LETTER pubs.acs.org/JPCL
pH-Jump Overshooting Mateusz L. Donten and Peter Hamm* University of Zurich, Institute of Physical Chemistry, Winterthurerstrasse 190, CH-8057 Zurich, Switzerland ABSTRACT: Acidbase systems are commonly expected to equilibrate on a time scale much faster than any other chemical reaction, so their composition can be deduced from the corresponding pKa or pKb values. In a pH-jump experiment done on a multi-acid/base pair system, it was found that it takes tens of microseconds before an equilibrium is established. Within that time, the system is kinetically driven, reaching surprising states very different from its final equilibrium; for example, carboxylate groups were protonated in the presence of hydroxyl ions. SECTION: Kinetics, Spectroscopy
A
cidbase reactions are among the very fundamental categories of chemical transformations and are typically thought of in terms of pH and the following chemical equilibrium H3 Oþ þ OH T 2H2 O Neutralization of OH, H3Oþ, or both is diffusion limited, as such is “fast” with a bimolecular rate of ∼1011 M1s1.14 Thanks to the Grotthuss mechanism,5 pH equilibration is commonly assumed to be faster than other diffusion controlled processes. The proton transfer step is studied in details on the ultrafast time scale by Nibbering, Pines, and coworkers6,7 as well as Bakker and coworkers.810 Despite the event of a single proton transfer happening on a picosecond time scale, the bimolecular rate of H3Oþ/OH recombination translates into a typical time scale of 100 μs needed to establish an equilibrium at neutral pH 7. Similar to temperature jumps,11,12 pH jump experiments are a modern experimental technique for time-resolved investigation of many proton driven chemical phenomena1317 and their follow up processes such as protonation-induced protein folding.1821 Using this example, folding of small to midsized proteins involves steps on time scales that extend from nanoseconds to 100 μs, that is, in the same range as H3Oþ/OH equilibration, and so the neutralization reaction can be confused with a process intrinsic to the studied molecule. In a pH jump experiment, a fast photochemical reaction triggered by a short laser pulse is used to release rapidly protons in either a reversible (i.e., with a photoacid) or irreversible (i.e., as a caged proton) manner. The further fate of the protons can be predicted by straightforward kinetic models based on equilibrium constants and kinetic rates of the involved acidbase reactions. Gutman and coworkers provided the first experimental evidence of pronounced nonequilibrium states in such pH-jump experiments using photoacids as proton source.2226 Photoacids r 2011 American Chemical Society
change their pKa in the electronically excited state.27 The pH jump is reversible in this case as the photoacid relaxes back to the ground state on typically a pico- to nanosecond time scale. Nevertheless, the proton can be transferred to another base within the lifetime of the excited state, from which it returns to the ground-state photoacid, which is a stronger base, on a much slower time scale of hundreds of microseconds. Such a relaxation process cannot be explained solely in terms of equilibrium constants but needs a kinetic description. In the present Letter, we systematically study a system with competing acid/base pairs in dependence of the starting pH using ortho-nitrobenzaldehyde (oNBA) as phototriggerable proton source and acetate as a prototype base. Using an irreversible proton source instead of a photoacid removes the competing process of reprotonation of the ground state of the latter and renders the lifetime of the pH-jump to last forever. This allows for a much more direct observation of a sequence of proton transfer processes between competing bases (acetate and OD in this case) in dependence of their concentrations, which furthermore can be adjusted independently from that of the proton source. oNBA is a well-studied and commonly used caged proton.14,1821 Its reaction kinetics is highly solventdependent28,29 and recently was examined in water,30 where nitrosobenzoic acid is formed already 7 ps after irradiation with UV (266 nm) light in its protonated form. The proton is then released after a few tens of nanoseconds. These experiments set a firm understanding of how the proton is delivered to the solution. Here we study its further fate in a multibase system by timeresolved UVpumpIRprobe spectroscopy. Received: May 5, 2011 Accepted: June 7, 2011 Published: June 07, 2011 1607
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The Journal of Physical Chemistry Letters The experiment was performed on a saturated aqueous solution (D2O) of oNBA (∼8 mM)18 and with 10 mM sodium acetate. The latter was 13C-isotope-labeled in the carboxyl group to avoid spectral overlap with the carboxyl group of nitrosobenzoic acid formed in the oNBA photoreaction (Figure 1, upper plots). The solution was prepared with acetic acid and titrated
Figure 1. Experimental data. Upper panel: FT-IR of the isotope-labeled acetic acid and acetate anion and pump probe spectra recorded for these groups following a pD jump initiated in a pD 11.4 solution. Lower panel: Kinetic traces extracted from the pumpprobe spectra showing the carboxyl group band evolution between 1640 and 1670 cm1 for samples of different initial pD. Signals were normalized to the intensity of the oNBA aldehyde bleach to compensate for the minor fluctuations in the pump power and jet thickness that affected the number of released protons.
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with NaOH to the desired initial pD.31 The pD in D2O was measured with a standard glass electrode and corrected for the H/D isotope effect with the known offset value of þ0.4 pD unit.32 Solutions with the following pD values were prepared: 6.4, 8.4, 10.4, 11.4, 11.9, 12.4, 12.9, and 13.4. We carefully verified by titration experiments that no contamination of any other acid/ base pair was present in the solution. For recording the timeresolved spectra, the pumpprobe setup described in refs 30, 33, and 34 was used. In brief, two femtosecond laser systems were electronically synchronized to generate a pump-pulse at 266 nm and a tunable broad band IR probe pulse in the desired frequency range of the carboxyl and carboxylate bands, with a time delay that was continuously scanned from 10 ps to 10 μs with ∼10 ps of time resolution.34 On the basis of the concentrations and equilibrium constants (i.e., pKa values) of the various acid/base pairs present in the solution, the following final outcome of the pH-jump is anticipated. The pump pulses generated ∼3 mM of nitrosobenzoic acid (estimated from the number of absorbed photons within the illuminated volume and the 50% quantum yield of the photo reaction30). Nitrosobenzoic acid is a much stronger acid (pKa = 2.9, estimated from the dissociation rate, see Table 1 below) than acetic acid (pKa = 4.8, which is corrected to pKa = 5.3 for the deuterated compound35), so we expect most of them to deprotonate and the protons to be transferred to the acetate ions. However, at basic pD (exceeding pD 11 corresponding to [OD] > 1 mM), the amount of OD ions is sufficient to neutralize a big part or all released protons. OD is widely known to be the strongest base with the pKa of water, 15.7,4 and even higher for deuterated water, 16.5.4 Consequently, the pD jump will not be nearly as large as it could be based on the total concentration of released protons. The sequence of events, that is, which of the competing bases is protonated first, will depend on the kinetics of the process. Figure 1 (upper panel) shows the IR spectra of the carboxyl and carboxylate groups. The FTIR spectra (upper plots) were used to assign the marker bands of the isotopically labeled acetic acid and its base. Also shown is their time-evolution upon oNBA photoproduct dissociation. Protonation of acetate is seen as the positive band appearing around 1655 cm1 (CH313COOD) and a corresponding bleach at 1520 cm1 (CH313COO). The beach signals at 1540 cm1 and 1700 cm1 originate from the NO2 and COH groups of the reactant, oNBA.30 The spectra were acquired with an initial pD ∼11.4. It can be noticed that under these basic conditions the protonation of acetic acid is not
Table 1. Reactions Included in the Kinetic Mode of the AcidBase System with Their Equilibrium and Kinetic Constantsa kþ
reaction (1) NBD þ D2 O a NB
þ D3 O
þ
(2)ACD þ D2 O a AC þ D3 O þ
pK a ¼ 2:9
1 10 s
8 10 M1 s1 b
pK a ¼ 5:3 35
2.8 105 s1 c
8 1010 M1 s1 37
2.6 106 s1 c
8.4 1010 M1 s1 4
6 102 s1 c
1 1010 M1 s1 e
(3) D2 O þ D2 O a OD þ D3 O þ pK a ¼ 16:5 (4)
AC þ D2 O a ACD þ OD
k
8 1 b
4
pK ¼ 11:2 d
10
a
Note: Water is printed in gray because by convention it was included in the pK values with a constant concentration of 55 M and is not treated explicitly in the kinetic rate constants (convention commonly used for pKa and pKb). Nonstandard symbols: ACD, deuterized acetic acid; AC, acetate ion; NBD, deuterized nitrosobenzoic acid; NB, nitrosobezoate ion. b Forward rate kþ fitted to experimental data, back reaction rate k assumed to be the same as for acetate protonation (reaction 2). The pKa is calculated from kþ and k (eq 3). c Calculated based on the equilibrium constant and the rate constant for the opposite reaction (eq 3). d Calculated based on pKa of ACD and D2O (pKa(4) = pKa(3) pKa(2)). e Free parameter of the kinetic model adjusted to match the experimental. 1608
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Scheme 1. AcidBase Reactions Following the Photoreaction of oNBA
permanent, and the peak and bleach pair disappear again on a longer time scale of a few hundred nanoseconds. To examine this phenomenon closer, we repeated the experiment starting from different initial pD values. From the recorded time-resolved spectra, the dynamics of the carboxyl group marker band was extracted to kinetic traces presented in Figure 1 (lower panel). For near-neutral samples (pD: 6.4, 8.4, and even 10.4), the expected time evolution of the signal is observed. The band intensity grows monotonically, reflecting the kinetics of the acetate protonation. An exponential fit to these traces gave a rise time of ∼10 ns, which matches the proton release from nitrosobenzoic acid,36 as deduced from the carboxyl band overlaying the oNBA aldehyde bleach (positive signal at 1725 cm1 in Figure 1). Given the low concentration of the acetate as compared with that of water, the vast majority of protons will initially be passed onto water molecules, forming D3Oþ ions as a mediator, and only somewhat later to acetate. Experiments done with HPTS (8-hydroxy-1,3,6-trisulfonate-pyrene) as fast photoacid and with higher concentrations of acetate ions gave 8 1010 M1 s1 for its protonation rate.37 In the case of a 10 mM acetate solution, this translates into a diffusion-controlled reaction time of ∼1 ns. Hence, proton release from nitrosobenzoic acid and not the diffusion of the D3Oþ ions is the ratelimiting step, so we observe the same time constant (within experimental error) for protonation of acetate. In the case of the alkaline samples (pD 11.4 and higher), a qualitatively different behavior is observed. The released protons are first located on the acetic base, giving rise to the 1655 cm1 carboxyl band, but somewhat later, this band disappears again (Figure 1). From the pD dependence of the second step, we conclude that it reflects recombination with OD. As the most striking observation of this Letter, we find that acetate is first protonated despite the presence of the much stronger base OD. Scheme 1 summarizes the possible proton transfer pathways, including the direct recombination of acetic acid with hydroxyl ions. This reaction scheme can be verified by building the corresponding kinetic model on the level of diffusion controlled bimolecular reactions, which includes all four acid/base equilibria: d½D3 Oþ þ ð2Þ þ ¼ kð3Þ ½D3 O ½OD k ½D3 O ½AC dt ð2Þ
ð1Þ
þ þ kþ ½ACD kð1Þ ½D3 O ½NB þ kþ ½NBD
Figure 2. Kinetic model calculations. Upper panel: Time evolution of concentration of all acids and bases in the solution and of the ionic product of water in a sample of initial pD 11.4. Lower panel: Time evolution of the acetic acid concentrations following a pD-jump in samples of different initial pD. Abbreviations: ACD, deuterized acetic acid; AC, acetate ion; NBD, deuterized nitrosobenzoic acid; NB, nitrosobezoate ion. Initial conditions and parameters for the simulation: cNBD = 2.5 mM, cAC = 10 mM, pKa values and rate constants as in Table 1.
d½D3 Oþ þ ð2Þ þ ¼ kð3Þ ½D3 O ½OD k ½D3 O ½AC dt ð2Þ ð1Þ þ þ kþ ½ACD kð1Þ ½D3 O ½NB þ kþ ½NBD d½OD þ ð4Þ ¼ kð3Þ ½D3 O ½OD k ½OD ½ACD dt ð4Þ þ kþ ½AC d½NBD ð1Þ þ ¼ kþ ½NBD þ kð1Þ ½D3 O ½NB dt d½NB ð1Þ þ ¼ þ kþ ½NBD kð1Þ ½D3 O ½NB dt d½ACD ð2Þ þ ¼ þ kð2Þ ½D3 O ½AC kþ ½ACD dt ð4Þ kð4Þ ½OD ½ACD þ kþ ½AC d½AC ð2Þ ¼ kð2Þ ½D3 O½AC þ kþ ½ACD dt ð4Þ þ kð4Þ ½OD½ACD kþ ½AC
ð2Þ
Here, kþ(i) and k(i) refer to forward and backward reaction rates, respectively, of reaction i in Table 1. (Water concentration does not appear explicitly in the equations because by convention it is included in the quasi-first-order rate constants.) The model builds on equilibrium constants and kinetic rates known from literature (Table 1). The nitrosobenzoic acid deprotonation rate 1609
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(10 ns)1 was obtained in a direct fit of its carboxyl band intensity at 1725 cm1, and its back-rate was assumed to be the same as that of acetate. (The exact number of that back-rate is not critical for the outcome of the simulation because nitrosobenzoic acid almost completely deprotonates under the conditions of the experiment.) The formation of contact pairs between nitrosobenzoic acid and acetate ions, in which the proton could be passed directly in a first-order kinetic process,37 was neglected in the modeling, given the millimolar concentration of both reactants. Furthermore, proton release from the ketenenitronic acid intermediate (as reported in ref 38) was ignored because the lifetime of that intermediate (7 ps) is by far too short and because the IR spectral signature proves that the final product (nitrosobenzoic acid) is formed in its protonated form.30 The only unknown in the reaction scheme (Scheme 1) was the direct recombination rate of CH3COOD with OD. Its equilibrium constant was calculated from the pKa values for water and acetic acid, and its reaction rate was considered to be the only free fitting parameter of the model. It turns out that the disappearance of acetic acids critically depends on the inclusion of this parameter, and we obtain a semiquantitative agreement with the experimental data for 1 1010 s1 M1 (Figure 2).39 The kinetic model reproduces the transient protonation of the acetate ions in dependence of the starting pD and thus gives an simple explanation of the effect. The bimolecular rate constants for the two neutralization reactions of CH3COO and OD are approximately the same, 8 1010 M1 s1.4,37,17 This is because these rates are limited by the diffusion of a proton through the Grotthuss mechanism and as such are essentially a constant (within much less than one order of magnitude). In simple words, the released protons react with any available base on a “first-foundfirst-protonate basis”, independent from the thermodynamic driving force. The latter can be calculated from the rates pKa ¼ log k log kþ
ð3Þ
but is essentially only determined by the deprotonation rate kþ of the acid and hardly by the relatively universal proton recombination rate k of the corresponding base. At not too high pD (i.e., around pD = 11), equilibration kinetics does not drive the system toward direct recombination between the hydronium ions and the strongest base. However, as a consequence of the kinetic character of the phenomenon, acetic acid deprotonation accelerates with increasing initial solution pD (Figure 1). It starts to compete with acetate protonation at similar time scales for a pD exceeding 12 (i.e., [OD] > 10 mM). It is interesting to note that the reaction rate for direct CH3COOD and OD neutralization is eight times slower than the typical recombination rate of D3Oþ with either OD 1,2 or acetate.37 Because the diffusion constants of hydronium and hydroxyl (which are the more mobile reactants determining the neutralization rate for one and the other case) differ by only a factor 2, the remaining difference must be related to details of the reaction mechanism such as contact distance and reactive volumes, which strongly depend on the charges of the reactants that differ between the two cases. The transient appearance of acetic acid in the more basic samples could be called pH jump overshooting in which the released protons form hydronium ions and protonate the acetate before neutralizing with the stronger but lower concentrated hydroxyl base. Because of its higher concentration, acetate is in a quasi-equilibrium with D3Oþ on a time scale faster than the
D3Oþ/OD equilibrium, and hence the vibrational spectrum of acetate might be considered to be an IR-indicator reporting on the transient pH of the sample. Although the meaning of the pH (or pD) is still rigorously defined as log(H3Oþ) (or log(D3Oþ)) under these nonequilibrium conditions, it largely looses its significance because it no longer allows one to determine the concentrations of the various acids or bases in the solution. The importance of pH comes from its direct translation to pKa equilibria (like in case of buffering systems) and from its correspondence with the equivalent base parameters like pOH (pOD) and pKb. None of these concepts is valid if the system is not in equilibrium. For example, the ionic product of water can significantly exceed its pKw (Figure 2, upper panel). Hence, this experiment illustrates the importance of a kinetic description of acidbase systems. The equilibrium state may be reached as late as hundreds of microseconds after initiation of the reaction. In this time regime, crucial for the investigation of many chemical reactions, discussing the proton transfers just in terms of strong and weak acids and bases fails. Awareness of that fact should find appreciation not only in pH jump experiments but also for reactions with constant proton release. If certain steps in such a reaction scheme occur in the microsecond range, then a stationary state is reached, in which the competing reactions will match their rates rather than equilibrate according to their thermodynamic driving forces. The simulations show that the dynamics of pH equilibration in a multibase system depend on the total concentration of the bases and, more importantly, on their ratio. In a first approximation, the total concentration impacts how fast equilibration is completed and the limit to which the pH jump will be neutralized. The ratio of concentrations will determine whether pH-jump overshooting will occur. With the concentrations in our current experiment, acetate anions are the only effectively available bases under nearneutral or slightly basic conditions (pH e10), and neutralizing them finishes the reaction. The trace amount of OH (or OD) is neutralized on a very long time scale that is expected to reach 100 μs to milliseconds based on the bimolecular rate constant. However, its impact on the signal is too small for detection and is not seen in Figure 1. Nevertheless, if one were to do these experiments at lower concentrations of both acetate and generated protons, then the pH range in which pH-jump overshooting is detectable would decrease accordingly, and the time scale of equilibration would increase by the same factor. pH-jump overshooting may also be taken advantage of by allowing for pH jumps under conditions where the delivered protons would normally neutralize with another base (such as hydroxyl ions). In such cases, the delivered protons may still be used to protonate transiently the investigated molecule. The pH jump can be controlled by choosing the concentrations of competing bases so that the proton is transferred to the stronger (interfering) base only after the investigated reaction took place.
’ AUTHOR INFORMATION Corresponding Author
*E-mail:
[email protected]. Fax: þ41 44 635 68 13.
’ ACKNOWLEDGMENT We thank Alexander Rodenberg for fruitful discussions. The work has been supported by the Swiss National Science Foundation under grant 200021-124463/1. 1610
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’ REFERENCES (1) Atkins, P. W. Physical Chemistry; Oxford University Press: Oxford, U.K., 1989. (2) Natzle, W. C.; Moore, C. B. Recombination of Hþ and OH In Pure Liquid Water. J. Phys. Chem. 1985, 89, 2605–2612. (3) Rice, S. A. Diffusion-Limited Reactions; Bamford, C. H., Tipper, C. F. H., Compton, R. G., Eds.; Comprehensive Chemical Kinetics 25; Elsevier: New York; pp 1617. (4) Eiden, M. Proton Transfer, Acid-Base Catalysis, and Enzymatic Hydrolysis, Part I: Elementary Processes. Angew. Chem., Int. Ed. 1964, 3, 1–19. (5) Agmon, M. The Grotthuss Mechanism. Chem. Phys. Lett. 1995, 244, 456. (6) Mohammed, O. F.; Pines, D.; Dreyer, J.; Pines, E.; Nibbering, E. T. J. Sequential Proton Transfer Through Water Bridges in Acid-Base Reactions. Science 2005, 310, 83–86. (7) Mohammed, O. F.; Pines, D.; Nibbering, E. T. J.; Pines, E. BaseInduced Solvent Switches in Acid-Base Reactions. Angew. Chem., Int. Ed. 2007, 46, 1458–1461. (8) Siwick, B. J.; Cox, M. J.; Bakker, H. J. Long-Range Proton Transfer in Aqueous Acid-Base Reactions. J. Phys. Chem. 2008, 112, 378–389. (9) Cox, M. J.; Bakker, H. J. Parallel Proton Transfer Pathways in Aqueous Acid-Base Reactions. J. Chem. Phys. 2008, 126, 174501. (10) Cox, M. J.; Timmer, R. L. A.; Bakker, H. J.; Park, S.; Agmon, N. Distance-Dependent Proton Transfer along Water Wires Connecting Acid-Base Pairs. J. Phys. Chem. A 2009, 113, 6599–6606. (11) Balakrishnan, G.; Weeks, C. L.; Ibrahim, M.; Soldatova, A. V.; Spiro, T. G. Protein Dynamics from Time Resolved UV Raman Spectroscopy. Curr. Opin. Struct. Biol. 2008, 18, 623–629. (12) Chung, H. S.; Shandiz, A.; Sosnick, T. R.; Tokmakoff, A. Probing the Folding Transition State of Ubiquitin Mutants by Temperature-JumpInduced Downhill Unfolding. Biochemistry 2008, 47, 13870–13877. (13) Gutman, M.; Huppert, D.; Pines, E. The pH Jump: A Rapid Modulation of pH of Aqueous Solutions by a Laser Pulse. J. Am. Chem. Soc. 1981, 103, 3709–3713. (14) Abbruzzetti, S.; Grandi, E.; Viappiani, C.; Bologna, S.; Campanini, B.; Raboni, S.; Bettati, S.; Mozzarelli, A. Kinetics of Acid-Induced Spectral Changes in the GFPmut2 Chromophore. J. Am. Chem. Soc. 2005, 127, 626–635. (15) Genosar, L.; Cohen, B.; Huppert, D. Ultrafast Direct Photoacid Base Reaction. J. Phys. Chem. A 2000, 104, 6689–6698. (16) Pines, D.; Nibbering, E. T. J.; Pines, E. Relaxation to Equilibrium Following Photoacid Dissociation in Mineral Acids and Buffer Solutions. J. Phys.: Condens. Matter 2007, 19, 065134. (17) Adamczyk, K.; Premont-Schwarz, M.; Pines, D.; Pines, E.; Nibbering, E. T. J. Real-Time Observation of Carbonic Acid Formation in Aqueous Solution. Science 2009, 326, 1690–1694. (18) Causgrove, T. P.; Dyer, R. B. Nonequilibrium Protein Folding Dynamics: Laser-Induced pH-Jump Studies of the HelixCoil Transition. Chem. Phys. 2006, 323, 2–10. (19) Abbruzzetti, S.; Sottini, S.; Viappiani, C.; Corrie, J. E. T. AcidInduced Unfolding of Myoglobin Triggered by a Laser pH Jump Method. Photochem. Photobiol. Sci. 2006, 5, 621–628. (20) Miksovska, J.; Larsen, R. W. Photothermal Studies of pH Induced Unfolding of Apomyoglobin. J. Protein Chem. 2003, 22, 387–394. (21) Abbruzzetti, S.; Viappiani, C.; Small, J. R.; Libertini, L. J.; Small, E. W. Kinetics of Local Helix Formation in Poly-L-glutamic Acid Studied by Time-Resolved Photoacoustics: Neutralization Reactions of Carboxylates in Aqueous Solutions and Their Relevance to the Problem of Protein Folding. Biophys. J. 2000, 79, 2714–2721. (22) Yam, R.; Nachliel, E.; Gutman, M. Time-Resolved Proton Interaction. Methodology and Kinetic Analysis. J. Am. Chem. Soc. 1988, 110, 2636–2640. (23) Gutamn, M.; Nachliel, E. Time-Resolved Dynamics of Proton Transfer in Proteinous Systems. Annu. Rev. Phys. Chem. 1997, 48, 329–356. (24) Gutman, M.; Nachliei, E.; Gershon, E. Effect of Buffer on Kinetics of Proton Equilibration with a Protonable Group. Biochemistry 1985, 24, 2937–2941.
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(25) Gutman, M.; Nachliel, E.; Gershon, E.; Giniger, R.; Pines, E. Ph Jump: Kinetic Analysis and Determination of the Diffiusion-Controlled Rate Constants. J. Am. Chem. Soc. 1983, 105, 2210–2216. (26) Checover, S.; Nachliel, E.; Dencher, N. A.; Gutman, M. Mechanism of Proton Entry into the Cytoplasmatic Section of the Proton-Conducting Channel of Bactiriorhodopsin. Biochemistry 1997, 36, 13919–13928. (27) Pines, E.; Huppert, D. The pH-Jump: A Relaxational Approach. J. Phys. Chem. 1983, 87, 4471–4478. (28) Laimgruber, S.; Schreier, W. J.; Schrader, T.; Koller, F.; Zinth, W.; Gilch, P. The Photochemistry of o-Nitrobenzaldehyde as Seen by Femtosecond Vibrational Spectroscopy. Angew. Chem., Int. Ed. 2005, 44, 7901–7904. (29) Yip, R. W.; Sharma, D. K. The Reactive State in the Photorearrangment of o-Nitrobenzaldehyde. Res. Chem. Intermed. 1989, 11, 109–116. (30) Donten, M. L.; Hamm, P.; VandeVondele, J. A Consistent Picture of the Proton Release Mechanism of oNBA in Water by Ultrafast Spectroscopy and Ab Initio Molecular Dynamics. J. Phys. Chem. B 2011, 115, 1075–1083. (31) The minor addition of protons through nondeuterated reagents (NaOH, oNBA, CH313COOH) was neglected because their concentrations were kept on a millimolar level. (32) Covington, A. K.; Paabo, M.; Robinson, R. A.; Bates, R. G. Use of the Glass Electrode in Deuterium Oxide and the Relation between the Standardized pD (paD) Scale and the Operational pH in Heavy Water. Anal. Chem. 1968, 40, 700–706. (33) Hamm, P.; Kaindl, R. A.; Stenger, J. Noise Suppression in Femtosecond Mid-Infrared Light Sources. J. Opt. Lett. 2000, 25, 1798 . (34) Bredenbeck, J.; Helbing, J.; Hamm, P. Continuous Scanning from Picoseconds to Microseconds in Time Resolved Linear and Nonlinear Spectroscopy. Rev. Sci. Instrum. 2004, 75, 4462–4466. (35) Paabo, M.; Bates, R. G.; Robinson, R. A. Dissociation of Acetic Acid-d4 in Deuterium Oxide from 5 to 50° and Related Isotope Effects. J. Phys. Chem. 1966, 70, 2073–2077. (36) The difference of a factor of 2 between the proton release time given in ref 17 and obtained in this fit is most likely caused by fitting different marker bands (here the carboxyl group of the photoproduct, in ref 17 the carboxylate band rising after the dissociation) and should be treated as within the experimental error, given the fact that bands strongly overlap and sit on strong backgrounds. (37) Nibbering, E. T. J.; Fidder, H.; Pines, E. Ultrafast Chemistry: Using Time-Resolved Vibrational Spectroscopy for Interrogation of Structural Dynamics. Annu. Rev. Phys. Chem. 2005, 56, 337–367. (38) Aburuzzetti, S.; Carcelli, M.; Rogolino, D.; Viappiani, C. Deprotonation yields, pKa, and aci-Nitro Decay Rates in Some Substituted o-Nitrobenzaldehydes. Photochem. Photobiol. Sci. 2003, 2, 796–800. (39) The deviations between the measured and simulated traces at pD 11.4 and 11.9 at long times reflect baseline drifts in the experimental data
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