The Properties of Oxygen Investigated with Easily Accessible

In the Classroom

The Properties of Oxygen Investigated with Easily Accessible Instrumentation The “One-Photon-Two-Molecule” Mechanism Revisited Manfred Adelhelm, Natasha Aristov, and Achim Habekost* Department of Chemistry, Padagogische Hochschule Ludwigsburg, Reuteallee 46, D-71634 Ludwigsburg, Germany *[email protected]

Oxygen has spectacular and unusual properties (1). Students are generally familiar with this gas as making up about 20% of the atmosphere, being required for combustion, participating in many oxidation reactions, and being a colorless gas. The revelation that it is blue as a liquid and paramagnetic is surprising. The observation that, under other conditions (as the product of a chemical reaction), oxygen emits red light, provides further amazement. Over the years, several methods of demonstrating the paramagnetic and optical properties of oxygen have been published in this Journal (2-6). Using these demonstrations is an elegant way of introducing or reinforcing the concepts of molecular orbital (MO) theory. Most recently, a demonstration using a small, strong, neodynium magnet was published (7). The availability of these magnets allows the demonstration of “magnetic” oxygen in most classrooms. The experiments described here show that the observation of the optical properties of oxygen is also easily possible with instruments that are commonly accessible, for example, handheld spectrophotometers and digital cameras. A review of the literature led us to consider the usual didactical discussion associated with the presentation of the paramagnetism and optical phenomena of oxygen. For example, often the blue color of liquid oxygen and its paramagnetism are presented in such a way to suggest that these two properties are intimately connected to one another. In fact, the structure of liquid oxygen is not simple, given that its open valence shell gives rise to intermolecular interactions stronger than van der Waals forces (8). The earliest measurements of oxygen's magnetic properties implied that some interactions among oxygen molecules must be occurring (9). That is, part of the liquefied oxygen is paramagnetic and consists of oxygen molecules and clusters of oxygen molecules with parallel aligned spins, but it is not a priori given that this is also the “blue” part. On the contrary, the blue part of the oxygen is most likely to consist of diamagnetic oxygen dimers (10). Thus, it is not the paramagnetism per se that is consistent with the observation of a blue color, but rather the deviation of the observed paramagnetism from the theoretically predicted value, which is usually not shown in classrooms. While using the paramagnetism of oxygen to demonstrate the reliability of MO theory for predicting electronic structure is legitimate, understanding the blue color (due to photon absorptions) of the liquid and the red chemiluminescence goes beyond the predictions of MO theory for O2 molecules. The strongest absorption and emission bands near 630 and 570 nm (in the red and the yellow-green) correspond to transitions 40

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between the ground state and first excited state of oxygen, but only if two ground-state molecules are promoted and relaxed to two excited-state molecules. This has been nicely presented as an exercise in spectral interpretation for the general chemistry laboratory in this Journal (11). A key point to understanding the absorption spectrum was the observation of the same features both for gaseous oxygen at high pressures and for large oxygenlayer thicknesses, on the order of the thickness of the earth's atmosphere.1 Inter-molecular interactions, more likely in condensed or high-pressure phases and more likely to be seen in a gas reservoir as huge as the atmosphere, were proposed, and later confirmed, to be the origin of the observed absorption bands (12-15). In fact, the earliest spectroscopists considered the existence of an O4 molecule to be likely (16). Nevertheless, in later interpretations, a “one-photon-two-molecule” mechanism was invoked, where the two molecules are interacting transiently as a “collision pair” (6, 17, 18). Much work has been done to understand the properties of O2 dimers and the likelihood of the formation of a tetraoxygen molecule in both the gas and various condensed phases of oxygen. Interest remains, not just because of the theoretical challenge of isolating and characterizing this interesting molecule, but also for practical reasons. It had been thought that oxygen polymers might serve as highly energetic materials (19). Theoretical calculations showed, however, no such long-lived O4 species to be available, as the barrier to dissociation to two O2 molecules was predicted to be low, about 20 kJ/mol (20). Even octaoxygen, O8, has been observed in solid oxygen (21). The tetraoxygen molecule, (O2)2/O4, continues to be the object of attention, for example, as a key metastable component in atmospheric processes such as the production of ozone, the de-excitation of vibrationally hot O2 (22, 23), and in the absorption of solar radiation (15). It also plays a role in chemiluminescent processes now commonly used in analytical methods (24, 25). In view of the results of scattering experiments, sophisticated spectroscopy, and extensive theoretical calculations, we argue in this article that the one-photon-two-molecule terminology for the absorption and the chemiluminescence spectra of oxygen be abandoned. Demonstrations Paramagnetism Our method of preparing liquid oxygen follows that of Shakhashiri (4, 26). We collect about 20-50 mL of liquid

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Vol. 87 No. 1 January 2010 pubs.acs.org/jchemeduc r 2009 American Chemical Society and Division of Chemical Education, Inc. 10.1021/ed800008g Published on Web 12/18/2009

In the Classroom Table 1. Lines Observed in the Absorption and Emission Spectra of Liquid Oxygen Relative Intensity (indicated by the number of x's) Absorption Transition

Transition Wavelength/nm

Spectrophotometer

Handheld Spectroscope (digital camera)

Handheld Spectroscope (VideoCom)

2 3Σg- (v = 0) f 2 1Δg (v = 0)

634

xxx

xxx

xxx

2 3Σg- (v = 0) f 2 1Δg (v = 1)

578

xxxxxx

xxxxxx

xxxx

534

xx

xx

x

478

xxx

xx

xx

446

x

380

xx

3

Σg-

1

(v = 0) f 2 Δg (v = 2) P 2 3Σg- (v = 0) f 1Δg þ 1 gþ (v = 0) P 2 3Σg- (v = 0) f 1Δg þ 1 gþ (v = 1) 3 1P þ 1P þ 2 Σg (v = 0) f g þ g (v = 0) P þ 3 1 1P þ 2 Σg (v = 0) f g þ g (v = 1) 2

362 Transition Wavelength/nm

xx Monochromator/ Photomultiplier Tube

Handheld Spectroscope (digital camera)

2 3Σg- (v = 0) r 2 1Δg (v = 0)a

634

xxxxx

xxxxx

2 3Σg- (v = 0) r 2 1Δg (v = 1)a

578

xxx

Emission Transition

a

v: level for both oxygen molecules.

Recording the Spectrum with a Spectrophotometer A small transparent Dewar, ∼10 cm length and 4 cm diameter, is precooled with liquid nitrogen, filled with liquid oxygen prepared as above, and placed into the beam path of a spectrophotometer (for example, a double-beam UV-vis Perkin-Elmer 555) in place of the usual cuvette.2 The oxygen remains liquefied for about 30 min, which is sufficient time to take a visible spectrum (Figure 2, panels A and B). Recording the Spectrum with a HandHeld Spectroscope and Camera

Figure 1. Electronic configurations of the three lowest electronic states of oxygen. For a more thorough treatment of the electronic structure of oxygen, see refs 11, 16, 27, and 28.

oxygen and direct it in a stream past a 0.25 T permanent magnet. To remove ice crystals, we pour the oxygen through a cone lined with filter paper that has been pierced in the center with a needle. For larger audiences, a video camera and monitor are used to enhance viewing ability. Optical Properties It is not difficult to assign the observed absorption and emission lines (Table 1) to well-documented transitions (18) among the lower electronic states of oxygen (Figure 1). The spectra were calibrated against a spectrum of Nd(NO3)3 (not shown here). Blue Color The blue color of liquid oxygen can be plainly seen, but can be more precisely investigated with spectrophotometers or, by each student individually, with handheld spectroscopes. Note that the spectra can be measured only to 350 nm, since Duran “test tube” glass is not transparent to UV.

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A small volume, 2-3 cm3, of liquefied oxygen is poured into a tall test tube (D50 Duran glass). Alternatively, a transparent (unsilvered) Dewar can be used in place of the test tube. The test tube is held up to a strong light source (generally, bright sunlight is sufficient) and observed through a handheld spectroscope (A. Kruss Optronic HS 1504, available through Leybold didactic, catalog no. 667-339). Ice buildup is removed as needed with a paper towel saturated with ethanol. The lens opening of a digital camera is placed directly behind the ocular of the spectroscope to take a photograph of the spectrum (Figure 2A). We used a Nikon CoolPix S4 because the diameter of its lens opening is exactly the same as that of the ocular on the spectroscope, saving us the labor of shielding the camera lens from stray light. As can be seen from Figure 2A, only low resolution is possible at the wide lens openings necessary to see the black absorption lines. This is a powerful demonstration of the complementary behavior of absorbed and scattered light. Viewing what is perceived by the naked eye as a blue liquid, one sees through a spectroscope red, (yellow-)green, and blue. Students' attention needs to be drawn to the black lines breaking up the continuity of the spectrum, that is, the light wavelengths that have been absorbed by the liquid oxygen. Their absence from the spectrum causes the color to appear blue. This concept is reinforced by examining the reverse process of emission of the red and yellowgreen photons in the chemiluminescence process discussed below. An elegant alternative is to use a Leybold didactic 337 47 VideoCom CCD camera to record the spectrum (Figure 2C). A light source irradiates a Dewar containing the liquid oxygen. It is

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In the Classroom

Figure 2. (A) Liquid oxygen absorption spectrum as seen through a handheld spectroscope. (B) UV-vis spectrum of liquid oxygen recorded with a double-beam spectrophotometer (Perkin-Elmer 555) showing features at 630, 578, 548, 446, 380, and 362 nm. Measurements beyond 350 nm are prevented by the Duran glass. (C) The absorption spectrum seen through the handheld spectroscope as recorded by the Leybold didactic VideoCom.

then spectrally analyzed in a (mounted) handheld spectroscope and projected by a focusing lens (f = 150 mm) onto the VideoCom. It is necessary to work in the dark. Chemiluminescence of Singlet Oxygen The reaction between hypochlorite anion and hydrogen peroxide in basic conditions will produce excited-state singlet oxygen, as first reported by Mallet in 1927 (29). After a brief induction period of about 10 s, an intense red emission is observed. This emission was spectrally resolved in two ways. The spectrum recorded with a prism monochromator and photomultiplier tube (Type 1 P 28A) is shown in Figure 3A. The monochromator is driven by a step motor. The analog signal was input to a Cassy interface (Leybold didactic) and then processed using the Cassy-Lab software. The image photographed with a digital camera aimed through the ocular of a handheld spectroscope pointed at the reaction vessel is shown in Figure 3B. To calibrate the wavelength axis, a mercury vapor spectrum (from the laboratory ceiling lighting) is shown for comparison (Figure 3C). 42


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Figure 3. (A) Spectrum of the Mallet reaction recorded with the prism monochromator, photomultipler tube, and Cassy interface. (B) Photograph of the Mallet reaction chemiluminescence spectrum taken with a digital camera through the ocular of a handheld spectroscope. (C) Mercury emission of the laboratory ceiling lighting taken as a calibration spectrum.

Discussion: Experimental and Theoretical Evidence for an Oxygen Dimer Other discussions of the history of the search for O4 can be found in refs 19, 30 and 31. The earliest work on the magnetism of liquid oxygen and its absorption spectrum suggested the existence of O4 molecules (9, 12, 16). Since then, oxygen dimers have been made (the references cited represent only some of the literature on this subject) under “soft” conditions, as in a gas cell at cold temperatures (15, 23, 32-35), in O2 þ O2 molecular beam scattering (10), and in molecular beam expansions (36, 37), and metastable excited states of (O2)2* have been produced (19, 38). Evidence for oxygen dimers has been observed in the atmosphere (13). That some O4 species exists that is more than just a molecule pair is implied, for example, by the observation of the reaction 2O2 f O3 þ O formation at energies lower than the bond dissociation energy of O2 (36). In fact, rovibrational spectra have been obtained for the [O2(3Σg-)v=0]2 f [O2(1Δg)v=0]2 transition in van der Waals dimers, an indication that these states of the (O2)2/O4 species exist for longer than just collisional interactions (34).


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In the Classroom

Only van der Waals complexes with some spin coupling between the O2 monomers can explain the experimental observations made thus far. For the dimer formed under soft conditions (molecular-beam scattering), the relative contribution to the O2-O2 binding energy from spin-spin interactions between two ground-state monomers has been found to be about 15% of that from the van der Waals interaction (10). Reviewing and summarizing the data of all these sources, the binding energies of the dimer formed from ground-state O2 are on the order of 10-20 meV and the O2-O2 separation is about 3-3.6 Å: those that correlate to excited-state O2 molecules have well depths of about 6-8 meV and separations of 3.2 - 4.0 Å. (For comparison, standard reference tables give the bond energy of a single O-O bond as about 1.5 eV and the bond length as about 1.5 Å.) A metastable state is seen at 4.1 eV above the separated ground-state O2 and found to be made from groundstate O2 and a highly excited 1Σu- state of O2. This state must have a lifetime on the order of microseconds to survive to detection (38). High-energy repulsive states of (O2)2 have been inferred in photoelectron detachment studies of O4- that correlate to 23Σg- and to 1Δg þ 3Σg- (39). A dimeric species must be responsible for the chemiluminescence in the Mallet reaction owing to the relaxation of two 1Δg O2 to two ground-state O2. As in the absorption, this transition is allowed only when two molecules are coupled together. Weakly bound states of (O2)2 with both O2 in 1Δg have been identified spectroscopically and predicted theoretically, as mentioned above. That such a dimer is being produced is reasonable in light of Khan and Kasha's (18) determination that only excited-state oxygen molecules can be produced in the Mallet reaction. The possibility of reactive processes concomitant with electronic transitions, that is, that the oxygens trade partners, has been evaluated. An interesting result obtained in an ion-molecule collision, neutralization, dissociative reionization reaction indicates that, for metastable excited states of (O2)2, no exchange reaction takes place. The decomposition products of the dissociation of 16O218O2 are 16O2 and 18O2; no mixed species, 16O18O, was observed (38). For reactions involving the ground states of the oxygen monomers, one notices that only the π* electrons (using, for simplicity, the notation for the separated oxygen monomers) can delocalize in this simple picture, the σp and π remain in the original molecules (top part of Figure 4). In fact, delocalization of the π* could even lead to a strengthening of the original O-O double bonds in the dimer, since it withdraws some antibonding electron density from the original bonds. Thus, a reactive process can be ruled as being unlikely. In the top part of Figure 4, we show a simple “frontier molecular orbital” picture of the interaction of the antibonding π* orbitals of two oxygen molecules in an “H” geometry, in which the oxygen molecules are parallel to each other, D2h symmetry group. Other limiting geometries of approach are shown in the bottom left of Figure 4. Of these, ab initio calculations have yielded stable structures only for the H and X approaches. Singlet, triplet, and quintet states are possible from 3Σg- þ 3Σg-. For these, an H geometry is the lowest energy species for the singlet and triplet, while the quintet is most stable in an X geometry, D2d, where the bond axes of the separated monomers are at an angle to each other (40). The singlet states of the dimer that dissociates to the 1Δg þ 1Δg excited states of O2 are equally stable in H or X geometries (34). This reconciles the early postulation of a “floppy” but “somewhat rigid” complex (32).


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Figure 4. (Top panel) Possible bonding combinations of the two π* orbitals of two oxygen molecules to form a 1B1u state of O4 in D2h symmetry (H geometry). The z-axis is perpendicular to the plane of the paper. To make the B3g orbital, the π* lobes are mixed above and below the plane of the paper; for the B2u, they are in-plane. (Bottom left panel) The limiting geometries of approach of two oxygen molecules: for H and X, the approach axis is shown out of the plane of the paper as indicated by the wide skewed arrow, and for T and L approach, the collision geometry is in the plane. (Bottom right panel) The calculated lowest energy geometry for cyclic, puckered, O4 and a further considered, but eliminated, covalently bound O4 species (10, 34, 40, 41).

Although theoretical calculations indicate its existence as a possibility (20, 42), a completely covalently bound, cyclic, O 4 molecule (bottom right of Figure 4) has not been observed. It is expected to be a puckered ring of D2d symmetry, with bond lengths similar to that in O2. There can be at least two reasons for this: First, formation of cyclic tetraoxygen in the gas phase is endothermic by about 4.27 eV (but it is bound by a low barrier to redissociation to 2O2 molecules, on the order of 0.5 eV) (30). Second, statistical mechanical calculations of the phase space available to freepair, metastable-pair, and true dimeric (covalently bound) states of oxygen showed relative fractions of about 51%, 47%, and 2%, respectively, near -183 °C, the boiling point of oxygen, and with the covalently bound-state fraction dropping off to zero above this temperature (43). Thus, cyclic O4 is not a likely product in the gas or liquid phases. Alternative limiting structures for O4 have been considered and eliminated, for example, a branched D3h shown in the bottom right of Figure 4. An extensive molecular dynamics study of liquid oxygen (8) showed that the electronic structure of O2 in the liquid is not different from the gas, the energy levels are mostly broadened but hardly shifted. Thus, it is reasonable to extrapolate the results for O4 in the liquid to those in the gas phase. For the liquid, it was also found that the most stable geometry in O4 is the H structure. Magnetic saturation was observed at O2-O2 separations less than 3.1 Å and a persistence of magnetic alignment was observed to separations as large as 4.4 Å.


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In the Classroom

Conclusion Considering all the experimental and calculational data collected thus far, we recommend no longer referring to a “onephoton-two-molecule” mechanism when discussing the absorption and chemiluminescence spectra of oxygen. Rather, they are due to spin-coupled and van der Waals-coupled metastable bound states. The oxygen molecules interact strongly enough with one another so that the spin and symmetry forbidden transitions between 3Σg- and 1Δg of the monomers become allowed in the coupled dimer between 2(3Σg-) and 2(1Δg). Acknowledgment We would like to thank Steffen Rieger for his photographic services and Regina Hornstein and Hans Strecker for their technical assistance. Notes 1. It has been estimated that oxygen dimers formed in the atmosphere contribute about 1-2% of the total absorption of solar radiation (15). They alone cannot be responsible for the blue color of the sky. The most important reason for the sky being blue is Rayleigh scattering. 2. Standard cuvettes are too small and not thermally insulated; the liquid oxygen evaporates too quickly.

Literature Cited 1. For a historical perspective, we refer the reader to Michael Faraday's discussion of oxygen: Faraday, M. Philos. Mag. 1847, 31 (Ser. 3), 409. 2. Lethbridge, J. W.; Davies, M. J. Chem. Educ. 1973, 50, 217. 3. Saban, G. H.; Moran, T. F. J. Chem. Educ. 1973, 50, 217. 4. Shakashiri, B. Z. Chemical Demonstrations; University of Wisconsin Press: Madison, WI; 1989, Vol. II, pp 147-152; originally published in Shakhashiri, B. Z.; Direen, G. E.; Williams, L. G. J. Chem. Educ. 1980, 57, 373–374. 5. Shakhashiri, B. Z.; Williams, L. G. J. Chem. Educ. 1976, 53, 358–361. 6. Shimada, H.; Yasuoka, T.; Mitzuzawa, S. J. Chem. Educ. 1990, 67, 62–63. 7. Mattson, B. J. Chem. Educ. 2007, 84, 1296–1298. 8. Oda, T.; Pasquarello, A. Phys. Rev. B 2004, 70, 134402–134420. 9. Lewis, G. N. J. Am. Chem. Soc. 1924, 46, 2027–2032. 10. Aquilanti, V.; Ascenzi, D.; Bartolomei, M.; Cappelletti, D.; Cavalli, S.; de Castro Vitores, M.; Pirani, F. J. Am. Chem. Soc. 1999, 121, 10794–10802 and references therein. 11. (a) Nyasulu, F.; Macklin, J.; Cusworth, W., III. J. Chem. Educ. 2002, 79, 356–359. See also (b) Puttemans, J. P.; Jannes, G. J. Chem. Educ. 2004, 81, 639. (c) Tudela, D.; Fernandez, V. J. Chem. Educ. 2003, 80, 1381. (d) Tudela, D.; Fernandez, V. J. Chem. Educ. 2004, 81, 197. 12. Ellis, J. W.; Kneser, H. O. Z. Phys. 1933, 86, 583–591. 13. Pfeilsticker, K.; Bosch, H.; Camy-Peyret, C.; Fitzenberger, R.; Harder, H.; Osterkamp, H. Geophys. Res. Lett. 2001, 28, 4595– 4598.

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The Properties of Oxygen Investigated with Easily Accessible

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