Thermodynamic Properties of n-Alkylamine Hydrobromides in Water
Achnowlea'gmmts. 1;t is a pleasure to thank Professor
L. De Maeyer ,and Dr. L. Hellemans for numerous stimulating discussions and helpful comments. Thanks are due to Mr. F. Nauwelaerfa for his excellent experimental work and to Mr.
The ratio Au/uo is obtained to i 0 . 3 % rap to the limit of sensitivity.13 For pure water, uo was calculated from cpo at different temperatures.14 Temperature Control. The temperature of all instruments was kept constant by means of a closed-loop thermostat.l5 Their stability, 3=0.0005" at 25", decreased to &0.005" at 5". The absolute value of each experimental temperature was measured with a calibrated thermistor probe (Yellow Spring) and believed to be accurate to f0.03".
Results The experimental (d - do), AHI*, and ALT/SOin the region 0.01 < m, < 1 are given in the Appendix (see paragraph at end of text regarding supplementary material). All data obtained with flow instruments Sre the average of a t least triplicate measurements. At low Concentrations in the premicellar region, q5v and +e are given by &, = &,' AVc1" -I- B$'c (7)
+
and $c = +c
+ Acc1'2 4- B 0 ) is more then compensatled by the loss in free space between the solute surface and hydrophobic shell (AV < 0).l In other words, the molar expansibility of the water molecules in the hydrophobic cosphere must be larger than that of the molecules in the bulb., and this difference in molar expansibility of the mollecules must exceed the decrease in volume for the fusion of ice in the cosphere during a unit change in temperature. The observed large and positive $EO of n-RepNhI&%r and n-OctNHsBr are therefore normal when compared with other hydrophobic systems. Solute-Solute Interactions. Information on solute-solute interactions can be obtained from the concentration dependence of th’e thernodynamic functions. Since the Debye-Hucbel limiting law, and probably the distance of closest approach, are the same for all the homologous salts, the extra ion-ion interactions in the prem.icellar region are given by the parameters &‘, &’, and BE’. With &, it IS impossible to represent the deviation from the DHLL with a s:ingle adjustable parameter. With all solutes, there is an initial decrease in $L - Ar,(dom)1/2, and with the hydrophobic solutes there is an eventual increase. The initid negative contribution seems to be present with all solutes (hJydrophobic or hydrophili~)~3,30 and is probably not Eitructural in origin; it seems to disappear when transfer function:; from W20 to D203 or from H20
300
R N HJBr IH20,25 ‘G 1
200
c
L E
-
100
7
*i
E
*
0 -In
0
h D
9
,
4
-100
1
I
-2 0 0
-300
A,, 1
2
3
NUMBER
4
5
8
7
,6
OF CARBON ATOM$
Figure 7. Excess enthalpies at 0.1 m for n-alkylamine hydro-
bromides and ammonium acetate in water at 25”. to urea-water mixtures30 are considered. One suggestion that has been put forward52 is that it is related to the repulsive forces (core potential or distance of closest approach). Levine and on the other hand, prefer to attribute the negative deviation from the Debye-Huckel law to pairwise interactions. It is difficult however to see how these shorter-range electrostatic interactions will lead to a negative contribution to +L while the long-range coulombic forces (Debye-Huckel limiting law) lead to a positive contribution. Despite the complexity of #I,, relative trends can be still studied by comparing ( b ~- AL(dom)’/’ at a fixed concentration (0.1 m). Such a plot is shown in Figure 7 . The concentration dependence of the thermodynamic properties hydrophobic solutes can be interpreted through the overlap of c ~ s p h e r e s . ~The * hydrophobic hydration per ion is decreased if the solutes share part of their cospheres. Therefore, the excess functions should be of opposite sign to those a t infinite dilution. This implies that, as ‘,~ the chain length of the RNH3Br salts increases, B ~ F &‘, and BE’ should become more negative and the excess enthalpy more positive. This is generally observed a t 25” (see Tables I1 and TI1 and Figure 7). However, with the higher members of the series, Bv’ and Bc’ seem to become relatively constant (BE’ of n-OctNH3Br is even much less negative than the others). Despite the larger uncertainty, &’ of n-HepNH3Br and ntive. On the other hand, the excess enthalpies in Figure 7 increase only slowly with chain length up to n HexNHaBr, then rapidly with the higher homologs. The regularity in the standard apparent molal quantities with chain length is therefore not observed with the excess functions. This could mean that the solute-solute interactions of the higher homologs are different than the others, that new interactions occur (e.g., dimerization), or that the model used for the interpretation of hydrophobic-hydrophobic interactions is wrong. It will be necessary to examine the The Journal of Physical Chemistry, Yo/. 78, No. 72, 1974
1224
P.-A. Leduc, J.-L. Fortier, and J, E. Desnoyers
temperature dependence of these functions before drawing any further conclusiom. Temperatun, Effects. If structural interactions are the main contribution to +GO - C,(IN) and $Eo, then we would expect i l decrease of these effects as the temperature increases The temperature dependence of $Go is shown in Figure 3. With the structure-breaking solutes, NH*]ESr, NPE4Ac, andl the three lower members of the homologous series, & * does in fact increase with temperature, but with the higher homologs it goes through a maximum rather than decrease as expected. A similar behavior i.s also o b r ~ r v e dwith tetraalkylammonium halides27 and with non:yltrimethyliammonium bromide.26 It could be that the decrease in c,(STR) i s compensated to some exteat by an increase in mCp(TN), but it is difficult to explain why cp(STR) should not have a large temperature dependence. A clue to the int>erpretationof (d(Pco/dT), might come from the high temperature behavior of simple electrolytes in water. The +Go of most alkali halides go through a The behavior of maximum in the region 60 to 80°.20,55*56 the hydrophobic saks are therefore similar to those of hydrophilic salts except that the maximum in d C o occurs a t lower temperatures, However, to our knowledge, no satisfactory explanation has yet been given to account for the (a+co/aT), of aqueais alkali halides a t high temperatures. Another pos;sible approach to the interpretation of the temperature di:pendence of c,, (STR) could be through an analogy with t:ransfer. functions. The main contribution to Cp(S?'R) is the passage of water molecules from the bulk to the solut,e ~ c o s p h e r e s . ~On ~ , the ~ ~ other hand, for the transfer from 'li.120 to D&13 or from H2O to urea-water mixtures,42 the heat capacity of transfer is reflecting primarily the difference between the heat capacities of the pure soivents or mixed solvents. Transposing this to (d$cQ/laT)p, it could be that this function is reflecting primarily the variation of the heat capacity of pure water with temperature. E the high heat capacity of water was correctly attrihaated tcr the shift in water equilibrium during a rise in temperature, ,&henthe heat capacity of liquid water shou1.d be higher at 0 than at 100". Experimentally the heat capacity of water goes through a minimum and is approximately the same at 0 and 100". Therefore, until we reach a better uadewtanding of the temperature dependence of the hcak capacity of pure water, it might be diffi~5? cult to interprel, ( a & O / a a ) , The volumw q b I , T o vary linearly with temperature, as seen from Figure 4. 'This implies that, within experimental uncertainty, is independent of temperature. Therefore, as with (STR), E(STR) hardly changes with temperature. Here again, no simple explanation can be offered. The temperature dependence of Bv' and Bc' is shown in Figure 5. At hi;& temperatures, t,he parameter & ' follows the order expected from the structural interaction model .54 However, a t low temperatures, the opposite order is observed. At so, Be' of n-OctNH3Br even becomes positive. Similar trends occur with tetraalkylammonium bromides.27 Positive BC' parameters are also observed with hydrophobic solutes having hydrogen bonding groups such a s - 0 H 3 3 and COO-"5,33The effect observed therefore seems genixal even tbough there is no obvious explanation for this behavior. A plausible explanation is that there are extra weak interactions between the solutes that are masked by stronger interactions with most functions.
c,,
The dournai of Physical Chemistry, Vol. 78, No. 12, 7974
I
178
I
- - - .- - -
1
can,7Nn3&
c rr-=
2000
/
C.M.C I I
/-
/
/ 1500 r
8-
100 0
50 0
0
I 02
0.4
06
0.8
1.0
Concentration dependence of +v, &, and gL less the DHLL slopes of n-octylamine hydrobromide in water at 25". Figure 8.
However, if these interactions, although small, are very temperature dependent, they may become very important with functions related to the third temperature derivative of the chemical potential. Hydrophobic Bonding. Micelle formation is often used as a model system for hydrophobic The alkyl chains will come together in order to minimize the overall hydrophobic hydration. The volumes, heat capacities, and enthalpies of n-OctNHaBr are compared up to I M in Figure 8. Micelles begin to form at about 0.2 M;I this corresponds to the minimum in 4" - AVe1I2as well as to the increases more sharply and region where (PL where (Pc - &,c1I2 drops very rapidly. Therefore, the heat of micellization is positive (endothermic), the heat capacity of micellization negative, and the volume of micellization positive.1 It is not easy to evaluate these thermodynamic functions of micellization since the micellization region is not sharp. Micelles are being formed or grow in size, over a wide concentration range (0.2-1 M ) . The order of magnitude of these functions can be approximated from the functions a t 1 M minus those a t 0.2 M and give AV, = +5.2 cm3 moland ACm = -278 J K- mol-I. On the same basis AHm would be about 1000 J mol-I, but its value is more doubtful in view of the large concentration dependence of 4~ in the premicellar region. The signs of these functions are all consistent with the decrease in hydrophobic hydration during association; there are fewer hydrogen bonds being formed ( A H > 0, A S > O), the free space around the solute reappears (A V > O), and the positive contribution of Cp(STR) to +co of hydrophobic monomers is lost. It should be noted also that A@, is large compared with AH,. This means that A H , is very temperature dependent and is in fact positive only a t low temperature; it would probably become negative above 30". Acknowledgment. We would like to thank P. Picker for his collaboration in the initial part of this project, C. Jolicoeur for useful discussion of the interpretation, and A. Hade of Universiti. du Quebec 5 Montrbal for the loan of a digital densimeter. We are grateful to the National Research Council of Canada and to the Quebec Ministry of Education for financial support of this project. One of us,
Thermodynamic Properties of n-Alkylamine Hydrobromides in Water
1225
J. -L. F., is also thankful to the National Research Council for the award of a scholarship.
E. M. Arnett, W. B. Kover. and J. V. Carter, J. Amer. Chem. SOC., 91, 4028 (1969). R . Arnaud, L. Avedikian, and J.-P. Morel, J. Chim. Phys., 76, 45 I1972). - -, G. M. Musbally, G. Perron, and J. E. Desnoyers, J. CoNold. Sci., submitted for publication. G. Perron and J. E. Desnoyers, manuscript in preparation. C. V. Krishnan and H. L. Friedman, J. Phys. Chem., 74, 2356 (1970). J . Greyson and H.Snell, J. Phys. Chem., 73, 3208 (1969). R. B. Cassel and W. - Y . Wen, J. Phys. Chem., 78,1369 (1972). T. S. Sarma and J. C. Ahluwalia. J. Phys. Chem., 76, 1866 (1972). P. R. Philip, J. E. Desnoyers, and A. Hade, Can. J. Chem., 51, 187 (1973), J. E. Desnoyers, R. Page, G. Perron, J. -L. Fortier, P. -A. Leduc, and R. F. Platford, Can. J. Chem., 51, 2129 (1973). C. de Visser and G. Somsen, J. Chem. SOC., Faraday Trans. 7 , 69, 1440 (1973). J. G. Aston and C. W. Ziemer, J. Amer. Chem. SOC., 68, 1405 (1946). J. C. Southard, R. T. Milner, and S. B. Hendricks, J. Chem. Phys., 1, 95 (1932). S. S. Chang and E. F. Westrum, Jr., J. Chem. Phys., 36, 2420 f19621. I , P. R Philip, G. Perron, and J. E. Desnoyers, Can. J. Chem., in press. H. S. Frank and N. W. Evans, J. Chem. Phw., 13.507 (1945). H. S. Frank and W. - Y . Wen, Discuss. Faraday sot., '24, 133 (1957). P. R. Philip and J . E. Desnoyers, J. Solufion Chem., 1, 353 (1972). In the two-state water model presented, the shift in equilibrium contribution to the structural interaction was unfortunately overlooked. N. Desrosiers, 6 . Perron, J. J. Mathieson, B. E. Conway. and J. E. Desnoyers, J. Solution Chem., accepted for publication. J. E. Desnoyers and G. Perron, J. Solution Chem., 1, 199 (1972). D . D . Eley, Trans. FaradaySoc., 35, 1281 (1939). E. Lindman, H . Wenmerstrorm, and S. Forsen. J. Phys. Chem., 74, 754 (1970), C. Jolicoeur and H . L. Friedman, Ber. Bunsenges. Phys, Chem., 75, 248 (1971). R. A. Pierotti, J. Phys. Chem., 67, 1840 (1963); 69, 281 (1965). D. N. Glew, J. Phys. Chem., 66, 605 (1962); see also M. Lucas. ibid.., 76. (19721. ~ .4030 , 6. E. Conway and L.' H. Laliberte, J. Phys. Chem., 72, 4317 (1968); see also Trans. Faraday SOC.,66, 3032 (1970). E. E. Conway and R. E. Verrall, J. Phys. Chem,, 70, 3952 (1966). R. H. Wood, R . A. Rooney, and J. N. Braddock. J. Phys. Chem., 73, 1673 (1969) J -L Fortier, P -A Leduc, P Picker, and J E Desnayers, J Solution, in press. A. S.Levine and R . H. Wood, J. Phys. Chem., 77,2390 (1973). J. E. Desnoyers, M. Arel, G. Perron, end C . Jolicoeur. J. Phys. Chem., 73, 3346 (1969). M. Eigen and E, Wicke, Z.Electrochem., 55, 354 (1951): C. M. Criss and J. W. Cobble, J. Amer. Chem. Soc., 83, 3223 (1961). It has been pointed out by one of the referees that the anomalous temperature dependence of #c0 could arise from the difference between C, and C., It seems that C, may be more directly related to structural effects than C, since Cv of pure water decreases monotonically with increasing temperature as expected from a decrease in structure. Unfortunately, the conversion of #cvoto @cVorequires and While approximate #EO have been knowledge of measured in the present paper, are not known far the present systems. However, such measurements are presently under study in our laboratory and this possible approach to the treatment of heat capacities will certainly be investigated. P. Mukerjee, Advan. Colloid lnterface Sci., 1, 241 (1967). \
Supplementary Material Available. The Appendix, which will appear following these pages in the microfilm edition of this volume of the journal, contains tables of the actual density, enthalpy of dilution, and volumetric specific heat data of the systems studied. Photocopies of the supplement,sry material from this paper only or microfiche (105 x 148 mm, 24X reduction, negatives) containing all of thc supplementary material for the papers in this issue may be obtained from the Journals Department, American Chemical Society, 1155 16th St., N.W., Washington, D. C. 201036. Remit check or money order for $5.00 for photocopy crr $2.00 for microfiche, referring to code number JPC-74-1211. References aiid :Notes J . E. Desnoyers and M. Arel, Can. J. Chem., 45, 359 (1967). J. E. Desnbyers, M. Arel, and P.-A. Leduc, Can. J. Chem., 47, 547 llRACli . J. E. Desnoyer!;, R . Francescon, P. Picker, and C. Jolicoeur, Can. J. Chem., 491, 34630 (1971). P. Picker, Leduc, F. R . Philip, and J. E. Desnoyers, J. Chem. Thermodyn., 3, lSJl (1971). P -A. Leduc ancl J E.Desnoyers, Can. J. Chem., 51, 2993 (1973). P. Picker, I?. Tremblay, and C. Jolicoeur, J. Solution Chem., in pres G. S.Kell,J. C h ' m . Eng. Data., 12, 66 (1967). P. Picker and d.-L. Fortier, manuscript in preparation; see also P. Picker, Can. Res. Deveiop., 7, No. 1, 11 (1974). P. Picker, C. Jolicoeur, and J. E. Desnoyers, J. Chem. Thermodyn., 1, 469 (1969). P. Picker, C. Jolicoeur, and J. E. Desnoyers. Rev. Sci. Instrum., 39, 676 (1968). J. -L. Fortier, P. -A. Leduc, P. Picker, and J . E. Desnoyers, J. Solution Chem., 2 , 4E7 (1973) H. S. Harned and B. B. Owen, "The Physical Chemistry of Electrolytic Solutions," Reinhold, New York, N. Y . , 1958, Chapters 3 and 8. J. -I_. Fortier, P. -A. Leduc, and J. E. Desnoyers, J. Solution Chem., in press H . F. Stimson, Amer. J. Phys., 23, 614 (1955). P. Picker, C. J(,!icoeur, and J. E. Desnoyers, J. Chem. Educ,, 45, 614 (1968). ~, L. A. Dunn, rraits. Fara'daySoc., 64, 2951 (1968). C. G. Matmberg and A. A. Maryotte, J. Res. Mat. Bur. Stand., 56, 1 (1956). G N. Lewis and M. Randall, "Thermodynamics," K. S. Pitzer and L. Brewer, Ed., McGraw-tiill, New York, N. Y . , 1961, Chapter 25 and Appendix 4, P. H. Bevington, "Data Reduction and Error Analysis for the Physical Sciences," FAcGraw-Hill, New York. N. Y., 1969. H. Ruterjans, F . Schreiner, U. Sage, and Th. Ackermann, J. Phys. Chem., 73, 986 (1969) E. P Whitlcw and W . A. Feising, J. Amer. Chem. SOC., 66, 2028 (1964). J , Konicek and I. Wadso, Acta Chem. Scand., 25, 1541 (1971); see also K . Kuijanc, J. Suuskuusk, and I . Wadso, J. Chem. Thermodyn., in press D. M. Alexaridei and 0 . .I. T. Hill, Aust. J. Chem., 22, 347 (1969). j
~
(40) (41)
(42) (43) (44) (45) (46) (47) (48)
~~~
(49) (50) (51) (52) (53) (54) (55) (56) (57)
(58)
The Journalof Physical Chemisfry, Vol. 78, No. 12, 1974
