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The binary complex between hydrogen peroxide and ozone: A matrix isolation study
Chemical Physics 293 (2003) 203–209 www.elsevier.com/locate/chemphys
The binary complex between hydrogen peroxide and ozone: A matrix isolation study Anders Engdahl, Bengt Nelander * Chemical Physics, Chemical Center University of Lund, P.O. Box 124, S-22100 Lund, Sweden Received 4 April 2003; in final form 11 June 2003
Abstract The interaction between hydrogen peroxide and ozone has been studied in argon matrices using infrared spectroscopy. The spectra show that hydrogen peroxide forms a weak hydrogen bond to one of the end oxygens of ozone. A comparison is made with other ozone complexes. Ó 2003 Elsevier Science B.V. All rights reserved.
1. Introduction In the last three decades, numerous spectroscopic investigations of small molecular aggregates in molecular beams have been published. Surprisingly few of these have involved ozone complexes. Gillies et al. used microwave spectroscopy to find the structures of the ozone complexes of ethylene [1], acetylene [2], and water [3]. Considerably more work has been done using infrared spectroscopy to study complexes in inert matrices. In 1967, Schwager and Arkell prepared HOCl and HOBr [4] by UV irradiation of argon matrices containing ozone and HCl or HBr, respectively. In the same year, Weissberger et al. [5] prepared CO3 from CO2 and O3 . These experiments showed that
* Corresponding author. Tel.: +46462228183; fax: +4646222 4119. E-mail address: [email protected] (B. Nelander).
matrix isolated ozone complexes could be used as precursors for interesting oxidation products, a possibility which has been utilized in numerous publications [6–17]. In some of these, the presence of a complex has been inferred from the appearance of one or two satellites on the precursor or ozone fundamentals. In a few cases the ozone complexes has been studied in more detail. Andrews et al. have presented data for the ozone complexes of hydrogen fluoride [18], hydrogen cyanide [19], ammonia [20], phosphine [20], arsine [21], and CF3 I [22], where 16 O3 , 18 O3 and scrambled 16;18 O3 were used in order to get structural information. Schriver et al. made infrared spectroscopic studies of ozone [23], the interaction of ozone with water [24], methanol [25] and bromine [26] in argon matrices, also making use of isotope data. Nord made an exploratory infrared spectroscopic study of ozone complexes in nitrogen matrices [27]. All ozone complexes studied so far have been found to be weak, and in most cases only some aspects of their structures are known.
0301-0104/$ - see front matter Ó 2003 Elsevier Science B.V. All rights reserved. doi:10.1016/S0301-0104(03)00314-8
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Work by R€ as€ anen et al. [28] has demonstrated that the hydrogen peroxide–urea complex is an excellent source of hydrogen peroxide for matrix isolation work. They have made detailed studies of the photochemistry of hydrogen peroxide in noble gas matrices [29–31]. They have also studied the hydrogen peroxide complexes of nitrogen [32] and carbon monoxide [33]. Ault and coworkers have published a series of infrared spectroscopic studies of the interaction between hydrogen peroxide and bases [34,35] and acids [36]. We have for some time been engaged in a study of complexes between peroxy radicals and small stable molecules. Since hydrogen peroxide is a byproduct in the preparation of peroxy radicals in matrices, we have studied the dimerisation of hydrogen peroxide [37,38] and the water–hydrogen peroxide complex [39]. The present paper is the result of a study of the photochemistry of the ozone hydrogen peroxide complex, which was found to be a precursor to hydrogen trioxide [40].
2. Experimental Hydrogen peroxide was obtained by sweeping needle valve regulated argon over the urea–hydrogen peroxide adduct, kept a fixed temperature, as described by Pettersson et al. [28]. It was then deposited on a cold CsI window kept at 17 K by a Leybold RDK-320 closed cycle cooler. In different experiments the temperature of the adduct was varied between 10 and 20 °C to obtain spectra with only monomer present and spectra with significant concentrations of higher aggregates. Deuterated hydrogen peroxide was prepared as described in [28]. Ozone was prepared from O2 by a tesla discharge in a closed volume. It was condensed with liquid nitrogen, pumped on and then used to prepare suitable argon mixtures by standard manometric techniques. All spectra were recorded with a Bruker 113v FTIR spectrometer at 17 K. A resolution of 0.5 cm1 and an average of 512 scans was used in the mid infrared while in the far infrared it was sufficient to use a resolution of 1 cm1 and 128 scans. The fourth harmonic (266 nm) from a Continuum NY 81-20C YAG laser was used to irradiate the samples.
3. Assignment The spectrum of matrix isolated ozone is given by Schriver-Mazzuoli et al. [23]. Hydrogen peroxide in argon matrices has been studied by Pettersson et al. [28]. When ozone and hydrogen peroxide were codeposited in argon matrices, a set of new bands appeared, which were absent when only ozone or hydrogen peroxide was present in the matrix. The bands had constant intensity ratios in different experiments, and decreased at the same rate, when the matrix was irradiated with 266 nm radiation, suggesting that they are due to a single species. The new bands were all found close to a fundamental of ozone or hydrogen peroxide, indicating that the species contains ozone and hydrogen peroxide. When the matrix was irradiated, the intensities of the new bands decreased at approximately the same rate as the intensities of the ozone monomer bands, much faster than the intensities of the hydrogen peroxide dimer and monomer bands. Since some of the bands are too strong to be due to 1:2 or higher complexes we are convinced that they should be assigned to a 1:1 complex between ozone and hydrogen peroxide. The corresponding set of bands has also been observed for DOOD. The coupling between the OH bonds of HOOH is very weak, it has not been possible to resolve the symmetric and antisymmetric OH stretches of matrix isolated HOOH [28]. The OH stretch of HOOD almost coincides with the OH stretch of HOOH [28], while there seems to be a small shift between the OD stretches of DOOD and HOOD [28]. In the spectra after deposition, we only observe one OH stretch which can be assigned to the ozone complex. However, one OH stretch of complexed HOOH appears to be hidden under the HOOH monomer band. As can be seen from a comparison between Figs. 1 and 2 the OH stretch of the monomer (at 3588 cm1 in Fig. 1) decreases significantly during a short irradiation while the antisymmetric bend of the monomer is unaffected by the irradiation (1273 cm1 band in Fig. 2). It therefore seems clear that there is a significant HOOH–O3 contribution to the band at 3588 cm1 . The symmetric HOO bend of hydrogen peroxide is a very weak band, which has not been observed for
A. Engdahl, B. Nelander / Chemical Physics 293 (2003) 203–209
Fig. 1. The OH stretch region of hydrogen peroxide. Upper curve, HOOH, middle curve HOOH and O3 after deposition, lower curve HOOH–O3 after 11 min of photolysis. The peaks at 3581.6 and 3576.9 cm1 are due to hydrogen peroxide dimers.
Fig. 2. The antisymmetric bending region of hydrogen peroxide. Same experiments as in Fig. 1. The peak at 1293.5 and 1290.3 cm1 is due to HOOH–HOOH.
matrix isolated HOOH (or DOOD). When HOOH forms a hydrogen bond to another molecule, it is expected to become the bound HOO bend of complexed HOOH. In relatively strongly bound complexes, such as the hydrogen peroxide dimer it becomes strong enough to be observable. We have not been able to find the bound HOOH bend for
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the hydrogen peroxide ozone complex (HOOH– O3 , DOOD–O3 ). For HOOD and ozone we observe a band at 998.3 cm1 , which is likely due to the free DOO bend of DOOH–O3 . We have not been able to find the bound DOO bend of HOOD–O3 but it is expected above 1000 cm1 and may be hidden by the ozone absorption. The free HOO bend of HOOD–O3 and the bound HOO bend of DOOH–O3 are expected not far from 1350 cm1 where the CsI window gives serious base line problems in experiments with strongly oxidizing compounds. In D-experiments with very small concentrations of HOOH the bound OH stretch of the ozone complex shifts from 3561.4 cm1 to 3560.9 cm1 . In these experiments the HOOH concentration is very low and the major contribution to all observed OH stretches is due to HOOD. The shift therefore suggests that the bound OH stretch of DOOH–O3 is 0.5 cm1 lower than that of HOOH– O3 . In a D-experiment where the HOOD concentration was significantly higher than the DOOD concentration we noted a separate peak at 2630.1 cm1 , close to the bound OD stretch of DOOD–O3 at 2629.0 cm1 . In experiments with a higher D/H ratio, the 2629.0 cm1 band had a varying, additional absorption on its high wave number side. We therefore believe that the bound OD stretch of HOOD–O3 differs by 1.1 cm1 from the corresponding DOOD–O3 band. No bands which could be assigned to the ozone–hydrogen peroxide complex were observed in the far infrared. The assignment is summarized in Table 1.
4. Discussion Table 2 gives the complex shifts of the ozone fundamentals of number of different ozone complexes. Inspection of these shifts indicates that the bending vibration of ozone undergoes a small blue shift in all complexes. The symmetric stretching vibration has been observed only in a few cases. The shifts are shifts of the order of 1 or 2 cm1 except for HF which gives a 12.0 cm1 shift [18]. (Note that the unperturbed ozone frequencies
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Table 1 Observed bands of the hydrogen peroxide–ozone complex (cm1 ) Ozone bands O3 a
O3 –HOOH
O3 –DOOD
1105.1 703.6 1039.7 2108.7
1109.5 710.7 1034.7 2118.1
1109.5 710.9
HOOHb; c
DOODb;c
HOODb;c
Monomer bands m1; m5
3588.0
2646.1 2645.5
m6
1271.0
951.7
3587.2 2646.1 2645.5 981.2
m1 m2 m3 m1 þ m3
2117.1
Hydrogen peroxide bands
Complex bands m5 m1 m6
HOOH–O3
DOOD–O3
HOOD–O3
DOOH–O3
3588 3561.4 1280
2629.0 955.7
2630.1
3560.9 998.3
a
Ref. [23]. Ref. [28]. c Ref. [37]. b
Table 2 Shifts of the ozone intramolecular fundamentals for different complex partners (cm1 ) HOOHa m1 m2 m3 m1 þ m3
4.4 7.1 )5.0 9.4
H2 Ob 2 4 6 5.8
HFc 12.0 9.6 )13.0
HCNd 6.1 7.8 6.9
PH3 e
NH3 f
)2.0 0.8 )2.6
2.3 5.6
C2 H4 f
)1.3 )2.6
HBrf 4.8 )2.8
20
Blue shifts are given as positive numbers, red shifts are negative. (The position of the bands of free ozone are given in Table 1.) a This work. b Ref. [24]. c Ref. [18]. d Ref. [19]. e Ref. [20]. f Ref. [27].
given in Table 1 are taken from [23], while the complex shifts of the ozone fundamentals in Table 2 are taken directly from the different references. In some cases, the values of the unperturbed ozone frequencies used by different authors differ some-
what, less than 1.5 cm1 , from those of [23].) The antisymmetric stretching vibration is blue shifted in complexes which are believed to have the complex partner interacting symmetrically with the two end oxygens of ozone. The exceptions are the
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PH3 complex which has a 2.7 cm1 red shift [20] and the ethylene complex, which has a 2.6-cm1 red shift [27]. In complexes where the complex partner interacts mainly with one of the end oxygens the antisymmetric OO stretch has a red shift. For instance the large red shift of the HF fundamental in the HF–ozone complex (Table 3) indicates that HF forms a hydrogen bond to ozone, isotope data show that it interacts mainly with one of the end oxygens [18], and the antisymmetric stretch of ozone is red shifted 13 cm1 . The hydrogen bonded structure is supported by a recent ab initio calculation [41]. The water–ozone complex is known to be a non-hydrogen bonded complex [3,24] with the water molecule symmetrically between the two ozone oxygens (see Fig. 3). The OD stretch of the HDO complex is red shifted by 13 cm1 while the OH is blue shifted 4 cm1 [24], possibly an indication of an extremely weak interaction between one of the hydrogens and ozone. The CH stretch of HCN bound to ozone is red shifted 29.4 cm1 . Mielke and Andrews [19] have used experiments with mixtures of 16 O and 18 O ozone to show that hydrogen cyanide interacts symmetrically with the two oxygens and the antisymmetric OO stretch of ozone shifts 7.8 cm1 to the blue [19]. The qualitative similarity between the complex shifts of the ozone complexes of HDO and HCN suggest that the two complexes have similar structures. In the halogen–ozone complexes isotope data show that the halogen interacts with one of the end oxygens
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Fig. 3. Structures of the water and hydrogen peroxide complexes of ozone. The structure of the water complex was taken from [3]. The structure of the hydrogen peroxide complex is inferred from the experimental data of this work.
of ozone [26] and for instance in the bromine complex, the antisymmetric OO stretch is red shifted 9.0 cm1 . The presently available observations allows us to partition the ozone complexes in three classes, one where the antisymmetric O–O stretch is blue shifted, a second containing the ethylene and PH3 complexes and a third with a red shifted O–O stretch and indications that the interaction is
Table 3 Intramolecular fundamentals of hydrogen peroxide, water and hydrogen fluoride complexed with ozone, compared to the corresponding bands of the free components (cm1 )
m1 m2 m3 m5 m6 a
Ref. [28,37]. This work. c Ref. [46]. d Ref. [24]. e Ref. [47]. f Ref. [18]. b
HOOHa
HOOH–O3 b
H2 Oc
H2 O–O3 d
HFe
FH–O3 f
3588.0
3561.4
3638.0 1589.1 3734.3
3632.5 1592.5 3726.5
3919.5
3802.6
3588.0 1271.0
3588 1280
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preferentially with one of the end oxygens. In the cases where there is some structural information available, the complex partner of the blue shifting complexes interacts symmetrically with the end oxygens of ozone. The OH shifts of the hydrogen peroxide bound to ozone, clearly show that hydrogen peroxide forms a single hydrogen bond to ozone, one OH stretch is shifted and the other unshifted. As can be seen from the ozone–hydrogen cyanide complex, the interaction could still be symmetrically with the end oxygens of ozone. However, the ozone bands are shifted qualitatively similar to the ozone shifts of the ozone–hydrogen fluoride complex, indicating a singly hydrogen bonded structure. It should be noted that one of the OH stretches of complexed hydrogen peroxide is unshifted. When hydrogen peroxide acts as a hydrogen bond acceptor, as in the hydrogen peroxide dimer, the OH of the acceptor oxygen is red shifted. We therefore believe that the free OH of the complex points away from the ozone part of the complex (Fig. 3 gives a schematic picture of the complex structure suggested by our data) A comparison between the data of Goebel et al. on the HX shifts of hydro halides binding to hydrogen peroxide [36] and data for the corresponding water complexes [42,43] shows that hydrogen peroxide is a weaker hydrogen bond acceptor than water. The OH shifts of hydrogen peroxide [34,35] and water [44,45] forming hydrogen bonds to ammonia ()351 and )203 cm1 , respectively) and dimethyl ether ()234 and )95 cm1 , respectively) show that hydrogen peroxide forms stronger hydrogen bonds than water. This difference in hydrogen bond donor strength is apparently sufficient to switch the ozone complex from a non hydrogen bonded to a hydrogen bonded structure. Calculations by Lundell et al. for the HOOH– NN [32] and HOOH–CO [33] complexes allows us to estimate the binding energy of the HOOH–O3 complex. The dissociation energies of the HOOH– CO and HOOH–NN complexes are 4.1 and 2.0 kJ/ mol, respectively. (Using the calculated frequencies of the respective papers to correct for the changes in zero point vibration energy.) The OH shifts of the two complexes are )40.1 [33] and )5.7
cm1 [32] respectively. Our observed OH shift for the ozone complex, )26.6 cm1 , then suggests a dissociation energy of approximately 3 kJ/mol. For the ozone–water complex Gillies et al. [3] gave a calculated well depth of 2.9 kJ/mol. Since no vibration frequencies were given it is not directly comparable to our estimate, but after correction for zero point vibration energy it is likely to be at most half of our estimate for HOOH–O3 . Acknowledgements This work was supported by Vetenskapsr adet and by Carl Tryggers Stiftelse.
References [1] J.Z. Gillies, C.W. Gillies, R.D. Suenram, F.J. Lovas, W. Stahl, J. Am. Chem. Soc. 111 (1989) 3073. [2] J.Z. Gillies, C.W. Gillies, F.J. Lovas, K. Matsumura, R.D. Suenram, E. Kraka, D. Cremer, J. Am. Chem. Soc. 113 (1991) 6408. [3] J.Z. Gillies, C.W. Gillies, R.D. Suenram, F.J. Lovas, T. Schmidt, D. Cremer, J. Mol. Spectrosc. 146 (1991) 493. [4] I. Schwager, A. Arkell, J. Am. Chem. Soc. 89 (1967) 6006. [5] E. Weissberger, W.H. Breckenridge, H. Taube, J. Chem. Phys. 47 (1967) 1746. [6] L. Andrews, in: L. Andrews, M. Moskovits (Eds.), Absorption Spectroscopy of Molecular Ions and Complexes in Noble Gas Matrices, North-Holland, Amsterdam, 1989. [7] L. Andrews, F.K. Chi, A. Arkell, J. Am. Chem. Soc. 96 (1974) 1997. [8] M. Hawkins, L. Andrews, J. Am. Chem. Soc. 105 (1983) 2523. [9] R. Withnall, L. Andrews, J. Phys. Chem. 89 (1985) 3261. [10] R. Withnall, L. Andrews, J. Phys. Chem. 91 (1987) 784. [11] R. Withnall, L. Andrews, J. Phys. Chem. 92 (1988) 2155. [12] Z. Mielke, M. Hawkins, L. Andrews, J. Phys. Chem. 93 (1989) 558. [13] B.W. Moores, L. Andrews, J. Phys. Chem. 93 (1989) 1902. [14] L. Schriver, D. Carrere, A. Schriver, K. Jaeger, Chem. Phys. Lett. 181 (1991) 505. [15] V. Raducu, D. Jasmin, R. Dahoo, P. Brosset, B. GauthierRoy, L. Abouaf-Marguin, J. Chem. Phys. 101 (1994) 1878. [16] E.Y. Misochko, A.V. Akimov, C.A. Wight, J. Phys. Chem. A 103 (1999) 7977. [17] R.J.H. Clark, L.J. Foley, S.D. Price, J. Phys. Chem. A 104 (2000) 10675. [18] L. Andrews, R. Withnall, R.D. Hunt, J. Phys. Chem. 92 (1988) 78. [19] Z. Mielke, L. Andrews, J. Phys. Chem. 94 (1990) 3519.
A. Engdahl, B. Nelander / Chemical Physics 293 (2003) 203–209 [20] R. Withnall, M. Hawkins, L. Andrews, J. Phys. Chem. 90 (1986) 575. [21] L. Andrews, R. Withnall, B.W. Moores, J. Phys. Chem. 93 (1989) 1279. [22] L. Andrews, M. Hawkins, R. Withnall, Inorg. Chem. 24 (1985) 4234. [23] L. Schriver-Mazzuoli, A. Schriver, C. Lugez, A. Perrin, C. Camy-Peyret, J.-M. Flaud, J. Mol. Spectrosc. 176 (1996) 85. [24] L. Schriver, C. Barreau, A. Schriver, Chem. Phys. 140 (1990) 429. [25] C. Lugez, A. Schriver, R. Levant, L. Schriver-Mazzuoli, Chem. Phys. 181 (1994) 129. [26] M. Bahou, L. Schriver-Mazzuoli, A. Schriver, P. Chaquin, Chem. Phys. 105 (1997) 105. [27] L. Nord, J. Mol. Struct. 96 (1982) 37. [28] M. Pettersson, S. Tuominen, M. R€as€anen, J. Phys. Chem. A 101 (1997) 1166. [29] L. Khriachtchev, M. Pettersson, S. Tuominen, M. R€ as€ anen, J. Chem. Phys. 107 (1997) 7252. [30] S. Pehkonen, M. Pettersson, J. Lundell, L. Khriachtchev, M. R€ as€ anen, J. Phys. Chem. A 102 (1998) 7643. [31] L. Khriachtchev, M. Pettersson, S. Jolkkonen, S. Pehkonen, M. R€ as€ anen, J. Chem. Phys. 112 (2000) 2187. [32] J. Lundell, S. Pehkonen, M. Pettersson, M. R€as€anen, Chem. Phys. Lett. 286 (1998) 382.
209
[33] J. Lundell, S. Jolkkonen, L. Khriachtchev, M. Pettersson, M. R€as€anen, Chem. Eur. J. 7 (2001) 1670. [34] J.R. Goebel, B.S. Ault, J.E. Del Bene, J. Phys. Chem. A 105 (2001) 6430. [35] J.R. Goebel, B.S. Ault, J.E. Del Bene, J. Phys. Chem. A 104 (2000) 2033. [36] J.R. Goebel, K.A. Antle, B.S. Ault, J.E. Del Bene, J. Phys. Chem. A 106 (2002) 6406. [37] A. Engdahl, B. Nelander, G. Karlstr€ om, J. Phys. Chem. A 105 (2001) 8393. [38] A. Engdahl, B. Nelander, Phys. Chem. Chem. Phys. 4 (2002) 2140. [39] A. Engdahl, B. Nelander, Phys. Chem. Chem. Phys. 2 (2000) 3967. [40] A. Engdahl, B. Nelander, Science 295 (2002) 482. [41] H. Tachikawa, S. Abe, T. Iyama, Inorg. Chem. 40 (2001) 1167. [42] L. Andrews, G.L. Johnson, J. Chem. Phys. 79 (1983) 3670. [43] A. Engdahl, B. Nelander, J. Chem. Phys. 84 (1986) 1981. [44] B. Nelander, L. Nord, J. Phys. Chem. 86 (1982) 4375. [45] A. Engdahl, B. Nelander, J. Chem. Soc. Faraday Trans. 88 (1992) 177. [46] A. Engdahl, B. Nelander, J. Mol. Struct. 193 (1989) 101. [47] L. Andrews, G.L. Johnson, J. Phys. Chem. 88 (1984) 425.
The binary complex between hydrogen peroxide and ozone: A matrix isolation study Anders Engdahl, Bengt Nelander * Chemical Physics, Chemical Center University of Lund, P.O. Box 124, S-22100 Lund, Sweden Received 4 April 2003; in final form 11 June 2003
Abstract The interaction between hydrogen peroxide and ozone has been studied in argon matrices using infrared spectroscopy. The spectra show that hydrogen peroxide forms a weak hydrogen bond to one of the end oxygens of ozone. A comparison is made with other ozone complexes. Ó 2003 Elsevier Science B.V. All rights reserved.
1. Introduction In the last three decades, numerous spectroscopic investigations of small molecular aggregates in molecular beams have been published. Surprisingly few of these have involved ozone complexes. Gillies et al. used microwave spectroscopy to find the structures of the ozone complexes of ethylene [1], acetylene [2], and water [3]. Considerably more work has been done using infrared spectroscopy to study complexes in inert matrices. In 1967, Schwager and Arkell prepared HOCl and HOBr [4] by UV irradiation of argon matrices containing ozone and HCl or HBr, respectively. In the same year, Weissberger et al. [5] prepared CO3 from CO2 and O3 . These experiments showed that
* Corresponding author. Tel.: +46462228183; fax: +4646222 4119. E-mail address: [email protected] (B. Nelander).
matrix isolated ozone complexes could be used as precursors for interesting oxidation products, a possibility which has been utilized in numerous publications [6–17]. In some of these, the presence of a complex has been inferred from the appearance of one or two satellites on the precursor or ozone fundamentals. In a few cases the ozone complexes has been studied in more detail. Andrews et al. have presented data for the ozone complexes of hydrogen fluoride [18], hydrogen cyanide [19], ammonia [20], phosphine [20], arsine [21], and CF3 I [22], where 16 O3 , 18 O3 and scrambled 16;18 O3 were used in order to get structural information. Schriver et al. made infrared spectroscopic studies of ozone [23], the interaction of ozone with water [24], methanol [25] and bromine [26] in argon matrices, also making use of isotope data. Nord made an exploratory infrared spectroscopic study of ozone complexes in nitrogen matrices [27]. All ozone complexes studied so far have been found to be weak, and in most cases only some aspects of their structures are known.
0301-0104/$ - see front matter Ó 2003 Elsevier Science B.V. All rights reserved. doi:10.1016/S0301-0104(03)00314-8
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Work by R€ as€ anen et al. [28] has demonstrated that the hydrogen peroxide–urea complex is an excellent source of hydrogen peroxide for matrix isolation work. They have made detailed studies of the photochemistry of hydrogen peroxide in noble gas matrices [29–31]. They have also studied the hydrogen peroxide complexes of nitrogen [32] and carbon monoxide [33]. Ault and coworkers have published a series of infrared spectroscopic studies of the interaction between hydrogen peroxide and bases [34,35] and acids [36]. We have for some time been engaged in a study of complexes between peroxy radicals and small stable molecules. Since hydrogen peroxide is a byproduct in the preparation of peroxy radicals in matrices, we have studied the dimerisation of hydrogen peroxide [37,38] and the water–hydrogen peroxide complex [39]. The present paper is the result of a study of the photochemistry of the ozone hydrogen peroxide complex, which was found to be a precursor to hydrogen trioxide [40].
2. Experimental Hydrogen peroxide was obtained by sweeping needle valve regulated argon over the urea–hydrogen peroxide adduct, kept a fixed temperature, as described by Pettersson et al. [28]. It was then deposited on a cold CsI window kept at 17 K by a Leybold RDK-320 closed cycle cooler. In different experiments the temperature of the adduct was varied between 10 and 20 °C to obtain spectra with only monomer present and spectra with significant concentrations of higher aggregates. Deuterated hydrogen peroxide was prepared as described in [28]. Ozone was prepared from O2 by a tesla discharge in a closed volume. It was condensed with liquid nitrogen, pumped on and then used to prepare suitable argon mixtures by standard manometric techniques. All spectra were recorded with a Bruker 113v FTIR spectrometer at 17 K. A resolution of 0.5 cm1 and an average of 512 scans was used in the mid infrared while in the far infrared it was sufficient to use a resolution of 1 cm1 and 128 scans. The fourth harmonic (266 nm) from a Continuum NY 81-20C YAG laser was used to irradiate the samples.
3. Assignment The spectrum of matrix isolated ozone is given by Schriver-Mazzuoli et al. [23]. Hydrogen peroxide in argon matrices has been studied by Pettersson et al. [28]. When ozone and hydrogen peroxide were codeposited in argon matrices, a set of new bands appeared, which were absent when only ozone or hydrogen peroxide was present in the matrix. The bands had constant intensity ratios in different experiments, and decreased at the same rate, when the matrix was irradiated with 266 nm radiation, suggesting that they are due to a single species. The new bands were all found close to a fundamental of ozone or hydrogen peroxide, indicating that the species contains ozone and hydrogen peroxide. When the matrix was irradiated, the intensities of the new bands decreased at approximately the same rate as the intensities of the ozone monomer bands, much faster than the intensities of the hydrogen peroxide dimer and monomer bands. Since some of the bands are too strong to be due to 1:2 or higher complexes we are convinced that they should be assigned to a 1:1 complex between ozone and hydrogen peroxide. The corresponding set of bands has also been observed for DOOD. The coupling between the OH bonds of HOOH is very weak, it has not been possible to resolve the symmetric and antisymmetric OH stretches of matrix isolated HOOH [28]. The OH stretch of HOOD almost coincides with the OH stretch of HOOH [28], while there seems to be a small shift between the OD stretches of DOOD and HOOD [28]. In the spectra after deposition, we only observe one OH stretch which can be assigned to the ozone complex. However, one OH stretch of complexed HOOH appears to be hidden under the HOOH monomer band. As can be seen from a comparison between Figs. 1 and 2 the OH stretch of the monomer (at 3588 cm1 in Fig. 1) decreases significantly during a short irradiation while the antisymmetric bend of the monomer is unaffected by the irradiation (1273 cm1 band in Fig. 2). It therefore seems clear that there is a significant HOOH–O3 contribution to the band at 3588 cm1 . The symmetric HOO bend of hydrogen peroxide is a very weak band, which has not been observed for
A. Engdahl, B. Nelander / Chemical Physics 293 (2003) 203–209
Fig. 1. The OH stretch region of hydrogen peroxide. Upper curve, HOOH, middle curve HOOH and O3 after deposition, lower curve HOOH–O3 after 11 min of photolysis. The peaks at 3581.6 and 3576.9 cm1 are due to hydrogen peroxide dimers.
Fig. 2. The antisymmetric bending region of hydrogen peroxide. Same experiments as in Fig. 1. The peak at 1293.5 and 1290.3 cm1 is due to HOOH–HOOH.
matrix isolated HOOH (or DOOD). When HOOH forms a hydrogen bond to another molecule, it is expected to become the bound HOO bend of complexed HOOH. In relatively strongly bound complexes, such as the hydrogen peroxide dimer it becomes strong enough to be observable. We have not been able to find the bound HOOH bend for
205
the hydrogen peroxide ozone complex (HOOH– O3 , DOOD–O3 ). For HOOD and ozone we observe a band at 998.3 cm1 , which is likely due to the free DOO bend of DOOH–O3 . We have not been able to find the bound DOO bend of HOOD–O3 but it is expected above 1000 cm1 and may be hidden by the ozone absorption. The free HOO bend of HOOD–O3 and the bound HOO bend of DOOH–O3 are expected not far from 1350 cm1 where the CsI window gives serious base line problems in experiments with strongly oxidizing compounds. In D-experiments with very small concentrations of HOOH the bound OH stretch of the ozone complex shifts from 3561.4 cm1 to 3560.9 cm1 . In these experiments the HOOH concentration is very low and the major contribution to all observed OH stretches is due to HOOD. The shift therefore suggests that the bound OH stretch of DOOH–O3 is 0.5 cm1 lower than that of HOOH– O3 . In a D-experiment where the HOOD concentration was significantly higher than the DOOD concentration we noted a separate peak at 2630.1 cm1 , close to the bound OD stretch of DOOD–O3 at 2629.0 cm1 . In experiments with a higher D/H ratio, the 2629.0 cm1 band had a varying, additional absorption on its high wave number side. We therefore believe that the bound OD stretch of HOOD–O3 differs by 1.1 cm1 from the corresponding DOOD–O3 band. No bands which could be assigned to the ozone–hydrogen peroxide complex were observed in the far infrared. The assignment is summarized in Table 1.
4. Discussion Table 2 gives the complex shifts of the ozone fundamentals of number of different ozone complexes. Inspection of these shifts indicates that the bending vibration of ozone undergoes a small blue shift in all complexes. The symmetric stretching vibration has been observed only in a few cases. The shifts are shifts of the order of 1 or 2 cm1 except for HF which gives a 12.0 cm1 shift [18]. (Note that the unperturbed ozone frequencies
206
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Table 1 Observed bands of the hydrogen peroxide–ozone complex (cm1 ) Ozone bands O3 a
O3 –HOOH
O3 –DOOD
1105.1 703.6 1039.7 2108.7
1109.5 710.7 1034.7 2118.1
1109.5 710.9
HOOHb; c
DOODb;c
HOODb;c
Monomer bands m1; m5
3588.0
2646.1 2645.5
m6
1271.0
951.7
3587.2 2646.1 2645.5 981.2
m1 m2 m3 m1 þ m3
2117.1
Hydrogen peroxide bands
Complex bands m5 m1 m6
HOOH–O3
DOOD–O3
HOOD–O3
DOOH–O3
3588 3561.4 1280
2629.0 955.7
2630.1
3560.9 998.3
a
Ref. [23]. Ref. [28]. c Ref. [37]. b
Table 2 Shifts of the ozone intramolecular fundamentals for different complex partners (cm1 ) HOOHa m1 m2 m3 m1 þ m3
4.4 7.1 )5.0 9.4
H2 Ob 2 4 6 5.8
HFc 12.0 9.6 )13.0
HCNd 6.1 7.8 6.9
PH3 e
NH3 f
)2.0 0.8 )2.6
2.3 5.6
C2 H4 f
)1.3 )2.6
HBrf 4.8 )2.8
20
Blue shifts are given as positive numbers, red shifts are negative. (The position of the bands of free ozone are given in Table 1.) a This work. b Ref. [24]. c Ref. [18]. d Ref. [19]. e Ref. [20]. f Ref. [27].
given in Table 1 are taken from [23], while the complex shifts of the ozone fundamentals in Table 2 are taken directly from the different references. In some cases, the values of the unperturbed ozone frequencies used by different authors differ some-
what, less than 1.5 cm1 , from those of [23].) The antisymmetric stretching vibration is blue shifted in complexes which are believed to have the complex partner interacting symmetrically with the two end oxygens of ozone. The exceptions are the
A. Engdahl, B. Nelander / Chemical Physics 293 (2003) 203–209
PH3 complex which has a 2.7 cm1 red shift [20] and the ethylene complex, which has a 2.6-cm1 red shift [27]. In complexes where the complex partner interacts mainly with one of the end oxygens the antisymmetric OO stretch has a red shift. For instance the large red shift of the HF fundamental in the HF–ozone complex (Table 3) indicates that HF forms a hydrogen bond to ozone, isotope data show that it interacts mainly with one of the end oxygens [18], and the antisymmetric stretch of ozone is red shifted 13 cm1 . The hydrogen bonded structure is supported by a recent ab initio calculation [41]. The water–ozone complex is known to be a non-hydrogen bonded complex [3,24] with the water molecule symmetrically between the two ozone oxygens (see Fig. 3). The OD stretch of the HDO complex is red shifted by 13 cm1 while the OH is blue shifted 4 cm1 [24], possibly an indication of an extremely weak interaction between one of the hydrogens and ozone. The CH stretch of HCN bound to ozone is red shifted 29.4 cm1 . Mielke and Andrews [19] have used experiments with mixtures of 16 O and 18 O ozone to show that hydrogen cyanide interacts symmetrically with the two oxygens and the antisymmetric OO stretch of ozone shifts 7.8 cm1 to the blue [19]. The qualitative similarity between the complex shifts of the ozone complexes of HDO and HCN suggest that the two complexes have similar structures. In the halogen–ozone complexes isotope data show that the halogen interacts with one of the end oxygens
207
Fig. 3. Structures of the water and hydrogen peroxide complexes of ozone. The structure of the water complex was taken from [3]. The structure of the hydrogen peroxide complex is inferred from the experimental data of this work.
of ozone [26] and for instance in the bromine complex, the antisymmetric OO stretch is red shifted 9.0 cm1 . The presently available observations allows us to partition the ozone complexes in three classes, one where the antisymmetric O–O stretch is blue shifted, a second containing the ethylene and PH3 complexes and a third with a red shifted O–O stretch and indications that the interaction is
Table 3 Intramolecular fundamentals of hydrogen peroxide, water and hydrogen fluoride complexed with ozone, compared to the corresponding bands of the free components (cm1 )
m1 m2 m3 m5 m6 a
Ref. [28,37]. This work. c Ref. [46]. d Ref. [24]. e Ref. [47]. f Ref. [18]. b
HOOHa
HOOH–O3 b
H2 Oc
H2 O–O3 d
HFe
FH–O3 f
3588.0
3561.4
3638.0 1589.1 3734.3
3632.5 1592.5 3726.5
3919.5
3802.6
3588.0 1271.0
3588 1280
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A. Engdahl, B. Nelander / Chemical Physics 293 (2003) 203–209
preferentially with one of the end oxygens. In the cases where there is some structural information available, the complex partner of the blue shifting complexes interacts symmetrically with the end oxygens of ozone. The OH shifts of the hydrogen peroxide bound to ozone, clearly show that hydrogen peroxide forms a single hydrogen bond to ozone, one OH stretch is shifted and the other unshifted. As can be seen from the ozone–hydrogen cyanide complex, the interaction could still be symmetrically with the end oxygens of ozone. However, the ozone bands are shifted qualitatively similar to the ozone shifts of the ozone–hydrogen fluoride complex, indicating a singly hydrogen bonded structure. It should be noted that one of the OH stretches of complexed hydrogen peroxide is unshifted. When hydrogen peroxide acts as a hydrogen bond acceptor, as in the hydrogen peroxide dimer, the OH of the acceptor oxygen is red shifted. We therefore believe that the free OH of the complex points away from the ozone part of the complex (Fig. 3 gives a schematic picture of the complex structure suggested by our data) A comparison between the data of Goebel et al. on the HX shifts of hydro halides binding to hydrogen peroxide [36] and data for the corresponding water complexes [42,43] shows that hydrogen peroxide is a weaker hydrogen bond acceptor than water. The OH shifts of hydrogen peroxide [34,35] and water [44,45] forming hydrogen bonds to ammonia ()351 and )203 cm1 , respectively) and dimethyl ether ()234 and )95 cm1 , respectively) show that hydrogen peroxide forms stronger hydrogen bonds than water. This difference in hydrogen bond donor strength is apparently sufficient to switch the ozone complex from a non hydrogen bonded to a hydrogen bonded structure. Calculations by Lundell et al. for the HOOH– NN [32] and HOOH–CO [33] complexes allows us to estimate the binding energy of the HOOH–O3 complex. The dissociation energies of the HOOH– CO and HOOH–NN complexes are 4.1 and 2.0 kJ/ mol, respectively. (Using the calculated frequencies of the respective papers to correct for the changes in zero point vibration energy.) The OH shifts of the two complexes are )40.1 [33] and )5.7
cm1 [32] respectively. Our observed OH shift for the ozone complex, )26.6 cm1 , then suggests a dissociation energy of approximately 3 kJ/mol. For the ozone–water complex Gillies et al. [3] gave a calculated well depth of 2.9 kJ/mol. Since no vibration frequencies were given it is not directly comparable to our estimate, but after correction for zero point vibration energy it is likely to be at most half of our estimate for HOOH–O3 . Acknowledgements This work was supported by Vetenskapsr adet and by Carl Tryggers Stiftelse.
References [1] J.Z. Gillies, C.W. Gillies, R.D. Suenram, F.J. Lovas, W. Stahl, J. Am. Chem. Soc. 111 (1989) 3073. [2] J.Z. Gillies, C.W. Gillies, F.J. Lovas, K. Matsumura, R.D. Suenram, E. Kraka, D. Cremer, J. Am. Chem. Soc. 113 (1991) 6408. [3] J.Z. Gillies, C.W. Gillies, R.D. Suenram, F.J. Lovas, T. Schmidt, D. Cremer, J. Mol. Spectrosc. 146 (1991) 493. [4] I. Schwager, A. Arkell, J. Am. Chem. Soc. 89 (1967) 6006. [5] E. Weissberger, W.H. Breckenridge, H. Taube, J. Chem. Phys. 47 (1967) 1746. [6] L. Andrews, in: L. Andrews, M. Moskovits (Eds.), Absorption Spectroscopy of Molecular Ions and Complexes in Noble Gas Matrices, North-Holland, Amsterdam, 1989. [7] L. Andrews, F.K. Chi, A. Arkell, J. Am. Chem. Soc. 96 (1974) 1997. [8] M. Hawkins, L. Andrews, J. Am. Chem. Soc. 105 (1983) 2523. [9] R. Withnall, L. Andrews, J. Phys. Chem. 89 (1985) 3261. [10] R. Withnall, L. Andrews, J. Phys. Chem. 91 (1987) 784. [11] R. Withnall, L. Andrews, J. Phys. Chem. 92 (1988) 2155. [12] Z. Mielke, M. Hawkins, L. Andrews, J. Phys. Chem. 93 (1989) 558. [13] B.W. Moores, L. Andrews, J. Phys. Chem. 93 (1989) 1902. [14] L. Schriver, D. Carrere, A. Schriver, K. Jaeger, Chem. Phys. Lett. 181 (1991) 505. [15] V. Raducu, D. Jasmin, R. Dahoo, P. Brosset, B. GauthierRoy, L. Abouaf-Marguin, J. Chem. Phys. 101 (1994) 1878. [16] E.Y. Misochko, A.V. Akimov, C.A. Wight, J. Phys. Chem. A 103 (1999) 7977. [17] R.J.H. Clark, L.J. Foley, S.D. Price, J. Phys. Chem. A 104 (2000) 10675. [18] L. Andrews, R. Withnall, R.D. Hunt, J. Phys. Chem. 92 (1988) 78. [19] Z. Mielke, L. Andrews, J. Phys. Chem. 94 (1990) 3519.
A. Engdahl, B. Nelander / Chemical Physics 293 (2003) 203–209 [20] R. Withnall, M. Hawkins, L. Andrews, J. Phys. Chem. 90 (1986) 575. [21] L. Andrews, R. Withnall, B.W. Moores, J. Phys. Chem. 93 (1989) 1279. [22] L. Andrews, M. Hawkins, R. Withnall, Inorg. Chem. 24 (1985) 4234. [23] L. Schriver-Mazzuoli, A. Schriver, C. Lugez, A. Perrin, C. Camy-Peyret, J.-M. Flaud, J. Mol. Spectrosc. 176 (1996) 85. [24] L. Schriver, C. Barreau, A. Schriver, Chem. Phys. 140 (1990) 429. [25] C. Lugez, A. Schriver, R. Levant, L. Schriver-Mazzuoli, Chem. Phys. 181 (1994) 129. [26] M. Bahou, L. Schriver-Mazzuoli, A. Schriver, P. Chaquin, Chem. Phys. 105 (1997) 105. [27] L. Nord, J. Mol. Struct. 96 (1982) 37. [28] M. Pettersson, S. Tuominen, M. R€as€anen, J. Phys. Chem. A 101 (1997) 1166. [29] L. Khriachtchev, M. Pettersson, S. Tuominen, M. R€ as€ anen, J. Chem. Phys. 107 (1997) 7252. [30] S. Pehkonen, M. Pettersson, J. Lundell, L. Khriachtchev, M. R€ as€ anen, J. Phys. Chem. A 102 (1998) 7643. [31] L. Khriachtchev, M. Pettersson, S. Jolkkonen, S. Pehkonen, M. R€ as€ anen, J. Chem. Phys. 112 (2000) 2187. [32] J. Lundell, S. Pehkonen, M. Pettersson, M. R€as€anen, Chem. Phys. Lett. 286 (1998) 382.
209
[33] J. Lundell, S. Jolkkonen, L. Khriachtchev, M. Pettersson, M. R€as€anen, Chem. Eur. J. 7 (2001) 1670. [34] J.R. Goebel, B.S. Ault, J.E. Del Bene, J. Phys. Chem. A 105 (2001) 6430. [35] J.R. Goebel, B.S. Ault, J.E. Del Bene, J. Phys. Chem. A 104 (2000) 2033. [36] J.R. Goebel, K.A. Antle, B.S. Ault, J.E. Del Bene, J. Phys. Chem. A 106 (2002) 6406. [37] A. Engdahl, B. Nelander, G. Karlstr€ om, J. Phys. Chem. A 105 (2001) 8393. [38] A. Engdahl, B. Nelander, Phys. Chem. Chem. Phys. 4 (2002) 2140. [39] A. Engdahl, B. Nelander, Phys. Chem. Chem. Phys. 2 (2000) 3967. [40] A. Engdahl, B. Nelander, Science 295 (2002) 482. [41] H. Tachikawa, S. Abe, T. Iyama, Inorg. Chem. 40 (2001) 1167. [42] L. Andrews, G.L. Johnson, J. Chem. Phys. 79 (1983) 3670. [43] A. Engdahl, B. Nelander, J. Chem. Phys. 84 (1986) 1981. [44] B. Nelander, L. Nord, J. Phys. Chem. 86 (1982) 4375. [45] A. Engdahl, B. Nelander, J. Chem. Soc. Faraday Trans. 88 (1992) 177. [46] A. Engdahl, B. Nelander, J. Mol. Struct. 193 (1989) 101. [47] L. Andrews, G.L. Johnson, J. Phys. Chem. 88 (1984) 425.