2003
NOTES
homologous series of aliphatic acids, alcohols, and amides show with longer chain length larger values of the B-coefficient. A high activation energy for viscosity has been interpreted as a n indication of liquid structure (for a discussion, see Glasstone, Laidler, and Eyring12). The effect of temperature upon the viscosities of solutions of ethanol, methanol, sucrose, and acetone is greater13 than its effect upon solutions of acetamide, urea, and formic acid This behavior is also in agreement with less solveiit order in the second group of solutions. The above interpretation of the viscosity data is supported by positive values calculated for the entropy of dilution of urea solution^.'^ Positive entropies of similar magnitude, otbserved for solutions of many salts, have been interpreted by Frank and Robinson’s in terms of a disruption of solvent structure.
Discussion The manner in which urea and ainides affect water structure may be analogous to the effect of ions, proposed by Gurneyg and Frank and Wen.” Like a n ion, urea may orient tlie nearest-neighbor solvent molecules, preventing their participation in hydrogen-bonded water clusters, and leading to a region of disorder about the solvated solute molecule. Disruption of the Eiolvent structure may affect the interpretation of certain properties of urea solutions. Recognizing an effect of urea upon the water activity could reconcile enthalpies of aniide hydrogen bond formation, which were calculated from the deviation from ideality of amid? solutions, 16v1’ with recent direct spectrophotometric measureinents.l8 Also, the solubilization of hydrocarbons and the denaturation of proteins by urea may reflect a disruption of the water (12) S. Glasstone, K. J . Laidler, and H. Eyring, “Theory of Rate Processes,” McGraw-Hill Book Co., New York, N. Y . , 1941. (13) Calculated from data in the International Critical Tables, a t concentrations of l0-20%. For example, E,,, is 7500 and 4500 cal./mole a t 10’ for 20% iethanol and 15% urea solutions, respectively: the values change only slightly with concentration in this range. (14) For 4 m urea the entropy of dilution is 0.034 e.u., calculated from the data of G. Scatchard, W. J. Hamer, and S. E. Wood, J . Am. Chem. SOC.,60,3061 (1938), and F. T. Gucker, Jr., and H. B. Pickard, ibid., 62, 1464 (1940). Negative entropies were calculated for acetic acid, methanol, ethanol, and acetone, in agreement with a structuremaking effect of these solvents. (15) H . S. Frank and A. L. Robinson, J . Chem. Phys., 8, 933 (1940). (16) J. A. Schellman, Concpt. rend. Trav. Lab. Carlsberg, 29, 223 (1955). (17) S. J. Gill, J. Hutson, J . R. Clopton, and M. Downing, J . Phya. Chem., 65, 1432 (1961).
(18) I. M . Klotz and J. 8. Fransen, J . Am. Chem. SOC., 84, 3461 (1962).
structure, since the solubility properties of hydrocarbons in water result in large part from an ordering of the solvent.
Potentiometric Determination of Nitrates in Benzene
by G. Scibona and B. Scuppa Comitato Nazionale Energia Nucleare, Industrial Chemistry Dicision, Centro S t u d i Nticleari, C A S A C C I A , Rome, Italy (Received October I S , 1068)
Studies of indicator electrodes and potentiometric titrations in hydrocarbon solvents have been the subject of several researches. 1-8 Attenipts have been made to check the validity of Nernst’s law for hydrogen ion or chloride ion in these solvents using glass or silver chloride electrode^.^-^ An extension of these studies to solutions containing other types of ions is limited by the availability of suitable electrodes. I n the present investigation a nitrate electrode has been developed to carry out the potentiometric determination of nitrate ion in benzene solution of tertiary and quaternary ammonium nitrates. The dependence of the e.1ii.f. of the cell on the concentration of the salt supports the proposal that the measured potentials are indeed electrochemical potentials. The electrode is described in Fig. 1. Purified mercury and iiioiiohydrate mercurous nitrate, prepared as by were triturated together in the presence of a little benzene to give a dense mercury-saturated paste. Mercury was introduced into the clean and dry cell a t A, the paste was transferred to the mercury pool B, and a benzene solution of an alkylainmonium nitrate of required concentration covered the paste surface. (1) L. Fischer, G. Winkler, and G. Jander, Z . Elektrochem., 62, 1 (1958). (2) J. T. Stock and W. C. Purdy. Chem. Rm., 57, 1159 (1957). (31 H. B. van der Heijde and E. A. M. F. Dahmen, A n a l . Chem. Acta, 16, 376 (1957). (4) L. Lykken, P. Porter. H. D . Ruliffson, and F. C. Tuemmler, Ind. E n g . Chem., A n a l . Ed., 16, 219 (1944). (5) (6) (7) (8)
A. Gemant, J . Chem. P h y s . , 12, 79 (1944). A. Gemant, ibid., 10,723 (1942). A. Gemant, ibid., 13, 146 (1945).
A. Gemant, “Ions in Hydrocarbons,” John Wiley and Sons, Inc., New York, N. Y.,1962. (9) P. P,ascal, “Nouveau Trait; de chimie Minerale,” Vol. V , Masson et Cie Editeurs, Paris, 1962.
V o l u m e 68, Number 7 J u l y , 1964
NOTES
2004
The bridge between the two half-cells was a quartz capillary of about 1-inn1 i.d. supported by the glass stoppers C. In the cell the left branch and the bridge were filled with the solution of higher concentration. The cell was placed upon a polystyrene foam block within a n earthed metal box. Connections from plati-
D
Tri-n-heptylammonium nitrate and methyltri-n-octylanimoniurn nitrate were used in separate experiments. The graph of Fig. 2a shows the relation between the obtained potential in mv. and log CIICz for both the
C 40
1
E=’-
1 1
2
3
4
Figure 2. ( a ) E.m.f.-concentration plots for tertiary ( 0 ) and quaternary (0)ammonium nitrates in benzene; ( b ) and ( e ) e.m.f.-conductivity plots for tertiary ( 0 )and quaternary (0)ammonium nitrates.
nitrates. The curves are reproducible within lo%, clearly showing the ability of this electrode to carry out the potentiometric determination of the nitrates. In Sernst’s relation for the concentration cell with reversible anion
Figure 1. Half-cell used for e.m.f. measurements.
nurn wires (D) led to a 3letroh:n E 336 potentiograph (maximum admissible electrode resistance = 1000 Mohms). Two days of aging was necessary to obtain stable and reproducible potentials. The potentiometric determination was carried out by changing stepwise the concentration Cz of the benzene solution of the alkylanimoniurn nitrate in the right half-cell. The higher concentration CL in the left half-cell was held constant during the measurement. T h e Journal of’ Physical Chemistry
( t = transport number of the nonreversible ion, (C)i = ionic concentration, y = activity coefficients ; indices 1 and 2 referring t o the two half-cells) it is necessary to use the ionic activit,y. Since the ion concentrations are very low the activity coefficient y may be taken as unity. Therefore, the nitrate ion concentration (C), may be related to the salt concentration C by (C) i = g / K C with K the equilibrium constant for the ion pair formation. By substituting the ratio of the square roots of the salt, concentration in place of the ratio of the ionic concentration (1)becomes
NOTES
2005
E
=
5% log
c1 -
A Gas-Phase Electron Diffraction
C,
Study of cis-Dibromoethylene
t, = 0.65 for the tertiary and quaternary ammonium
nitrates is calculated from the slopes of Fig. 2a. Equation 2 is valid only in isodielectl-ic media in which K is constant since K depends on the dielectric coiistant D.I0 However, by increasing the concentration from 0.04 to 0.1 during the experiment, D increases from 2.3 to 3. To account for this effect it should be necessary to know the K values as a function of D. I) must be noted that by using the ratio of the conductivity in (caused by ion migration) of the solution in place of the ratio of the ionic concentration, as recently suggested,5a7we have the relation
E
=
WZl
11% log mz
which is followed by the experimental curve shown in Fig. 2b. The cation transport numbers, t, =: 0.085, for both tertiary and quaternary ammonium nitrates in the concentration range 0.04-0.1 are obtained from Fig. 2b. The disagreement with the transport numbers previously calculated shows the falilure of relation (2). A salt concentration effect for the t-amnioniurn ni-. trate cation transport number is evident (Fig. 2c). A value of 1, = 0.13 in the concentration range 0.040.4 M t-ammonium nitrate is obtained. The relatively low values for the cation transport number are understandable on the basis of their large cation size compared with that of the nitrate anion. The dependence of t, for the t-ammonium salt on the concentration range may be explained with the decrease of the nitrate transport number t ~ o , ,as a consequence of triple ion formation and of the formation of a higher order of aggregation with an increase of the concentration. A check of the m values a t two different frequencies will be necessary, since the frequency independence of m is a proof that the contribution of rotation of dipolar groups to the conductivity is negligible. Acknowledgment. The authors are greatly indebted to A. Gemant for many valuable suggestsions and for a critical review of the manuscript. ~
(10) R 31.Fuoss and C. Kraus, J . Am. Chem. Soc.. 5 5 , 1019 (1933).
by 11ichael I. Davis, Harvey A. Kappler, Department of Chemistry, T h e University of T e x a s , A u s t i n 12, Texas
and David J. Cowan Department of Physics, T h e Cniversity of Texas, A u s t i n 12, Texas (Received November 27, 1.965)
The cis and trans isomers of the 1,Zdihalogenoethylenes have been the object of a great deal of chemical curiosity, on account of the cis isomer being, somewhat unexpectedly, the more stable.' It was felt that some further light might be thrown on the subject from an evaluation of the molecular dimensions of a t least one of the isomers. The intensity data for cis-dibromoethylene were collected by Cand. Real Almenningen, with the University of Oslo diffraction u n k 2 Reliable intensity data could only be obtained out to about s = 20 k - l . This was due to the comparatively large valbe of I , = 0.12 for the Br-Br pair and to the significant difference in atomic phase changes for carbon and bromine; both factors introducing a marked damping effect upon the molecular scattering. The preliminary handling of the experimental intensity data followed much the same procedure as that outlined by Bastiansen and S k a ~ i c k e . ~ The more conventional approach would have been to carry out the structural analysis on the radial distribution curve. There proved, however, to be a number of objections to the adoption of that line of procedure. I n order to obtain a radial distribution curve, which is built up of approximately Gaussian peaks, it is necessary to multiply the experimental intensities by a function
ZlZ,/(ZL - F(s)J(Z, - F ( s ) , )
(1)
where i and j are two of the atomic species present in the molecule, and F ( s ) is the appropriate atomic scattering factor. Repardless of how many differing clement pairs there may be present in a particular molecule, it is necessary to find a form for the given function that will satisfy all of them. The function Z , / ( Z , F ( s )J does not vary appreciably among atoms of closely (1) IZ. M. Soyes and R G . Diekinson, J Am Chem. SOC.,6 5 , 1427
(1943). (2) 0. Bastiansen, 0. Hassel, and E. Risberg, Acta Chem. Scand., 9, 232 (1955). (3) 0. Bastiansen and N. Skancke, Adaan. Chem. P h y s . , 3, 323 (1960).
Vol71me 68, Number 7
J u l y , 1964
