ISOTOPIC EXCHANGE BETWEEN POTASSIUM IODIDEAND BENZYLIODIDES
Aug. 20, 1959
Hafnium tetrafluoride did not exhibit any phase changes within the range of temperatures studied. Equipment limitations prevented extending the heat-content measurements to a temperature sufficiently high for determining the melting point and the heat of fusion of this compound. Rubidium fluoride had one phase change, that of melting, within the temperature range studied. The results for this compound, in both solid and liquid states, appear normal in comparison with other alkali halides. The melting point was found to be 1048'K. and the heat of fusion 5490 cal. per mole.
4173
The heat-content equations presented below were derived by the method of least squares, using the experimental data from Table I. The average deviations are indicated by the figures in parentheses. The heat-content equation of the solid R b F was derived to fit conditions : C, = b
HT - Hzia,l, = u + bT + cT' + dT-' (1) + 2cT - d F 2 where C, = 12.04 a t 283.15OK.5 (2) 0 = a + 6T + cTZ + dT' a t 273.15'K. (3)
The heat-content equation of HfF4 was made t o fit only conditions 1 and 3, since no heat-capacity TABLE I1 data were available a t low temperatures. HEATCOSTENTS( C A L . MOLE-') AXD ENTROPY IXCREMENTS Least square solutions were obtained by means ( C A L . DECREE-''MOLE-')ABOVE 273.5"K. of a Univac-BO electronic computer.6 Entropy RbF---HfFa--increments were calculated by means of the HT ST HT Sr T, H273.15 Sria.is H273.11 Szis.ia method suggested by Kelley.' The heat-content 1.86 305 1.06 675 298.15 data are represented by the equations
---
7--
OK.
400 450 500 550 600 650 TOO 750 800 850 900 950 1000 1048 1048 1050 1100 1150 1200
1544 2175 2823 3489 4175 4881 5609 6356 7148 7912 8723 9553 10408 11252(S) 16742(L)
4.65 6.14 7.50 8.77 9.96 11.06 12.14 13.17 14.19 15.17 16.10 17.00 17.88 18.70 23.94
...
...
17668 18613 19554
24.80 25.63 26.43
3276 4621 6023 7344 8832 10344 11886 13431 14990 16565 18147 19717 21321 .,. . , .
22912 24529 ...
..
9.33 12.49 15.44 17.96 20.54 22.92 25.20 27.33 29.34 31.35 33.15 34.84 36.48
... ... 38.03 39.54
...
...
Table I1 presents smooth heat-content and entropy data for the two substances a t even 50' intervals in the 400 to 1200°K. range.
[CONTRIBUTION
FROM THE
DEPARTMENT OF
R b F (solid) HT - H m . 1 ; 1.206 X R b F (liquid)
=
-2078
lo5
+ i.96GT + 4.605 X
10-3T2 -
T-' (273.15 to 1048°K.; +0.5'/l-)
+
HT -
H2ia.16 = 61518 - 1.13!13 X 1O'T 4.416 X 10-4Tz - 3.50677 X lo7 2"-' (1048 to 120OoK.; 0.5' :I) HfF4 HT - Hm.16 = -11884.9 3.1904 X 10' T 3.739 X 10-4T2 8.9889 X 105 T-' (273.15 to 1103.3°K.; i0. 7 $ )I
*
+
+
+
The data show that average deviations of the experimental heat contents from the calculated values are greater a t the lower than at the higher temperatures. f.5) J. W. Br@nsted,Z . Elektvochem., 20, 554 (19141. (6) R . W. Smith, Jr., Supervisory Mathematician, Region V. Bureau of Mines, U. S. Department of the Interior, Pittsburgh, Penna. ( 7 ) K. K . Kelley, Bulletin 476, Bureau of hlines, U. S. Department o f the Interior, Berkeley, California.
UNIVERSITY. ALABAMA
CHEMISTRY O F THE
UNIVERSITY OF
SEIV
Isotopic Exchange between Potassium Iodide and Benzyl Iodides.
MEXICO]
Solvent Effects1
BY J . A. LEARYAND MILTONKAHN RECEIVED MARCH 12, 1959 The rates of exchange between benzyl iodide and potassium iodide and between p-nitrobenzyl iodide and potassium iodide, in acetone, are represented by the rate laws: R = 9.5 X 1010e-13~720/RT(KI)(B~I) and R = 2 X 1 0 1 2 e - 1 4 ~ 3 ~ 0 ~ R T ( K I ) ( ~ - N O ~ BzI), respectively, where the units of R are mole liter-' set.-'. The rates of exchange, a t no, between beniyl iodide and potassium iodide in acetone-ethanol, acetone-water, acetone-phenol and acetone-carbon tetrachloride mixtures and between #-nitrobenzyl iodide and potassium iodide in acetone-ethanol mixtures have been studied. The inhibitory effect of the hydroxylic compounds has been correlated with their tendency to solvate the iodide ion through hydrogen-bond formation.
Introduction was found to proceed via an S S mechanism ~ with an The isotopic exchange between benzyl iodide easily measurable rate a t room temperature.2 and potassium iodide, in methanol and ethanol, However, under the same conditions, the reaction appeared to be extremely rapid in absolute ace(1) T h i s paper is a portion of t h e dissertation presented by J. A. (3) tone. Leary in partial fufillment of t h e requirements f o r t h e degree of Doctor of Philosophy in t h e Graduate School of t h e University of New Mexico. June, 1956. Presented before t h e Physical and Inorganic Division of the American Chemical Society in Atlantic City, September, 1956.
(2) P . Stillson and M. K a h n , THISJ O U R N A L , 76, 3579 (1953). ( 3 ) P . Stillson, Doctoral Thesis, T h e University of New Mexico, 1951.
4174
J. A. LEARYAND
This paper deals with the role of the solvent in the aforementioned reaction. Isotopic exchange between benzyl iodide and potassium iodide and between p-nitrobenzyl iodide and potassium iodide in acetone and acetone-ethanol mixtures was investigated. Also, the effect of water, phenol and carbon tetrachloride on the rate of exchange between benzyl iodide and potassium iodide in acctone was studied. The variation of reaction rate with composition of solvent has been correlated with the tendency for the hydroxylic solvent t o solvate the iodide ion through hydrogen-bond formation. Experimental
Vol. s1
MILTON KAHN
separated reactants were prepared by tlic inctliod of Iethanol>>acetone.lo Also, the position 14.00' 0 .... .... 0.99 of ethanol with respect to water is in accord with 0.0610 .70 13.92 0.0661 0.2557 the intermolecular hydrogen bond energies of 13.86 ,1136 .2046 ,0976 .60 G.2 and 4.5 kcal./mole, respectively." Assuming, 13.60 ,336 ,3846 .1384 ,353 then, that the organic iodide exchanges a t a much .188 13.20 ,672 ,2563 ,2768 lower rate with hydroxylically solvated iodide ion 12.38 1.315 ,529 ,2412 ,071 than with purely acetonated iodide ion and that all See Table I. species of solvated organic iodide react rapidly, the observed initial rapid decrease in the specific TABLE V reaction rate k with increase in hydroxylic comRESCLTSOF BEXZYLIODIDE EXCHANGE EXPERIMESTS AT pound concentration can be attributed to the 0.0' I N A4CETONE-cARBON TETRACHLORIDE SOLVENT MIX- reversible reaction
TABLE IV EXPERIMENTS AT RESULTSOF BENZYLIODIDEEXCHANGE 0.0" IN ACETONE-PHENOL SOLVENT MIXTURES
TURES
(C3H60),
(CClr),
14.00" 13.29 12,GO 11.20 a See Table
0
.LI
JI
0,531 1,063 2.120 I,
(BzI), 1 'I
(y,
....
....
0.2532 3730 1688
0.2692 ,2580 ,2592
k , I. mole -1
sec. - 1
0.99 1.01 1.02 0.93
respectively. Table VI contains the results of an investigation of the p-nitrobenzyl iodide exchange reaction in acetone-ethanol mixtures a t 0". Because the benzyl iodide exchange reaction was found ~ reaction in acetone, to proceed via an S N type methanol, ethanol2 and acetone-ethanol3 mixtures, it is presumed that the exchange reactions in the various solvent mixtures reported here are also first order with respect to each reactant.
I-.An
4- W Z H=
H m .I-.An-m
+ WZA
(5)
where I-.A,, is an iodide ion solvated by n molecules of acetone and H,,.I--A,-, represents an iodide ion solvated by m molecules of hydroxylic compound and n-m molecules of acetone. At sufficiently low concentrations ( 7 -1.4M) of hydroxylic compound, m is presumed to have only values of 0 and 1. The equilibrium constant K is defined for equation 5 as
where the parentheses refer to molar concentrations. It is assumed that over the aforementioned concentration range of hydroxylic compound the (10) P . D . Bartlett and H. J. Dauben, Jr., T H I S J O U R N A L , 62, 1339 (1940). (11) L. Pauling, "The N a t u r e of the Chemical Bond," Cornell IJniversity Press, Ithaca, New York, 1948, p . 333.
T. H. SIDDALL, I11
4176
ratio of the activity coefficients of all the species is constant. I t follows that
1 2 r----
Carbon Tetrachloride
1
I
where 2(1-) = (I-.-I,,)
+ (H.I-.d,,-,)
= (KI)
In line with the foregoing comments, the rate of exchange R for the benzyi iodide exchange reaction is now expressed as I
r d , E ~ r g Cht,irr., 61, 4 1 (l!l.j!jj,
A S for the extraction reaction of uranium with the compounds that were tested. I t was felt that values for AI1 and A S might throw some light in a fundamental way on the effect of altering the substituents around the phosphorus atom. Experimental The thermodynamic quantities were calculated from the variation in tlie extraction coefficient for uranium over the temperature range, 0 to 5 0 " . I n two caws the range was extended t o >80°. T h e standard conditions chosen were 1.OO 11f nitric acid in the aqueous phase and 0.050 31 extractant diluted with n-dodecane in the organic phase. These conditions gave extraction coefficients of conveniently measurable qize. From the work of h.lcKay m d co-~orkeriSand from work it] ( 5 ) H . A . C. RlcKay and 1'.V . Henlg, R r r . trnz'. , - / r i m , , 76, 730 I l!rali)
