Electron pair
An electron pair in chemistry consists of two valence electrons that either share an orbital between two atoms to form a covalent bond, known as a bonding pair, or remain unshared on a single atom, referred to as a lone pair.[1] These pairs are fundamental to understanding chemical bonding and molecular structure, as they arise from the tendency of atoms to achieve stable electron configurations through sharing or possession of electrons.[1]
In covalent bonding, bonding electron pairs occupy the space between atomic nuclei, attracting them together while the electrons are repelled by each other and by the nuclei.[1] Lone pairs, in contrast, reside entirely on one atom and do not participate in bonding but influence the molecule's polarity, reactivity, and geometry by exerting repulsive forces on surrounding electrons.[1] The concept of electron pairs was pioneered by Gilbert N. Lewis in his 1916 work on the octet rule, which posits that atoms strive to surround themselves with eight valence electrons—equivalent to four electron pairs—to mimic the stable configuration of noble gases.[1]
Electron pairs play a central role in the valence-shell electron-pair repulsion (VSEPR) theory, which predicts molecular shapes by assuming that electron pairs around a central atom arrange themselves to minimize mutual repulsion.[2] In VSEPR, each bonding or lone pair constitutes an electron domain, and the geometry is determined by the number of these domains: for example, two domains yield a linear structure, while four domains result in tetrahedral electron-pair geometry.[2] Lone pairs occupy more space than bonding pairs due to less nuclear attraction, often distorting bond angles, as seen in molecules like ammonia (NH₃) with its trigonal pyramidal shape.[2]
Beyond bonding, electron pairs are key in Lewis acid-base theory, where a Lewis base donates an electron pair to a Lewis acid, which accepts it to form a coordinate covalent bond.[3] This framework, introduced by Lewis in 1923, broadens the understanding of reactions involving electron transfer, such as in coordination chemistry and catalysis, emphasizing the electron pair as the basic unit of reactivity.[3]
Definition and Fundamentals
Definition
In 1916, American chemist Gilbert N. Lewis introduced the concept of the electron pair in his seminal paper "The Atom and the Molecule," proposing that a chemical bond forms through the sharing of two electrons between atoms.[4] This idea revolutionized the understanding of atomic interactions by shifting focus from ionic models to shared electron arrangements.[5]
An electron pair is defined as two electrons that occupy the same molecular orbital or are closely associated with a single atom. These electrons typically possess opposite spins, enabling them to coexist in the same spatial region without violating fundamental quantum rules.[6]
Electron pairs play a central role in chemistry by underpinning covalent bonding, where they facilitate the attraction between atoms and promote overall molecular stability.[7] They are essential for explaining valence, as atoms tend to form structures that complete electron pairs to achieve lower energy configurations.[8]
Quantum Mechanical Basis
The quantum mechanical foundation of electron pairs lies in the Pauli exclusion principle, formulated by Wolfgang Pauli in 1925, which dictates that no two electrons in an atom or molecule can share the same set of four quantum numbers: the principal quantum number n, the azimuthal quantum number \ell, the magnetic quantum number m_\ell, and the spin quantum number m_s.[9] This principle ensures that each atomic or molecular orbital can accommodate at most two electrons.[6] For pairing to occur within the same orbital, the electrons must possess opposite spin orientations, with one having m_s = +\frac{1}{2} (spin up) and the other m_s = -\frac{1}{2} (spin down), as identical spins would violate the exclusion rule.[10]
Spin pairing in electron pairs produces a singlet state with zero net spin angular momentum, where the individual magnetic moments of the electrons cancel each other out.[11] This cancellation results in diamagnetic properties for systems with fully paired electrons, as there is no unpaired spin to interact with an external magnetic field and generate attraction.[12] In contrast, unpaired electrons exhibit a net magnetic moment due to their nonzero total spin, leading to paramagnetic behavior where the atoms or molecules are weakly attracted to magnetic fields.[12]
Electron pairs occupy orbitals in accordance with quantum rules, filling them from lower to higher energy levels while adhering to the Pauli principle. Core electron pairs reside in inner, lower-energy orbitals close to the nucleus, such as the 1s orbital in atoms, where they are tightly bound and shielded from external influences. Valence electron pairs, conversely, occupy the outermost orbitals, like the 2s and 2p in carbon, and are more reactive due to their higher energy and weaker nuclear attraction. This occupancy pattern stabilizes atomic and molecular structures by maximizing orbital filling without violating quantum constraints.[6]
Types of Electron Pairs
Bonding Pairs
A bonding pair consists of two electrons shared between the nuclei of two atoms in a covalent bond, allowing each atom to achieve a more stable electron configuration by filling its valence shell.[13] This shared pair is distinct from non-bonding electrons and forms the basis of covalent bonding in molecules composed of nonmetals.[14]
In covalent bonds, the number of bonding pairs determines the bond multiplicity: a single bond involves one bonding pair, a double bond involves two bonding pairs, and a triple bond involves three bonding pairs.[15] As the number of bonding pairs increases, so does the bond order, which correlates with greater bond strength and shorter bond lengths; for instance, triple bonds are typically the strongest and shortest among these.[16] A classic example is the hydrogen molecule (H₂), where a single bonding pair forms between the two hydrogen atoms, resulting in a bond order of one.[17] Similarly, in the oxygen molecule (O₂), two bonding pairs create a double bond, yielding a bond order of two.[18]
Bonding pairs directly contribute to the bond order of a molecule, defined as the number of shared electron pairs between bonded atoms, which enhances overall molecular stability by increasing the energy required to break the bond.[16] Higher bond orders from multiple bonding pairs thus promote greater molecular integrity, as seen in diatomic molecules where the accumulation of bonding pairs lowers the system's potential energy relative to separated atoms.[19]
Lone Pairs
A lone pair, also known as a non-bonding pair, consists of two electrons localized on a single atom that are not involved in covalent bonding with other atoms. These electrons occupy the valence shell of the atom and remain paired due to their mutual repulsion and the stability provided by the Pauli exclusion principle.[20]
In the valence shell, lone pairs reside in hybrid or atomic orbitals surrounding the central atom, contributing to the overall electron density and influencing the atom's reactivity. Unlike bonding pairs, which are shared between atoms, lone pairs are fully associated with one nucleus, leading to higher local electron density that can participate in coordination chemistry or hydrogen bonding. For instance, the nitrogen atom in ammonia (NH₃) possesses one lone pair in an sp³ hybrid orbital, while the oxygen atom in water (H₂O) has two lone pairs, both in sp³ orbitals, which enhance the molecule's polarity.[21][22]
Lone pairs occupy a larger effective volume in the valence shell compared to bonding pairs because they experience repulsion solely from the host atom's nucleus, without the counterbalancing attraction from another atom's nucleus. This expanded spatial requirement arises from the electrons being delocalized over a broader region around the atom, resulting in stronger repulsive interactions with adjacent electron domains.[23]
Representation in Chemical Models
Lewis Structures
Lewis dot symbols, introduced by Gilbert N. Lewis in 1916, represent the valence electrons of an atom as dots surrounding its elemental symbol, providing a visual depiction of the electrons available for bonding.[4] These symbols illustrate the octet tendency, where atoms seek to achieve eight valence electrons, akin to noble gas configurations, through sharing or lone pairs.[24]
Lewis structures extend this notation to molecules and ions by showing the arrangement of valence electron pairs, distinguishing between bonding pairs (shared between atoms) and lone pairs (unshared on a single atom).[25] The primary rules for constructing these structures involve calculating the total number of valence electrons from all atoms (adjusted for charge in ions), pairing electrons into single bonds between atoms, and distributing the remainder as lone pairs or multiple bonds to satisfy the octet rule for as many atoms as possible, particularly second-period elements.[24] Formal charges are then assessed to evaluate structure validity, calculated as valence electrons minus non-bonding electrons minus half of bonding electrons.[25]
The step-by-step process for drawing Lewis structures begins with determining the total valence electrons by summing contributions from each atom and adding electrons for negative charges or subtracting for positive ones.[26] Next, arrange the atoms with the least electronegative (often central) atom in the middle and form single bonds between it and surrounding atoms using two electrons per bond.[25] Distribute the remaining electrons as lone pairs to complete octets, starting with terminal atoms; if electrons are insufficient, convert lone pairs into multiple bonds (double or triple) on the central atom.[24] Finally, compute formal charges on each atom and select the structure with the lowest charges or most stable distribution, allowing for expanded octets (more than eight valence electrons) for elements beyond the second period, particularly central atoms in hypervalent molecules.[26]
For methane (CH₄), the carbon atom serves as the central atom with four valence electrons, each hydrogen contributes one; the total of eight electrons forms four bonding pairs as single C–H bonds, resulting in no lone pairs and all atoms achieving their octet or duet.[24]
H
|
H-C-H
|
H
H
|
H-C-H
|
H
In carbon dioxide (CO₂), the total valence electrons are 16; oxygen atoms are terminal, connected to central carbon with double bonds (O=C=O), each oxygen bearing two lone pairs to complete its octet, while carbon achieves octet through the four bonding pairs.[25]
.. ..
:O::C::O:
.. ..
.. ..
:O::C::O:
.. ..
Molecular Orbital Description
In molecular orbital (MO) theory, atomic orbitals from constituent atoms combine through linear combinations to form molecular orbitals that extend over the entire molecule, providing a delocalized description of electrons rather than localized pairs around atoms.[27] These molecular orbitals are classified as bonding (lower energy, stabilizing the molecule), antibonding (higher energy, destabilizing), or non-bonding (energy similar to atomic orbitals, neither stabilizing nor destabilizing). Each molecular orbital can accommodate a pair of electrons with opposite spins, adhering to the Pauli exclusion principle, and orbitals fill according to the Aufbau principle and Hund's rule.[28] This framework, developed by Robert S. Mulliken and others in the 1930s, treats electron pairs as occupying these delocalized orbitals to explain bonding and molecular stability.[29]
Bonding electron pairs occupy bonding molecular orbitals, contributing to covalent bonds such as sigma (σ) and pi (π) bonds. Sigma bonds form from the end-to-end overlap of atomic orbitals (e.g., s-s or p-p along the internuclear axis), creating a bonding σ orbital with maximum electron density between nuclei, while pi bonds arise from the side-to-side overlap of p orbitals, resulting in a bonding π orbital with electron density above and below the bond axis.[28] These paired electrons in bonding orbitals lower the overall energy of the molecule compared to separated atoms. Lone pairs, in contrast, occupy non-bonding molecular orbitals, which do not significantly alter the bond order and are localized primarily on individual atoms, akin to the lone pairs in Lewis structures but described quantum mechanically as delocalized over the atomic region.[30]
A simple example is the molecular orbital diagram for the hydrogen molecule (H₂), the simplest diatomic homonuclear molecule. The 1s atomic orbitals from each hydrogen atom combine to form a σ_{1s} bonding orbital (lower energy) and a σ^*{1s} antibonding orbital (higher energy); the two valence electrons occupy the σ{1s} orbital as a bonding pair with opposite spins, yielding a bond order of 1 and a stable single bond. This pattern extends to other homonuclear diatomic molecules like N₂, where multiple bonding pairs fill σ and π orbitals to form a triple bond, while the electrons corresponding to the lone pairs occupy nearly non-bonding molecular orbitals derived primarily from the atomic 2s orbitals.[28]
Theoretical Frameworks
Valence Shell Electron Pair Repulsion Theory
The Valence Shell Electron Pair Repulsion (VSEPR) theory posits that the geometry of a molecule is determined by the repulsion between electron pairs in the valence shell of the central atom, which arrange themselves to minimize this repulsion and achieve the lowest energy configuration. Introduced by Ronald J. Gillespie and Ronald S. Nyholm, the model treats both bonding pairs (shared between atoms) and lone pairs (unshared on the central atom) as electron domains that exert repulsive forces on one another, with the strength of repulsions following the order: lone pair–lone pair > lone pair–bonding pair > bonding pair–bonding pair. This arrangement maximizes the separation between domains, leading to predictable spatial orientations.
The electron domain geometry refers to the arrangement of all electron pairs around the central atom, independent of whether they are bonding or lone pairs. For a central atom with two electron domains, the geometry is linear (bond angle 180°), as seen in molecules like CO₂. Three domains result in a trigonal planar arrangement (bond angles 120°), while four domains form a tetrahedral geometry (bond angles 109.5°). Five domains adopt a trigonal bipyramidal structure (axial angles 90°, equatorial 120°), and six domains yield an octahedral arrangement (all angles 90°). These ideal geometries assume equivalent repulsion and provide the foundational framework for molecular shape prediction.
Molecular geometry, however, is the arrangement of only the atoms (excluding lone pairs), which can distort the electron domain geometry due to the larger spatial requirements of lone pairs. In notation such as AX₃E (where A is the central atom, X a surrounding atom, and E a lone pair), three bonding pairs and one lone pair lead to a trigonal pyramidal molecular shape, as in ammonia (NH₃), where the lone pair occupies a tetrahedral position, compressing bond angles to approximately 107°. Similarly, for AX₂E₂, two bonding pairs and two lone pairs result in a bent geometry, exemplified by water (H₂O), with a bond angle of about 104.5° deviating from the ideal tetrahedral 109.5°. In contrast, BF₃ (AX₃) exhibits a trigonal planar molecular geometry with no lone pairs on boron, maintaining 120° bond angles. These distortions arise because lone pairs repel bonding pairs more strongly, pushing atoms closer together.[31]
| Number of Electron Domains | Electron Domain Geometry | Example Molecule (Molecular Geometry if Applicable) |
|---|
| 2 | Linear | CO₂ (linear) |
| 3 | Trigonal planar | BF₃ (trigonal planar) |
| 4 | Tetrahedral | CH₄ (tetrahedral); NH₃ (trigonal pyramidal, AX₃E); H₂O (bent, AX₂E₂) |
| 5 | Trigonal bipyramidal | PCl₅ (trigonal bipyramidal) |
| 6 | Octahedral | SF₆ (octahedral) |
Octet Rule
The octet rule is a fundamental guideline in chemistry stating that atoms of main group elements tend to gain, lose, or share electrons to achieve eight valence electrons in their outermost shell, equivalent to four electron pairs.[32] This principle, introduced by Gilbert N. Lewis in his 1916 paper on atomic structure and bonding, provides a simple model for predicting the formation of stable molecules through covalent or ionic interactions. It emphasizes the role of electron pairs in achieving chemical stability, where shared pairs form bonds and unshared pairs remain as lone pairs on atoms.
The rationale for the octet rule lies in the exceptional stability of noble gas electron configurations, which feature a filled valence shell of eight electrons (ns²np⁶ for elements beyond helium).[32] Atoms of other elements react to emulate this inert configuration because it minimizes energy and maximizes electron repulsion balance in the valence shell. For instance, second-period elements like carbon and oxygen commonly form compounds where each atom attains this octet, leading to predictable bonding patterns in simple molecules.
Despite its utility, the octet rule admits several exceptions that highlight its limitations as a heuristic rather than a strict law. Odd-electron molecules, such as nitric oxide (NO), possess an unpaired electron and thus cannot achieve octets for all atoms, resulting in radical character.[33] Incomplete octets arise in electron-deficient compounds like boron trifluoride (BF₃), where the central boron atom is surrounded by only six valence electrons due to its limited capacity for bonding.[33] Conversely, expanded octets occur in hypervalent molecules such as sulfur hexafluoride (SF₆), where the central sulfur holds 12 valence electrons, facilitated by d-orbital involvement in larger atoms.[33]
The octet rule applies primarily to main group elements (groups 1, 2, and 13–18), where valence electrons occupy only s and p orbitals.[32] For transition metals, adherence varies significantly due to the presence of d orbitals, often leading instead to the 18-electron rule in coordination compounds for analogous stability. In Lewis structures, these electron pair arrangements visually depict how atoms approach or deviate from the octet to form stable entities.[32]
Applications in Chemistry
Molecular Geometry and Polarity
The arrangement of electron pairs around a central atom, governed by the Valence Shell Electron Pair Repulsion (VSEPR) theory, dictates the three-dimensional molecular geometry, which in turn determines whether a molecule exhibits a net dipole moment and is thus polar or nonpolar.[34] In VSEPR, bonding pairs and lone pairs repel each other, positioning themselves to minimize repulsion and thereby defining the overall shape.[35] This geometric configuration is crucial because it affects the distribution of electron density and the vector sum of individual bond dipoles.
Molecular polarity arises from the asymmetry in charge distribution: in symmetric geometries, individual bond dipoles cancel out, resulting in a nonpolar molecule, whereas asymmetric geometries lead to a net dipole moment and polar character.[36] For instance, carbon dioxide (CO₂) adopts a linear geometry with two bonding pairs arranged symmetrically around the central carbon atom, causing the polar C=O bonds to point in opposite directions and cancel, yielding a nonpolar molecule overall.[37] In contrast, water (H₂O) has a bent geometry due to two bonding pairs and two lone pairs on the oxygen atom, creating an asymmetric arrangement where the O-H bond dipoles do not fully cancel, producing a polar molecule with a bent bond angle of approximately 104.5°.[38]
Further examples illustrate the role of electron pair arrangements in polarity. Methane (CH₄) features a tetrahedral geometry from four bonding pairs around carbon, with high symmetry that cancels all C-H bond dipoles, making it nonpolar.[39] Ammonia (NH₃), however, exhibits a trigonal pyramidal geometry due to three bonding pairs and one lone pair on nitrogen, introducing asymmetry that results in a net dipole moment and polar nature, with a bond angle of about 107°.[40] Lone pairs contribute to this by occupying more space and distorting the electron density away from bonding regions, often enhancing overall molecular polarity in asymmetric structures.[41]
Reactivity and Bonding
Electron pairs play a central role in determining chemical reactivity, particularly through the behavior of lone pairs and bonding pairs in various reaction mechanisms. Lone pairs on nucleophilic species, such as the oxygen in hydroxide ions or nitrogen in ammonia, act as electron donors in bimolecular nucleophilic substitution (SN2) reactions, where they attack the electrophilic center of an alkyl halide from the backside, leading to inversion of configuration and displacement of the leaving group.[42][43] In contrast, bonding pairs in covalent bonds undergo cleavage during reactions: homolytic cleavage evenly divides the shared electron pair, producing two radicals each with an unpaired electron, while heterolytic cleavage assigns the entire pair to one fragment, generating a cation and anion./09%3A_Free_Radical_Substitution_Reaction_of_Alkanes/9.01%3A_Homolytic_and_Heterolytic_Cleavage)[44]
The strength of bonds involving electron pairs is quantified by bond dissociation energies (BDEs), which increase with the number of shared pairs; for instance, a C-C single bond (one pair) has a BDE of approximately 346 kJ/mol, a C=C double bond (two pairs) 602 kJ/mol, and a C≡C triple bond (three pairs) 835 kJ/mol, reflecting greater stability and higher energy required for homolytic cleavage./Chemical_Bonding/Fundamentals_of_Chemical_Bonding/Bond_Energies)[45] This trend underscores how multiple bonding pairs enhance bond robustness, influencing reaction pathways that favor breaking weaker single bonds over multiple bonds. In radical reactions, unpaired electrons drive reactivity differently from paired electrons, as radicals form via homolytic cleavage and propagate chains through single-electron transfers, contrasting with the pair-based mechanisms in ionic or polar processes./16%3A_Radical_Reactions/16.02%3A_Overview_of_Single-Electron_Reactions_and_Free_Radicals)[46]
Electron pair availability is a key determinant of Lewis acidity and basicity, where bases donate lone pairs to acids that accept them, forming coordinate covalent bonds as seen in acid-base reactions like ammonia (NH₃) with boron trifluoride (BF₃), where nitrogen's lone pair coordinates to boron's empty orbital.[47][48] Greater accessibility of lone pairs, as in less sterically hindered or more electronegatively favorable species, enhances basicity, while electron-deficient sites with available orbitals increase acidity./Acids_and_Bases/Acid/Lewis_Concept_of_Acids_and_Bases) This framework explains reactivity in coordination chemistry and extends to fulfill octet requirements in transient reaction intermediates.