effects of solvent polarity on the acid dissociation constants of benzoic acids

Effects of Solvent Polarity on the Acid Dissociation Constants of Benzoic Acids

JOSEPH T. R U B I N O ~ AND WILLIAM s. BERRYHILL Received August 14, 1985, from the School of Pharmacy, University of North Carolina at Chapel Hill, Chapel Hill, NC 27514. publication October 29, 1985.

Accepted for

Abstract 0 The pKa values ofbenzoic acid, pmethylbenzok, andp aminobenzoic acid (PABA) were determined by potentiometric titration in mixtures of 0-0.5 volume fractions of various cosolvents and water. The differences between the aqueous and semiaqueous plc, values were similar for the three solutes at a particular cosolvent-water mixture for most of the cosolvents studied. The largest differences occurred in the dirnethyl sulfoxide (Me2SO)-water system, where the plc, changes were larger for PABA than the other two solutes. The data are analyzed by a consideration of both electrostatic and nonelectrostatic medium effects. The electrostatic medium effect was calculated from the Born Equation while any residual pKa change was attributed to nonelectrostatic medium effects. The residual effects were found to correlate well with indexes of solvent hydrogen bond acceptor ability such as HBA and p- values. These results provide a rationale for the use of two solvent polarity indexes for more accurate estimates of pKa values of weak electrolytes in semiaqueous solvent systems. Analysis of solubility data of the salt and acid forms of PABA and benzoic acid in ethanol-water and Me,SO-water mixtures suggests that the higher activity of the anionic form of PABA in Me,SO-water mixtures is primarily responsible for the large pKa changes observed for that solute.

The importance of the effects of solvent composition on the dissociation of weak electrolytes has been recognized for a number of years. The poor aqueous solubility of many weak electrolyte drugs has necessitated the use of water-miscible cosolvents to prepare solutions of these compounds for potentiometric determination of the pK,. The advantages and disadvantages of this procedure have been reviewed by Benet and Goyan' and by Albert and Sejeanta2 Processes involving partition phenomena such as membrane transport, drug extraction, and chromatography are highly dependent upon the fraction of ionized and un-ionized drug present in solution. Drug stability3 and solubilitp may also be highly dependent on the fraction of ionized compound present in solution. The widespread use of cosolvents in liquid drug formulations often requires an understanding of the behavior of weak electrolytes in semiaqueous solvents. Thus, it is often desirable to predict the influence of solvent composition on the acid dissociation constant, K,, of weak electrolyte compounds.

The effect of solvent composition on the acid dissociation constant has been expressed in terms of the medium effect, myi, on each species in so1ution:6.E

(1)

where the subscripts w and s refer to water and the mixed solvent, and H+, A, and HA refer to the hydrogen ion, the dissociated, and the undissociated form of the drug, respectively. The medium effect is the ratio of the activity of each species in water to the mixed solvent and is therefore a measure of the free energy of transfer of each species from water to mixed solvent. Equation 1 can be divided into

&(w) - - mm+m3/A

K(S) m m -

electrostatic (el) and nonelectrostatic (n) terms and can be rewritten?

m?H+m?A

log m?HA

The electrostatic contribution to the medium effect represents the ability of the solvent to separate charged species in solution. It can be represented by the Born equation:'

(i -6) (3)

where N is Avogadro's number, e is the electronic charge, R is the gas constant, T is the temperature, Z is the charge on the individual species, r is the radius of each species, and E is the solvent dielectric constant. For a given solute in cosolvent water mixtures, 2, r, and eW will be constants and eq. 3 can be used as the basis for relating pK,(s) to 1/ee. Although the Born equation may be useful in describing such relation- ships in a particular cosolvent-water system, it may not be effective in relating pK&) to solvent polarity when cosolvents of different polarity and structure are considered to- gether.7

The nonelectrostatic medium effect represents the differences between the cohesive and adhesive forces between the solute and solvent species, assuming all solute species to be uncharged.8 Thus all species of the same size and bonding characteristics should have the same noneledrostatic medium effect in a given solvent. The concept is conceptually identical to nonelectrolyte solubility theory, e.g., regular solution theory. The primary difference is, of course, the fact that the charge on the solute species cannot be totally ignored. It has been suggested that ions will induce consider- able solvent structuring, however, this effect may be compar- atively minimal for large ions.8

The present study was conducted to determine the effects of electrostatic and nonelectrostatic solvent polarity indexes in the prediction of solute pK, in various semi-aqueous solvent systems. Benzoic acid and two of its derivatives, p-aminobenzoic acid (PABA) and p-methylbenzoic acid (toluic acid) were chosen to examine the effect of the solute substituent groups on the relative changes in pK, induced by the cosolvent. A previous study9 was concerned with the prediction of the solubilities of poorly water-soluble drugs in cosolventiwater mixtures from physicochemical parameters of the solvent

182 /Journal of Pharmaceutical Sciences Vol. 75, NO. 2, February 1986

0022-3549/86/02OO-O 1 82$0 l.OO/O 0 1986, American Pharmaceutical Association

mixtures, including the solubility parameter, 8, hydrogen bond donor group density, HBD, and hydrogen bond acceptor group density, HBA. These polarity indexes, as well as Kamlet and Taft pvalues10 were used to investigate the relationship between the nonelectrostatic medium effects, as defined in eq. 2, and solvent polarity.

Experimental Section Materials-The cosolvents were all reagent grade or higher.

These included ethyl alcohol (ETOH; AAPER Alcohol and Chemical Co., Shelbyville, KY); dimethyl sulfoxide (Me2SO; MCB Reagents, EM Science, Gibbstown, NJ); dimethylacetamide (DMA; Aldrich Chemical Co., Milwaukee, WI); propylene glycol (PG; Aldrich Chem- ical Co., Milwaukee, WI); polyethylene glycol 400 (PEG 400; Aldrich Chemical Co., Milwaukee, WI); and glycerol (Aldrich Chemical Co., Milwaukee, WI). Water was prepared by double distillation in a glass still and was deaerated under reduced pressure prior to use. Solvent mixtures were prepared as described previously11 by combining 10,20, 30, 40 or 50 mL of water with 90, 80, 70, 60, or 50 mL, respectively, of cosolvent. Titrant solutions were composed of 0.100 M potassium hydroxide (Sigma Chemical Co., St. Louis, MO) in the appropriate cosolvenbwater mixture. Benzoic acid (Fisher Scientific Co., Fair Lawn, NJ), PABA, as well BB their sodium salts (Aldrich Chemical Co., Milwaukee, "I), p-toluic acid (Sigma Chemical Co., St. Louis, MO), p-nitroaniline (Aldrich Chemical Co., Milwaukee, WI), p-nitrophenol (Aldrich Chemical Co., Milwaukee, WI), p-ni- troanimle (Aldrich Chemical Co., Milwaukee, WI), and NJV-diethyl p-nitroaniline (Frinton Labs, Vineland, NJ) were used as received from the suppliers.

Methods-The pK, values were determined by potentiometric titration using a modification of the methods described by Albert and Sejeant. The temperature of the benzoate solutions was maintained at 25 2 0.2"C by circulating water from a constant temperature water bath (Fisher Scientific Co., Fair Lawn, NJ) through 112 in. (1.27 cm) 0.d. Tygon tubing which was coiled around a 100-mL beaker. The tubing was held in place by duct tape and the assembly placed on a stirring hot plate which was used to agitate the solution during the titrations. A 1/2 in. (1.27 cm) thick piece of Styrofoam insulation was placed between the hot plate surface and the bottom of the beaker to provide insulation for this exposed area. All titrations were performed under a constant stream of nitrogen gae. A digital type pH meter (Fisher model 825 MP pH meter), glass pH electrode (Fisher Model #13-639-3 universal glass electrode), and calomel reference electrode (Fisher model X13-639-52 calomel reference electrode) were used. An automatic temperature probe (Fisher model X13-636-752 Automatic temperature probe) was used to continuously monitor the temperature throughout the titration. Titrations were performed in duplicate.

The pK, values were evaluated by the graphical technique described by Benet and Goyan.1 Six to twelve data points were used for each pK, determination. The "apparent" pH readings obtained during the titrations in mixed solvents were adjusted according to the method of Van Uitert and Haas.12 This involved the measure- ment of the pH of solutions of the mixed solvents containing 1.00 x lo-' M HCl. The difference between the observed readings and the expected value of 3.00 WBB used as the correction factor. The pK, values were further corrected for the "salt effect" using the Debye- Hiickle limiting law.

Kamlet and Taft fi values were determined as described by the authors.10 The methods were repeated for those solvents previously studied as well as those for which p values had not been reported. The frequency shifts for the p-nitroaniline versus N,N-diethyl-p- nitroaniline system were used in the determination of the P values, 88 this scale tended to be more sensitive to the structure of the solvents.

Values for the polarity indexes of solvent mixtures, P,, were estimated using a linear combination rule:

where P is the particular polarity scale being calculated, f is the volume fraction of cosolvent, and the subscripts c and w refer to cosolvent and water, respectively.

The values of HBA were calculated as described previously by the following formula: [(# nonbonding electron pairs)(density of solvent)/ molecular weight of solvent] x 1000. Molecular radii were estimated from the calculated molar volumes13 by assuming spherical mole- cules (H' = 1.22 A, benzoate = 3.46 A, PABA = 3.63 A, toluate = 3.78 A).

Solubilities were determined by equilibrating excess solute with solvent for 48 h in a constant temperature water bath maintained at 25 f 05°C. The samples were rotated end-over-end during this time period. The solute concentrations were determined spectrophotomet- rically (model DU-40 Beckman Instrument Co., Irvine, CA) after being diluted to an appropriate concentration by 0.1 M potassium hydroxide. Analysis of the titration data and the various correlations were performed using the Subprogram Regression found in SPSS- X."

Results and Discussion Table I lists the values of the correction factors used in the

adjustment of the apparent pH readings as well as the values of the polarity indexes for the neat solvents. These empirical correction factors are necessary when performing pH measurements in semiaqueous solvents using an electrode system possessing a liquid junction. The method was originally used for dioxane-water mixtures12 and has been validated for use in alcohol-water systems by Bates et al.16 A modification of the method has been used by Pashankov et al.16for pH measurements in acetonitrile-water mixtures. The primary assumption in the use of the empirical correction factors is that strong electrolytes a t dilute concentrations will be completely ionized in all the solvents studied. This assumption is considered valid when the solvent dielectric constant is above 39.12 Corrections for changes in the activity coeffi- cient of hydrogen ions in various solvents were considered negligible in these studies due to the low ionic strength and relatively high polarities of the solvents.

The pK, values for benzoic acid, p-toluic acid, and PABA are presented in Tables 11-N. Duplicate values were most often in agreement by 0.02 pKa units or less, except where standard errors are indicated in parentheses. The plots of 2 versus ZIH+ll showed good linearity in all solvents except DMA were a constant curvature in the data was persistent.

Table I-Values of Polarlty Indexes and PH Correctlon Factors of Solvents

pH Correction Factors

0.1 0.2 0.3 0.4 0.5

HBA Volume Fraction Cosolvents Solvent 8' 66 P

Water 78.5 23.4 Ethyl alcohol 24.3 12.7 Propylene glycol 32.0 12.6 Glycerol 42.5 17.7 Polyethylene glycol 400 13.6 11.3 Dimethylacetamide 37.8 10.8 Dimethyl sulfoxide 46.7 12.0

0.13 111.0 - - - - - 0.66 34.0 0.04 0.03 0.07 0.14 0.23 0.61 54.8 0.00 -0.04 -0.04 0.00 0.00 0.52 82.2 -0.05 -0.09 -0.16 -0.21 -0.28 0.62 50.8 0.03 0.13 0.27 0.50 0.81 0.79 32.3 0.27 0.44 0.61 0.78 0.95 0.70 28.2 0.08 0.12 0.24 0.36 0.50

'Taken from ref. 21. 'Taken from ref. 22,

Journal of Pharmaceutical Sciences I183 Vol. 75, No. 2, February 7986

Table IC-pK, Valuea, Electrosbtlc,m and Rerldualb ApK. Valuer tor Benzolc Acld In Cosolvent-Water Mlxturer

f PYS Ap&(obs) ApK,(el)' Ap&(resid) " 0.0 0.1 0.2 0.3 0.4 0.5

0.1 0.2 0.3 0.4 0.5

0.1 ' 0.2' 0.3' 0.4 0.5

0.1 0.2 0.3 0.4 0.5

0.1 0.2 0.3 0.4 0.5

0.1 0.2 0.3 0.4 0.5

4.17 4.29 4.49 4.81 5.16 5.47

4.25 4.41 4.62 4.86 5.16(.04)

4.64 4.76 4.86 5.03 5.28

4.29 4.43(.03) 4.69 5.00( .04) 5.28

4.21 4.31 4.49 4.61 4.81

4.35 4.63 4.89 5.27 5.66

Ethyl Alcohol 0.00 0.00 0.12 0.12 0.32 0.27 0.64 0.45 0.99 0.65 1.30 0.91

0.08 0.10 0.24 0.23 0.45 0.37 0.69 0.53 0.99 0.72

0.47 0.09 0.59 0.19 0.69 0.31 0.86 0.45 1.11 0.60

0.12 0.07 0.26 0.15 0.52 0.23 0.83 0.33 1.11 0.43

0.04 0.08 0.14 0.17 0.32 0.27 0.44 0.38 0.64 0.51

0.18 0.15 0.46 0.34 0.72 0.57 1.10 0.85 1.49 1.21

Propylene Glycol

Dimethylacetamide

Dimethyl Sulfoxide

Glycerol

Polyethylene Glycol

0.00 0.00 0.05 0.19 0.34 0.39

-0.02 0.01 0.08 0.16 0.27

0.38 0.40 0.38 0.41 0.51

0.05 0.1 1 0.29 0.50 0.68

-0.04 -0.03

0.05 0.06 0.13

0.03 0.12 0.15 0.25 0.28

~~~~ ~ ~ ~~ ~ ~ ~ ~ ~ ~

'See eq. 3. "Observed pK, - calculated pK,. 'Data not included in regression analyses in Table V.

Table IlCpK. Valuer and Electroatatlc' and Realdualb ApK. Valuer for PABA In Corolvent-Water Mlxturea

f PK, Ap&(obs) APK'(W Ap&(resid) " Ethyl Alcohol

0.0 4.91 0.00 0.00 0.00 0.1 5.01 0.10 0.12 -0.02 0.2 5.23 0.32 0.27 0.05 0.3 5.55 0.64 0.44 0.20 0.4 5.93 1.02 0.65 0.37 0.5 6.26 1.35 0.90 0.45

0.1 4.97 0.06 0.10 -0.04 0.2 5.17 0.26 0.22 0.04 0.3 5.37 0.46 0.36 0.10 0.4 5.57 0.66 0.53 0.13 0.5 5.90 0.99 0.71 0.28

0.1 5.20 0.29 0.09 0.20 0.2 5.32 0.41 0.19 0.22 0.3 5.45 0.54 0.31 0.23 0.4 5.71 0.80 0.44 0.36 0.5 5.99 1.08 0.59 0.49

0.1 5.01 0.10 0.06 0.04 0.2 5.23 0.32 0.14 0.18 0.3 5.47 0.56 0.23 0.33 0.4 5.80 0.89 0.32 0.57 0.5 6.26 1.35 0.43 0.92

Propylene Glycol

Dimethylacetamide

Dimethyl Sulfoxide

'See eq. 3. "Observed p& - Calculated p&.

Table IV-pK. Valuer and Electrortalcm and Resldualb ApK. Valuer for pTolulc Acld In Corolvent-Water Ylxhrrer

f PK, Ap&(obs) Ap&(el)' Ap&(resid)" Ethyl Alcohol

0.0 4.41 0.00 0.00 0.00 0.1 4.49 0.08 0.12 -0.04 0.2 4.75 0.34 0.26 0.08 0.3 5.11 0.70 0.44 0.28 0.4 5.48 1.07 0.64 0.43 0.5 5.73 1.32 0.89 0.43

0.1 4.52 0.1 1 0.10 0.01 0.3 4.83 0.42 0.36 0.06 0.5 5.33 0.92 0.71 0.21

0.1 4.87 0.46 0.09 0.37 0.3' 5.04 0.63 0.30 0.33 0.5 5.45 1.04 0.59 0.45

0.1 4.50(.03) 0.09 0.06 0.03 0.3 4.86 0.45 0.23 0.22 0.5 5.52 1.11 0.42 0.69

regression analyses in Table V.

Pmpylene Glywl

DifnethyI8~et8fnide

Dimethyl Sulfoxide

'See eq. 3. "Obsewed p& - Calculated pK,. 'Data not included in

This most often occurred when titrating benzoic or toluic acid in 0.10-0.30 volume fractions cf, of DMA. The nonlinearity could not be attributed to the sources of error d immed by Albert and Serjeantl or Benet and Goyan2 which include errors in standardization of the pH meter, loss of solution during the titration, errors in titrant volume, etc. Repre- sentative plots are illustrated in Fig. 1 for benzoic acid in f = 0.5 cosolvent-water mixtures.

An additional verification of the methods used is provided by the PK, values of benzoic acid in Me2SO-water mixtures. The values obtained at f = 0.3 and 0.5 are in reasonably good agreement with the values determined by Fiordiponti et 81.17 at mole fractions of 0.1 cf- 0.32) and 0.2 cf- 0.5) Me2S0, i.e., 4.67 and 5.21, respectively. Their study was performed using an electrode system without transference.

Figure 2 illustrates the change in PK, from water to various volume fractions of cosolvents and water for the three solutes studied. In most cases the change in PK, is similar for the three solutes in a given cosolvent-water system. This is particularly true of the amphiprotic coeol- vent-water systems such as ETOH and PG (Fig. 2A). The most obvious differences occur in mixtures of aprotic cosolvents and water, MezSO and DMA (Fig. 2B). The lines in Fig. 2 were drawn using the benzoic and toluic acid data in order to illustrate these differences. The PK, change is generally greater for PABA in MeaO-water mixtures than for benzoic or toluic acid. This was not the case for the DMA-water mixtures where the change in PK, was much higher for toluic and benzoic acids than for PABA at f = 0.1-0.3 DMA. However, due to the uncertainty of the PK, values of benzoic and toluic acids at these solvent compositions the results should be considered cautiously. Nonetheless, the pKa values for PABA do not follow the same trends at f = 0.4 and 0.5 DMA as in the other aprotic codvent, MezSO. This ie unexpected due to the ability of DMA to strongly solvate amphiprotic so1utes.B In addition, DMA possesses a much lower dielectric constant than MezSO and, although both cosolvents might be expected to have approximately equal nonelectrostatic medium effects, the lower ability of DMA to separate charged species in solution should result in compar- atively higher changes in PK,. It should be noted that the empirical correction factors for the DMA-water mixtures are considerably higher than those used for the remaining coeol- vent-water systems. This may be indicative of significant

184 / Journel of Phanneceutical Sciences Vol. 75, No. 2, Februety 1986

10 L 8

12 ’I 0 2 4 6 8

W+l (x 109

Flgure l-Representative plots of Z versus Z[W] benzoic acid in f = 0.5 cosolvent-water mixtures. Key: (0) glycerol; (0) PEG 400; (A) PG; (m) DMA; (0) Me2SO; (A) ETOH.

changes in the activity of hydrogen ions or an adverse effect on the electrode operation in the DMA-water system. The use of these procedures in performing acid-base titrations in this solvent system should therefore be considered cautiously:

By combining eqs. 2 and 3 it can be suggested that those factors attributed to nonelectrostatic solvent effects can be examined from a knowledge of the observed medium effect and the solution of eq. 3. Thus, if it is assumed that the electrostatic effects of the solvent are described by the Born equation, the residual APK, between the observed values and those predicted from eq. 3 should be an estimate of the nonelectrostatic solvent effect. The observed, predicted, and residual APK, values are presented for each solute in Tables II-N. The residual values were examined for their relationship to solvent polarity indexes for each solute. The results of the residual ApKo, values versus the polarity index regression are presented in Table V. Good correlations were observed between those polarity indexes which reflect the hydrogen bond accepting ability of the mixed solvent, HBA and p. The use of an additional term describing the hydrogen bond donating ability was found not to improve the correlation to a significant extent. It was also observed that the use of the p” term provided an improved estimate of the nonelectrostatic medium effects. In general, the HBA scale was somewhat superior in predicting the nonelectrostatic effect than the p value. The greater significance of HBA and p in predicting the nonelectrostatic effects is most likely due to the greater prevalence of solvent-associated complexes which have been postulated to take place in semiaqueous and nonaqueous aolvents.18 The ability of the cosolvent to associate with both

Volume Fraction of the Cosolvent

Flgure 2-The PK, versus the volume fraction of cosolvent for (A) ETOttwater (open symbols) and PG-water (closed symbols); and (B) DMA-water (open symbols) and Me2SO-water (closed symbols). Key: (0, m) benzoic acid; (A, A) toluic acid; (0, 0) PABA.

the dissociated and undissociated forms of the solute would most frequently occur via hydrogen bonding, and the activity of each solute species would be most likely related to these properties of the solvent.

As these results suggest, the prediction of the PK, of a weak electrolyte in semiaqueous solvent mixtures can be estimated from a combination of the dielectric constant and the hydrogen bond acceptor terms such as HBA. Estimation of flu values in various mixed solvents can be achieved by calculation of the electrostatic medium effect using the Born equation and determination of the nonelectrostatic effect from the residual ApKu in one or more mixed solvents and water. Alternatively, two or more flu values determined in mixed solvents can be used to estimate the flu values in

Journal of Pharmaceutical Sciences / 18s Vol. 75, No. 2, February 1986

Table V-Summary of Regresslon Analyses Resldual p& versus Solvent Polarity Indexes.

Benzoic Acid Toluic AcidC PABAd ~~

33.33(3.1)6-’ 35.1 3(5.01)6-’ 39.46(5.75)6-’ -1.47(0.15) - 1 H(0.25) - 1.72(0.28) r = 0.8838 r = 0.8822 r = 0.8258

-0.002(0.0002)6-2 +0.99(0.09) r = 0.8641

1.69(0.16)p -0.29(0.04) r = 0.8856

3.20(0.25)@ -0.09(0.02) r = 0.9154

-0.002(0.0003)6-2 +1.07(0.14) r = 0.8677

1.74(0.23)/3 -0.27(0.07) r = 0.8939

3.15(0.40)@ -0.07(0.04) r = 0.9032

- 0.002(0.0004)6-2 + 1 .18(0.15) r = 0.8062

1.94(0.27)p -0.31 (0.08) r = 0.8359

3.52(0.47)@ -0.08(0.05) r = 0.8498

112.71(5.74)HBA-‘ 110.78(9.16)HBA-’ 127.90(12.18)HBA-’ -1.06(0.06) - 1.04(0.10) -1.20(0.14) r = 0.9620 r = 0.9553 r = 0.9131

a Numbers in parentheses are SEM. n = 33. n = 16. n = 23.

other mixed solvent systems by analysis of pKJs) versus l/q and l/HBA using a multiple linear-regression program. In most cases the ApK, should be estimated to within 0.1 pK, units or less.

These procedures utilize estimates for the values of the polarity scales based on a linear combination rule. Although some error is inherent in such estimations due to nonideal mixing of the cosolvent and water,lB the procedure appears satisfactory for the purpose of the estimation of the pKa values.

It has been suggestede that for large ions m y A t m y m ; the assumption will be more valid the larger the size of the ion. Thus, the nonelectrostatic effect of a particular solvent should be constant for a given solute series according to eq. 2. The solutes examined were chosen to study the effects of nondissociable functional groups on the medium effect. In particular, PABA and p-toluic acid were chosen as homo- morphs with different bonding characteristics, i.e., PABA is able to act as both a hydrogen bond donator and acceptor both in the undissociated and anionic state while benzoic and toluic acid will only act as hydrogen bond acceptors in the anionic state. In these two cases it might be expected that

# ,,,FHA. However, the results in Tables 11-V suggest that for most of the solvent mixtures, the medium effect is constant for the three solutes in a particular solvent mixture. This is evident from the similarity of the A S , values in most of the solvent mixtures for the three solutes as well as the similarity in the coefficients of the regression analyses. Grunwald and Berkowitz20 observed this phenomenon for several carboxylic acids in ethanol-water mixtures. These data suggest that this is the case for structurally related solutes in other amphiprotic solvent-water mixtures such as propylene glycol-water.

As noted previously, the largest differences in pKa changes between the three solutes were evident in the MezSO-water system. The ApKa values for benzoic and toluic acid were similar while those of PABA were significantly higher. Based on eq. 1, the larger ApKa values for PABA could be due to a smaller mm value, a larger myA value, or a combination of the two. The solubility of the salt and acid form of PABA and benzoic acid were therefore examined in MezSO-water and ETOH-water mixtures to provide a clue as to the relative changes in activity of the dissociated and undissociated species. These data are presented in Table VI. It can be seen that the ratio of the solubility of the salt to the

Table VI-Solubllltles (mglmL) of Salt and Acld Forms of Benzoate and PABA

0.5 Ethyl alcohol 0.5 Me$O Me2SO/ETOH. Benzoic Acid

Salt Acid SalVAcid

Salt Acid SalVAcid

PABA

288 k 0.0 174 ? 6.0 0.60 108 f 4.0 137 2 4.0 1.26

2.67 1.27

186 ? 3.5 433 f 0.0 2.33 65.5 2 4.4 182 ? 8.3 2.78

2.84 2.38

undissociated acid form of the solute is larger in MezSO- water mixtures for PABA than for benzoic acid when each solute is compared to itself in ETOH-water mixtures. In addition, the ratios of the sodium salt in f = 0.5 MezSO/f = 0.5 ETOH and the acid in f = 0.5 MezSO/f = 0.5 ETOH for each solute reveal that the solubility of both the dissociated and undissociated form of PABA are more soluble in f = 0.5 MezSO than in f = 0.5 ETOH. Both m~~ and m~ appear to be influenced by the solvent; however, it is also apparent that the medium effect on the anion has the predominant effect on the fla change. This is probably due to the ability of the anionic forms of both PABA and benzoic acid to be solvated by the ethanol hydroxyl group. Only the anionic form of PABA is strongly solvated by MezSO. Thus, the inability of the benzoate ion to form hydrogen bonds with MezSO will result in a lower activity of the ion in this solvent system.

References and Notes 1. Benet, L. Z.; Goyan, J. E. J . Pharm. Sci. 1967,56, 665. 2. Albert, A.; Sejeant, E. P. “The Determination of Ionization

Constants”; Cha man and Hall: New York, 1984; chap. 2. 3. Connors, K. A.; L i d o n , G. T.; Kennon, K. “Chemical Stability

of Pharmaceuticals: A Handbook for Pharmacists”; Wiley-hter- science: New York, 1979.

4. Martin, A.; Swarbrick, J.; Cammarata, A. “Physical Pharmacy”; Lea & Febiger: Philadelphia, 1983; ,p 300.

5. Bates, R. G. “Determination of pH ; John Wiley & Som: New York, 1964; p 195.

6. P o p ch, 0.; Tomkins, R. P. T. “Nonaqueow Solution Chemis- try’;%hn Wiley and Sons: New York 1981; chap. 5.

7. Robinson, R. A.; Stokes, R. H. “Electrolyte Solutions”; Butter- worths; London, 1968; p 356.

8. Alfenaar, M.; DeLi 9. Rubino, J. T.; Blangard, J.; Yalkowsky, S. H. J . Pharm. Sci. in

10. PfEiet, M. J.; Taft, R. W. J . Am. Chem. Soc. 1976,98,377. 11. Rubino, J. T.; Blanchard, J.; Yalkowsky, S. H. J . Parent Sci.

12. Van Uitert, L. G.; Haas, C. G. J . Am. Chem. Soc. 1953, 75,451. 13. Fedors, R. F. P?1 . E 14. Norusis, M. J. &SS?fIntroductory Statistics Guide”; McGraw-

15. Bates, R. G.; Paabo, M.; Robinson, R. A. J . Phys. Chem. 1963,67,

ey, C. L. Rec. Tmu. Chim. 1967,86, 929:

Technol. 1984,38, 215.

. Sci. 1974,14, 147.

Hill: New York, 1983. .“on IOJJ.

16. Pashankov, P. P.; Ziklov, P. S.; Budevsky, 0. B. J. Chrumutog.

17. Fiordiponti, P.; Rallo, F.; Rodante, F. Gozz. Chim. Ztal. 1974, 1981,209,149.

104,649.

All

1149.

4939.

dak Co.: Rochester, NY, 1975.

18. Fritz, J. S. “Acid-Base Titrations in Nonaqueous Solventd’;

19. So$, D. L.; Bitter, R. G.; Webb, J. G. J . Pharm. Sci. 1963,52,

20. Grunwald, E.; Berkowitz, B. J. J . Am. Chem. Soc. 1951, 73,

21. “Physical Properties of Some Organic Solvents”; Eastman Ko- 22. Barton, A. F. M. Chem. Rev. 1975, 75, 541.

and Bacon: Boston, 1973; chap. 2.

Acknowledgments Supported in part b the Pharmacy Foundation of North Carolina,

Inc. and the Burrougis Wellcome Co.

186 / Journal of Pharmaceutical Sciences Vol. 75, No. 2, February 1986

effects of solvent polarity on the acid dissociation constants of benzoic acids

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