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we will put it in the 2px orbital. There is a gap in the table between Be and B. The reason for this will be clear when we discuss the fourth row below. The next element is carbon (C, 6) with six electrons. Now Hund’s Rule comes into play, and we put the sixth electron into the 2py orbital following the layout of Figure 11.3. Nitrogen is next (N, 7). Following Hund’s Rule, the seventh electron of N goes into the 2pz orbital so none of the electrons in the p orbitals are paired. Oxygen (O, 8) has eight electrons. The eighth electron must pair because the first seven electrons put two electrons in the 1s, two electrons in the 2s, and one electron in each of the 2p orbitals. To avoid spin pairing requires putting the eighth electron in the 3s orbital, which is much higher energy. So as in Figure 11.3, the eighth electron goes in the 2px orbital. Fluorine (F, 9), has its ninth electron go into the 2py orbital. Finally, Neon (Ne, 10) completes the n = 2 row or shell with 10 electrons. The 10th electron goes in the 2pz orbital.

Closed Shell Configurations

The electron configuration for neon is shown in Figure 11.5. No additional electrons can go in the second shell (n = 2 orbitals) without violating the Pauli Principle. As will be discussed, the elements, He, Ne, Ar, Kr, etc., that run down the last column on the right-hand side of the Periodic Table are special. These elements are called the noble gases. They all have closed shells, that is, the next element with one more electron goes into an orbital with the n quantum number one unit larger, which is substantially higher in energy.

FIGURE 11.5. The electron configuration for neon (Ne, 10). The second shell is complete.

Atoms Want to Form Closed Shell Configurations

We are now ready to use the energy level diagram, Figure 11.1, and the three rules for placing electrons in the energy levels to understand the structure of the Periodic Table and the properties of the atomic elements. The following chapters investigate in considerable detail what holds atoms together to make molecules. However, a great deal can be learned from an amazingly simple rule: Atoms will gain or lose electrons to obtain the nearest closed shell configuration. The closed shell electron configurations are the configurations of the noble gases that comprise the right-hand column of the Periodic Table. A closed shell configuration is particularly stable. The noble gases, also called the inert gases, have the closed shell configuration and are essentially chemically inert. The noble gases with small atomic numbers, helium, neon, and argon, do not form chemical compounds at all. The higher atomic number noble gases can be forced to form a small number of compounds under specialized conditions. Atoms other than the noble gases change in ways so that they achieve a stable closed shell electron configuration.

There are two ways that an atom can change the number of electrons it has to achieve a closed shell configuration. The first is to become a positive ion (cation) or negative ion (anion). The atom gives up one or more electrons and becomes positively charged (cation), or the atom takes on extra electrons, and becomes negatively charged (anion). The other is for an atom to share electrons with one or more other atoms. When two or more atoms share electrons, it is as if each atom has the electrons it needs. So an atom with fewer electrons than the number needed for the next closed shell configuration obtains the correct number, but so do the other atoms that are involved in the sharing. When atoms share electrons to get to the next closed shell number of electrons, the sharing holds the atoms together. The sharing makes the energy of the combined atoms lower than the energies of the individual open shell atoms. The lowering of the energy bonds the atoms together. This type of chemical bond is called a covalent bond. Covalent bonds are responsible for most of chemistry. The detailed nature of the covalent bond is presented for the simplest molecule, the hydrogen molecule in Chapter 12, and