Negligible Isotopic Effect on Dissociation of Hydrogen Bonds

Article abstract of DOI:10.1021/acs.jpcb.5b12652

Isotopic effects on the formation and dissociation kinetics of hydrogen bonds are studied in real time with ultrafast chemical exchange spectroscopy. The dissociation time of hydrogen bond between phenol-OH and p-xylene (or mesitylene) is found to be iden

Full text of DOI:10.1021/acs.jpcb.5b12652

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Article
Negligible Isotopic Effect on Dissociation of Hydrogen Bonds
Chuanqi Ge, Yuneng Shen, Gang-hua Deng, Yuhuan Tian,
Dongqi Yu, Xueming Yang, Kaijun Yuan, and Junrong Zheng
J. Phys. Chem. B, Just Accepted Manuscript • DOI: 10.1021/acs.jpcb.5b12652 • Publication Date (Web): 11 Mar 2016
Downloaded from http://pubs.acs.org on March 13, 2016
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Negligible Isotopic Effect on Dissociation of Hydrogen Bonds
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Chuanqi Ge1,2¶,Yuneng Shen2,5¶, Gang-Hua Deng2, Yuhuan Tian2, Dongqi Yu1,
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Xueming Yang2, Kaijun Yuan2*, Junrong Zheng2,3,4
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1School of Physics and Electronic Technology, Liaoning Normal University, Dalian
116029, China
2State Key Laboratory of Molecular Reaction Dynamics, Dalian Institute of the
Chemical Physics, ChineseAcademy of Sciences, Dalian 116023, Liaoning, China
3Department of Chemistry, Rice University, Houston, Texas, USA
4College of Chemistry and Molecular Engineering, Beijing National Laboratory for
Molecular Sciences, Peking University, Beijing 100871, China
5University of the Chinese Academy of Sciences, Beijing 100049, China
¶The authors have equal contribution.
*To whom correspondence should be addressed:
kjyuan@dicp.ac.cn; junrong@rice.edu
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Abstract
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Isotopic effects on the formation and dissociation kinetics of hydrogen bonds are
studied in real time with ultrafast chemical exchange spectroscopy. The dissociation
time of hydrogen bond between phenol-OH and p-xylene (or mesitylene) is found to
be identical to that between phenol-OD and p-xylene (or mesitylene) in the same
solvents. The experimental results demonstrate that the isotope substitution (D for H)
has negligible effects on the hydrogen bond kinetics. DFT calculations show that the
isotope substitution doesn’t significantly change the frequencies of vibrational modes
that may be along the hydrogen bond formation and dissociation coordinate. The zero
point energy differences of these modes between hydrogen bonds with OH and OD
are too small to affect the activation energy of the hydrogen bond dissociation in a
detectible way at room temperature.
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1. Introduction
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Hydrogen bonds are of fundamental importance in nature. It plays an
important role in defining the three dimensional structures, dynamics, and functions
of chemical and biochemical molecular systems.1-6 Traditionally, hydrogen bonds
exist between electronegative atoms (e.g., O, N, F and Cl) and H atoms covalently
bound to similar electronegative atoms.7 Recently this concept has been extended to
unusual acceptors such as π systems which this work focuses on.8 Most of these two
types of hydrogen bonds are substantially weaker than a typical covalent bond and
have a wide range of binding energies from about 4 to 40 kJ/mol.9-11 Therefore, they
can frequently dissociate and reform at room temperature (RT). This reversible nature
of hydrogen-bond formation play a critical role in biological processes which depend
upon hydrogen-bonding, such as DNA replication, protein folding, drug-receptor, and
other recognition interactions.4,11 To comprehensively understand their mechanisms
involved in the dynamics, obtaining hydrogen-bond formation and dissociation rates
is desired. Because of the advent of ultrafast IR techniques, researchers can measure
these kinetic processes of relatively weak hydrogen bonds (<5 kcal/mol) at RT under
thermal equilibrium conditions with adequate time resolution.9,11-30 A series of studies
on the lifetimes of hydrogen bonds between phenol-OD (and silanol-OD) and
hydrogen bond acceptors using the method of ultrafast 2D-IR vibrational echo
chemical exchange spectroscopy show that the dissociation rates of the weak
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hydrogen bonds are strongly correlated with their formation enthalpies9,13-15,31,32
which can be described by an equation similar to the Arrhenius equation.
,
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The majority of the studies mentioned above were performed on systems with
deuterated hydrogen bonds (OD…acceptor).9,13-15,32 Studies at the same quantitative
level on regular hydrogen bond (OH…acceptor) dissociation kinetics have not been
conducted. Though chemically very similar, deuterated and regular hydrogen bonds
can have different dissociation kinetics. It is well known that many processes
involving hydrogen bonds or the transfer of a proton have different kinetics if the
proton is replaced by a deuterium. For example, the dissociation of C-H can be
several times faster than that of C-D33, and the vibrational relaxation of OD is much
longer than that of OH in water solution.34-36 If the zero point energies of vibrational
modes associated with hydrogen bond dissociations are significantly affected by the
D-H replacement, isotope effects observed in regular chemical reactions can also
occur in hydrogen bonds.
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In this work, we performed both experiments and ab initio calculations to
study the isotope effects on hydrogen bond dissociation kinetics. Ultrafast chemical
exchange spectroscopy was applied to monitor the real time formation and
dissociation dynamics of both deuterated and regular hydrogen bonds. Ab initio
calculations were performed to compute the binding energy, zero point energy, and
normal mode analyses of the hydrogen bond systems. Both experimental and
theoretical studies suggest that isotope effects at room temperature are negligible in
these systems investigated: 0.6 wt% phenol or phenol-OD in the mixed p-xylene/CCl4
(1:7 wt) solvent, 0.6 wt% phenol or phenol-OD in the mixed mesitylene/CCl4 (1:7 wt)
solvent. An issue needs to be clarified here. The donor/acceptor interaction of the π
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H-bonds investigated here has two contributions: (1) the H-bond between H and the π
electrons; and (2) the interaction between the lone electron pairs of oxygen atoms and
the π electron orbitals. If the latter was dominant, our experimental results could not
reflect the isotope effect of H-bond kinetics. Luckily, previous and current
experiments have suggested that the interaction between the effect of lone electron
pairs of oxygen atoms and the π electron orbitals on affecting the donor/acceptor
dissociation/association kinetics is negligible, compared to that of the H-bond
between H and the π electrons13,14,37. In a series of phenol/benzene, phenol/tolunene,
phenol/p-xylene, and phenol/mesitylene complexes, the H-bond donor (phenol) is the
same. The acceptors are different with more methyl groups adding to the benzene ring:
benzene has no methyl group, toluene with one, p-xylene with two, and mesitylene
with three. It is well known that the methyl group is electron donating, which can
enhance the electronic density of the benzene ring and increase its H-bond accepting
ability. On the other hand, more methyl groups must weaken the ability of the
benzene orbitals accepting the lone electrons of oxygen of phenol because electrons
repulse each other. If the interaction of benzene orbitals and the lone electrons of
oxygen of phenol is dominant, one would observe that the dissociation is fastest in the
phenol/mesitylene complex and slowest in the phenol/benzene complex in the order
of phenol/mesitylene (fastest) >phenol/p-xylene > phenol/toluene > phenol/benzene
(slowest). If the H-bond is dominant, the opposite trend would be observed.
Experimentally, the dissociation of phenol/benzene is fastest and that of
phenol/mesitylene is slowest with the dissociation time constants: 9ps
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(phenol/benzene), 15ps (phenol/tolunene), 22ps (phenol/p-xylene), and 31ps
(phenol/mesitylene). The results indicate that the π H-bond dissociation/association
kinetics is dominated by the H-bond and that of the interaction of benzene orbitals and
the lone electrons of oxygen of phenol is negligible.
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2. Experimental methods
The experimental setup has been described elsewhere.38-42 Briefly, the
experimental setup consisted of a Ti:Sapphire oscillator (~400 mw, ~800nm with a
band width ~45 nm,<35 fs, ~76 MHz, LIGHTHOUSE PHOTONICS sprout
combining QUANTRONIX Ti-Light), two Ti:Sapphire regenerative amplifiers and
two optical parametric amplifier (OPA) systems. The oscillator pumps the two
Ti:Sapphire regenerative amplifiers,90% power entering a picosecond amplifier(~3.5
W, ~800 nm with a band width ~0.6 nm, ~1.3 ps, 1 kHz, QUANTRONIX Integra-C)
and 10% to a femtosecond amplifier (~3.5 W, ~800 nm with a band width ~26 nm,
~40 fs, 1 kHz, QUANTRONIX Integra-C) respectively. We use ~2.8Wof the
picosecond amplifier to pump a picosecond OPA(TOPAS-800) plus non-collinear
difference frequency generator (NDFG) system to producing ~1.4ps (vary from
1.0~1.5ps in different frequencies) mid-IR pulses with a bandwidth ~15cm-1 in a
tunable frequency range from 500cm-1 to 4000cm-1 with energy1~40µJ/pulse at 1
KHz.We also use ~2.8Wof the femtosecond amplifier to pump another femtosecond
OPA (TOPAS-Prime) and NDFG system producing ~140fs mid-IR pulses with a
bandwidth ~200cm-1 in a tunable frequency range from 500cm-1 to 4000 cm-1 with
energy1~40µJ/pulse at 1 KHz. In the 2D IR experiments, the ps IR pulse excites the
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vibrations of the molecules and its excitation power is adjusted based on need. The fs
IR pulse is the probe beam. It is frequency resolved by a spectrograph and then is
collected on a 2×32 pixel mercury cadmium telluride (MCT) detector (Infrared
Associates)yielding the ωm axis of a 2D spectrum. The optical resolution used was
4-7 cm-1 (depending on the used grating). Scanning the pump frequency yields the ωτ
axis of the spectrum. Two polarizers are inserted into the probe beam pathway, one is
located immediately behind the sample to selectively measure the parallel or
perpendicular polarized signal relative to the pump beam and the other is before the
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sample rotating the polarization of the probe pulse 45 with respect to that of the
°
pump pulse. The entire system is computer controlled. The IR pump probe signal P(t)
is collected by measuring the transmission of the probe beam through the sample by
chopping the pump beam at 500 Hz. For a given delay time t (the delay between the
IR pump and the IR probe pulse was changed by changing the optical path length
travelled using a translation stage), the IR pump probe signal is defined by
P(t)=[Ipum-on - Ipump-off](t)/Ipump-off=∆I(t)/I where I is the transmission of the probe beam.
P = P + 2× P
Vibrational lifetimes are obtained from the rotation-free signal
,
life
//
⊥
P
where
P
and
are parallel and perpendicular data, respectively.
⊥
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Phenol was purchased from SIGMA-ALDRICH. P-xylene was purchased
from Tianjin Kemiou Chemical Reagent Co., Ltd.. Mesitylene was purchased from
Tianjin Guangfu Fine Chemical Research Institute. Carbon tetrachloride was
purchased from Xilong Chemical Industry Incorporated Co. Ltd..These samples were
used without further purification.
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To obtain the deuterated hydroxyl (OD) of the phenol, phenol was dissolved in
methanol-OD (1g/10g) and stirred using magnetic stirring apparatus for 0.5 h. After
the deuterium exchange completed, the solvent was then drew away by pump.
Meanwhile the sample was under vacuum. The procedure was repeated many times
until the compounds with >90% deuteration of the OH were obtained. The
experimental liquid samples were contained in sample cells with Teflon spacers that
provide suitable sample thickness according to their optical densities. The cell
windows were made of two ~2 mm CaF2 and provided sufficiently low group velocity
dispersion. Experiments were performed at RT.
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Density functional theory (DFT) calculations43 of the structures of the
intermolecular hydrogen bonded complexes were performed by using (DFT/B3LYP)
method with the 6-311+G(d,p) or 6-31G* basis set level. All the DFT calculations
were performed using Gaussian 09 program package. In our calculations, we don't
include the surrounding solvent. Therefore all results reported in our paper are for the
isolated molecules.
3. Results and discussions
3.1.FTIR spectra and the complex/free concentration ratios
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(B)
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frequency (cm-1
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frequency (cm-1
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3650
)
Figure 1.FTIR absorption spectra of (A)the OD (hydroxyl H replaced with D) stretch
and (B)the OH stretch of phenol in CCl4 (free species, green curve), in p-xylene
(complex species, red curve), and in the mixed p-xylene/CCl4 solvent, which displays
absorptions for both free and complex species (black curve).FTIR absorption spectra
of (C) the OD stretch and (D) the OH stretch of phenol in CCl4 (free species, green
curve), in mesitylene (complex species, red curve), and in the mixed mesitylene/CCl4
solvent (free and complex species, black curve).
The studied systems are 0.6 wt% phenol-OD or phenol in the mixed
p-xylene/CCl4 (1:7wt) solvent, and 0.6 wt% phenol-OD or phenol in the mixed
mesitylene/CCl4 (1:7wt) solvent. In the solutions, the solutes have very low
concentration, which ensures that resonant energy transfer between the phenol-OD or
phenol species is negligible. In these solvent mixtures, the OD or OH stretching band
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separates into two distinguishable peaks, a broad resonance at low energy
corresponding to phenol molecules that donate a hydrogen bond to the π-system of a
neighboring aromatic molecule (the structures are presented in fig. 4&5) and a
narrower high frequency resonance assigned to molecules with no hydrogen bonding
partner. Experimental evidence for the formation of complexes is the shift of the OD
(or OH) stretch frequency of phenol-OD (or phenol) to a lower frequency in p-xylene
or mesitylene compared to that in CCl4. Fig.1 (A) shows the FTIR spectra of
phenol-OD in pure CCl4, pure p-xylene, and a mixed p-xylene/CCl4 solvent. When
phenol-OD is dissolved in CCl4, the solute is free and has no hydrogen bonding
partner. The OD stretch is at 2666 cm-1 and its peak is relatively narrow with the
half-width of 14 cm-1. When phenol-OD is dissolved in p-xylene, solute molecules
form a π hydrogen-bonded complex with solvent molecules. The frequency of OD
stretch redshifts to 2621 cm-1 and its fwhm becomes 30 cm-1. When phenol-OD is
dissolved in mixed p-xylene/CCl4 solvent, both free phenol-OD and complex exists
with thermal equilibrium. Fig.1 (B) shows the FTIR spectra of phenol in the same
solvents as phenol-OD in fig.1 (A). The free OH stretch is at 3610 cm-1 with
half-width of 14.5 cm-1. The OH stretch of complex is at 3544 cm-1 with half-width of
39 cm-1. Fig.1 (C) shows the FTIR spectra of phenol-OD in pure CCl4, in pure
mesitylene, and in a mixed mesitylene/CCl4 solvent. The free OD stretch is at 2666
cm-1, while the complex OD stretch is at 2614 cm-1. Fig.1 (D) shows the FTIR spectra
of phenol in the same solvents as phenol-OD in fig.1 (C). The free OH stretch is at
3610 cm-1, and the complex OH stretch is at 3533 cm-1.
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From the above FTIR spectra, we can also obtain the complex/free
concentration ratios. The method is given in the below. The transition dipole moment
of complex (free) phenol in pure solvent is assumed to be the same as complex (free)
in the mixed solvent. Then we can use them (solutes in pure solvents) to divide the
transition dipole moment and obtain the molar ratio of complex/free in mixed solvent,
IFmTIR =kC1µ2
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which is based on the fact that the FTIR signal of mixed solvent is
and
IFpTIR =kC2µ2
that of pure solvent is
, where I is the integration of the band, k is
constant, C1 is the unknown concentration of complex or free in mixed solvent, C2 is
the known concentration of complex or free in pure solvent and µ is the transition
dipole moment. Based on this procedure, the complex/free concentration ratios are
calculated to be 1.17, 1.22, 1.2 and 1.19 in phenol-OD/p-xylene/CCl4 (0.6 wt%, 1:7
wt), phenol/p-xylene/CCl4 (0.6 wt%, 1:7 wt), phenol-OD/mesitylene/CCl4 (0.6 wt%,
1:7 wt), phenol/mesitylene/CCl4(0.6 wt%, 1:7 wt) solutions, respectively (The
detailed calculation is provided in SI.) From the results, we find that the equilibrium
constant of the solute is almost same with that of deuterated one within experimental
uncertainty.
3.2. 2DIR spectra of chemical exchange
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2675
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0 ps
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free
complex
2675
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2615
2595
10 ps
16 ps
dissociation
formation
2565 2590 2615 2640 2665 2680
2565 2590 2615 2640 2665 2680
Figure 2.Time dependent 2DIR vibrational spectra (measured for parallel
polarization) of hydroxyl stretch of phenol-OD in the mixed mesitylene/CCl4 solvents
(0.6 wt%, 1:7 wt). The red contours on the diagonal (positive) are from the 0-1
vibration transitions, and the blue contours off the diagonal(negative, displayed
partly) are from the 1-2 vibrational transitions. Each spectrum has been normalized
to its largest peak for each time. Each contour is a 10% change. At 0 ps, there are two
peaks on the diagonal (red). With time increasing, dissociation and formation of the
phenol-OD-mesitylene complex causes two additional red peaks on the off-diagonal
to grow in. Between 0 ps and 5ps, the peaks change shape because of spectral
diffusion. Therefore, the time for spectral diffusion is<5 ps.
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Fig. 2 displays four 2DIR spectra (only the 0-1 transition region is shown) of
phenol-OD in a mesitylene/CCl4 mixture for a very short time (0 ps) and three long (5,
10, 16 ps) reaction period. At 0 ps, the red peaks (the 0-1 vibration transitions of the
free phenol-OD and the phenol-OD-mesitylene complex) appear only on the diagonal
positions in the 2DIR spectrum, showing that dissociation and formation of the
complex has not occurred and the complex or free species in the sample are
unchanged. The red peaks correspond to the absorption bands shown in fig. 1(C) for
the mixed solvent. As time increases, off-diagonal peaks appear due to chemical
exchange. At long time, 16 ps, extensive dissociation and formation of the complex
has occurred, then causes additional peaks which can be seen clearly to grown in. The
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peak labeled as “dissociation” has its
(initial frequency) at the lower frequency,
ωτ
showing that the initial structure was complex phenol-OD. Its
(final frequency) is
ωm
at higher frequency, which shows the final structure is free phenol-OD. This
off-diagonal peak arises from complex species converting to free species. The peak
labeled as “formation” is the exact opposite of “dissociation”. The molecules in the
two diagonal peaks are the same species, and they have either remained the same or
have exchanged an even number of times during the chemical exchange process.
At short time, it is shown that the 2D vibrational spectrum is elongated along
the diagonal. This is caused by inhomogeneous broadening.9,13-15,44 By 5 ps they have
become symmetrically rounded about the diagonal. The change in the 2D line shape is
caused by spectral diffusion.9,13-15,44
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The time dependent growth of the off-diagonal peaks provides direct data for
us to inspect the chemical exchange process. From fig.2 we can see that the two
off-diagonal peaks grow in at the same rate, showing that the rate of formation equals
the rate of dissociation. As a result, the system is in equilibrium. However, it is
difficult for us to directly compare the chemical exchange dynamics of the four
systems with the 2DIR spectra (from the following sections, we can know that their
chemical exchange times are too close to see difference in their 2DIR spectra). So we
only present the 2DIR spectra for the phenol-OD/mesitylene/CCl4 system to illustrate
the chemical exchange process.
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3.3. Chemical exchange kinetic model
Usually, we should analyze the peak volumes of the 2DIR spectra to extract
the chemical exchange kinetic.9,13,45 However, if the chemical exchange rate is much
slower than spectral diffusion, we can use the peak intensities of 2DIR spectra to
obtain the chemical exchange dynamics.9,13 In this way, we can determine the
chemical exchange rate more easily, and it only introduce small uncertainty. In our
experiments we use peak intensities to fit chemical exchange times (the following
sections will show that chemical exchange times of our studied systems are much
longer than spectral diffusion). We use a kinetic model which include all dynamic
processes.13,14,38,39,45,46 The method is illustrated schematically for the complex species
(CC) converting to free species (CF).
kd
kC
kF
⎯⎯→
←⎯⎯ CC
CF ⎯⎯→
[
]
[
]
←⎯⎯
ka
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where [CC] denotes the population of excited complex molecules, and [CF] denotes
the population of excited free molecules. It should be noted that dependence on
orientational relaxation of population is eliminated using rotation-free operation
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(
, see the details in experimental methods section). kd represents the
⊥
P = P + 2× P
life
//
rate constant for dissociation of the complex and
k
a represents the rate constant for the
association of the complex. and are the vibrational lifetimes of [CC]and[CF],
1 kC
1 kF
respectively, which can be single or biexponetial.38,39 When we separate one species
into two subgroups, the weighing factor of each subgroup is determined by the
prefactor of the bi-exponential. Each subgroup has a single-exponential-decay lifetime
time. Each subgroup can exchange population with other species, but the subgroups
can’t exchange population with each other (this follows the assumed physical picture
of bi-exponential: the sub-components can be considered as independent species). In
the model, the effects of dynamics such as spectral diffusions, vibrational decays and
vibrational exchanges within each species are simply treated as apparent vibrational
decays as experimentally measured. This is because not only the relaxation of the
diagonal peaks is complex, but also the chemical exchange process is independent of
the origin of the relaxation. Therefore, in our experiments we focus on the
off-diagonal peaks which arise from the chemical exchange, and this treatment will
not affect the value of chemical exchange lifetime we want to obtain.
According to the scheme showed above, we have kinetic equations:
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⎧d CC
[
]
]
= −k CC + k CF − k CC
[
]
[
]
[
]
]
⎪
d
a
C
⎪
⎨
⎪
dt
d CF
[
eq. 1
= −k CF + k CC − k CF
[
]
[
]
[
a
d
F
⎪
dt
⎩
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In the equations, the time dependent populations were obtained from the
chemical exchange measurements (the method is described in the supporting
information). In calculating the data, the apparent lifetimes determined by
pump/probe experiments were used as initial input values and allowed them to vary at
most 20%. Because the systems are in equilibrium, the rate of formation equals the
rate of dissociation.13 Therefore, the rate ratio
is determined by the complex/free
kd ka
concentration ratio (D). The dissociation rate constant kd of complex is independent of
concentration, and therefore, it is reported and discussed. is the association
ka = Dkd
rate constant. The association rate is ka [free]. Because the concentration ratio (D) is
known, we can assume the single unknown parameter is kd
.
3.4.Hydrogen bond lifetimes of hydroxyl and deuterated hydroxyl for
solute/solvent complexes are equal
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1.0
0.8
0.6
0.4
0.2
0.0
1.0
(A)
(B)
0.8
Free peak (diagonal)
Complex peak (diagonal)
0.6
0.4
0.2
0.0
Free peak (diagonal)
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Complex peak (diagonal)
Exchange peaks (off-diagonal)
Exchange peaks (off-diagonal)
0
10
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50
0
10
20
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50
T ime (ps )
T ime (ps )
1.0
0.8
0.6
0.4
0.2
0.0
1.0
0.8
0.6
0.4
0.2
0.0
Complex peak (diagonal)
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(C)
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Free peak (diagonal)
Exchange peaks (off-diagonal)
Exchange peaks (off-diagonal)
0
10
20
30
40
50
0
10
20
30
40
50
T ime (ps )
T ime (ps )
Figure 3.Time dependent data (symbols) of the 2 off-diagonal and 2diagonal peak
intensities in the 0-1 portion of the pump probe vibrational spectra for (A) phenol-OD
in the p-xylene/CCl4 mixture, (B) phenol in the p-xylene/CCl4 mixture, (C) phenol-OD
in the mesitylene/CCl4 mixture and (D) phenol in the mesitylene/CCl4 mixture. In each
figure, we use the solid curves calculated from the chemical exchange kinetic model
with one adjustable parameter, kd, to fit the experimental data. Other parameters in
the model are determined experimentally. Using these parameters, we can obtain
1/kd=19±2ps, 20±3p, 31±2ps and 33±3ps for(A), (B), (C) and (D), respectively.
Figure 3A shows the peak intensity data for the 0-1 transition region which
consists of four peaks: the two diagonal peaks for the complex and free phenol-OD
species, and the two off-diagonal peaks (complex species convert to free species and
free to complex). With the time increasing, the off-diagonal peaks grow in because of
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complex dissociation and formation. By inputting the known constants in the kinetic
equations, we can fit the four peaks with a single adjustable parameter, the complex
dissociation time, 1/kd. For the phenol-OD/p-xylene/CCl4system, the input parameters
used are listed in table 1. All the vibrational lifetimes in table 1 are experimentally
determined in individual solution with pure solvent and all behave like biexponential.
As can be seen from fig. 3A, the calculated solid curves fit in very well with the time
dependent data of all of the peaks. From the fits, the dissociation time(hydrogen bond
lifetime) for the phenol-OD/p-xylene complex is obtained and is listed in table 1.
Table 1.The complex/free concentration ratios, the vibrational lifetimes and its
corresponding weighing factors of each subgroup and the dissociation rate constants
of complexes in the four mixtures. (A), (B), (C) and (D) correspond to the solutions
in fig.3, respectively.
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Complex species
Free species
Complex/free
concentration
ratio D
Fast process Slow process Fast Process Slow process 1/kd
1/kC-fast
AC-fast
1/kC-slow AC-slow
1/kF-fast
AF-fast
1/kF-slow
AF-slow
(ps)
(ps)
(ps)
(ps)
(ps)
(A)
(B)
(C)
(D)
1.17
1.22
1.2
4.6
5
0.454 12.3
0.546
0.2
11
3.5
8
0.3
20
20
14
0.7
19±2
20±3
31±2
33±3
0.8
20
0.51
0.3
0.49
0.7
4
0.5
16
0.5
1.19
4
0.55
12.3
0.45
3
0.465 21
0.535
When analyzing the data in Figure 3, we assumed that the equilibrium
constants and the exchange rates were not disturbed by vibrational excitation. This is
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supported by the following evidence. Such as both off-diagonal peak intensities
(triangles and squares in Figure 3) have the same time dependence, the rates of
complex formation and dissociation are equal. Namely, the aggregate populations of
free and complex species in the mixtures are not affected by the excitation sequence.
Other further tests by comparing time evolution of the 1-2 and 0-1 transition regions
showing that the systems are in equilibrium have been presented for the
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phenol/benzene system13 and many other systems9,14,15
.
For the phenol/p-xylene/CCl4system (Figure 3B), the input parameters used
are listed in table 1. As shown in fig. 3B, the calculated solid curves of the free
diagonal peak misses a little with the time dependent data of that peak. However, the
calculated solid curves of the other three peaks fit in very well with the time
dependent data of those peaks. It should be noted that not only the free diagonal peak
data in mixed solvent has a recurrence on its experimental curve at 12 ps, but also the
diagonal peak data in pure solvent has the same phenomenon. Details are provided in
SI. The recurrence is very interesting, but its origin is not known at this moment
which is subject to future studies. The recurrence has little effect on the determination
of the dissociation time, because the dissociation time is mainly determined by the
growth of the off-diagonal peaks.15 From the fits, the dissociation time for the
phenol/p-xylene complex (hydrogen bond lifetime) is obtained and listed in table 1.
It is interesting that the measured chemical exchange rates show that the
hydrogen bond lifetimes of the two systems (OH…p-xylene and OD…p-xylene) are
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equal within experimental uncertainty. There is a negligibly small isotopic effect on
chemical exchange dynamics of phenol solutions.
8
9
Using the same method, we also obtain the dissociation times for the
phenol-OD/mesitylene/CCl4 system (Figure 3C) and the phenol/mesitylene/CCl4
system (Figure 3D).The used input parameters and the results are listed in table 1. As
can be seen from fig. 3C&D, the calculated solid curves fit in very well with the time
dependent data of all of the peaks. The chemical exchange experiment results again
show that the hydrogen bond lifetimes of the two systems are equal within
experimental uncertainty.
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3.5.DFT calculation results for solute/solvent complexes of hydroxyl and
deuterated hydroxyl are almost identical
The gas-phase structures for hydrogen bonds have been determined by DFT
calculations. The results show that the phenol donate a hydrogen bond to the π-system
of a neighboring aromatic molecule in the mixed solvents and all of the complexes
have a similar T-shaped structure. In addition, the T-shape structure is identical for
both deuterated and regular hydrogen bonds. The structures of phenol/p-xylene and
phenol/mesitylene complexes are displayed in figure 4&5. In a previous work31, for
the phenol/benzene complex, the T-shape structure has been confirmed by high-level
electronic structure calculations and full molecular dynamic simulations of phenol in
the benzene/CCl4 solvent. Because of the consistency between the simulations and the
electronic structure calculations for the phenol/benzene system, the electronic
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structure calculations for the other complexes provide a reasonable description of
their structures.9,13,14
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(a)
(b)
Figure 4.Two views of the hydrogen bond structure of the phenol-p-xylene complex or
phenol-OD-p-xylene complex calculated with DFT at B3LYP/6-31+G(d,p) level for
isolated molecules (calculations show that the structures of the two complexes are
almost identical). The binding energies of the two complexes are same, and are
predicted to be -2.7 kcal/mol without any solvent molecules (with zero point energy
correction).
(a)
(b)
Figure 5.Two views of the hydrogen bond structure of the phenol-mesitylene complex
or phenol-OD-mesitylene complex calculated with DFT at B3LYP/6-31G* level for
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isolated molecules(calculations show that the structures of the two complexes are
almost identical). The binding energies of the two complexes are same, and are
predicted to be-3.9 kcal/mol without any solvent molecules (with zero point energy
correction)
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We then used calculations to analyze the vibrational modes of the hydrogen
bonded complexes to obtain the directions and the zero point energies of the
vibrations, investigating the principles leading to the observed negligible isotope
effect on hydrogen bond dissociations. It is well known that, if a vibration is along the
1
reaction coordinate, its zero point energy (the zero-point energy is defined by
,
hν
2
where
ν
is the 0-1 transition vibrational frequency.) can affect the activation energy
of the reaction, leading to a remarkable isotope effect on the reaction kinetics.47The
idea can be simply illustrated in Figure 6.
Dissociation
E
ΔHH
ΔHD
VX-H(0)
VX-D(0)
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Figure 6.Illustration of isotope effect caused by the zero point energy difference
between X-H and X-D stretches for the dissociation of X-H and X-D bonds. The lower
zero point energy of X-D stretch results in a larger dissociation enthalpy and a slower
kinetics. The tunneling effect is not considered here.
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Table 2.Vibrations of phenol-p-xylene and phenol-OD-p-xylene complexes that can
be involved in hydrogen bond dissociation
Vibrations (cm-1)
Phenol/p-xylene
1
2
3
4
5
6
7
12.33
12.28
15.92
15.89
30.50
30.42
38.14
37.79
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43.32
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310.46
308.59
phenol-OD/p-xylene
Table 3.Vibrations of phenol-mesitylene and phenol-OD-mesitylene complexes that
can be involved in hydrogen bond dissociation
Vibrations (cm-1)
1
2
3
4
5
6
7
8
9
13.51
13.49
21.32
21.28
28.52
28.48
37.62
37.35
52.44
52.02
82.39
81.87
225.64
225.22
233.46
232.99
237.85
235.78
Phenol/mesitylene
phenol-OD/mesitylene
The frequencies of vibrations (0-1 transition) that can be involved in hydrogen
bond dissociation of phenol/p-xylene, phenol-OD/p-xylene, phenol/mesitylene, and
phenol-OD/mesitylene are listed in tables 2&3 (the corresponding pictures for
vibrations are provided in SI.). The largest frequency difference between deuterated
and regular hydrogen bonds is 2.07 cm-1 (No.9 in table 3), corresponding to a zero
point energy difference 1.04 cm-1. The calculated phenol/mesitylene hydrogen bond
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dissociation enthalpy is 3.9 kcal/mol (1362 cm-1). At RT, according to the Arrhenius
equation, the zero point energy difference only causes 0.5% difference in dissociation
rate which is about 0.2ps for the dissociation (33ps) of phenol/mesitylene. The
difference 0.2ps is smaller than our experimental uncertainty 2-3 ps, and therefore
was not detected in the experiments. Our calculations also show that the total zero
point energy difference between phenol/mesitylene and phenol-OD/mesitylene is
0.04/kcal/mol. Even if the total zero point energy completely contributes to the
hydrogen dissociation activation energy, at room temperature the energy difference
only causes a rate difference of 2ps which is still within our experimental uncertainty.
The result for the phenol/p-xylene and phenol-OD/p-xylene is similar. The calculated
total zero point energy difference between them is only 0.03kcal/mol. The energy
difference only causes a rate difference of 1ps. Here, it is worth to emphasize that the
calculated complexes in the gas phase may not have the same structures as those in
real liquids. However, the main purpose of the calculations is to examine whether the
H/D replacement will cause any frequency shifts of those low frequency vibrations
when one phenol and one benzene derivative form a complex. Therefore, whether the
calculated structures are close to real structures is expected not to affect the
conclusions drawn from current calculations that much. In fact, MD simulations show
that structure of a phenol/benzene complex calculated in this way is actually close to
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.
In summary, the ab initio calculations show that the values of zero energy of
vibrations that can contribute to the dissociations of regular and deuterated hydrogen
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bonds are very similar. The differences are too small to cause detectible dissociation
rate changes in our experiments. The experimental results also indicate that the
hydrogen bond strengths are very similar for the OH and OD complexes studied in
this work. This is somewhat different from that D2O has stronger H-bonds than
H2O.50 However, DFT calculations performed in this work cannot provide
information to address this issue. Future theoretical studies by considering nuclear
quantum effects can be very helpful51.
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4. CONCLUDING REMARKS
In this work, isotopic effects on the formation and dissociation kinetics of
hydrogen bonds are studied in real time with ultrafast chemical exchange
spectroscopy. The dissociation time 20ps (33ps) of hydrogen bond between
phenol-OH and p-xylene (or mesitylene) is found to be identical to that between
phenol-OD p-xylene (or mesitylene) in the same solvents. The experimental results
demonstrate that the isotope substitution (D for H) has negligible effects on the
hydrogen bond kinetics. DFT calculations show that the isotope substitution doesn’t
significantly change the frequencies of vibrational modes that may be along the
hydrogen bond formation and dissociation coordinate. The zero point energy
differences of these modes between hydrogen bonds with OH and OD are very small.
The change of the hydrogen bond dissociation rate caused by the energy difference is
smaller than 2ps at room temperature, which is within our experimental uncertainty.
We expect the conclusion that the isotope effect is very small for hydrogen bond
dissociation kinetics at room temperature is not only limited in the phenol systems
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investigated here. It should be general for a large number of other hydrogen bonds if
the zero point energy difference is the only reason responsible for the isotope effect,
because the zero point energy changed by deuteration is too small, compared to the
thermal energy at room temperature. It is interesting to note that the H-bond strength
can vary upon isotope exchange in some systems, which may also affect its
dissociation and association kinetics.50 This topic will be subject to future studies.
Conflicts of interest
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The authors declare no competing financial interest.
Associated content
Supporting Information
The calculation parameters of the figures containing kinetic model calculations.
Acknowledgement
This work is supported by the National Natural Science Foundation of China (No.
21373213), the Chinese Academy of Sciences, and the Ministry of Science and
Technology.J. R. Zheng is supported by the AFOSR Award No. FA9550-11-1-0070,
the Welch foundation under Award No.C-1752, the David and Lucile Packard
Foundation for a Packard fellowship and the Alfred P. Sloan Foundation for a Sloan
fellowship.
REFERENCES
(1) Jeffrey, G. A.: An introduction to hydrogen bonding; Oxford university press:
New York, 1997; Vol. 12.
(2) Schuster, P.; Zundel, G.; Sandorfy, C.: The hydrogen bond: Recent
Developments; North-Holland: Amsterdam, 1976.
(3) Pimentel, G. C.; McClella.Al: Hydrogen Bonding. Ann. Rev. Phys. Chem.
1971, 22, 347-385.
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