INSTRUCTOR: In the modern theory of the hydrogen atom or any one electron ion, all subshells with the same principle quantum number n have the same energy. Theory tells us, however, that an atoms containing more than one electron, the repulsive interactions of the electrons with one another cause some shells of the same principle quantum number to differ in energy. The one electron in the hydrogen atom in its lowest energy or ground state occupies the 1s orbital. In helium, atomic number 2, there are two electrons. Where does the second electron go? The first rule is that electrons occupy the lowest available orbital. The second rule follows from the Pauli exclusion principle. Any orbital can hold at most two electrons of opposite spin. These rules tell us that in helium, both electrons occupy the 1s atomic orbital. With increasing atomic number, the electrons occupy orbitals of increasing energy. With carbon we encounter a new question. Because there are three degenerate 2p orbitals that is orbitals of equal energy, Hund's Rule tells us that with two or more orbitals of the same energy, we should place electrons singly in each with parallel spins until all are half full. With nitrogen, we arrive at a half-filled subshell. A half-filled subshell is associated with special stability. As we come to oxygen having used up all the vacant 2p orbitals, we must pair electrons. With neon, we come to the end of the row and the stable octet of electrons in the outermost shell.