FRANK J. MILLEROAND W. DROST-HANBEN
1758
Apparent Molal Volumes of Ammonium Chloride and Some Symmetrical Tetraalkylammonium Chlorides at Various Temperatures by Frank J. Millero and W. Drost-Hansen Contribution N o . 909 from the Institute of Marine Sciences, University of Miami, Miami, Florida $3149 (Received November 16, 1967)
The apparent molal volumes, ~ v ’ s ,of (0.05 m) NH4C1, Me4NC1, Et4NC1, n-PrrNC1, and n-BudNC1 have been determined at one-degree intervals from 20 to 40” from precision density measurements. The apparent molal expansibilities, +E’S, have been calculated from the 4~ values at various temperatures and have been equated to the infinitedilution molal expansibility, B. The .??values of the RrNCl’s are not a linear function of molecular weight of the cation, RaN+. The @ , of T\TH4C1is low, apparently due to the ability of the NHr+ ion to form an “icelike” species in solution. The Eo’s of Pr4NC1and Bu4NC1appear to be high compared to the iower molecular weight R4NCl’s. This nonlinear behavior is interpreted as being due to expansibility changes in the structure of water caused by the R4N+ cations, EO(struct). This structural effect decreases as the temperature increases and increases as the size of R4N+ increases. The similarity of Bo properties of R4Nf cations and the aliphatic alcohols is discussed, and it is postulated that the abnormal properties of R4N+ salts may be normal for solutes able to cause hydrophobic bonding.
Introduction
Results
This study is a continuation of a program’ of measuring the apparent molal volumes of aqueous salt solutions at various temperatures. Interest in the symmetrical tetraalkylammonium salts stems from the abnormal behavior of (Pv, the apparent molal volume found by other workers.2-11 Also, since these large symmetrical cations are not hydrated in the normal sense (Le., electrostrictive hydration), we hoped to determine the absolute expansibility of the C1- ion by a method similar to that used by Conway, et u1.,536in determining the absolute Poof C1- ion. Precision density measurements were made on dilute solutions (0.05 m) of NH4C1, Xe4NC1, EtdNCI, n-P4NC1, and n-BurNC1 a t one-degree intervals from 20 to 40”. The apparent molal volume, (PV, and apparent molal expansibility, (PE, have been calculated from the density data.
The densities of the various solutions have been measured at l ” intervals from 20 to 40”. The densities of the solutions were fit to equations of the form
Experimental Section NH4C1 used was reagent grade Baker Analyzed. All the tetraalkyIammonium chlorides were obtained from Eastman Organic Chemicals. All the salts were purified by several recrystallizations according to the methods listed by Conway, et aL5 Each salt was dried in YUCUO for at least 1 week before use. The solutions were made by weight with degassed, doubly distilled water. The magnetic float densitometer used to make the density measurements has been described elsewhere.I2 The constant-temperature bath in which the densitometer is submersed was controlled to ~ 0 . 0 0 1 ”with a Hallikainen regulator. The Journal of Physical Chemistry
=
A
+ Bt + Ct2 + Dt’
(1) The constants for this equation are listed in Table I along with the root-mean-square (rms) deviations. The thermal expa,nsion coefficient, asoln =-(bdaoln/bt)1/ d,,ln, of the solutions can be calculated from dsoln
asoin=
-(B
+ 2Ct + 3Dtz)/ds0ln
(2)
The apparent molal volumes listed in Table I1 were calcuIated from the equation
+-
100OAdaoln dHzodsoinm
(Pv = -__-
fi/1
(3)
dsoln
(1) F. J. Millero and W. Drost-Hansen, presented at the 154th National Meeting of the American Chemical Society, Chicago, Ill., Sept, 1967. (2) W. Y. Wen and S. Saito, J . Phys. Chem., 68, 2639 (1964). (3) W. Y. Wen and S . Saito, ibid., 69, 3569 (1965). (4) W. Y. Wen and K. Nara, ibid., 71, 3907 (1967). (5) B. E. Conway, R. E. Verrall, and J. E. Desnoyers, Trans. Faraday SOC.,62, 2738 (1966). (6) B. E. Conway, R. E. Verrall, and J. E. Desnoyers, Z. Phys. Chem. (Leipzig), 230, 157 (1966). (7) B. E. Conway and R. E. Verrall, J . Phys. Chem., 70, 3952
(1966); 70, 3961 (1966). (8) J. E. Desnoyers and M.Arel, Can. J . Chem., 45, 359 (1967). (9) F. Franks and H. T. Smith, Trans. Faraday Soc., 63, 2586 (1967). (10) H. E. Wirth, J . Phys. Chem., 71, 2922 (1967). (11) L. G . Hepler, J. M. Stokes, and R. H. Stokes, Trans. Faraday SOC.,61, 20 (1965). (12) F. J. Millero, Rev. Sci. Znstrum., 38, 1441 (1967).
1759
APPARENT MOLALVOLUMES OF AMMONIUM CHLORIDE
from the apparent molal volumes: they are listed in Table 11. The average @E a t 25 and 30" and the average deviation from the mean are compiled in Table IV; also listed are the results of other worker^^*^ for comparison. Since OE is not strongly dependent upon concentrationg and the measurements were made on dilute solutions, we have equated OE with Eo,the partial molal expansibility at infinite dilution. The difference between $E and the true value of Eo is well
Table I: Constants for Density Equation ~ O B B , - IOBC, ~ V D , Rms dev, g/ml deg g/ml degz g/ml degz ppm
A, g/ml
Salt
"4'21 -Me4NC1 EtrNCl
1.001355 1.000385 1.000145 0'999879 0.999312
n-Pr4NC1 n-BurNC1
26.171 17.652 28*121 25'061 44.662
6.5780 6.3473 6.6932 6'6391 7.4151
2.4349 2.1656 2.5212 2'4989 3.3222
0.4 1.2
l.o 1'5 0.9
~
Table I1 : Apparent Molal Volume and Apparent Molal Expansibility of NH&l and Some Symmetrical Tetraalkylammonium Chlorides a t Various Temperaturesa Temp, "C
20 21 22 23 24 25 26 27 28 29 30 31 32 33 34 35 36 37 38 39 40 a +V
,--NH&l
(0.06970 m)-
-MerNCl
QV
QE
QV
36.1.56 36.1.74 36.1.93 36.212 36,231 36.236 36,269 36.275 . 36.309 36.314 36.334 36.335 36.375 36.396 36.431 36.452 36.474 36.495 36.517 36.538 36.560
(0.05268 m)-
107.15 107.21 107.27 107.39 107.51 107.58 107.63 107.73 107.78 107.89 107.94 108.11 108.22 108.23 108.36 108.46 108.55 108.69 108.77 108.88 109.02
0.018 0.019 0.019 0,019 0.005 0.033 0.006 0.034 0.005 0 020 0.021 0.020 0.021 0.035 0.021 0.022 0.021 0.022 0.021 0.022 I
is given in milliliters per mole, and
+E
yEtrNC1 (0.05369 m)-
Salt
Molality
value
NH&l MedNCl Et4NCl n-Pr4NC1 n-BurNCl
0.06970 0.05268 0.05369 0.05042 0.05341
36.24 107.58 166.50 229.13 293.31
QV
QE
+V
QE
0.06 0.06 0.12 0.12 0.07 0.05 0.10 0.05 0.11 0.05 0.17 0.11 0.01 0.13 0.10 0.09 0.14 0.08 0.11 0.16
166.04 166.11 166.23 166.34 166.42 166.50 166.62 166.70 166.82 166.91 167.04 167.16 167.35 167.42 167.59 167.71 167.81 167.98 168.12 168.26 168.40
0.07 0.12 0.11 0.08 0.08 0.12 0.08 0.12 0.09 0.13 0.12 0.19 0.07 0.17 0.12 0.10 0.17 0.14 0.14 0.14
228.45 228.58 228.72 228.81 228.97 229.13 229.23 229,36 229.57 229 69 229.87 230.00 230.12 230.29 230.50 230.66 230.80 230.97 231.10 231.31 231.52
0.13 0.14 0.09 0.16 0.16 0.10 0.13 0.21 0.12 0.18 0.13 0.12 0.17 0.21 0.16 0.14 0.17 0.13 0.21 0.21
292.31 292.48 292.67 292.89 293.12 293.31 293.51 293.74 293.98 294.27 294.53 294.80 295.07 295.34 295.60 295.86 296.15 296.43 296.72 296.97 297.19
0.17 0.19 0.22 0.23 0.19 0.20 0.23 0.24 0.29 0.26 0.27 0.27 0.27 0.26 0.26 0.29 0.28 0.29 0.25 0.22
I
is given in milliliters per mole per degree.
within the experimental error (k0.03 ml/mol deg); thus equating @E and Eois justifiable.
ml/mol-Lit. value
36.36 107.69 166.57 232.06 292.83
Discussion of Results Ref
14 5 5 5 5
The densities of water used in these calculations were taken from Tilton and Taylor's work.13 The results at 25' can be compared to other workers5,l4by using the equation +v = Ovo
rn-BurNC1 (0.05341 m)-
QE
and Some Symmetrical Tetraalkylammonium Chlorides a t 25"
Our
(0.05042 m)--
QV
Table I11 : Apparent Molal Volume of c-&,
-n-PrrNCl
QE
+ SV;' + bc
(4)
This comparison is shown in Table 111. The apparent molal expansibilities +E = d+/dW can be calculated
Recently, various ~ o r k e r s ~ J ~have - ' ~ attempted to divide the partial molal volume of ions at infinite dilution into two contributions Po(ion) = P ( i n t )
+ Po(e1ect)
(5)
where Po(int) is the intrinsic volume of the ion plus the volume due to void space and Po(elect) is the decrease in volume due to electrostriction. For ions that have strong structural effects on water (e.g., Li+, F-, Be2+, (13) L. W. Tilton and J. K. Taylor, J. Res. Nut. Bur. Stand., 18, 205 (1937). (14) B. B. Owen and S. R. Brinkley, Chem. Rev., 29, 461 (1941). (16) L. Hepler, J. Phys. Chem., 61, 1426 (1957). (16) P. Mukerjee, ibid., 65, 740 (1961). (17) E. Glueckauf, Trans. Faraday SOC.,61, 914 (1965).
Volume 72,Number 6 May 1968
FRANK J. R'IILLERO AND W. DROST-HANSEN
1760
Table IV : Apparent Molal Expansibility of NH&l and Some Symmetrical Tetraalkylammonium Chlorides ,
+E, ml/mol
Salt
deg
25'
350
n-PrrNC1
0.018 f:0.007 0.079 f 0.028 0.100 i0.020 0.141 f 0.027
n-BurNC1
0.223 f:0.028
0.023~0.003 0.108 =k 0.030 0.136i0.027 0.166 f 0.029 0.266 f:0.015
NH&1 MedNC1 EtdNCl
These values are for tetraalkylammonium bromides; however, since .D(Cl-)
and Mg2+)17JS another term, added to eq 5. PO(ion)= Vo(int)
VO
(struct), must be
+ Po(elect) + Po(struct)
(6)
Since the term Po(int) is large for most ions compared to the other terms of eq 6, it is difficult to separate the individual terms of Po(ion). If one differentiates eq 6 with respect to temperature, the partial molal expansibility of an ion at infinite dilution, Eo(ion), is given by the equation Eo(ion) = Eo(int)
+ Eo(e1ect) + EO(struct)
(7)
where EO(int) is the expansibility due to void space (since the intrinsic size of the ion can be assumed to be independent of temperature), Eo(elect) is the expansibility due to electrostriction (solute-solvent interactions), and Eo(struct) is the expansibility due to structural changes in water. Since the large R4N+ ions have little or no electrostriction ( Po(elect) 0 and Eo(elect) 0), it was of interest to determine whether Eo(ion) for R4N+ could be related to some function of these ions, for instance po(ion) and n general relationship for Eo(ion) for other ions could be inferred. It has been shown by others5s6J9that Voof R4N+ halides is a linear function of the molecular weight of the cation5f6and V0(R4N+)is a linear function of the radius cubed.l9 From the linear function of V0(R4N+) halides, other^^,^ have been successful in determining the absolute Vo of C1-, Br-, and I- (by plotting 7 0 (R4N+) halides vs. the molecular weight of the cation and extrapolating to zero) or Vo(H+),that agree within experimental error with those calculated by other methods.20 From the results of Eo of the RdSCl's at 25" (Figure l), it can be clearly seen that a linear relation is not obtained when one plots EO(RdNC1) vs. the molecular weight of the cation, R&+. 4 smooth curve is obtained with the Eoof Pr4NC1 and Bu4NC1 appearing to be high compared to the lower molecular weight, R4NCl's. At 30", however, a linear relation is obtained for Me4;1-C1, EtdNCI, and Pr4NC1 with Eo of Bu4NCI appearing to be high. The high values of Eo of Pr4NC1 and Bu&Cl at 25" indicates that the structural term, Eo(struct),is causing a larger value for Eothan expected
-
The Journal of Physical Chemistry
-
Other workers'"
Ref
...
-
0.11 (15') 0.13 (15') 0.14 (15') 0.25 (15O), 0.24 (25')
9 9
9 9,2
B(Br-) the comparison is justified.
by comparison with the lower molecular weight R4NCl's. The Eoresults at 30" indicate that Eo(struct) decreases with temperature with only Bu4NC1 appearing to be high. D(struct) disappears as the temperature is increased and increases as the size of the R4N+ increases. Although Pr&+ and BukN+ chlorides clearly show the structural term, Eo(struct), this does not mean that the effect is not important for the lower molecular weight RdN+ salts (especially a t lower temperatures). The low Eo of NH.421 is apparently due to the ability of NH4+ to form an "icelike" structure which has a negative expansibility similar to water below 4" (or ice becoming liquid at 0'). It should be noted that the adiabatic molal compressibility, Ro = -(bPO/bP)v, behavior of the R4X+ bromides and chlorides' is similar to the Eo behavior. The KO of the R4N+ salts are generally higher (less negative) than R"J of salts, and R'I decreases ( V o increases) as V0(R4K+) becomes larger. Although a linear increase in KO was not found for all the R4S+ ions7 (as for Eo for our measurements), it does appear that Ro of Pr4N+ and Bu&+ is small (Pois larger) compared to RIedK+ and Et4N+. The high value for Eo and i f 0 for Pr&+ and Bu&+ compared to the other R4X+ cations suggests that Vo(struct) for the RdN+ cations is negative. Since the structural effects disappear as the temperature or pressure is increased, this negative contribution is reduced causing Eo and Roof Pr&+ and Bu4X+to appear high. Further evidence for Bo(struct) having a negative value for the RIN+ cations can be shown by comparing the 7 0 of R4N+cations with the aliphatic alcohols and other ions. First, if one calculates V0(R4X+) from radii obtained from molecular models, one obtains a value that is smaller than the measured value.g This difference increases in magnitude with increasing ionic size. For simple ions, where this effect is due to electrostriction, the reverse is trueaJ6J7 ( L e . ) Vo(elect) decreases with increasing ionic radius). This behavior is similar, however, to that observed for the aliphatic (18) F. J. Millero, J. Phys. Chem., 71, 4567 (1967). (19) R. Zana and E. Yeager, ihid., 71, 4241 (1967). (20) For comparison of VO(H+) obtained by various methods, Bee P. Mukerjee, ihid., 70,270 (1966).
1761
APPARENTMOLALVOLUMESOF AMMONIUM CHLORIDE
E"fsertl
0
50
100
150
250
200
MM! of Cation Figure 1. Molal expansibility, EO, of "$1 (solid circles) and 25" (open circles).
and some RaNCl salts us. molecular weights of cations at 35"
alcohols.21 Second, when one plots P(RdN+) cations us. r3 (the crystal radius), a slope of 2.3 is ~ b t a i n e d . ' ~ This slope is lower than the theoretical value of 2.52 ( 4 ~ N / 3 )and also the value obtained for other ions (4.75). l 9 Thus the Boof the R,N+ is lower than expected from comparison to other ions. Since Bo(elect) is very small or zero fior the R4N+ cations (with the exception of Jle4N+), the low value of Po can be attributed to Vo(struct) being negative. The VO(Me4N+) is lower when compared to the higher molecular weight R4N+cations, indicating that this cation has a negative contribution due to electrostriction, Vo(e1ect).19 This is in agreement, with Kay and Evans' conductance and viscosity B-coefficient studies.22 The similarity of the Bobehavior of the R4N+cations with the behavior observed for the aliphatic alcohols (ROH) can also be demonstrated as follows. 1. The Eo of the RIN+salts increase with increasing temperature (Table IV) in a manner similar to the behavior of aliphatic alcohols (MeOH, EtOH, PrOH, BuOH, PeOH, i-BuOH, and se~-BuOH).~3+The EO'S of electrolytes (salts), however, decrease with increasing temperature.26,26 2. Comparison of the Eo of RIN+ cations and aliphatic alcohols23(shown in Figure 2 at 30") indicates that a common line is obtained when plotted us. Po, with the exception of BUN+. Since Eo(elect) is very small or zero for the R4N+ cations and the alcohols, the Eo of these solutes can be attributed to Eo(int) and EO(struct). It appears that Eo(int) is proportional to VO and as the volume increases Eo(struct) becomes impor-
tant. The apparent deviation of the BueN+ cation compared to the other solutes may be real or due to a gradual increase of Eo(struct). From the comparison between electrolyte and nonelectrolyte, Pobehavior as a function of temperature, one might postulate that the r0(R4N+) salts increase with increasing temperature like nonelectrolytes and do not go through a maximum near 40-60" like other electroIyte~.27-~~ From the recent m e a ~ u r e m e n t sit, ~appears ~ that up to 65' the Bo of the R4NBr salts do not go through a maximum. As mentioned in the Introduction, we hoped to obtain the absolute expansibility of the C1- ion, Eo(Cl-), by plotting E0(R4NC1)us. the molecular weight of the cation and extrapolating to zero. Although the linearity does not exist for all the RhNCl salts, it is possible to estimate Eo(Cl-) = 0.048 ml/mol deg a t (21) F. Franks and D. J. G. Ives, Quart. Rev. (London), 21, 1 (1966). (22) R. L. Kay and D. F. Evans, J . Phys. Chem., 70, 2325 (1966). (23) M. E. Friedman and H. A. Scheraga, J . Phys. Chem., 69, 3795 (1965). (24) D. M. Alexander, J . Chem. Eng. Data, 24, 252 (1959). (26) R. E. Gibson and 0. H. Loeffler, J . Amer. Chem. SOC.,63, 443 (1941). (26) H. S. Harned and B. B. Owen, "The Physical Chemistry of Electrolytic Solutions," 3rd ed, ACS Monograph No. 137, Reinhold Publishing Corp., New York, N. Y., 1958. (27) A. J. Ellis, Chem. Commun., 21, 802 (1966). (28) A. J. Ellis, J . Chem. Soc., 1579 (1966). (29) A. J. Ellis, ibid., 660 (1967). (30) S. Schiavo, B. Scrosati, and A. Tommasini, Ric. Sci., 37, 211 (1967). Volume 72, Number 6 Ma21 1968
FRANK J. MILLEROAND W. DROST-HANSEN
1762 I
I
I
1
1 nBu,N'
.2
E" .1
0 0
50
100
150
200
250
?' Figure 2. Molal expansibility, Eo,of RdN+ cations and ROH alcohols (M. E. Friedman and H. A. Scheraga, Chem., 69, 3795 (1965)) us. Vo or RdN+ cations and ROH alcohols at 30'.
J. Phys.
25" using the results for R/Ie4NC1and EtXC1. Using m a t i ~ n ,salting-in ~~ and induced cationthe Eo(HC1) = 0.034 ml/mol deg20 at 25", EO(H+) = cation i n t e r a c t i o n ~ . ~ - ~Franks ,~ and Smith9 have -0.014 ml/mol deg. From the Bo(R4NBr)gat 5 and recently discussed these various explanations and sug25", it is possible to calculate vo(Br-) = 30.32 and gested that induced cation-cation interaction is the 31.02 ml/mol, respectively, by the method of Conway, preferred explanation. This cation-cation interaction et aZ.6~~From these two values of Bo(Br-), the absolute presumably is due to the overlapping of extensive Eo(Br-) = 0.035 ml/mol deg can be calculated. Using hydration structures and not directly caused by coulomEO(KBr) = 0.060,' EO(KC1) = 0.089,' and EO(HC1) = bic interactions. Since the measurements described in 0.034 ml/mol deg,2e B0(H+) is found to be equal to this paper were not made as a function of concentration, 0.010, Thus, Bo(H+) is between -0.010 and -0.014 we cannot use our results to discuss the solute-solute interactions of the RIN+ salts. From the similarity ml/mol deg; ie., Bo(H+)is small and negative. From of the Poproperties of RIN+ cations with the aliphatic the recent measurements of Bo of R4NBr salts over a alcohols at infinite dilution, it appears that the apparent wide range of temperatures (15-66°),30 we obtain a value for Eo(H+) within these limits at 25", although abnormal properties of the R4N+ cations may, in fact, be normal for solutes that can cause hydrophobic the experiments were not made in very dilute solutions b ~ n d i n g . ~Presently there are no theoretical means nor with the precision of Franks and Smith's work.g of estimating the limiting slope (qh vs. c) for nonelecTo adjust the ,!?)(ions)(listed in our previous paper') to trolytes and, furthermore, very few measurements have the average Eo(H+),0.012 should be added to the anions been made of nonelectrolytes in dilute solutions with and subtracted from the cations. the accuracy of the present study. Using the absolute VO(Br-) = 31.02, BO(KC1) = I n our future work we plan to measure the $v of 26.89,31 Vo(KBr) = 33.738 and P(HC1) = 17.82 ml/ electrolytes and nonelectrolytes as a function of conone obtains Vo(H+) = -6.19 ml/mol at 25". This value is lower than that obtained by others;20 centration and temperature (over a large range, 0-100") with hopes of obtaining a better understanding of however, this may be due to the negative Po(struct) of solute-solvent , solute-solute, and solvent-solvent inthe R4N+ ions, which increases in magnitude as the size teractions in aqueous solution. of the RIN+ increases (the negative contribution of Bo(struct)for the R4K+salts would cause the value of Bo(halide) to be larger than the true absolute value (31) F. Vaslow, J . P h y s . Chem., 70, 2286 (1966). (32) L. A. Dunn, Trans. Faraday SOC., 62, 2348 (1966). and this would likewise give a value for BO(H+) that is (33) R. M. Diamond, J . Phys. Chem., 67, 2513 (1963). lower than that obtained by others).*O (34) B. J. Levien. A u s t . J. Chem., 18, 1161 (1965). Previous studies seem to indicate that the R4N+ (35) R. L. Kay and D. F. Evans, J . P h y s . Chem., 70, 366 (1966). salts strongly influence the structure of (36) F. Franks and H. T. Smith, ibid., 68, 3581 (1964). The negative deviations of 7 0 from the limiting law (37) H. S. Frank and W. Y . Wen, Discussions Faraday Soc., 24, 133 and the other thermodynamic anomalies of the RIN+ (1957). salts have been attributed to ion pairing,10~11J3-36 (38) 5. Lindenbaum and G. E. Boyd, J . P h y s . Chem., 6 8 , 911 (1964). hydrophobic bonding,a6"iceberg effect,"37 micelle for(39) B. E. Conway and R. E. Verrall, ibid., 70, 1473 (1966). T h e Journal of Physical Chemistry
ELECTRICAL CONDUCTIVITY OF TETRAALKYLAMMONIUM HALIDESOLUTIONS Acknowledgment. The authors wish to thank Dr. R. L. Kay for his helpful comments and suggestions and Dr. F. Franks who made his manuscript available to us
1763
prior to publications. The authors wish to acknowledge the support of the Office of Saline Water for this study.
The Electrical Conductivity of Aqueous Tetraalkylammonium Halide Solutions under Hydrostatic Pressure by R. A. Horne and R. P. Young Arthur D . Little, Inc., Cambridge, Massachusetts OOdl4
(Received November 80, 1967)
The electrical conductivities of 0.10 aqueous solutions of the salts R4N+il-, where R is CHI, C2Hs,CaH7, n-C4H9,and n-CaHn and A- is C1-, Br-, and I-, have been measured at 4 and 25” over the pressure range 1 atm to 4000 kg/cm2. The results of these measurements support the viewpoints that (1) the hydrophobic hydration atmospheres of these ions are fundamentally different from the coulombic hydration envelopes surrounding “normal” ions such as the alkali metal cations, in particular being much more stable with respect to hydrostatic pressure, and (2) cation-cation interactions are important in solutions of the larger ions in this series.
Introduction I n order to account for cert,ain properties of liquid water and aqueous solutions, Wickel has proposed that, in addition to the “free” or monomeric water and the flickering, H-bonded Frank-Wen clusters, there exists a “third state” which he postulates consists of dimers, trimers, tetramers, etc., even though such small aggregates were specifically discounted in Nemethy and Scheraga’s2 quantitative development of the FrankWen3 theory. Similarly, in order to account for certain interface phenomena, we have been obliged to postulate a “third state”’ which we have called simply the p structure to distinguish it from the a structure of the Frank-Wen clusters and the hydration atmospheres of normal ion^.^'^ This p structure, which is largely present at interfaces, is less susceptible to destruction by the application of hydrostatic pressure than is the a structure of the Frank-Wen clusters.6 The solution properties of the tetraalkylammonium ions have long been recognized as being different from those of “normal” cations such as the alkali metal ions, and Wickel also discussed this “hydration of the second kind.” These ions are heavily hydrated and powerfully enhance the structure of the water surrounding them as reflected in their very large viscosity B coefficients, ranging from +0.10 for (CH3)4N+ to almost $0.90 for (C3H7)4N+ as compared to about +0.14 for Li+, the strongest structure maker of the alkali metal cations.e A comparison of the dependence of the B
coefficients on ionic radius of the tetraalkylammonium ions with that of the alkali metal cations (Figure 1) illustrates that the hydration of the two families of cations are profoundly different, the former exhibits hydrophobic-the latter, coulombic h y d r a t i ~ n . ~We have further proposed that the structure of the hydrophobic hydration of the tetraalkylammonium ions is the p f ~ r m . ~Inasmuch ,~ as the p form is more stable with respect to temperature than the a form, the Walden products for the tetraalkylammonium cations are notably more invariant with respect to t e m p e r a t ~ r e . ~ , ~ (1) E. Wicke, Angew. Chem., 5 , 106 (1966).
(2) G . Nemethy and H. A. Scheraga, J . Chem. Phys., 36, 3382, 3401 (1962). (3) H. S. Frank and W. Y. Wen, Discussions Faraday SOC.,24, 133 (1957). (4) R. A. Horne, “The Structural Forms of Liquid Water a t Interfaces and Near Biopolymers,” Technical Report No. 22, Arthur D. Little, Inc., May 31, 1966; Office of Naval Research Contract No. Nonr-4424(00). (5) R. A. Horne, A. F. Day, R. P. Young, and N. T. Yu, “Interfacial Water Structure: The Electrical Conductivity of Aqueous Electrolyte Permeated Particulate Solids under Hydrostatic Pressure,” Technical Report No. 23, Arthur D. Little, Inc., Sept 30, 1966; Office of Naval Research Contract No. Nonr-4424(00); Electrochim. Acta, in press. (6) R. L. Kay, T. Vituccio, C. Zawoyski, and D. F. Evans, J. Phys. Chem., 70,2336 (1966). (7) R. A. Horne, “Hydrophobic Hydration,” Technical Report No. 26, Arthur D. Little, Inc., Sept 30, 1966; Office of Naval Research
Contract N o . Nonr-4424(00). (8) R. A. Robinson and R. H. Stokes, “Electrolyte Solutions,” 2nd ed, Butterworth and Co. Ltd., London, 1959. (9) R. L. Kay and D. F. Evans, J . Phys. Chern., 70,2325 (1966). Volume 7,9, Number 6 May 1968
