Chemical Bonding
Introduction
Atoms rarely exist as independent particles in nature, but instead are made up of combinations of atoms that are held together with chemical bonds. A chemical bond is an attraction between the nuclei and valence electrons of different atoms. When atoms bond, their valence electrons are rearranged in ways that make the atoms more stable. Atoms will gain, lose, or share electrons to fill their outermost level with eight electrons
(Octet rule), and have the electron configuration of a noble gas. This rearrangement of electrons determines their type of bonding.
Types of Bonding
Ionic Bonds
Ionic bonds are formed between positive and negative ions that combine so that their outermost level is filled.
In other words, one ion will lose electrons and another ion will gain electrons so that the numbers of positive and negative charges are equal. For example, a group 1 metal, with one valence electron, will combine with a group 17 nonmetal, with 7 valence electrons, to fulfill their octet and form an ionic compound.
The chemical formulas of ionic compounds show the simplest ratios of the compound’s combined ions that make the compound neutral. This is called the formula unit. For example, in the ionic compound calcium fluoride, two fluoride anions (F ), each with a charge of 1- must balance the 2+ charge of each calcium cation
(Ca 2+ ). See the figure to the right.
Covalent Bonds
Covalent bonds are different than ionic bonds in that they result from the sharing of electron pairs between atoms. Atoms are still trying to achieve the octet rule, but instead of losing or gaining electrons they share them instead. These types of bonds occur between nonmetals and form covalent compounds. The smallest unit of a covalent compound is the molecule. A molecule is a neutral group of atoms that are held together.
Covalent bonds can be further classified as polar covalent or nonpolar covalent bonds so as to describe the sharing of their electrons. Equal sharing of electrons results in nonpolar covalent bonds, and unequal sharing

of electrons results in polar covalent bonds. Polarity comes from an atom’s electronegativity. Think back to this past unit, electronegativity was a measure of the ability of an atom to attract electrons. Who loved the electrons the most? nonpolar polar
The difference in electronegativity of the atoms determines the bonds they are going to form. The electronegativity of the less electronegative element is subtracted from that of the more electronegative element. The difference is then used to determine the type of bond. The following scale illustrates the breakdown when the difference in electronegativity is calculated:
0 - 0.3 0.31 - 1.69 1.7 – 4.0
Non-polar covalent polar-covalent
Comparing Ionic and Covalent Compounds
Ionic
The force that holds ions together in ionic compounds is a very strong overall attraction between positive and negative charges. In a covalent compound, even though the bonds are strong, they are much weaker than in an ionic compound. This difference gives rise to different properties in the two types of compounds. When an ionic compound is dissolved in water it will conduct electricity and a covalent compound will not. Another difference is that ionic compounds generally melt and boil at much higher temperatures than covalent compounds. Ionic compounds tend to be hard because the bonds holding them together make it difficult for one layer to move relative to another.
Metallic Bonds
Chemical bonding is different in metals than it is in ionic or covalent compounds. This difference is reflected in the unique properties of metals. The highest energy levels of most metal atoms are occupied by very few electrons. The many “vacant” orbitals allow the atom’s outer electrons to roam freely throughout the metal.
These mobile electrons form a sea of electrons around the metal atoms. The chemical bonding that result from the attraction between metal atoms and surrounding sea of electrons is called metallic bonding. See figure below.
The freedom of electrons to move in a network of metal atoms accounts for the high conductivity characteristic of all metals. In addition, metals can absorb a wide range of light frequencies which results in in the excitation of the metal atoms’ electrons to higher energy levels. When the electrons drop back down to their ground state they emit light which gives metals their luster (shine). Most metals are also easy to form into desired shapes. Two important properties related to this characteristic are malleability (hammered into sheets) and ductility (pulled to produce a wire). The malleability and ductility of metals are possible because metallic bonding is the same in all directions throughout the solid.