Odd-Electron Molecules

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Source: OpenStax Chemistry 2e

Odd-Electron Molecules (OpenStax Chemistry 2e)

We call molecules that contain an odd number of electrons free radicals. Nitric oxide, NO, is an example of an odd-electron molecule; it is produced in internal combustion engines when oxygen and nitrogen react at high temperatures.

To draw the Lewis structure for an odd-electron molecule like NO, we follow the same five steps we would for other molecules, but with a few minor changes:

  1. Determine the total number of valence (outer shell) electrons. The sum of the valence electrons is 5 (from N) + 6 (from O) = 11. The odd number immediately tells us that we have a free radical, so we know that not every atom can have eight electrons in its valence shell.
  2. Draw a skeleton structure of the molecule. We can easily draw a skeleton with an N–O single bond:
    N–O
  3. Distribute the remaining electrons as lone pairs on the terminal atoms. In this case, there is no central atom, so we distribute the electrons around both atoms. We give eight electrons to the more electronegative atom in these situations; thus oxygen has the filled valence shell:
    A Lewis structure shows a nitrogen atom, with one lone pair and one lone electron single bonded to an oxygen atom with three lone pairs of electrons.
  4. Place all remaining electrons on the central atom. Since there are no remaining electrons, this step does not apply.
  5. Rearrange the electrons to make multiple bonds with the central atom in order to obtain octets wherever possible. We know that an odd-electron molecule cannot have an octet for every atom, but we want to get each atom as close to an octet as possible. In this case, nitrogen has only five electrons around it. To move closer to an octet for nitrogen, we take one of the lone pairs from oxygen and use it to form a NO double bond. (We cannot take another lone pair of electrons on oxygen and form a triple bond because nitrogen would then have nine electrons:)
    A Lewis structure shows a nitrogen atom, with one lone pair and one lone electron double bonded to an oxygen atom with two lone pairs of electrons.

Related Topic: Writing and Balancing Chemical Equations

Source:

Flowers, P., Theopold, K., Langley, R., & Robinson, W. R. (2019, February 14). Chemistry 2e. Houston, Texas: OpenStax. Access for free at: https://openstax.org/books/chemistry-2e

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Related External Link:

Studies of a Large Odd‐Numbered Odd‐Electron Metal Ring: Inelastic Neutron Scattering and Muon Spin Relaxation Spectroscopy of Cr8Mn