1202
C. w u ,M. nI. BIRKY,AKD L. G. HEPLER
needs to be large enough to interact with one acetylenic hydrogen atom. I n the self-hydrogenation of ethylene to form chemisorbed ethyl groups, the following is known about the nature of the adsorption sites: (a) The ethyl group is probably held “end-on” to the surface by a carbon atom. This is confirmed by the fact that the ratio of the optical densities of the bands due to CH3and CH2is similar to that in molecules such as ethyl bromide. Since ethylene is much less acidic than acetylene and can behave as a weak base, ie., a proton acceptor, the sites on the alumina which are holding the ethyl groups can be said to have some acidic character . (b) The sites active in the formation of -C2H6 groups appear to be remote from the surface OH groups, as only -CzDs groups were found on adding CzDd to a surface containing only OH groups. This also implies that the adsorbed species are not mobile, as discussed in (d) below. This is particularly interesting since the self-hydrogenation of ethylene in these surfaces must be a t least a bimolecular process in which extensive bond formation and bond breakage occur. (c) In agreement with the above, removal of all the OH groups from an alumina by evacuation a t 1000° did not change the species formed after admitting ethylene. (d) There must be a t least two pairs of adjacent sites available, to act simultaneously, for the self-hydrogenation of ethylene to take place. Conceptually, we can envisage two ethylene molecules adsorbing side by side. Both open up their double bonds, and one donates hydrogen to the other. This would leave a stable -CZH6 group held strongly on one side to the surface, and a residue of nominal composition -C2H3. Possibly, a second ethylene molecule would then adsorb adjacent to the (presumably) reactive -C2H3group, giving -C2H6 and a residual group of low hydrogen to carbon ratio. If the latter process takes place, one has to invoke three adjacent pairs of sites, if all the adsorbed species are immobile. Little direct information is available on the last point, except that it seems very likely that any reactive hydrogen-containing mobile adsorbed species would exchange with surface OD groups. No exchange
Vol. 67
has been observed between CzH4 and OD groups, and between C2D4 and OH groups in any experiment. Finally, we are led to the conclusion that the sites for ethylene self-hydrogenation seem to be slightly affected by pre-adsorbed acetylene, and, conversely, that the end-on adsorption of acetylene is unaffected (to within & 5%) by the pre-adsorption of ethyl groups. In view of the fact that aoetylene exchanges with one type of OD group, and ethylene does not exchange at all with OD groups, even while undergoing self-hydrogenation, it is to be expected that the sites are independent, and displaced physically from each other on the surface. An even more striking example of the physical separation between the sites is provided by the experiment using pre-adsorbed deuterioacetylene. Even though subsequently added ethylene formed -CzH6, no evidence was found for the existence of -CZD6 or -C2H4D species in the adsorbed phase formed from C2H4. This evidence suggests that all of the adsorbed species are immobile, and that the end-on acetylene is adsorbed near to one type of OH group. While some donation of hydrogen must take place during the formation of ethyl groups by self-hydrogenation, neither of the (at least two) molecules involved in this rearrangement, is in a reactive condition while close to an OH group. Relatively large patches of the surface must be involved in self-hydrogenation of ethylene, while if an acetylene molecule impinges on its potential site with the correct orientation, the formation of end-on held acetylene can take place immediately. Both these ideas are consistent with the fact that acetylene adsorption is very fast, in most cases within 10 min., while the formation of -CzH6 is a process usually taking over 12 hr. to form significant amounts of adsorbed material. Also of interest is the fact that on a given sample there seem to be a definite number of sites which rapidly form end-on acetylene. Increasing the pressure of the acetylene, and waiting for several days gave no increase in the number of end-on species. With ethylene, in marked contrast, the rate of formation of ethyl groups is slow, and a t constant CZH4 pressure the alumina continues to form -C&& groups for the longest times which have been used (8 days).
THERMOCHEMISTRY OF SOATE BROMINE AND IODINE SPECIES I N AQUEOUS SOLUTION BY CHING-HSIEN Wu, MERRITTM. BIRKY,AND LORENG. HEPLER Departments of Chemistry, University of Virginia, Charloffesville,Virginia, and Carnegie Institute of Technology, Pittsburgh, Pennsylvania Received November 6, 1968 Calorimetric determinations of the following have been made: (a) AH of solution of Bre(1iq) in water, ( b ) AH of reaction of Rrz(1iq) with OH-(aq), ( c ) AH of reaction of Rrz(1iq) with excess I-(aq), ( d ) AH of solution of Iz(c) in excess I-(as), and (e) AH of reaction of KIOs(c) with excess I-taq). Results of these experiments have been used with data from the literature for calculation of thermodynamic properties of Brttaq), Idas), OBr-(aq), HOBr(aq), I08-(aq), and KIOdc).
Introduction Most of the enthalpy data for bromine and iodine species in aqueous solution are based on calorimetric experiments done more than fifty years ago, or on temperature variations of equilibrium constants. Results of
different investigqtors are in poor agreement for several species. We have, therefore, undertaken calorimetric measuremests of the following: (a> AH of solution of Brdliq) in water, (b) AH of reaction of Brdliq) with OH-(aq), (c) AH of reaction of Br2(liq) with excess
THERMOCHEMISTRY OF BROMINE AND IODINE SPECIESIN AQUEOUS SOLUTION
June, 1963
1203
I-(aq), (d) A H of solution of Iz(c)in excess I-(aq), and (e) AH of reaction of KI03(c) with excess I-(aq). Results of these measurements have been used with data from the literature for thermochemiaal calculations. Experimental
with temperature according to the integrated van’t Hoff equation
Most of the measurements reported here were made with a calorimeter similar to that described by O’Hara, Wu, and Hep1er.l The calorimeter used for some of our measurements differs from that described earlier in that the calorimeter heater and thermometer were enclosed in a glass coil rather than between concentric silver and copper cylinders as in the earlier calorimeter. A few of the measurements reported here were made in another calorimeter that has been de~cribed.~?3 All calorimetric measurements were made a t 25.0 i 0.3” with 950 ml. of water or solution in the calorimeter. Bromine was handled as previously de~ c r i b e d . ~None of the reactions took longer than five minutes to reach completion. Reagent grade bromine from Baker and Adamson and also from Mallinckrodt (containing less than 0.3y0 Clz aE principal impurity) was used without further purification. Reagent grade iodine was sublimed twice and stored in a desiccator over phosphorus pentoxide. Analytical grade KIOa was recrystallized and then dried at 120’. Other chemicals were of C.P. grade and were used without further purification except for drying. Stock solutions were made and standardized by common methods.
where m, and ys represent the molality and the activity coefficient of bromine in saturated solution. Substituting (1) in (2)gives
Results and Calculations Results of measurements of the heat of solution of liquid bromine in 0.1 M perchloric acid are given in Table I. These measurements were carried out in dilute acid to prevent hydrolysis of BrPto HOBr and Br-. The standard heat of formation of Brz(aq) given in NBS Circular 5005corresponds to a value of 1.1kcal./mole for AHo of solution, based upon old work of BertheKot, Pickering, and Thomsen. TABLE I HEATSOF SOLUTION OF BROMINE IN DILUTE AQUEOUS ACID Moles BrP/ 950 ml. soln.
AH (kcal./mole)
0.00601 .03185 .03,519 ,04619
-0.15 - .19 - .23 - .15 - .27
.04840
AH (av.) = - 0.20 kcal./mole
Av. dev. = 0.04 kcal./mole
Because there is no significant dependence of AH of solution of Brs(liq) on the concentration of the final solution, we take the standard heat of solution of Brz(liq) to be the average of the values reported in Table I with estimated maximum uncertainty of hO.08 kcal./ mole. Therefore, AHfO = -0.20 kcal./mole for Brz(as> It is also possible to estimate AHo for solution of BrZ(1iq) from solubility data for bromine in! water.6 Following earlier calculation^^^^ we have taken the activity coefficient y to depend on concentration na as in
-
In y = km (1) The activity of bromine in saturated solution varies (1) W. F. O’Hara, C. H. Wu, and L. G. Hepller, J . Chem. Educ., 38, 512 (1961). (2) R . L. Graham and L. G. Hepler, J . Am. Chem. rSoc., 78, 4846 (1956). (3) C. N. Muldrow and L. G. Hepler, ibid., 79, 4045 (1957). (4) L. G. Hepler, J. S. Sweet, and R. A. Jesser, ibid., 82, 304 (1960). (5) “Selected Values of Chemical Thermodynamic Propertieti,” Circular 500, National Bureau of Standards (1952). (6) A. Seidell, “Solubilities of Inorganic and M e h i Organic Compounds,” Vol. 1, D. Van Nostrand, New York, N. Y., 1940. (7) M. A. Paul, J . Am. Chem. &e.. 78, 2513 (1953). (8) L. P. Fernander and L. G. Hepler, J . Phys. Chem., 63, 110 (1959).
ln mays =
AHo -+C
RT
We also write the general equation 3 with ma* and T* and designate thizi specific equation 3’. Then subtraction of (3’) from (3) and rearrangement gives m,
T - T* (m, - ms*)(TT*)
l--
k 2.303
(4) Equation 4 has been applied by letting m,* be the solubility at 273’ andl setting T* = 273’. The slope and intercept of a graph of (log ms/ms*)/(m,- ma*) against (T - T * ) / ( m , -. nzS*)(TT*)lead to AHo = -0.17 kcal./mole and k = -3.59. The concentration of the saturated solution at 25’ is 0.219 molal, so ys = 0.46. Rearrangement of equation 3 gives AHo
log ma -k 2.303RT
=
C
- -km,
2.303
+
2.303
(5)
A graph of the left side of (5), taking AHa = -200 cal./ mole, against m,has slope 1.SO, leading to IC = -3.46 and ys = 0.42. Use of equation 1 with the same k a t different temperatures implies that AH of dilution of aqueous bromine is small. Interpretation of the solubility of bromine in water in terms of the equilibrium BrZ(1iq)
=
Brf(aq)
implies that the solubility of mater in bromine is negligible. Taking ys and ms to be 0.44 and 0.219, we calculate for the standard free energy of solution and formation of Brz(aq) in water at 298’ AFO
=
-RT In mays= 1.38 kcal./mole
Combining AFo with AHo gives ASo = -5.3 cal./deg. mole for solution of bromine in water. Since the standard entropy of Brz(liq) is 36.38 cal./deg. we calculate that the standard partial molal entropy of Brr (as) is 31.1 cal./deg. mole. We have measured the heat of solution of Br2(liq) in aqueous sodium hydroxide (0.25-1.3 M) where the principal reaction is Brz(1iq)
+ 20H-(aq)
=
Br-(aq)
+ OBr-(aq) + H20
(6)
Data are given in Table 11. The average A H reported in Table I1 cannot be attributed entirely to reaction 6 because of the side reaction (9) D. L. Hildebrand, W. R. Kramer, R. A. McDonald, and D. R . Stull. J . Am. Chem. Soc., 80, 4129 (1958).
C. Wu, M. M. BIRKY,AND L. G. HEPLER
1204 Brz(liq)
+ 20H-(aq)
Vol. 87
gives
=
Brz(aq)
+ 20H-(aq)
=
Br-(aq) TABLE I1 HEATSOF SOLUTION OF Brz(liq) I N AQUEOUS XaOH AIoles Brz/ 950 ml. s o h .
(koal./mole Brz)
0.00758 .00940 .01021 .01069 .01131 ,01480 .01483 .02341 AH (av.) = -9.92 kcal./mole
-9.81 -9.97 -9.85 -9.90 -10.02 -9.97 -9.98 -9.83 Av. dev. = 0.07 kcal./mole
AH
Using S B S 6 heats of formation of OH-(aq), Br-(aq). and H20 and 18.3kcal./mole for AHfOof bromate ion,1° we calculate AHo = - 12.7 kcal./niole of Bra for reaction 7. Representing the measured heat given in Table I1 by AHmJwe write
-
AH, = (1 - f)AH, - 12.7f
(8) where f designates the fraction of the Brz that reacts according to equation 7. Following McDonald and Cobble,ll we find that f is small and approximately 0.01 for all of our experiments. Taking A H , = -9.92 kcal./ mole and f E 0.01, we calculate from (8) that AH6 = = -9.89 kcal./mole. Since reaction 6 involves equal numbers of ions in dilute solution with -1 charges among reactants and products, we assume that heats of dilution cancel and take AHGo = -9.9 f 0.4 kcal./mole. Combining this value of AH6' with XBS6 data for the species in reaction 6 leads to AHrO = -22.6 kcal./mole for the standard heat of formation of OBr-(aq). McDonald and Cobblell found -23.0 kcal./mole for this heat of formation. Kelley and TartarlZhave reported AFgO= 11.7 kcal./ mole, AHgO= 6.8 kcal./mole, and AXgo = -16.6 cal./ deg. mole for the ionization of hypobromous acid a t 298'K.
HOBr(aq)
=
H+(aq)
+ OBr-(aq)
(9) It is possible that considerable uncertainty should be attached to these values because the data obtained by Kelley and Tartar also lead to AC: = -215 cal./deg. mole for ionization of aqueous hypobromous acid, as compared to a usual value of about -40 cal./deg. mole, for such reactions. Liebhafsky13has measured K a t several temperatures for the reaction Brdaq)
+ H20 =
+
+
HOBr(aq) H+(aq) Br-(aq) (10) His equation for log K as a function of temperature leads to AFloo = 11.23 kcal./mole, AHloO = 13.1 kcal./mole, and Axloo = 6.3 cal./deg. mole a t 298OK. Combining reactions 9 and 10 with 2H+(aq)
+ 20H-(aq)
= 2H20
(11)
(10) H. C. Mel, W. L. Jolly, and W. M. Latimer, J . Am. Chem. Soc., 7 6 , 3827 (11) J. E. McDonald and J. W. Cobble, J . Phtm. Chem., 66, 2014 (1961). (12) C. M. Kelley a n d H.V. Tartar, J . Am. Chem. Soc.. 78, 5782 (1956). (13) H. A. Liebhafsky, zbid., 56. 1500 (1934).
(1953).
+ OBr-(aq) + H2O
(12)
Similar combination of the thermodynamic data above for reactions 9, 10, and 11 gives AFrzO = -15.3 kca1.l mole, AHlzo= -8.8 kcal./mole, and ASlzo = 28.2 cal./ deg. mole. Combining our calorimetric AHeoand AHo of solution of liquid bromine in water gives AHlzo= -9.7 kcal./mole. Since all errors in our calorimetric results are less than one kcal./mole and our AH6O differs by less than 0.5 kcal./mole from that of McDonald and Cobble,ll we conclude that the calorimetric AHlz' = -9.7 kcal./mole is the more reliable value. Since we have no calorimetric value for AHQ,we use the value given by Kelley and TartarI2 with the above heat of formation of OBr-(aq) to calculate AH# = -29.4 kcal./mole for HOBr(aq). Combining AF1Z0 = - 15.3 kcal./mole with our data for Brz(aq) with SBS5data for the other species in reaction 12 leads to AFfo = -7.8 kcal./mole for the standard free energy of formation of OBr-(aq). Csing this free energy with the free energy of ionization of HOBr(aq) given by Kelley and T a r t a P leads to AFfO = - 19.5 kcal./mole for HOBr(aq). For further caleulations we obtain ASlZo = 18.8 cal./ deg. mole by combining AF120 = - 15.3 kcal./mole with our calorimetric AHlz0 = -9.7 kcal./mole. Using this AS12O with our entropy of Br2(aq) and NBS6 entropies of the other species in reaction 12 leads to 8.8 cal./deg. mole for the standard partial molal entropy of OBr-(aq). Since we have no calorimetric value for AHo of ionization of HOBr(aq), we use the results of Kelley and TartarlZ with values reported above for OBr-(aq) to calculate 25.4 cal./deg. mole for the standard partial molal entropy of HOBr(aq). We have determined the heat of reaction of bromine with excess aqueous iodide. Most of the iodine in the final solutions mas in the form of 13-(aq) so the principal reaction was Brz(1iq)
+ 31-(aq)
=
I3-(aq)
+ 2Br-(aq)
(13)
To obtain AH13' from measured heats, corrections must be made for the small amounts of Iz(aq) present in the final solutions. Davies and Gwynne14determined K a t several temperatures, leading to AHo = -3.8 kcal./mole for Idaq)
+ I-(aq)
=
L-(aq)
(14)
Our calorimetric results are summarized in Table I11 where the actual measured heats are listed under Q (all heats mere exothermic). Corrections for reaction 14 are listed under Q14 and are almost negligible. Thus, the particular value of AH14O is not important for evaluation of AH12. Since reaction 13 involves the same number of equally charged ions on each side and all solutions were quite dilute, we assume that heats of dilution cancel and take AH13' to be the average of the values listed in Table 111. Combining AH130 = -29.36 f 0.3 kcal./molewith NBS6 data for Br-(aq) and I-(aq) gives AHrO = - 11.7 kcal./ mole for 18-(aq). (14) il1. Davies and E. Gwynne, ibid., 74, 2748
(1952).
THERMOCHEMISTRY OF BROMINE AND IODINE SPECIESIN AQUEOUS SOLUTION
June, 1963
TABLEI11 HEATSOF REACTION OF Br2(liq)WIT" AQUEOUS KI 950 ml. soln.
hIoles KI/ 950 ml. Boln.
H C1
0.7423 1.2022 1.4197 1.4970
0.2302 .3778 ,3464 .1832
G. Brz(liq)/
AH13
(1W
6, (oal.)
(cal.)
(ked./ mole)
0.0239 .0191 .0239 .0095
135.76 221.11 259.79 275.93
0.10 .10 .13 .27
-29.26 -29.41 -29.26 -29.49
Qir
We also have measured the heat of solution of Iz(c) in excess I-(as) where the principal reaction is L(c)
+ I-(aq)
Is-(aq)
(15) Corrections similar to those discussed above for presence of Iz(aq)are summarized with our data in Table IV. =
TABLEIV REACTION OF L(c) WITH AQUEOUS IODIDEION G. Iz(c)/ 950 ml. soln.
Moles KI/ QBO ml. soln.
HC1
(W
Q (cal.)
(cal.)
(ked./ mole)
0.9896 1.2924 0.6438 0.8466 0.4958
0.1777 .0917 ,1149 .1603 ,1521
0.0192 .0239 ,0239 .0239 .0965
-5.68 -6.67 -3.44 -4.44 -2.83
0.11 .29 .11 .10 .07
1.43 I .24 1.31 1.30 1.41
AHin &I4
Assuming that heats of dilution cancel for the charged species in dilute solution on each side of reaction 15, we take AH1bO = 1.34 =t0.2 kcal./mole. Combining our AH16' with AHfOof I-(as) from the NBS6 gives A H f o = - 12.0 kcal./mole for 13-(aq). We take AHfo= - 11.9 kcal./mole of Is-(aq) to be the best value based on all of our data. Stern and Pa~schier'~ have measured the heat of solution of Iz(c) in 2 M K I and reported AH = 0.89 kcal./ mole. On the basis of this AH and estimated heat of dilution corrections they calculated AHfO = - 12.7 kcal./mole for Is-(aq). Since reaction 15 involves the same number of ionic charges on each side, me have assumed that heats of dilution cancel and have calculated AHfo= -12.5 kcal./mole for 13-(aq) from their AH. Because heats of dilution for our dilute solutions introduce less uncertainty than for the concentrated iodide solutions used by Stern and Passchier,15 we believe our standard AHfo= -11.9 kcal./mole for 13-(aq) js more reliable, From the solubility of iodine6,I6in water a t :So, we calculate the standard free energy of solution of Iz(c) to be 3.93 kcal./mole. Therefore, AFfO = 3.93 kcal./mole for Iz(aq). From the temperature coefficient of the solubility we calculate AH16O = 5.2 kcal./mole and AXlao = 4.4 cal./deg. mole for Idc) = Iz(aq) (16) Combining this A816' with NBS6 entropy of 12(c) leads to 32.3 cal./deg. mole for the standard partial molal entropy of Iz(aq). Combining our calorimetric AH15' = 1.3 kcal./mole with the above AH16' gives AHI~O= --3.9 kcal./mole, in (15) J. H. Stern and A. A. Passchier, J. Phys. Chem., 6 6 , 752 (1962). (16) L. I. Katzin a n d E. Gebert, J. A m . Chem. Soc., 77, 5814 (1955).
1205
good agreement with the heat calculated from the determination by Davies and G ~ v y n n eof~ the ~ temperature coefficient for reaction 14. Using Davies and Gwynne's14 value of K14 = 768 with our free energy of formation of Iz(aq) and the NBS6 free energy of formation of I-(aq), we calculate AFfo = - 12.35 kcal./mole for 13-(aq). Similarly, we calculate 58.5 cal./deg. mole for the standard partial molal entropy of 13-(aq). We have also investigated the heat of reaction of K I o 3 with HI. Table V presents both experimental and calculated data for these experiments. I n each experiment two reactions took place in the calorimeter KI03(c)
+ 6H+(aq) + 81-(as) 313-(aq) + 3Hz0 + K+(aq) =
(17)
and reaction 14. In Table V, Q represents the heat measured in the calorimeter, Q14has the meaning given before, and AH17 is the heat of reaction 17. Assuming that heats of dilution of triiodides do not differ from those of the corresponding iodides in dilute solutions and combining heats of dilution of K I and H I with our values for AH,,, me obtain AH170 = -73.4 f 0.8 kcal./mole of KIO3. TABLEV REACTION OF KIOs(c) WITH AQUEOUS HI G. KIOs(o)/ Q50 ml. soln.
Moles HI/ 9rjO ml. soln.
0 9696 .5930 ,4684 .2547 .2179
0.05577 .11163 .055'77 ,055'77 .167:30
(Gal.)
(cal.)
AHIT (kcal./ mole)
331.52 202.08 159.23 85.91 75.78
3.45 0.46 .85 .38 .95
-73.94 -73.10 -73.14 -72.50 -74.52
Q
&I4
Combination of the above AH170 with our AHfO of 13-(aq) and NBSS data for I-(as), HzO, and K+(aq) leads to AHfO = - 120.3 kcal./mole for KI03(c). Then the previously determined heat of solution of KI03(c)l7 leads to AHfO = -53.7 kcal./mole for IOs-(aq), compared to earlier values of -54.6 kcal./molelL and -55.0 kcal./mole.6 Our values of standard free energies and heats of formation and partial molal entropies are summarized in Table VI. TABLEV I SUMMARY OF THERMODYNAMIC PROPERTIES (298'K.) SZQ
Species
AHfO (koal./mole)
AFfO (kcal./mole)
(cal./deg./ mole)
Brz(aq) OBr-(aq) HOBr( aq) Iz(aq) Ia 7a s ) 1 0 3 7 aq) KIodc)
-0.20 .-22.6 .-29.4 5.2 .- 11.9 .-53.7 - 120.3
1.38 -7.8 -19.5 3.93 -12.35 -31.5 -100.5
31.1 8.8 25.4 32.3 58.5 28.4l7 36.26
Acknowledgment,,-This work was supported by the National Science Foundation. We thank J. G. Spencer for making available the data in Table V and Professor Henry Frank for his helpful comments. (17) J. G. Spencer and 1,. G. Hepler, J. Phys. Chem., 64, 499 (IQ60).
