Odd molecules

The study of the reactions of the simple free radicals begun by Bodenstein and Lind in 1906 on the kinetics of gas phase reactions showed that the reactions of H2 with CI2 and Bt2 were complex processes/ and a radical chain mechanism for these reactions (equations 14-18) was proposed in 1919 by Christiansen, Herzfeld, and Polanyi/ The theoretical basis for understanding these reactions in terms of free radicals was presented by G.N. Lewis in 1916, with the theory of the electron pair bond, and free radicals, or odd molecules / Further studies on chain reactions including the extension to explosions in gaseous systems were made by Hinshelwood and by Semenovwho shared the Nobel Prize in 1956. 

Lewis hi his It tG paper and in his book on valence emphasized the fact that there exist only a few stable molecules and complex ions (other than those containing atoms of the transition elements) for which the total number of electrons is odd. He pointed out that in general an odd molecule, such as nitric oxide or nitrogen dioxide, would be expected to use its unpaired electron to form a bond with another such molecule, and that the monomeric substance should accordingly be very much less stable than its dimer and he stated that the method by winch the unpaired electron is firmly held in the stable odd molecule v/as not at that time understood. Since then the explanation of the phenomenon has been found, as the result of the... 

Nitric Oxide.—Nitric oxide is the most stable of the odd molecules. For the first of the two structures I and II... 

No structural studies have as yet been made for other simple odd molecules, NO, CIO, IO , which may contain three-electron bonds. The nitrosodisulfonate ion, [ON(SO ) ]—, which has been shown by magnetic measurements of the potassium salt to be an odd ion, prob-... 

Perhaps one of the most interesting applications of electron affinities and the Born-Haber treatment is the calculation of stabilities of odd molecules such as NeF, NaF2, and CaF (Exercises 15-17). 

Over a dozen oxides of the halogens have boon characterized, many of them quite unstable. Perhaps the most important are chlorine dioxide, CJ02, and iodine pentoxide, I2O5. Chlorine dioxide (boiling point 11° C) is an odd molecule (p. 62), but apparently it shows no tendency to dimerize. Although it has been used as an antiseptic in water purification and as a bleach, it must be handled in diluted form for it is explosive when alone. It is formed, along with HCIO4, when chlorates are treated with concentrated sulfuric acid, but a safer preparation involves reduction of a chlorate with oxalic acid. 

In the normal alkanes (Table 31) one finds the alternation of the increase of melting point also well known in the fatty acids. This is a consequence of the difference in crystal structure of even and odd molecules resulting from a difference in symmetry (centre or plane through the middle). 

Chlorine dioxide is an odd molecule that is, a molecule containing an odd number of electrons. It was pointed out by G. N. L.ewis in 19J6 that odd molecules (other than those containing transition elements) are rare, and that they are usually colored and are always paramagnetic (attracted by a magnet). Every electronic structure that can bjC written for chlorine dioxide contains one unpaired electron. This unpaired electron presumably resonates among the three atoms, the electronic structure of the molecule being a resonance hybrid ... 

Perhaps the best known of the inorganic free radicals is the substance nitric oxide, one of the very few "odd molecules that are known to exist. Its method of preparation is a typical one. Briefiy, an equilibrium or stationary state is established at elevated temperatiures and then the temperature is suddenly reduced to a low enough point so that the equilibrium is frozen. It is then possible to study the chemical behavior of the particles at leisure. Among free radicals, nitric oxide is peculiar in that it behaves as a stable molecule at ambient temperatures. Its stability (18) may be interpreted as resulting from resonance between the structures +N—0 and N—0+. In the liquid and solid states, combination occiurs and nitric oxide exists entirely or almost entirely as the dimer. It seems probable that trapping NO in a matrix would be a useful way of studying the effectiveness of this method of stabilizing radicals. 

FIGURE 17.5 Change of C-C bond lengths with m for the /j = odd molecules. 

One of the products of a simple reaction between an odd molecule and a saturated one is another odd molecule. This is more likely to... 

Notes on the addition reactions of nitric oxide. Nitric oxide is an odd molecule, with an odd number of electrons. Probably because of this fact it is unusually active in forming coordination compounds. Examples of such coordination compounds and complex ions are (FeNO)++, [Co(NH3)6NO]++, CuNOCls-, FeNOCls, AINOCI3, Fe(CN)BNO", and the nitrosyl carbonyls, such as Co(CO)3NO. Many of these complexes are unstable and decompose on heating. They appear to be formed by the donation of either one or three electrons from the NO molecule thus in the nitroprusside ion, Fe(CN)5NO , produced by the action of nitric acid on a ferrocyanide, the nitric oxide is considered to contribute three electrons to the iron atom, leaving the latter in the ferrous rather than the ferric condition. Likewise the existence... 

Ozone is one of those odd molecules that negotiate with the octet rule. Its electronic structure can be written following the rule as two structures, which are depicted in the bottom part of Scheme 10.R. 1, with a double-headed arrow between them. Recall that the double-headed arrow means that the two structures contribute equally to the description of the molecule, which has thereby four mobile electrons (recall benzene in Scheme 4.R.3). It turns out, however, that quantum mechanics (QM) calculations predict that these two descriptions are minor, and the major description of ozone is the structure that is seen in the scheme to possess two regular 0—0 bonds and one long bond between the two terminal oxygen atoms. The bond lengths of the molecule show that this is not a short bond like the other two. 1 have put the reference of a scientific book, not as an actual source of reading material, but rather just to show you that this electronic structure is discussed somewhere in these terms. 

Red and blue salts discovered by C. Wurster as intermediates in the oxidation of paradiamines, were formulated as CgHuNaBr and CioHisNgBr, respectively. Bernthsen found that the red salt had quinonoid properties and formulated it as a quinone imide. Willstatter and J. Piccard regarded it as a meriquinonoid compound analogous to quinhydrone, but later work by L. Michaelis points to its being an odd molecule with a deficiency of one electron on one nitrogen atom. 

Looking at the structures of alkanes in the figure we can see that while the molecules with an odd number of carbon atoms appear symmetrical (both terminal carbon-carbon bonds being oriented upwards - black lines), where the structures with an even number are asymmetrical in the sense that the terminal carbon-carbon bonds are oriented upwards on the left end of the chain, but downwards on the right end (blue lines). Although the correct explanation is not simple, this example demonstrates how macroscopic properties correlate with microscopic structures. We can assume that odd molecules in the condensed state shall be packed differently from even molecules. 

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