Marie-Madeleine Marciacq-Rousselot-and Michel Lucas
1056
Nuclear Magnetic Resonance Chemical Shift of the Water Proton in Aqueous Allcoholic Solutions at Various Temperatures. Some Thermodynamic Properties of 'These Solutions Marie-Madeleine Marciacq-Rousselot Laboratoire de Physique Experimentale Molectilaire, Universite Paris V I , paris, France
and Michei Lucas" Departemenf de Genie Radioactif, Centre d'Etudes Nucleaires, 92260 Fontenay Aux Roses, France (Received October 6 , 1972) Pubhication costs assisted by Commissariat a I'Energie Afornique
In the present article, we report the results of nmr measurements of the chemical shift of the water proton in solutions of methyl, ethyl, n-propyl, n-butyl, and tert-butyl alcohol at various temperatures and concentrations. At low temperature the downfield shift of the water proton is consistent with an increase in water structure promotion by the solutes. At high temperatures, the upfield shift i s consistent with a weakening of the water structure by the solutes. These results are also qualitatively consistent with what ma> be expected of the influence of hard-sphere molecular solutes on the water structure. The enthalpy of tyansfer of some alcohols from light to heavy water a t 51" is given, but the classical interpretation of this type of measurements is questioned. Finally the enthalpy of transfer of t-BuOH from HzO to tBuOH aqueous solutions is also given and discussed according to current interpretations based on the overlap of hydration spheres around a hydrophobic solute. These interpretations are also questioned.
~ntr~uction A consideration osf the influence of hard-sphere molecular solutes on water structure has yield the following interpr.etati0n.l At temperatures lower than 4", the temperature at which the pure water expansion coefficient is equal to zero, a hard-sphere solute enhances the bonds between water mollecules, and the bigger the solute the more important is the enhancement. When the temperature i s raised this enhancement gradually diminishes and a t temperatures higher than 4" the solute weakens the bonds between water molecules. This effect increases with the solute size and the temperature. In order to test this hypothetical behavior, we report here the results of nmr measurements of the chemical shift of the water piroton in solutions of methyl, ethyl, npropyl, n-butyl, and tert-butyl alcohol a t various temperatures and solute concentrations. Alcohols are different from hard-sphere nonpolar molecules, but less dissiimilar solutes as hydrocarbons or rare gases are not soluble enough to allow measurements .to be performed and it is hoped that the influence of the alcohols on water structure will be somewhat similar to that of hard-sphere solu1;es. In addition, some results of the molar heat of transfer from HzO to DzO for the alcohols a t 51", and from HzB l o t-BuON aqueous solutions at 2 and 60" are also reported and the interpretation of these !measurement discussed in relation with what is expected for a hard-sphere solute. Experimental Section Chemicals. Reagent grade alcohols were refluxed over CaO and then fractionally distilled. The middle fraction was collected. Nmr Measurements. Nmr spectra were obtained with a Varian A 60 spectrometer operating a t 60 MHz and equipped with the 'V 6040 temperature controller. The The Joiirnai of Physical Chemistry, Vo/. 77, No. 8, 7973
chemical shift between the peak due to the water protons and a peak due to protons in the CH3 group of the alcohols was measured. The CH3 protons were used as an in,ternal standard. It has been shown for t-BuOH aqueous solutions that similar results are obtained with an external standard as chloroform or an internal standard.2 Heat of Transfer. The calorimeter and procedure have already been described in detaiL3 The heat of solution of small quantities of liquid alcohol in H20, DzO, or aqueous t-RuOH is measured at given temperatures and the heat of transfer deduced by difference.
Results In Figure 1 the chemical shift (as ordinate) is plotted against alcohol molality (as abscissa) for different alcohols and temperatures. The temperature in general was controlled within k2". The reproducibility of measurements was within 0.5 Hz. Bars of error are shown in the plots. Figure 2 shows the plots of the enthalpy of transfer from light to heavy water for the different alcohols at 25 and 51". Results a t 25" are taken from the l i t e r a t ~ r e Figure .~ 3 shows the plots of the heat of solution of t-BuOH solutions a t 2 and 60" against the alcohol molality. Bars of error are shown in the figure. Discussion 1. Nmr Measurements. Using an internal standard, we have found that at 0 and 40" the molal chemical shift caused by t-BuOH is respectively 2.5 and 0.8 Hz downfield. Corresponding numbers found in ref 2 are respectively 3.0 and 1.0 Hz using an external standard. M. Lucas, J. Phys. Chem., 76,4030 (1972). R. G. Anderson and M. C. R. Symons, Trans. Faraday SOC., 65, 2550 (1969). M. Lucas, Bull. SOC.Chim., 2902 (1970).
G . C. Kresheck, H. Schneider, and H. A. Scheraga. J . Phys. Chem.,
69,3132 (1965).
1057
Nmr Chemical Shift of the Water Proton
!
----c-----+-----
1
m
I
m
t
1
m Z
2
Figure 1. Plots oi tho chemical shift between the Deak due to the water proton and a peak due to protons in the CH3 groups against t h e alcohol molality.
A
A H Kcal/mole
I-% Kcal /mole
I
/ I
0
-1
0.5 25
51
t.
Figure 2. Plots ot the enthalpy of transfer from H 2 0 to D 2 0 for the alcohols at 25 and 51 '.
Another difficulty arises from the contribution of the alcohol hydroxyl p r o h i to the peak due to water protons. The position is given by the mean of the individual water and alcohol resonancm weighted with respect to the reiative number of each type of protons. At -6" two separate proton resonances were observed in solutions for which the t-Bu0Y-I mole fraction is higher than 0.2, with a separation increasing slightly from 23 Hz a t the mole fraction 0.6 to 26 Hz at the mole fraction 0.2, the water proton resonance lying to highej. field. At 25" the separation is about 24 Hz for mole fraction between 0.6 and 0.3.2If a downfield shift from 30 to 60 Ilz is assumed for the t-BuOH hydroxyl resonance at the lowest alcohol concentrations, then the t-BuOH contribution to the observed chemical shift should vary froin 30 m/110 to 60 m/110 where m is the t-BuOH molality in the solution. From a nmr study a t 60 MHz of some ROH-D20 mixtures, Cernicki has concluded that the resonances from n-PrOH and EtlOH hydroxyl protons and HDO-M20 protons do not coacervate until the alcohol mole fraction is lower than O.l.& At 26" the separation decreases from 55 Mz a t a 0.9 n-PrOH mole fraction to 30 Hz a t a 0.1 n-
-2
-3
-I
1
1 2 m 3 Figure 3. Plots of the enthalpy of solution for liquid t-BuOH at 60" and solid t-BuOH at 2" in aqueous I-BuOH solutions against 0
the alcohol molality.
PrOH mole fraction in the PrOH-D20 mixture, from 50 to 20 Hz for the corresponding EtOH-D20 mixtures, and from 55 to 40 Hz when the n-BuOH mole fraction in the BuOH-DzO solution decreases from 0.9 to 0.5. In addition, for the t-BuOH-DzO mixture, the separation de(5) B Cernicki, Ber. Bunsenges Phys. Chem , 69, 5 7 (1965!, 7 0 , 154 (1966)
The Journal of Physical Chemistry, Vo/. 77, No. 8, 1973
Marie-Madeleine Marciacq-Rousselotand Michel Lucas
1058
modal shift 3
--
90 -fo 0.25
0
-3
- 20
10
20
70
to
100
of t h e alcohols molal chemical shift against the water temperature Figure 4. Plots
creases from 25 Wz a t a 0.4 alcohol mole fraction to 22 Hz a t a 0.2 mole fraction, thus being very similar to that found with t-BuOZ-[-H20 mixtures. In all cases the water proton resonance lies to higher fields. Therefore 60 rn/110 is probably an upper limit for the alcohol hydroxyl proton contribution to the chemical shift measured in our study. The results given in Figure 1 are summarized in Figure 4 which shows plots of the molal shift caused by alcohols a t various temperatures. It is apparent that as the temperature is raised, the shifts downfield a t low temperature become upfield at high temperatures. The shift for the water proton only s$ould be more upfield than the shifts plotted in Figure 4, the difference being between 0.2 and 0.6 unit. A difficulty of another sort arises from the interpretation of nmr measurements. An upfield shift is caused for water protons by addition of a structure breaker as CsI or by higher temperatures. Therefore if the addition of a solute to water a t a given temperature causes the water shift to move upfield, this solute may be ascribed as a structure breaker. However, iff one assumes that the H bond has some covalent character, then the direction of the shift (upfield with increasing solute concentration) can be explained by an increase in the covalent character in the presence of the solute with an accompanying increase in the proportion of organized water in the solution.6 On the other hand, if an electrostatic model is assumed for the H bond, an upfield shift on solute addition must be interpreted as a decrease in the H bonding in the water, that is the solute is a structure breaker. According to this model the plots in Figure 4 may be interpreted as showing the solutes acting as structure formers a t low temperature and structure breakers at high temperature. It ib difficult to interpret the results if the partially covalent model IS assumed since it would then result that the alcohols are structure breakers at low temperature and structure formers at high temperatures, a conclusion d ~ f ~ n cto u ~accept. t It is possible tu compare the influence of an alcohol on the water structure to that of corresponding hydrocarbons, as deduced from a water two-state model devised by Frank and frank^.^ In this model water is represented as a two species mixture of dense and bulky (fully hydrogen bonded) constituents. The dissolved hydrocarbon is represented as dissolving separately in these constituents as if they were two “pphar;es.” The mole fraction of the bulky and dense water species is f o and 1 - fo, respectively, and the fraction of the hydrocarbon in the bulky “phase” is go and that in the dense “phase” is 1 - go. Equation 7 in ref 7 provides a simple measure of the structure promoting or breaking produced by the hydrocarbon solute, as the difThe Jowrnai of Physical Chcvnistry, Voi. 77, No. 8, 1973
15
25
t.
35
Plots of go - f a for hydrocarbons in water against the water temperature. Figure 5.
ference between the two numbers go and fo. An interpretation of this equation is as follows: if the hydrocarbon is more soluble in the bulky “phase” than in the other, then the addition of the hydrocarbon to water shifts the equilibrium between the two water species toward the more hydrogen bonded. The variation of the two numbers with temperature and hydrocarbon size is given by eq 13 and 14 of ref 7 . From Frank and Franks’ data, we have computed the values of go and f o in a small temperature range around 25”,assuming constant heat of transfer (since the heat capacity of transfer is not known). Plots of 60 - fo are shown in Figure 5 for CH4, C2136: C&Ig, and C4H10. The solute is considered a structure former when go - fo is positive and a structure-breaker wben go -. fo is negative. The plots show that the bigger the solute, the more it acts as a structure breaker above 27”. However, at low temperatures, all hydrocarbons are probably structure promoters. The explicit conclusion of Frank and Franks is that a negative entropy of solution into water is not evidence of water structure promotion by the solute, but rather that water behaves differently from a nonpolar solvent. An implicit conclusion is that a high partial molar heat capacity for a solute is not evidence for water structure promotion by this solute since butane with a partial molar heat capacity of about 130 cal/mole “K is considered a structure breaker.8 This needs to be outlined since a high heat capacity for BueNBr in water is still considered by many workers as evidence of water structure promotion by this salt. The comparison between Figures 4 and 5 shows that the influences of alcohols and hydrocarbons on the water structure is qualitatively similar in that an increase in temperature changes the alcohol behavior, being a structure former at low temperature and a structure breaker at high temperature. The temperatures at which the alcohols’ structural influences change cannot be determined in view of the uncertainties in the molar chemical shifts and the contributions of alcohol hydroxyl protons to the shifts. 2. Heat of Transfer of the Alcohols from HzO t o DzO. This type of measurements is often used as a probe for the structural influence of a solute on wadt.ar.9 According to (6) (7) (8) (9)
J. Clifford and B. A. Pethica, Trans. Faraday Soc., 60, 1483 (1964). H. S. Frank and F. Franks, J. Chem. Phys.. 48,4746 (1968). T. J. Morrison, J. Chem. SOC., 3814 (1952). C. V. Krishnan and H. L. Friedman, J. Phys. Chem.. 74, 2356 (1970).
1059
Nmr Chemical Shift of the Water Proton
30
MtOH
10
40
I O
for t-BuOH aqueous BuOH nriolality. arid plotr; of the difference
solutions against tbetween t h e expansion coefficient of the solution and that of pure water.
60
40
Figure 6. Plots of L
Friedman's interpretation, negative values are taken as evidence for water structure promotion. Then the plots in Figure 2 may be interpreted as showing that a t 25" the bigger the alcohol the more it behaves a3 a structure former. At 51°, all alcohols have the same influence on the water structure and possible a t higher temperatures the alcohols with tho larger alkyl chain could act as structure brealrers. However, a different interpretation is possible. Eley has derived the relation 2 0 ( a T / @ ) Vbetween z the molal internal energy of solution for a solute from the gas state, the solute molar volume V 2 in solution, the expansion coefficient a , and rhe isothermal Compressibility coefficient @ ofthe pure solvent.l* Then the enthalpy of transfer for a hard-sphere solute from H:!O to E20 should he
Since @D20,~ H z O and , Vz(H20), Vz(Dz0) are probably not much different but uBzO is much smaller than aHzO at room temperature," LH should be negative for a hardsphere solute. When the temperature i s further raised the expansion ccefficients for ET00 arid D20 are less different, so the enthalpy of transfer may become less negative. Then the sign of the experimental enthalpy of transfer should be only the consequence of the fact that .alcohols behave somewhat as hard-sphere solutes, arid it would not give any information on the structural properties of the solute. 3. Heat of 'i'ransjsr of t-HuOH t o Aqueous t-BuOH Solutions. The difference AH - LIP,where AH is the molar heat of' solution of small amounts of liquid (or solid) t BuOH in the aqueous t - uOH solution of molality m and LFP the heat of solution in pure water, is plotted at two temperatures against m in Figure 6. A H - - i\No is also equal to the relative partial heat content Lz. In the same figure is also plotted the difference CY - c y o between the of the aqueous alcohol solution expansian Coefficient and N O the water expansion coefficient. Plots a t 4" are deduced from the influerice of the solute on the TMD (temperature of maximum. density) of its solution.12 Plots a t 45" are deduced from a few measurements of the densities of aqueous t-BuOH solutions at 40 and 50" made in this
120
90
I
i/
t-BuOH
120
100 120
100
1
0
I
a
0.025
x
0.05
0
0.050
X
Figure 7. Plots of (pcp2 - Cpwo for various alcohols at various temperatures against the alcohol mole fraction in the solution. laboratory. The initial negative variation of at 4" clearly shows the enhancement of the water structure by the addition of small quantities of t-BuOH. The further increase in CY - cyo may possibly be ascribed to the increasing disruption of water structure as more alcohol i s added to water. At 45" the steep increase of CY - cyo as t-BuOH is added may possibly be ascribed to the fact that the alcohol enhances the water structure to a much smaller extent. The interpretation of AH - A P for a hard-sphere solute is straightforward since we have shown that a close proportionality exists between AH - Ahr" and CY (At 4" the relation between the two quantities is AH = ( a - aO)T/@VZwhere @ and are respectively the compressibility of the solution and the alcohol partial molal volume.) Then AH - AHo should be positive at high temperature and negative a t 4" for a hard-sphere solute in aqueous t-BuOH solutions. In addition, since a t 4" a hard-sphere solute decreases the expayion coefficient of its solution with water,l it results that Lz for such a solute should be negative a t high dilution. When the solute is not a hard-sphere solute but t-BuOH, the relation be-
vz
(IO) D. D. Eley, Trans. FaradaySoc., 35, 1281 (1939). (11) D. Eisenberg and W. Kauzrnann, "The Structure and Properties of Water," Oxford, Clarendon Press, 1969, p 187 (12) G . Wada and S. Umeda, Bull. Chem. Sac. Jap., 35,646 (1962). (13) (a) M. Lucas and A. Feillolay, J. Phys. Chern., 75, 2330 (1971); (b) A. Feillolay and M. Lucas, ibid., 76,3068 (3972). The Journal of Physical Chemistry. Vol. 77, No. 8, 7973
1060
G. A. Takacs and G. P. Glass
tween z 2 and a - a' is not followed since z 2 is slightly positive a t 2" for low t-BuOH molalites, but LZis smaller at 2"than a t 60". The usual interpretation of a positive value for t 2 is entirely different. The overlap of the more hydrogen-bonded water sphere arounld a structure-forming solute as its concentration is increased is interpreted as producing a positive contribution to z2.14 Since all other effects on r";z are neglected, a positive E 2 is taken as evidence for water structure promotion, by BurNBr by instance. In our opinion too much importance is given to this particular contribution caused by the overlap of the hydration spheres. For instance, O( -. a' a t 4" for aqueous t-BuOH solutions may be equated to - - a ~ i n+ a2m2 where --aim may be interpreted as showing the solute influence on the expansion coefficient if no other interaction as solute-solute interaction, overlap of hydlration sphere . . . is present. The term a2m2 incorporates all these effects. Then the contribution of the overlap to 122 should be a t most proportional to a2m2, provided that other effects are absent. Although this contribution to E2 for a water structure former is certainly positive, it cannot be assumed in our opinion that a positive L z is evidence for water structure promotion in view of all the other effects which may exist. In any case such an interpretation would lead to the conclusion that a t low molalities t-BuOH is a strong structure former a t Go", a weak one a t 2" and possib_ly a structure breaker at lower temperatures if a negative LZ is found as the extrapolation of data a1 i! and 60" suggests. This concluslon is difficult to accept. An interesting consequence of the variation of A H ilHo with the temperature is that the partial molal heat capacity of t.BuOI-I around 300 should a t first increase with its molality and then should decrease when the molality is further increased.
Figure 7 shows the plots of +cP2 - epwo against the alcohol mole fraction in the aqueous solution at various temperatures for some alcohols, where +cp2 is the alcohol apparent molar heat capacity and cpwo the pure water heat capacity. These plots are drawn from cp data in the 1 i t e r a t ~ r e .The l ~ values of cPzo - cpWohave been computed from Hill's datas6 and the liquid alcohols heat capacities.17 For all alcohols except MeQH and EtQH these plots show a similar behavior for the variation of 4cp2 with the concentration and support our findings for the variation of the t-BuQH partial molar heat capacity. We do not interpret high solute heat capacity as evidence for water structure promotion by this solute, but rather as showing that the solute structural influence changes much with the temperature.I8 Then an increasing cPz is in our opinion related to a change of the solution structure with the temperature more important than the pure water structural changes, since the solute is assumed to be a structure former a t low temperature and a structure breaker at high temperature, a t moderate concentration. A possibly similar idea is expressed by Frank: " . . . An appropriate heat capacity contribution must also be expected, arising from the transformation, with rising temperature, of bonded to non bonded species."lg
(14) R . H . Wood, H. L. Anderson, J. D. Beck. J. R. France, W. E. de Vry, and L. J. Soltzberg, J. Phys. Chem, 71, 2149 (1967). (15) W. S. Knight, doctoral dissertation, Princeton, 1962. (16) D. M. Alexander and D. J. T.Hili, Ausi. J. Chem., 22, 347 (1969). (17) (a) Landolt-Bornstein, "Physikalische-chemische Tabellen," Voi. 2. Part 4, Julius Springer, Berlin, 1960. ( b ) F. L. Oetting, J , Phys. Chem., 67, 2757 (1963). (18) M. Lucas and A. de Trobriand, C. R. Acad. Sci., Ser C, 274, 1361 (1972). (19) H. S. Frank in "Water a Comprehensive Treatise," Vol, 1, F, Franks, Ed., Plenum Press, New York, N. Y . , 1972, p 543.
Reactions of Hydrogen Atoms and Hydroxyl Radicals with Hydrogen Bromide C. A. Takacs and G . P. Glass* Department of Chemistry, Rice University, Houston, Teras 77001 (Received October 10, 1972) Pb'blkation costs assisted by The Petroleum Research Fund
A fast discharge-flow system was used to study reactions of hydrogen atoms and hydroxyl radicals with hydrogen bromide a t 295 K. The reactions were followed by monitoring the epr spectra of H(2S1,2), Br(2P3,~),and OH(2n3,2) a t a number of different reaction times. A fluorinated halocarbon coating, applied to the flow tube, was found to be extremely effective in eliminating wall recombination of Br(2P3/2). Rate constants of (3.4 & 0.8) x 10-12 and (5.1 & 1.0) x 10-12 cm3 molecule'l see-I were obHz + Br and OH HBr H20 + Br, respectively. An unsuccesstained for the reactions H + HBr ful search was made for Br(2PI,z).
-
Introduction Hydrogen bromide is known to inhibit combustion of hydrogen and hydrocarbon fuels. Numerous studies have demonstrated its effect on flame propagation, flamability, and explosion lirnits~~-7 It i s generally accepted that the The Journal of Physicai Chemistry, Vo/. 77, No. 8, 1973
+
-
inhibition results from reactions of HBr with chain centers which are important for flame propagation; for example, with H and OH. However, absolute rate constants for D . R. Clark, R , F . Simmons, and 66, 1423 (1970).
D.A .
Smith, Trans. Faraday SOC.,
