Monatsh Chem (2016) 147:1009–1014 DOI 10.1007/s00706-016-1728-4
ORIGINAL PAPER
Solubility and acidic constants at 25 C in NaClO4 aqueous solutions of 1-(2-hydroxyphenyl)ethanone Emilia Furia1 • Giovanni Sindona1 • Antonio Tagarelli1
Received: 4 November 2015 / Accepted: 7 March 2016 / Published online: 31 March 2016 Springer-Verlag Wien 2016
Abstract The acid–base equilibria of 1-(2-hydroxyphenyl)ethanone (HY) were studied, at 25 C, by potentiometric measurements with a glass electrode in NaClO4 media for ionic strength ranging from 0.51 to 3.5 mol/kg. The solubility (S) of HY was determined in the same range of ionic strength in NaClO4 media. The total concentration (CL) of ligand was ranged from 1 9 10-3 mol/kg to 20 9 10-3 mol/kg; the experimental data were explained by assuming, for HY, the species HY2- besides Y-. Elaboration of data, according to the specific interaction theory, had provided the constants valid in the infinite dilution reference state: HY ¢ Y- ? H?, log K1-1 = - 10.11 ± 0.05 and 2 HY ¢ HY2- ? H?, log K2–1 = - 7.85 ± 0.05. The uncertainties were given as 3r. The interaction coefficients (kg/mol) of ionic species Y- and HY2- with the medium counter ion were e (Y-, Na?) = 0.50 ± 0.05 and e (HY2-, Na?) = 0.59 ± 0.05, respectively. Graphical abstract 1
ZH = K1-1 × h-1 / (1 + K1-1 × h-1)
1.05 3.5
Dimer formation evidence 0.2
0.5 0 8.5
9.5
0 8
10
-log [H+]
12
& Emilia Furia
[email protected] 1
Department of Chemistry and Chemical Technologies, University of Calabria, Arcavacata di Rende, CS, Italy
Keywords Solvent effect Dimer formation Acidity Constant ionic medium method
Introduction In the last years, we have been interested in a systematic study concerning the complexation of metal cations with 2-hydroxybenzoic acid (salicylic acid) and its derivatives [1–8]. Salicylic acid is both the simplest model for humic acids present in soil and an outstanding antirheumatic and antifungal substance; from this point of view its derivatives, such as 1-(2-hydroxyphenyl)ethanone (HY, Scheme 1), display the same properties. Research in literature [9] showed that data were scanty and measurements were carried out in non-aqueous solvents because of the low solubility of ligand, while it seems interesting to evaluate the complexing power in water to estimate its biological properties. The aim of this work was the evaluation of solubility and acidic constants of 1-(2hydroxyphenyl)ethanone at 25 C in NaClO4 aqueous solutions, as constant ionic medium at four different concentrations (i.e., 0.51, 1.05, 2.20, and 3.50 mol/kg). The constant ionic medium method, proposed by Biedermann and Sille´n [10], has proved to be indispensable in equilibrium studies of complicated ionic reactions. This method, which consists of using, as a solvent, concentrated solutions of inert salts, has remarkable effectiveness for the minimization of variations in activity coefficients. The equilibrium constants, calculated from measurements in a given medium, are strictly valid only in that medium. This means that for practical applications, such as the speciation of metal ions in a natural system, it is necessary to determine the equilibrium constants in each of the conditions prevailing in the ocean, in biological fluids, and in surface and
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E. Furia et al.
Scheme 1
O
OH
ground waters. The investigation on the acidic constants of HY in aqueous solutions showed the formation, besides Y-, of the dimeric species HY2-. These results were unexpected, taking into account the low concentration of the ligand, but not unlikely. It is well-known that carboxylic acids tend to associate into dimers as a consequence of intermolecular hydrogen bond formation [11]. As concerns salicylic acid, these associations were subject of a great number of investigations [11–15], while there was no evidence on the dimer’s formation of HY in aqueous solution. The potentiometric data, in combination with solubility measurements of ligand in pure water and in NaClO4 solutions, were treated with the specific ion interaction theory (SIT) [16, 17], to calculate the acidic constants at the infinite dilution reference state (K1-1 and K2-1) and the interaction coefficients of the ligand’s anions (Y- and HY2-) with the medium counter ion Na?. Since the theory is formulated in terms of molal units, constants and other quantities in the following treatment are expressed on the molal scale. The conversion factors from molarity to molality were assumed from Baes and Mesmer [18].
Fig. 1 Absorption spectrum of 1-(2-hydroxyphenyl)ethanone in aqueous solution
standard solutions, taking into account that the absorbance (Ak) may be expressed as Eq. (1): Ak ¼ lek S
ð1Þ
where l was the optical path and ek was the molar absorptivity. The reproducibility of solubility data was 0.5 %. The activity coefficients of the neutral species HY (cHY) were obtained from the relationship between the solubility in H2O and those in I mol/kg NaClO4 [19] [Eq. (2)]: logcHY ¼ logðS =SÞ
ð2Þ
where S is the solubility of 1-(2-hydroxyphenyl)ethanone at the infinite dilution reference state. Results are reported in Table 1. As can be seen, 1-(2-hydroxyphenyl)ethanone had displayed a salting-in behavior as the solubility was increased monotonically with the ionic strength. Potentiometric data
Results and discussion Solubility of 20 -hydroxyacetophenone The evaluation of the activity coefficients of 1-(2-hydroxyphenyl)ethanone at different levels of ionic strength was required to determine the maximum concentration level of ligand and to evaluate the acidic constants at infinite dilution reference state by means of the specification interaction theory [16, 17]. NaClO4 solutions were equilibrated with an excess of 1-(2-hydroxyphenyl)ethanone at 25 C under continuous stirring. After separation through a glass filter G4 and appropriate dilution, the absorbance of the various solutions were measured. Three replicates were run for each point. A typical absorption spectrum is presented in Fig. 1. As can be seen, there were three bands centered at 215, 252, and 325 nm. The most intense peak at 215 nm seems the most appropriate, however, its reproducibility was inadequate. The peak at 252 nm was considered more satisfactory for accurate determinations. The absorbance at 325 nm was not used since it was less intense than that at 252 nm. The solubility (S) was obtained by interpolation of the absorbance from a calibration curve, constructed with
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Experimental data were explained according to the following equilibria [Eqs. (3) and (4)]: HY Y þ Hþ K11
ð3Þ
þ 2HY HY 2 þ H K21
ð4Þ
The primary data (EG, VT, h, CL, CA, and CB) were treated with both graphical and numerical methods. The Table 1 Solubility of 1-(2-hydroxyphenyl)ethanone in I mol/kg NaClO4 I/mol kg-1 NaClO4
S/mmol kg-1
-log cHY
0.00
13.7 ± 0.4
0
0.10
15.3 ± 0.5
0.05 ± 0.02
0.25
15.68 ± 0.05
0.06 ± 0.01
0.51
16.6 ± 0.3
0.08 ± 0.02
1.05
17.3 ± 0.2
0.10 ± 0.01
1.51
19.1 ± 0.3
0.14 ± 0.02
2.20
19.8 ± 0.3
0.16 ± 0.02
3.50
21.2 ± 0.4
0.19 ± 0.02
Solubility and acidic constants at 25 C in NaClO4 aqueous solutions of 1-(2-hydroxyphenyl)…
1011
Table 2 Survey of ZH vs. pH in 0.51 and 1.05 mol/kg NaClO4 pH
pH
ZH CL = 15.25 mM CL = 1.89 mM I = 0.51 mol/kg
8.51
I = 1.05 mol/kg
1.25 9 10-2
1.59 9 10-2
8.45
8.54
1.34 9 10
-2
8.73
8.76
2.25 9 10-2
8.74
8.77
-2
2.93 9 10
3.79 9 10-2
8.75
8.78 8.93
3.04 9 10-2
8.97
4.85 9 10-2
9.07 9.12
ZH CL = 14.94 mM CL = 4.95 mM
3.51 9 10
-2
5.37 9 10
-2
8.90 9.03
6.02 9 10-2
9.04
6.66 9 10-2 7.76 9 10-2
9.27
4.06 9 10-2 5.18 9 10-2 -2
9.13
6.31 9 10
9.21
7.44 9 10-2 8.27 9 10-2
9.26
-2
-2
9.33
8.54 9 10
9.29
8.58 9 10
9.45
0.1041
9.35
9.71 9 10-2
9.53
0.1165
9.39
0.1051
9.58
0.1231
9.41
0.1085
9.66
0.1420
9.46
0.1199
9.51
0.1314
9.70
0.1632
9.74
0.1610
9.52
9.82
0.1803
9.56
0.1429
9.93
0.2488
9.60
0.1544
9.97 10.10 10.22 10.51
0.5574
10.59
0.6022
10.77
0.1774
0.3442
9.75
0.2006
0.4484
9.82
0.22392
9.88
0.2474
9.94
0.2711
10.00
0.2950
10.05
0.3190
0.7955
10.15
0.3680
graphical treatment had the advantage of providing a realistic evaluation of errors on the equilibrium constants. From set of the potentiometric and analytical data, the ZH function was obtained, which was defined as the average number of protons per ligand [Eq. (5)]: ð5Þ
where CH is the analytical protonic excess. By assuming the existence of the only equilibrium 3, the function ZH may be expressed according to Eq. (6): ZH ¼ K11 h1 =ð1 þ K11 h1 Þ
0.1741
0.6962
0.8136
ZH ¼ ðh CH Þ=CL
0.1508
9.76
0.7426
11.00 11.05
9.64 9.68
0.3924
10.32
10.87
0.2664 0.3288
10.13
0.1279
ð6Þ
A summary of ZH vs. pH at the ionic strengths investigated is reported in Tables 2 and 3. By introducing
10.25
0.4182
10.34
0.4698
the normalized variable (u = K1-1 9 h-1) the model function was obtained [Eq. (7)]: ZT ¼ u=ð1 þ uÞ
ð7Þ
The points ZH vs. log [H?] (Fig. 2), at different CL, should fall on a unique curve. As can be seen in Fig. 2, the majority of the experimental points fell on a unique curve which tends to 1. A careful inspection of the graph showed that small but systematic deviations from the model including equilibrium 3 exclusively were observed. These deviations were an indication that some additional species were present. A preliminary estimation of the constant K1-1 was obtained by fitting the points coinciding within the experimental error
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E. Furia et al.
Table 3 Survey of ZH vs. pH in 2.20 and 3.50 mol/kg NaClO4 pH
pH
ZH CL = 18.35 mM CL = 10.77 mM I = 2.20 mol/kg
I = 3.50 mol/kg 7.92 9 10-3
8.64 8.70
9.12 9 10
9.13
8.77
1.61 9 10-2
9.41 2.31 9 10
-2
3.87 9 10
-2
9.59
-2
8.98 9.14
2.71 9 10
9.27
0.0528
9.72 9.91
5.43 9 10
-2
10.09
0.0688
9.50
7.01 9 10
10.21
9.64
8.60 9 10-2
10.27
9.74
0.1019
10.37
9.84
0.1180
10.54
9.92
0.1341
10.61
9.99
0.1503
10.75
10.06
0.1667
10.82
0.1073
9.97
2.91 9 10-2 -2
4.17 9 10
5.43 9 10-2 -2
7.97 9 10
9.25 9 10-2 0.1054
10.15 -2
9.82
0.0165
10.01
9.34 9.45
4.04 9 10-3
8.53
-3
8.78
ZH CL = 20.96 mM CL = 5.77 mM
0.1183 0.1313 0.1443 0.1705
10.46
0.1452
10.67
10.13
0.1971
10.88
10.32 10.49
0.2755
10.94 10.99
10.63
0.4371
11.05
10.69
0.4713
11.11
0.3600
10.70
0.1969 0.2236 0.2506 0.2778 0.3061 0.3345 0.3634 0.3929 0.4231 0.4541 0.4861
0.4770
1
ZH
0.2
1.05 3.5
deviations, was obtained by numerical treatment [21]. Results were reported in Table 5. No direct comparison with K1-1 and K2-1 values obtained by others can be made due to differences in the experimental conditions [22, 23].
0 8.5
0.5
9.5
Extrapolation to infinite dilution reference state
0 8
10
-log [H+]
12
Fig. 2 ZH as a function of -log [H ?] for I = 1.05 and 3.5 mol/kg NaClO4
with model curves [20]. The best agreement was found with the constants given in Table 4. The uncertainty on the graphical values was evaluated taking into account the shift along x axes that still gave an acceptable fit. The probable composition of the species, existing in minor concentration and responsible for the
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The magnitude of the constants valid at zero ionic strength was evaluated by assuming the validity of the SIT. According to the theory the activity coefficients (ci), of the i species with charge zi can be expressed at 25 C in aqueous solutions according to Eq. (8): X eði; kÞmk ð8Þ logci ¼ z2i D þ where D = 0.51(I0.5)/[1 ? 1.5(I0.5)] and e was the specific interaction coefficient of i with species k of molality mk. Interaction coefficients depend on the ionic strength but the variation in the range 0.5 B I B 3.5 molal is sufficiently low that they may be assumed as constants. For this reason e(H?, ClO4-) was assumed from Ref. [16]. As a further simplification, interaction coefficients of ions with the
Solubility and acidic constants at 25 C in NaClO4 aqueous solutions of 1-(2-hydroxyphenyl)… -7.5
Table 4 Graphical values of log K1-1 -1
I/mol kg
NaClO4
Log K1-1
0.51
-10.4 ± 0.2
1.05
-10.41 ± 0.01
2.20
-10.7 ± 0.2
3.50
-11.1 ± 0.1
Table 5 Survey of the log K1-1 and log K2-1 values by numerical methods I/mol kg-1 NaClO4
(log K1-1 ± 3r)
1013
Y2-1
-8
-8.5
-9
(log K2-1 ± 3r) -9.5
0.51
-10.1 ± 0.2
-8.0 ± 0.2 -8.73 ± 0.09
1.05
-10.550 ± 0.004
2.20
-11.17 ± 0.05
-9.2 ± 0.2
3.50
-11.672 ± 0.008
-9.44 ± 0.03
-10
I /m 0
0.5
1
1.5
2
2.5
3
3.5
4
Fig. 4 Graphical extrapolation at the infinite dilution reference state of log K2-1
-10
Y1-1
Table 6 Results of extrapolation to zero ionic strength -10.5
-11
-10.11 ± 0.05 -7.85 ± 0.05
e(Y-, Na?)
0.50 ± 0.05
e(HY2-, Na?)
0.59 ± 0.05
Conclusion
-11.5
-12
Log K1-1 log K2-1
I /m 0
0.5
1
1.5
2
2.5
3
3.5
4
Fig. 3 Graphical extrapolation at the infinite dilution reference state of log K1-1
same charge type are nearly zero. The variation of the various equilibrium constants determined in this work can be expressed as reported below [Eqs. (9) and (10)]: þ logK11 ¼ logK11 2D þ ½e Hþ ; ClO 4 þ eðNa Y Þm logcHY ð9Þ þ logK21 ¼ logK 21þ 2Dþ ½e H ; ClO4 þ e Na ; HY2 m 2 logcHY
ð10Þ
Hence, plots on the known terms of Eqs. (9) and (10) as a function of I (Y1–1 and Y2–1 in Figs. 3 and 4, respectively) result in straight lines in which the constants at zero ionic strength, log K1-1 and log K2-1, were the intercepts while e(Na?, Y-) and e(Na?, HY2-) were the slopes. From the best lines through the points the constants and the specific interaction coefficients given in Table 6 were found.
The solubility and the acidic constants of 1-(2-hydroxyphenyl)ethanone were determined, at 25 C, in NaClO4 solutions of ionic strength ranging from 0.51 to 3.5 mol/kg and at the infinite dilution reference state. These values were used to estimate the salting effect of NaClO4 on the neutral molecule as well as the interaction coefficients e(i, k). Results obtained in this work can be used to evaluate the complex formation equilibria between HY and bioavailable metal cations.
Experimental Sodium perchlorate, perchloric acid, and sodium hydroxide stock solutions were prepared and standardized as described previously [7]. 1-(2-Hydroxyphenyl)ethanone (98 %, Sigma-Aldrich) was used without further purification. All solutions were prepared with bidistilled water; CO2 was removed from the water before the preparation of the solutions using the nitrogen gas. UV–Vis measurements The spectrophotometric measurements were conducted with a Varian Cary 50 Scan UV–Visible Spectrophotometer.
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E. Furia et al.
Absorbance values between 190 and 400 were measured every 2 nm. The temperature of the cell-holder was kept at (25.0 ± 0.3) C by a Grant circulating water bath. Matched quartz cells of thickness 1 cm were employed. The absorbance (Ak), was recorded to ±1 9 10-4 units. The formulations of the parameters and the acquisition of the data were managed with the aid of a computer connected to the instrument.
determination of the EG value, carried out by measuring the electromotive force of the cell (G) depending on the acidity decreased by addition of OH- ions, coulometrically generated at constant current intensity with the electrolysis circuit (C): Pt=TS=AuxiliaryElectrodeþ
In these conditions the free hydrogen ion concentration (h), was expressed by Eq. (12):
Potentiometric measurements
h ¼ CA CB The cell arrangement was similar to the one described by Forsling [24] and Ag/AgCl electrodes were prepared according to Brown [25]. The alkaline glass electrode was manufactured by Metrohm and acquired a constant potential within 20 min after the addition of the reagents. The titrations were carried out with a programmable computer controlled data acquisition switch unit 34,970 A supplied by Hewlett and Packard. The EMF values were measured with a precision of ±10-5 V using an OPA 111 low-noise precision DIFET operational amplifier. Test solutions were purged with nitrogen gas, taken from commercial cylinders and purified by passage through 1 M H2SO4, 1 M NaOH, bidistilled water and I mol/kg NaClO4. The protolysis equilibria were studied in a thermostatic bath at the temperature of (25.0 ± 0.1) C through measurements carried out in the form of potentiometric titrations using an alkaline glass electrode sensitive to variations of the free hydrogen ion concentration (h), in solutions in which the ionic strength was maintained constant by adding I mol/kg NaClO4 (i.e., 0.51, 1.05, 2.20, and 3.50). Test solutions (TS), had the following general composition: TS ¼ CL mol=kgHY; CA mol=kgHClO4 ; CB mol=kgNaOHðI CA CB Þmol=kgNaClO4 where CA and CB were the analytical concentrations of HClO4 and NaOH, respectively, and CL was the analytical concentration of 1-(2-hydroxyphenyl)ethanone. The total concentration (CL), of ligand was ranged from 1 9 10-3 mol/kg to 20 9 10-3 mol/kg while pH ranged from 8 to 11.5. The free hydrogen ion concentration (h), was measured with cell (G): Reference Electrode=TS=Glass Electrode
ðGÞ
The electromotive force of the cell (G) can be written according to Eq. (11): EG ¼ EG þ 59:16logh þ Ej
ð11Þ
where EG was constant for each series of measurements and Ej was the liquid junction potential which was a function of [H?] only. The values of j (mV/M) at the various I were taken from literature [1]. Each titration was divided into two parts; the first was necessary for the
123
ðCÞ
ð12Þ
The EG data were treated with the Gran method [26] to refine the analytical concentration CA and the standard potential EG. Taking into account the Eq. (13): EG ¼ EG 59:16logh
ð13Þ
h was varied to obtain constant values of EG within the experimental uncertainty (±0.1 mV). In the second part of each titration, measurements were performed in the presence of ligand for the determination of the acidic constants. In this case the acidity was gradually decreased by adding well-known volumes (VT), of NaOH standard solution. For each level of investigated ionic strength two titrations were performed.
References 1. 2. 3. 4. 5. 6. 7. 8. 9. 10. 11. 12. 13. 14. 15. 16. 17. 18. 19. 20. 21. 22. 23. 24. 25. 26.
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