As we saw earlier, hydrogen and helium are created by the filing of the 1s orbital. Lithium and beryllium arise from the placing of the third and fourth electrons in the 2s orbital, and boron through neon represent the subsequent filling of the three 2p orbitals with six more electrons. This completes the second row of the table. The next lowest orbitals are the 3s and the three 3p, and filling these orbitals produces the third row of the table. It is common to write the electronic structure of an atom by listing the orbitals in order of increasing n and I values, and indicating the number of electrons in each state with a superscript. As we have seen before, hydrogen has the arrangement 1s, and helium 1s. Lithium is 1s 2s, nitrogen is 1s 2s 2p, and phosphorus, just below nitrogen in the third row, is 1s 2s 2p 3s 3p, with the same s p outer electronic configuration. In the "Electronic Structures of the Elements" shown opposite and on pages 17, 18, 19 and 20, orbitals being filled are depicted with a white background, and completely filled orbitals are in color. An outer s p arrangement with full s and p orbitals is the hallmark of a noble gas (see table on previous page)

The energy gap between a p level and the next higher s level is always greater than the gaps between other nearby levels. This can be seen when the levels are stacked vertically, as on page 14. An unusually large amount of energy is required to add another electron to an atom that already has a filled set of s p orbitals, which is one reason why the noble gases are so unreactive.