THE RHODIUM-CHLORINE SYSTEM AT HIGH TEMPERATURE' BY WAYNEE. BELL,M. TAGAMI, AND ULRICH MERTEN John J a y Hopkins Labolatory for Pure and Applied Science, General Atomic Division of General Dynamics Corporation Sun Diego, California Received September 2.5, 1961
The rhodium-chlorine system has been studied over the temperature range 700 to 1500" and over the chlorine-pressure range 0.01 to 1.0 atm. Results show that solid RhCL is the o:ly stable condensed chloride under the conditions studied. The dissociation pressure of the chloride reaches 1 atm. at 970 . Chlorine-pressure-dependence data indicate that RhClz and RhC13are the important gaseous Epecies. At 1504' and 1atm. chlorine pressure, I)RhCla = 6.5 X atm. and I)RhCla = 15 x 10-3 atm. Below 970°, where RhCL(s) is the condensed phase, the vapor pressures fall off rapidly with decreasing temperature, and a t SOZ", PRhCiz = 4 x 10+ atm. and 1)RhCIa = 1.7 x 10-6 atm.
Introduction The rhodium-chlorine system has been studied in the temperature range 700 to 1500° and the chlorine-pressure range 0.01 to 1.0 atm. This work was undertaken to identify condensed phases and vapor species, obtain dissociation pressure and vapor pressure data, and calculate thermodynamic values for the various reactions and species involved. I n previous investigations of the rhodiumchIorine system a t high temperature, Wohler and Muller2measured dissociation pressures and studied condensed chloride phases, and Puche3 measured dissociation pressures. Experimental Dissociatian Pressure Studies.-Dissociation pressures were measured by both the static and the transpiration methods as described in detail in a previous paper.4 I n the transpiration method, gas flow rates were around 2 ml. STP/min., which is in the range where measured pressures mere found to be independent of flow rate. Vapor Pressure Studies.-Vapor pressures were determined by the transpiration method. Except for the use of a radioactive-tracer method of analysis, techniques were essentially the same as those used in previous work.6 Mullite reaction tubes were used. Chlorine served as the carrier gas and was collected in KI solution and determined volumetrically. The vapor was collected in a condensing region a t the edge of the furnace. Slight chlorine corrosion of the mullite occurred at the highest temperatures used. At 1500°, the mullite radiation shield (which weighed about 2.5 g.) lost about 10 mg. during an experiment, whereas at 1000° the loss in weight was negligible. Flow rates ranged from 0.01 to 0.1 mmole Clz/min., depending on temperature and pressure conditions. The results of flow-rate studies given in Table I show that these flow rates were in the range where the mole ratio (moles Rh:mole Clz) was independent of flow rate. The rhodium metal used m-as irradiated in an elertron linear accelerator by means of bremsstrahlung emitted from a platinum converter under bombardment by 25MeV. electrons. The radionuclides 220-d Rh102 and 4.5-d Rh'O1 were produced by (7,n) and (y,2n) reactlons, respectively. Initially, the activity of the RhlO1 was about three times that of the RhlOZ; hoxever, when the countlng was done a t the end of the experimental w-ork, the Rh'O' activity was negligible. Two different metal samples mere used having activities of about 400 and 1500 c.p.m./mg., respectively, under our counting conditions. As a check on the analytical technique, each of the active metal samples iyas used in a vapor pressure measurement a t 1102'; the (1) This work was supported in part by the U. S. Atomic Energy Commission under Contract AT(04-3)-164. ( 2 ) L. Wohler and W. Muller, Z. anorg. u. allgem. Chem., 149, 125 (1925). (3) F. Puohe. Ann. Chem., 9, 233 (1938). (4) W. E . Bell, M. C. Garrison. and U. Merten, J . Phys. Chem., 64, 145 (1960). (8) W . E. Bell, G. Merten, and Af. Tagami, ibid., 66, 510 (1961).
EFFECTOF Temp., OC.
852 852 852 902 902 1301 1301 1301 1301
TABLE I FLOW RATEON MOLERATIO
System pressure, atm.
0.497 .492
.500 .496 .495 .977 .977 .978 .975
Flow rate [mmoles Rh)/ min. 1
(CL
+
0.016 .038 044 .050 .118 ,029 .092 .098 .33
Rh:Cla. mole ratio ( X 109
0.167 .19a
.I74 .82 .88 11.0 10.8 11.6 11.5
two vapor pressure values obtained were found to agree. y-Ray spectra of the samples agreed with spectra reported for Rhi02 and Rhlo1. To determine the quantity of rhodium condensed, the mullite condensing region (see Fig. 1 of ref. 5 for diagram of reaction tube) was crushed, placed in a plastic vial, and counted utilizing a NaI well-counter and a single-channel analyzer. Optimum counting conditions were obtained by counting integrally above 0.38 Mev. (The main yenergy peak of Rh'02 is 0.46 Mev.) Statistical counting errors were less than 2% standard deviation. For standards, samples of the radioactive metal were weighed out, mixed with crushed mullite, and counted in the same manner as the unknowns. Activities of the unknowns ranged from about 50 to 6000 c.p.m. The background stabilized at about 85 c.p.m. At the beginning of the present rork, attempts were made to develop a chemical method of analysis; however, no convenient means was found for dissolving the condensate. General .-The materials used were rhodium sponge (Johnson Matthey, 99.995% purity); RhC1~3Hz0(Fisher, 99.9% rhodium purity); chlorine gas (Matheson, 99.85% minimum purity); and argon gas (Liquid Carbonic, 99.9% minimum purity). The gases were dried and purified as described in an earlier paper.6 Tube furnaces and Pt, Pb-10% Rh thermocouples were used as described previously.6 Temperature uncertainties are believed to range from 1 2 ' a t 800' to +4" a t 1500O.
Results and Discussion Condensed Phase Studies.-Crystals made by reacting rhodium sponge and chlorine gas directly a t 800° for three days showed a chlorine content of 49.6870 by weight as determined gravimetrically by hydrogen reduction. Dehydration of commercial RhCI3.3H,O in a chlorine stream at 600' yielded an anhydrous product which analysis showed to contain 50.43% chlorine. Gavis and Sienko,B by prolonged chlorination a t 800°, obtained a product which contained 50.6070 chlorine. On comparing these values with the 50.830/, theoretical value for RhC13, it is apparent that the chloride formed was RhC&(s). (6) J. Gavis and M. J. Sienko, J . Am. Chem. Soc., 77, 4983 (1985).
49 1
THERHODIUM-CHLORIXE SYSTEMAT HIGHTEMPERATURE
llarch, 1962
X-Rajr powder patterns of samples of RhCL(s) formed a t 800' showed the compound to be isomorphous with the black form of RuCln(s) and the violet form of CrC13(s). The latter two compounds previously were shown to be isomorphous by Stroganov and O~chinnikov.~ 'To test for lower chlorides, a mixture was made in which rhodium and chlorine were in the atom ratio 1 to 1. This mixture was sealed in an evacuated quartz tube, annealed overnight a t 800°, quenched, and then examined by X-ray techniques. Diffraction lines only for RhCl,(s) and Rh(s) were found. As a further test for lower chlorides, the chlorine coiitent of the samples used in the dissociationpressure work (discussed below) was varied over a wide range. There was no apparent effect on the dissociation pressure data. On the basis of these results, we conclude that the stable condensed chloride is solid RhC& and that no other stable condensed chlorides exist within the temperature and chlorine-pressure ranges studied. Wohler and &fuller,z on the basis of gravimetric aiialyses of chloride fractions separated by flotation methods, claimed the existence of RhCL(s), RhCl?(s), and RhCl(s) under conditions similar to ours. However, since they did not demonstrate homogeneity of the phases separated, and in light of the pl'esent results, their conclusions regarding RhCl,(s) and RhCl(s) would appear to be in error. Dissociation Pressures.-The dissociation pressure data. measured over the range 722 to 963' by two different methods, are plotted in Fig. 1. The data show a linear relationship between log p and 1/ T within experimental error. Extrapolation of tlie line shows the dissociation pressure to be 1atm. at 970". The chlorine content of the samples used in the static method gradated from 50.4%, essentially the content of RhClj, to less than one chlorine atom per rhodium atom. Transpiration experiments at, 800 and a t 843' were continued to complete decomposition of the chloride, and in each case only one pressure plateau was observed. These results show that the degree of chloriiiation had no apparent effect on the dissociation-pressure data and suggest that RhC18 was the only solid chloride involved. From the slope of the line in Fig. 1, we calculate, a t the mean temperature of the measurements (llOO°K.) AHolloo = -67.8 f 2.0 Imd./mole
and from this AB0
ASOIIOO= -
-
T
- 732 R In p
=
-54.5 i 2.0 e.u.
for the reaction Rh(s)
+ 53 C1z = RhClz(s)
(1)
From a rough rule given by Kubaschewski and Evans,8 we estimate AC, for reaction 1 to be 4.5 (7) E. V. Stroganov and K. V. Ovchinnikov, Veslnzk Lenzngrad Unzv., 12, No. 22, Ssr. Fzz. i Khim., No. 4,152 (1957). (8) 0. Kubaschewski and E. L. Evans, "SIetallurgical Thermochemibtry," 3d ed., Pergamon Press, New Yolk, N. Y.,1958.
1.0
0.5
0.2
i
-!2 W
a 3 VI
w
0.1
P K
w
E
a
9
0 I
0.05
0.02
0.01
0.8
1.0
0.9 IPK
Fig. 1.-Dissociation
x lo3. pressure of RhC13(s).
k l . 0 cal./mole OK., and assuming it to be constant over the range 298 to llUO'K., we calculate A H o 2 S= ~ -71.5 3~ 3.0 kcal./mole A80zss =
log pc1.2 =
-60.4 f 3.0 e.u.
-A 620 T
- 1.510 log T
+ 13.20
Combining ASo298 with standard entropies Xo298 Rh = 7.53 f 0.05 e.u. and XoZ98Clz = 53.29 f 0.01 e.u. given by Kelley and King,9 we obtain RhCla(s) = 27.1 & 3.0 e.u. This value may be compared with the experimental values Sop95RuC13(s) = 30.5 i 2.5 e.u., obtained by Bell, Garrison and M e r t e ~and ~ , ~S0298CrCl,(s) = 29.4 =k 0.2 e.u., reported by Kelley and Kingg and based on low temperature heat capacity data. The value for SO298 RhCL(s) also may be compared with 38.0 * 6.0 e.u., estimated by Brewer, et aZ.,'O and with 33.2 e.u., estimated by Latimer's rules." Wohler and Muller2 reported dissociation-pressure data for RhC13,RhC12, and RhCl (as mentioned above, we find 110 evidence for the lower chlorides), (9) X. X. Xelley and E. G. Xing, C . S. Bur. Mines Bull. 592, 1961. (10) L. Brewer, et al., in "The Chemlstry and Metallurgy of Miscellaneous Materials: Thermodynamics," National Nuclear Energy Series, Div. I V , Vol. IQB, AIoGraw-Hi11 Book Co., Inc., New York, N. Y . , 1950. (11) W. Kf. Latimer, J . Am. Chem. 8 6 e . , 73, 1480 (1951).
W. E. BELL,M. TAOAMI, AND U. MERTEN
492
Vol. 66
1.0
20 0.50
2 z
a \
0
t
W J
r
W 0
3
a \.
W -l
. . I
0.20
0
5.0
W 0
3
VI
i
8 -
3
3-mg
VI
x
13
8, Pw -a
9" VI 4 VI
z:iz
w'
0.050
K 3 VI VI Y K
1.0
J
a
3 K
w
a
09
W
E
0.10
z c
2.0
- x
2
902'C
c
IO
0.50
8020C
0.020
n
K
0
0
0
g 3 0.010
0.20
Fig. 3.-Effect
/
/ 0.10
0.20
0.10
0.50
1.0
I 0.10
I
0.20
I
0.50 CHLORINE PRESSURE (ATY.),
1.0
of chlorine pressure on vapor pressure at 802, 852 and 902'.
pressure values 10.8, 10.6, 11.4, and 11.2 X low3 atm. obtained in flow rate studies (see Table I). Fig. 2.-Effect of chlorine pressure on vapor pressure at The point a t 1102' and 0.99 atm. chlorine pressure 1102, 1301 and 1504'. is the average of the values 4.98 and 4.96 X atm. obtained using metal samples of different and Puche3 reported data for RhC1, and RhC12. specific radioactivity. The uncertainty of the The data which Wohler and Muller report for RhCl individual points is about &5% a t the higher temfall fairly well on our curve in Fig. 1; however, the peratures and 1 atm. chlorine pressure and increases data which they and Puche report for RhCL and to about *15% as the temperature and chlorine RhC1, fall above our curve. pressure are lowered. The curves of Fig. 2 were Identification of Vapor Species.-Since Rh(s) and drawn to be consistent not only with the experiRhCl,(s) are the stable condensed phases in the mental points, but also with each other. temperature range and the chlorine-pressure range I n Fig. 2, the slopes of the isotherms taken a t 1 studied, the solid-vapor equilibria to be con- atm. chlorine pressure are 1.36 (1504'), 1.44sidered are (1301°), and 1.50 (1102'). Those of Fig. 3 are 1.4 (902') and zero (902,852, and 802'). Thus, y/2 zRh(s) Clz = RhxCl,(g) (2) 1.5 and (y - 3s)/2 S 0 ; y = 3 and x = 1. It therefore is apparent that the principal vapor y - 3x zRhCla(s) --2C1, = RhzCldg) species is RhCh. (3) Comparison of the experimentally determined From the equilibrium constant for eq. 2, we obtain (solid) curves and the dashed curves (drawn with a slope of 1.5) in Fig. 2 shows that the experilog p ~ i , ~= r i$~log p a 2 -k log K mental curves decrease in slope with increasing A similar expression is obtained for eq. 3, and we temperature and decreasing chlorine pressure. evaluate y/2 and (y - 3x)/2 by studying the effect The effect apparently is the result of a contribution from a lower chloride. This means that the exof chlorine pressure on vapor pressure. Figures 2 and 3 show the isotherms obtained. perimental curves are actually the sum of linear The break in the 902' isotherm a t 0.36 atm. pressure-dependence curves representing RhCla fixes the equilibrium dissociation pressure of Rh- and a t least one unknown lower chloride. By successive approximations, we were able to C&(s) and agrees with the dissociation-pressure data in Fig. 1. The points in Fig. 2 and 3 repre- resolve each of the experimental curves into two sent the results of individual experiments with the linear curves having slopes of 1.5 and 1.0, respecfollowing exceptions. The point at 1301' and 0.98 tively. From the latter value, we find that y/2 atm. chlorine pressure is the average of the vapor- S 1.0 and y = 2 ; thus, the lower chloride species CHLORINE PRESSURE (ATM.).
+ +
THERHODIUM-CHLORINE SYSTEM AT HIGH'TEMPERATURE
March, 1962
appears to be RhClz (in this case, we assume that x = 1). Table I1 gives individual partial-pressure values for RhC13and RhC12taken from the resolvedilinear curves a,t 1 atm. chlorine pressure and a t pressures corresponding to those used in the experimental work. To show how the data correlate, the sums of the individual partial pressures are compared with experimentally determined total vapor-pressure values. The agreement appears satisfactory. i
TABLE I1 COMPARI~ON OB EXPERIMENTAL AND RESOLVED VAPOR PRESSURES Resolved vapor pressures, a h . X 103 PClP
(atm.)
+
PRhCln PRhCl8
s-
-
Exptl. vapor pressures, b atrn. X 103
PRhClaa
0.100 .243 .485 ,972
1.10 3.33 8.1 20.6
1.(E4 3.56 8.3 20.3
1.oo
0.63 1.53 3.03 6.1 6.3
1504' 0.468 1.80 5.1 14.5 15.2
0,099 ,246 .490 ,98 1.00
0.238 0.59 1.17 2.35 2.40
1301' 0.275 1.08 3.08 88 9.0
0.51 1.67 4.25 11.2
0.53 1.76 4.02 11.0
..
..
1102O 0.056 0.138 0.194 0.197 0.100 0.69 ,137 0.54 0.68 .246 ,273 1.55 1.82 1.78 .492 .55 4.45 5.00 4.98 .99 .56 4.50 .. I .oo eTaken from resolved linear curves at the appropritte chlorine pressures. * Experimental vapor pressures, assuming one vapor molecule per rhodium atom condensed.
..
Although the results correlate very nicely if RhClz(g) is taken to be the lower-chloride gaseous species, we cannot entirely eliminate the possibility that the deviations from a 1.5 slope of the experimental curves of Fig. 2 may have resulted from a rhodium- bearing species formed from chlorine corrosion of the mullite. The species RhCl(g) appears to be eliminated by the fact that the curves of Fig. 2, because of their large radius of curvature, do not resolve into two linear curves having slopes of 1.5 and 0.5, respectively. Temperature Dependence of Vapor Pressures.Since RhClz and RhC1, appear to be the important vapor species under our experimental conditions, the solid- vapor equilibria to be considered are
+ Clz = RhCl,(g) 3 Rh(s) + 5 Clz = RhCls(g)
RhCla(s) = RhCl,(g)
% %!
w
x
3
%
P
y
Rh(s)
+ 51 C1,
RhCl,(s) = RhCla(g)
2.0
-
1.0
-
0.5
-
o.2
-
0.1
-
0.05
-
2.0,
v)
..
-
ed
PRhCIZa
..
5.0
493
(4) (5) (6)
(7)
When Rh(s) is the stable condensed phase, the observed vapor pressure must be the sum of contributions from reactions 4 and 5 ; and when Rhc13(s) is the stable condensed phase, the observed pressures must be the sum of contributions from
g
P
0
a
3
0.02 0.01
-
0.005
-
0.002 nnoi -.--. L 0.5
Fig. 4.--Individual
I
Q6
I 0.7
I
I\
0.e
09
1.0
IPK x io3. partial pressures of RhClz and RhClt.
reactions 6 and 7. Reaction 4 is related to (6) and reaction 5 is related to (7) by reaction 1. Figure 4 shows the individual partial pressures of RhClz and RhC4 as a function of 1/T. Curves A and C represent vapor pressures of the two species in equilibrium with Rh(s) a t 1 atm. chlorine pressure (reactions 4 and 5 ) , and curves B and D represent the pressure data for RhC13(s) a t 1 atm. chlorine pressure (reactions 6 and 7). These curves were arrived at in the following manner. The three RhCh pressure values a t 1 atm. chlorine pressure, obtained as described above and included in Table 11, were plotted and curve A was drawn through the points. Curve B was drawn to intercept curve A a t 1/T = 0.804 (970°, the dissociation temperature of the solid chloride a t 1 atm. chlorine pressure) so that the slope changes by -67.2 kcal./mole, the heat of reaction 1 a t 970'. Partial pressure values for RhC12, shown in Table 111, were taken from curves A and B at the appropriate 1/T values and were subtracted from the experimentally observed vapor pressures to obtain the partial pressure values for RhCh shown in Tab@111. These were plotted in Fig. 4 and curves C and D were drawn through the points. The curves were drawn as straight lines, since the data are not considered sufficiently accurate to warrant taking into account heat capacity effects over the short range of temperature involved. Thermodynamic values for reactions 4 through 7 are summarized in Table IV. The A H o I 2 k 3 and
LEONARD KHrzIS AXD ELSIEQULYAS
494
AS01243 values for reaction 4 were obtained from curve A. These values were combined with the AHo12d3 and ASOIZ~~ values for reaction 1 to obtain AHo12d3and ASo1243 values for reaction 6. The latter values are represented by curve B. The AH01243 and AS01243values for reactions 5 and 7 were obtained from curves C and D. The AC, values of Table IV were obtained from C, values estimated from values for similar halides given by Kelley,12 and the AH0298 and ASoaQ8values were calculated by assuming AC, t o be constant over the temperature raiige involved. TABLE 111 OBSERVED TOTALVAPOR PRESSURES RESOLVEDINTO INDIVIDUAL PARTIAL PRESSURES OF RhCh AND RhCls AT 1 ATM.CHLORINE PRESSURE Temp., OC.
Obsd. total pressures,a atm. X 108
Resolved individual pressures, atm. X l o 3 PRhCIZb
PRhCl3
802 0.0175 0.0004 0.0171 852 IO88 .003 085 902 .394 .018 .376 956 1.95 .ll 1.84 1003 3.10 .25 2.85 1102 5.00 .57 4.43 7.8 1.20 6.6 1202 1301 11.4 2.3 9.1 1401 15 8 3.9 11.9 1504 21.3 6.5 14.8 5 Experimental vapor pressures, assuming one vapor molecule per rhodium atom condensed. The 956" value is 1.96, the average of three experimental values-1.90, and 1.99 X atm. The 1003, 1202, and 1401" values are from individual experiments. A slight correction was required to bring these values to 1 atm. chlorine pressure. The remaining values mrere taken from Figs. 2 and 3 . b From the RhClz curve of Fig. 4. CColumn 2 minus column 3.
Vol. 66
TABLE IV THERMODYNAMIC DATAFOR v.4PORlZATION REACTIONS Rh(s)
+ CL = RhCln(g)
AH'uta = +29.2 f 2.0
kcal./mole = +6.4 f 2.0 e.u. AC, = - 1.2 cal./mole "K. (estimated) AH02sg = +30.3 f 3.0 kcal./mole AS043s= +8.1 f 3.0 e.u.
AS01243
RhCls(s) = RhClz(g)
+ 1/2CIz
AH'im = +96.4 f 4.0
kcal./mole = +60.5 =!= 4.0 e x . AC, = -5.7 cal./mole "K. (estimated) AHozga = +101.8 f 5.0 kcal /mole ASozss= 4-68.6 f 5.0 e.u.
.
Rh(s)
+ 3/2cIz
=
RhCls(g)
+14.9 f 1.0 kcal. /mole AS'i243 = +0.2 =k 1.0 e.u. AC, = - 1.1 cal./mole 'K. (estimated) AH'29s = +16.0 f 2.0 kcal./mole = +1.8 =k 2.0 e a .
AHOma
RhCla(s) = RhCla(g)
AHom, = f79.5 f 2.0
kcal. /mole ASnIz4~= +51.9 f 2.0 e.u. AC, = -5.6 cal./niolc OK.
(estimated) AH02m = +84.7 f 3.0
kcaI./mole AXOpss = +59.8 f 3.0e.u.
-67.2 f 2.0 kcal./mole and AX01243 = -54.0 f 2.0 e.u. obtained from the dissociation pressure measurements. Combining ASo2g8 values for reactions 4 and 5 with standard entropies Xo298Rh(s) = 7.53 f 0.05 e.n. and So298C1, = 53.29 f 0.01 e.u. given by Kelley and King19we obtain SO298 RhClz(g) = 68.9 f 4.0 e.u. and SO298 RhC13(g) = 89.3 f 4.0 e.u. These values may be compared, respectively, with the estimated values of 79.4 and 83.1 e.u. calculated from an empirical equation given by Kubaschemki and Evans.8 They also may be compared with the following experimental values for analogous gaseous species: SO298 CrClz(g) = 76.0 e.u. and 8'298 CrC13(g) = 86.1 e.u. (found by DoerneP), and SOggr RuC13(g) = 95.1 e.11. (found by Bell, Garrison, and SIerter~'~). By combining the AH0m3and ASOIZGvalues for Acknowledgments.-The authors are indebted to reaction 5 with the respective AH01243 and AX0m3 R E. Inyard for performing part of the experinienvalues for reaction 7, we obtain AH0j243 = -64.6 tal work and to J. &I. Dixon for performing the f 3 kcal./mole and ASo1243= -51.7 f 3.0 e.u. for X-ray analyses. reaction 1. These values agree, within the un(13) H A. Doerner, U. S. Bur. hlines Tech. Paper 577, 1937. certainty of the data, with the values AHOW = (14) W, E. Bell, M. C . Garrison, and C. Merten, J . Phys. Chem., 65, 517 (1961).
(12) K. K. Kelley, U. S. Bur. Mines Bull. 584, 1960.
EFFECTS OF ELECTROLYTES ON ROTATORY DISPERSION OF AQUEOUS TARTRATE SOLUTIONS' B Y LEONARD I. KATZIK$ S D
ELSIE
GULYAS
The Chemistry Dzoision, Argonne Sational Laboratory, Argonne, Illinois Recezved September 27. 1961
Rotatory dispersion data over the range 6500-2650 8., for aqueous tartaric acid systems containing HC1 and chlorides of sodium, lithium, calcium, praseodymium, or thorium, and for alkaline tartrate, are fitted to a two-term Drude equation. The wave length parameters obtained are found t o be in fair agreement with the absorption spectra. The data are compatible with the hypothesis of direct interaction between the cations and the OH groups of the un-ionized tartaric acid.
One of the facts that continually has to be kept in mind in working with solutions of salts in water is that the ions formed interact with the water. (1) Based on work performed under the auspices of the U. S. Atomic Eneray Commission, Presented in part a t Smnrioan Chemical Society National Meeting, Chicago, September 3-8, 1961.
The type of interaction of most significance with regard to magnitude of energy effects, and influence on chemical behavior, is that of cations with the non-bonding electrons of the water OH groups. Such interaction may give a "ordinate bond Of varying degrees of strength. This mode is, in
