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Liquid Oxygen is Blue, Why? The electronic properties of oxygen (i.e. the distribution of electrons) give rise to some interesting optical properties - this section is not for the light hearted. Firstly the gas is optically transparent with no real absorption at the standard atmospheric pressures of earth, however, in the condensed phase the liquid shows a strong blue colour. The boiling point for liquid O2 is -183°C. The closest thing that comes to mind is a bottle of Bombay Sapphire Absorption Spectrum If one condenses O2 gas via a cooling loop immersed in a liquid nitrogen trap (-196°C), it is possible to generate liquid oxygen. Bear some thoughts toward safety in doing this as (i) liquid oxygen is a cryogenic liquid, (ii) it is a potent oxidant and (iii) in liquid form is highly concentrated at 19M. (translation:- If something burning comes in contact with the liquid oxygen there could be a large explosion!) So that warning aside, the you then need a dewared cuvette and a spectrophotometer. The electronic absorption spectrum of liquid oxygen is blue due to the 2 strong absorption bands at 630 and 577 nm. The electronic bands appearing in O2 are somewhat complex. The electronic state of the ground state is and this nomenclature relates to the angular momentum of the diatomic molecule and the fact that the ground state is a triplet state. The +/- superscript and g/u subscript confer terms (symmetric, anti-symmetric, gerade, ungerade) that relate to the symmetry of the wave function of a diatomic molecule. The first excited state is singlet state with a different angular momentum and is designated .
. The energy of this transition is 1268.65 nm (0.9772 eV or 7882.4 cm -1) and is actually not resolved in the spectra as it was a NIR transition from the ground to excited state 0 vibrational state (0,0) and was outside of the range of my HP diode array. However the vibrational 1,0 band of this transition is resolved at 1064 nm, as is the 2,0 and 3,0 vibrational bands at 922 nm and 817 nm respectively. The second excited state of O2 is the singlet state with a energy of 1.6266 eV (13120.9 cm -1 or 762.14 nm) above the ground state. Both forms of "singlet" oxygen are particularly reactive, and the species in the first excited state is molecule of concern in biology. It is obvious in the above spectrum that there are other bands in the electronic spectrum of O2 than the ground state (not directly seen) and first excited state . These additional transitions arise from a "one-photon two-molecule" processes that result from the high concentrations of O2 molecules found in the condensed phase or at high pressures. The series of bands 300-700 nm occur at higher energies than the first and second excited states. These bands arise from the simultaneous excitation of two O2 molecules. The electronic configurations of oxygen is then explained in terms of a molecular orbital with the following schemes.
The ground state is a triplet state with two unpaired and parallel electron spins. This parallel pair accounts for the paramagnetic properties of oxygen. As discussed in the electrochemical section there is significant thermodynamic potential for O2 to react, yet it is the triplet state of O2 that causes "spin forbidden" reactions to proceed slowly. Typically most molecules are singlet states and for a triplet molecule to react with a singlet molecule there needs to be spin inversion. The quantum mechanical rules of chemistry dictate that this can not happen so the reaction relates triplet to singlet occur slowly, unless a catalyst is used. The first excited state has two paired electrons that occupy the same molecular orbital (note there is two molecular orbital's that these electrons can occupy designated: + and - ). The second excited state is also is a singlet state and is rapidly deactivated to the in condensed phase. This rapid deactivation accounts for the very weak absorption intensity in the figure above. Potential Energy Diagram Spectra of Diatomic Molecules (1950) G. Hertzberg The above figure is the potential energy diagram of oxygen showing the electronic and vibrational energies. The "Spectra of Diatomic Molecules" source book has much of the information contained in the text described herein. A solar rainbow was seen to grace my desk above this open page one morning. So I could only imagine that it was meant to be included in this page.
At higher energy it is also possible to excite electrons from the orbital into the and levels giving rise to the absorptions at 280 nm and 202 nm that are referred to as "Herzberg continuum" and the "Schumann-Runge bands". These absorptions bands account in part for the UV absorption by the atmosphere (O3 having an additional role) and arise from triplet states. Excitation above ~200 nm leads to dissociation of diatomic oxygen into atomic oxygen. The absorption of O3 is strongest 240-270 nm and extends a further ±40 nm. This absorption feature, the "Hartley band" is the protective UV screening layer in the atmosphere. At wavelengths longer than 300 nm the Hartley band weakens and from 310-350 nm a temperature dependent band termed the Higgins band appears . Astronomy of the Middle Atmosphere (1986) G. Brasseur and S. Solomon (Ed) Atomic Oxygen Emissions : Aurora The ionization of oxygen molecules in the atmosphere by UV radiation causes dissociation to oxygen atoms: O(1S), O(3P) or O(1D). These atoms then tend to emit their absorbed radiation via fluorescence and chemiluminescence from specific electronic transitions resulting in colourfull emission in the atmosphere . The emissions of oxygen species play an important role in the emissions of the Aurora and in the airglow of the night sky. The excited states of the oxygen are also important in atmospheric chemistry of O2 and O3. The strong red and green emissions in aurora arise from oxygen (OI) emissions which can be readily seen and photographed. Many other emissions are also detectable but may be very weak.
http://earthobservatory.nasa.gov/Features/ISSAurora/iss_aurora.phphttp://earthobservatory.nasa.gov/Features/ISSAurora/iss_aurora.php
