Mercury (Hg), known as quicksilver, is the only metal that exists as a liquid at standard temperature and pressure. This heavy, silvery element has an atomic number of 80. Understanding the arrangement of its outermost electrons is key to explaining Mercury’s distinct behavior compared to other metals. The specific number of these valence electrons dictates how Mercury forms bonds and interacts with other substances.
The Role of Valence Electrons in Chemistry
Valence electrons are the electrons located in the outermost shell or energy level of an atom. These electrons are the primary participants in all chemical bonding, including the formation of ionic and covalent compounds. The number and arrangement of these electrons govern an element’s ability to gain, lose, or share electrons with other atoms. Atoms react to achieve a stable configuration, often by having a complete set of eight valence electrons, known as the octet rule.
Calculating Mercury’s Valence Electrons
Mercury is a member of Group 12 on the periodic table, classified as a d-block transition metal. Its atomic number, 80, means a neutral Mercury atom contains 80 electrons. Determining the valence electrons requires examining its ground state electron configuration.
The condensed electron configuration for Mercury is \([Xe] 4f^{14} 5d^{10} 6s^2\). The largest principal quantum number identifies the outermost energy level, which is the sixth shell (\(n=6\)). This shell contains the valence electrons.
Specifically, the \(6s^2\) orbital holds the two electrons in this outermost shell. Although the \(5d^{10}\) and \(4f^{14}\) orbitals are chemically relevant, they are considered filled inner shells when determining the standard valence electron count. Since these \(d\) and \(f\) orbitals are stable and completely full, they are counted as core electrons.
Therefore, Mercury has two valence electrons, residing in the \(6s\) orbital. This count of two is consistent with its position in Group 12 of the periodic table, alongside Zinc and Cadmium.
How Valence Electrons Influence Mercury’s Reactivity and Oxidation States
The two valence electrons in the \(6s\) orbital are readily lost, leading to Mercury’s most common and stable oxidation state, \(\text{Hg}^{2+}\). Losing both \(s\) electrons achieves a stable configuration with a filled \(5d^{10}\) subshell. This stability accounts for the predominance of mercuric compounds and makes Mercury less chemically reactive than Group 2 elements, which also have two valence electrons.
Mercury also exhibits a less stable \(+1\) oxidation state, which exists as the dimeric mercurous ion, \(\text{Hg}_2^{2+}\). This structure results from the partial sharing of the two \(6s\) valence electrons, forming a covalent bond between two Mercury atoms. The resulting \(\text{Hg}_2^{2+}\) ion acts as a single unit with an overall \(+2\) charge.
Relativistic Effects and the Liquid State
The tight binding of the \(6s\) valence electrons to the nucleus contributes to Mercury’s low melting point and liquid state. This phenomenon is enhanced by relativistic effects in heavy atoms. This strong attraction means the electrons are less available to participate in the metallic bonding that typically holds solid metals together, resulting in weak interatomic forces.