Acid–base and coordination properties of 2-phenyl-3-hydroxy-4-quinolones in aqueous media

Abstract

The acid–base and coordination properties of 2-phenyl-3-hydroxy-4(1H)-quinolone (1) and 1-methyl-2-phenyl-3-hydroxy-4(1H)-quinolone (2) were characterized by potentiometric, UV-Visible and fluorescence titrations in water containing 5 or 30% vol MeCN and in a micellar solution of a cationic surfactant. The first dissociation constants (pK_a1) corresponding to OH deprotonation of 1 and 2 are about 10 and ligand 1 undergoes a second NH deprotonation with a pK_a2 about 12, which is reduced to 10.4 in the presence of a cationic surfactant. More detailed complexation studies were performed with more soluble ligand 1, which forms stable complexes of 1 : 1 and 1 : 2 compositions with Fe(III), Cu(II), Zn(II), Pb(II) and Me₂Sn(IV) cations in neutral solutions. The most unusual behavior is observed with Zn(II), which strongly promotes NH deprotonation of ligand 1 with formation of the Zn(L)₂²⁻ complex at a pH about 8. The formation of this complex is confirmed by the results of ¹H NMR titrations in DMSO-d₆. Binding of all cations is accompanied by the appearance of a new absorption band in the range 385–405 nm with concomitant disappearance of the band at 350–360 nm in the free ligand. Interactions of 1 and 2 with Zn(II) and Me₂Sn(IV) are accompanied by strong and selective fluorescence enhancements with the blue shift of the emission bands allowing ratiometric detection of these cations. Complexation with transition and heavy metal ions as well as with lanthanides induces fluorescence quenching. Ligand 2 is characterized by X-ray crystallography.

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Introduction

Quinolone compounds, both natural and synthetic, possess important biological and medicinal properties including anticancer and antimicrobial activities.¹ A particular group of these compounds constitutes 3-hydroxy-4-quinolones, among which 2-heptyl-3-hydroxy-4-quinolone (PQS) serves as a quorum sensing signal molecule^2,3 and 2-phenyl-3-hydroxy-4-quinolones such as 1 and 2 are known to possess anticancer activity, immunosuppressive properties and to act as enzyme inhibitors.⁴ 3-Hydroxy-4-quinolones can be considered as aza-analogues of 3-hydoxyflavones. In particular 1 and 2 share common structural features with 3-hydroxyflavone (3) and 3-hydroxy-4-pyridinones such as Deferiprone (4). The presence of quinolone fragment brings to these compounds interesting fluorescence properties⁵ typical for flavones and the presence of 3-hydroxyl group allows the metal ion coordination. The presence of nitrogen atom in 1 and 2 instead of oxygen in 3 should increase the negative charge on carbonyl group and consequently enhance the affinity to metal ions. Indeed reported log β₁ = 14.6 for complexation of Fe(III) with 2-methyl-3-hydroxy-4-quinolone³ is larger than log β₁ = 13.3 for complexation with 3⁶ and is approaching that for 4 (log β₁ = 15.01),⁷ which is one of the most powerful iron(III) chelators. Also calculations of the electrostatic potentials for the ground state of 1 and 3 using DFT at the PBE0/G-31+G(d,p) level of theory (Fig. S1, ESI†) predict a significantly increased negative charge on 4-carbonyl group of 1 as compared to 3 with very little change in the electronic density on 3-OH group.

Although 3 possesses rather modest metal ion affinity studied mostly in organic solvents^8,9 it finds applications for fluorescence determination of different metal ions such as Al(III) and organotin(IV) compounds¹⁰ and metal complexes of 3-hydroxyflavonoles were tested for their biological activities.¹¹ Not surpringly metal complexes of 3-hydroxy-4-quinolones, which should be more powerful ligands have attracted significant attention. A study of cytotoxicity as well as DNA binding and cleavage activity of mixed ligand Cu(II) complexes of 1 showed promizing results.¹² Cytotoxic activities of Ru^II(arene) complexes of 1 were studied in comparison with related 3-hydroxyflavone complexes.¹³ The Zn(II) complex of 1 was prepared and characterized by its cytotoxic activity and spectroscopic properties.¹⁴ Crystal structures of several Cu(II),^12,15 and a Zn(II)¹⁴ complexes of 1 were reported demonstrating expected type of metal ion coordination through deprotonated 3-hydroxyl and 4-carbonyl groups. At the same time no solution studies of complexation of 1 and 2 or their derivatives were reported yet. Such studies are of interest for several reasons. First, the biological activity of 3-hydroxy-4-quinolones is related to their ability to bind intracellular inorganic cations.¹⁶ Second, due to their intense fluorescence they can find applications for sensitive and possibly selective analytical detection of metal ions.¹⁷ Third, the metal complexes of 3-hydroxy-4-quinolones may by themselves serve as optical sensors for anions (for anion sensing with flavonol complexes see ref. 18 and 19).

In this paper we report a detailed study of acid–base and co-ordination properties of 1 and 2 in aqueous media. A very low solubility of both compounds in water represented a serious obstacle for experimental measurements. With more soluble 1 it was possible to obtain reproducible spectrophotometric and fluorescence results with 5% vol MeCN, although for potentiometric titrations it was necessary to increase the organic co-solvent content to 30% vol. With less soluble 2 results obtained under such conditions were poorly reproducible. We found, however, that both ligands could be studied without problems in the presence of cationic surfactant hexadecyltrimethylammonium bromide (HTAB) at 5 mM concentration, which is well above its critical micelle concentration. Using of micellar media in this case not only solve the solubility problems, but is of interest by itself because being poorly soluble and highly hydrophobic compounds 1 and 2 and their analogs applied in vivo would be most probably bound to cellular membranes or proteins. In this context it is worth noting that liposomal solubilization of 3-hydroxyquinolones was suggested for their in vitro and in vivo testing²⁰ and PQS as well as other related quorum sensing molecules form micelles with very low critical micelle concentrations (9 μM for PQS) in aqueous solutions.²¹

Results and discussion

Crystal structures of the ligands

The ligand 1 has been characterized by the crystal structure previously.¹² Here we report in addition the crystal structure of 2, Fig. 1. Compound 2 crystallizes in triclinic P1̄ space group containing one molecule in the asymmetric unit. A comparison with 1 reveals one significant conformational difference: the torsion angles for the N–C1–C11–C12 fragment in 1 and 2 are 38.96° and 64.22°, respectively. The larger angle in 2 can be attributed to the steric hindrance induced by the N-methyl group. The C1–C2 and C3–O2 distances are significantly smaller in 2 than in 1, indicating a stronger quinoid character in the former molecule. As expected, the carbonyl and hydroxyl groups constitute a hydrogen bond acceptor and donor moieties in both molecules (Fig. 1B). Crystal data and structure refinement for 2 are shown in Table S1 (ESI†).

Fig. 1 (A) ORTEP diagram for 2. Ellipsoids are shown at the 50% probability level. (B) Hydrogen bond interactions in 2.

Acid dissociation constants

Due to solubility problems (see Introduction) the acid dissociation constants of 1 were determined in 5% and in 30% vol MeCN as well as in 5 mM HTAB while the dissociation constants of 2 were measured only in 5 mM HTAB solution. In the following discussion the medium containing 5% MeCN will be labeled as “aqueous solution” and the medium containing 5 mM HTAB as “micellar solution” for brevity.

Fig. 2 shows the absorption spectra of 1 recorded at variable pH. In both media increase in pH induces appearance of a new red shifted absorption band around 400 nm and disappearance of the band at 355–360 nm. In aqueous solution (Fig. 2A) these changes occur with preservation of four isosbestic points indicating co-existence of only two forms of the compound, but in micellar solution (Fig. 2B) the isosbestic points are not preserved. The absorbances vs. pH profiles at fixed wavelengths shown in inset of the Fig. 2B clearly demonstrate at least two deprotonation processes.

Fig. 2 UV-Visible spectra of 40 μM 1 at variable pH (25 °C and 0.05 M NaCl): (A) in 5% MeCN, pH 8.3–11.4; (B) in 5 mM HTAB, pH 6.5–11.5. Insets show absorbance vs. pH profiles at selected wavelengths; solid lines are the theoretical fitting profiles to the eqn (1) or (2). Arrows show directions of the spectral changes on increase in pH.

The fitting of the titration results to HypSpec allowed us determination of the respective pK_a1 and pK_a2 values given in Table 1.

Table 1 Dissociation constants of 1 and 2 obtained from UV-Vis and fluorescence pH-titrations at 25 °C and 0.05 M NaCl

Ligand	Medium	pKa1		pKa2		Ref.
		UV-Vis	Fl	UV-Vis	Fl
a HTAC is hexadecyltrimethylammonium chloride.
1	5% MeCN	10.2 ± 0.1	10.8 ± 0.1			This work
1	5 mM HTAB	9.1 ± 0.1	9.15 ± 0.06	10.4 ± 0.1	9.7 ± 0.1	This work
2	5 mM HTAB	9.65 ± 0.08	9.97 ± 0.05			This work
3	50% EtOH/water	9.99				6
3	5 mM HTAC^a	6.72				18
4	Water	9.82				7

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The results of fluorescence pH-titrations of 1 in both media are shown in Fig. 3A and B. In aqueous solution two emission bands are observed: one at 494 nm and another one at 412 nm. This type of the emission spectrum was interpreted as a result of excited state intramolecular proton transfer with the band at longer wavelengths belonging to emission from the tautomeric form with proton transferred from 3-OH group to 4-carbonyl group and the band at shorter wavelength belonging to “normal” structure.^5a In micellar solution the intensity of the emission from the “normal” form strongly decreases. Deprotonation of the ligand quenches the fluorescence and again the titrations profiles correspond to mono-deprotonation in aqueous solution, but double deprotonation in micellar solution clearly seen in the pH-profile of the fluorescence intensity at 470 nm (inset in Fig. 3B). The respective pK_a values are given in Table 1.

Fig. 3 Fluorescence spectra of 40 μM 1 at variable pH: (A) in 5% MeCN, (B) in 5 mM HTAB. Insets show fluorescence vs. pH profiles at selected wavelengths. Arrows show directions of the spectral changes on increase in pH. The excitation wavelength is 373 nm (A) and 385 nm (B).

Fig. 4 shows spectrophotometric and fluorescence pH-titrations of 2 in micellar solution. The changes in UV-Vis spectra induced by deprotonation of this ligand are similar to those observed with 1, but fluorescence of 2 shows a more complex trend: the band at 427 nm disappears while the band at 520 nm becomes more intense. This observation agrees with the absence of tautomeric forms in the deprotonated molecule.

Fig. 4 UV-Visible (A) and fluorescence (B) spectra of 40 μM 2 at variable pH (6.5–11.5) in 5 mM HTAB. Insets show absorbance or fluorescence vs. pH profiles at selected wavelengths. Arrows show directions of the spectral changes on increase in pH. The excitation wavelength is 385 nm.

The results of all spectroscopic titrations are summarized in Table 1 together with some relevant literature data. Spectrophotometrically determined pK_a1 of 1 in aqueous solution is close to pK_a values reported for 3 and 4. A larger value found from fluorescence titration probably involves a contribution from excited state dissociation with shifted pK*_a value. In micellar solution one observes decreased pK_a values due to stabilization of the anionic deprotonated form of 1 by the positive surface charge of the cationic micelle. Similar effect was reported for 3 (see Table 1, lines 4 and 5). Surprisingly we observed a second deprotonation process for 1 in micellar solution, which can be attributed to the deprotonation of N–H group (Scheme 1). Previously the N-deprotonation of 3-hydroxy-2(1H)-pyridinone ligands was observed in their Fe(III) complexes,²² but for a free ligand this type of dissociation is unusual. However, the quantum chemical calculations using a PBE0/6-31+G(d,p)/IEF-PCM level of theory (two water molecules were included in the model in order to take into account explicit interactions) predicted that NH deprotonation in 1 indeed is quite feasible with calculated pK_a1 = 9.1 and pK_a2 = 10.5 close to experimental values obtained in micellar solution.

Scheme 1

Metal-ion complexation studies

Interactions of the ligands with metal ions were studied by potentiometric and spectroscopic titrations. With more soluble 1 potentiometric titrations requiring relatively high concentrations of components above 1 mM were performed in 30% vol MeCN. With 2 the precipitation occurred even in this medium. Titrations in the presence of HTAB, which prevented precipitation, did not give reproducible results. Therefore with this ligand only spectroscopic titrations in micellar solution employing less than 0.1 mM 2 were performed.

The results of potentiometric titrations were analyzed in terms of traditional pqr scheme expressed by eqn (1) and (2), where L is a completely deprotonated dianionic form of the ligand and M is the metal ion. The overall binding constants and the appropriate pK_a values of free 1 and complexes are collected in Table 3.

pM + qL + rH ⇄ M_pL_qH_r (1)

β_pqr = [M_pL_qH_r]/[M]^p[L]^q[H]^r (2)

Titration of the free ligand 1 confirms the existence of two deprotonation processes with pK_a1 = 10.66 close to the value found in 5% MeCN (Table 1) and very high pK_a2 = 11.9 not detected spectroscopically without surfactant.

In case of Zn(II) two complexes are formed: one of the M(LH)⁺ composition with mono-deprotonated ligand, which correspond to the reported crystal structure of the isolated Zn(II) complex of 1,¹⁴ and another one of the ML₂²⁻ composition with doubly deprotonated ligand (see Scheme 2). The first complex is dominating at pH about 7 and the second complex is dominating at pH equal and above 8 (see Fig. S2 in ESI† for the calculated species distribution profile).

Scheme 2 Types of metal complexes of 1 (LH₂) established by potentiometric titrations.

In order to prove formation of the complex with doubly deprotonated ligand interaction of 1 with Zn(II) was followed by ¹H NMR in DMSO-d₆. The spectrum of the mixture of 1 with ZnCl₂ at molar ratio 2 : 1 in DMSO (Fig. 5A, bottom) coincides with the spectrum of free ligand indicating the absence of interaction. Additions of 0.5 and 1 equivalents of Et₃N respective to 1 to the mixture induce changes in the signals of aromatic protons and a downfield shift with broadening in the signal of NH proton (spectra 2 and 3, Fig. 5A). At the same time the signals of Et₃N appear at the positions corresponding to protonated amine (signals of CH₂ and CH₃ groups of free base are observed at 2.41 and 0.93 ppm while corresponding signals of the protonated form at 3.08 and 1.17 ppm respectively). This behavior is consistent with predominant formation of the complex with a mono-deprotonated form of the ligand. When more than 1 equivalent of Et₃N is added the signal of NH proton starts to disappear and the signals of ethyl groups of Et₃N start to move upfield towards positions of the signals of free base, although still remain closer to the signals of the protonated form. Assuming that the observed chemical shifts of methyl and methylene groups of triethylamine are the weighted averages of the signals of free base and protonated forms, as should be in the case of the fast exchange, we calculated the concentration of the protonated form as a function of total triethylamine concentration shown in Fig. 5B as a profile of [Et₃NH⁺] vs. [Et₃N]_Total. It demonstrates that formation of Et₃NH⁺ reaches the limiting concentration of 20 mM, which is exactly twice the total concentration of 1. This observation together with disappearance of the NH signal proves the complete deprotonation of the ligand in the presence of Zn(II) at the molar ratio 1 : 2 and therefore formation of the complex of the ML₂²⁻ type.

Fig. 5 (A) ¹H NMR spectra of the mixture of 5 mM ZnCl₂ and 10 mM 1 in DMSO-d₆ in the presence of increased amounts of Et₃N. Intervals of chemical shifts containing solvent signals (2.5 ppm for DMSO and 3.3 ppm for traces of water) are eliminated. (B) Concentration of the protonated form of Et₃N as a function of total Et₃N concentration added to a mixture of 1 and Zn(II).

Results for Cu(II) showed that in this case the formation of a 1 : 1 complex was negligible. The predominant form in neutral solution is a 1 : 2 complex of the composition M(LH)₂ which at high pH is transformed with pK_a = 10.7 into M(LH)L⁻ containing one mono and one doubly deprotonated ligands. This pK_a is only 1.2 units lower than the pK_a2 of the free ligand. At the same time formation of the complex ZnL₂²⁻ implies much stronger reduction in pK_a2 of 1 on coordination to Zn(II). The type of equilibria involved in formation of Zn(II) complexes does not allow one to estimate the pK_a2 of the coordinated ligand, but the species distribution profile in Fig. S2 (ESI†) indicates that it must be below 8. It seems therefore that more electrophilic Cu(II) cation surprisingly is less effective in inducing the second deprotonation of the ligand than Zn(II).

With Me₂Sn²⁺ complexes of three types are observed (Table 2). The simple 1 : 1 complex of the type M(LH)⁺ is much more stable with this cation than with Zn(II) and is a dominating species in acid and neutral solutions. With excess of the ligand and pH about and above 7 the M(LH)₂ complex is the dominating species, although it has lower stability than the respective Cu(II) complex (see Table 2). At higher pH this complex is transformed with pK_a = 10.77 into M(LH)L⁻ species indicating similar effects of Me₂Sn(IV) and Cu(II) on the acidity of NH group.

Table 2 Protonation constants of ligands and stability constants of metal complexes determined potentiometrically in 30% vol MeCN, 0.05 M NaCl at 25 °C

Species	log βpqr	Equilibrium	log K
LH	11.90 ± 0.01	L + H = LH	11.90
LH2	22.56 ± 0.01	LH + H = LH2	10.66
ZnLH	18.94 ± 0.04	Zn + LH = Zn(LH)	7.04
ZnL2	22.0 ± 0.1	Zn + 2L = Zn(L)2	22.0
CuL2H2	46.1 ± 0.1	Cu + 2LH = Cu(LH)2	22.4
CuL2H	35.4 ± 0.1	Cu(LH)2 = Cu(LH)L + H	−10.7
SnMe2LH	25.12 ± 0.09	SnMe2 + LH = SnMe2(LH)	13.22
SnMe2L2H2	45.47 ± 0.08	SnMe2 + 2LH = SnMe2(LH)2	21.67
		SnMe2(LH) + LH = SnMe2(LH)2	8.45
SnMe2L2H	34.7 ± 0.1	SnMe2(LH)2 = SnMe2(LH)L + H	−10.77

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Spectroscopic, UV-Vis and fluorescence, titrations of 1 were performed with a large set of metals at pH 7 with the purpose to estimate the selectivity of complexation and to evaluate the applicability of 1 for optical sensing of metal ions. The same reaction media, 5% MeCN and 5 mM HTAB were employed as for the spectroscopic pK_a determinations. Although measurements were performed in media different from that employed for potentiometric titrations the complexation stoichiometry should be similar and also there should be at least an approximate agreement between observed stability constants (K_obs) experimentally determined from spectroscopic data at a fixed pH value and the values of K_obs calculated for a given pH from potentiometric results. As will be shown below such agreement indeed generally is confirmed.

Among tested cations additions of alkaline-earth cations (Mg²⁺ and Ca²⁺) and Bi(III) did not change neither UV-Vis no fluorescence spectra of 1. Visual effects of other 11 cations on florescence of 1 at pH 7 are compared in Fig. 6. “Naked eye” detectable effects are observed with Fe(III) and Cu(II) inducing the fluorescence quenching and with Zn(II) and Me₂Sn(IV) inducing blue shifts of the emission bands.

Fig. 6 Visual effects of metal ions (40 μM) on florescence of 40 μM 1 at pH 7 in 5% MeCN. Excitation at 356 nm.

Fig. 7 illustrates titration results with Zn(II) in aqueous solution. Essentially similar results were obtained in micellar solution (Fig. S3, ESI†). The fitting of the titration results to HypSpec revealed formation of a 1 : 1 complexes with the logarithms of the observed stability constants K_obs given in Table 3.

Fig. 7 UV-Visible (A) and fluorescence (excitation at 370 nm) (B) spectra of 40 μM 1 in 5% MeCN at pH 7 and variable concentration of Zn(II). Insets show absorbance or fluorescence vs. Zn(II) concentration profiles at selected wavelengths. Arrows show directions of the spectral changes on increase in Zn(II) concentration.

Table 3 Observed stability constants (log K_obs) and complexation-induced changes in absorption and emission maxima of 1 at pH 7 in 0.05 M NaCl at 25 °C^a

Metal ion	5% vol MeCN					5 mM HTAB
	log K11	log K12	λmax abs^b	λmax fluor	I/I0^c	log K11	log K12	λmax abs^b	λmax fluor	I/I0^c
a Mean values between stability constants determined by spectrophotometric and fluorescence titrations; relative error less or equal to ±0.05.b The longest wavelength absorption band.c The ratio of fluorescence intensities at saturation and in the absence of metal ion at the wavelength of the emission maximum of the complex.
None			353	494				359	508
Mg^2+	No interaction
Ca^2+	No interaction
Mn^2+	2.45		400	494	0.920	2.04		406	507	0.942
Fe^3+	5.97	9.71	415	Quenching	0	6.33	11.83	396	Quenching	0
Co^2+	3.08		402	495	0.590	3.41		405	506	0.717
Ni^2+	3.55		402	494	0.318	3.60		405	506	0.457
Cu^2+		12.29	405	Quenching	0		13.21	404	Quenching	0
Zn^2+	3.76		384	471	2.894	3.54		398	472	3.581
Zn^2+ (pH 8)	4.32		385	472	3.063	4.02	9.61	409	473	3.105
Cd^2+	2.73		401	488	1.066	3.02		403	503	1.010
Me2Sn^2+	4.41		422	452	12.316	4.4	9.8	394	452	8.354
Hg^2+	<2		398	494	0.646	<2		359	507	0.952
Pb^2+	5.68	9.62	383	488	0.173	4.39	10.47	384	502	0.065
Eu^3+	4.75		390	493	0.130	4.25		386	506	0.097
Bi^3+	No interaction

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Since potentiometric results indicated a change in the stoichiometry at increased pH, titrations were repeated also at pH 8. In aqueous solution only a tighter binding of the same 1 : 1 stoichiometry was observed, but in micellar solution expected change to predominantly 1 : 2 metal to ligand complexation was confirmed (Fig. 8): the UV-Vis titration plot (Fig. 8A, inset) clearly shows the saturation at 1 : 2 molar ratio and the fluorescence (Fig. 8B) initially drops down due to formation of a less fluorescent 1 : 2 complex and then with excess of the metal ion appears a more intense band at 471 nm characteristic for the 1 : 1 complex.

Fig. 8 UV-Visible (A) and fluorescence (excitation at 376 nm) (B) spectra of 40 μM 1 in 5 mM HTAB at pH 8 and variable concentration of Zn(II). Insets show absorbance or fluorescence vs. Zn(II) concentration profiles at selected wavelengths. Arrows show directions of the spectral changes on increase in Zn(II) concentration.

In the pH range 7–8 the actual forms of the reactants are Zn²⁺ and LH₂ and the reaction of complex formation for the M(LH)⁺ species is given by the eqn (3).

Zn²⁺ + LH₂ = Zn(LH)⁺ + H⁺ (3)

The respective expression for K_obs takes the form of the eqn (4).

log K_obs = log β₁₁₁ − log β₀₁₂ + pH (4)

With stability constants given in Table 2 one obtains log K_obs = 3.38 and 4.38 at pH 7.0 and 8.0 respectively, which are reasonably close to the experimental values in both aqueous and micellar solutions given in Table 3.

Similar analysis for formation of the complex ML₂²⁻ predicts that the respective expression for K_obs takes the form of the eqn (5).

log K_obs = log β₁₂₀ − 2 log β₀₁₂ + 4pH (5)

It follows from the eqn (5) that log K_obs = 8.88 and 4.88 at pH 8 and 7 respectively. The experimental K_obs at pH 8 in micellar solution is even larger than predicted (Table 3), but at pH 7 the complex is not detected at all. This agrees with extremely sharp pH-dependence of K_obs in accordance with eqn (5).

Titration results for Cu(II) demonstrated similar to Zn(II) changes in UV-Vis spectra, but complete quenching of the fluorescence in both media (Fig. S4, ESI†). Fitting of the concentration profiles confirmed formation of only 1 : 2 complexes in agreement with results of potentiometric titration. The expression for K_obs for formation of expected in this case M(HL)₂ complex takes the form of the eqn (6), which predicts log K_obs = 14.98 at pH 7.0. Somewhat lower experimental values (Table 3) can be attributed to a difference in solvent composition.

log K_obs = log β₁₂₂ − 2 log β₀₁₂ + 2pH (6)

Titration results for Me₂Sn(IV) are shown in Fig. S5 (ESI†) (aqueous solution) and 9 (micellar solution).

The spectral changes observed in the presence of Me₂Sn(IV) resemble those in the presence of Zn(II), but are more pronounced in the fluorescence. This is manifested also in visually much more pronounced blue shift in the fluorescence of 1 in the presence of this cation (see Fig. 6). Fitting of the titration results by using the HypSpec revealed formation of 1 : 1 complexes in aqueous solution, but both 1 : 1 and 1 : 2 complexes in micellar solution with K_obs collected in Table 3. The pattern of the fluorescence changes in micellar solution resembles that observed with Zn(II) at pH 8 (cf. Fig. 8B). The green fluorescence of the free ligand initially disappears due to formation of a less fluorescent 1 : 2 complex and only with excess of the metal ion one observes appearance of the more intense blue band at 452 nm characteristic for the 1 : 1 complex (see inset in Fig. 9B). Although the stability constant for Me₂Sn(HL)⁺ complex is 10⁶-fold larger than that for Zn(HL)⁺ (Table 2), the K_obs values for these complexes differ less than by one order of magnitude (see Table 3). The reason for this effect is a strong hydrolysis of Me₂Sn²⁺ cation in neutral solutions. In accordance with reported hydrolysis constants in water (pK_a1 = 3.12, pK_a2 = 5.33, log β₁₀₋₂ = −8.45)²³ at pH 7 the cation is completely transformed into dihydroxo complex and the actual reactions of the formation of M(HL)⁺ and M(HL)₂ complexes are (7) and (8) respectively.

SnMe₂(OH)₂ + LH₂ + H⁺ = SnMe₂(LH)⁺ + 2H₂O (7)

SnMe₂(OH)₂ + 2LH₂ = SnMe₂(LH)₂ + 2H₂O (8)

Fig. 9 UV-Visible (A) and fluorescence (excitation at 375 nm) (B) spectra of 32 μM 1 in 5 mM HTAB at pH 7 and variable concentration of Me₂SnCl₂. Inset in A shows absorbance vs. Me₂SnCl₂ concentration profile at a fixed wavelength. Inset in B shows visual effects of increased concentrations of Me₂SnCl₂ (0; 0.25; 0.5; 0.75; 1 and 3 equivalents) on the fluorescence of 1. Arrows show directions of the spectral changes on increase in Me₂SnCl₂ concentration.

The corresponding values for K_obs at pH = 7 are log K_obs = log β₁₁₁ − log β₀₁₂ − log β₁₀₋₂ − pH = 4.01 for SnMe₂(LH)⁺ and log K_obs = log β₁₂₂ − 2 log β₀₁₂ − log β₁₀₋₂ = 8.80 for SnMe₂(LH)₂. Comparison with data in Table 3 shows that both calculated K_obs are reasonably close to the respective experimental values.

Spectroscopic titrations of 1 with other metal ions are illustrated in Fig. S6–S13 (ESI†) and the K_obs values together with spectral characteristics of the complexes determined from these results are collected in Table 3. An inspection of the Table 3 reveals the following general features of the complexation processes.

The general stability sequence for divalent transition metal cations follows the Irving–Williams series Mn(II) < Co(II) < Ni(II) < Cu(II) > Zn(II). Very low observed stability for Hg(II) most probably results from its strong hydrolysis and strong binding of this cation to halide anions from the reaction medium. Relatively small K_obs for Fe(III), which forms much more stable complex than Cu(II) with ligand 4,⁷ also is a result of strong hydrolysis of Fe(III) in neutral solution. Relative affinities of cations Cu(II), Zn(II), Cd(II) and Pb(II) are similar to those reported for mimosine and related ligands.²⁴ The micellar medium affects very little the stability of 1 : 1 complexes, but promotes formation of 1 : 2 complexes. This effect can be attributed to increased local concentration of the hydrophobic ligand in the micellar pseudo-phase.²⁵

In UV-Vis spectra of 1 the complex formation induces appearance of a new band in the range 385–405 nm with concomitant disappearance of the band at 353 nm (359 nm in micellar solution). Similar spectral change occurs on the deprotonation of the ligand. All metal ions besides Zn(II) and Me₂Sn(IV) induce fluorescence quenching particularly strong with Cu(II) and Fe(III). This is a typical behavior for transition metals. Coordination with Zn(II) induces blue shift and strong enhancement of the fluorescence. The effect of Me₂Sn(IV) depends on the stoichiometry of the complex: formation of a 1 : 2 complex leads to fluorescence quenching, but 1 : 1 complex has strong blue fluorescence. Similar behavior was observed previously with diphenyltin(IV) complexes of 3.¹⁸ The reason of the absence of the fluorescence of the 1 : 2 complex is not clear yet.

Interactions of metal ions with 2 were studied for comparative purposes with Zn(II), Cu(II) and SnMe₂(IV) only in micellar solution. UV-Vis titrations with Zn(II) at pH 7 and 8 (Fig. S14, ESI†) indicate formation of 1 : 1 complexes at both pH with log K = 3.6 and 4.2 respectively. No contribution of a 1 : 2 complex with this ligand incapable to produce a doubly deprotonated dianionic form is detected at increased pH. Strong 20-fold fluorescence enhancement with a blue shift in the emission maximum is observed, Fig. 10A. The fluorescence vs. Zn(II) concentration profile fits to the 1 : 1 complexation equilibrium with the same stability constant as that from UV-Vis results.

Fig. 10 Fluorescence (excitation at 372 nm) spectra of 40 μM 2 in 5 mM HTAB at pH 7 and variable concentration of Zn(II) (A) or Me₂SnCl₂ (B). Insets show fluorescence vs. metal ion concentration profiles at selected wavelengths. Arrows show directions of the spectral changes on increase in metal ion concentration.

Interaction of 2 with Me₂Sn(IV) was accompanied by spectral changes similar to those observed with Zn(II), Fig. S14 (ESI†) and 10B. Only a 1 : 1 complex was formed with log K₁₁ = 4.2 at pH 7.0. Additions of Cu(II) induced a complete quenching of the fluorescence of 2. Only one complex of the 1 : 2 stoichiometry was observed with log K₁₂ = 12.7 at pH 7. Thus, towards these two cations, 2 behaves similarly to 1. An important feature of the results shown in Fig. 10 is that together with a strong increase in the fluorescence intensity about 470 nm there are certain wavelengths (425 nm for Zn(II) and 406 nm for Me₂Sn(IV)) where the fluorescence is not changed. This makes possible highly sensitive ratiometric detection of these two cations.

Conclusions

The first dissociation constants of 1 and 2 corresponding to OH deprotonation of 2-phenyl-3-hydroxy-4-quinolones are close to those for 3-hydroxy-flavone and 1,2-dimethyl-3-hydroxy-4-pyridinone, however in contrast to 3-hydroxypyridinones the ligand 1 undergoes second NH deprotonation with pK_a about 12. Comparison of results obtained in water containing 5% vol MeCN and in micellar solution of a cationic surfactant demonstrates certain “micellar effect” manifested in promotion of the second deprotonation of free ligand and preferable formation of 1 : 2 rather than 1 : 1 metal complexes. In neutral solutions ligands 1 and 2 have highest affinities to Fe(III), Cu(II) and Pb(II) cations in terms of observed stability constants. The spectral characteristics of the complexation processes demonstrate possible sensitive and selective ratiometric fluorescence detection of Zn(II) and organotin(IV) compounds.

Experimental section

Materials and physical measurements

Commercially available starting materials, components of buffer solutions (CHES, MOPS, MES) and solvents were used as supplied. Infrared spectra were determined on an FTIR/FIR spectrometer 400 Perkin-Elmer 1600 series. Elemental analyses were carried out on a Perkin-Elmer 2400 CHNS Elemental Analyzer. Melting points were measured on an Electrothermal 9200 apparatus. ¹H and ¹³C NMR spectra were recorded at room temperature on a 300 MHz Varian unity spectrometer. Chemical shifts (ppm) are relative to (CH₃)₄Si. Mass spectrometry (MS-FAB+) was obtained by using a Thermo-Electron DFS. Measurements of pH were carried out using an Orion model 710-A research digital pH meter. Fluorescence and UV-Vis absorption spectra were measured on a FluoroMax spectrofluorometer from HORIBA Scientific and Thermo Scientific Evolution diode array UV-Vis spectrophotometer, respectively, equipped with a thermostated cell compartment (recirculating water bath at 25 ± 0.1 °C). Compounds 1 and 2 were synthetized as previously described.²⁶

2-Phenyl-3-hydroxy-4(1H)-quinolone (1). ¹H NMR (300 MHz; DMSO-d₆; Me₄Si): δ 11.64 (br. s, 1H, NH), 8.17 (d, J = 7.8 Hz, 1H, H-5), 7.81 (d, J = 7.1 Hz, 2H, H-12), 7.74 (d, J = 8.4 Hz, 1H, H-8), 7.61–7.52 (m, 4H, H-7, H-13, H-14), 7.28 (t, J = 7.8 Hz, 1H, H-6); ¹³C NMR (75 MHz; DMSO-d₆; Me₄Si): δ 170.5, 138.5, 138.3, 132.8, 132.1, 131.1, 129.7, 129.7, 128.8, 124.9, 122.5, 122.3 and 118.9. IR (ν_max/cm⁻¹) 3197 (O–H stretching), 2922, 2850, 2808 (C–H stretching), 1645 (C=O stretching). Anal. calc. for C₁₅H₁₁N₁O₂ DMF: C 69.66, H 5.85, N 9.03; found: C 69.71, H 5.23, N 9.34. MS (FAB, m/z) 237 [M]⁺; m.p. 276–279 °C.

1-Methyl-2-phenyl-3-hydroxy-4(1H)-quinolone (2). ¹H NMR (300 MHz; DMSO-d₆; Me₄Si): δ 8.15 (d, J = 8.5 Hz, 2H, ArH), 7.81 (d, J = 7 Hz, 2H, Ar′H), 7.72 (d, J = 8 Hz, 1H, ArH), 7.57 (m, 4H, ArH, Ar′H), 7.27 (t, J = 7 Hz, 1H, Ar′H), 355 (s, 3H, N–CH₃); ¹³C NMR (75 MHz; DMSO-d₆; Me₄Si): δ 169.7, 139.4, 139.3, 137.3, 132.8, 131.7, 130.3, 129.5, 129.0, 125.6, 123.65, 122.6, 117.5 and 37.5. IR (ν_max/cm⁻¹) 3387 (O–H stretching), 2925, 2862, 2811 (C–H stretching), 1693 (C=O stretching). Anal. calc. for C₁₆H₁₃N₁O₂: C 76.48, H 5.21, N 5.57; found: C 74.24, H 6.65, N 5.10. MS (FAB, m/z) 251 [M]⁺; m.p. 273–276 °C.

Spectrophotometric and fluorometric titrations

All titration experiments were performed at 25 °C and ionic strength 0.05 M created either by buffer or NaCl. The experiments were performed either with 5% vol MeCN or with 5 mM hexadecyltrimethylammonium bromide (HTAB). An aliquot of 1 mM stock solution of 1 or 2 in acetonitrile was added to a 5 or 50 mM MOPS, MES or CHES aqueous buffered solution in appropriate pH intervals, allowing equilibration for 15 minutes before titrations. The final content of acetonitrile was less than 1% when HTAB was used. The program HypSpec version 1.1.33 was used to calculate all equilibrium constants.

Potentiometric titrations

Potentiometric titrations were performed in a 10 mL thermostatted cell kept under nitrogen at 25 ± 0.1 °C with 0.05 M NaCl as background electrolyte in 30% vol MeCN with 3 mM 1 and molar ratio metal : ligand = 1 : 2. Experimental details and procedure for the electrode calibration were the same as in ref. 27 for potentiometric titrations in aqueous DMSO. The autoprotolysis constant of water pK_w = 14.6 ± 0.1 was determined in 30% MeCN in agreement with reported value in this medium.²⁸ The program Hyperquad 2008²⁹ was used to calculate all equilibrium constants.

X-ray crystallography

Crystals of 2 were grown by slow evaporation from a saturated DMSO solution of 2 at room temperature. Single-crystals of 2 were studied with Oxford Diffraction Gemini “A” diffractometer with a CCD area detector (λ_MoKα = 0.71073 Å, graphite monochromator, T = 293 K) source equipped with a sealed tube X-ray source. Unit cell constants were determined with a set of 15/3 narrow frame/runs (1° in ω) scans. Structure solution and refinement were performed with SHELX-2013 software,³⁰ and Mercury Crystal Structure Visualization software was used for molecular visualization.³¹ WinGX environment program set³² was used to prepare material for publication. Full-matrix least squares refinement was carried out by minimizing (Fo² − Fc²)². All non-hydrogen atoms were refined anisotropically. H atom attached to the oxygen atoms was located in a difference map. H atoms attached to C atoms were placed in geometrically idealized positions and refined as riding on their parent atoms, with C–H = 0.98–0.99 Å and U_iso (H) = 1.2U_eq (C) for methyl group.

Computational methods

Quantum chemical calculations were obtained by using DFT and TD-DFT with Polarizable Continuum Model³³ as performed in the Gaussian 09 code,³⁴ using a PBE0/6-31+G(d,p)/IEF-PCM (water) level of theory to determine the optimized molecular geometry of 1 and 3. Then, a frequency analysis corroborates that this geometry corresponds to an energy minimum, finding no imaginary frequencies. Zero point vibrational energies (ZPVE) were considered to account for thermal and entropic effects during pK_a calculations. As a first step in the analysis of the electron charge distribution in the molecules, the electrostatic potentials were computed to compare the local charge distribution in these molecules. To compute the pK_a values of 1 we use the Born–Haber method.³⁵ In order to determine the ΔG_solv the water solvent was modeled by an implicit (IEF-PCM) explicit solvent model (IE). In the IE approach two water molecules were included in order to model explicit interactions and its positions were fully optimized as well.

Acknowledgements

Arturo Jiménez-Sánchez thanks DGAPA-UNAM for the postdoctoral fellowship.

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Footnote

† Electronic supplementary information (ESI) available: Electrostatic potential mapped onto total electron density for 1 and 3, crystallographic data for compound 2, CCDC 1401533, spectrophotometric and fluorescence titrations of 1 with metal ions. For ESI and crystallographic data in CIF or other electronic format see DOI: 10.1039/c5ra10217k

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