The properties of elements can be generally explained in terms of electron occupancy at the highest energy level in the atom. These electrons are usually called valence electrons. Elements whose atoms have relatively few valence electrons are typically metals, while those with relatively large numbers of valence electrons are typically nonmetals. This is particularly true for the representative elements (main group elements). For example, alkali metal and alkaline earth elements are typically metals whereas oxygen family elements and halogens are typically nonmetals.
Atoms form molecules and gain stability (lowered potential energy) through covalent bonding by sharing one or more electron pairs. Electron pair sharing is typical of nonmetals and is best illustrated for the hydrogen molecule, a case of "equal sharing." Two isolated hydrogen atoms have only one electron each. As the atoms approach each other, the electron of one atom is attracted by the nucleus of the other atom, and vice versa. This mutual attraction by two nuclei for an electron pair gives rise to the covalent bond, with an excess of attraction over repulsion. Only "like" atoms (e.g., same electronegativity) form such equal-sharing, nonpolar bonds. Most covalent bonds are not "equal sharing" since there is frequently a difference in electronegativity. In these cases the bond is called a polar covalent bond. This means that the bond has a positively-charged and negatively-charged end. However, the charges are not nearly as large as the charges on ions in ionic solids. It is common practice to show the bonding in molecules in terms of Lewis-dot formulas, named in honor of G.N. Lewis. In these formulas, the elemental symbols represent the nucleus and all the electrons except valence electrons. These are represented as dots-one dot means one electron, two dots two electrons, and so on. The structures are consistent in most simple cases with the octet rule.
Now consider the opposite extreme. When a metal and nonmetal form a compound, most often the nonmetal attracts electrons more strongly than the metal (i.e., it has a larger electronegativity). In such a case the electron pair is "taken over" by the more electronegative atom to form a negatively-charged ion. The metal atom, by virtue of losing an electron, acquires a positive charge to form a positively-charged ion. This is essentially what happens between an alkaline metal and a halogen. If two electrons are "transferred" as between an alkaline earth element and an oxygen family element, then 2+ and 2- ions are formed. In such cases the bonding is called ionic bonding and the stability (lowered potential energy) is due to the mutual attraction between oppositely-charged ions in the solid crystals these compounds form.
The bonding categories considered thus far-covalent and ionic-explain the structure and properties of many substances. But metals, covalent network solids, and molecular solids are somewhat unique and require additional attention.
Bonding in metals, called metallic bonding, involves valence electrons. These electrons are loosely held by any one atom and collectively form a "sea of valence electrons" that can be used to explain many metallic properties, e.g., metallic luster, malleability, electrical conductivity, etc. The electrons are loosely held since each atom has several unoccupied valence orbitals; it is relatively easy for the electrons to move about. In this manner the electrons allow atoms to slide past each other and be "worked" (hammered) into shapes and drawn into wires (evidence of malleability and ductility). The mobile electrons in appropriate circumstances move and conduct electricity and heat. The metallic bond is the third and last of the generally-recognized types of chemical bonds.
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