PROTONATION OF ARYLSULFONATES IN AQUEOUS SULFURIC ACID
2279
The Protonation of Arylsulfonates in Aqueous Sulfuric Acid. Apparent pKa Values of Azo Dye Sulfonic Acids from Solubility Measurements by R. L. Reeves and R. S. Kaiser Research Laboratories, Eastman Kodak Company, Rochester, New York 14660
(Received December 18, 1968)
A number of azo dye sulfonatesare converted into protonated species in 0.1-3 M aqueous HzS04 with accompanying decreases in solubility of about 100-fold. Above ca. 3 IM HzS04, the solubilities become independent of acidity. Plots of log solubility against the HOacidity function conform quantitatively to the relationship
derived for an assumed acid-base equilibrium yielding a sparingly soluble acid. The derived pK, values are more positive by several log units than values for simple arylsulfonic acids estimated by other workers from solubility and other measurements. Similar large decreases in solubility of the salts occur in neutral solutions of NaClOl over a range of ionic strengths comparable to those of the acid solutions. Our results suggest that protonation of the azo sulfonate anions in the acid solutions occurs by salting out of a neutral protonated species, so that the apparent basicities of the salts may be a reflection of the tendency of the acids to be saltedout of a medium of varying ionic composition and structure. If this interpretation is correct, then there is a close parallel between the indicator acidity of the medium and the ability of the medium to “salt out” the electrically neutral species of the organic dyes. Although organic sulfonic acids are known to be strong acids in ionizing solvents, estimates of pK, values are very meager and give no consistent picture of acid strength. The leveling effect of water and alcohols has made direct measurement of pK, values in these solvents impossible, and most estimates of acidities of sulfonic acids are based on comparisons between these and other strong acids in relatively acidic solvents. Thus, benzenesulfonic acid is considered to be a stronger acid than hydrochloric acid in such solvents,’ and if one assumes that the two acids retain their relative acidities in aqueous solution, the aqueous pK, for benzenesulfonic acid should be more negative than -7.2 The pK, values for p-toluene- and p-xylenesulfonic acids have been estimated to be -4.12 and -3.10, respectively, from solubility measurements in aqueous sulfuric acid solution^.^ On the other hand, the pK, values for the sulfonic acid groups in t a ~ r i n e meth,~ enesulfonic acidJ6 and 4-(4-methoxyphenylazo) benzenesulfonic acid6 are estimated to be -0.3, -0.6, and 1.0, respectively. We have found that dye sulfonates 1-8 are converted into the corresponding sulfonic acids at surprisingly low acidities with accompanying large decreases in solubility. The decreases in solubility are directly proportional to the increase in acidity at low acidity, and apparent pK, values for an acid-base equilibrium have been derived. Evidence is presented to suggest that the solubility pK, values may not be thermodynamic pK, values which measure the intrinsic basicities of the sulfonate groups. Instead, the apparent values may be artifacts which just reflect the changes in the
+
activity coefficients of the dye ions with the changing medium. Our results suggest that there is a close parallel between the acidity and salting ability of aqueous sulfuric acid.
Experimental Section Materials. Dyes 2, 3, and 7 were the same as those used previously.’ Dye 1 was prepared in the usual way and after recrystallization from water as the sodium salt, it was converted into the crystalline sulfonic acid by acidification of a water solution with sulfuric acid. Anal8 Calcd for Cl6Hl1N~NaO4S.H2O: C, 52.2; H, 3.66; N, 7.6; S, 8.7. Found: C, 52.4; H, 3.3; N, 7.4; S, 8.7. Calcd for C~&II~NZO~S.HZO: C, 55.5; H, 4.0; N, 8.1; S, 9.2. Found: C, 55.5; H, 4.1; N, 8.0; S, 9.4; €490 (HzO), 2.11 X lo4. Dyes 4 and 5 were prepared from diazotized p trifluoromethylaniline. Dye 4 was recrystallized three times from methanol-ethanol, chromatographed on polyamide with DMF as eluent to remove traces of a (1) L. P. Hammett, “Physical Organic Chemistry,” McGraw-Hill Book Co., Inc., New York, N. Y., 1940,p 261. (2) R.P. Bell, “The Proton in Chemistry,” Cornel1 University Press, Ithaca, N.Y., 1959,p 89. (3) J. Horyna, Collect. Czech. Chem. Commun., 24, 1596 (1959). (4) E. J. King, J . Amer. Chem. SOC.,75, 2204 (1953). (5) K. N. Bascombe and R. P. Bell, J. Chem. Soc., 1096 (19.59). (6) J. F. Bunnett and E. Buncel, J . Amer. Chem. Soc., 83, 1117 (1961). (7) R. L. Reeves, ibid., 88, 2240 (1966). (8) Elemental analyses were carried out in the Microanalytical Laboratory of the Kodak Research Laboratories. Estimated precision is *0.3% for carbon, hydrogen, nitrogen, and sulfur on standard samples giving no interferences.
Volume 75, Number 7 July 1969
2280
R. L. REEVESAND R. S.KAISER
5
1, X = R = H 2, X = OMe, R = Na 3, X-CI, R=Na 4, X=CF3, R a N a
6, X=NOz 7, X-H
'OH
OH
S0,Na 8
9
0
&Qf
S03Na
d 10
colored impurity, and recrystallized two more times as the sodium salt from methanol-ethanol. Anal. Calcd for C I ~ H I O F ~ N ~ N ~C,O ~48.8; S : H, 2.4; N, 6.7. Found: C, 48.8; H, 2.6; N, 7.2; e484 (K,o), 2.34 x 104. Dye 5 was recrystallized twice from ethanol-water, chromatographed on polyamide to remove colored impurities, and recrystallized two more times as the sodium salt from ethanol-water. Anal. Calcd for C I ~ H I O F ~ N ~ W ~ O $ . H C,~ O 46.8; : H, 2.8; N, 6.4. Found: C, 46.8; H, 2.4; N, 6.5; e483 (H~o),2.66 x 104. Sulfonate 6 was obtained by the nitration of 7 with concentrated nitric acid at 110" for 10 min.O The product was recrystallized from ethyl acetate. Anal. Calcd for ClzHdVa0&3.2.5H20: C, 40.9; H, 4.0; N, 11.9; S, 9.1; H20, 12.8. Found: C, 40.6; H , 3.8; N, 11.9; S, 8.9; HzO, 10.3. Compounds 7, 8, 9, and 10 were commercial samples and were recrystallized before use. Solubility Measurements. Saturated solutions were prepared in two ways, which gave identical results. I n the first, a water solution containing a relatively high concentration of dye sulfonate salt was diluted to the desired acidity by addition of a stock sulfuric acid solution from a buret, so that the acid dye precipitated. This method gave very flocculent precipitates that were difficult to separate from the supernatant solutions on which the measurements were made. I n the second method, solid dye sulfonate salt was shaken in contact The Journal of Ph&cal Chemialry
with the aqueous acid solution for sufficiently long times to ensure equilibration in the solid phase as well as equilibration between solid and solution. With this method, separation of the two phases was easier. The mixtures were shaken continuously for several days in a room a t 24.5" and allowed to equilibrate several more days in a constant-temperature bath at 25". The samples were then centrifuged at 8000 rpm in a Serval1 centrifuge for 10 min, during which time the sample temperature remained at 25 f 1". A portion of the supernatant solution was withdrawn with a hypodermic syringe fitted with a 25-gauge needle. Aliquots were diluted quantitatively, and their absorptions were measured with a Gary Model 14 spectrophotometer. All the dye solutions obeyed Beer's law at the analytical concentrations. I n those cases where the molar absorptivity varied with acidity, interpolated values of the absorptivity at the acidity of the diluted samples were obtained from plots of e vs. [HzS04]. The precipitates could not be removed by filtration because the saturated solutions were very dilute and the dyes were adsorbed from solution onto paper and glass.
Results States of Aggregation. Dyes often form aggregates in solution if the concentrations are sufficiently high. Most of the dyes studied here are sufficiently insoluble in aqueous electrolyte solutions (Tables I and 11) that the dissolved species are molecularly dispersed over the entire dye concentration range up to saturation. The following accepted spectral criteria were used to rule out aggregation: (1) adherence to Beer's law (constant molar absorptivity at the absorption maximum to within *l'%); (2) the absence of new bands formed at the expense of old ones as the dye concentration is increased; and (3) the absence of any band broadening or changes in the band shape as the dye concentration is changed. Absorption curves of the dye solutions M up were obtained at concentrations from 2.5 X to saturation in 0.5 M sulfuric acid and at several higher acidities appropriate to the substrate in question. Dyes 7 and 8 showed evidence for aggregation at dye concentrations near saturation at some acidities. Solutions of Orange I1 (8) showed significant deviations from Beer's law in 0.66 and 2.0 M acid. Solutions of 7 showed similar deviations at concentrations below saturation in 0.66 and 3.5 M acid, The solubilities of sulfonates 9 and 10 are higher at all acidities than the solubilities of the dyes by several orders (Table I). Solutions of these sulfonates gave significant deviations from Beer's law at concentrations below saturation a t all acidities investigated. The remaining dyes appeared, by the spectral criteria cited, to be molecularly dispersed in the solution phases of the acid solutions at all dye concentrations (9) Janovsky, Monatsh., 4, 276 (1883) ; 271.
through Beilstein, 16,
PROTONATION OF ARYLSULFONATES IN AQUEOUS SULFURICACID
228 1
Table I : Solubilities of Sulfonates in Aqueous Sulfuric Acid a t 25"
Compd 1
IHzSOrl,
S,
[HiSOd,
SI
M
M X 106
Compd
M
M X 10s
Compd
1945 73.5 32.7 23.6 16.4" 6.78 4.42 1.56 0.765 0 . 423" 0.334a 0.2540 0.235
2
0.0 0.125 0.250 0.325 0.500 0.660 0 850
290 11.1 5.09 4.08 2.34 1.71 1.19 0.880 0.548 0.457 0.307 0.244 0.203 0.222
3
0.00 0.250 0.500 0.660 0.850 1.40 1.75 2.25 2.85 3.50 4.00 4.50 5.00
I
1.10 1.40 1.75 2.10 2.45 3.00 3.50
[Hisor],
8,
M
M X 106
0.0 0.125 0.250 0.500 0.850 1.10 1.40 1.75 2.10 2.45 2.85 3.50 4.00
716 8.62 4.06 2.05 1.07 0.798 0.622" 0.353 0.278 0.236 0.190a 0.124a 0,125'
M X 104
4
0.125 0.250 0.660 1.40 1.92 2.45 2.85 4.00 4.50 5.50 6.50
17.0 9.65 3.94 1.71 1.14 0.638 0.494 0.164 0.139 0.100 0.093
5
0.0 0.050 0,100 0.250 0.660 0.920 1.40 1.92 2.45 3.50 4.00 4.50
M X 108
7
0.345 0.480 0.500 0.660 0.850 1.40 1.75 2.45 2.85 3.50 4.00 4.50 5.00
242 114 64.1 25.5 13.9 4.79 2.95 1.22 0.820 0.467 0.280 0.2215 0.188
112 26.6 18.4 7.41 3.050 1.86 1.44 1.04 0.653= 0.568" 0.603 0.640"
6
0.660 1.40 1.75 2.45 3.50 4.50 5.50 6.50
M X 10
M X 10'
8
0.500 0.545 0.660 0.800 0.850 1.40 1.75 2.45 2.85 3.50 4.00 4.50
8844 437 9.78 3.13 3.25 0.952 0.554 0.271 0.191 0.149 0.141 0.164
73.3 25.3 14.9 8.06 3.54 2.39 1.48 1.52
9
0.850 1.25 1.75 2.45 3.25 4.50 6.42
9.84 8.33 7.54 5.14 4.11 2.56" 1.36"
[HISO'], wt %
55.8 60.4 64.5 68.0 71.4 74.7 78.0
1.570 1.06 0.940" 0.563 0.413 0.362 0,4725
M X 10a
10
a
0.250 0.660 1.4W 1.75 2.45 3.50 4.50 5.50
3.46 4.02 3.77 3.56 3.05 2.16 1.78 1.73
Mean value of replicate determinations.
Volume 73, Number 7 July 1969
R. L. REEVESAND R. S. KAISER
2282
below saturation. No deviations from Beer's law were found for these dyes over the entire solubility range at several acidities. We compared the visual absorption bands of dyes 1 through 5 by superimposing normalized curves of saturated solutions with those of very dilute M ) . I n every case, the two solutions (2-5 X curves were congruent within the combined experimental error of 0.008 absorbance unit for spectrophotometry and sample preparation. Table I shows that the solubilities of dyes 1 to 3 change dramatically between 0.125 M acid and pure water. Solutions of all the dyes studied appeared to be aggregated in concentrated solutions in water, as evidenced by deviations from Beer's law a t concentrations below saturation. For this reason, no attempt is made to interpret the large decreases in solubility found in going from pure water to the most dilute acid studied since the dissolved species are not the same in the two media. For the same reason, no effort is made to calculate numerical values for activity coefficients from the solubilities in water and the acid solutions.
Table I1 : Solubilities of Dyes in Aqueous NaClOa Solutions at 25" (NaClOd), M
0.100 0.375 0.500 0.975 1 .oo 1.98 2.00 3.50
(a) ( M X 10s)
6.48 1.30
Ho
(8)
(M
x
(6)
109
16.9
(M
x
108)
3.11
...
...
4.23
0.512
...
2.44
0.328
0.455
...
...
...
1.53
0.219
0.499
0.858
...
...
0.665
...
Figure 1. Logarithmic plots of solubility in aqueous HzSOc at 25" against Ho: 0, sulfonate 5 ; A, sulfonate 6.
...
proportion to the acidity as measured by the tendency of weakly basic aromatic amines to be protonated in the same medium. Such proportionality suggests that the solubility changes result from an acid-base interaction. Table 111: Apparent pK, Values from Solubilities at 25"
Solubilities in Aqueous HzS04 and NaC104. The solubility data are recorded in Tables I and 11. With increasing acidity the solubilities of the dyes decrease markedly and then become independent of acidity. Those dyes having electron-withdrawing para substituents (3-6) could be studied at higher acidities than the others without significant protonation of the azo nitrogens.' The acidity-independent solubilities of these dyes are true limiting solubilities. The apparent limiting solubilities of 2 and 7 may be slightly in error since, a t the acidities at which the solubilities become independent of acid concentration, a small fraction of the dissolved dyes is present as the azonium ions. This fraction is comparable to the error in the solubility 5%) at these acidities and is, theremeasurements (I fore, not serious. Plots of log solubility against the HOacidity function are shown in Figure 1 for two typical unassociated sulfonates. The slopes of the acidity-dependent segments of the plots are given in Table 111. These slopes show that the log solubility values decrease in direct The Journal of Physical Chemistry
d log SldHo
P K ~
-1.8 -0.9 -1.1 -2.0 -0.7 -1.8
1.1 1.0 1.0 0.8 1.0 0.9 ~~~
~~
~
We considered a model in which soluble sulfonate salt is converted into a sparingly soluble protonated form, HA. Kn
HA
+ HzO JrA- + HaOf
(1)
By this model HAsolid
aHA(so1n)
HAso~n
=
(2)
const = C
and
8=
CA-
-k
CHA
(3)
PROTONATION OF ARYLSULFONATES IN AQUEOUS SULFURIC ACID
2283
where S = solubility
S =
KaCHAf H A
aH +fA -
+
CHA =
KaaHA
+ HA
-
aH+fA-
~
(4)
fHA
where K, is the acid dissociation constant of HA. If u H + is expressed in terms of the Hammett acidity function, ho, then
where f B and f B H + are the activity coefficients of the unprotonated and protonated forms of the Hammett indicator. At high acidities the solubility will be independent of acidity and equal to C/fHA = S l i m . At low acidities, the second term will be negligibly small compared to the first and
A plot of log S against Ho should be linear with unit slope at low acidities if the last term in eq 5' is constant. Our data for dyes 1-6 conform very well to the behavior predicted by eq 5 . The solubilities of the arylazosulfonates 7 and 8 also follow the predicted relationship over a limited range of acidities. Apparent pKa values were determined from the Ho value at which the intercept of the linear aciditydependent segment of the log S plots crosses the extension of the acidity-independent segment (Figure 1). These values are given in Table 111. The values for dyes 1-4 do not fall into the order of basicities that would be predicted from normal substituent effects and that were previously found for the protonation of the azo nitrogen^.^ If CHA is assumed to be constant at all acidities, it can be replaced by the values of the limiting solubilities ( S l i m ) found at high acidities, and eq 4 becomes
which is the form used by Horynae3 Plots of log [(X/Xlim) - 11 against HO for dyes 1-6 were linear throughout and had slopes close to unity. The same apparent pKa values were obtained by either treatment. A plot of log S against HOis shown in Figure 2 for saturated solutions of Orange I1 (8). The initial decrease in solubility of this dye is much steeper at low acidities than was found for dyes 1-6. Since we have shown that this dye forms aggregates at low acidity a t which the solubility is high, the initial precipitous de0.2 and Ho - 0.2 crease in solubility between H o probably represents a region of acidity in which highly soluble aggregates are being converted into smaller aggregates or dye in molecular dispersion. The region of change between HO - 0.3 and Ho - 1.0 has nearly unit slope and is undoubtedly the result of the same sort
+
Figure 2. Logarithmic plot of the solubility of Orange I1 (8) in aqueous HzSOca t 25' against Ho.
of phenomenon that causes the similar change in the nonaggregated dyes.l0 The same sort of plot is obtained for the azobenzene (7)) which forms aggregates in the more concentrated solutions at low acidities. I n the region of intermediate change, log S varies with H o with a slope of one. A limiting solubility for 7 could not be obtained before significant protonation of the azo nitrogens occurred, although the solubility tends toward a limiting value. Figure 3 shows a comparison of the solubility changes for dyes 2, 3, and 5 as a function of ionic strength in aqueous sulfuric acid and in sodium perchlorate solutions. The ionic strengths of the acid solutions were calculated from the tabulation of concentrations of ionic species given by Robertson and Dunf0rd.l' The results show that the solubility of the dye salts is also decreased severely by a given ionic strength of a neutral salt. Whereas the solubility curves for the two electrolyte solutions are qualitatively similar, quantitative differences do exist. The largest difference between the two electrolyte solutions is with dye 5 at an ionic strength of 2 M , where the solubilities differ by a factor of seven. This is of the order to be expected from specific electrolyte effects on the activity coeffi(10) The minimum at Ho - 1.8 does not represent a true limiting solubility since significant concentrations of the szonium ion are present at this acidity. We have estimated a p K , for dissociation of the azonium ion of Orange I1 from spectral changes and obtained a value of -2.2 on the Ho acidity scale. The slope, d log [(BH+)/ (B)]/dHo, however, is only 0.89. (11) E.B. Robertson and H. B. Dunford, J . Amer. Chem. SOC.,86, 5080 (1964).
Volume 78, Number 7 July 1969
2284 cients of an uncharged organic molecule. Thus, the activity coefficient of 2,4-dinitrochlorobenzene is six times greater in a solution of sodium sulfate at ionic strength of 2 M than in a solution of sodium perchlorate at the same ionic strength.12 This suggests that electrolyte concentration may be the principal factor in decreasing the solubility in the acid solutions and that the sulfonates may be converted into the sulfonic acids through a salting-out mechanism. Plots of log (1,'s) against ionic strength of the acid solutions were curvilinear, with the greatest curvature a t ionic strengths less than one. Thus, a simple Setchenow relationship does not hold for the dyes, and the activity coefficients of the dyes respond to changes in ionic strength in a somewhat different way than do small neutral organic molecules. Finally, the solubilities of two arylsulfonates that are not azo dyes (9 and 10) were measured to see what effect the presence of azo groups might have in causing the solubility changes. Both of these materials are associated at concentrations near saturation so the solubility data are difficult to interpret without ion association constants. The solubilities are not decreased by added acid in the same way as the solubilities of the azo dye sulfonates. We cannot distinguish between basic chemical differences and the formation of stable aggregates. Spectral Studies. Absorption curves of 2.5 X M solutions of dyes 1-5 were measured in water and in dilute sulfuric acid a t acidities up to 2.45. Only in the case of 2 are measurable concentrations of the azonium species present in 2.45 M acid so that any spectral changes at lower acidities are due to some other interaction. We find small changes in the molar absorptivities of the ultraviolet bands at these low acidities, without any wavelength shifts. There are, in addition, complex changes in the absorptivities of the visible bands in the solutions of 4 and 5 between water and 2.45 M acid. Figure 4 shows a plot of absorptivity against Hofor 4 and 5 at the wavelengths of three absorption maxima. The decreases in E are not accompanied by broadening of the bands; the changes reflect differences in the areas under the curves. No isosbestic points are found. The most significant changes in a occur between water and 0.125 M acid. Bunnett and Bunce16have reported a 48% increase in the absorbance of the long-wavelength band of 444methoxypheny1azo)benzenesulfonate in going from water to 5% aqueous sulfuric acid and have suggested that this change is the result of covalent protonation of the sulfo group. We have not observed such a change M ) solutions of our sample of in very dilute (2.5 X this compound.7 Jaff6 and his coworkers have also observed small changes in molar absorptivities of nonionic14 and of cationic az~benzenes,'~ in aqueous sulfuric acid, which tends to rule out any specific role of the sulfo group in causing these small variations. The Journal of Physical Chemistry
R. L. REEVESAND R. S. KAISER
o-2'1
1
2
3
CI (MI
Figure 3. Logarithmic plots of solubilities a t 25" in aqueous HzSOl and in aqueous NaClOd for sulfonates 2, 3, and 5: 0, HzSO~;A, NaC104.
Discussion An unambiguous chemical interpretation of the results cannot be made from the solubility data alone, The solubility changes are a measure of the changes in the mean activity coefficients of the dyes. We have shown that the changes can be treated thermodynamically in terms of an acid-base equilibrium. The mechanism of the change is another problem. This involves the question of whether protonation occurs primarily because of intrinsic basicity or because the high concentration of electrolyte competes so effectively for water of hydration that the dye ions become insufficiently solvated to remain dissociated. I n the first case, the over-all free-energy change should reflect largely potential energy changes in the organic molecule, and systematic changes in polar substituents should cause the pK, values to change systematically in a predictable direction. On the other hand, if protonation occurs by a salting-out mechanism, the derived pK,, values will reflect more strongly the free-energy changes of the whole system. Factors other than potential energy changes within the molecule would become important. These factors include competitive hydration, ionic size and shape, and the hydrophobic volume of the organic moiety of the organic ion. We believe our present results are best interpreted in terms of the salting-out mechanism. The fact that the acid solutions and solutions of neutral salts cause comparable changes in the activity coefficients of the dyes is the most compelling evidence for this conclusion. I n addition, we find that the derived pK, values do not (12) C.A. Bunton and L. Robinson, J . Amer. Chem. SOC.,90, 5965 (1968). (13) H. S. Harned an,:' B. B. Owen, "The Physical Chemistry of Electrolyte SoIutions, ReinhoId Publishing Gorp., New York, N. Y.,1958,p 532. (14) H.H. Jafft5, 8. J. Yeh, and R. W. Gardner, J . Mol. Spectrosc., 2 , 120 (1958). (15) M. Isaks and H. H. Jaff6, J . Amer. Chem. Soc., 86, 2209 (1964).
PROTONATION OF ARYLSULFONATES IN AQUEOUS SULFURIC ACID I
2285
l -
I
B-.
*
*H+*(z 1)HzO
K%
+ HZO (9) BHanHz0 + HzO (10)
B Q . * H + *( z 2)HzO B-...H+.(n
+
K7l
1)H20
The sum of equilibria 7-10 is the classical Bronsted protonation equilibrium
+
B - * z H ~ O H+.yH20 J_ BH-nHzO
+ (Z + y - n)H2O
(11)
The acid dissociation constant for equilibrium 11 is the product of the reciprocals of the individual KZ
1- 1
I
70
0.4
0.8
0
-0.4
I
-06
2
-I
HO
Figure 4. Effect of acidity on the molar absorptivities of solutions of 4 (dashed lines) and 5 (solid lines) at the wavelengths: 0, 486 nm; A, 358-362 nm; 0, 298 nm; 0, 482-486 nm; A, 285 nm; H, 263-267 nm.
change with polar substituents in a systematic way but instead change rather randomly with such structural changes (compare 1 with 4 and 2 with 3). The closeness of the values may be understood when one considers that the sizes of the dyes are similar except for 6, and all should be salted-out by approximately the same amount a t a given electrolyte concentration. Consideration of a salting mechanism for protonation of sulfonates can also help explain the large difference in pK, values between toluene- and xylenesulfonic acids as determined by solubility changes. a A decrease in acidity of one pK unit is quite large to be explained in terms of the polar effect of an additional methyl group, but the change seems more reasonable when interpreted in terms of differences in molecular size and shape and the effect of these differences on ion hydration. The conversion of the organic salts into covalently protonated species by a salting mechanism probably involves solvated ion pairs along the way. The most general form of such a scheme has been set forth by Haldna for nonionic weak bases.lB KO
+
B - ~ z H ~ O H+*yHzO
+ + y - z)HzO
(7)
+ HzO
(8)
B-- * *H+*zHzO (Z Ki
B-.
*
*H+.xH20
B-* * H + * ( z 1)HzO
I n this scheme, the weakly basic anion becomes covalently protonated when the mineral acid level becomes high enough to compete successfully for the water needed to maintain the organic salt in a dissociated state. The larger the hydrophobic volume of the organic ion, the greater will be the solvation needed to keep it dissociated and the lower will be the critical “acidity” a t which it forms ion pairs and finally becomes covalently protonated. This interpretation might explain why the solubility pK, values for the smaller toluene- and xylenesulfonic acids are so much more negative than the values for the larger dye sulfonic acids. The scheme represented by eq 7-10 actually involves a competition between the ionic base B- and added electrolyte for water molecules, so it seems equally applicable to solutions of neutral salts, in which H+ is replaced by M+. The proton may shift the equilibria to the right more effectively because of its compactness and tendency to hydrate so strongly. Efforts to calculate indicator acidities of aqueous sulfuric acid solutions by considering proton hydration have been s u ~ c e s s f u l717-19 .~~ Our results suggest that the hydration of the indicator should also be considered and that a competition model might be more appropriate. The compact, highly hydrated inorganic ions have an effect on the activity coefficients of the large, poorly hydrated dye ions much different from that of increasing concentrations of arylsulfonic acids and sulfonates upon their own activity coefficients. Thus, the mean activity coefficients of a series of alkyl-substituted arylsulfonic acids and sulfonates decrease by factors of (16) (a) Yu. L. Haldna, Reakta. Sposobnost Org. Soedin., 1 (l), 184 (1964); (b) Yu. L. Haldna and J. K. Rodima, ibid., 3 (2), 169 (1966); ( 0 ) Yu. L. Haldna, H. J. Kuura, and L. E.-J. Erreline, ibid., 4 (3), 678 (1967); (d) Yu. L. Haldna, Zh. Fiz. Khim., 41, 2106 (1967). (17) K. N. Bascombe and R. P. Bell, Discussions Faradag Soc., 24, 168 (1957). (18) P. A. H. Wyatt, ibid., 24, 162 (1957). (19) E. Hagfeldt, Acta Chem. Scand., 17, 785 (1963).
Volume 78, Number 7 July 1969
R. L. REEVESAND R. S. KAISER
2286 less than ten in the range, 0.1-5.0 M,20 whereas the activity coefficients of our dye ions incyease by factors of 100 in the mineral acid solutions. All the sulfonates are half-protonated a t acid concentrations between 2 and 4 M . Robertson and Dunford have calculated that at 2 M H2SO4the concentration of water involved in the hydration of proton species is approximately 25 M.” The total water concentration at this acidity is 51.5 M. The characteristic “aqueous” structure of the medium must be altered considerably, therefore, even in 2 M acid. Robertson and Dunford point out that the greatest change in the properties of aqueous sulfuric acid solutions occurs a t 2 i 1 M acid. Thus far, the question of the nature of the protonation measured by the solubility pK, values has been avoided. We have no way of knowing whether the dyes precipitated from the acid solutions exist in the solid phases as the azosulfonic acids (11) or as azonium sulfonate zwitterions (12). We are certain that except in the cases cited, the concentrations of azonium ions in solution are below the level of spectrophotometric
bOBH 11
$0,12
detection in the range of acidities where the solubility pKa values are obtained. We have measured the pK, values for the azonium species previously by a spectrophotometric method and have found that the spectrophotometric pK, values for the dyes studied here are more negative than the values from the solubility measurements by 1.3-2.0 pK units.’ The present findings do not alter the significance of the spectrophotometric pK, values for the azonium species. The assumption that the solubility pK, values measure the dissociation of the sulfonic acid species (11) is not free of difficulties. The pK, values found here are more positive by 2-3 pK units than the values found for simple arylsulfonic acids. Whereas the electrondonating influence of the OH group makes the sulfo group more basic, the polar effect of the strongly electron-withdrawing phenylazo groups21should largely counterbalance this effect. As pointed out above, the large molecular size of the dye sulfonates may account
The Journal of Physical Chemistry
for the difference in pK, values if protonation is occurring by a salting-out mechanism. We cannot say whether the presence of arylazo moieties in the molecules has any specific effect’ other than one of increasing molecular size. Direct comparison of solubilities of the unassociated dyes with those for 9 and 10 is impossible from the present data because the latter compounds are aggregated in the solution phase at saturation. The present results are significant when compared with data for monoionic weak organic bases in aqueous sulfuric acid. A variety of experimental methods indicate that weak oxygen bases undergo association with solvated protons a t much lower acidities than those needed to protonate the oxygen atom covalently. 22-24 The estimated “pK,” values for all these protonated complexes fall within the range 0 to -1, which is remarkably close to the apparent pKa values of our ionic bases. We believe this supports the contention that such weak bases become protonated because of changes in the medium rather than any intrinsic basicity. The results given here cannot answer all the questions they raise without further work. The significant findings are the fact that protonation of some weakly basic ionic organic indicators appears to occur by a salting-out mechanism in aqueous sulfuric acid and that protonation by this mechanism parallels the indicator acidity as measured by nonionic indicators. The results provide additional insight into the factors giving rise to the high acidity of the medium as well as to why different types of indicators generate different acidity function scales.
Acknowledgment. The authors are indebted to Dr. Karl Tong and to Dr. William West of the Kodak Research Laboratories for helpful discussions of this work. (20) 0.D. Bonner and 0. C. Rogers, J. Phys. Chem., 64, 1499 (1960). (21) M.Syz and H. Zollinger, Helv. Chim. Acta, 48, 383 (1966). (22) (a) V. A. Pal’m, Yu. L. Haldna, and A. J. Talvik in “The Chemistry of the Carbonyl Group,” S. Patai, Ed,, Interscience Publishers, Inc., London, 1966, Chapter 9; (b) V. A. Pal’m and Yu. L. Haldna, Dokl. Akad. Nauk SSSR, 135, 667 (1960); (0) Yu. L. Haldna and R. K. Pass, Zh. Fiz. Khim., 38, 1629 (1964). (23) S. Nagakura, A. Minegishi, and K. Stanfield, J . Amer. Chem. Soc., 79, 1033 (1967). (24) (a) C. F. Wells, Trans. Faraday SOC.,62, 2816 (1966); (b) C. F.Wells, ibid., 63, 147 (1967).
