Chemical bonds represent the attractive forces holding atoms together to form molecules and compounds. Atoms interact to achieve a more stable electron configuration, typically by filling their outermost electron shells. The way atoms achieve this stability, primarily through interactions involving their electrons, determines the type of bond formed.
Ionic Bonds
Ionic bonds typically form between a metal atom and a non-metal atom. This bond involves a complete transfer of one or more electrons from the metal atom to the non-metal atom. The metal atom, having lost electrons, becomes a positively charged ion (cation), while the non-metal atom gains these electrons, becoming a negatively charged ion (anion). The strong electrostatic attraction between these oppositely charged ions constitutes the ionic bond. For example, in sodium chloride (table salt), a sodium atom donates an electron to a chlorine atom, forming Na+ and Cl- ions held together by this attraction.
Ionic compounds possess distinctive properties. They generally exhibit high melting and boiling points due to the significant energy required to overcome the strong electrostatic forces holding the ions in a crystalline lattice. Many ionic compounds can conduct electricity when molten or dissolved in water, as their ions become mobile and can carry an electric charge.
Covalent Bonds
Covalent bonds primarily form between two non-metal atoms. Instead of transferring electrons, atoms in a covalent bond achieve stability by sharing electrons. This sharing allows each atom to effectively complete its outer electron shell.
The number of electron pairs shared between atoms can vary. A single bond involves the sharing of one pair of electrons, a double bond involves two shared pairs, and a triple bond consists of three shared electron pairs. The sharing of electrons can sometimes be unequal, leading to a slight charge separation within the molecule, a concept known as polarity. This unequal sharing occurs when one atom attracts the shared electrons more strongly than the other, resulting in partial positive and negative charges on the atoms.
Covalent compounds generally display different characteristics compared to ionic compounds. They often have lower melting and boiling points because the intermolecular forces, which are attractions between separate molecules, are typically weaker than the electrostatic forces within ionic compounds. Most covalent compounds do not conduct electricity, as their electrons are localized within the bonds and are not free to move. Common examples include water (H2O), where oxygen shares electrons with two hydrogen atoms, and carbon dioxide (CO2), which features double bonds between carbon and oxygen.
Metallic Bonds
Metallic bonds occur exclusively within metals. In this bonding arrangement, metal atoms release their outermost valence electrons, which then become delocalized. These electrons are not associated with any single atom or bond but rather form a mobile “sea” that moves freely throughout the entire metallic structure.
The metallic bond is the attraction between the positively charged metal ions (consisting of the nucleus and inner electrons) and this surrounding sea of negatively charged, delocalized electrons. This electron sea model explains many characteristic properties of metals. For instance, the free movement of electrons accounts for the electrical and thermal conductivity observed in metals.
The delocalized electron sea also provides metals with their characteristic malleability, allowing them to be hammered into thin sheets without breaking, and ductility, enabling them to be drawn into wires. The interaction of light with the mobile electrons contributes to the lustrous appearance of metals. Examples of substances held together by metallic bonds include common elements such as copper, iron, and gold.