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Electron Shells and Chemical Bonding : We know from our previous study of atoms that an atom consists of a positively charged nucleus surrounded by moving, negatively charged electrons. According to the shell model of atom, electrons behave as if they were arranged in concentric shells around the nucleus. As shown in Figure 1, there are at least seven shells available to the electrons in any atom. 

Each of these shells has a limited capacity for the number of electrons it can hold. The first shell can hold up to 2 electrons, while the second and third shells can each hold up to 8 electrons. The fourth and fifth shells can each hold 18 electrons, and the sixth and seventh shells can each hold 32 electrons. Figure 2 shows how this model applies to the first four elements of group 18, the noble gases. Electrons, being negatively charged, are attracted to the positively charged nucleus. They occupy the innermost shells first, where they are closest to the nucleus and possess minimum potential energy. 

Outer shells only get filled once the inner shells have reached their capacity for electrons. It is the exposed electrons in the outermost occupied shell that are the ones most responsible for an atom’s chemical properties, including its ability to form bonds with other atoms.a chemical bond is an electrostatic force of attraction between atoms that holds them together. To indicate their importance, an atom’s outer-shell electrons are called valence electrons and they are said to occupy the atom’s valence shell.

If we are going to keep track of the bonding behavior of an atom, we need to keep track of its valence electrons. We do this by depicting valence electrons as a series of dots surrounding an atomic symbol. The atomic symbol represents the nucleus and the atom’s inner-shell electrons. This notation is called the electrondot structure or, sometimes, a  Lewis dot symbol in honor of the American chemist G. N. Lewis who first proposed the concept of shells. Figure 3 shows the electrondot structures for the representative elements of the periodic table. When you look at the electron-dot structure of an atom, you immediately know two important things about that element that relate to its bonding behavior. You know how many valence electrons it has and how many of these are paired. Chlorine, for example, has three sets of paired electrons and one unpaired electron, and carbon has four unpaired electrons:

Paired valence electrons are relatively stable, or resistant to change. They sually do not form chemical bonds with other atoms. For this reason, electron pairs in an electron-dot structure are called  nonbonding pairs. Valence electrons that are unpaired, by contrast, have a strong tendency to participate in chemical bonding. By doing so, they become paired with an electron from another atom. The most stable electron arrangement for an atom is reached when all its valence electrons are paired so that its outermost occupied shell is filled to capacity. How can an atom with an unfilled valence shell attain a completely filled valence shell? It can share electrons with another atom or transfer electrons to another atom through bonding. The three types of chemical bonds discussed in this chapter ionic bonds, covalent bonds, and metallic bonds all result from either a transfer or a sharing of unpaired valence electrons. The noble gases are referred to as the inert gases because they are chemically nonreactive they almost never bond with other atoms. What is the reason for this? The valence shells of all noble gases are already filled to capacity, so they do not increase their stability by bonding. Atoms tend to bond with one another so that they achieve the same type of electron configuration as the noble gases. This tendency is called the octet rule:

Other arrangements of electrons can also increase an atom’s stability, especially for the transition metals. Nevertheless, the octet rule is the best indicator of stability for the main-block elements of the periodic table. Consider the bonding behavior of sodium, Na, as an example of the octet rule. We can see from Figure 3 that sodium, being a group 1 element, has one valence electron. If an atom has only one or only a few electrons in its valence shell, it will tend to lose its outer-shell electrons so that the next shell inward, which is filled, becomes the outermost occupied shell. Then, the atom will have a filled valence shell. Sodium readily gives up the single electron in its third shell. This makes the second shell, which is already filled to capacity, the outermost occupied shell.

Conceptual Integrated Science