Polyhedral skeletal electron pair theory for cluster compounds, including transition metals and main group elements such as boron including Wade's rules for polyhedral cluster compounds, including transition metals and main group elements and mixtures thereof.
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.
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 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 counting employing both approaches, though it is important to be aware that most chemical species exist between the purely covalent and ionic extremes.
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.
E.g. in period 2: B, C, N, O, and F have 3, 4, 5, 6, and 7 valence electrons, respectively.
E.g. in period 4: K, Ca, Sc, Ti, V, Cr, Fe, Ni have 1, 2, 3, 4, 5, 6, 8, 10 valence electrons respectively.
Add one for every halide or other anionic ligand which binds to the central 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.
Calculate the number of electrons of the element, assuming an oxidation state
e.g. for a Fe2+ has 6 electrons
S2- has 8 electrons
Add two for every halide or other anionic ligand which binds to the metal through a sigma bond.
Add two for every lone pair bonding to the metal (e.g. each phosphine ligand can bind with a lone pair). Similarly Lewis and Bronsted acids (protons) contribute nothing.
For unsaturated ligands such as alkenes, count the number of carbon atoms binding to the metal. Each carbon atom provides one electron.
The numbers of electrons "donated" by some ligands depends on the geometry of the metal-ligand ensemble. Perhaps the most famous example of this complication is the M-NO entity. When this grouping is linear, the NO ligand is considered to be a three-electron ligand. When the M-NO subunit is strongly bent at N, the NO is treated as a pseudohalide and is thus a one electron (in the neutral counting approach). The situation is not very different from the η-3 vs. η-1 allyl. Another unusual ligand from the electron counting perspective is sulfur dioxide.
Examples of electron counting
CH4, for the central C
neutral counting: C contributes 4 electrons, each H radical contributes one each: 4+4(1) = 8 valence electrons
conclusion: ionic counting indicates a molecule lacking lone pairs of electrons, therefore its structure will be octahedral, as predicted by VSEPR. One might conclude that this molecule would be highly reactive - but the opposite is true: SF6 is inert, and it is widely used in industry because of this property.
neutral counting: Fe contributes 8 electrons, the 2 cyclopentadienyl-rings contribute 5 each: 8 + 2(5) = 18 electrons
ionic counting: Fe2+ contributes 6 electrons, the two aromatic cyclopentadienyl rings contribute 6 each: 6 + 2(6) = 18 valence electrons on iron.
conclusion: Ferrocene is expected to be an isolable compound.
Please Note: These examples show the methods of electron counting, they are a formalism, and don't have anything to do with real life chemical transformations. Most of the 'fragments' mentioned above do not exist as such; they cannot be kept in a bottle: e.g. the neutral C, the tetraanionic C, the neutral Ti, and the tetracationic Ti are not free species, they are always bound to something, for neutral C, it is commonly found in graphite, charcoal, diamond (sharing electrons with the neighboring carbons), as for Ti which can be found as its metal (where it shares its electrons with neighboring Ti atoms!), C4- and Ti4+ 'exist' only with appropriate counterions (with which they probably share electrons). So these formalisms are only used to predict stabilities or properties of compounds!