2244
HARRYP. HOPKINS,JR.,CHING-HSIENWu, AND LORENG. HEPLER
of the optical densities a t 7300 in ethanol solution of anthracene and in acetone solution of anthracene, = 1.0 for the absolute reference value being GZ:!,Fo' the solvated electron in ethanol from our previous work.lG The use of this ratio G(anthracene-) =
D730oaCetone
X 1.0
(V)
07300
assumes that the extinction coefficient of the anion in acetone is the same as in ethanol. The data give G (anthracene-) = 1.7 f 0.6 molecules/100 e.v. Most of the uncertainty originates in the value for GE~~~?O1. In pulse-irradiated acetone, without any additive,
a weak transient absorption is observed, starting at about 6000 A. and increasing to about 3200 A., the transmission limit. The decay is complex, showing the presence of several absorbing species.
Acknowledgments. We are greatly indebted to Mr. Douglas Harter for his technical assistance in this work. The linear accelerator was operated by Mr. B. E. Clifft and Mr. E. Backstrom. We are grateful to Dr. Myran C. Sauer, Jr., with whom we have had frequent discussions. (16) I. R. Taub, D. A. Harter, M. C. Sauer, Jr., and L. M. Dorfman, J. Chem. Phys., 41, 979 (19134).
Thermochemistry of Aqueous Aminosulfonic Acids.
Sulfamic and
Sulfanilic Acids and Taurine
by Harry P. Hopkins, Jr., Ching-Hsien Wu, and Loren G. Hepler Department of Chemistry, Carnegk Institute of Technology, Pittsburgh, Pennsylvania (Received December $8,1964)
Results of calorimetric determinations of heats of solution and of ionization of sulfamic acid, sulfanilic acid, and taurine are presented and used in subsequent thermodynamic calculations. For ionization of the aqueous acids we find AH" = 0.25, 4.2, and 9.99 kcal. mole-' and AS" = -3.8, -0.7, and -7.95 cal. (deg. mole)-' for sulfamic acid, sulfanilic acid , and taurine, respectively. These small negative values for entropies of ionization provide supporting evidence that the aqueous acids exist in the zwitterion form. We calculate Sz" = 56.4 and 48.4 cal. (deg. mole)-' for the standard partial molal entropies of aqueous taurine and the aqueous anion derived from taurine, respectively.
In addition to the inherent biological importance of some aminosulfonic acids, such as taurine, these substances are of interest because of their structural relation to amino acids. Like the amino acids, the aminosulfonic acids can exist in solution as zwitterions. Acidities of the aminosulfonic acids vary widely, sulfamic acid being of special interest because it is an "almoststrong" electrolyte. This paper is concerned with our The Jaurnal of Physical Chmiatry
calorimetric investigations of these acids and subsequent thermodynamic calculations.
Experimental The calorimeter used is patterned after one previously described,' except that a Mueller G-2 bridge and (1) W. F. O'Hara, C.-H. Wu, and L. G. Hepler, 619 (1961).
J. Chem. Educ., 38,
THERMOCHEMISTRY OF AQUEOUS AMINOSULFONIC ACIDS
HS galvanometer have been used with a nickel wire resistance thermometer for temperature measurements. Also, the resistance thermometer and calibration heater are contained in a glass spiral filled with mineral oil rather than wound on a silver cylinder. All of the calorimetric work reported here was carried out with 950 or 960 ml. of water or aqueous solution in the calorimeter with mean reaction temperature of 25.0 f 0.2". Sulfamic acid (HS03NH2) was purified according to the method described by Sisler, Butler, and Audrieth.2 The crystals were dried in a vacuum desiccator, broken up with a glass rod, dried in an oven at 40" for 4 hr., and replaced in the desiccator. Certified reagent grade sulfanilic acid monohydrate from Fisher Scientific Co. was recrystallized from water and dried in an oven a t 100-105" for 2-3 hr., ground, and returned to the oven for 2-3 hr. The resulting sulfanilic acid (p-n'H2CsH4S03H)was stored for a t least 1 week in a desiccator with 98y0 H2S04 before analysis and use. Titration with standard NaOH indicated a purity of 99.9%. Taurine, NH2CH2CH2S03H,received from Fisher Scientific Co. as Highest Purity material, was recrystallized twice from water. Kjeldahl nitrogen analysis indicated that the recrystallized material contained ll.Ol~o nitrogen (theoretical: ll.19yo).
2245
errors, and extrapolation to infinite. dilution) is less than *0.06 kcal. mole-'. The e.m.f. studies of King led to K values at several temperatures from which he dalculated that AH" = 10.00 or 10.04 kcal. mole-'. Combination of our AH' with King's AGO = 12.36 kcal. mole-' gives AS" = -7.95 cal. (deg. mole)-'. King reports AS" = -7.9 cal. (deg. mole)-'.
Table I : Heats of Neutralization of Aqueous Taurine by Aqueous NaOH Mole0 of taurine/950 ml. of Hz0
- AHn, kcal. mole-'
0.009655 0,013930 0.016635 0.020431 0.021865
3.34 3.36 3.42 3.39 3.35
Heats of solution of taurine in water are reported in Table 11. The solubility of taurine has been reported6 to be 0.837 m, but the composition of the solid phase in equilibrium with saturated solution was not specified. Our Kjeldahl nitrogen determination indicates that the solid phase is anhydrous taurine. We have estimated the activity coefficient of aqueous taurine in saturated solution to be 1.32 by analogy with the data for betaine Results and Discussion reported by Smith and Smith.' For the standard free King3 has determined pK = 9.0614 for ionization of energy of solution we calculate AGO = -0.061 kcal. aqueous taurine as indicated by mole-'. This quantity is combined with our AHo = +NH3CK2CHzSO3-(aq)= 5.78 kcal. mole-' to yield AS" = 19.6 cal. (deg. mole)-' for the standard entropy of solution. From this ASo H+(aq) N H z C H ~ C H ~ S O ~ - (1) (~~) and the standard entropy of crystalline taurine,* we Heats of ionization were determined by measuring the calculate S2" = 56.4 cal. (deg. mole)-' for the standard heats of reaction of aqueous taurine (HT) with aquepartial molal entropy of aqueous taurine. ous NaOH to yield water and the anion of taurine The preceding entropy of aqueous taurine can be com(T-). Samples (10-ml.) of 5.15 M NaOH were mixed bined with our ASo = -7.95 cal. (deg. mole)-' for with 950 ml. of HzO containing a known amount of disionization of aqueous taurine as in eq. 1 to yield 32" = solved HT. A reaction equation for this process is 48.4 cal. (deg. mole)-' for the standard partial molal entropy of the aqueous anion of taurine, n'H2CH2CH2HT(aq) OH-(concd) = T-(aq) HzO(l) (2) SO3-(aq). Separate determination of the total heat associated with The entropy of ionization of aqueous taurine is conbreaking the bulb containing the concentrated XaOH siderably less negative than is characteristic of acids and diluting the NaOH (our values in agreement with literature values) permitted us to calculate heats for
+
+
HT(aq)
+ OH-(aq)
+
=
T-(aq)
+ HzO(1)
(3)
These values, designated AHn, are listed in Table I. Combination of AH," = -3.35 kcal. mole-' with AH" = 13.34 kcal. mole-' for the heat of ionization of water4J leads to AH' = 9.99 kcal. mole-' for the ionization reaction (1). We estimate that the total uncertainty (from possible sample impurities, calorimetric
(2) W. C. Fernelius, Znorg. Syn.,2, 178 (1946). (3) E. J. King, J . A m . Chem. SOC.,7 5 , 2204 (1953). (4) J. Hale, R . M . Iaatt, and J. J. Christensen, J . P h y s . Chem., 6 7 , 2605 (1963). (5) C. E. Vanderaee and J. A . Swanson, ibid., 6 7 , 2608 (1963). (6) J. B. Dalton and C. L. A. Schmidt, J . Biol. Chem., 109, 241 (1935). (7) E. R.B. Smith and P. K . Smith, i b i d . , 117, 209 (1937). (8) H . M. Huffman and S. W. Fox, J . A m . Chem. Soc., 62, 3464 (1940).
Volume 69,Number 7
July 1966
2246
HARRYP. HOPKINS,JR.,CHING-HSIEN Wu, AND LORENG. HEPLER
which is rearranged to
Table 11: Heats of Solution of Taurine
AHobsd
Moles of taurine/950 ml. of H20
A H , k e d . mole-1
0.007166 0.009470 0.009608 0.010790 0.012951 0.019194
5.78 5.78 5.72 5.74 5.72 5.75 AH'
= 5 . 7 8 f 0.05
that exist in a neutral rather than zwitterion form in solution, in agreement with King's3 argument that the ionization process should be described as in eq. 1. Calorimetric determination of the heat of ionization of aqueous sulfamic acid is necessarily less direct than for weaker acids such as taurine since solutions of sulfamic acid in water contain nonnegligible amounts of hydrogen and sulfamate ions as well as un-ionized sulfamic acid. One approach to the desired AH' of ionization makes use of previously reported heats of solution of crystalline sulfamic acid in water while another approach makes use of those data and our newly determined heats of solution of sulfamic acid in aqueous sodium hydroxide. When crystalline sulfamic acid is dissolved in water, processes represented by both of the following equations take place HS(c) HS(c)
=
=
HS(m')
H+(m)
AH4
+ S-(m)
(4)
AH5
(5)
where HS and S-represent sulfamic acid and sulfamate ion and m and m i indicate their concentrations. The previously reportedg heats of solution of sulfamic acid can be expressed as AHobsd
= aAH5
(I - CY)AH~
(6)
where a represents the degree of ionization calculated from the known amount of acid dissolved in a known amount of water and the ionization constant reported by King and King.'O Activity coefficients used in this calculation are subsequently discussed. In order to use eq. 6 it is necessary to combine eq. H+(m)
+ S-(m)
=
H+(aq)
+ S-(aq)
AHl (7)
7 and AH, with eq. 5 and AHSto obtain HS(c)
=
H+(aq)
+ S-(aq)
The Journal of Physical Chemistry
=
AH~O+ (*AH40CY
AHs
(8)
- a)~H4
(9)
In eq. 9 we have written AHaO to emphasize that this is a standard heat of solution and have assumed that the heat of dilution of the neutral HS(aq) in dilute solution is zero so that AH4 = AH~O. Heats of dilution represented by AH, have been taken to be equal to the already known heats of dilution of aqueous hydrochloric acid at the same concentrations." We, have followed Icing and Kinglo in taking 6 = 3.85 8. in the extended Debye-Huckel equation and have used the activity coefficients calculated in this fashion with the ionization constant also reported by King and King for evaluating by successive approximations the values of CY needed for use in eq. 9. A graph of (AHobsd CYAH,)/CY us. (1 - a ) / a has been constructed and the slope and intercept used to evaluate AH," = 4.53 and AH4' = 4.33 kcal. mole-'. The same calculation has also been carried out with activity coefficients of hydrogen and sulfamate ions taken equal to those of aqueous HCl at the same ionic strength leading to AHs' = 4.52 and AHlO = 4.35 kcal. mole-'. Combination of AH*' and AH,' gives AHlo' = 0.19 kcal. mole-' for the heat of ionization as in
+
HS(aq)
=
H+(aq)
+ S-(aq)
(10)
Another method of determining the heat of ionization of aqueous sulfamic acid has made use of the data in Table I11 for heat of reaction of crystalline sulfamic acid with aqueous NaOH represented by AHll and eq. 11. In making use of these data, we also use the preHS(c)
+ OH-(aq)
=
S-(aq)
+ H20
AHll (11)
viously reportedg heats of solution of sulfamic acid in water (already designated AHobsd in this paper) and heats of ionization of water (AH,) at ionic strengths corresponding to those of the solutions in our experim e n t ~ . Combining ~ these AH quantities gives us
in which a-values are those appropriate to the solutions that would have been obtained had the amounts of sulfamic acid listed in Table I11 been dissolved in 960 OH-(m) = ml. of water, and AH,,, refers to HS(aq) S-(m') H20.
+
+
in which (as) indicates that the preceding species are a t infinite dilution. Equation 6 now becomes AHobsa = CY(AH~ - AH,) f (1
-I- ~ Y A H I a
~~~~~
(9) C.-H. Wu and L. G. Hepler, J . Chem. Eng. Data,7, 536 (1962). (10) E. J. King and G. W. King, J . Am. Chem. Soc.. 74, 1212 (1952). (11) "Selected Values of Chemical Thermodynamic Properties," National Bureau of Standards, Circular 500, U. S. Government Printing Office, Washington, D. C., 1952.
2247
THERMOCHEMISTRY OF AQUEOUS AMINOSULFONIC ACIDS
Table I11 : Heats of Solution of Sulfamic Acid in 0.050 M NaOH
Table IV : Heats of Solution of Sulfanilic Acid in Water
Moles of sulfamic acid/960 ml. of Holn
AH'', kcal. mole-1
AHn, kcal. mole-'
0.008126 0.006472 0.03126 0.02791 0.02613
-8.94 -8.96 -8.85 -8.87 -8.84
-13.28 -13.35 -13.26 -13.40 -13.21
The simplest treatment, and all that is justified by the accuracy of the AH, values listed in Table 111, is to take an average A H , = -13.30 (av. dev. = 0.08) kcal. mole-' a t an average ionic strength of about 0.05 combined with A H , = 13.44 kcal. mole-' at this same ionic strength to give ilHlo = 0.14 kcal. mole-' for the heat of ionization of aqueous sulfamic acid. Consideration of the possible errors in experimental heats (ref. 9 and Table 111) and errors introduced by uncertainties in values of a (based on K and calculated activity coefficients) indicates that our first value (AHlo' = 0.19 kcal. inole-') is more certain than our second (AHloO = 0.14 kcal. mole-') for ionization of aqueous sulfamic acid. King and Kinglo calculated AH" = 0.27 and 0.46 kcal. mole-' from values of K a t different temperatures based on two sets of activity coefficients. Precise analysis of uncertainties in these four AHlo' values (which are not entirely independent of each other) is impossible, but it appears reasonable to select AHloO = 0.25 f 0.15 kcal. mole-'. Combination of this heat of ionization with the average free energy of ionization (AC&o0 = 1.38 kcal. mole-') reported by King and Kinglo leads to the entropy of ionization Afi'loO = -3.8 cal. (deg. mole)-'. This small negative value of the entropy of ionization provides further evidence that undissociated sulfamic acid is largely in the zwitterion form in aqueous solution. The heat of ionization of sulfanilic acid has also been calculated from calorimetric data. Heats of solution of sulfanilic acid in water are listed in Table IV. These data have been used as in eq. 9 for sulfamic acid in calculating AH" = 3.6 kcal. mole-' for the solution of sulfanilic acid to yield un-ionized aqueous sulfanilic acid and AH" = 7.6 kcal. mole-' for solution to give ions as in eq. 4 and 8 already written for sulfamic acid. Combination of these heats gives A H o = 4.0 kcal. mole-' for ionization of aqueous sulfanilic acid. A more accurate evaluation of the heat of ionization of aqueous sulfanilic acid has been obtained from combination of the heats listed in Table I V with those for solution of sulfanilic acid in aqueous NaOH listed in Table V. These data have been treated by an equa-
Moles of sulfanilic acid/950 ml.
A H , kcal. mole-'
0.005234 0.005947 0.006441 0,008483 0,009376 0.01127 0.01617 0.01984
4.72 4 62 4.60 4 52 4 47 4 41 4 35 4.26
tion like (12) and have yielded the A H , values also listed in Table V. The ionization constant at 25" reported by JZacLaren and Swinehart12has been used in this calculation and also in the calculation discussed in the preceding paragraph.
Table V : Heats of Solution of Sulfanilic Acid in 0.05 M NaOH Moles of sulfanilic ac,d,960 ml of soin,
0 008150 0 01089 0 01275 0 01440 0.01722
kcal mole-l
-5 65 -5 63 - 5 61 - 5 64 -5.63
AHn kcal mole-l
-9 10 -9 12 -9 16 -9 18 -9.18
Although there appears to be a dilution effect in the values of A H , listed in Table V, we are not at all sure this effect is real since uncertainties in A H , increase as concentration decreases owing to the increasing value of a and the decreasing magnitude of the actual calorimetrically measured heat. We therefore take an average value of A H , = -9.15 kcal. mole-' and combine with the heat of ionization of water4 at the same average ionic strength to obtain AH = 4.3 kcal. mole-' for ionization of aqueous sulfanilic acid. Estimation of heats of dilution and combination of the resulting AH" = 4.2 kcal. mole-' with that of the less accurately known value given earlier in this paper indicates that our "best" value for the heat of ionization is AH" = 4.2 f 0.2 kcal. mole-'. JIacLaren and Swinehartl2 have reported an equation for log K as a function of temperature, from which we have calculated AH" = 4.29 kcal. mole-'. We are unable to estimate the uncertainty to be associated with this value. Combination of our A H o = 4.2 kcal. mole-l with the (12) R. 0. MacLaren and D. F. Swinehart, J . Am. Chem. SOC.,73, 1822 (1951).
Volume 69. A'umber 7 J u l y 1965
MARYL. KILPATRICK, MARTINKILPATRICK, AND JOHNG. JONES
2248
Table VI : Thermodynamics of Ionization in Aqueous Solution at 298'K. AGO,
ASo, cal./
mole mole
mole)
AHor kca1.l kcal./
+HtNSOs-
-
H+
+ HzNSOa-
1.38 0.25 -3.8
4.41 4 . 2 +HaNCtH,SOI-
- E++
(deg.
HrNCyHd3Oa-
12.36 9 . 9 9
-0.7
-7.95
standard free energy of ionization, AGO = 4.41 kcal. mole-', calculated from the reported12 ionization con-
stant a t 25", gives AS" = -0.7 cal. (deg. mole)-' for the standard entropy of ionization of aqueous sulfanilic acid. This small negative entropy value provides supporting evidence for the conclusion of MacLaren and SwinehartI2 that sulfanilic acid in aqueous solution is in the zwitterion form. A summary of thermodynamics of ionization of sulfamic acid, sulfanilic acid, and taurine is given in Table
VI. Acknowledgments. We are grateful to the National Institutes of Health for financial support of this research and to Professor E. J. King for his helpful comments.
The Nitration of Nitrobenzene with Nitronium Fluoroborate in Hydrogen Fluoride
by Mary L. Kilpatrick, Martin Kilpatrick, and John G. Jones1 Department of Chemistry, Illinois Institute of Technology, Chicago, Illinois (Received December 28, 1964)
60616
In this paper, the nitration of nitrobenzene by solutions of nitronium fluoroborate in hydrogen fluoride is shown to be of the first order with respect to both reactants. The rate constant is enhanced as the ionic strength of the medium rises, and an equation describing this behavior can be based on the Brfinsted equation. Activation parameters indicate that the slowness of the nitration compared with its counterpart in sulfuric acid as solvent is due to a higher activation energy not entirely offset by a more favorable entropy of activation. This is explained in terms of solvation.
The investigation here described was begun in 1958, by which tiine several nitroniuni salts were known. I t was thought a good idea to employ one of these in some ionizing solvent, so as to measure the rate and order of attack of the nitroriiuni ion itself on an aromatic substrate without having to take into account the pre-equilibria which produce nitronium ions from nitric3 acid.3 If a suitable salt-solvent systeni were chosen, one would be able to weigh out nitronium ions directly into solution. Hydrogen fluoride was chosen as the ionizing solvent because of its low viscosity4 and high dielectric The Journal of Physical Chemistry
c o n ~ t a n t . ~Handling experience had already been gained16and it was known to be a good solvent for (1) Taken in part from the Ph.D. Thesis of J. G. Jones, Illinois Institute of Technology, June 1964. (2) (a) D. R. Goddard, E. D. Hughes, and C. K. Ingold, Nature, 158, 480 (1946); (b) A. H. Woolf and E. J. Emelbus, J . Chem. SOC.,1050 (1950); (c) D. R. Goddard, E. D. Hughes, and C. K. Ingold, ibid., 2559 (1950) ; (d) M. Schmeisser and S. Elischer, 2.Naturforsch., 7b, 583 (1952); (e) E. E. Aynsley, G. Hetherington, and P. L. Robinson, J . Chem. SOC.,1119 (1954). (3) (a) G. M. Bennett, J. C. D. Brand, and G. Williams. ibid., 869 (1946); (b) R. A. Marcus and J. M. Fresco, J. Chem. Phys., 27, 564 (1957); (c) F. G. Bordwell and E. W. Garbisch, J . Am. Chem. Soc., 82, 3588 (1980).
