Electron counting

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Electron counting is a formalism used for classifying compounds and for explaining or predicting electronic structure and bonding.[1] Many rules in chemistry rely on electron-counting:

Atoms that do not obey their rule are called "electron-deficient" when they have too few electrons to achieve a noble gas configuration, or "hypervalent" when they have too many electrons. Since these compounds tend to be more reactive than compounds that obey their rule, electron counting is an important tool for identifying the reactivity of molecules.

Contents

Counting rules

Two styles of electron counting are popular and both give the same result. The neutral counting approach assumes the molecule or fragment being studied consists of purely covalent bonds. It was popularized by M.L.H. Green along with the L and X ligand notation.[2][3] It is usually considered easier especially for low-valent transition metals.

The "ionic counting" approach assumes purely ionic bonds between atoms. It rewards the user with a knowledge of oxidation states, which can be valuable. One can check one's calculation by employing both approaches, though it is important to be aware that most chemical species exist between the purely covalent and ionic extremes.

Neutral counting

  • Locate the central atom on the periodic table and determine the number of its valence electrons. One counts valence electrons for main group elements differently from transition metals.
  • Add one for every halide or other anionic ligand which binds to the central atom through a sigma bond.
  • Add two for every lone pair bonding to the metal (e.g. each Lewis base binds with a lone pair). Unsaturated hydrocarbons such as alkenes and alkynes are considered Lewis bases. Similarly Lewis and Bronsted acids (protons) contribute nothing.
  • Add one for each homoelement bond.
  • Add one for each negative charge, and subtract one for each positive charge.

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