NOTES
746
Vol. 63
equalization: the electronegativities (potentials) of two atoms are equalized during the formation of a chemical bond. The electronegativity X of an atom is a function of the net atomic charge q. Developing X(p) in powers of q and disregarding higher terms
known that different screening constants must be used for different properties (ionic size, ionization potential, electron-affinity, etc.).ll I n fact, HartreeI2 recently proposed the use of the term screening parameter instead of screening constant because the quantity s is not constant for variations of any of the variables of which it might be considered a X ( d = X(0) ddX/dd(O) (1) function. Thus, as a first approximation we may suppose Screening constants for electronegativities are that there is a linear relationship between electro- not known. However, if we use Mulliken's defininegativity and charge. P a ~ l i n g , Hais~insky,~ ~ tion, X = I E , and since the ionization potential and Daudel and Daudel'O have shown how to calcu- ( I ) is generally much greater than the electron-aflate (dX/dq) from the electronegativity scale. These finity ( E ) , we may use the ionization potential authors admitted that the screening constant of one screening constants calculated by Kohlrausch'* valence electron for another is independent of their without introducing very serious errors. electronic wave functions. A more accurate treatWe have followed this procedure for the calcument is as follows: lation of the charge distribution of the hydrogen When one goes from an element of atomic num- halides, using the corrected electronegativity values ber 2 to the next one, (Z l),the effective nuclear of Mulliken's scale as given by Pritchard and charge increases by 1 - Sk, where sli is the screening Skinner.' Applying equations 2 and 3 me obtained constant of the differentiating electron of the ele- the values: AXH* = 2.62; AXF- = 2.40; AXci1). When the element 2 loses one va- = 1.92; AXBr- = 2.39; AXl- = 2.42. ment (Z lence electron the effective nuclear charge increases Table I shows the values of q obtained from equaby si, the screening constant of the lost electron. tion G and the values of the dipole moments (/Jralod) Hence the increase in the electronegativity of an obtained by multiplying the charge q by the exatom B caused by a positive unitary charge result- perimental interatomic distances H-X. This is ing from the complete removal of the i electron of strictly correct only if the pure covalent function atom B is QH - x describes a state of zero dipole moment.14 For the molecules H X the theoretical calculations nre contradictory,'6016 and our values are in good Similarly the decrease in the electronegativity of an agreement with the experimental values of the diatom A caused by a unitary negative charge re- pole moments17-'9 shown in the second column of sulting from the complete acceptance of the j elec- Table I. tron by atom A is given by TABLE I
+
+
+
+
(3)
where sj and s h are, respectively, the screening constants of the differentiating electrons of elements (2 1) and 2. From equation 1
+
n
x,(-np)
XA(O) -
=
AXj-
(4)
j=1
=
XB(O)
+
m
Axi'
(5)
i= 1
I n these equatioiis n and m are the oxidation numbers of atoms A and B, and q is the fractional charge (in electronic units) the atoms will acquire when their oxidation numbers are, respectively, - 1 and 1. From the principle of electronegativity equalization it follows that
+
9 =
XA(0) - XB(0)
2 AXj- + 2 AXii
j= 1
HF HC1 HBr HI
(e1ect;onic-1 units
0.32
.16 .10 .05
Acknowledgments.-We
BexDti
D 1.91 1.03 0.78 0.38
CLonlodl
D 1 39
1 01 0.68 0.38
wish to thank Professor
G. Beck and Dr. J. Danon for valuable discussions.
and XB(fmg)
Conipound
(6)
i=l
It is seen that in principle it is possible to calculate the charge distribution of simple molecules from electronegativity data. However, a difficulty arises in the choice of suitable screening constants. The screening constant of one valence electron for another varies with the electronic wave function and with the atomic number. Besides, it is well (9) M. Haissinsky, J . Phys., 7 , 7 11946). (10) P. Daudel and R . Daudel, ibid., 7 , 12 (1946),
(11) L. Pauling and J. Sherman, 2. Krzst., 81, 1 (1932). (12) D. R. Hartree, Rev. Mod. Phgs., 30, 63 (1958). (13) K. W. F. Kohlrausoh, Acta P h y s . Austr., 3 , 452 (1949). (14) C. A. Coulson, Proc. Roy. SOC.(London), A207, F Y (1951). (15) D. Z. Robinson, J . Chem. Phys., 17, 1022 (1919). (16) P. N. Soliatz, zbzd., 22, G95 (1954). (17) C. T. Zahn, P k y s . Reu., 27, 455 (1926). (18) R. P. Bell and I. E. Coop, Trans. Faraday Soc., 34, 1209 (1938). (19) N. B. Hannay and C. P. Smyth, J . A m . Chem. Soc., 6 8 , 171 (1946).
BOILING POINTS OF NORMAL PERFLUOROALKANES BY WILLIAM POSTELNEK Materials Laboratory, Il'righl Air Development Center, United Statea Air Force, Wright-Patterson Air Force Base, Ohio Receiued August 88, I968
Previously, no simple equations have been proposed which relate the number of carbon atoms of fluorocarbons to their boiling points. However, many empirical relationships have been proposed for calculation of boiling points of normal hydro-
.
May, 1959
NOTES
747
TABLE I BOILINGPOINTS OF NORMAL PERFLUOROALKANES (2' = OK. ) No. of carbon atoms
(4
I 2 3 4
5 G
7
No. of T(obsd.)
1452 1948 2344 2376 271.36 2744 2777 303. l8 3209 3301° 381 l1 33312 35513 355.5'4
T(ca1cd.)
142 104.2
carbon atoms
(4
8
9 237.2 273.4
10
304.8
11 12 13 14 15 1G
333.2
T(obsd.)
37713 380L1 395.5-6. 513 398. 316 4001' 417. 216 42311 433. 816 4G6-911
5 11-1 316
T(ca1cd.)
370,6 400 418.8 436.2 452.4 467. G 481.8 495.3 508
357.2
carbons. Notable among these is that of Egloff, Sherman and Dull,' who observed this relationship to exist (equation 1) T = a log (n + b ) + k (1) where T is the boiling point in degrees Kelvin, n is the number of carbon atoms in a hydrocarbon molecule and a, b and IC are empirical constants. Equation 2 was derived from evaluation of the observed boiling points of normal alkanes and subsequent curve-fitting. T 745.42 log ( n 4.4) - 416.31 (2) By application of equation 1 to the observed boiling points of normal perfluoroalkanes, and by curve-fitting according to the method of least squares, the coiistsiits a, b and k were evaluated to give the proposed equation 3 T = 540.87 log (n + 3) - 183.67 (3) Boiling points mere calculated for normal perfluorocarbons from CPB to C16F34using equation 3. Calculated boiling points are compared with observed boiling points in Table I. Since an appreciable variation exists in boiling points reported for the same compound by different investigators, it is not possible to effectively assess the accuracy of equation 3 at this time, and the accuracy of this
+
(1) G . Egloff, J. Sherman a n d R. B. Dull, THISJOURNAL, 44, 730 i1940). (2) J . D. Calfee a n d L. A. Bigelow, J . Am. Chem. Sac., 69, 2072 (1937). (3) 0. Ruff a n d 0. Bretsclineider, Z . anorg. allgem. Chem., 210, 173 (1933). (4) R. N. Haszeldine, J . Chem. Sac., 3559 (1953). (5) P. Lebeau a n d A. Damiens, Compt. rend., 191, 939 (1930). (6) R . D. Fowler, J. M. Hamilton, et al., I n d . Eng. Chem., 39, 375 (1947). (7) .I. H. Siinons a n d L. P. Block, J . Am. Chem. Sac., 59, 1407 (1937). ( 8 ) L. L. Burger and G. H. Cady, ibid., 73, 4243 11951). 19) M. Hauptschein a n d A. V. Grosse, % b i d . ,74, 4454 (1952). (10) R. N. Haszeldine a n d E. G. Walaschewski, J . Chem. Soc., 3607 11953). (11) A. F. Benning and J. D. Park, U. S.P a t e n t 2,490,764 (1949). (12) J. H. Simons and L. P. Block, J . Am. Chem. Soc., 61, 2962 (1939). (13) M. Stacey, R o y . T n s f . Cliem. Gt. B i t . and rreland, 1 (1948). (14) C.D. Oliver, e. Elurnkin and C . W. Cunningham, J . Am. Chem. Soc., 73, 5722 (1951). (15) R. N. Haszeldine, J . Chem. Sac., 3617 (1950). (16) W. B. J3urford, R . D. F o m l ~ r el , al., Ind. Eng. Chem., 39, 319 (1917).
equation can be said to be in the order of +3.5". However, good agreement is noted to exist between the calculated boiling point and a t least one of the observed boiling points in most of the cases studied. The chief reason for the variation in the observed boiling points as reported for a given fluorocarbon probably can be attributed t o isomerization which might have occurred during the fluoriliation process to give branched compounds which were not detectable by analytical means available during the early periods of fluorocarbon research. The 3 degree anomaly for the first member of the series was not unexpected, since Egloff observed an 18 degree difference between the observed and calculated boiling point of methane. The boiling points for n-CI2F26,n,-C14F30, and nC16F32 which have not been reported at this time are predicted to be 452.4, 481.8 and 495.3"R., respectively. DIPOLE MOMENTS OF SOME AMINE EXTRACTANTS I N BENZENE BY W. J. MCDOWELL A N D KENNETH A. ALLEN Oak Ridge National Laboratory, Oak Ridge, Tennessee' Received A u g u s t 29, 1958
Previous studies of the long chain amines in liquid-liquid extraction systems have shown that the sulfate and bisulfate salts of these extractaiits do not behave as would be expected for monomeric solutes.2 Recent light, scattering measurements3 have shown, indeed, that the sulfates of most of the amines investigated are aggregated, but to widely varying degrees, depending on chain branching and class. In the present investigation, the dielectric behavior of these solutes was examined as a further means of elucidating their structures. Experimenta1 Capacitance messurements were made in the range 15 Kc. (1) Operated for t h e U. S. Atomic Energy Commission b y Union Carbide Nuclear Company. (2) I
