Atoms are the fundamental building blocks of all matter, composed of a dense nucleus (containing protons and neutrons) surrounded by a cloud of negatively charged electrons. When atoms interact, they engage in chemical bonding to form new substances, such as molecules or compounds. This combination process is driven by the behavior of the outermost electrons and results in the vast diversity of materials observed in the physical world.
Why Atoms Seek Stability
Atoms combine to achieve a lower, more stable energy state, a condition governed by their outermost, or valence, electrons. These electrons determine an atom’s reactivity. For most elements, maximum stability is achieved when this shell contains eight electrons, a principle known as the Octet Rule. Atoms attempt to gain, lose, or share electrons to reach this configuration, mirroring the electron arrangement of the noble gases.
The lightest elements, particularly hydrogen and helium, follow the Duet Rule, seeking stability with two valence electrons in their first shell. Atoms with incomplete shells readily engage in chemical reactions to minimize their overall energy. The change in energy that occurs during bond formation, typically a release of energy, signifies the creation of a more stable system.
The Exchange of Electrons (Ionic Bonds)
Atoms achieve stability through the complete transfer of one or more valence electrons, resulting in an ionic bond. This bonding occurs between a metal and a non-metal, which have greatly differing tendencies to hold electrons. The metal atom readily gives up its electron(s) to empty its outer shell, becoming a positively charged ion called a cation.
Conversely, the non-metal atom accepts the transferred electron(s) to complete its octet, becoming a negatively charged ion called an anion. For example, a sodium atom transfers its single valence electron to a chlorine atom, creating \(text{Na}^{+}\) and \(text{Cl}^{-}\). The resulting bond is purely an electrostatic attraction that holds the oppositely charged ions together.
The Sharing of Electrons (Covalent Bonds)
When two non-metal atoms combine, they often have a similar attraction for electrons, making a full transfer unfavorable. Instead, they form a covalent bond by sharing one or more pairs of valence electrons between their nuclei. This shared electron density satisfies the stability requirements of both atoms, allowing them to complete their outer electron shells. The number of electron pairs shared determines the type of bond formed: single (one pair), double (two pairs), or triple (three pairs).
For instance, the nitrogen molecule (\(text{N}_{2}\)) is held together by a triple bond, where six electrons are shared. Polarity arises when one atom exerts a stronger pull on the shared electrons than the other, a difference known as electronegativity. This causes the electron density to be unevenly distributed, creating a partial negative charge near the more attractive atom and a partial positive charge near the less attractive one.
How Bond Type Determines Structure
The difference in how ionic and covalent bonds are formed leads to major differences in the resulting structures and properties. Ionic compounds are not composed of individual molecules but form an extended three-dimensional arrangement known as a crystal lattice. In this rigid structure, ions are held together by powerful, continuous electrostatic forces.
This strong attraction throughout the lattice requires significant energy to break, giving ionic compounds characteristic properties such as very high melting and boiling points; sodium chloride, for example, melts at 801 °C. Covalent compounds, however, typically form discrete, separate molecules, like water (\(text{H}_{2}text{O}\)). The forces holding the atoms together within the molecule are strong, but the forces between separate molecules are comparatively weak. Consequently, covalent compounds tend to have much lower melting and boiling points and can exist as gases, liquids, or soft solids at room temperature.