Donor–Acceptor Complexes between Ammonia and Sulfur Trioxide

Article pubs.acs.org/JPCA

Donor−Acceptor Complexes between Ammonia and Sulfur Trioxide: An FTIR and Computational Study Karolina Haupa, Andrzej Bil, and Zofia Mielke* Faculty of Chemistry, University of Wrocław, Joliot Curie 14, 50-383 Wroclaw, Poland S Supporting Information *

ABSTRACT: The complexes of ammonia with sulfur trioxide have been studied using FTIR matrix isolation spectroscopy and DFT/B3LYP calculations with the aug-cc-pVTZ basis set. The NH3/SO3/Ar matrixes were prepared in two different ways. In one set of experiments the matrix was prepared by the simultaneous deposition of the NH3/Ar mixture and SO3 vapor from the thermal decomposition of K2S2O7. In the second set of experiments thermolysis products of sulfamic acid were trapped in an argon matrix. Both methods of matrix preparation led to the formation of the H3N·SO3 electron donor−acceptor complex that was characterized earlier. In the matrixes comprising thermolysis products of sulfamic acid, in addition to H3N·SO3, the H3N−SO3···NH3 complex (IID) was also identified. The identity of the complex was confirmed by comparison of the experimental and theoretical spectra of H3N−SO3···NH3 and D3N−SO3···ND3. The performed calculations show that in H3N−SO3···NH3 the two N atoms and the S atom are collinear; the two S−N bonds are nonequivalent, one is much shorter (2.230 Å) than the other one (2.852 Å). In the AIM topological analysis, the interaction energy decomposition and topological properties of the electron localizability indicator (ELI-D) allowed us to categorize the stronger N−S bond in the IID complex as a dative bond and to assume that the fragile N−S bond is a consequence of a weak electron-donor−acceptor interaction. The calculations indicate that the identified IID complex corresponds to a local minimum on the PES of the NH3/SO3 system of 2:1 stoichiometry. The (NH3)2SO3 complex, IIHB, corresponding to a global minimum is 7.95 kcal mol−1 more stable than the IID complex. The reason that the IID complex is present in the matrix but not the IIHB complex is discussed.

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INTRODUCTION Sulfamic acid has been extensively studied for many years due to its interesting structural properties and wide application in industry.1−3 It is well established that it exists in all phases as a donor−acceptor adduct (EDA) in the zwitterionic form, +H3N· SO3−.4−10 The adduct was characterized in the gas phase using microwave spectroscopy5,6 and in a low-temperature nitrogen matrix with the aid of infrared spectroscopy.7,8 The crystal structure of sulfamic acid was determined by both X-ray and neutron diffraction techniques.9,10 The14N−1H dipolar tensors of sulfamic acid have been determined from a solid state singlecrystal study.11 The performed microwave studies confirmed the C3v symmetry of the adduct and demonstrated the sensitivity of its structure to the environment; in particular, the N−S distance was found to be 0.186 Å longer in the gas phase than in the crystalline environment. The dipole moment of this complex also shows a significant enhancement in the solid state (9.6 D)12 relative to the gas phase (6.2 D).6 The structural properties of sulfamic acid have also been extensively studied by quantum chemistry methods;13−27 the H3N·SO3 adduct has drawn particular attention as a prototype EDA complex. Ab initio quantum mechanical methods have been employed in an attempt to explain the influence of the environment on the structural and electronic properties of this complex.19,26,27 In early work the self-consistent reactant field (SCRF) method at the Hartree−Fock level of theory was © XXXX American Chemical Society

applied to model the structure of the H3N·SO3 adduct in the condensed phase.19 The calculated charge transfer was found to increase from 0.28 to 0.36 e from the gas phase to the condensed phase. The gas-phase and SCRF charge-constrained calculations at the HF/6-311++G(d,p) level performed for the H3N·SO3 complex showed that solvation leads to the strengthening of both the charge transfer and the covalent components of the S−N bond in this complex.20 The density functional method along with the polarizable continuum model (PCM) for solvation was also employed to explain the origin of the decrease in S−N distance and an increase in the binding energy, dipole moment, and charge transfer from the gas phase to the condensed phase.26 In addition to its interesting structural properties the H3N· SO3 complex has been postulated to be a possible nucleation site for the formation of atmospheric aerosols.28 The possible complexes derived from the NH3/SO3/H2O system have been characterized using DFT and ab initio methods.29−31 It has been shown that the H3N·SO3 complex was progressively stabilized by the addition of subsequent water molecules which formed H3N·SO3···(H2O)n (n = 6, 9, 12) complexes.31 The N− S bond distance decreased monotonically with the addition of Received: August 14, 2015 Revised: October 6, 2015

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entrapping its decomposition products in solid argon led to matrixes of composition SO3/ND3/Ar. The gas mixtures were condensed onto a gold-plated copper block kept at 17 K. The temperature was maintained by means of a closed -cycle helium refrigerator (Air Products, Displex 202). Infrared spectra were recorded at 11 K by means of a Bruker 113v spectrometer employing a liquid-nitrogen-cooled MCT detector (4000−550 cm−1) and coadding 256−512 scans at a resolution of 0.5 cm−1. The spectra were recorded after deposition and after annealing. The matrixes were warmed to 30, 35, and 39 K for ca. 10 min and cooled to 11 K. K2S2O7 (>99%, Sigma) and sulfamic acid (99.999%, Aldrich) were commercially available. Gaseous ammonia was obtained by dropping a 50% water solution of NaOH into a bulb containing solid NH4Cl; the liberated gaseous ammonia was dried by passing through the tube filled with CaCl2. The deuterated aminosulfonic acid, D3N·SO3, was obtained by triple recrystallization of H3N·SO3 from D2O and was dried above phosphorus pentoxide, P4O10. Computational Details. The Gaussian 09 program39 was used for geometry optimization and wavenumber calculations. The structures of the NH3 and SO3 monomers and the ammonia/sulfur trioxide complexes of 1:1, 1:2, and 2:1 stoichiometry were fully optimized, and their corresponding energies were calculated using the DFT/B3LYP/aug-cc-pVTZ method. It has been demonstrated earlier40,41 that the B3LYP method coupled with the aug-cc-pVTZ basis set is generally sufficient for the direct calculations of many molecular properties of small inorganic molecules including sulfur oxides. The structures of the H3N·SO3 and H3N·SO3·NH3 complexes were also optimized by the B3LYP-D342 and MP2 methods with the aug-cc-pVDZ basis set. All optimized structures discussed in the main body of the paper correspond to real minima on the potential energy surface as evidenced by the positive values of the calculated harmonic wavenumbers. The stability of a wave function was confirmed for all complexes. Additional information on the nature of interactions in the complexes formed was obtained from the topological analysis of the electron density. In Bader’s theory of atoms in molecules (AIM)43 the classification of the interactions is based on the analysis of the numerical values of the electron density (ρ), the Laplacian of the electron density (∇2ρ), the total energy density (H), and the ellipticity (ε) at the (3, −1) bond critical point (BCP), which is a saddle point on the electron density (i.e., a minimum with respect to the interatomic path and a maximum along the perpendicular directions). The bond can be additionally characterized by the value of the delocalization index, which can be interpreted as a bond order.44,45 A negative value of ∇2ρ at the bond critical point characterizes a bond as shared (the electron density is locally concentrated), whereas a positive value defines a closed-shell type of interaction (the electron density is locally depleted). From a chemical point of view the former case, together with a large value of electron density, is typical of a covalent bond, while the latter is characteristic of ionic or hydrogen bonds, where the bonding is driven by electrostatic interactions. The topological analysis of the electron density in the complexes identified in the studied matrixes has been performed using the QTAIM method and the AIMALL package.46 Another topological tool applied in this paper, the electron localizability indicator (ELI-D),47,48 uses many constructs similar to those defined by Bader but yields a chemical

each H2O molecule. The studies strongly suggested that H3N· SO3 may act as an effective nucleation agent for the formation of atmospheric aerosols and cloud particles. The interesting structural properties of sulfamic acid as well as its possible role in aerosols formation inspired our recent studies on sulfamic acid aggregation by the help of the ESI MS method and DFT calculations.32 Large singly and doubly charged clusters of sulfamic acid were identified in the negative ion ESI mass spectrum of sulfamic acid. In an attempt to obtain spectroscopic characteristics of the small sulfamic acid aggregates (n = 2, 3) we entrapped the vapor above the sulfamic acid in solid argon and measured its infrared spectra. Much to our surprise, we discovered that, in addition to the H3N·SO3 adduct, the matrix comprised the H3N−SO3···NH3 complexes in which two separate ammonia molecules were attached to sulfur trioxide. Electron donor−acceptor (EDA) complexes are of importance in several phenomena in biology and chemistry and have been the subject of many studies. Among the best characterized EDA complexes are those between SO3 or SO2 and ammonia or amines.5,7,33−35 There is thermodynamic and spectroscopic evidence that sulfur dioxide forms with ammonia an adduct of 2:1 stoichiometry, SO2· (NH3)2 in addition to SO2·NH3.35−37 To our knowledge, the SO2·(NH3)2 adduct is the only known EDA complex in which one sulfur atom serves as an electron acceptor for two electrondonor molecules. In this paper we present the results of the infrared matrix isolation studies and quantum chemistry calculations that led to the identification of the SO3·NH3 and H3N−SO3···NH3 adducts and to the elucidation of their structure and bonding situation.

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EXPERIMENTAL SECTION Infrared Matrix Isolation Studies. The SO3/NH3/Ar matrixes were prepared in two different ways. In the first set of experiments the SO3 vapor was deposited simultaneously with the argon matrix gas or NH3/Ar mixture in such a way that SO3 was mixed with argon or a gaseous mixture inside the cryostat. The concentration of NH3/Ar varied in the range of 1:100− 1:1000. The source of SO3 was solid K2S2O7 placed in a small quartz heater whose nozzle exit was facing the deposition mirror of the cryostat. Prior to deposition, a few short degassing cycles were performed by heating the solid K2S2O7 sample to about 140 °C, slightly below the decomposition temperature of K2S2O7. For deposition, the temperature of the heater was raised to 170−180 °C. The concentration of the SO3/Ar matrix varied in the range of 1:100−1:1000 as estimated from the relative intensities of the bands due to (SO3)n aggregates that were reported earlier in the literature.38 In the second set of experiments the source of SO3 and NH3 was the thermal decomposition of sulfamic acid. In these experiments, solid crystalline sulfamic acid, H3N·SO3, was placed in a quartz heater mounted inside the vacuum chamber of the cryostat. Prior to deposition the solid sample was heated to ca. 130 °C for degassing and water removal. Small amounts of water in the solid sulfamic acid sample led to the presence of sulfuric acid in the matrixes. During deposition the solid H3N· SO3 was heated to temperatures in the range of 160−180 °C; its thermolysis products were diluted with the argon gas in the vacuum chamber. The concentration of decomposition products, NH3 and SO3, in the argon matrix was controlled by the temperature of the solid sample and by varying the flow of argon at a rate from ca. 5 to 15 mmol h−1. Some experiments were performed with deuterated sulfamic acid D3N·SO3; B

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The Journal of Physical Chemistry A interpretation based on core, bonding, and nonbonding electron pairs.49 This method allows us to split the molecular space into core and valence attractor basins, from which we can calculate integral properties (such as an average electron population), and is particularly useful for covalently bonded systems. A topological analysis of ELI-D(r) has been performed using the DGrid program.50 The analysis of the bonding situation in the discussed complexes has been supported with the description given by natural bond order (NBO) analysis51 as well as interaction energy decomposition proposed by Su and Li52 and implemented in Gamess.53 Within the framework of this partition scheme, the total interaction energy of a dimer or a trimer is expressed as a sum of the electrostatic, repulsion, polarization, exchange, and so-called dispersion contributions, while the last two terms are defined using the changes in the exchange and correlation functionals on going from monomers to a supermolecule.52

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RESULTS AND DISCUSSION Infrared Spectra. In Figure 1 the spectra of the SO3/NH3/ Ar matrixes obtained by deposition of the SO3 + NH3/Ar mixtures (trace c) are compared with those obtained by entrapping in solid argon the sulfamic acid decomposition products (traces d−f). The spectra of the NH3/Ar and SO3/Ar matrixes are also shown (traces a and b). A comparison of the studied spectra with those reported in the literature indicates that the matrixes obtained by the two methods contain SO3 and NH3 monomers, sulfur trioxide, and ammonia dimers and trimers38,54,55 as well as traces of the ammonia−water complexes56 (due to a small amount of water contamination in the studied matrixes). The relative concentration of the monomers, dimers, and trimers was dependent on the conditions of the experiment (temperature of the sample, argon flow). No decomposition products of ammonia or sulfur trioxide (in particular, sulfur dioxide) were identified in the studied matrixes. In all recorded spectra of the SO3/NH3/Ar matrixes, a group of seven bands appeared, I, that was observed in the spectra of neither the SO3/Ar nor NH3/Ar matrixes. In the spectra of matrixes containing sulfamic acid thermolysis products, in addition to group I, another group of bands, II, consisting of five absorptions was observed. The relative intensities of the bands belonging to groups I and II were constant, within experimental error, in all recorded spectra. Two bands due to group II that appeared in the NH stretching region were accompanied by additional absorptions, IIC. The wavenumbers of the I, II, and IIC band sets are collected in Tables 1 and 2. The relative intensities of bands I with respect to those of the II and IIC bands were dependent on the temperature of the crystalline sulfamic acid sample during deposition and on matrix annealing. The lower temperature of the sample during deposition favored the formation of the complexes characterized by II and IIC absorptions whereas its temperature increase favored the formation of complex I (compare traces d and f in Figure 1). Matrix annealing led to a strong increase in bands I at the cost of bands II and IIC (traces d and e). In the spectra recorded a short time after deposition (ca. 5−15 min) only weak bands II appeared, and prolonged matrix deposition and/or temperature growth of the sample led to the appearance of bands I (Figure S1, Supporting Information). Absorptions II and IIC exhibited the same behavior with respect to the temperature of the solid sample during deposition and to matrix annealing; however, the mutual

Figure 1. Spectra of the argon matrixes obtained by the deposition of SO3 vapor (above solid K2S2O7) diluted with an NH3/Ar = 1/300 mixture (c) or by entrapping thermolysis products of sulfamic acid (d−f). Spectra d, f correspond to matrixes obtained at different deposition temperatures of sulfamic acid (160, 165 °C). Spectrum e was recorded for matrix d after its annealing to 35 K for 10 min. For comparison, the spectra of SO3/Ar and NH3/Ar = 1/300 matrixes are also shown. Labels * and + indicate the bands due to NH3−H2O and H2SO4.

intensities of II and IIC varied in various experiments, indicating that they do not belong to the same species. The infrared spectra of the thermolysis products of D3N(H3N)·SO3 entrapped in solid argon indicated that the deuteration ratio exceeded 90%. The spectra of the matrixes doped with thermolysis products of the deuterated sample and the spectra of the annealed matrixes allowed us to identify the deuterium counterparts of bands I and II observed in the C

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Table 1. Comparison of the DFT/aug-cc-pVTZ Calculated Wavenumbers with the Experimental Ones for the H3N·SO3 and D3N·SO3 Complexesa H3N·SO3 ν νcalc νobs νcalc νobs νcalc νobs νcalc νobs νcalc νobs νcalc νobs νcalc νobs

D3N·SO3 Int

3587 3388.7 3455 3281.6 1636 1589.4 1337 1359.7 1204 1246.0 1037 1064.4 722 800.0

[7]

112 27 58 1565 524 1354 160 1317 16 1072 15

ν νcalc νobs νcalc νobs νcalc νobs νcalc νobs νcalc νobs νcalc νobs νcalc νobs

2646 2531.4 2464 2375.2 1187 1139.1 1335.5 1357.2 919 961.6 1041 1068.3 589 651.8

Int

Δν

approx assignment

62

−941 −857.3 −991 −906.4 −449 −450.3 −1.5 −2.5 −285 −284.4 +4 +3.9 −133 −148.2

νasNH3/νasND3

31 22 516 98 44 15

νsNH3/νsND3 δasNH3/δasND3 νas SO3 δsNH3/δsND3 νs SO3 ρNH3/ρND3

The literature data for the H3N·SO3 complex, the calculated and observed wavenumbers shifts after deuteration (Δν = νH − νD) and the calculated intensities (Int, km mol−1) for the particular modes are also given.

a

Table 2. Comparison of the DFT/aug-cc-pVTZ Calculated Wavenumbers with the Experimental Ones for the H3N·SO3·NH3 and D3N·SO3·ND3 Complexesa H3N·SO3·NH3 ν II νcalc νobs νcalc νcalc νcalc νobs νcalc νobs νcalc νobs νcalc νobs νcalc IIC νobs νobs

3603 3455.2 3599 3474 3469 3360.2 1343 1381.4 1132 1181.4 1032 1041.2 1022

D3N·SO3·ND3 Int

ν

73

52

νcalc νobs νcalc νcalc νcalc νobs νcalc νobs νcalc νobs νcalc

119

νcalc

18 1 13 482 161

Int

Δν

2657 2582.0 2650 2479 2473 2460.1 1342 1374.0 866 885.3 783

42

νas(NH3)s/νas(ND3)s

107

−946 −873.2 −949 −995 −996 −900.1 −1 −7.4 −266 −296.1 −249

1032

15

9

δs(NH3)w + νs(SO3)+/νs(SO3)

13 0 12 470 109

approx assignmentb

νas(NH3)w/νas(ND3)w νs(NH3)w/νs(ND3)w νs(NH3)s/νs(ND3)s νas(SO3) δs(NH3)s/δs(ND3)s δs(NH3)w + νs(SO3)/δs(ND3)w

νas (NH3)s (complexed) νas (NH3)w (complexed)

3479.2 3379.1

a The calculated and observed wavenumbers shifts after deuteration (Δν = νD − νH) and the calculated intensities (Int, km mol−1) for the particular modes are also given. bSuperscripts s and w indicate strongly and weakly bonded ammonia molecules.

(2d), and 6-311G(d)) resulted in S−N distances in the range from 1.954 to 1.912 Å.19 The experimental S−N distance and OSN angle determined from the microwave spectra are equal to 1.957 Å and 97.6°. The best agreement between the experimental N−S distance and that from HF calculations was obtained for small basis sets, most probably due to error cancellation. Bands I, appearing in the spectra of the matrixes obtained by both methods, are assigned with confidence to the H3N·SO3 complex on the basis of the following arguments. First, the observed wavenumbers of bands I are in accord with the calculated ones for this complex (Table 1). Second, the observed wavenumber shifts of the corresponding absorptions in the spectra of the H3N·SO3 and D3N·SO3 experiments also show good agreement with the calculated ones as presented in Table 1. The observed isotopic shift ratios, νasNH3/νasND3 and

spectra of the SO3/NH3/Ar matrixes. In Tables 1 and 2, the identified wavenumbers of bands I and II in the experiments with hydrogenated and deuterated sulfamic acid are compared. In Figure 2, the spectra of the matrixes doped with the D3N· SO3 thermolysis products are presented in the region of the ND3 and SO3 stretching vibrations. H3N·SO3 Complex. The DFT/B3LYP/aug-cc-pVTZ-optimized structure of the H3N·SO3 adduct in the gas phase is presented in Figure 3. The complex has a C3v staggered conformation as reported earlier.5 The calculated S−N distance and OSN angle are equal to 2.088 Å and 96.8°, respectively. The DFT/B3LYP/6-31+G(d) and 6-311+G(2df,p) optimizations on the isolated adduct, reported earlier, gave bond lengths of 2.100 and 2.074 Å.26,27 For comparison, the ab initio calculations at the Hartree−Fock (HF) level with various basis sets (6-31G(d), 6-31G(d,p), 6-31+G(d), 6-31G(2d), 6-31+GD

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the δsNH3 wavenumbers (1246.0, 1317 cm−1 in the spectra of the argon and nitrogen matrixes, respectively). This discrepancy is due to the known strong effect of the nitrogen matrix on this vibration. In addition, we identified the absorptions of the complex in the region of the NH3 stretching vibrations (νsNH3 3281.6 cm−1, νasNH3 3388.7 cm−1) which were not reported earlier. H3N−SO3···NH3 Complex: Structure and Infrared Spectra. The presence of ammonia, sulfur trioxide monomers, and their aggregates as well as the formation of the H3N·SO3 complex in the deposited matrixes led us to the conclusion that the set of bands II corresponds to higher-order 1:2 or 2:1 complexes between ammonia and sulfur trioxide. Two structures were optimized, using the DFT/B3LYP/augcc-pVTZ method, for the ammonia/sulfur trioxide complex of 2:1 stoichiometry: IIHB and IID. The structures are presented in Figure 4. (See also Figure S2 and Table S1 in the Supporting Information.)

Figure 2. Spectrum of the argon matrix, obtained by entrapping thermolysis products of deuterated sulfamic acid, D3N·SO3, measured directly after deposition (a) and after matrix annealing to 39 K for 15 min (b).

Figure 3. DFT/B3LYP/aug-cc-pVTZ-optimized structure and molecular graph for the H3N·SO3 complex (a). The smallest sphere represents bond critical points (BCPs). The geometrical parameters (larger font) and AIM atomic charges for all atoms (smaller font) are shown. Localization domains of the electron localizability indicator plotted for ELI-D(r) = 1.15. (b) The basin representing the lone electron pair on the N atom, V(N), is marked in green.

Figure 4. DFT/B3LYP/aug-cc-pVTZ-optimized structures of the (H3N)2SO3 and H3N·SO3·NH3 complexes.

In the IIHB complex one of the NH groups of the H3N·SO3 adduct is attached to the nitrogen atom of the second ammonia molecule, forming a relatively strong N−H···N bond. The interaction of the adduct with NH3 leads to the shortening of the N−S distance from 2.088 to 1.975 Å. In the other structure, IID, the second NH3 molecule is attached to the S atom of H3N·SO3 opposite to the first ammonia molecule. The three heavy atoms are collinear, and the complex is stabilized by the double interaction of the sulfur atom with each of the two nitrogen atoms. One of the two S−N distances (2.230 Å) is shorter than the other one (2.852 Å), which indicates that the

νsNH3/νsND3, are equal to 3388.7/2531.4 = 1.41 and 3281.6/ 2375.2 = 1.38, respectively. In turn, the isotopic ratios for δsNH3/δsND3 are equal to 1589.4/1139.1 = 1.39 and 1246.0/ 961.6 = 1.30. And third, bands I identified at 1589.4, 1359.7, 1246.0, and 1064.4 cm−1 have values very close to the corresponding bands of the H3N·SO3 complex isolated in the nitrogen matrix (Table 1).7 The largest difference appears for E

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from the identification of two bands corresponding to the δs(NH3) motion of the two ammonia molecules. The calculations result in 1132 and 1032 cm−1 wavenumbers for the more strongly and more weakly bonded ammonia molecules. The corresponding bands are identified at 1181.4 and 1041.2 cm−1, being ca. 207 and 67 cm−1 shifted toward higher wavenumbers from the corresponding band of the NH3 monomer (974 cm −1 ). In the H 3 N·SO 3 complex the corresponding vibration occurs at 1246.0 cm−1 and is more perturbed than that in the IID complex. This is due to the shorter S−N distance (2.088 Å) and stronger EDA interaction in the 1:1 complex than in the IID complex. The stretching vibrations of ammonia molecules are less sensitive to interaction with SO3 than the bending vibrations; the calculations predict very similar νs(NH3) and νas(NH3) wavenumbers for the two bonded ammonia molecules in the IID complex (Table 2). The calculated νs(NH3) and νas(NH3) wavenumbers for complex IID are larger than the corresponding wavenumbers for complex I, which is in accord with the experimental data (compare Tables 1 and 2). The bands IIC at 3379.1 and 3479.2 cm−1 occur in the vicinity of bands II at 3360.2, 3455.2 cm−1, respectively. It is hard to identify the product they belong to on the basis of only two absorptions. The conformer of the H3N·SO3·NH3 complex other than IID cannot be completely excluded; however, the calculated wavenumbers for IID (true minimum) and for the conformer corresponding to the lowest in energy scale TS (in which the strongly and weakly bonded NH3 molecules are in eclipsed and staggered conformations relative to SO 3, respectively) are very similar. The largest difference for the NH stretches equals 5.4 cm−1, which is in contradiction to the identified bands. Alternatively, the two absorptions may correspond to the H3N·SO3·NH3 complex which is perturbed by the matrix environment or the other molecule present in the matrix cage. We have also optimized the structure of the ammonia/sulfur trioxide complex of 1:2 stoichiometry. Its structure is shown in Figure S3 in the Supporting Information; the SO3 molecule is attached to the H3N−SO3 complex in such a way that one of the oxygen atoms interacts with one of the NH groups of ammonia and the orientation of the two SO3 molecules is similar to that in the (SO3)2 dimer of C1 symmetry.38 The calculated wavenumbers are collected in Table S3. The obtained spectra prove that this complex is not present in the studied matrixes in a concentration large enough for its detection. The calculations predict six bands due to the SO3 stretching vibrations in the 1400−1250 cm−1 region (four very intense bands and two less intense ones) while only one absorption belonging to group II was identified in this region. Moreover, the identified II wavenumbers do not agree with the calculated ones for the NH3(SO3)2 complex. Bonding Situation in H3N·SO3 and H3N−SO3···NH3. In Figures 3 and 5 the atomic charges obtained by the AIM method are shown for the H3N·SO3 and H3N−SO3···NH3 complexes, respectively. The extent of charge transfer is conveyed by the net charge on the SO3 subunit as determined from the sum of the atomic charges. The AIM charges computed from the electron density obtained at the DFT/B3LYP/aug-cc-pVTZ level reveal that the amount of electron transfer from the NH3 group to the SO3 moiety in the 1:1 complex is equal to 0.28 e. For the H3N− SO3···NH3 complex the obtained data indicate a transfer of 0.20 e from the more strongly bonded and 0.05 e from the more

two ammonia molecules are unevenly bonded, one much more strongly than the other one. The strongly bonded NH3 molecule is in the staggered conformation, and weakly bonded NH3 is in an eclipsed conformation with respect to SO3. Structure IIHB is a global minimum on PES and is 9.92 kcal mol−1 lower in total energy than the IID structure (7.95 kcal mol−1 more stable in terms of binding energy). In addition to IID, the calculations showed four stationary points on PES sharing an N−S−N structural motif. The energy difference between IID and other structures is in a narrow range up to 0.3 kcal mol−1, while no other true minimum apart from IID was found. (See the details in Supporting Information, Figure S2 and Table S1.) Interestingly, the structure with D3h symmetry (N−S bonds are equivalent), which is a transition structure 0.18 kcal mol−1 higher in energy than IID at the B3LYP level, has turned out to be a true minimum at the B3LYP-D3 level, where semiempirical dispersion corrections are included. Such a structure is, however, inconsistent with the experimental spectrum discussed below and is also disproved by MP2 calculations which predict it to be a transition structure (Table S4, Supporting Information), as B3LYP does. Although at the B3LYP level we found only one minimum structure adopting the N−S−N bond pattern, due to the fact that the loosely bonded NH3 molecule leads to a set of transition structures only slightly different in total energy from IID, the accuracy of the computational methods does not allow us to claim unambiguously which of the possible conformations is adopted by the complex present in the matrix under the experimental conditions. The calculated harmonic wavenumbers for the IID hydrogenated complex, H3N−SO3···NH3, and its deuterated analogue, D3N−SO3···ND3, are compared with the experimental values in Table 2. The full sets of wavenumbers of the IIHB and IID structures are presented in Tables S2 and S3 in the Supporting Information. The comparison of the identified wavenumbers belonging to group II with the calculated ones for the IIHB structure allowed us to exclude the presence of this complex in the studied matrixes. For IIHB, the calculations predict a very strong band, due to the N−H stretching vibration of the N−H···N bond, in the 2950−2700 cm−1 region (the calculated harmonic wavenumber is equal to 2965.8 cm−1, I = 1000 km mol−1). Such an absorption was not observed in this region in the studied spectra. The other calculated wavenumbers for complex IIHB also do not agree with the bands II wavenumbers (see Table S3). Another strong argument against the IIHB complex is a strong decrease of bands II after matrix annealing which evidences decomposition of the complex to which the bands are due. As one can see in Table 2, the wavenumbers of the bands II, identified in the H and D experiments, agree reasonably well with those calculated for the hydrogenated and deuterated complexes of the IID structure. The deviations of the calculated wavenumbers from the observed ones are similar to those occurring for the 1:1 H3N·SO3 complex. Three counterparts of the ν(NH3) vibrations and one counterpart of the ν(SO3) one were identified in the spectra of matrixes doped with thermolysis products of deuterated sample. The isotopic shift ratios for the ammonia vibrations are equal to νas(NH3)s/ νas(ND3)s = 3455.2/2582.0 = 1.33, νs(NH3)s/νs(ND3)s = 3360.2/2460.1 = 1.33, and δs(NH3)s/δs(ND3)s = 1181.4/885.3 = 1.34 (superscript s indicates the more strongly bonded ammonia molecule). Strong evidence for the presence of two unevenly bonded ammonia molecules in the IID complex comes F

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reported for H3N−BH3 (0.40). However, one must remember that the numerical value of ∇2ρBCP is quite sensitive to the method and basis set applied. These values are on the borderline between shared (covalent) and closed-shell interactions, as might be expected for a dative bond. The covalent contributions to the N−S bond in H3N·SO3 are highlighted by the negative value of the energy density at BCP59 (−0.041) and also by the bond order of 0.384 calculated from the bond delocalization index. The N−S bond in H3N− SO3···NH3 is slightly weaker and less covalent as the energy density at BCP is only −0.020 and the order of the bond is 0.286. The ∇2ρBCP value for the weaker N···S bond in the 2:1 complex is quite small (0.05) and is comparable to the value reported for HCN−BF3 (0.05). For weak interactions characterized by low values of electron density at BCP the value of ∇2ρBCP is typically small, and its sign is of no importance. Also the energy density at BCP is close to zero, and the bond order calculated from the delocalization index is small (0.065). For all N−S bonds in H3N·SO3 and H3N−SO3···NH3 the local symmetry of the electron density distribution is axial, as indicated by the value of the ellipticity (εBCP = 0). These considerations suggest that the degree of charge transfer and covalent character in the stronger N−S bond of H3N−SO3···NH3 is somewhat less than in the H3N·SO3 and H3N·BH3 complexes. The second N−S bond in the 2:1 complex is a weak one of closed-shell character and is comparable to that in the HCN−BF3 complex. The nature of N−S bonds in the complexes of interest which emerges from AIM analysis agrees well with the results of the interaction energy decomposition, which are collected in Table 4. The total interaction energy of H3N and SO3 in the H3N· SO3 complexes results from a subtle interplay of all terms, while the leading attractive electrostatic term (−76.73 kcal mol−1) is comparable to the polarization contribution (−64.69 kcal mol−1). The polarization term conceptually represents “orbital relaxation energy” on going from the monomer orbitals to the supermolecule orbitals, and its relatively large value suggests that the orbitals undergo a significant change upon complex formation.52 Interestingly, the total interaction energy in H3N− SO3···NH3 calculated with respect to the two NH3 and SO3 moieties is −22.17 kcal mol−1, which indicates a slightly weaker interaction than that in H3N·SO3. Clearly, the formation of an adduct between H3N·SO3 and an additional NH3 molecule perturbs the N−S bond. Again, a polarization term is an important contribution to the total energy. To compare individual contributions from the N−S and N···S bond formation in the H3N−SO3···NH3 complex, we considered a hypothetical two-step process where NH3 and SO3 form a H3N−SO3 complex adopting geometrical parameters from the H3N−SO3···NH3 followed by addition of the second ammonia

Figure 5. Molecular graph and AIM atomic charges for the H3N− SO3···NH3 complex. The smallest sphere represents bond critical points (BCPs). (a) Localization domains of the electron localizability indicator plotted for ELI-D(r) = 1.15 (b).

weakly bonded NH3 molecules to SO3. For comparison with the literature data we have also calculated the extent of charge transfer from the NBO atomic charges. It is equal to 0.25 e in the 1:1 complex, which is close to the value calculated at the HF/6-31+G(2d, p) level (0.28 e).19 Both values are lower than the one (0.36 e) determined from the analysis of the nitrogen experimental quadrupole coupling constant;5 however, one must remember that an atomic charge in a molecule is not a quantum chemical observable. For the 2:1 complex the NBO determined values are equal to 0.17 and 0.02 e for the more strongly and more weakly bonded NH3 moieties, respectively. In order to gain further information on the electronic structure and strength of the two complexes, we have carried out topological analyses defined within the framework of Bader’s theory of atoms in molecules.43 The molecular graphs representing the nets of bonds in H3N·SO3 and H3N−SO3··· NH3 are presented in Figures 3a and 5a, and corresponding numerical values are collected in Table 3. As expected, in both adducts the formation of an N−S bond or bonds is confirmed by the presence of the respective (3, −1) bond critical point in the electron density. The value of the electron density at BCP, ρBCP, of the N−S bond in H3N−SO3 (0.098 e/bohr3) is relatively small (which indicates a bond weaker than a typical covalent one) but is greater than the ρBCP value for the stronger N−S bond in H3N−SO3···NH3 (0.072 e/ bohr3). These two ρBCP values are much greater than ρBCP for the weaker N−S bond of the latter complex (0.019 e/bohr3). The ρBCP value for N−S in H3N·SO3 is close to that of H3N− BH3 (0.10 e/bohr3), which is a prototypical system with a strong B−N dative bond,57 whereas the ρBCP value for the longer N···S bond in the 2:1 complex is comparable to that for HCN−BF3 (0.02 e/bohr3).58 The Laplacians, ∇2ρBCP, slightly positive for H3N·SO3 (0.003) and for the stronger N−S bond in the 2:1 complex (0.055), are slightly less positive than that

Table 3. N−S Bond Distance [Å] and Selected Data from AIM Properties of NH3−SO3 and NH3−SO3···NH3 Complexesa complex

bond

q(SO3)

q(NH3)

R

ρBCP

∇2ρBCP

HBCP

εBCP

DI

NH3−SO3 NH3−SO3···NH3 NH3−SO3···NH3

N−S N−S N···S

−0.275 −0.252 −0.252

+0.275 +0.202 +0.050

2.088 2.230 2.852

0.098 0.072 0.019

0.003 0.055 0.049

−0.041 −0.020 +0.001

0 0 0

0.384 0.286 0.065

q(X), total charge on the X moiety; ρBCP, electron density at the bond critical point [a.u.]; ∇2ρBCP, Laplacian of the electron density [a.u.]; HBCP, energy density [a.u.]; εBCP, ellipticity; DI, bond order (delocalization index).

a

G

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Table 4. Interaction Energy Decomposition [kcal mol−1] Calculated for the NH3−SO3 and NH3−SO3···NH3 Complexesa NH3−SO3···NH3 electrostatic exchange repulsion polarization dispersion total

NH3−SO3

NH3−SO3···NH3

NH3−SO3

[NH3−SO3]···NH3

−76.73 −45.66 171.58 −64.69 −9.85 −25.35

−67.90 −38.16 140.24 −44.54 −11.81 −22.17

−55.00 −32.09 117.25 −41.77 −8.28 −19.89

−11.36 −7.08 24.64 −4.94 −3.54 −2.28

a

Column three contains the energy decomposition with respect to the two NH3 and SO3 moieties. Column four contains the energy decomposition within the NH3−SO3 moiety. Column five contains the energy decomposition with respect to NH3−SO3 and NH3.

value of the bond order, and a polarization contribution to the interaction energy comparable to the electrostatic one. The bond is formed due to an electron transfer from the N lone pair to the S−O antibonding orbital of the acceptor molecule as illustrated by a reasonable value of the NBO second-order stabilization energy. All of these facts allow us to categorize the stronger N−S bond as a dative bond. The weaker N−S bond is a closed-shell one; however, it falls outside any further concepts. A low value of the interaction energy between H3N−SO3 and NH3 (−2.28 kcal mol−1) might suggest a vdW interaction; however, the absolute values of the interaction energy components are larger than for typical vdW complexes. Small but not negligible values of the electron transfer from NH3 to SO3 and the value of the NBO secondorder donor−acceptor stabilization energy as well as the characteristic structural motif with the lone pair on the NH3 molecule directed toward the S atom acceptor might suggest a weak electron donor−acceptor complex. Note that the presence of the second NH3 molecule in the H3N−SO3··· NH3 complex influences the geometrical parameters of the H3N−SO3 moiety such as the N−S bond length (2.23 Å vs 2.09 Å for the equilibrium structure of H3N·SO3) or the OSOO torsion angle (166.3 vs 156.6° for the equilibrium structure of H3N·SO3). Some Remarks on the Origin of the H3N−SO3···NH3 Complex. The identification of this complex in the studied matrixes raises a question about its origin. The obtained experimental data do not allow us to give a definitive answer to this question. However, they lead to a hypothesis that the complexes are formed in the sample during its warming-up, close to the temperature of crystal decomposition. Then, they are sublimed and entrapped in the matrix. There are at least four experimental facts that lead to such a suggestion. First, the IID complexes are not formed when the matrix is obtained in a traditional way by the simultaneous deposition of SO3 diluted with the NH3/Ar mixture, even with a large excess of NH3 in the matrix. Second, the most stable 2:1 complexes, (NH3)2· SO3, IIHB, were not identified in the matrixes in spite of the fact that IIHB is ca. 8 kcal mol−1 more stable than IID. Third, the number of IID complexes trapped in the matrix decreases with prolonged matrix deposition time and at higher temperatures for the sulfamic acid sample during its deposition. Finally, the IID complexes entrapped in the matrix decompose into the H3N·SO3 adduct and NH3 when the matrix is annealed to 35 K. These four experimental facts seem to exclude the possibility of formation of the IID complex in the gas phase or on the surface of the matrix. The thermochemical data for the formation of the H3N·SO3 complex as well as the products of reactions H3N· SO3 + NH3 → H3N−SO3···NH3 and H3N·SO3 + NH3 → (H3N)2SO3 are collected in Table 5.

molecule. The interaction energy between moieties in H3N− SO3 is −19.89 kcal mol−1, and the data from its decomposition is collected in the fourth column of Table 4. Subsequently, we analyze the interaction between this moiety and the second NH3 molecule, which leads to the value of −2.28 kcal mol−1. The respective decomposition energy data are collected in column five. The numerical values of the interaction energies calculated in such a way agree well with the results from the AIM analysis as the total interaction energy of H3N−SO3···NH3 is dominated by the interaction energy within the strongly bonded NH3−SO3 moiety. In the case of the ammonia molecule weakly bonded to H 3 N−SO 3 , the attractive contributions to the interaction energy are dominated by the electrostatic term whereas the polarization term is much smaller but not negligible. Although the total interaction energy is relatively small, the absolute values of particular contributing terms are larger than for a typical van der Waals adduct.52 The difference between the stronger and weaker N−S bonds in H3N−SO3···NH3 is also clearly manifested in the results of NBO analysis. Both bonds are related to the interaction between the orbital representing a lone pair on a particular nitrogen atom and the antibonding S−O orbital in the acceptor SO3 molecule, while the second-order stabilization energy due to the donor−acceptor charge transfer is 40.77 kcal mol−1 for the stronger bond and only 4.11 kcal mol−1 for the weak one. The spatial localization of the bonding and nonbonding electron pairs visualized using topological properties of the electron localizability indicator calculated for H3N·SO3 and H3N−SO3···NH3 complexes are depicted in Figures 3b and 5b. As expected, in both complexes the lone electron pair on N atom is directed toward the S atom in the SO3 acceptor. In the first complex the electron density integrated over the V(N) basin representing the electron pair engaged in the formation of the N−S dative bond yields 1.99 e. The stronger N−S bond in the H3N−SO3···NH3 complex is represented by two basins, the first formally classified as a lone pair on the N atom V(N) populated by 1.76 e and the V(S) basin populated by 0.25 e, which highlights the partial covalent character of the N−S bond. The lone pair on the second nitrogen atom is directed toward the S atom, but the basin representing the lone pair does not share the surface with the basin representing the atomic core of the sulfur atom. The results of the computational study discussed above allow us to conclude that the N−S bonds in the H3N−SO3···NH3 complex are different with respect to their strength and covalent contributions, even though both are formed by NH3 molecules approaching the S atom with their lone electron pairs. The stronger bond, although of closed-shell type (∇2ρBCP > 0), has some covalent contribution, as illustrated by the negative value of the electron density at BCP, a reasonable H

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Laplacian, ∇2ρBCP, also have similar values for the weak S−N bonds in these two complexes (0.017 and 0.052 versus 0.019 and 0.049 in the SO2 and SO3 complexes, respectively).

Table 5. DFT/B3LYP/aug-cc-pVTZ-Calculated Binding Energies [kcal mol−1] and Gibbs Free Energies [kcal mol−1] for the Binary Association Reactions Resulting in the Formation of Ammonia/Sulfur Trioxide Complexes of 1:1 and 2:1 Stoichiometry reaction

ΔEbind

ΔEdys

ΔG(298 K)

ΔG(453 K)

NH3 + SO3 → NH3−SO3 NH3−SO3 + NH3 → NH3−SO3···NH3 NH3−SO3 + NH3 → (NH3)2−SO3

−18.18 −1.16

14.77 0.73

−5.81 5.55

−0.74 8.38

−11.00

8.68

−0.90

3.41

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CONCLUSIONS The H3N·SO3 (I) and H3N−SO3···NH3 (IID) complexes were identified in argon matrixes doped with thermolysis products of sulfamic acid. In contrast to the thoroughly studied 1:1 H3N· SO3 complex, the H3N−SO3···NH3 complex has been identified and characterized for the first time. It has not been identified in the matrixes obtained by the simultaneous deposition of the NH3/Ar mixture and SO3 vapor. The DFT/B3LYP/aug-cc-pVTZ calculations indicate that the H3N−SO3···NH3 complex corresponds to a local minimum on PES of the ammonia/sulfur trioxide system of 2:1 stoichiometry. It is 7.95 kcal mol−1 less stable than the (NH3)2SO3 complex (IIHB) corresponding to a global minimum on the PES of this system. In the IIHB complex one of the NH groups of the H3N·SO3 adduct serves as a proton donor for the nitrogen atom of the second ammonia molecule, forming a relatively strong N−H···N bond (R(H···N) = 1.862 Å). In the H3N−SO3···NH3 complex the second ammonia molecule is attached to the sulfur atom of the 1:1 adduct in such a way that the S atom and the two N atoms of the ammonia molecules are collinear. One of the two ammonia molecules is more strongly bonded (R(N−S) = 2.230 Å) than the other one (R(N−S) = 2.852 Å). The identification of the two bands in the IR spectra of the complex corresponding to the δs(NH3) motion of the ammonia molecules provides strong evidence for the presence of two unevenly bonded moieties. The AIM charges computed from the electron density obtained at the DFT/B3LYP/aug-cc-pVTZ level indicate a transfer of 0.20 e from the more strongely bonded and 0.05 e from the more weakly bonded NH3 molecules to SO3 in the IID complex as compared to the transfer of 0.28 e in the H3N·SO3 adduct. In AIM topological analyses, the interaction energy decomposition and topological properties of the electron localizability indicator (ELI-D) allow us to categorize the stronger N−S bond in the IID complex as a dative bond. The corresponding analyses performed for the weak N···S bond in IID suggest a weak electron-donor−acceptor interaction.

The calculated Gibbs free energy for the H3N·SO3 + NH3 → H3N−SO3···NH3 reaction explains why the IID complexes decompose after matrix annealing to 35 K. In Figure 6, we

Figure 6. van’t Hoff plot for the H3N·SO3 + NH3 association reaction in the range of 11−298 K. Thermochemical data are calculated at the DFT level.

present the van’t Hoff plot, the logarithm of association constant ln K plotted vs inverse absolute temperature T−1 for the H3N·SO3 + NH3 → H3N−SO3···NH3 association reaction. The logarithm of the association constant of IID is negative in the temperature range of 454−35 K and becomes positive below 35 K. When the temperature of the complex increases to ca. 35 K or above, the ΔG value becomes positive and its subunits tend to stay apart, so the complex may dissociate into NH3 and H3N·SO3. Hydrogen-bonded complex IIHB should be more stable than H3N−SO3···NH3 at all temperatures, and ΔG of the association process, H3N·SO3 + NH3 → (H3N)2SO3, approaches 0 at ca. 330 K. We must remember, however, that the formation of complexes under the experimental conditions is a nonequilibrium process and metastable complexes are often detected in low-temperature matrix experiments. H3N·SO3·NH3 and H3N·SO2·NH3. The only other report on the complex involving the double EDA interaction with sulfur atoms concerns the adduct formed between ammonia and sulfur dioxide. The anhydrous reaction between NH3 and SO2 has been known for a very long time.36 Measurements of vapor pressures over a solid formed at various NH3:SO2 ratios indicated that with excess NH3, at temperatures below 15 K, the 2:1 adduct is formed in addition to the 1:1 adduct.37 Ab initio calculations performed for the 1:1 and 1:2 complexes resulted, respectively, in Cs and C2v structures, with the SO2 plane almost perpendicular to the N−S axis. In contrast to H3N−SO3···NH3, in the H3N·SO2·NH3 adduct the two S−N bonds are equivalent and have a length (2.852 Å) comparable to that of the weaker S−N bond in the SO3 complex (2.816 Å). The calculated amount of charge transferred from one ammonia molecule to SO2 equals 0.029 e and is also comparable to that of the SO3 complex. The topological parameters, namely, the electron density, ρBCP, and the

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ASSOCIATED CONTENT

S Supporting Information *

The Supporting Information is available free of charge on the ACS Publications website at DOI: 10.1021/acs.jpca.5b07936. Spectra of the argon matrixes. Optimized structures referring to the (NH3)2SO3 formula. DFT/B3LYP/augcc-pVTZ-optimized structure of the H3N(SO3)2 complex. relative energies of structures referring to the (NH3)2SO3 formula with respect to the IID minimum;. DFT/B3LYP/aug-cc-pVTZ-calculated harmonic wavenumbers and intensities of the H3N·SO3, H3N·SO3· NH3, (NH3)2SO3, and NH3(SO3)2 complexes. (PDF)

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AUTHOR INFORMATION

Corresponding Author

*Phone: +48 071 3757 475. E-mail: zofi[email protected]. wroc.pl. Notes

The authors declare no competing financial interest. I

DOI: 10.1021/acs.jpca.5b07936 J. Phys. Chem. A XXXX, XXX, XXX−XXX

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(20) Cioslowski, J.; Martinov, M. Effects of Solvation on Chemical Bonding: An Electron-Flow Analysis. J. Chem. Phys. 1995, 103, 4967− 4974. (21) Mo, Y.; Gao, J. Polarization and Charge-Transfer Effects in Lewis Acid−Base Complexes. J. Phys. Chem. A 2001, 105, 6530−6536. (22) Singh, A. K.; Mishra, P. C. Electronic Structures and Properties of Some Charge Transfer Complexes: Usefulness of Hybridization Displacement Charge. J. Mol. Struct.: THEOCHEM 2002, 584, 53−67. (23) Remko, M. Theoretical Study of Molecular Structure and GasPhase Acidity of Some Biologically Active Sulfonamides. J. Phys. Chem. A 2003, 107, 720−725. (24) Solimannejad, M.; Boutalib, A. Computational Investigation of SO3−NH3‑nXn (n = 0−3; X = F, Cl) Interactions. J. Phys. Chem. A 2004, 108, 10342−10345. (25) Starikov, A. G.; Minyaev, R. M. Electronic and Geometric Structures of Sulfamic Acid in the H3NSO3 · (A)n Clusters (A = NH3, H2O, HF; n = 0, 1, 2, 3): Ab Initio Calculations. Russ. J. Inorg. Chem. 2005, 50, 398−406. (26) Ida, B. N.; Fudacz, P. S.; Pulsifer, D. H.; Standard, J. M. A Gasand Condensed-Phase Density Functional Study of Donor−Acceptor Complexes of Sulfur Trioxide. J. Phys. Chem. A 2006, 110, 5831−5838. (27) Venter, G. A.; Dillen, J. A Computational Study of the Effect of an External Electric Field on the Geometry of Selected Donor− Acceptor Complexes. J. Mol. Struct.: THEOCHEM 2009, 915, 112− 121. (28) Lovejoy, E. R.; Hanson, D. R. Kinetics and Products of the Reaction SO3+NH3+N2. J. Phys. Chem. 1996, 100, 4459−4465. (29) Hunt, S. W.; Brauer, C. S.; Craddock, M. B.; Higgins, K. J.; Nienow, A. M.; Leopold, K. R. Microwave Observation of H3N-SO3··· H2O Using a Concentric, Dual Injection Nozzle Source. Chem. Phys. 2004, 305, 155−164. (30) Larson, L. J.; Tao, F. − M. Interactions and Reactions of Sulfur Trioxide, Water, and Ammonia: An ab Initio and Density Functional Theory Study. J. Phys. Chem. A 2001, 105, 4344−4350. (31) Pawlowski, P.; Okimoto, S. R. Structure and Stability of Sulfur Trioxide-Ammonia Clusters with Water: Implications on Atmospheric Nucleation and Condensation. J. Phys. Chem. A 2003, 107, 5327− 5333. (32) Pszona, M.; Haupa, K.; Bil, A.; Mierzwicki, K.; Szewczuk, Z.; Mielke, Z. Clustering of Sulfamic Acid: ESI MS and Theoretical Study. J. Mass Spectrom. 2015, 50, 127−135. (33) Sass, C. S.; Ault, B. S. Matrix-Isolation Infrared Spectroscopic Study of Sulfur-Dioxide Amine Complexes. J. Phys. Chem. 1984, 88, 432−440. (34) Ault, B. S. Matrix Isolation Studies of Reactive Intermediate Complexes. Rev. Chem. Intermed. 1988, 9, 233−269. (35) Wong, M. W.; Wiberg, K. B. Structures, Bonding, and Absorption Spectra of Amine-Sulfur Dioxide Charge-Transfer Complexes. J. Am. Chem. Soc. 1992, 114, 7527−7535. (36) Scott, W. D.; Lamb, D. Two Solid Compounds which Decompose into a Common Vapor. Anhydrous Reaction of Ammonia and Sulfur Dioxide. J. Am. Chem. Soc. 1970, 92, 3943−3946. (37) Landreth, R.; de Pena, R. G.; Heicklen, J. Thermodynamics of the Reactions (NH3)n·SO2(s) ↔ NH3(g) + SO2(g). J. Phys. Chem. 1974, 78, 1378−1380. (38) Givan, A.; Loewenschuss, A.; Nielsen, C. J.; Rozenberg, M. FTIR and Computational Studies of Pure and Water Containing SO3 Species in Solid Argon Matrices. J. Mol. Struct. 2007, 830, 21−34. (39) Frisch, M. J.; Trucks, G. W.; Schlegel, H. B.; Scuseria, G. E.; Robb, M. A.; Cheeseman, J. R.; Scalmani, G.; Barone, V.; Mennucci, B.; Petersson, G. A.; et al. Gaussian 09, Revision D.01; Gaussian, Inc.: Wallingford, CT, 2013. (40) Raymond, K. S.; Wheeler, R. A. Compatibility of CorrelationConsistent Basis Sets with a Hybrid Hartree-Fock/ Density Functional Method. J. Comput. Chem. 1999, 20, 207−216. (41) Denis, P. A. Basis Set Requirements for Sulfur Compounds in Density Functional Theory: a Comparison between CorrelationConsistent, Polarized-Consistent, and Pople-Type Basis Sets. J. Chem. Theory Comput. 2005, 1, 900−907.

ACKNOWLEDGMENTS A grant of computer time from the Wrocław Center for Networking and Supercomputing (WCSS) is gratefully acknowledged.

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REFERENCES

(1) Santmeyer, R.; Aarons, R. Kirk-Othmer Encyclopedia of Chemical Technology; Interscience: New York, 1991. (2) Benson, G. A.; Spillane, W. J. Sulfamic Acids and Its N Substituted Derivatives. Chem. Rev. 1980, 80, 151−186. (3) Spillane, W. J.; Malaubier, J. B. Sulfamic Acids and Its N and OSubstituted Derivatives. Chem. Rev. 2014, 114, 2507−2586. (4) Benoit, R. L.; Boulet, D.; Frechette, M. Solvent Effect on the Solution, Ionization, and Structure of Aminosulfonic Acids. Can. J. Chem. 1988, 66, 3038−3043. (5) Canagaratna, M.; Phillips, J. A.; Goodfriend, H.; Leopold, K. R. Structure and Bonding of the Sulfamic Acid Zwitterion: Microwave Spectrum of +H3N−SO3−. J. Am. Chem. Soc. 1996, 118, 5290−5295. (6) Canagaratna, M.; Ott, M. E.; Leopold, K. R. Determination of the Dipole Moment of H3N−SO3 in the Gas Phase. Chem. Phys. Lett. 1997, 281, 63−68. (7) Sass, C. S.; Ault, B. S. Infrared Spectra and Structure of the Matrix-Isolated 1:1 Complexes of Sulfur Trioxide with Ammonia and the Methylamines. J. Phys. Chem. 1986, 90, 1547−1551. (8) Sass, C. S.; Ault, B. S. Infrared Matrix Isolation Study of the Interaction of Sulfur-Trioxide with Soft Lewis Bases. J. Phys. Chem. 1987, 91, 551−554. (9) Sass, R. L. A Neutron Diffraction Study on the Crystal Structure of Sulfamic Acid. Acta Crystallogr. 1960, 13, 320−324. (10) Bats, J. W.; Coppens, P.; Koetzle, T. F. The Experimental Charge Density in Sulfur Containing Molecules. A Study of the Deformation Electron Density in Sulfamic Acid at 78 K by X-ray and Neutron Diffraction. Acta Crystallogr., Sect. B: Struct. Crystallogr. Cryst. Chem. 1977, B33, 37−45. (11) Harbison, G. S.; Kye, Y. S.; Penner, G. H.; Grandin, M.; Monette, M. 14N Quadrupolar, 14N and 15N Chemical Shift, and 14 N-1H Dipolar Tensors of Sulfamic Acid. J. Phys. Chem. B 2002, 106, 10285−10291. (12) Coppens, P.; Gururow, T. N.; Leung, P.; Stevens, E. D.; Becker, P. J.; Yang, Y. W. Electron Population Analysis of Accurate Diffraction Data. 7. Net Atomic Charges and Molecular Dipole-Moments from Spherical-Atom X-ray Refinements, and the Relation Between Atomic Charge and Shape. Acta Crystallogr., Sect. A: Cryst. Phys., Diffr., Theor. Gen. Crystallogr. 1979, A35, 63−72. (13) Douglas, J. E.; Kenyon, G. L.; Kollman, P. A. The AmmoniaSulfur Trioxide Interaction. An Ab Initio Study. Chem. Phys. Lett. 1978, 57, 553−556. (14) Cruickshank, D. W. J.; Eisenstein, M. The Role of d Functions in Ab-Initio Calculations: Part 1. The Deformation Densities of H3NSO3 and SO3−. J. Mol. Struct. 1985, 130, 143−156. (15) Kaliannan, P.; Vishveswara, S.; Rao, V. S. R. Sulphamic Acid in Zwitterion Form-An Ab Initio SCF MO Study. Curr. Sci. 1985, 54, 1174−1176. (16) Kaliannan, P.; Vishveswara, S.; Rao, V. S. R. Ab initio SCF-MO Study of the Molecular Structures of Aminomethanol, Aminesulfonic Acid and N-methyl-sulfamate. J. Mol. Struct.: THEOCHEM 1983, 105, 359−374. (17) Hickling, S. J.; Woolley, R. G. An Ab Initio Hartree-Fock Study of the Zwitterion of Sulphamic acid, +H3NSO3−. Chem. Phys. Lett. 1990, 166, 43−48. (18) Woolley, R. G. Ab Initio Geometries and Frequencies for Aminosulphuric acid, H2NSO2(OH), and the Sulphamate Anion, H2NSO3−. J. Mol. Struct.: THEOCHEM 1991, 228, 329−333. (19) Wong, M. W.; Wiberg, K. B.; Frisch, M. J. Solvent Effects. 2. Medium Effect on the Structure, Energy, Charge Density, and Vibrational Frequencies of Sulfamic acid. J. Am. Chem. Soc. 1992, 114, 523−529. J

DOI: 10.1021/acs.jpca.5b07936 J. Phys. Chem. A XXXX, XXX, XXX−XXX

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The Journal of Physical Chemistry A (42) Grimme, S.; Antony, J.; Ehrlich, S.; Krieg, H. A Consistent and Accurate Ab Initio Parametrization of Density Functional Dispersion Correction (DFT-D) for the 94 Elements H-Pu. J. Chem. Phys. 2010, 132, 154104. (43) Bader, R. F. W. Atoms in Molecules: A Quantum Theory; Oxford University Press: Oxford, U.K., 1995. (44) Bader, R. F. W.; Stephens, M. E. Spatial Localization of the Electronic Pair and Number Distributions in Molecules. J. Am. Chem. Soc. 1975, 97, 7391−7399. (45) Fradera, X.; Austen, M. A.; Bader, R. F. W. The Lewis Model and Beyond. J. Phys. Chem. A 1999, 103, 304−314. (46) Keith, T. A. AIMAll, Version 11.12.19; TK Gristmill Software: Overland Park, KS, USA, 2011; aim.tkgristmill.com. (47) Kohout, M.; Pernal, K.; Wagner, F. R.; Grin, Y. Electron Localizability Indicator for Correlated Wavefunctions. I. Parallel-Spin Pairs. Theor. Chem. Acc. 2004, 112, 453−459. (48) Kohout, M. A Measure of Electron Localizability. Int. J. Quantum Chem. 2004, 97, 651−658. (49) Kohout, M. Bonding Indicators from Electron Pair Density Functionals. Faraday Discuss. 2007, 135, 43−54. (50) Kohout, M. DGrid, version 4.6; Radebeul, 2011. (51) Reed, A. E.; Curtiss, L. A.; Weinhold, F. Intermolecular Interactions from a Natural Bond Orbital, Donor-Acceptor Viewpoint. Chem. Rev. 1988, 88, 899−926. (52) Su, P.; Li, H. Energy Decomposition Analysis of Covalent Bonds and Intermolecular Interactions. J. Chem. Phys. 2009, 131, 014102. (53) Schmidt, M. W.; Baldridge, K. K.; Boatz, J. A.; Elbert, S. T.; Gordon, M. S.; Jensen, J. H.; Koseki, S.; Matsunaga, N.; Nguyen, K. A.; Su, S.; et al. General Atomic and Molecular Electronic Structure System. J. Comput. Chem. 1993, 14, 1347−1363. (54) Süzer, S.; Andrews, L. FTIR Spectra of Ammonia Clusters in Noble Gas Matrices. J. Chem. Phys. 1987, 87, 5131−5140. (55) Bondybey, V. E.; English, J. H. Infrared Spectra of SO3 Polymers and Complexes in Rare-Gas Matrices. J. Mol. Spectrosc. 1985, 109, 221−228. (56) Nelander, B.; Nord, L. Complex Between Water and Ammonia. J. Phys. Chem. 1982, 86, 4375−4379. (57) Smith, E. L.; Sadowsky, D.; Phillips, J. A.; Cramer, C. J.; Giesen, D. J. A Short Yet Very Weak Dative Bond: Structure, Bonding, and Energetic Properties of N2-BH3. J. Phys. Chem. A 2010, 114, 2628− 2636. (58) Jonas, V.; Frenking, G.; Reetz, M. T. Comparative-Theoretical Study of Lewis Acid-Base Complexes of BH3, BF3, BCl3, AlCl3 and SO2. J. Am. Chem. Soc. 1994, 116, 8741−8753. (59) Cremer, D.; Kraka, E. Chemical Bonds Without Bonding Electron Density? Does the Difference Electron-Density Analysis Suffice for a Description of the Chemical Bond? Angew. Chem., Int. Ed. Engl. 1984, 23, 627−628.

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DOI: 10.1021/acs.jpca.5b07936 J. Phys. Chem. A XXXX, XXX, XXX−XXX