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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

Elements and the Periodic Table - The terms element and atom are often used in a similar context. You might hear, for example, that gold is an element made of gold atoms. Generally, element is used in reference to an entire macroscopic or microscopic sample, and atom is used when speaking of the submicroscopic particles in the sample. The important distinction is that elements are made of atoms and not the other way around. How many atoms are bound together in an element is shown by an elemental formula. For elements in which the basic units are individual atoms, the elemental formula is simply the chemical symbol: Au is the elemental formula for gold, and Li is the elemental formula for lithium, to name just two examples. For elements in which the basic units are two or more atoms bonded into molecules, the elemental formula is the chemical symbol followed by a sub script indicating the number of atoms in each molecule. For example, elemental nitrogen, commonly consists of molecules containing two nitrogen atoms per molecule. Thus N2 is the usual elemental formula given for nitrogen. Similarly, O2 is the elemental formula for oxygen (two oxygen atoms per molecule), and S8 is the elemental formula for sulfur (eight atoms per molecule). 

The periodic table is a listing of all the known elements with their atomic masses, atomic numbers, and chemical symbols. Recall from our introduction to the atom in Chapter 9that the total mass of an atom is called its atomicmass. This is the sum of the masses of all the atom’s components (electrons, protons, and neutrons). A special unit is used for atomic masses. This is the atomic mass unit, amu, where 1 atomic mass unit is equal to 1.661 x 10-24 g , which is slightly less than the mass of a single proton. Also recall from Chapter 9that the atomic number of an element is the number of protons in an atom of that element. It is also equal to the number of electrons in the neutral atom. Besides atomic numbers, atomic masses, and chemical symbols, there is much more information about the elements in the periodic table (Figure 1). The way the table is organized in groups tells you a lot about the elements’ structures and how they behave. Look carefully at Figure 2. It shows that metals make up most elements.

Metals are defined as those elements that are shiny, opaque, and good conductors of electricity and heat. Metals are malleable, which means they can be hammered into different shapes or bent without breaking. They are also  ductile, which means they can be drawn into wires. All but a few metals are solid atroom temperature. The exceptions include mercury, Hg; gallium, Ga; cesium, Cs; and francium, Fr. These metals are all liquids at a warm room temperature of 30°C (86°F). Another interesting exception is hydrogen, H. Hydrogen acquires the properties of a liquid metal only at very high pressures (Figure 3).

Under normal conditions, hydrogen behaves as a nonmetallic gas. The nonmetallic elements, with the exception of hydrogen, are on the right of the periodic table. Nonmetals are very poor conductors of electricity and heat, and they may also be transparent. Solid nonmetals are neither malleable nor ductile. Rather, they are brittle and shatter when hammered. At 30°C (86°F), some nonmetals are solid (such as carbon, C). Other nonmetallic elements are liquid (such as bromine, Br). Still other nonmetals are gaseous (like helium, He). 

Six elements are classified as metalloids: boron, B; silicon, Si; germanium, Ge; arsenic, As; tin, Sn; and antimony, Sb. You’ll see them between the metals and the nonmetals in the periodic table. The metalloids have both metallic and nonmetallic characteristics. For example, they are weak conductors of electricity. This makes them useful as semiconductors in the integrated circuits of computers. Note in the periodic table how germanium, Ge (number 32), is closer to the metals than to the nonmetals. Because of this positioning, we can tell that germanium has more metallic properties than silicon, Si (number 14), and is a slightly better conductor of electricity. So we find that integrated circuits fabricated with germanium operate faster than those fabricated with silicon. Because silicon is much more abundant and less expensive to obtain, however, silicon computer chips remain the industry standard.