Paramagnetism of oxygen

Unlike Lewis s theory, molecular orbital theory can account for the existence of electron-deficient compounds and the paramagnetism of oxygen. 

Figure 6.132 demonstrates the assembly of such a modular analyser system. The different analysers are based upon infrared absorption (see Figs 6.130 and 6.131), thermal conductivity, ultraviolet absorption (e.g. for C6H6, Cl2, 03, Hg (as a vapour), S02, H2S, N02, and toluene QH5-CH3), paramagnetism of oxygen and electrochemical cell for oxygen. 

The paramagnetism of oxygen is an anomaly in terms of the Lewis theory, although it is predicted by a more comprehensive theory that we will look at later. There are, however, a few other molecules that we would expect to be paramagnetic simply because they contain an odd number of valence electrons. The most well known example is nitric oxide, NO. Since oxygen has four and nitrogen has five outer electrons, the total number of valence electrons is nine, and magnetic measurements show that one of these is unpaired. 

These two unshared electrons are lesponsible for the paramagnetism of oxygen. ... 

An early triumph of molecular orbital theory was its ability to account for the observed paramagnetism of oxygen, O2. According to earlier theories, O2 wi expected to be diamagnetic, that is, to have only paired electrons. 

Then K = BF = Xgd, in a vacuum, or to good approximation in nitrogen gas. In air a correction for the paramagnetism of oxygen should be made ... 

The blue question mark suggests that there is some doubt about the validity of structure (10.9), and the source of the doubt is illustrated in Figure 10-3. The structure fails to account for the paramagnetism of oxygen—the O2 molecule must have unpaired electrons. Unfortunately, no completely satisfactory Lewis structure is possible for O2, but in Chapter 11, bonding in the O2 molecule is described in a way that accounts for both the double bond and the observed paramagnetism. 

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