JAMES F. SKINNER AND RAYMOND lli. Fuoss
2998
Electrostriction in Polar Solvents. 11182
by James F. Skinner and Raymond M. Fuoss Contribution No. 1769 f r o m the Sterling Chemistry Laboratory, Yale University, New Haven, Connecticut (Received M a y 26, 1964;
Viscosities and densities in acetonitrile of solutions (maximum concentration 3 g./lOO nil.) of a variety of compounds satisfy the equations 7 = ro(l Ac"' Bc) and p = po g ( 1 - v,p0) where r is viscosity of the solution, yo is that of the solvent, p and po are the corresponding densities, c is molar concentration, g is weight concentration, and us is the volume in solution of the solute. For ideal hydrodynamic spheres, 400B = M'U,, where 111 is molecular weight of solute. This relationship is closely approximated by tetraphenyllead, -tin, -silane, and -methane. Triphenylaniine, -phosphine, -arsine, and -stilbine contribute less to viscosity than one would expect from their volumes. By coniparing mono-, di-, and trinitrobenzenes and -toluenes, it is shown that the total scalar dipole field (rather than the net molecular moment) determines the effect on viscosity; For example, p- and nz-dinitrobenzene have practically equal B-values. Results with the three nitroanilines and p-phenylenediamine suggest that hydrogen bonding to acetonitrile occurs with these compounds. Smaller diffusion constants are found for compounds which show large B-values, confirming the hypothesis that local dipole fields tighten solvent structure around polar molecules.
+
In the aprotic solvent, acetonitrile, viscosity Bcoefficients and the corresponding density coefficients were found to be additive for solutions of 1-1 electrolytes.3 For large ions, such as those of tetrabutylammonium tetraphenylboride, the B-coefficient could be calculated from the density-concentration coefficient on the assumptions (I) that the latter coefficient measures the volume of the solute in the solution and ( 2 ) that the B-coefficient (times concentration) has the Einstein value of five-halves the volume fraction. For smaller ions, the observed B-coefficient was found larger than the Einstein value calculated from the density coefficient, while for neutral molecules, especially those containing no permanent dipoles, the viscosity coefficient was smaller than that predicted. The hypothesis was made that, in the former case, electrostriction in the solvent near the ion enhanced the viscosity, while, in the latter case, the electroneutral nzolecules could more easily slip through holes in the solvent. The primary purpose of this paper is to present measurements on solutions of dipolar solutes ; some neutral and electrolytic solutes will also be considered for comparison. It will be shown that polarity in a niolecule increases its contribution to solution visT h s Journal of Physical Chemistry
+
+
cosity, presumably due to local dipole-dipole attraction between solvent and solute. The effect is proportional to the summed polarity. For example, B for o-dinitrobenzene is 0.243 and for p-dinitrobenzene (which has zero net moment) is 0.266. This and similar results show that the effect on viscosity is due to short-range forces around the individual dipoles in the molecule, rather than to its long-range average field.
Experimental Methods and apparatus have already been described. All measurements were made in acetonitrilea at 25.00", except those on the tetraphenyl series (D, E, F, G), which were made in carefully dried and fractionated benzene (f.p. 5.51"). Some of the solutes were used as received, and some were further purified. Their description follows; a code for identification is given in Table I. For later (1) This paper is based on part of a thesis presented by James F. Skinner to the Graduate School of Yale University in partial fulfillment of the requirements for the degree of Doctor of Philosophy. (2) Grateful acknowledgment is made to the donors of the Petroleum Research Fund, administered by the -4merican Chemical Society, for partial support of this work. (3) D. F.-T. Tuan and R. M .Fuoss, J . P h w . Chem., 67, 1343 (1963)
ELECTROSTRICTION IN POLAR SOLVESTS
2999
use, the specific volumes (+0.005 ml. Ig.) of the pure solutes a t 25" are given in the last column of the table. Triphenylsulfonium tetraphenylboride was preciloitated by mixing a 5% aqueous solution of triphen:ylsulfonium chloride (Aldrich) with an equivalent soILution of sodium tetraphenylboride. It was thoroughly washed with water and dried a t room temperature. It decomposes a t 110". S o solvent satisfactory for recrystallization was found. Diphenyliodonium iodalte was prepared4 from iodosobenzene and iodoxybenzene, and diphenyliodonium tetraphenylboride was prepared from a solution of the iodate by precipitation with sodium tetraphenylhoride, also in aqueous solution. The salt was purified by dissolving 1g.in 3 nd. of ethanol, followed by 30 ml. of hot acetone. After filtering the hot solution, 200 ml. of distilled water was added. Fine colorless needles were obtained, which were dried at 70" and 1 mm., m.p., 140". Matheson nhexadecyltrimethylaimmonium broimide was used as received.
Table I : Catalog of Compounds Spec1 fi c volLme
A B C D
E F G H
I J K L M N 0
P
Q
R
S T U
Ph3S.B Fh, PhzI.BPh4 n-ClaHa7NMe3 .BPh, PhrPb Ph& PhrSi Ph,C PhaSb Ph3As Ph3P Ph3CH Ph3Y Mesitylene 2,4-Dinitrotoluene o-Dinitrobenzene m-Dinil robenzene p-Dinitrobenzene 1,3,5-Trinitrobenzene Phthalonitrile
V
1,2-Dimethyl-3-nitrobenzene 1,3-Diniethyl-2-nitrobenzene 1,4-Din~ethylnitrobenzene
W X Y Z
p-Phenylenediamine o-Ni trottniline rn-Nitroaniline p-nitro aniline
0 856 0 807 0 890 0 605 0 670 0 835 0 850 0 675 0 770 0 865 0 895 0 860 1 165 0 665 0 640 0 635 0 615 0 595 0 750 0 880 0 910 0 890 0 880 0 690 0 700 0 700
Triphenylstibine (I< and K Laboratories) was recrystallized from petroleum ether (3.5 g. in 15 ml.), m.p. 53.6'. Triphenylarsine (K and K) was also re-
crystallized from petroleum ether (2.0 g. in 10 ml.), m.p. 60.1". Triphenylphosphine (Aldrich) mas recrystallized from diethyl ether (5.9 g. in 15 ml.), m.p. 80.1". Triphenylamine (Matheson) was recrystallized from ethyl acetate (2.8 g. in 12 ml.), m.p. 126.2". Triphenylmethane (Xatheson) was recrystallized from 95yc ethanol (3.5 g. in 10 ml.), m.p. 93.8". Tetraphenyllead (K and K) was recrystallized from benzene (3.0 g. in 30 ml.), m.p. 231". Tetraphenyltin (Matheson) and tetraphenylsilane (K and K) were recrystallized from pyridine; m.p. 231 and 239") respectively. Tetraphenylmethane (K and K) was recrystallized from benzene (3.5 g. in 140 ml.), n1.p. 284". 1,3,5-Trinitrobenzene (Eastman) was recrystallized from acetone by chilling a 50y0 solution to -80". Phthalonitrile was recrystallized from ethanol (4 g. in 30 ml.), m.p. 137". o-Dinitrobenzene (K and K) was recrystallized from 95% ethanol (5 g. in 50 ml.), m.p. 118". o-Nitroaniline was recrystallized from water (2.5 g. in 350 ml.) and dried a t 5 p at room temperature, m.p. 73.8'. m-Nitroaniline was recrystallized from water (4.3 g. in 150 ml.) and also dried at 5 p a t room temperature, m.p. 114". Me-
sitylene, 1,2-dimethyl-3-nitrobenzene, 1,3-dimethyl-dnitrobenzene, 1,4-dimethylnitrobenzene, 2,4-dinitrotoluene, p-phenylenediamine, m-dinitrobenzene, and p-dinitrobenzene were used as received from Eastman. Results and Discussion Up to concentrations of several weight per cent, specific volumes of many two-component systems can be reproduced to a precision of about 0.02% by the linear equation 2,
=
(1 - w)vo
+ vsw
(1)
where uo is the specific volume of the solvent, w is the weight fraction of solute, and is the specific volume of the solute in the solution. Excluding aqueous and other hydrogen-bonded solutions, us is always positive and usually is only a little larger than the specific volume of the pure solute. Translating into densities, (I) becomes ~1~
P :=
PO
== PO
+ g(1 -
(2)
UsPo)
+ (MC/1000)(1 - U@")
(3)
where g is weight concentration (g. of solute/nd. of solution), d l is molecular weight of solute, and c is molar conceiitration (moles/l. of solution). The molar volume V in solution is JIr, ml./mole. For (4) H J. Lucas and E R Kennedy, O r g . S y n , 22, 52 (1942)
Volume 68, S u m b e r 1 0
October, i.g(s4
tJ.mm F. SKINXER ASD RAYMOND AI. Vvoss
3000
deterniining vs from density data, it is convenient to define a function F ( p , g ) , where F ( p , g ) = (PO - P
+ 9)//PO
(4)
= “9
Then a plot of F against g determines i’8 as the slope of the line through the origin and the experiniental points. Since all our data gave linear plots, the large body of data may be compactly sunimarized by giving the values of v,; these are collected in Table 11. They reproduce our observed densities within 0.0270 up to g = 0.02-0.03.
Table 11: Constants for Density and Viscosity Equations U8
B
776 697 904 581 714 875 864 0 696 0 778 0 874 0 900 0 862 1 171 0 698 0 672 0 655 0 629 0 573 0 870 0 862 0 877 0 877 0 814 0 725 0 733 0 721
1 80 1 67 1 24 0 85 0 80 0 75 0 72 0 40 0 38 0 35 0 37 0 33 0 031 0 279 0 243 0 261 0 266 0 37 0 201 0 163 0 148 0 158 0 224 0 232 0 245 0 294
Key
A B C
I1 E F G
H I J K L 11 N 0 P
Q
R S T
u T.’
w x Y Z
0 0 0 0 0 0 0
R,
LOSB’ I.1
3 09 2 75 3 41 1 65 1 86 2 33 2 23 1 14 1 23 1 32 1 53 1 36 0 26 1 53 1 44 1 55 1 58 1 74 1 57 1 08 0 98 1 05 2 07 1 70 1 77 2 13
4 08 3 74 5 03 4 97 4 97 4 89 4 80 4 60 4 55 4 49 4 45 4 39 3 83 3 60 3 55 3 53 3 49 3 66 3 53 3 72 3 75 3 76 3 27 3 41 3 42 3 40
4, 5 64
5 41 5 41 5 09 4 4 4 3 3
3 4 3
0 3
3 3 3 3 3 2 2 2 3 3 3 3
99 92 87 99 91 81 01 75 79 53 39 46 49 89 18 96 87 91 29 33 39 62
Q 1 1 1 1 1 1 1 0 0 0 0
0 0
60 63 50 14 05
02 03 65 63 60 68 63 088 88 86 95 01 21
0 0 0 1 1 0 72 0 50 0 45 0 48 1 02 0 94 0 97 1 18
Similarly, a t low concentrations, solution viscosities are reproduced by the Jones-Dole equation5 7 70 =
1
+ -4~’’’ + BC
(5)
where A can be theoretically coniputedfi from single ion conductances or be empirically determined by graphical nicthods. (For nonelectrolytes, A is zero, of course.) For a system which iiiay be represented by spheres in a hydrodynamic continuum, Einstein7 showed that the viscosity increment depends on the volume of the solute, so that
RC = l O O O B g ’ S P The .Journal of Physical Chemistry
=
5+ 2
(6)
where 4 is the volume fraction of solute.s I’or such an ideal system
400BiM = v , (7) and viscosities are predictable from densities, and vice versa. For large solute particles, such as tetraphenylniethane or the ions of tetrabutylanimoniuin tetraphrnylboride,3 the observed coefficients do in fact satisfy ( 7 ); smaller ions, however, give higher viscosities than one mould expect from their volumes, while neutral, and especially nonpolar, molecules give lower viscosities. For analysis of the viscosity data, it is convenient to define a viscosity function G(7,g) by the equation G(7,g)
=
0.4(q’qO - 1 - Ac”’) = 0.4Bc
(9)
= Qvsg
where Q is unity for a n ideal system. system
Q
(8)
=
400B/Mv,
For a real (10)
where B and v, are the observed values. Over our working range Of concentrations (g 5 0.03 g . ’ ml.), G was a linear function of g or e. The viscosity data are suinniarized by the B and Q values given in Table 11. (For the electrolytes A, B, and C, the coefficient A is zero within our experimental error; the Be term is very much larger than the square root term and effectively swamps it.) We consider first the electrolytes. For triphenylsulfoniuin tetraphenylboride, B = 1.80. Subtracting Tuaii’s value3 of 0.68 for the anion, the viscosity coefficient B + for the cation is 1.12. This is very much larger than the values found for quaternary ions ( B + = 0.33 for lTeJ+. 0.68 for Bu&+); the positive charge in the Ph3S+ ion is not sterically shielded as is the charge in the quaternary anmioniuni ions, and interaction with the polar acetonitrile iiiolecules is correspondingly greater. The effect of the bare charge is emphasized if we compare the value 1.12 with the neutral model compound triphenylmethane for which B = 0.37. The volumes of Ph3CH and Ph3S+ must be very nearly the same ; the ion-dipole interaction (5) G. Jones and M . Dole, J . Am. Chem. Soc., 51, 2950 (1929). (6) H. Falkerihagen and 51. Dole, Physik. Z . , 30, 611 (1929). (7) A. Einstein. Ann. Phyaik., 19, 289 (1906); 3 4 , 591 (1911). (8) As the coefficient of concentration expressed in moles/l., B has awkward dimenjions. I’hysicaliy more significant is lO3L3/MM,the coefficient of weight concentration g, and still more useful intuitively is lO3B/.Ma,, the coefficient of volume fraction (ml. of solute/ml. of solution). For similar reasons, eq. 2 ia preferred to eq. 3 for density. For both viscosity and density, it is the volume of solute which is
physically significant, not the number of moles.
ELECTROSTRICTION IN POLAR SOLYESTS
3001
in the latter produciss twice the viscosity increment due t o the Einstein volume effect. Diphenyliodonium tetraphenylboride is similar: B = 1.67, whence B+ = 0.99. For diphenylmethane and diphenyl ether, Tuan3 finds 0.21, so here the unshielded charge on the iodine produces four times the viscosity increment in addition to that due t o the volume of the ion. Pcrhaps part of the large B for this ion niay be due to its approximately cylindrical shape, but the model cornpounds must have similar geometry. The effect of shape may, however, be shown in the case of n-hexadecyltriniethylanimoniuni bromide, for which B = 1.24. Using B- = 0.25 from Tuan’s result with tetrabutylaninioniuni bromide, B + here is 0.99, which is considerably larger than 0.68 for Bu4N+ which has about (16 LIS. 19) the same number of carbon atoms, but symmetrically arranged. The charge is also Less shielded in the hexadecyltriniethylainnioiiiu~ii ion. From these results and those of Tuan, it is clear that the viscosity coeficient B measures at least three properties of an ion in a given solvent: its volume, its shape, and the intiensity of its field. Studies of selected series of electrolytes promise to show a number of interesting correlations between B ’ X and structure. The neutral series tetraphenyllead, -tin, -silane, and -methane, especially the last three, behave practically like ideal hydrodynamic spheres in benzene. The radii, calculated from us are 5.09, 4.99, 4.92, and 4.87 respectively; their volumes are therefore about an order of magnitude greater than the volume of a beiizene niolecule, and that this solvent appears like a cotitinuuiii to the tetraphenyl molecules is not at all surprising. That the B values decrease froin 0 . 6 2 for lead to 0.716 for carbon as the central atom is merely the numerical consequence of the fact that B is the coefficient of concentration in moles I . The Q-values, which eliminate the trivial effects of nioleaular weight, are 1.14, 1.05, 1.02, and 1.03. The triphenyl compounds (H-L inclusive) present a different picture; they all increase viscosity by coiisiderably less than one mould expect from their volumes. Their &-values average to 0.64, which is nearer to the values for the diphenyl series (Q = 0.50 for diphenyl ether and diphenylniethane) than to those for the tetraphenyl series just discussed. One similarity, which niay be only incidental, between the di- and tri- series should be mentioned , the Q-values for the compounds with onc proton are higher than for the other members of the series: Q(PhzSH) = 0.71 and Q(Ph,CH) == 0.68. The triphenyl series, while less effective than Stokes spheres in producing viscosity, are more effective than the one irialkyl compound measured (trin-butylamine, for which Q = 0.26 in acetonitrile).
x.,
The volume effect is, however, present in the neutral series PhZCH2, Ph,CH, Ph4C, for which Q = 0.49, 0.68, and 1.03, in that the sequence of coefficients a t least increases regularly with increasing size. In Table I1 are also given values of the radii of the spheres equivalent to the solutes, as calculated from us (kin the table) and from the viscosity, assuming the validity of eq. 7 ( R , in the table). For the three electrolytes, cationic radii R+ are tabulated ; these were obtained from our coefficients using Tuan’s values3 for the BPI??- and bromide ions. When Q = 1, R, = R,. As before, we assume that only R , has physical significance. All the latter are in good agreement with expected values. But paralleling the Qvalues, the viscosity radius is too large for the electrolytes, and too small for most of the other compounds. Finally, we consider the substituted benzenes. With no polar group, the viscosity coefficient B smaller than the Einstein value, as shown by niesitylene. From vus,we obtain R , = 3.83 as the radius of the equivaleiit sphere calculated from the density, which is quite reasonable. But from viscosity, if we assume Bc == 5+,’2, R, calculates to 0.79 which is obviously absurd. The value of Q is only 0.088, as compared with unity for ideal spheres. Just as soon, however, as a polar group such as the nitro group is attached to the ring, a significant increase in the viscosity coefficient appears. The three niononitro conipounds (T, U, V) have Q near 0.5; perniuting methyl groups around the polar group has very little influence on the coefficients. With two nitro groups, Q practicallg doubles. Again, insertion of a methyl group has little effect; in fact, Q for 2,4-dinitrotoluene is a little smaller than For nz-dinitrobenzene. This might nieaii that the methyl group adjacent to the 2-nitro group sterically blocks attraction of solvent dipoles. The same steric argument might be applied also to the ortho, meta, para sequence, where there is a small but systematic increase of Q as the two nitro groups get farther apart. The same sequence is observed with the nitroanilines, compounds X, Y, and 2 . That dipole strength is more significant than the nature of the dipole is shown by the similarity between phthalonitrile and the dinitro series; the hionients of the nitrile and the nitro groups are nearly the same. With three nitro groups, Q increases still more. These results are summarized in Fig. 1 where Q is plotted against the scalar sum of the nionients in the niolecule (NOz, 4.0; SH:, 1.5; CH,, 0.3). The point marked “ P A ) ’ is for picric acid.3 For no iiionient, there is a small voliinie effect (cT. points for niesitylene and xylene,), and then Q increases about linearly with the summed moments. The viscosity-creating potential of a moleT’ol~~rne68. ,\-umber
10
October, 19Gq
JAMES F. SKIKXER A N D RAYMOND AI. Fuoss
3002
2.0
0
0
B I
Figure 1. Dependence of viscosity effect on total polarity.
cule is clearly the result of short-range forces; Fig. 1 proves this point. The most direct proof. however, is the large values of Q for p-dinitrobenzene and 1,3,5trinitrobenzene. Both of these niolecules have zero net moments (just like xylene and mesitylene), but still they give large viscosity increments. We conclude that the viscosity increase is due to local dipoledipole interaction between individual polar groups and solvent niolecules, independent of the net moment of the whole molecule. It will be noted that the three nitroanilines lie above the line in Fig. 1, as does the point W for p-phenylenediamine. This means that these compounds increase viscosity even more than corresponds to their total polarity. Hydrogen bonding between the amino group and acetonitrile, which would increase the volume of the kinetic entity, would account for this increase. JTe plan to investigate the nitrophenols which should show a similar effect if hydrogen bonds are involved. For comparison with the earlier r e s ~ i l t s the , ~ coefficient B is plotted against the molar volume Jlv, in Fig. 2 . Points for ideal systems would lie on the solid line representing ey. 7. The lhree electrolytes. despite their large volume, lie considerably above the line, showing much stronger interaction than tetrabutylainnionium tetraphenylboride (which is almost ideal). The tetrapheriyl series gives a clusttr of points nearly on the line, showing that the viscosity effect of these compounds is alinost a pure volume effect; in other words, for these compounds the solvent behaves like a continuuni. The tripheliyl series, however, gives a cluster below the line, showing that these compounds can slip through acetonitrile more readily than spheres equivalent in voluiiie to those deduced from the density nieasurenients. The nitro compounds on this plot form a confused cluster near the origin: as already shot^ in Fig. 1, they are best understood on the basis of total polarity. The Journal of Phusical Chemzstry
Figure 2. Comparison of viscosity and density parameters: key in Table I.
The observed correlation between the summed scalar moments of a niolecule and its capability to increase viscosity in a polar solvent supports the hypothesis that short-range dipole-dipole forces constrict the solvent around the solute molecule. We would then expect a n analogous correlation between suinrned moments and diffusion coefficients for molecules of the same shape and size. In Table 111 are summarized
Table 111: Diffusion Constants in Acetonitrile a t 25" 1060
C
Mesitylene
2 772 2 658
0 0594 0 1362 p-Nitroaniline
2 459 2 422
0 0252 0 0549 1,3,5-Trinitrobenzene 0 0249 0 0551 0 0850
2 194 2 154 2 108
the results for diffusion of mesitylene, p-nitroaniline, and 1,3,5-trinitrobenzene in acetonitrile a t 25" (concentration c in moles, 1.). The measurements were made by the Gouy interference methodg; we are grateful to Professor P. A. Lyons for the use of his apparatus and t o A h . H. E. Meissner for making the measurements. The data, extrapolated linearly to (9) L. J. Gosting, E. 1LI. Hansen, G. Kegeles, and ICI. 8. Morris, Rec. Sci. Insti.., 2 0 , 209 (1949).
COXDUCTANCE OF COPPERm-BENZENEDISULFONATE HEXAHYDRATE
zero concentration, give the following values for the liniiting diffusion coefficients : mesitylene, 2.85 X p-nitroaniline, 2.49 X and lj3,5-trinitrobenzene, 2.22 X As expected, the polar molecules diffuse more slowly than the nonpolar mesitylene.
3003
The differences in 105Dare 0.36 and 0.133, respectively, for p-nitroaniline and trinitrobenzerie compared to mesitylene. Thus both viscosity coefficients and diffusion constants show the effects of local iiiolecular interaction.
Conductance ad' Copper m.-BenzenedisnlfonateHexahydrate in N-Methylpropiconamide from 20 to 4OC'l
by Thomas B. Hoover National Bureau of Standards, Washington, D . C.
(Received M a y $6, 1964)
The conductance of copper in-benzenedisulfonate hexahydrate in N-methylpropionaniide was measured a t 5" intervals froin 20 to 40°, and in the concentration range of 3 X M . Viscosity, conductance, and solubility observations indicate that water to 1 X of crystalliza,tion does not remain associated with the electrolyte in solution. The FuossOnsager conductance equation represents the data satisfactorily, although there is a barely significant, temperature-dependent contribution froin higher order terms in concentration. The ion-size parameter, &, increases from 3.0 to 4.5 with increasing temperature, while the mean hydrodynamic (Stok.es) radius is 4.9 8. The limiting equivalent conductance is 25% larger than that of potassiuin chloride in the same solvent.
Introduction Despite the reniairkably high dielectric constant of N-methylpropionainkde (NMP) , previous conductance measurenients2 indicated that potassium chloride was appreciably associated in this solvenl,. That conclusioin was based primarily on the unrealistically small values of the ion-size parameter, d J , needed to fit the conductance data to the Fuoss-Onsager equation for strong electrolytes. Both the ion-pair association of potassium chloride and the large effect of salts on the viscosity of solutions in NBIP were indicative of pronounced ioiisolvent interactions. Such effects were expected to be enhanced by more highly charged ions; hence, the present conductance study of the 2-2 electrolyte, copper n~-benzenedisulfcinate (CuBDS) was undertaken. This salt is not only readily soluble in NRlP but previous
conductance measurements3 have shown that it is fully dissociated in aqueous solution, in contrast to most 2-2 salts. Dawson and co-workers4have measured conductances of a nuinber of multivalent electrolytes in the related solvent, K-metjhylacetaniide. Apart froin some anoinalies that were attributed to traces of acetate ion in the solvent, the salts all behaved as typical strong electrolytes. (1) Presented, in part, before the Division of Physical Chemistry at the 145th National Meeting of the American Chemical Society, New Tork, N. T..September 9-13, 1963. (2) T. E. Hoover, J . P h y s . Chem., 6 8 , 876 (1964). (3) G. Atkinson, M . 1-okoi, and C. J. Hallada, J . A m . Chem. Soc., 8 3 , 1570 (1961). (4) L. R. Dawson, J. W. Vaughn, G. It. Lester, >I. E. Pruitt, and P. G. Sears, J . Ph:ys. Chem., 67, 2i8 (1963).
Volume 68, Sumber 10 October, 1864
