A chemical compound is a substance formed when two or more different chemical elements are chemically bonded together. These combinations create entirely new substances with properties distinct from the original elements. Atoms join to form compounds to achieve a state of greater stability. This stability is attained through attractive forces, known as chemical bonds, which hold the atoms together in a defined arrangement.
The Driving Force: Chemical Bonds
At the heart of chemical bonding lies the behavior of electrons, particularly those in the outermost shell of an atom, called valence electrons. Atoms interact to achieve a more stable electron configuration, often resembling the full outer electron shells of noble gases. This tendency drives atoms to either gain, lose, or share valence electrons with other atoms. These attractive forces hold atoms together within a compound, giving it unique identity and properties.
Ionic Bonds
Ionic bonds typically form between a metal and a non-metal atom, driven by a significant difference in their attraction for electrons. One atom completely transfers one or more valence electrons to another, becoming a positively charged ion (cation), while the other atom gains electrons and becomes a negatively charged ion (anion). These oppositely charged ions are strongly attracted through electrostatic forces, holding them together in a crystal lattice structure. A common example is table salt, sodium chloride (NaCl), where sodium donates an electron to chlorine, forming Na+ and Cl- ions. Compounds with ionic bonds often exhibit high melting and boiling points, reflecting the strength of these electrostatic attractions.
Covalent Bonds
Covalent bonds form when two atoms share one or more pairs of valence electrons, typically occurring between two non-metal atoms. This sharing allows both atoms to achieve a stable electron configuration, creating a strong attractive force that holds the atoms together. Depending on the number of electron pairs shared, covalent bonds can be classified as single, double, or triple bonds, as seen in water (H2O) with single bonds or carbon dioxide (CO2) with double bonds. Covalent compounds generally have lower melting and boiling points compared to ionic compounds, and their properties are influenced by the specific arrangement and sharing of electrons.
Metallic Bonds
Metallic bonds are found in metals, where valence electrons are delocalized and shared among a lattice of positively charged metal ions. This arrangement is often described as a “sea of electrons” that are free to move throughout the entire metallic structure. This unique electron mobility is responsible for many characteristic properties of metals, explaining why they are excellent conductors of electricity and heat. Their malleability (ability to be hammered into thin sheets) and ductility (ability to be drawn into wires) also stem from this electron sea model. Examples include copper and iron; alloys, which are mixtures of metals, also rely on metallic bonding to hold their constituent atoms together.
Intermolecular Forces: Holding Compounds Together as a Whole
While chemical bonds (ionic, covalent, metallic) hold atoms within a compound, intermolecular forces are weaker attractions between separate molecules. These forces are electrostatic attractions between existing molecules, not involving electron sharing or transfer. Common types include van der Waals forces, such as London dispersion forces (temporary attractions from instantaneous electron shifts) and dipole-dipole interactions (occurring between polar molecules with permanent partial positive and negative charges). Hydrogen bonding is a strong type of dipole-dipole interaction, occurring when hydrogen is bonded to highly electronegative atoms like oxygen or nitrogen. These forces influence a compound’s physical properties, such as its melting point, boiling point, and state at room temperature; for example, water’s high boiling point is attributed to strong hydrogen bonds.