on the new experimental values of the ionic equivalent conductivities of the Pr,N+ ion in organic solvents published after 1971, we found that the value of the real dimension of this ion was equal to 3.33 i 0.10 A9 in contrast to the value 3.35 A which was reported earlier.' values of The dotted lines in Figure 2 correspond to the f3R'N+Elnst unhydrated Pr,N+ and Bu,N+ ions, 0.234 and 0.360 crn3/rnol, respectively, BS calculated above. Israel Oceanographic & Limnological Research Institute P.O.6. 8030, Haifa, Israel
Borls S. Krumgalz
Received July 6, 1978
Use of Hammcatt Indicators for Acidity Measurements in Zeolites
Sir: T h e classiic method of acid strength measurements in catalysts by noting the color changes of adsorbed dye indicators (Hammett indicators) during titration with an organic base (amines) in a nonaqueous solvent1S2 has become a well-established procedure for catalyst evaluation. It was shown t h a t this method can be applied t o acidity measurements in zeolites3 since parameters such as type of cation exchanged, degree of exchange, and cation distribution within the Si/Al framework all greatly influence the color changes of the indicators used.4 I t has been suggested" that acidity measurements depend more 0022-3654/79/2083-0765$0 1.OO/O
+7
+5
+3
+I
0-1
-3
-5
-7
-9
PK,
Figure 1. n-Butylamine titration of Na-Y (100% exchange), Na-M (100% exchange), and H-Y zeolites in dry benzene using Hammett indicators.
strongly on the size of the organic base than of thc indicator molecules, in t h a t the comparably smaller base molecules move inside the channels, whereas the large indicator molecules remain in accessible sites. The actual interaction of the indicator molecules, Le., color changes with acids or bases, should then only take place a t a fraction of the acidic centers. If this is really the case, experimentally determined butylamine titers should be orders of magnitude lower than those ~ b s e r v e d and , ~ the whole titration process should proceed quite fast, its speed depending upon the movement of the small base molecules. Reaching the end point with Hammett indicators during the titration with faujasites, however, often took some days, which indicates a slow diffusion process of the dye niolecules into the channels. T h e process was considerably accelerated by immersing the test tubes in an ultrasonic tank.4 Indicators used were from Eastman Kodak, as follows: neutral red (pK, = +6.8), 4-(pethoxyphenylazo)-mphenylenediamine (pK, = +5.0), 4-phenylazo-1-naphthylamine (pK, = +4.0), 4-o-tolylazo-o-toluidin (pK, = +2.0), 4-phenylazodiphenylamine (pK, = +1.5), 2-nitrodiphenylamine (pK, = -2.1), dicinnamalacetone (pK, = -3.01, chalcone (pK, = -5.6), and anthraquinone (pK, = -8.2). Figure 1 describes the formulas of some of the indicators used, the atomic distances of their planar projection being calculated with the help of ref 6. In the present work the results of n-butylamine titration in ion-exchanged faujasite Y (Na-Y, H-Y) are compared with those of H mordenite (Zeolon-H) and Na mordenite (100% exchanged). T h e adequately pretreated (calcined, H-M a t 500 "C, H-Y a t 290-300 "C, a t torr) zeolites (0.1 g) were placed in sealed test tubes and titrated with n-butylamine in dry benzene using a microburet for this purpose. Figure 2 shows the results of these measurements. Whereas the acidity of both Na-Y and Na-M is nearly zero, H--Y exhibited a very strong acidity with about 2.0 mequiv of n-butylaminelg of zeolite for the most acidic centers (equivalent to 90% H2S04, see ref 7). "Total" acidity a t pK, = +6.8 (neutral red) was found to amount to 2.8 mequiv of n-butylaminelg of zeolite. Therefore, more than two-thirds of all the acidic centers of H-Y are very strong. On the other hand, determination of the acidity of the H mordenite was impossible, since most of
t2 1979 Arnericari Chemical
Society
