Sodium (Na) is a silvery-white alkali metal and one of the most common metals on Earth. It is familiar as sodium chloride (common table salt) and is a fundamental component in biological systems and industrial processes. As a highly reactive element, sodium readily participates in chemical reactions to achieve a stable electronic state. The number of chemical bonds sodium typically forms is determined by the arrangement of electrons surrounding its nucleus.
Sodium’s Atomic Setup and the Drive for Stability
The sodium atom has an atomic number of 11, containing 11 protons and 11 electrons. Its electron configuration is \(text{1s}^2text{2s}^2text{2p}^6text{3s}^1\). The first two inner shells are full, but the outermost third shell contains only a single electron in the \(text{3s}\) orbital.
This lone electron is the valence electron and dictates sodium’s chemical interactions. Atoms strive to achieve an outer shell of eight electrons (the Octet Rule) for stability. For sodium, the most direct path is losing this single valence electron, leaving the underlying shell as the new, full outer shell. This makes sodium extremely reactive and dictates its bonding behavior.
The Standard Answer: Forming a Single Ionic Bond
In the vast majority of its compounds, sodium achieves stability by surrendering its single \(text{3s}\) valence electron. The loss of this electron results in the formation of the positively charged sodium cation, \(text{Na}^{+}\). This ion possesses 11 protons and 10 electrons, giving it a net charge of \(+1\). Since sodium only has one electron available for transfer, its capacity for forming a chemical bond is restricted to one.
This single electron transfer leads to the formation of an ionic bond, which is a powerful electrostatic attraction. The \(text{Na}^{+}\) cation is drawn to a negatively charged ion (anion), such as the chloride ion (\(text{Cl}^{-}\)) in table salt. The energy required to remove a second electron from the stable \(text{Na}^{+}\) ion is prohibitively high, confirming that sodium is virtually never found as a \(text{Na}^{2+}\) ion.
Sodium in Different Environments: Metallic Structure and Coordination
While the single ionic bond defines sodium’s chemistry with non-metals, the concept of “bonding” changes when considering pure sodium metal. In a solid block of sodium, the bonding is metallic, a non-directional force where the single valence electrons from all atoms are pooled together to form a “sea” of delocalized electrons. This electron sea surrounds positive \(text{Na}^{+}\) cores, holding the metal structure together without forming discrete, localized bonds.
In this metallic structure, each sodium atom is surrounded by eight nearest neighbors in a body-centered cubic arrangement, giving it a coordination number of eight. A different structural arrangement is seen when the \(text{Na}^{+}\) ion is dissolved in water or crystallized in a salt like \(text{NaCl}\). In solid \(text{NaCl}\), each \(text{Na}^{+}\) ion is structurally surrounded by six \(text{Cl}^{-}\) ions, a coordination number of six. This coordination number describes the number of neighboring ions arranged around the central ion in a crystal lattice, which is distinct from the single ionic bond that formed the \(text{Na}^{+}\) ion.