Lone pair
A lone pair is a pair of valence electrons occupying an orbital on an atom that is not involved in covalent bonding and thus remains localized on that single atom, often represented by two dots in Lewis structures.[1] These nonbonding electron pairs play a crucial role in determining molecular geometry through the valence shell electron pair repulsion (VSEPR) theory, where they exert stronger repulsive forces on surrounding bonding pairs due to their higher electron density and occupation of larger spatial regions compared to shared bonding electrons.[2][3] In Lewis acid-base chemistry, lone pairs enable atoms to act as Lewis bases by donating electrons to electron-deficient species, forming coordinate covalent bonds essential for coordination compounds and reactions involving metal ions.[4] This electron donation capacity also contributes to intermolecular forces such as hydrogen bonding, where lone pairs on electronegative atoms like oxygen or nitrogen interact with hydrogen atoms bound to similar electronegative atoms, influencing properties like boiling points and solubility in molecules such as water or ammonia.[5] Furthermore, the presence and orientation of lone pairs affect molecular polarity, as they create regions of high electron density that can lead to dipole moments even in otherwise symmetric structures.[6] Beyond basic molecular structure, lone pairs are integral to advanced chemical phenomena, including their stereochemical activity in solid-state materials where they can distort coordination geometries to enhance properties like ion conductivity or ferroelectricity in compounds containing elements such as lead or bismuth.[7] In organic and inorganic synthesis, lone pairs facilitate nucleophilic attacks and resonance stabilization, underscoring their fundamental importance across chemical disciplines.[8]Fundamentals
Definition and Characteristics
A lone pair, also known as an unshared or nonbonding pair, consists of two valence electrons localized on a single atom that are not involved in covalent bonding with another atom.[9] These electrons occupy a specific orbital in the atom's valence shell, contributing to the fulfillment of the octet rule, which posits that atoms tend to achieve a stable configuration with eight valence electrons.[10] The concept of lone pairs was introduced by Gilbert N. Lewis in his 1916 seminal paper, where he described the octet rule and covalent bonding as involving shared electron pairs, with unshared pairs completing the octet on atoms such as oxygen and nitrogen.[10] Lone pairs exhibit characteristics similar to bonding pairs in that they occupy hybrid or valence orbitals, but they generate stronger repulsive forces in molecular geometry due to their higher localized electron density compared to the more delocalized density in bonding pairs.[11] This increased repulsion arises because lone pair electrons are held closer to the nucleus of the central atom, without being shared across an interatomic bond.[11] Lone pairs play a key role in calculating an atom's formal charge, a measure used to assess the electron distribution in Lewis structures. The formal charge is given by the formula: \text{Formal charge} = (\text{valence electrons}) - (\text{nonbonding electrons}) - \frac{1}{2} (\text{bonding electrons}) where nonbonding electrons include those in lone pairs. This calculation, formalized by Linus Pauling, helps identify the most stable resonance structures by minimizing formal charges. Representative examples illustrate these properties: in the water molecule (H₂O), the oxygen atom possesses two lone pairs, completing its octet alongside two bonding pairs to the hydrogen atoms; similarly, in ammonia (NH₃), the nitrogen atom has one lone pair, achieving its octet with three bonding pairs to hydrogen atoms.[10]Representation in Lewis Structures
In Lewis dot structures, lone pairs are depicted as pairs of dots positioned adjacent to the symbol of the atom possessing them, distinguishing these non-bonding electrons from bonding pairs, which are represented by lines connecting atoms. This convention, introduced by Gilbert N. Lewis in 1916, facilitates the visualization of valence electron distribution in molecules and ions.[12][13] The construction of Lewis structures follows the octet rule, particularly for elements in the second period (such as carbon, nitrogen, oxygen, and fluorine), where atoms seek to surround themselves with eight valence electrons to achieve stability akin to noble gases. After placing atoms and drawing bonds to connect them, remaining valence electrons are distributed as lone pairs to fulfill this octet for each atom, starting with terminal atoms.[14][15] To calculate the number of lone pairs on a central atom, subtract the number of electrons committed to bonding from the atom's valence electrons and divide by 2, as each lone pair consists of two electrons. For a neutral central atom with V valence electrons forming B single bonds, the formula is \frac{V - B}{2}. This method assumes standard single bonds; adjustments apply for multiple bonds or charged species.[16] Notation for lone pairs can vary slightly; while dots are standard, some organic chemistry representations use short dashes or lines for lone pairs to streamline sketches, especially in skeletal formulas where hydrogens and some lone pairs are implied. For elements beyond the second period (like phosphorus or sulfur), the octet rule can be exceeded, permitting expanded valence shells with 10, 12, or more electrons, which may accommodate additional bonding electrons and fewer or no lone pairs on the central atom.[17][18] Consider ammonia (NH₃) as an example: Nitrogen, with 5 valence electrons, forms three single bonds to hydrogen atoms (using 3 electrons from nitrogen), leaving 2 electrons as one lone pair, depicted as two dots above the nitrogen symbol in the structure H–N–H with the third H below and :: on N. Water (H₂O) follows similarly: Oxygen, with 6 valence electrons, forms two single bonds (using 2 electrons), resulting in 4 electrons forming two lone pairs, shown as :: on top and bottom of the O in H–O–H. In contrast, sulfur hexafluoride (SF₆) exhibits an expanded octet: Sulfur, with 6 valence electrons, forms six single bonds to fluorine atoms (using all 6 plus additional from fluorines, totaling 12 electrons around sulfur), yielding no lone pairs on sulfur, while each fluorine has three lone pairs to complete its octet.[13][19][18] Formal charge, computed as valence electrons minus lone pair electrons minus half the bonding electrons, serves as a tool to verify structure accuracy by minimizing charges on atoms.[20]Influence on Molecular Geometry
Valence Shell Electron Pair Repulsion Theory
The Valence Shell Electron Pair Repulsion (VSEPR) theory, introduced by Ronald J. Gillespie and Ronald S. Nyholm in 1957, serves as a foundational model for predicting molecular geometries by considering the spatial arrangement of electron pairs around a central atom. The core principle is that electron pairs in the valence shell—encompassing both bonding pairs (shared between atoms) and lone pairs (unshared on the central atom)—exert repulsive forces on one another, positioning themselves to achieve the minimum overall repulsion and thus the lowest energy configuration. This repulsion arises from the Pauli exclusion principle and electrostatic interactions among the negatively charged electron domains.[21] In VSEPR, an electron domain is defined as either a lone pair or a bonding pair (including single, double, or triple bonds, each counted as one domain). The total number of domains around the central atom, termed the steric number, dictates the basic electron pair geometry, independent of whether the domains are bonding or lone pairs. Lewis structures provide the starting point for identifying these domains. The standard geometries for different steric numbers are outlined below:| Steric Number | Electron Pair Geometry | Example Central Atom Configuration |
|---|---|---|
| 2 | Linear | Be in BeCl₂ (AX₂) |
| 3 | Trigonal planar | B in BF₃ (AX₃) |
| 4 | Tetrahedral | C in CH₄ (AX₄) |
| 5 | Trigonal bipyramidal | P in PCl₅ (AX₅) |
| 6 | Octahedral | S in SF₆ (AX₆) |