Chemical bonding is what holds atoms together to form molecules, compounds, and materials. Understanding bonds is fundamental to understanding why some substances are hard and brittle, others are soft and flexible, and still others conduct electricity. The type of bond determines the physical and chemical properties of the substance, from melting point to solubility to electrical conductivity.
Ionic Bonds
Ionic bonds form when one atom transfers electrons to another, creating oppositely charged ions that attract each other. This typically happens between metals (which lose electrons easily) and nonmetals (which gain electrons easily). Sodium chloride is the classic example: sodium loses one electron to become Na+, chlorine gains one to become Cl-, and the electrostatic attraction between them holds the crystal together. Ionic compounds have high melting points, are brittle, and conduct electricity when dissolved in water or melted (but not as solids, because the ions are locked in place). Use our Nomenclature Helper to look up common ionic compounds and their properties.
Covalent Bonds
In covalent bonding, atoms share electrons rather than transferring them. This is typical between nonmetals. A single covalent bond involves sharing one pair of electrons (H2), a double bond shares two pairs (O2), and a triple bond shares three pairs (N2). Covalent compounds generally have lower melting points than ionic compounds, are often liquids or gases at room temperature, and do not conduct electricity. The Bond Energy Calculator can estimate reaction enthalpies by comparing the energy needed to break bonds with the energy released when new bonds form.
Metallic Bonds
Metallic bonding is different from both ionic and covalent bonds. In a metal, the valence electrons are not bound to individual atoms but form a "sea of electrons" that is shared by all the metal cations. This delocalized electron model explains why metals are good conductors of electricity and heat, why they are malleable and ductile, and why they have characteristic luster. The strength of metallic bonding varies widely — tungsten has an extremely high melting point due to strong metallic bonds, while mercury is a liquid at room temperature.
Electronegativity and Bond Type
The electronegativity difference between two bonded atoms predicts the bond type. If the difference is greater than about 1.7, the bond is considered ionic. Between 0.4 and 1.7, it is polar covalent. Below 0.4, it is essentially nonpolar covalent. This is a simplified view — bond character exists on a continuum, not in rigid categories. Even "ionic" compounds have some covalent character, and "covalent" bonds between very different atoms have significant ionic character.