Authors: Maciej Śmiechowski, Ingmar Persson
Categories: Article
Source: Inorganic Chemistry
The strength of hydrogen bonding to and structure of hydrated oxometallate ions in aqueous solution have been studied by double difference infrared (DDIR) spectroscopy and large-angle X-ray scattering (LAXS), respectively. Anions are hydrated by accepting hydrogen bonds from the hydrating water molecules. The oxygen atom of the permanganate and perrhenate ions form weaker and longer hydrogen bonds to water than the hydrogen bonds in bulk water (i.e., they act as structure breakers), while the oxygen atoms of the chromate, dichromate, molybdate, tungstate, and hydrogenvanadate ions form hydrogen bonds stronger than those in bulk water (i.e., they act as structure makers). The oxometallate ions form one hydration shell distinguishable from bulk water as determined by DDIR spectroscopy and LAXS. The hydration of oxoanions results in X–O bond distances ca. 0.02 Å longer than those in unsolvated ions in the solid state not involved in strong bonding to counterions. The oxygens of oxoanions with a central atom from the second and third series in the periodic table and the hydrogenvanadate ion hydrogen bind three hydrating water molecules, while oxygens of oxoanions with a heavier central atom only form hydrogen bonds to two water molecules.
A limited number of methods are available to study the hydration of anions in aqueous solution due to weak hydration and broad distance distribution of the hydrating water molecules. The strength of the hydrogen bonds between an anion and surrounding water molecules is preferably studied by double difference infrared (DDIR) spectroscopy.^1^ Structures of hydrated anions in aqueous solution are preferably studied with large-angle X-ray and neutron scattering, LAXS^2^ and LANS,^3^ respectively. The most commonly applied method to study structures in solution, EXAFS, is normally not applicable as the X–(O···H−)O distance is long and the distance distribution is wide, which causes the contribution to the EXAFS function to be very small and such distances to be hardly observable.^4^ Furthermore, the hydrogen atom has too weak backscattering ability to be observed accurately. However, LAXS is a very suitable method as it is very sensitive to long distances with a wide bond distance distribution according to the LAXS equation.^2^ Numerous LAXS studies on aqueous solutions have been reported, while only a very limited number of LANS studies on hydrated chloride, bromide, nitrate, and perchlorate ions have been reported.^5,6^
Besides the experimental DDIR and LAXS methods, the hydration of
oxoanions has been studied by computational methods. Both static quantum
chemical computations and molecular dynamics (MD) simulations (based
on ab initio electronic representation or force-field
parametrization) have been applied for this purpose. Since the data
on nonmetallic oxoanions have been previously reviewed and discussed
in detail,^7^ this summary will focus on
the few computational investigations devoted specifically to hydrated
oxometallate ions.^8−11^ An ab initio MD study on the hydrated permanganate
ion in aqueous solution reports an intramolecular Mn–O bond
length of 1.62 Å and four water molecules at a mean Mn(−O)···Oaq distance of ca. 3.8 Å in a hydration shell characterized
as being weak overall.^8^ A detailed MD study
of the hydrated chromate ion found that it is a moderate structure
maker with a coordination number in water of 13–14, with every
chromate oxygen forming on average 2.6 hydrogen bonds to water molecules,
and a Cr–O bond length of 1.61 Å.^9^ An ab initio MD study of aqueous tungstate revealed
that WO4^2–^ hydrolyzes easily to [WO~n(OH)aq~ distance is 2.856 Å.^11^ The
structure parameters and the O–D stretching frequencies measured
by the DDIR method of the hydrated oxoanions in aqueous solution reported
so far are summarized in Table 1.4–n]^(2–n)–^ ions that further
coordinate up to two additional water molecules in order to reach
the W coordination number of 6.^10^ Finally,
static calculations on aqueous vanadate clusters found the intramolecular
V–O bond length is 1.766 Å in hexahydrate, while the V(−O)···O
Hydrated ions are regarded
as structure makers or breakers in aqueous solution depending on their
ability to form stronger or weaker hydrogen bonds, respectively, to
anions or to water molecules in the first hydration shell of metal
ions in comparison to the intermolecular hydrogen bonds in pure water.^25^ Anions with low charge density, as is the case
of most monovalent ions except fluoride and hydroxide, are regarded
as water structure breakers, while di- and trivalent anions are regarded
as water structure makers. Previous DDIR and structural studies of
hydrated anions in aqueous solution show that the O–D stretching
frequency of hydrating water molecules binding to structure breaking
ions is higher than that in pure water, 2510 cm^–1^, and the (X−)O···(H)–O distances ((X−)O···Oaq) are longer than the O(−H)···O(−H)
distances in bulk water (Oaq···Oaq) (Table 1). The O–D
stretching frequency of the water molecules binding to structure making
ions is lower than 2510 cm^–1^, and the (X−)O···Oaq distances are shorter than 2.89 Å. The X–O,
X(−O)···Oaq, and (X−)O···Oaq distances are used to calculate the X–O···Oaq bond angle from which the mean number of water molecules
hydrogen-bound to each oxoanion oxygen can be estimated with approximately
two or three water molecules at X–O···Oaq bond angles of 120 and 109.5°, respectively. It can
be observed that oxoanions with a central atom belonging to the second
and third row in the periodic table bind three water molecules, while
those lower in the periodic table only bind two (cf. Table 1). Whether an oxoanion hydrogen
binds two or three water molecules cannot be distinguished in the
O–D stretching frequencies. For example, the sulfate and selenite
ions bind three and two water molecules per oxygen, respectively,
but they have the same O–D stretching frequency (Table 1).
The aim of this study is to determine the hydrogen bond strength and structure parameters of the hydrated oxometallate anions in aqueous solution and to summarize the present knowledge of the hydration of anions in aqueous solution.
Sodium
permanganate monohydrate, NaMnO4·H2O (Aldrich,
97%+), sodium perrhenate, NaReO4 (Aldrich, 99.99%), sodium
chromate tetrahydrate, Na2CrO4·4H2O (Mallinckrodt, p.a.), sodium dichromate dihydrate, Na2Cr2O7·2H2O (Merck, p.a), sodium
molybdate dihydrate, Na2MoO4·2H2O (Merck, p.a.), sodium tungstate dihydrate, Na2WO4·2H2O (Merck, p.a.), and sodium vanadate,
Na3VO4 (Aldrich, 99.98%) were used as supplied.
Heavy water (D2O, 99.9% isotopic purity) was supplied by
Aldrich.
Stock solutions were prepared by dissolving
amounts (weighed with
0.1 mg precision) of respective sodium salts in double-distilled water.
For NaMnO4, Na2CrO4, Na2MoO4, and Na2WO4, a series of solutions
spanning the molality range ∼0.2–1.0 mol·kg^–1^ were prepared by dissolving weighed amounts of the
respective stock solution in double-distilled water. For the NaReO4 and Na3VO4 salts, only the stock solutions
were used for IR spectra measurement. Sample solutions containing
HDO were prepared by adding 4% (by weight) of D2O relative
to H2O and reference solutions by adding the same molar
amounts of H2O. Low fraction of D2O is preferable,
because the unfavorable equilibrium (H2O + D2O ⇄ 2 HDO, K = 3.86)^26^ must be shifted to the right. In the present case, more than 98%
of OD oscillators are found as HDO, with the additional benefit of
providing a convenient raw absorbance value of ∼1 in the νOD stretching vibration maximum. The final salt molalities
of the HDO solutions, for which the IR spectra were collected, are
reported in Table 2. The concentrations and absorption coefficients of the solutions
used in the large-angle X-ray scattering measurements are summarized
in Table 3. The pH
of all solutions except Na2Cr2O7 were
adjusted to pH 12 with sodium hydroxide to minimize the risk of protonation
of the anion. The composition of the vanadate solution at pH 12 is
dominated by the HVO4^2–^ ion, as the HVO4^2–^ ion is a very weak acid (VO4^3−^ is a strong base) with a Ka3 value of 13.27 in dilute aqueous solution,^27^ giving a composition of ca. 5% VO4^3–^ and 95% HVO4^2–^ of the studied solution.
The other anions studied are much weaker bases and are not protonated
at pH 12.
Spectrometer
FTIR spectra were recorded on a PerkinElmer
1720 X spectrometer. The measurement cell was equipped with CaF2 windows separated by a Teflon spacer giving a path length
of 0.0026 cm, as determined interferometrically. For each spectrum,
256 scans were collected with a resolution of 4 cm^–1^ and averaged. The temperature in the instrument chamber was 26.5
± 0.5 °C.
A large-angle θ–θ diffractometer was used to measure the scattering of Mo Kα radiation, λ = 0.7107 Å, from the free surface of the aqueous solutions. Ag Kα radiation, λ = 0.5609 Å, was used for the molybdate solution as Mo Kα radiation will cause substantial fluorescence radiation in this case. The solutions were contained in a Teflon cuvette inside a radiation shield with beryllium windows. After monochromatization of scattered radiation, by means of a focusing LiF crystal, the intensity was measured at 450 discrete points in the range of 1° < θ < 65° (the scattering angle is 2θ). A total of 100 000 counts was accumulated at each angle, and the whole angular range was scanned twice, corresponding to a statistical uncertainty of about 0.3%. The divergence of the primary X-ray beam was limited by 1 or 0.25° slits for different θ regions with overlapping some parts of the data for scaling purposes.
All data treatment was carried
out using the KURVLR program^28^ which has
been described in detail previously.^29^ The
experimental intensities were normalized to a stoichiometric unit
of volume containing one metal atom in the oxometallate anion, using
the scattering factors f for neutral atoms, including
corrections for anomalous dispersion, Δf′
and Δf″,^30^ and values for Compton scattering.^31,32^ For a better
alignment of the intensity function, a Fourier back-transformation
was applied to eliminate spurious (unrelated to any interatomic distances)
peaks below 1.2 Å in the radial distribution function.^33^ Least-squares refinements of the model parameters
were performed by means of the STEPLR program^34^ to minimize the error square sum U = Σ w(s)[iexp(s) – ical(s)]^2^.
The IR spectra were analyzed by the commercial GRAMS/32 v. 4.01 (Galactic Industries Corporation, Salem, USA) and RAZOR (Spectrum Square Associates, Inc., Ithaca, USA) run under GRAMS/32. Spectral data were analyzed using the affected spectra method attributed to Stangret et al.,^35,36^ which is conceptually similar to the one proposed independently by Lindgren and co-workers.^1,37,38^ This numerical procedure leads to the separation of the spectrum of solute-affected water from the bulk water, based on all spectra in the entire solution series and of pure HDO. It has already provided a substantial library of spectral data on ionic solutes, as reviewed previously in detail.^39,40^
The procedure is based
on the main assumption that water in the solution can be divided into
additive contributions of solute-affected (a) and unaffected bulk
(b) εc = εaca + εbcb where εa and εb are the
molar absorption coefficients (dm^3^ mol^–1^ cm^–1^) of affected and bulk water, respectively,
and c symbols the respective concentrations (mol
dm^–3^) and c = ca + cb. The length of sample
cell, l, is constant in the measurements. Rearranging
for εa ,we obtain
We now introduce the so-called “affected number”, N, equal to the number of moles of water affected per mole
solute. Using it and the known solute molality, m (mol kg^–1^), the concentrations can be expressed
the concentrations in eq 1 as c = 1/MV and ca = Nm/V, where M (kg mol^–1^) is the mean molar mass of
water in the sample (H2O + 4% D2O) and V (dm^3^) is the volume of the solution containing m moles of solute and 1 kg of water. Substituting into eq 1 the following equation
is obtained.
Equation 2 allows determination of the affected spectrum for a solution
of finite molality provided that the N parameter
is known. In this work, it is applied to the NaReO4 and
Na3VO4 solutions, for which only single stock
solutions were prepared. In contrast, the advantage of using full
solution series for the other solutes and an approximation of the
experimental spectra ε versus m at each discrete
wavenumber ν~i~ by the least-squares
method can be performed. This makes it possible to extrapolate the
spectrum of solute-affected water to the infinite dilution conditions
(m → 0) and eq 2 takes the form
Equations 2 and 3 still contain the hitherto unknown parameter N. As mentioned above, it is formally equal to the number of moles of water spectrally affected per mole solute. It should not be directly identified with the “hydration number” or “coordination number” (in the sense of the parameters obtained from diffraction experiments or molecular simulations). N is close to the hydration number derived from direct methods only when the solute-affected HDO band differs significantly from the bulk HDO band, in either position or half-width.^36,39^ Otherwise, the affected number is usually lower than the hydration number, and the solute-affected HDO spectrum carries structural and energetic information about the hydration shell in a “condensed” manner.^36^
The proper value of the affected number
can be found either via a band-fitting algorithm^35,36^ or chemometrically.^41,42^ In this work the former method
is used, which briefly consists of generating trial solute-affected
water spectra for given N values and fitting them
using the baseline, analytical bands, and the bulk water spectrum.
A mixed Gaussian–Lorentzian product shape is used for the bands.
All band parameters are kept unconstrained during the fit, with the
exception of the bulk HDO spectrum, for which only intensity is allowed
to vary. The maximum value of N, for which the solute-affected
water spectrum still contains a negligible amount of εb, is considered as the “true” value of N, and the corresponding εa spectrum is considered
the “true” affected water spectrum. The threshold value
for εb contribution is set at ≤0.5% of the
total integrated intensity of the εa spectrum. As
a result of the fitting procedure, both unknowns (i.e., N and εa) are thus obtained simultaneously.
and Discussion
The final affected
spectra in the νOD stretching vibrations range of
HDO are shown in Figure S1. The corresponding
affected numbers can be found in Table 4. As mentioned above, the affected spectra reflect
the state of the solute-affected water, and the N parameter is the number of thus perturbed HDO molecules that are
spectrally differentiable using the DDIR analysis.
It can immediately be seen that the spectra can be grouped into three distinct classes for monovalent, divalent, and trivalent anions, which are progressively red-shifted with increasing formal charge. This observation is in agreement with the previously found dependencies of the band position on the anion’s polarizing power (q/r, where q is the ionic charge and r is the ionic radius).^23,39,40^ Notably, the affected spectra for divalent oxoanions are almost identical across the entire group 6 of the periodic table.
The deconvolutions
of the affected spectra into analytical bands are shown in Figures 1 and 2. The physical significance of the component bands lies in
the representation of the elementary energetic states of water (HDO)
underlying the more complex shape of the full spectrum. Thus, OD oscillators
in different environments can be readily distinguished. Upon careful
examination, these deconvolutions are seen to contain some spurious
peaks (i.e., not corresponding to genuine HDO energetic states) that
function as either baseline correction, atmospheric CO2 absorption correction, or (in the case of Na3VO4) broad components of the extremely red-shifted affected νOH spectrum; this redshift is also apparent in the νOD range. Therefore, they were eliminated from further analysis,
and we proceeded with the genuine affected HDO spectra as seen in Figure 3.
Figure 1 Deconvolution of the solute-affected spectra as shown in Figure S1 into analytical bands (dashed lines) for (a) NaMnO
4, (b) NaReO4, and (c) Na3VO4. Spurious bands not belonging to the νODvibrations of HDO are excluded from further analysis and are indicated with an asterisk.
Figure 2 Deconvolution of the solute-affected spectra as shown in Figure S1 into analytical bands (dashed lines) for (a) Na
2CrO4, (b) Na2MoO4, and (c) Na2WO4.
Figure 3 (a) Solute-affected spectra for the studied sodium salts in the ν
ODstretching vibrations range of HDO after removal of the spurious peaks compared with the bulk HDO spectrum. (b) Separated anion-affected spectra for the studied sodium salts in the νODstretching vibrations range of HDO after removal of the Na^+^-affected spectrum contribution as described in the text compared with the bulk HDO spectrum.
The spectra in Figure 3a contain both anionic and cationic contributions.
However, the influence of Na^+^ on the HDO spectrum has been
previously studied in detail.^37,39,43,44^ In common with other alkali metals,
it shows a slight blueshift of the νOD band to ∼2540
cm^–1^ (i.e., +30 cm^–1^ from the
maximum of the bulk HDO band at ∼2510 cm^–1^). The influence of the cations on the HDO spectrum was extensively
reviewed previously,^39^ and their effect
arises primarily from modifying the solvation environment and structure
of water molecules in the first hydration shell, as investigated theoretically.^45^ Notably, a careful study of the affected HDO
spectrum of NaPF6, containing an extremely structure-breaking
hexafluorophosphate anion, led to total separation of the contributions
from anion and cation to the affected HDO spectrum, thus delivering
an unperturbed Na^+^-affected spectrum.^44^ This was already successfully used in the extraction of
the cationic contribution to εa in the case of NaOH.^46^
Since the affected HDO spectra for the studied oxoanions contain broad unresolved peaks that mix cationic and anionic contributions, especially in the case of divalent anions (cf. Figure 2), it is extremely difficult to ascribe the analytical component bands to genuine anionic influence. Therefore, the same procedure as that in ref (46) was The separated Na^+^-affected HDO spectrum from ref (44) was added as an analytical component to the existing fits, and the salt-affected HDO spectrum was refitted allowing only the intensity of the Na^+^ band to vary. Finally, the converged Na^+^ contribution was subtracted, thus providing the anion-affected HDO spectra as shown in Figure 3b. The most important parameters of these spectra are summarized in Table 2. As expected, the effect of subtraction of the Na^+^-affected HDO spectrum depends on the overall position of the salt-affected spectrum. For the monovalent anions, for which the spectra in Figure 3a are blue-shifted with respect to that of bulk HDO, the removal of the Na^+^ band reveals more strongly the structure-breaking character of the anion. In contrast, for the di- and trivalent anions, which are seen to red-shift the OD band position, the removal of this band induces further redshift. In any case, the elimination of the cationic influence is beneficial to a deeper understanding of the anionic part of the affected HDO spectrum.
It is informative
to compare the obtained results with available data for other oxoanions.
In the case of the monovalent ones, perchlorate seems to be a good
reference ion. It is well-known that the ClO4^–^ anion is one of the more structure-breaking ones in aqueous solution.
The maximum position of the νOD band (ν°)
is ∼2630 cm^–1^.^15,35,37−39,43^ Both permanganate and perrhenate are characterized with similar
ν° values. Divalent and trivalent anions have been hitherto
studied to a much smaller extent. However, the available limited data
set corresponds well with the present results. In particular, ν°
= 2477 and 2480 cm^–1^ for SO4^2–^ and SeO4^2–^, respectively,^18,20^ in perfect agreement with our results for the group 6 oxoanions.
Similarly, ν° = 2434 cm^–1^ for PO4^3–^,^22^ the only
trivalent anion studied to date, which is strikingly close to our
ν° = 2432 cm^–1^ for VO4^3–^ (see Table 4), and the affected HDO spectrum is likewise broad and contains
underlying features. The full width at half-height of the anion-affected
HDO spectrum increases with the increasing red shift of the νOD band. This dependence is linear to a good approximation
(Δ1/2 = −0.725ν° + 1972, R^2^ = 0.91) and was already detected for a series
of simple aqueous anions.^16^ Notably, the
results of the previous investigation for some particular anions are
in excellent agreement with the current data, viz., Δ1/2 = ∼75 cm^–1^ for ClO4^–^, ∼170 cm^–1^ for SO4^2–^, and 180 cm^–1^ and for SeO4^2–^.^18,20^
The vanadate anion clearly stands
out in the sense that the affected HDO spectrum contains underlying
peaks instead of forming a single broad band resembling the bulk HDO
spectrum shape. The peak positions at maximum for the three analytical
bands are 2324, 2432, and 2638 cm^–1^, in good agreement
with the band parameters previously found for PO4^3–^.^22^ On the basis of the cited work, the
2324 cm^–1^ band is ascribed to strong and short hydrogen
bonds from water to the phosphate anion, while the 2432 cm^–1^ component corresponds to the hydrating waters simultaneously donating
two hydrogen bonds to two oxygen atoms in VO4^3–^. The most blue-shifted component band merits further investigation.
At first glance, it is hard to expect severely weakened hydrogen bonds
in the vicinity of a strongly hydrated ion. However, the long known
concept of anticooperativity of the two hydrogen bonds donated by
the same water molecule provides an explanation.^47^ Namely, whenever a water molecule donates to a strong hydrogen
bond, the other hydrogen is engaged in a weak hydrogen bond. This
particular phenomenon has recently computationally been confirmed
for PO4^3–^, for which the short and strong
water–phosphate hydrogen bonds are accompanied by weak hydrogen
bonds donated by the same hydrating water molecules to the second
hydration shell.^48^ While this explains
the appearance of the strongly blue-shifted component band in the
anion-affected HDO spectrum of VO4^3–^,
we should mention another possible phenomenon, namely, the influence
of the cation on the outermost hydration sphere of the anion detectable
in the affected spectra via the formation of solvent-separated ion
pairs. This was extensively discussed in the other direction (i.e.,
the influence of the anion on the cation’s outermost hydration
sphere),^44^ but the same mechanism is expected
to arise here. We note that the two explanations are not mutually
exclusive.
A more direct comparison of HDO spectra with structural
data is made possible by transforming the molar absorptivity band
contour εa(ν) to the probability distribution
of the intermolecular oxygen–oxygen distance P(ROO). It is based on the generalization
of the empirical relationship noted previously between νOD/νOH band positions of HDO in solid hydrates
and the respective intermolecular distances determined by diffraction
methods.^49^ The details of the ε(νOD) → P(ROO) transformation are published elsewhere.^36,37^ In brief, it is based on the relation between ROO and νOD found in solid hydrates,^49^ which is of the form ROO = [16.01 – ln(2727 – νOD)]/3.73.
The band contour is then transformed to the probability distribution
by inverting this relation and taking the derivative according to
where the normalization constant C is chosen such that the probability distribution is properly normalized to unity.
The obtained probability distributions based on the anion-affected HDO spectra are shown in Figure 4 and are compared with bulk water. We note that intermolecular oxygen–oxygen distance distributions for oxoanions contain contributions stemming both from anion–water and water–water (in the hydration shell) hydrogen bonds. On the basis of the data in Figure 4, we can reiterate the strongly structure breaking character of the studied monovalent anions, contrasting of the mildly structure-making divalent anions of group 6 and the more-structure-making vanadate.
Figure 4 Intermolecular oxygen–oxygen distance distributions derived from the anion-affected HDO spectra shown in Figure 3 compared to the respective distribution in bulk HDO. The curves for the three divalent anions are almost superimposed.
The LAXS data show that the hydrated oxometallate ions bind one
shell of water molecules to the oxygen atoms. The M–O and M–(O)···Oaq distances have been determined very accurately as they are
not superimposed with other distances. However, for di- and trivalent
oxometallate ions, the (M−)O···Oaq distance is too close to the bulk Oaq···Oaq distance to be separated from each other, while for the
permanganate and perrhenate ions it is sufficiently longer to be separated
from of the bulk Oaq···Oaq contribution.
The mean Oaq···Oaq distance in
LAXS studies of fairly concentrated (ca. 1 mol·dm^–3^) aqueous solution is ca. 2.89(2) Å,^2,50^ while
it is slightly shorter in pure water, ca. 2.85 Å.^51^ For the divalent anions it was not possible
to separate the O···Oaq and Oaq···Oaq distances, and they were refined
as a mean value slightly shorter than 2.89 Å. This indicates
an O···Oaq distance of ca. 2.85 Å that
has been used in the calculations of the M–O–O bond
angles. All oxoanions in this study except the hydrogenvanadate hydrogen
bind on average two water molecules to each oxoanion oxygen, while
hydrogenvanadate ion binds three water molecules per vanadate oxygen.
This is in line with previous observations that oxoanions with a central
atom belonging to series 4 or higher only hydrogen bind two water
molecules to each oxoanion oxygen. The hydrogenvanadate ion is the
exception binding three water molecules per oxygen. The refined structure
parameters of the hydrated permanganate, perrhenate, chromate, dichromate,
molybdate, tungstate, and hydrogenvanadate ions in aqueous solution
as determined by LAXS are summarized in Table 5, and the fitting of the raw data and radial
distribution function (RDF) of the sodium perrhenate solution is shown
in Figure 5. The fittings
of the remaining solutions are shown in Figures S2–S7.
Figure 5 Top LAXS radial distribution curves for a 1.502 mol·dm^–3^ aqueous solution of sodium perrhenate. (Upper part) Separate model contributions (offset: 14) of the hydrated perrhenate ion (green line), the hydrated sodium ion (brown line) and aqueous bulk (light blue line). (Middle part) Experimental RDF: D(r) – 4πr^2^ρ
o(black line); sum of model contributions (red line); difference (blue line). Bottom Reduced LAXS intensity functions *s·*i(s) (black line); model s·icalc(s) (red line).
Due to the hydration, the M–O bond distances become almost 0.02 Å longer in aqueous solution than those in solid compounds where the oxometallate anions are not binding any metal ion other than alkali metal ions or organic cations or molecules; see Tables 6 and S1. We note that the charge neutralization of the oxygens of the oxometallate anions through the hydrogen bonding of the hydrating water molecules seems to be independent of the charge of the anion even though the hydrogen bond strength is different.
Anions are hydrated by accepting hydrogen bonds from the hydrating water molecules. Monovalent anions, except fluoride and hydroxide, form hydrogen bonds to water which are weaker and longer than the hydrogens bonds in bulk water (i.e., they act as structure breakers), while the fluoride, hydroxide, and di- and trivalent anions form hydrogen bonds stronger than those in bulk water (i.e., they act as structure makers). The anions form one hydration shell distinguishable from bulk water, as also determined by DDIR spectroscopy and LAXS. The hydration of oxoanions results in X–O bonds ca. 0.02 Å longer than those of unsolvated ions in the solid state not involved in strong bonding to counterions as metal ions except the alkali metal ions (Tables 6 and S1). The oxygens of oxoanions with a central atom from the second and third series in the periodic table and hydrogenvanadate hydrogen bind three hydrating water molecules, while oxygens of oxoanions with a heavier central atom only hydrogen bind two water molecules (Tables 1 and 5).
The support from the Swedish Research Council is gratefully acknowledged. We acknowledge funding from the statutory fund of the Chemical Faculty, Gdańsk University of Technology. The support to M.S. from the Visby Programme through the Swedish Institute for his stay in Uppsala is gratefully acknowledged.
The Supporting Information is available free of charge at https://pubs.acs.org/doi/10.1021/acs.inorgchem.0c00594.
The authors declare no competing financial interest.