1. Draw a legitimate Lewis structure consistent with the observed or hypothetic geometry of the ligand. If a non-linear geometry is observed, put a "lone pair" on the vertex atom. If a linear geometry is observed, be sure to use the correct sp-hybridization.See Rule 5 for the special case of bridging hydrogen and rule 6 for dimers and clusters.
2. Assign a formal charge to each atom of the ligand, according to the usual conventions. In assigning the formal charge to the atom(s) of the ligand which connect to the central metal, the electrons involved in covalent bonds should be split equally. A dative interaction (, where L is a ligand and M is a metal) implies that both of the electrons should be included in the ligating atom's formal charge count. For dative bonds, add 2 electrons to the ligand. For covalent bonds, add 1 electron to the ligand. Let the quantity be equal to the sum of the atomic formal charges over the entire ligand.
3. The number of electrons that any given ligand "donates" to the central atom's effective atomic number (EAN), defined by § is given by

4. The total EAN of an atom is given by

5. Bridging hydrides should be considered as

6. It is usually best to count each central atom individually when it is part of a dimer or a cluster. This is particularly necessary if the atoms are not symmetry related.
Here are a few illustrative examples. NO (nitrosyl), is of the form , so and . For , and . CO (carbonyl) is of the form , so and . NH₃ is of the form , so and .