April 20, 1959
HEATS
AND
ENTROPIES OF IONIZATION O F
entropy increment or that i t is spread out over a rather broad temperature interval. Valuesof C,, So, H" - H; and - (F" - H,")/T a t selected temperatures are listed in Table 11. The entropy and enthalpy increments were computed by numerical integration, using graphically interpolated values of heat capacity. The heat capacity values are considered to have a probable error of about 5y0 at 5"K., 1% a t 10°K. and 0.1% above 25°K. Values below 5°K. were extrapolated with a T 3 function. The effects of nuclear spin and isotopic mixing are not included in the entropy and free energy functions and the values listed in Table I1 are therefore the conventional ones to be used in calculations of chemical equilibria. The estimated probable error in the thermodynamic functions is 0.1% above 10O0K., but some of the values are given to an additional digit for comparison purposes. The present measurements accord better with the data by Parks and Kelley3 on specular hematite than on the finely divided powder but are 3% lower near 90°K. and within about 0.5% a t higher temperatures, while they are 2 to 6% lower than those on the Kahlbaum sample. Inasmuch as even
PHENOL AND STJBSTITUTED PHENOLS
1783
the specular hematite contained 0.5y0 HzO and 0.5% Si02 and no attempt was made to correct for these impurities, the agfeement must be considered satisfactory. Since Parks and Kelley3 used the data on the specular crystals in calculating the entropy, 21.5 f 0.5 e.u., the entropy increment over the common region (90 to 295°K.) agrees quite well with that found in the present investigation. Most of the difference of 0.6 e.u. is in the extrapolated portion as a consequence of the errors in their lowest temperature data. Acknowledgments.-The support of the Division of Research of the U. S. Atomic Energy Commission is acknowledged with gratitude. F. Grgnvold expresses his thanks to the University of Oslo for leave of absence. The authors also wish to thank Professor Lars Thomassen for the use of X-ray equipment, hir. Clinton F. Jefferson for chemical analyses and magnetic moment determination of the a-Fe203 sample, Miss Elfreda Chang and Mr. Roger Berg for assistance with the operation of the calorimeter and Mr. Wallace Sue for assistance with the calculations. ANNARBOR,MICHIGAN
[CONTRIBUTION FROM COBB CHEMICAL LABORATORY, UNIVERSITY OF VIRGIXIA]
Heats and Entropies of Ionization of Phenol and Some Substituted Phenols BY L. P. FERNANDEZ AND L. G. HEPLER RECEIVED OCTOBER 27, 1958 Standard heats of ionization of aqucous phenol, o-chlorophenol, p-chlorophenol, o-nitrophenol and p-nitrophenol have been determined calorinietrically at '298°K. These heats of ionization have been used with free energies, obtained from ionization constants, t o calculate standard entropies of ionization of these compounds in aqucous solution. A detailed quantitative interpretation of the thermodynamics of ionization of phenol and substituted phenols is impractical until more d a t a are available b u t some interesting comparisons of entropies of ionization are discussed.
This investigation of the heats and entropies of ionization of aqueous phenol and some substituted phenols was undertaken as part of a program aimed a t the accumulation of thermodynamic data on the ionization of certain selected organic acid; in water. I t is to be hoped and possibly expected that systematic accumulation of such data on selected acids, combined with similar data of other investigators, will lead to a detailed interpretation of the thermodynamics of aqueous acids in terms of the structure of the weak acids and their anions and of solute-solvent interactions. Experimental T h e solution calorimeter used in this investigation has been described in detail.'#* All of the calorimetric work reported in this paper was carried out with 950 ml. of water or solution in the calorimeter. All heats of reaction wid solution were investigated at 25.0 f 0.3". Baker C.P. phenol was distilled under reduced pressure with only the middle fraction saved for calorimetric experimeiits. Eastman Kodak Co. practical grade o-chlorophenol and p-chlorophenol also were purified by fractional distillation, The o-chlorophenol and P-chlorop!ienol for calorimetric experiments boiled a t 173' antl melted at 42.Y-14.8', respectively.
Sodium o-nitrophenoxide was prepared by adding o-nitrophenol (m.p. 45-5G0) t o a solution of NdOH in ethanol. T h e resulting phenoxide was recrystallized once from 95y0 ethanol and twice from absolute ethanol. I t then was dricd and stored for future analysis and use. T h e p-nitrophenol (m.p. 112-114') used for our calorimetric experiments was supplied by Eastman Organic Chemicals. Sodium p-nitrophenoxide dihydrate was prepared by adding qome of this p-nitrophenol t o a solution of NaOH in ethanol. T h e salt formed was recrystallized twice from ethanoikwater mixtures and washed with ethanol before drying and storing for future analysis and use. Sodium o-nitrophenoxide was analyzed by titration with aqueous HCl. T h e yellow o-nitrophenol formed as a product of the titration reaction partly precipitated out and partly remained in solution and tended to obscure the indicator color change a t the end-point. To eliminate this difficulty, 10 ml. of CCl, vias added t o the solution to extract and dissolve the o-nitrophenol, thus leaving the water layer nearly colorless. hlethyl red indicator then was added and the solution titrated to the indicator color change without removing the CClr layer. The titrations required 100.87; of the tllr oretical amount of HCI. Sodium p-rutrophenoxide dihydrate also was analyzed by titration wit!] aqueous HCI. This compound could serve as its own indicator, but the addition of one drop of methyl red solution made t!ie color change a t the end-point easier to see. The titrations requircd 101 ;1(' of the theoretical unouiit r ~ f €IC1
Results and Calculations lye have measured the heat of solution of phenol in water antl the heat of neutralization of phenol
L. P. FERNANDEZ AND L.G. HEPLER
1754
by 0.1039 d4 NaOH as in equations 1 and 2 . CsHjOII(C) = CsH,OH(aq) CsH,OH(c) OH-(aq) = CaHSO-(aq)
+
+
AH1 H20 AH2
(1) (2)
Standard heats of reactions 1and 2, AHlo and were obtained from the experimental heats at finite concentrations (given in Tables I and 11) b y graphically extrapolating to infinite dilution. Heats so obtained were AHIo = 3100 f 50 and AHZo = -4750 f SO cal./mole where f indicates our estimate of the maximum total uncertainty. HEATSOF
TABLE I SOLUTIOX OF PHENOL IN WATER
ITtJes CsHaOH/SSO ml.
AH1 (cal./mole)
0.005574 ,005919 01032 ,0108” ,01273 ,01431
3091 3113 3164 3091 3016 3103
+chlorophenol. Results of these calorimetric experiments are given in Tables 111, IV, V and VI. Graphical extrapolations to infinite dilution have yielded AHlO’ = -78 += 20, AHzO’ = -8950 80. AHlo’’ = 3825 f 60 and AHZo”= - 3880 f 60 cal./ mole. Standard heats of ionization of aqueous o-chlorophenol and p-chlorophenol (analogous t o AH30 and designated AHSO’ and AH3O”) have been calculated from these heats of solution and neutralization in the same way that AH30 was calculated. We have found AIg3” = 4630 80 and AH3°” = 5800 90 cal./mole for o-chlorophenol and p chlorophenol, respectively,
*
*
TABLE I11 HEATSOF SOLUTION OF O-CHLOROPHEUOL Moles o-CICeH~OH/950 ml
A H , ‘ (cal /mole)
0 009028 01346 05473 06043
- 67
TABLE I1 IIEATS
OF
SEUTRALIZATION
OF
PHENOL
BY
0.1039
AH2 (cal./mole)
0 004370 004009
- 4728
,0104R 01572 01735
0333; ,03GC,”
-4752 -4765 -4713 - 4702 -4733 -4716 -4T32
The standard heat of ionization of aqueous phenol as in equation 3 has been calculated from AHlo, AMao and the heat of ionization of water, AHw0.3 CsH50H(aq) = CbHSO-(aq) 4- H + ( a q ) AHaO (3) The relation 1 f 1 3 0 = AH20 - AHlo AHwoleads to = 5630 f 100 cal./mole. The estimated uncertxinty for AH3O is less than the sum (or the square root of the sum of the squares) of uncertainties in J.HlO, AHZO and AHwobecause AHI0and AH2O were determined under similar conditions in the same calorimeter so that systematic errors partly cancel each other. The ionization constant of aqueous phenol a t 25” has been investigated many times and we have chosen K 3 = 1.05 X References 4-7 describe representative investigations of this ionization constant. \Ye have calculated from Kz that AF30 = 13(i10 cnl./mole and have combined this free energy with AH3O to obtain AS3O = -26.7 cal./deg. mole. 1i.e have also measured heats of solution and neutralization of o-chlorophenol and p-chlorophenol. ’The calorimetric reactions were analogous to rexctions 1 and 2 for phenol and the heats of these rr:ictions have been designated AHI’, 4H2’, ABl‘’ :!rid LII2’’ where ’ refers to o-chlorophenol and ” to
+
P ~ p t ‘ eLI’. J . C a n a d y a n d K . J . Laidler, C o n . J . C h e m . , 34, HigK’. 7 ’ , c t i ~ s , F u t o h y .%IC, 62, 35 (1956).
, I(
Slirrnglixig a n d G. LV. I,ervis, T H I S. ? O U R N A L , 7 5 , 5709
IT Joilion a n d hi. Kilpatrick, i b i d . , 71, 3110 (1949). ( i l I:. G . Dordrrrll and G. D. Cooper, i b i d . , 74, 1058 (1952). (Cj C.
- 88 - 79 +-
-,a
TABLE IV l v
SaOI-I l I ~ ! v sCsHsOII/’930 ml.
,007862
VOl. 18
H E A T SOF SEUTRALIZATIOV OF O-CHLOROPHENOL B Y 0 1279
‘tr X ~ O H Moles o C l C s H ~ O H / Q i Oml
AHz’ (cal /mole)
0 002116 002496 004804 006606 01821
HEATSO F
-8915 - 8900 - 8979 - 8940 - 8851
TABLE 1’ ,b-CliLOROPHE\OL
SOLUTIOh OF
Moles p ClCsHICH/950 ml
A ? I i ” (cal /mole)
0 006661 008469 009627 01444 0251 1 03140 04375
3849 3829 3860 3938 4000 3998 4045
T A B L E1’1 HEATSOF NEUTRALIZATIOS OF p-CHLOROPHESOL Moles p-CICsHdOH X 102/950 ml.
1 1184
1 229 1 428 I.796 2 292 2 587 3 476 4 058
N a O H , .Vi
0.1279 ,1170 1170 1279 1407
1.599 1279 1727
Ionization constants a t 25’ for aqueous ochlorophenol and p-chlorophenol have been determined by Judson and Kilpatrick who have also reviewed earlier investigations. We have calculated from the constants given by Judson and Kilpatrick that the standard free energies of ionization of aqueous o-chlorophenol and p-chlorophenol are 11570 and 12800 cal./mole, respectively. These free energies and our heats of ionization (AH30’ and AH30”) have been used to calculate that the standard entropies of ionization of aqueous o-
HEATSAND ENTROPIES OF
April 20, 1959
ICNIZATION OF PHENOL AND SUBSTITUTED PHENOLS
chlorophenol and p-chlorophenol are both - 23.5 cal./deg. mole. The small solubility and slow rate of solution of o-nitrophenol made the calorimetric determination of the heat of solution of this compound difficult. Therefore, a procedure different from that followed for phenol and the chlorophenols was used for calorimetric determination of the heat of ionization of aqueous o-nitrophenol. The heat of solution of the salt, sodium o-nitrophenoxide, was measured in dilute NaOH in order to prevent any appreciable hydrolysis of the anion. The calorimetric reaction is given by equation 4 and the experimental
+
o-SOnCsHaOSa(c) = o-NO2C6H4O-(aq) Na+(aq) AH4
(4)
results are in Table VII. We have extrapolated t o infinite dilution to obtain AH4O = 2130 f 60 cal./mole.
1785
We have extrapolated from these heats to obtain AfJIO”’ = 5400 f 100 and AHzo“‘ = -3270 f 90 cal./mole and have calculated AH30”’ = AH20”l AfJlO”’ AHwo = 4830 f 130 cal./mole for the heat of ionization of aqueous P-nitrophenol a t 298’K.
+
TABLE IX HEATSOF SOLUTION OF NITROP PHENOL A!Ii”t
Moles p-NOzCsH4OHI950 ml.
(cal./mole)
0.006493 0.01198
5385 5423
TABLE X HEATS OF NEUTRALIZATION OF ~ - Y I T R O P H E SIS O 0.1170 L Jf NaOH Moles p-NOnCsH~OH/950ml.
A H i “ ‘ (cal./mole)
0,005166 0.01009
-3255 -3235
TABLE VI1 The heat of ionization of aqueous p-nitrophenol HEATSOF SOLUTION OF SODIUM 0-NITROPHENOXIDE also has been investigated by the method used for Moles o-NOnCaH10Pu’a/ 950 ml
NaOH, M X 103
A H , (cal./mole)
0 002674 ,006205 ,01042
3.823 5.310 2.655
2141 2187 2166
The heat of reaction of sodium o-nitrophenoxide with dilute HCl to form aqueous o-nitrophenol as in equation 5 also has been measured. We have o-SOzCsHiONa(c)
+ H + ( a q ) = O - N O L ! ~ H ~ O H +( ~ ~ ) Na+(aq) AH5
(5)
extrapolated the results of these experiments (given in Table VIII) to infinite dilution to obtain AH6o = -2525 f 50 c.d./mole. The heat of ionization of aqueous o-nitrophenol has been calculated as AH40 - AHso and found to be 4655 f 70 cal./mole. TABLE VI11 HEATSO F REACTION O F SODIUM 0-NITROPHENOXIDE 0 02444 .lf HC1 Moles o-XOgCaH40!‘ia X 10,’/950 ml
AHa (cal./mole)
3.423
-2526
6.596
-2513 - 2527
9,229
WITH
Biggs4 and Judson and KilpatrickE have reported ionization constants of aqueous o-nitrophenol a t 25” t h a t are in good agreement with each other. The average of their values, 6.0 X has been used to calculate the standard free energy of ionization of aqueous o-nitrophenol to be 9850 cal./deg. From this free energy and our heat of ionization we have calculated that the standard entropy of ionization of o-nitrophenol is -17.4 cal./deg. mole. Neither the solubility nor the rate of solution of p-nitrophenol causes undue difficulty in the calorimetric investigation of the heat of solution of p nitrophenol. Therefore, we have determined the heat of ionization of aqueous p-nitrophenol in the same way that we investigated phenol and the chlorophenols. Results of our heat of solution and heat of neutralization experiments (designated AH,”’ and AH2”’) are given in Tables IX and X.
o-nitrophenol. Heats of solution of sodium pnitrophenoxide dihydrate in dilute aqueous NaOH and in aqueous HCl (designated AH4‘ and AHs’, respectively) have been measured and the results of these experiments are given in Tables XI and XII. We have extrapolated to infinite dilution TABLE XI HEATSO F
SOLUTION O F SODIUhl ~ - ~ I T R O P I < E N O X I D E111-
HYDRATE Moles p-NOd2eH~ONa.2HzOX lOJ/ 950 ml.
AI1.I (cal./m(ile)
2,751 5,635
7804 7Gis
TABLE SI1 HEATSO F REACTIOX O F SODIUM I)-NITROPIfESOXIDE DIHYDRATE WITH n.2444 M m i Moles p-XOzCsH,ONa.2HzO X loa/ 950 ml.
A U , ’ (cal./molej
3.908 7,327
0216 3158
to obtain AH40‘ = 7750 f 180 and AI160‘ = 3150 f 120 cal./mole and have calculated AIIdn’” = AH40’ - AHSO’ = 4600 f 210 cal./mole. LT’e have taken 4700 cal./mole as the best value for the heat of ionization of p-nitrophenol and this value has been used in calculating the entropy of ionization to be -16.9 cal./deg. mole. For this calculation we used a free energy of ionization of p nitrophenol ( A P = 9750 cal./niolc) based on an average ionization constant ( K = 7.2 X 1(1-‘) taken from the results of Judson and KiIpatrickfl and Robinson and Biggs.* I t might be noted that had we used. AHwo= 13.36 kcal./mole as found by Pitzer” and others, most of the difference betweeii our two values for the heat of ionization of /.Initrophenol would disappear. Nevertheless, we have used AHWo= 13.50 kcal.,/mole because this recent calorimetric value is in accord with the heat calculated from the temperature coefficient of the ionization constant of water. (8) R. A . R o b i n s o n nnd A . (19%). (9) K .
S. Pitzer,
1. Digg?,
l’rtiirr
12ur,~,io??
THISJ U U R N A I . . 59. 2:3ti5 (I9:l;).
w
,
6 1 !In1
L. P FERNANDEZ AND L. G. HEPLER For easy comparison we have gathered in Table XI11 the ionization constants and heats and entropies of ionization at 25” of aqueous phenol and a number of substituted phenols. Included in Table XI11 are data on m-nitrophenol. The heat of ionization of m-nitrophenol has been determined calorimetrically by Cottrell, et ~ 1 . ’ ~We have calculated the free energy and entropy of ionization from the ionization constant reported by Judson and Kilpatrick6 and the heat reported by Cottrell, et a1.I0 It may also be noted that Cottrell, et U L . , ’ ~ have reported heats of ionization of o-chlorophenol and p-nitrophenol that are in good agreement with our heats of ionization of these compounds.
Vol. 81
result of losing a proton and acquiring a net negative charge. This problem is usually approached from the point of view of resonance, inductance, etc., effects on the energy or sometimes free energy difference between the acid and its anion. An equally important problem is consideration of energy efTects arising outside the acid and anion due to differences in the acidsolvent and anionsolvent interaction energies. This combination of problems makes theoretical investigation (or even correlation) of heats of ionization appear less promising a t present than investigation of entropies of ionization. When considering the entropy of ionization in detail, we must be concerned with entropy changes TABLE XI11 due to differences within the acid and its anion and THERMODYSAMICS OF IONIZATION OF AQUEOUSPHENOLS also with entropy changes due to differences in A H 0 (tal./ A S 0 fcal./deg. solute-solvent interactions. In this case the first Substance K m X 10’0 mole) mole) problem has been discussed by Pikerg who has Phenol 1.05 56.50 -26.7 shown that for a series of acids the “internal” o-Chlorophenol 33.4 4630 -23.5 entropies of the anions are less than those of the p-Chlornphcnol 4.18 5800 -23.5 corresponding acids by an amount that is very 600 4655 -17.4 o-Sitrophenol nearly constant. We conclude, therefore, that 4700 -16.9 p-Sitrophenol 720 most of the differences in entropy of ionization for vr-Sitrophenol 45.1 4705 -22.5 the acids in Table XI11 are due to differences in solutesolvent interactions. Discussion LVith this background in mind, it is interesting to The principal aim of the research reported in this paper has been to obtain experimental data that consider specifically the entropies of ionization of will further our attempts to deduce and make use of the nitrophenols. We conclude from these enrelations between acid strength and structure. tropies of ionization that the anion of m-nitrophenol The following discussion of our data is relevant to (relative to undissociated m-nitrophenol) is more effective in orienting neighboring water molecules that end. I t may first be noted that substitution of a than the anions of 0- or p-nitrophenol (relative to chlorine or nitro group for hydrogen in the 0-, undissociated 0- or p-nitrophenol). It seems likely m- or p-position yields an acid that is strongtr that this extra orienting effectiveness of the mthan phenol. Other more interesting observations nitrophenol anion results from a greater localization ccncern the relative importance of the heat and of negative charge on the phenolic oxygen as comentropy of ionization for the various phenols under pared to anions of u- and p- nitrophenol. This arguconsideration. .Along these lines, these several ment is in agreement with conventional ideas about observations are pertinent to the problem a t hand: the small tendency of metu substituents, as com(i) the heat of ionization of p-chlorophenol is more pared to ortho or para substituents, to take part in endothermic than that of phenol but, in spite of resonance with the phenolic function. In summary, it is well to emphasize that a satisthis, the free energy of ionization of 0-chlorophenol factory correlation of strengths and structures of is less positive (greater K ) than that of phenol organic acids in aqueous solutions must not be because of the considerable difference in the enbased solely on energy (or even free energy) tropies of ionization. (ii) IT-e note that m-nitrochanges or differences within the acid and its anion. phenol is a weaker acid than p-nitrophenol and oIt is equally important that differences in solutenitrophenol because of a considerable difference in entropy of ionization and not because of a more endo- solvent interaction be considered in as much detail thermic heat of ionization. Further, about two- as possible. These solute-solvent interactions are thirds of the increase in acid strength of p-nitro- presently most easily considered in terms of enphenol and o-nitrophenol over phenol is an entropy tropy changes. Acknowledgments.-n‘e are pleased to express effect. -1ny attempt to calculate or correlate heats of our gratitude to Professor T . I. Crowell and l f r . ionizatioii of a series of acids m u s t take account of Robert T. Kemp for helpful discussions and to the the energy changes that occur within the acid as a National Science Foundation for necessary financial support. f:U\ T I. Cottrell. G. ’A’ Drake, D . L. Levi, K . J. T u l l y and J. H. ‘\\‘i,ifenilen J
CI:,tn. SOL. 1016 (1948).
CHARLOTTESVILLE, VIRCISIA
