Excited-State Proton Transfer and Proton Reactions of 6

Oct 20, 2009 - ... School of Chemistry, Tel Aviv University, Tel Aviv 69978, Israel ..... In the literature, there are large numbers of studies report...
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Excited-State Proton Transfer and Proton Reactions of 6-Hydroxyquinoline and 7-Hydroxyquinoline in Water and Ice I. Presiado, Y. Erez, R. Gepshtein, and D. Huppert* Raymond and BeVerly Sackler Faculty of Exact Sciences, School of Chemistry, Tel AViV UniVersity, Tel AViV 69978, Israel ReceiVed: August 20, 2009; ReVised Manuscript ReceiVed: September 30, 2009

Time-resolved and steady-state emission spectroscopies as well as absorption UV-vis spectroscopy were employed to study the photoprotolytic cycle and other protic processes of the bifunctional 6-hydroxy- and 7-hydroxyqunoline molecules in methanol-doped ice over a wide range of temperatures. In ice at high temperatures of T > 173 K, the excited-state proton transfer rate decreases as the temperature decreases. The emission band of the H+NRO-* zwitterion, where the imine nitrogen is protonated and the hydroxyl is deprotonated, is observed. At T > 173 K, the formation rate of the H+NRO-* emission band is approximately that of the decay rate of the neutral form, NROH*. Below 173 K, the rate of the photoprotolytic process is much slower than the radiative and the nonradiative rates, and the excited-state proton transfer could not be clearly observed. Addition of a small concentration of acetic acid increases the proton transfer rate significantly at temperatures below 235 K. The reaction rate in the presence of acetic acid is temperature-independent over a wide range of temperatures (80-235 K). We propose as an explanation for this observation that there exists a direct proton transfer from the hydroxyl group to water-acetic acid complexes at temperatures below 235 K. SCHEME 1

Introduction 1–13

For many years, intermolecular excited state proton transfer (ESPT) to a solvent or to a base in a liquid solution, and more recently in ice,14–16 has been widely researched. In the past decades, we extensively studied the reversible photoprotolytic cycle of a photoacid. We used a proton transfer model that explains the reversibility and accounts for the diffusion assisted geminate recombination of the transferred proton with the deprotonated form of the photoacid.7,17,18 Intramolecular proton transfer processes take place between a proton donating group (acid) and a proton accepting group (base) in the same molecule. 3-Hydroxyflavone (3HF) is a prototype system that shows an intramolecular proton transfer reaction following an electronic excitation.12,19–26 Besides the enol (E) form observed under different conditions, other H-bonding species can appear at the ground state producing H-bonded complexes19–21 or even the anionic form of 3HF.12 Photoexcitation of E produces, in less than 50 fs, a proton transferred form with a zwitterionic character. In solution (e.g., in acetonitrile or alcohol), another component, with a time constant of 5-10 ps, has been observed.22–25 Another class of molecules that undergo the photoprotolytic process is the hydroxyquinolines and similar compounds. In this class of molecules, bifunctional hydroxyquinoline (HQ) is characterized by a weak acidic hydroxyl functional group and a weak basic imine functional group in the ground state.27–33 In the electronic excited state, the acidity of hydroxyl and basicity of imine groups are considerably enhanced. HQs are present as cationic (H+NROH*), anionic (NRO-*), and neutral (NROH*) forms in acidic, alkaline, and neutral aqueous solutions due to the release of a proton to the solvent or the cleaving of protons from solvent molecules. * Corresponding author. E-mail: [email protected]. Telephone: 9723-6407012. Fax: 972-3-6407491.

These forms of HQ are differentiated by the position of their lowest energy absorption bands.27,28,34,35 At pH 7, a small amount of tautomeric form of HQ has also been found to coexist with the normal form manifested in the appearance of its absorption band.27,28 For 6-hydroxyquinoline (6HQ), the neutral form emits around 380 nm, the cationic form at ∼450 nm, the anionic form at ∼490 nm, and the tautomer (T) at ∼585 nm.27,28,34,35 Similar band positions are found for 7-hydroxyquinoline (7HQ). In the current study, we focused our attention on the ESPT process, from the bifunctional 6HQ and 7HQ molecules, shown in Scheme 1, to water molecules in methanol-doped ice over a wide range of temperatures. We explored the possibility of a proton transfer process via a water bridge that can be formed in water and ice between the acidic hydroxyl and the basic imine nitrogen. Such a bridge upon excitation could, in principle, enhance the production of the zwitterion. The experimental results for both 6HQ and 7HQ show that the ESPT process and the production of the zwitterion that are well-studied in water are also efficient in methanol-doped ice. The temperature dependence of both processes clearly indicates a non-Arrhenius behavior. At temperatures below 173 K, the proton transfer rate is slower than the radiative and nonradiative processes, and

10.1021/jp908051t CCC: $40.75  2009 American Chemical Society Published on Web 10/20/2009

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Figure 1. Time-resolved emission of 6HQ at several temperatures in the range of 80-265 K in H2O and D2O samples containing 0.1% mole ratio of methanol.

consequently, it could not be clearly observed. We found that a small concentration of acetic acid increases the ESPT rate in ice at temperatures below 235 K, and it also enables the ESPT down to the lowest measured temperature, 80 K. Experimental Section We used the time-correlated single-photon counting (TCSPC) technique to measure the time-resolved emission of 6HQ and 7HQ. For sample excitations, we used a cavity dumped Ti:sapphire femtosecond laser, Mira, Coherent, which provides short, 80 fs pulses. The laser’s third harmonic (THG), operating over the spectral range of 260-290 nm, was used to excite both photoacid ice samples. The cavity dumper operated with a relatively low repetition rate of 500 kHz. The TCSPC detection system is based on a Hamamatsu 3809U photomultiplier and Edinburgh Instruments TCC 900 computer module for TCSPC. The overall instrumental response was about 35 ps (fwhm). The excitation pulse energy was reduced to about 10 pJ by neutral density filters. 6HQ and 7HQ (+95%) were purchased from Sigma and Meryer Chemical Company (China), respectively. For transient measurements, the sample concentrations were between 2 × 10-4 and 2 × 10-5 M. Deionized water had a resistance of >10 MΩ. Methanol of analytical grade was purchased from Fluka. All chemicals were used without further purification. The temperature of the irradiated sample was controlled by placing the sample in a liquid N2 cryostat with a thermal stability of approximately (1.5 K. Ice samples were prepared by first placing the cryogenic sample cell for about 20 min at a supercooled liquid temperature of about 260 K. The second step involved a relatively rapid cooling (5 min) to a temperature of about 250 K. Subsequently, the sample froze within a few minutes. To ensure ice equilibra-

tion prior to the time-resolved measurements, the sample temperature was kept for another 10 min at about 250 K. Results The absorption and emission spectra of 6HQ in pH-neutral solution is shown in Figure s1 in the Supporting Information. The main emission bands are that of the neutral species (∼380 nm), the cationic species (∼450 nm), and the zwitterion (580 nm). Figure 1 shows the time-resolved emission of 6HQ in H2O and D2O samples at several temperatures in the range of 80-265 K using the TCSPC technique. The aqueous sample was doped with 0.1% mole ratio of methanol to prevent the aggregation of 6HQ at the grain boundaries of the microcrystals of the polycrystalline ice. Ice is known to be a poor solvent, and the methanol doping prevents the photoacid molecules from being excluded from the ice bulk. The time-resolved signal was measured at 375 nm, which is close to the peak of the NROH* band (see Figure s1 in the Supporting Information). The signal decays at high temperatures in a bimodal pattern. The relative amplitude of the short-time component of the decay is much larger than the long-time component of the decay. The decay rate of the short-component at room temperature is about 3 × 1011 s-1, and it strongly depends on the temperature. At about 175 K, the decay rate is on the order of 2 × 108 s-1, which is faster than the radiative rate. We attribute the short-time component’s decay rate at the high temperature range of T g 173 K to the photoprotolytic process: NROH* h NRO-* + H+. The short-time component’s decay rate is susceptible to the isotope in the reaction, having a kinetic isotope effect (KIE) of 3 at high temperatures above 235 K in both liquid and ice.36 At temperatures below 173 K, the KIE is relatively weaker. A KIE value of 3 is a common value for protolytic processes involving hydroxyaryl photoacids

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Figure 2. Time-resolved emission of the NH+RO-* form of 6HQ measured at 550 nm in methanol-doped H2O sample.

such as 2-naphthol and its 2-naphtholsulfonate derivatives and other similar photoacids. The smaller value of the KIE at low temperatures may be the result of a low proton/deuteron transfer rate in comparison to the overall excited-state decay rate. When kPT < krad + knr, then the decay rate is less dependent on kPT, since k = krad + knr + kPT. The KIE on kPT in such a case affects the overall decay rate, k (the experimentally measured fluorescence decay rate) to a much lesser extent. Figure 2 shows the time-resolved emission of the H+NRO-* form of 6HQ in a methanol-doped H2O sample measured at 550 nm. The signals at T > 173 K show a distinct rise-time followed by a nearly exponential decay, and they also show that the lower the temperature, the longer the rise-time of the signal and its decay. At temperatures below 173 K, the signal measured at 550 nm does not show a long rise-time but rather a fast rise limited by the IRF of the system. The decay time at T < 173 K is very long and independent of the temperature. In general, the time-integrated steady-state emission spectrum shows at T < 173 K a large shift from a band with a peak at 570 nm at T > 173 K to a broad band with a peak at 530 nm. The temperature dependence of the decay time at T > 173 K is explained by a nonradiative process that probably does not involve a reaction with excess protons or hydroxyls in a simple form. The Effect of Acetate Ion and Acetic Acid on Proton Transfer. The role of the acetate ion in enhancing excited-state proton transfer reaction is extensively studied and is welldocumented in recent years. In most of the recent studies by Nibbering, Pines, and co-workers, femtosecond UV pump-IR probe and time-resolved spectroscopies were used to monitor the fast proton transfer from excited 8-hydroxy-1,3,6-pyrenetrisulfonate (HPTS) directly to an acetate ion and also to acetate bridged by water molecule complexes.10,37,38 In water, it was found that when an acetate ion is directly hydrogen-bonded to the hydroxyl group of HPTS, the ESPT process occurs within the time resolution of the optical system, that is, ∼100 fs. For bridged water-acetate complexes, where one water molecule or more acts as a bridge between the hydroxyl group of HPTS and the acetate ion, the rate decreases by more than 1 order of magnitude with respect to the direct transfer. In the current study, we used a similar approach to enhance the proton transfer rate of 6HQ and 7HQ in ice, especially at temperatures below 173 K. In several previous studies,39,40 we found that ESPT to methanol-doped ice for

Presiado et al. several photoacids is not effective below 173 K. In the current hydroxyquinoline study of 6HQ and 7HQ, we were unable to clearly detect an efficient ESPT process in a sample doped with 0.1% mole ratio of methanol below 173 K. Similar behavior is also observed for hydroxyaromatic acids such as 2-naphtholsulfonate derivatives. As we will show below, an addition of a small concentration of acetic acid increases the ESPT rate by at least 10-fold at temperatures below 173 K. Figure 3 shows the time-resolved emission of the NROH* form of 6HQ in water and in ice of three samples with methanol doping of 0.1% mole ratio excited at 285 nm and monitored at 375 nm. One of the 6HQ samples contains 10 mM acetic acid, the second contains 5 mM sodium acetate, and to the third no acid or base was introduced. In ice at 247 K, the 10 mM acetic acid sample has a long-time fluorescence tail with a large amplitude. At temperatures below 173 K, the sample containing 10 mM acetic acid decays faster than the water and the sodium acetate samples. We attribute the faster decay of the acidic solution to an efficient proton transfer from the hydroxyl group to either acetic acid or to an H2O molecule via an acetic acid molecule or an acetic acid dimer. At present, we do not understand why acetate ions are not active in the proton transfer process as opposed to the acetic acid molecules, whose presence enables the proton transfer process. The results of a more systematic study are presented in the next figures. Figure 4 shows the time-resolved emission of the NROH* of 6HQ in three H2O liquid and ice samples doped with 0.1% mole ratio of methanol. In each panel of the figure, we compare a H2O sample with two samples containing a mixture of sodium acetate (NaAc) and acetic acid (AcH). These samples form a buffer solution in the liquid state. The Henderson-Hasslebalch equation provides the relation between the pH of the solution and the ratio of the acid concentration to the concentration of its conjugate base:

pH ) pKa + log

[base] [acid]

(1)

The pKa of acetic acid in water is ∼4.75. The pH of the two samples shown in the figure is 4.75 and 5.35 for the 2 mM AcH + 2 mM NaAc and the 1 mM AcH + 4 mM NaAc samples, respectively. Bardez31 found that the ground state pKa for the following process is about 5.1 for 7HQ: Ka

H3O+ + NROH {\} NH+ROH + H2O

(2)

At pH values lower than the pKa, the ground state absorption shifts to the red and the emission intensity at 375 mn, when excited at 325 nm, strongly decreases. Therefore, increasing the proton concentration in the sample sharply decreases the NORH concentration, and as of consequence it also weakens the NROH* TCSPC signal intensity measured at 375 nm. As seen in the figure, for the buffered samples, the decay rate of the NROH* signal measured at 375 nm is much faster than the decay rate of the acid-free solution. In ice, the signal of the 2 mM AcH + 2 mM NaAc sample decays faster than that of the 1 mM AcH + 4 mM NaAc sample. Both samples decay at T < 173 K, much faster than the acidfree sample. The proton concentration in both samples is very low, that is, on the order of 10-5 M. In such low concentrations, the reaction between the proton and 6HQ in bulk ice should be rather slow in comparison with the radiative decay time. In

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Figure 3. Time-resolved emission of the NROH* form of 6HQ in water and in ice of three samples with methanol doping of 0.1% mole ratio excited at 285 nm and monitored at 375 nm.

Figure 4. Time-resolved emission of the NROH* of 6HQ in ice samples doped with 0.1% mole ratio of methanol in acetate-free and acetic acid/acetate buffered H2O samples.

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Figure 5. Time-resolved emission of the NROH* form of 7HQ measured at 375 nm in H2O and D2O ice samples doped with 0.2% mole ratio of methanol.

previous studies, we found that the proton diffusion constant, DH+, is 10 times larger in ice than in water.14–16 However, this high value of DH+ holds only for T > 240 K. The figure shows that the decay rate of the buffered sample at T < 173 K is faster than that in water, and that the 2 mM AcH sample decays even faster than the 1 mM AcH + 4 mM NaAc sample. We expected that the decay rate of the signal of the 1 mM AcH + 4 mM NaAc sample would be the fastest, since the Ac- concentration, which is the potential proton acceptor, is the highest in this sample. Time-Resolved Emission of 7HQ. In addition to 6HQ, we also experimented in the current ice study on 7HQ in methanoldoped ice. We expected that the position of the hydroxyl group would have a large impact on the photoprotolytic cycle. More specifically, water molecules can bridge between the acidic group (the hydroxyl) and the basic group (the imine nitrogen). Bardez31 suggested that two water molecules may successfully bridge the acidic and basic groups in the case of 7HQ.31 We therefore expected a large difference in the ESPT rate in ice between 6HQ and 7HQ. As we will show in the next figures, 6HQ and 7HQ undergo similar photoprotolytic processes. Based on the similarity of the experimental data, and assuming that such a bridge does not exist for 6HQ, it seems that in the case of 7HQ there is no evidence to support the existence of a water bridge between the hydroxyl group and the imine nitrogen. Figure 5 shows the time-resolved emission of the NROH* form of 7HQ in both H2O and D2O ice samples doped with 0.2% mole ratio of methanol-d. The solubility of 7HQ in water is smaller than that of 6HQ. Consequently, we had to double the methanol doping level. As seen in the figure, the proton transfer rate at T > 173 K is faster than the deuteron transfer rate. The KIE at T g 235 K is ∼3. Samples below 173 K show that the initial fast decay rate is independent of the isotope. All signals at all temperatures consist of short and long components. We explain the lack of a KIE on the decay rate of the NROH*

signal below 173 K by that the fluorescence decay rate is not strongly influenced by the photoprotolytic process. The overall fluorescence decay rate constant depends much more on other nonradiative processes as described in more detail in the case of 6HQ. We conclude that below 173 K the proton transfer rate is slow and therefore cannot be determined by the fluorescence decay rate of the NROH* signal. Figure 6a shows the time-resolved emission of the NROH* form of 7HQ in three ice samples at temperatures equal or higher than 222 K. Two samples contained sodium acetate and acetic acid in order to moderate the pH level and introduce specific well-studied proton acceptors to the sample, whereas the third sample was an acetate-free H2O sample for comparison. As seen in the figure, at all measured temperatures (T g 247 K), the initial decay rate, which signifies the proton transfer rate, is only slightly affected by the addition of 2 mM AcH + 2 mM NaAc or 1 mM AcH + 4 mM NaAc. The long-time nonexponential fluorescence tail’s amplitude and the average decay time strongly depend on the acetic acid/acetate concentration in the sample. Both buffered solutions have much faster decay rates and smaller amplitudes of the long-time fluorescence tail than those of the acetate-free sample (control experiment). The 2 mM AcH + 2 mM NaAc (pH ) 4.75) sample has a larger effect on the longtime fluorescence tail than the 1 mM AcH + 4 mM NaAc sample. Figure 6b shows the time-resolved emission of the NROH* form of 7HQ in the three samples shown in figure 6a at T e 197 K. At this temperature range (197-80 K), we found that both the short and long times of the signal are strongly dependent on the presence of the buffer in the sample. The shortand long-time components of the buffered samples decay faster in comparison to the acetate/acetic-acid-free sample. The fast short-time decay of both buffered samples is almost the same, and thus, some specific bridge between the hydroxyl group of 7HQ and the imine nitrogen may be operative. This bridge may

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Figure 6. Time-resolved emission of the NROH* form of 7HQ in three ice samples at (a) T g 222 K and (b) T e 197 K.

contain one or two acetate or acetic acid molecules. The longtime fluorescence tail’s amplitude and decay time show, as in the high temperature region (T g 197 K), that for the 2 mM AcH + 2 mM NaAc sample the amplitude is the smallest and the decay time is the shortest. This behavior is the opposite of the one observed for 6HQ, where the long-time tail’s effective lifetime is much longer in the buffered sample than in the acetic acid/acetate-free samples (see Figure 4).

The difference between the long-time component’s decay of 6HQ and 7HQ may indicate that the recombination of H+NRO-* with the proton to reform the NROH* form of 6HQ is a reversible reaction, whereas for 7HQ it is partially irreversible (a large kq); see Scheme 2. Such a large difference in the photoprotolytic cycle of two similar photoacids is also observed for 1-naphthol and its

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SCHEME 2

SCHEME 3

derivatives (the irreversible case) and 2-naphthol and its derivatives (the reversible case).16,39 Discussion The main findings of the current study on the photoprotolytic properties of 6HQ and 7HQ in methanol-doped water and ice are as follows: (1) We found that the ESPT process is efficient for both compounds in methanol-doped ice in the temperature range of 270-173 K. We observed the reaction rate by monitoring the decay rate of the time-resolved emission of the NROH* band, whose maximum is at ∼375 nm. The decay rate of this band approximately matches the rise-time of the time-resolved emission of the H+NRO-* band, whose peak is at 585 nm. For 6HQ, the amplitude of the clearly marked rise-time component is large, that is, g60% of the total signal, whereas for 7HQ the amplitude is smaller, that is, on the order of e40% of the total signal. (2) We found a relatively large KIE on the ESPT rate constant for both 6HQ and 7HQ of ∼3 in ice at T g 222 K. (3) Both of the Arrhenius plots of 6HQ and 7HQ in the temperature range of 80-270 K, which show the logarithm of the ESPT rate constant, kPT, versus 1/T, have a concave shape. The slope of ln kPT decreases as the temperature decreases (1/T increases). This type of non-Arrhenius behavior of kPT was observed in many of our previous studies on photoacids in methanol-doped ice.14–16,41–45 At temperatures below 173 K, the proton transfer rate is slower than the (radiative and nonradiative) excited-state decay and the efficiency of the H+NRO-* zwitterion formation process is low. Consequently, both processes cannot be followed by the TCSPC technique. We will discuss this general characteristic behavior of the ESPT process in methanol-doped ice in further detail. (4) We found that adding a small concentration of a few millimolars of acetic acid or a mixture of acetic acid and sodium acetate to moderate the sample’s pH enhances the proton transfer rate in methanol-doped ice for both 6HQ and 7HQ as well as for the deuterated samples. (5) In ice containing acetic acid, the proton transfer rate below 235 K is fast and nearly temperature-independent. This finding is also true for deuterated samples. The Temperature Dependence of 6HQ and 7HQ in Methanol-Doped Ice. The time-resolved emission signal of the NROH* band of both 6HQ and 7HQ in methanol-doped ice measured at 375 nm is nonexponential, and at T g 235 K it is bimodal. We treated the signal as if it arises from a model, where the proton transfer step is followed by a diffusion-assisted reversible geminate recombination process. The excited molecule can undergo an additional step, protonation of the imine nitrogen. This last step is not included in the model used to fit the decay of the neutral form measured at 375 nm. The overall process is shown diagrammatically in Scheme 3. The first step in the photoprotolytic cycle is a proton transfer to a nearby water molecule. Proton transfer to a nearby methanol molecule is negligible, since ESPT to a methanol molecule in pure liquid methanol at 296 K takes ∼200 ps to complete, that is, a rate constant of 5 × 109 s-1, a rate which is about 10 times

Presiado et al. slower than that of pure water. When the methanol concentration is increased 10-fold to 1.25% mole ratio, the proton transfer rate in methanol-doped ice is halved. These results clearly indicate that methanol is not as directly involved as the primary proton acceptor. The role of methanol is to preferentially solvate the hydroxyquinoline molecules in the bulk ice. When the methanol-free water sample of 6HQ or 7HQ freezes, its steadystate fluorescence intensity decreases by a factor of ∼100. This observation is explained by the tendency of ice to expel all impurities from its bulk, and consequently, the 6HQ and 7HQ molecules aggregate at the grain boundaries of the microcrystals of the polycrystalline material. The hydronium-NRO-* ion pair (see Schemes 2 and 3) dissociates and forms a diffusing proton that migrates within the ice structure. A large defect zone, which includes several methanol molecules as well as water molecules, is expected to surround the NRO-*. The proton may recombine with NRO-* to form the zwitterion or the protonated hydroxyl group:

H+ + NRO-* f H+NRO-* H+ + NRO-* f NROH*

(3)

Protonation of the imine nitrogen may occur by water cleavage prior to the proton transfer from the hydroxyl group as suggested by Bardez.31 The diffusion-assisted geminate recombination model of Pines et al.46 predicts that the time-resolved emission signal of the protonated hydroxyl form of a photoacid would consist of a fast nearly exponential decaying component, followed by a nonexponential long-time fluorescence tail. This pattern is observed for the time-resolved emission signal of the 375 nm band of both 6HQ and 7HQ. A quantitative analysis of the signal using the SSDP program of Krissinel and Agmon yields the proton transfer rate constant, kPT.47 For more details, please refer to the Supporting Information and previous studies.7,18,46 Table s1 in the Supporting Information provides the fitting parameters of the time-resolved emission of 6HQ in D2O at several temperatures in the range of 80-295 K. Figure 7 and Supporting Information Figure s7 show Arrhenius plots of ln kPT versus 1/T for 6HQ and 7HQ, respectively. Both figures show that the slope of ln kPT decreases as the temperature decreases. We fit the plot, shown in figure 7, to two models, which we have already used in the past for similar experiments on green fluorescent protein in ice.48 The first model is a modified one-dimensional proton tunneling, for which the intermolecular distance between heavy atoms is the main coordinate. In our case, this is the distance between the two oxygen atoms: the hydroxyl oxygen and the hydrogen-bonded water molecule oxygen. The distance is modulated by the intermolecular vibration. When the two oxygens reach a minimum distance from each other, where the amplitude of the oscillation is of ∼0.05 Å, the tunneling rate is more than 10 times faster than its rate at the other extreme, that is, at the maximum distance between the oxygens. The detailed model and the parameters that influence the rate are given elsewhere.49,50 Figure 7b shows the fit to the experimental data using the above-mentioned model. The fitting parameters are the attempt frequency, ν ) 1013 s-1, the intermolecular vibration frequency, ω ) 250 cm-1, the tunneling integral at the equilibrium position J ) 17.4, and its derivative with respect to the distance, J′ ) 52. The fit (solid line in the plot) is rather good in the temperature range of 270-173 K. Another plausible mechanism that may explain the nonArrhenius temperature dependence of the decay rate of the

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k2(T) ) k02 exp(-Ea2 /RT)

(4b)

where k10 . k20 and Ea1 . Ea2. At any given temperature, the total rate constant is given by

ktot ) k1 + k2

Figure 7. Arrhenius plot of ln kPT versus 1/T for 6HQ (a) in H2O and D2O samples and (b) in acid-free H2O sample, buffered ice sample containing 2 mM AcH and 2 mM NaAc and fit to the experimental data using the intermolecular one-dimensional vibration-assisted proton tunneling model. (c) Two coordinates model.

protonated NROH* form in ice in the temperature range of 79-270 K may include two processes that control the overall decay rate of the neutral form, rather than one. In this case, we can assign a rate constant for each coordinate (k1 and k2). The temperature dependence of each rate constant follows an Arrhenius law. Let us assume that k1 and k2 are substantially different from each other in both their activation energy and their pre-exponential factor

k1(T) ) k01 exp(-Ea1 /RT)

(4a)

(5)

In such a case (by choosing the precise parameters), the main channel at high temperatures is that of the rate constant k1, whereas at a low enough temperature we find a switchover, where k2 g k1. Figure 7c also shows the calculated curves of ktot, k1, and k2 (solid, broken, and dashed curves, respectively) and the values of ktot(T) as obtained from the best fit of the time-resolved emission for H2O and D2O samples. The fit is rather satisfactory at both high and low temperatures. We used for H2O k10 ) 3.0 × 1013 s-1, k20 ) 0.3 × 109 s-1, Ea1 ) 17 kJ/mol, and Ea2 ) 1 kJ/mol. The second “coordinate” of the overall decay rate of the neutral form may arise from a nonexponential nonradiative process of the excited NROH. In previous studies on ice, we found that the nonradiative decay of guanosine derivatives is nonexponential.51 Enhanced Proton Transfer in Ice Containing Acetic Acid. Figures 3, 4, and 6 and Figures s4 and s5 in the Supporting Information show the effect caused by the presence of a few millimolars of acetic acid on the rate of proton transfer from the hydroxyl group in 6HQ and 7HQ samples. As seen in the figures, the proton transfer rate increases for both 6HQ and 7HQ at T e 235 K. As aforementioned, in acetic-acid-free ice, we were unable to determine whether a proton transfer occurs in ice at T < 173 K. The reasons we think that the proton transfer process at T < 173 K is negligible are as follows: a very small KIE on the time-resolved emission of the NROH* form measured at 375 nm, the disappearance of the tautomeric (zwitterionic) emission band at 580 nm in these temperatures, and the significant presence of the cationic H+NROH* form emitting at 450 nm. The experimental result, which may indicate that proton transfer does occur below 173 K in acetic-acid-free ice samples, is the nonexponential decay profile of the NROH* signal (see Figure 1) with a short-time component of ∼2 ns. We propose that a phase transition in ice, from hexagonal ice to cubic ice, is a plausible explanation for these drastic spectroscopic changes at ∼173 K. This change in the ice structure drastically decreases the ESPT rate and thus stops the proton transfer process. In the literature, there are large numbers of studies reporting on the properties and structures of a metastable low-temperature phase designated “ice Ic” in which the oxygens are arranged in a cubic diamond structure rather than on the hexagonal lattice of the regular Ih.52 From a large number of studies reported in the referenced book on ice physics,52 it is clear that ice Ic is always metastable, and the temperature, at which it transforms to ice Ih, is determined by the process of molecular rearrangement. As published in the literature, the temperature of the phase transition between Ic and Ih ranges from 130 up to 200 K. Recently, it was shown that ice Ic forms when aqueous NH4SO4 droplets freeze homogeneously below 183 ( 1 K.53 An important mechanism for ice cloud formation in the Earth’s atmosphere is homogeneous nucleation of ice in aqueous droplets, and this process is generally assumed to produce hexagonal ice.54,55 However, there are some reports that the metastable crystalline phase of ice, cubic ice, may form in the Earth’s atmosphere.56–58 As seen in Figures s4b and s5 in the Supporting Information, the faster decay rate of the NROH* band in the presence of

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acetic acid is also accompanied by a large KIE for both 6HQ and 7HQ. The large KIE indicates that the acceleration of the decay rate in the presence of acetic acid is strongly related to the proton transfer from the hydroxyl group. The fast decay of the NROH* form in cubic ice at T < 173 K is not followed by the appearance of the final zwitterionic H+NRO-* form, whose peak is at 580 nm, and observed at T > 173 K. Instead, we observed at T < 173 K a high intensity fluorescence band at 450 nm. This band is assigned to the cationioc form, H+NROH*. The time-resolved emission signal measured at 450 nm is a superposition of the relatively short-lived NROH* signal and the long-lived H+NROH* emission signal, excited directly from its ground-state by the ∼200 fs excitation pulse at 285 nm. We were unable to detect a distinct rise-time of the 450 nm signal, and therefore, we suggest that a large portion of the 450 nm band intensity arises from a direct excitation of the groundstate population of this species. Figure 7b shows ln kPT of 6HQ versus 1/T for the buffered ice sample containing 2 mM AcH and 2 mM NaAc so that its pH was maintained at ∼4.75. The plot shows that the proton transfer rate constant at 235 K decreases by only about a factor of 3 as compared to its value at 268 K. In the range of 235-79 K, kPT is temperature-independent within the experimental error. We propose that a formation of a complex of the hydroxyl group of 6HQ and 7HQ and one or more acetic acid molecules may explain the sharp rise of the proton transfer rate in the presence of acetic acid. In the papers of Pines, Nibbering, and co-workers4,5,10,37,38 and the more recent paper of Bakker and co-workers,9 it was found that acetate ions and chloroacetate derivatives are excellent proton acceptors from HPTS in aqueous solutions. They found a wide range of water bridge lengths that assist the proton transfer from the hydroxyl group to the acetate, which could be located more than two water molecules away from the hydroxyl group of HPTS. In the current study, we found an accelerated proton transfer rate for both 6HQ and 7HQ in ice at T < 235 K in the presence of a few millimolars of acetic acid in the sample. This is a remarkable observation, since in the acetate-free sample the ESPT process has a strong temperature dependence, whereas below 173 K this phenomenon is not observed since kPT < kr + knr. Mehata59 studied the proton translocation along a hydrogen-bonded “molecular wire” in a 6HQ acetic acid complex. In his study, he presented steadystate absorption and emission spectra as well as time-resolved emission spectra of 6HQ in a benzene sample containing also a small concentration of acetic acid. They concluded that when this complex is excited, charge redistribution initiates the acid-base reaction in the electronic excited state via coupled electron-proton transfer and produces a photoproduct, that is, keto-tautomer, within 200 ps. The proton-transfer reaction takes place in a unidirectional fashion from the O-H to quinolinic -N- site of cis-6HQ along the O-H · · · AcH · · · AcH · · · N wire, and the resulting tautomer has a decay time of ∼1.4 ns. Summary and Conclusions The photoprotolytic process of the two bifunctional 6HQ and 7HQ molecules in water were studied by time-resolved and steady-state emission and UV-vis spectroscopies over a wide temperature range of 80-296 K. We found that in methanoldoped ice at T > 173 K the ESPT rate from the hydroxyl group and the formation rate of the H+NRO-* zwitterion’s emission band at any given temperature are similar for both 6HQ and 7HQ. Below 173 K, we were unable to observe the formation of the zwitterion with our spectroscopic techniques for both compounds. Several recent studies53,60,61 on water droplets

Presiado et al. freezing in the presence of inorganic and organic compounds show that at ∼200 K the ice formed is not in its hexagonal phase Ih, but rather in its metastable cubic ice Ic. We therefore propose that below 173 K the methanol-doped ice Ih samples we prepared transform into the cubic ice Ic. In this phase, the photoprotic molecule cannot transfer a proton within its excitedstate lifetime. In this study, we examined the possible formation of the zwitterion by a concerted mechanism, in which the released hydroxyl proton reaches the nitrogen via a water bridge of two or more molecules or a complex of water molecules in general. The experimental results of both HQs show that decay rates of their respective neutral forms’ emission band at 375 nm and the rise-time of the time-resolved emission of their respective zwitterions at 580 nm have similar absolute rates and temperature dependences. Based on the experimental results, we conclude that the concerted formation of the zwitterion in methanol-doped ice is not the main process when the neutral form is excited. We added a small concentration of acetic acid to the methanol-doped ice samples of 6HQ and 7HQ and also sodium acetate in order to moderate the drop in the pH level. In the presence of a small concentration of acetic acid, we found that there is a relatively fast ESPT rate, kPT = 5 × 109 s-1 at T < 173 K. The rate constant, kPT, is almost temperature-independent in the range of 222-80 K. As aforementioned, in acetic-acidfree samples, we were unable to observe the ESPT below 173 K. Our explanation to this observation is that there is a direct ESPT from the hydroxyl group to the acetic acid. Water may be the proton’s final destination, since the neutral forms of both 6HQ and 7HQ in samples containing acetic acid have a nonexponential long-time fluorescence tail at T < 173 K, and the diffusion-assisted geminate recombination model predicts such a nonexponential fluorescence tail. In many previous studies, it was found that the presence of acetic acid in aqueous solutions enhances the proton transfer rate of a photoacid. Pines, Nibbering, and co-workers10,37,38 and, more recently, Bakker and co-workers9 extensively studied the proton transfer rate from HPTS to acetate and chloroacetate in aqueous solutions containing high and moderate concentrations of these substances at room temperature by using fs UV pump-IR probe spectroscopy. They concluded that a wide range of water bridge complexes between the proton donor, the hydroxyl group, and the acetate enable the efficient ESPT process. Our experiments on ice doped with acetic acid show that proton transfer indeed occurs at T < 173 K, though we could not detect it in acetic-acid-free samples. Acknowledgment. This work was supported by grants from the Israel Science Foundation and from the James-Franck German-Israeli Program in Laser-Matter Interaction. Supporting Information Available: Steady-state emission and absorption spectra of 6HQ; time-resolved emission spectra of 6HQ and NH+RO-* and NROH* forms of 6HQ and 7HQ; Arrhenius plot of ln kPT versus 1/T for 7HQ; SSDP model fitting parameters. This material is available free of charge via the Internet at http://pubs.acs.org. References and Notes (1) Ireland, J. E.; Wyatt, P. A. AdV. Phys. Org. Chem. 1976, 12, 131. (2) Gutman, M.; Nachliel, E. Biochem. Biophys. Acta 1990, 391, 1015. (3) Tolbert, L. M.; Solntsev, K. M. Acc. Chem. Res. 2002, 35, 19. (4) Rini, M.; Magnes, B. Z.; Pines, E.; Nibbering, E.T. J. Science 2003, 301, 349.

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