The electronic configuration of an atom is given by listing its subshells with the number of electrons in each subshell, as shown in Table
1. Study the third column of complete electronic configurations carefully so you understand how electrons are added to the subshell of lowest energy until it reaches its capacity; then the subshell of the next energy level begins to be filled. The electrons in the highest numbered subshells are the
valence electrons, which comprise the valence shell of the atom.
For brevity, many chemists record the electron configuration of an atom by giving only its outermost subshell, like 4 s 1 for potassium or 4 s 2 for calcium. These electrons are most distant from the positive nucleus and, therefore, are most easily transferred between atoms in chemical reactions. These are the valence electrons.
For ions, the valence equals the electrical charge. In molecules, the various atoms are assigned chargelike values so the sum of the oxidation numbers equals the charge on the molecule. For example, in the H 2O molecule, each H has an oxidation number of +1, and the O is –2.
In Table, the common oxidation numbers in the last column are interpreted as the result of either losing the valence electrons (leaving a positive ion) or gaining enough electrons to fill that valence subshell. Table 2 compares three ions and a neutral atom.
The charges on the chlorine, potassium, and calcium ions result from a strong tendency of valence electrons to adopt the stable configuration of the inert gases, with completely filled electronic shells. Notice that the three ions have electronic configurations identical to that of inert argon. These ions and the atom of argon are known as isoelectronic.