Electronic Structures of Atoms and Ions
On this page you can practice constructing electronic structures of atoms and ions. When you press "New Problem", an atom or ion will appear in the left side of the page. Place the electrons in the appropriate orbitals in the left side of the page, following the rules for electronic structures. To indicate the degree of occupancy of a given orbital, do the following:
- If the orbital has no electrons leave the cell empty. If it has a single electron, enter a 1 in the cell. If it has a pair of electrons, enter a 2 in the cell.
- Anything other than empty, a 1 or a 2 will cause your answer to be ruled incorrect.
- Special cases
- Half or fully filled d subshells (24 or 29 electrons)such as Cr,Cu must be properly noted. In general, all s1d5(24 electrons) and s1d10(29 electrons) configurations will be considered as preferred, provided the species has a charge of 1+ or less. Thus Mn1+ is a special case, but Fe2+ is not.
- For neutrals and anions the 4s fills before the 3d. For cations, the 3d fills first.
- Since this is an exercise in determining electronic configurations, no effort has been made to limit the possible species to observed ions. Thus, don't be surprised to see Na3-.
- Be certain to follow Hund's rule and, for simplicity, fill the orbitals in partially filled subshells from left to right. Thus, if you have 4 electrons in a p subshell, enter them 2 1 1. If you enter other than this, the program will try rearrange your entries according to the above convention.