Equilibrium constants for a-complex formation between cyanide ion and 1,3,5-trinitrobenzene in alcoholic solvents

Transcription

1 Equilibrium constants for a-complex formation between cyanide ion and 1,3,5-trinitrobenzene in alcoholic solvents E. BUNCEL, A. R. NORRIS, W. PROUDLOCK, AND K. E. RUSSELL Departtnet~t of Chemistry, Queen's Utziuersity, Kingstoti, Ontario Received June 6, 1969 Equilibrium constants have been determined spectrophotometrically for the reaction of cyanide ion with 1,3,5-trinitrobenzene in methanol, ethanol, 11- and iso-propanol, and n- and t-butanol. The equilibrium constants at 25 'C vary from 39 1 mole-' for the reaction in methanol to mole-' with t-butanol as solvent. Enthalpy changes, determined fromequilibriummeasurements and by direct calorimetry, vary from 0 to kcal mole-' and the calculated entropy changes decrease from 7 to -26 cal deg-' mole-' as the solvent is changed from methanol to t-butanol. These results are interpreted on the basis that desolvation of the small cyanide ion is the primary factor influencing enthalpies and entropies of reaction. The equilibrium constant in t-butanol is comparable to the values observed for aprotic solvents such as chloroform and acetone. Canadian Journal of Chemistry, 47, 4129 (1969) Introduction Ex~erimental Materials The reaction of cyanide ion with 1,395-trinitro- 1, 3, 5 - ~ ~ i ~ was i t ~ ~ b ~ ~ ~ twice ~ ~ from ~ benzene in chloroform gives rise to a deep red ethanol and dried in vacua, solution which is fairly stable at low temperatures Tetraphenylarsonium cyanide was prepared by adding (1). Proton magnetic resonance (p.m.r.) studies a concentrated solution of sodium cyanide to an aqueous solution of tetraphenylarsonium chloride. The resulting (2, 3) establish the structure of the red product as tetraphenylarsonium cyanide was recrystallized twice from the a-complex 1 formed as in reaction [I ] water in the presence of sodium cyanide, care being taken that the solution temperature did not exceed 60 "c. The H CN compound was dried in vacuo and stored over silica gel O N Methanol and ethanol were dried over Drierite and [I I 0.,, fractionally distilled in the absence of atmospheric mois- The complex has absorption maxima in chloroform solution at 442 mp (E = 3.98 x lo4 1 mole-' cm-') and 558 mp (E = 2.24 x lo4 1 mole-' cm- '). Red solutions with similar absorption characteristics are observed in a variety of other solvents and it is considered that the a- complex 1 is formed in all cases (1, 4). The equilibrium constant in chloroform solution at 25 "Cis 3.35 x lo5 1 mole-' and in acetone solution 1.44 x lo5 1 mole-' (1). In this study we have investigated the effect of alcoholic solvents on the equilibrium constant and on the enthalpy and entropy changes in reaction [I]. The equilibrium constant for the reaction in t-butanol is very large and it is not feasible to estimate the enthalpy change from the variation of equilibrium constant with temperature. Direct calorimetric measurements of enthalpy change have therefore been made, using a solution calorimeter. ture. n-propanol, iso-propanol, tz-butanol, and t-butanol were dried over Drierite and passed through a column of neutral alumina. They were fractionally distilled from Linde 4A molecular sieves. Procedure Solutions of tetraphenylarsonium cyanide and 1,3,5- trinitrobenzene were made up in an atmosphere of dry nitrogen. All flasks were sealed with silicone rubber serum caps and wrapped in aluminum foil to protect the solutions from light. In a typical experiment, 5 ml of 1,3,5- trinitrobenzene solution was placed in a quartz cell of 2 cm path length and the cell stoppered with teflon stoppers. The solution was allowed to reach the temperature of the thermostated cell compartment in the spectrophotometer and ml solution of tetraphenylarsonium cyanide added by means of a calibrated "Repette" syringe. Spectra were determined using a Unicam SP800 spectrophotometer or occasionally Beckman DKI or DU spectrophotometers. The cell temperature was controlled to better than ko.5 "C in the temperature range 545 "C. Condensation on cell windows at temperatures below 25 'C was prevented by purging with dry air or nitrogen. Calorimetric measurements were made using a Guild Corporation Solution Calorimeter (Model 400) equipped with a base line compensator. Solvent (215 ml) was placed in the Dewar flask (capacity 250 ml). Normally

2 4130 CANADIAN JOURNAL OF CHEMISTRY. VOL. 47, g 1,3,5-trinitrobenzene was weighed into a tube whose base consisted of a thin-walled glass bulb, and the tube and breaker inserted through the teflon cover of the Dewar flask so that the glass bulb was immersed in the solvent. After a period of not less than 1 11, 2 electrical calibrations.of the heat capacity of the system were made. The glass bulb was then broken and the heat of solution of 1,3,5-trinitrobenzene in the solvent determined. To the resulting solution was added g tetraphenylarsonium cyanide using the same procedure. In some experiments the order of addition was reversed, the heat of solution of tetraphenylarsonium cyanide being determined in the first step. The molar ratio of 1,3,5-trinitrobenzene to tetraphenylarsonium cyanide was always greater than 4:l. The initial temperature of the solvent in the calorimetric studies was close to 25 "C with all solvents except t-butanol for which the initial temperature was about 26.5 "C. Results Preliminary spectroscopic studies were performed with ethanol as solvent. The concentration of 1,3,5-trinitrobenzene was low (about M) and a large excess of tetraphenylarsonium cyanide was added to give solutions which were to loe2 M in cyanide ion. A red color developed but the color faded rapidly to pale yellow even in the dark and no meaningful absorption spectrum could be obtained. Addition of a small amount of tetraphenylarsonium cyanide to a large excess of 1,3,5-trinitrobenzene in ethanol solution gave a red solution whose absorbance at 435 mu reached a maximum after a few minutes. The absorbance did not change significantly over a period of min. The visible spectrum showed maxima at 435 and 544 mp, the absorbance at 435 mp being nearly twice that at 544 mp (E~,,/E,,, = 1.76 for ethanol). Similar results were obtained with other alcohols; the wavelengths of the absorption maxima are given in Table I. The spectru each solvent is almost identical with that of he red substance formed in chloroform or acetone solution which has been shown to be the D-com- TABLE I Wavelengths of absorption maxima of cr-complex 1 in various alcohols at 25 "C "ln plex 1 (2, 3). The 1:1 stoichiometry has been confirmed in the case of n-propanol as solvent by the method of continuous variations. Equilibrium constants were determined by analysis of absorbance measurements at the short wavelength maximum for solutions which were initially 2-7 x lop5 M in tetraphenylarsonium cyanide and M in 1,3,5-trinitrobenzene (TNB). For the reaction [2 I TNB f CN- =$ TNBCN- the equilibrium constant [TNB. CN-] K= [TNB] [CN -1 If the absorbance at equilibrium in the 2 cm cell is given by A and the extinction coefficient for this wavelength is E, it follows that where [CN-lo, [TNB], are the initial concentrations of cyanide ion and 1,3,5-trinitrobenzene. Following the Rose-Drago method of analysis (9, plots were made of 1/K vs. assumed values of E for each set of initial concentrations of cyanide ion and 1,3,5-trinitrobenzene (Fig. I). Values of K and E were obtained from the points of intersection of the lines and error limits were based on average deviations from the mean of individual estimates of K and E. Equilibrium constants were in some cases also determined from absorbances hm,, hm,, Solvent (m~) Ow) Methanol Ethanol n-propanol iso-propanol n-butanol t-butanol FIG. 1. Rose-Drago plot for the reaction of cyanide ion with 1,3,5-trinitrobenzene in ethanol at 25 "C. K = 1265 t 20 1 mole-'; E = t mole-' cm-'.

3 BUNCEL ET AL.: EQUILlBRIUM CONSTANTS FOR a-complex FORMATION 4131 at the long wavelength maximum and these agreed within experimental error with values of K obtained at the short wavelength maximum. The equilibrium constants K of reaction [l ] at 25 "C and the extinction coefficients of the (3- complex at the lower wavelength maxima for the solvents methanol, ethanol, n-propanol, isopropanol, n-butanol, and t-butanol are given in Table 11. With methanol as solvent the low equilibrium constant results in a relatively high concentration of free cyanide, with the consequence TABLE I1 Equilibrium constants K for reaction [l] at 25 "C and extinction coefficients at lower wavelength maximum 10-~ & Solvent K (1 mole-') (1 mole-' cm-') Methanol Ethanol npropanol iso-propanol n-butanol t-butanol that at 25 "C the maximum absorbance begins to decrease significantly after a period of only 2 or 3 min. The rate offadingwaslower at 6 and 15 "C; maximum absorbance values were constant for over 5 min at these temperatures. Rose-Drago plots at 25 "C gave reproducible values of K, but the value of mole-' is probably less than the true value because of the'fade reaction. The maximum absorbances for reactions in ethanol, n-propanol, iso-propanol, and n-butanol are not subject to significant fading under the chosen experimental conditions and the equilibrium constants are hence assigned narrow limits of error. The equilibrium constant for the reaction in t-butanol is very large and its value is probably not known to better than +20% because complexing occurs to a large extent over the whole range of 1,3,5-trinitrobenzene concentrations employed. Enthalpies of reaction have been obtained from the variation of equilibrium constants with temperature over the range 6-25 "C with methanol as solvent, "C with ethanol, and "C with n-propanol and n-butanol. The values are recorded as AH,,' in Table 111. Enthalpies of reaction determined calorimetrically AHcal are also given in Table 111. The fraction of cyanide ion converted to complex varied from 83 to close to 100 % dependent on solvent and reaction conditions and the calorimetric enthal~ies of reaction are based on the amount of complex formed calculated from the equilibrium constants of Table 11. The maximum concentration of tetraphenvlar- - sonium cyanide used in the calorimetric measurements was 2 x low3 M and the ratio of concentrations of 1,3,5-trinitrobenzene to cyanide ion was always at least 4: 1 ; even so, some difficulties were experienced because of secondary reactions. With ethanol as solvent, for example, the temperature of the reaction system continued to rise slowly after equilibrium should have effectively been reached and the measured heat of reaction was dependent on the concentration of free cyanide ion in the final solution. The AHcal for ethanol was obtained by extrapolating the experimental values of AH to zero concentration of free cyanide ion. The reproducibility of individual values of AH was better than +2% and the larger limits of error are assigned in Table I11 because of the difficulties with secondary reactions. No solubility problems were encountered with the first 5 alcohols listed in Table 111. With t-butanol as solvent, however, the 1,3,5-trinitrobenzene dissolved slowly and precipitation of the TABLE I11 Thermodynamic data for formation of a-complex 1 at 25 "C -AGO -AH,,"* -AHca, - AS0 Solvent (kcal mole-') (kcal mole-') (kcal mole-') (cal deg-' mole-') -- Methanol k5 Ethanol k n-propanol k k3 iso-propanol Butanol ' ' k3 t-butanol O+_ 3 *AH 0 in methanol was calculated using equilibrium constants determined at 6 15 and 25 'C. All other AH.,Q values were ca1cula;;d using equilibrium constants determined at 25, 35, and 45 'C. The ~argkr &ror associated with AH.," in methanol reflects the greater uncertainty in the equilibrium constants in this solvent due to the presence of a slow secondary reaction.

4 4132 CANADIAN JOURNAL OF CHEMISTRY. VOL. 47, 1969 red complex occurred when the initial cyanide concentration was greater than 1.5 x M. Free energy changes, AGO, and entropy changes, AS0, for reaction [l] in the various alcoholic solvents are also given in Table 111. The standard state is 1 mole 1- '. Discussion Since the concentration of tetraphenylarsonium cyanide does not exceed M, it is probable that ion pair formation between the bulky tetraphenylarsonium ion and the cyanide ion is small in the alcohols studied. The ecluilibrium measurements are therefore considered to apply to the reaction between free cyanide ion and 1,3,5-trinitrobenzene. Differences of solvation of reactants and products may vary considerably from one solvent to another. The cyanide ion is small with a relatively localized negative charge and it should be strongly solvated by polar protic solvents such as methanol and ethanol (6). The charge on the o-complex is dispersed and hydrogen bonding between the complex and polar protic solvents is probably fairly weak. Formation of the o-complex thus requires desolvation of cyanide ion to occur and this is likely to be a major factor in determining enthalpy and entropy changes for reaction [I ] in methanol. At the other extreme, t-butanol solvates the cyanide ion only to a limited extent and solvation effects are likelv to be of minor importance. The highest equilibrium constant observed is 5 x lo5 1 mole-' for reaction [l] in t-butanol; there is a large negative enthalpy change and the effect of the enthalpy term on the equilibrium constant is only partially compensated by the large entropy decrease. Differences in solvation of reactants and products are expected to be small with t-butanol as solvent. It might therefore be anticipated that the reaction of cyanide ion with 1,3,5-trinitrobenzene to give the o-complex in a weakly interacting solvent would be characterized by a high equilibrium constant, a large enthalpy of reaction (-- 15 kcal mole-') and a large negative entropy change ( cal deg- ' mole-'). Values of AH and AS of this order of magnitude are associated with o-complex formation in chloroform and acetone'. The other alcoholic solvents possess OH groups which are considerably less shielded by 'A. R. Norris. Unpublished results. methyl groups and solvation by these alcohols has a considerable effect on the equilibrium constant. The equilibrium constant in methanol at 25 "C is some times smaller than in t-butanol; for the former, the enthalpy change is approximately zero and the entropy change is small and positive. The most important reason for the striking behavior of methanol as solvent is the solvation of the small cyanide ion. Apparently a large positive enthalpy change accompanies the reduction in solvation as the o-complex is formed. There is a simultaneous decrease in order, and, although methanol has some structure, the overall result is a much more positive entropy change for the reaction in methanol than in t-butanol. The equilibrium constants at 25 "C in ethanol, n-propanol, and n-butanol are all in the range mole- l at 25 "C. The heats of reaction are kcal mole- ' and the entropy changes - 9 +_ 3 cal deg mole- '. They thus show intermediate behavior between the 2 extremes of methanol and t-butanol. Desolvation of the cyanide ion appears to be the main factor causing a decrease in equilibrium constant from the value in t-butanol, but the effect is smaller than in methanol and there is a slight trend towards 2. larger equilibrium constants as the size of the alkyl group is increased. As judged by the equilibrium constant, the solvating effect of iso-propan01 is significantly weaker than that of the normal alcohols. It was anticipated that the equilibrium constant of mole-' would be associated with a heat of reaction of about -9 kcal mole- '. The lower observed value was reproducible in calorimetric studies but secondary reactions at higher temperatures prevented confirmation of this heat of reaction in equilibrium studies. The enthalpy and entropy changes recorded here for the reaction of cyanide ion with 1,3,5- trinitrobenzene in ethanol may be compared with an enthalpy change of 0.3 kcal mole-' and an entropy change of 16.3 cal deg-' mole-' for the formation of the cr-complex between ethoxide ion and 1,3,5-trinitrobenzene in ethanol (7). The larger entropy change in the latter case may imply that desolvation of ethoxide ion in the formation of its o-complex with 1,3,5-trinitrobenzene is greater than with cyanide ion. Ethoxide ion attacks 2,4,6-trinitrotoluene under some conditions to form almost exclusively the trinitro-

5 BUNCEL ET AL.: EQUILlBRlUM CONSTANTS FOR a-complex FORMATION benzyl anion (8). 7H3 OEt- -t * O ~ N ~ N 4- OEtOH Z The entropy change in this reaction is 27 cal deg mole-', and this large value has been interpreted by Caldin as being largely due to desolvation of ethoxide ion (7). We have attempted to determine whether any relationship exists between the observed solvent effects on reaction [l] and the various empirical solvent polarity parameters (9). Figure 2 shows a plot of -AGO for reaction [I] at 25 "C against ET which measures the energy of the intramolecular charge transfer transition in a pyridinium FIG. 2. Plot of -AGO for reaction of cyanide ion with 1,3,5-trinitrobenzene in various alcohols at 25 'C vs. the solvent parameter ET. N-phenol betaine (10). It appears that the results for cr-complex formation in the 6 alcohol solvents are correlated well by the parameter E,. The closeness of the correlation may be in part due to the fact that these solvents are all of the same general type. The significance of such correlations and the extension to other systems are being investigated. Acknowledgments The authors are grateful to the National Research Council of Canada and the Defence Research Board of Canada for financial support of this work. The assistance of Mrs. J. Scouten in some of the calorimetric studies is gratefully acknowledged. 1. A. R. NORRIS. Can. J. Chern. 45, 2703 (1967). 2. E. BUNCEL, A. R. NORRIS, and W. PROUDLOCK. Can. J. Chern. 46, 2759 (1968). 3. A. R. NORRIS. Can. J. Chern. 47, 2895 (1969). 4. E. BUNCEL, A. R. NORRIS, and K. E. RUSSELL. Quart. Rev. London, 22, 123 (1968). 5. N. J. ROSE and R. S. DRAGO. J. Arner. Chern. Soc. 81, 6138 (1959).- 6. R. A. ALEXANDER, E. C. F. KO, A. J. PARKER, and T. J. BROXTON. J. Arner. Chern. Soc. 90,5049 (1968). 7. E. F. CALDIN. J. Chem. Soc (1959). 8. E. F. CALDIN and G. LONG. Proc. Roy. Soc. London, Ser. A, 226, 263 (1955). 9. C. REICHHARDT. Angew. Chern. Intern. Ed. Engl. 4, 29 (1965). 10. K. DIMROTH and C. REICHHARDT. Palette, No. 11, 28 (1962).