Chemical Bonding and Molecular Structure

The One Thing

Chemical bonding arises because atoms combine to form arrangements of lower energy and greater stability, commonly through attainment of a noble-gas-like valence-shell configuration. The nature, geometry, polarity, strength, and magnetic behaviour of the resulting species depend on electron transfer or sharing, electron-pair repulsions, orbital overlap, resonance, and molecular-orbital occupancy.

Reactions, Processes and Experiments

Key Terms

• Chemical bond: An attractive force that holds atoms or ions together in a molecule or compound.
• Valence electrons: Electrons in the outermost shell of an atom that participate in chemical bonding.
• Octet rule: Atoms generally tend to gain, lose, or share electrons to obtain eight electrons in their valence shell, although exceptions exist.
• Ionic bond: The electrostatic attraction between oppositely charged ions formed after electron transfer, commonly from a metal to a non-metal.
• Covalent bond: A bond formed by the mutual sharing of one or more pairs of electrons between atoms.
• Coordinate bond: A covalent bond in which both shared electrons are donated by the same atom.
• Lewis structure: A diagram showing valence electrons, shared electron pairs, lone pairs, and bonds in a species.
• Formal charge: The hypothetical charge on an atom in a Lewis structure, calculated as FC = valence electrons − non-bonding electrons − one-half of bonding electrons.
• Electronegativity: The tendency of an atom in a molecule to attract the shared electron pair toward itself.
• Bond polarity: Unequal sharing of bonding electrons caused by a difference in electronegativity between bonded atoms.
• Dipole moment: A measure of charge separation in a bond or molecule; for a simple bond, μ = q × r.
• VSEPR theory: The Valence Shell Electron Pair Repulsion theory, which states that electron pairs around a central atom arrange themselves as far apart as possible to minimise repulsion.
• Bond pair: A pair of electrons shared between two atoms and involved in bonding.
• Lone pair: A pair of valence electrons located on one atom that does not participate directly in bonding.
• Repulsion order: Electron-pair repulsions generally follow the order lone pair-lone pair > lone pair-bond pair > bond pair-bond pair.
• Hybridisation: The intermixing of atomic orbitals of comparable energy on the same atom to form equivalent hybrid orbitals.
• sp hybridisation: Mixing of one s and one p orbital to form two hybrid orbitals with linear geometry and an angle of 180 degrees, as in BeCl2.
• sp2 hybridisation: Mixing of one s and two p orbitals to form three hybrid orbitals with trigonal planar geometry and bond angles of about 120 degrees, as in BF3.
• sp3 hybridisation: Mixing of one s and three p orbitals to form four hybrid orbitals arranged tetrahedrally, as in CH4.
• sp3d hybridisation: Mixing of one s, three p, and one d orbital to form five hybrid orbitals with trigonal bipyramidal electron geometry, as in PCl5.
• sp3d2 hybridisation: Mixing of one s, three p, and two d orbitals to form six hybrid orbitals with octahedral electron geometry, as in SF6.
• Sigma bond: A covalent bond formed by head-on overlap of orbitals along the internuclear axis; every single bond is a sigma bond.
• Pi bond: A covalent bond formed by sideways overlap of parallel p orbitals above and below the internuclear axis; it occurs in multiple bonds.
• Bond order: The number of bonds between two atoms; higher bond order generally means greater bond strength and shorter bond length.
• Resonance: A situation in which one Lewis structure cannot fully represent a species, so the actual structure is a resonance hybrid of two or more contributing structures.
• Hydrogen bond: A relatively strong intermolecular or intramolecular attraction involving hydrogen bonded to a highly electronegative atom, usually nitrogen, oxygen, or fluorine.
• Molecular orbital: An orbital formed by the combination of atomic orbitals and spread over the entire molecule.
• Bonding molecular orbital: A molecular orbital of lower energy than the combining atomic orbitals that increases molecular stability.
• Antibonding molecular orbital: A higher-energy molecular orbital that reduces stability and is represented with an asterisk, such as sigma-star or pi-star.
• Bond order in molecular orbital theory: Bond order = one-half × (number of electrons in bonding molecular orbitals − number of electrons in antibonding molecular orbitals).

Easily Confused

• Electron geometry and molecular geometry: Electron geometry counts both bond pairs and lone pairs, whereas molecular geometry describes the positions of atoms only.
• Bond pair and lone pair: A bond pair is shared between atoms; a lone pair remains on one atom and does not participate directly in bonding.
• Sigma bond and pi bond: A sigma bond forms through head-on overlap along the internuclear axis, whereas a pi bond forms through sideways overlap of parallel p orbitals.
• Single, double, and triple bonds: A single bond contains one sigma bond; a double bond contains one sigma and one pi bond; a triple bond contains one sigma and two pi bonds.
• Resonance structures and the resonance hybrid: Resonance structures are contributing Lewis representations, while the actual species is their more stable resonance hybrid.
• Ionic character and ionic bonding: Ionic character can occur to varying degrees across a continuous covalent-to-ionic scale, whereas ionic bonding refers to electrostatic attraction between oppositely charged ions.
• Paramagnetic and diamagnetic species: Paramagnetic species contain one or more unpaired electrons; diamagnetic species have all electrons paired.
• Bond order and molecular-orbital bond order: Bond order may refer generally to the number of bonds between atoms, whereas molecular-orbital bond order is calculated using BO = 1/2(Nb − Na).

What Gets Asked

• Constructing Lewis structures: Questions require total valence-electron counting, octet completion, placement of remaining electrons on the central atom, and formation of multiple bonds where necessary. Marks are lost by omitting lone pairs, miscounting electrons, or ignoring octet exceptions.
• Applying octet-rule exceptions: Questions may involve hydrogen and helium following a duet, incomplete octets in boron and beryllium, expanded valence shells in elements in the third period and beyond, or species with an odd number of electrons. The key error is applying the octet rule without recognising these exceptions.
• Predicting molecular shape and bond angle: Questions may use BeCl2, BF3, CH4, NH3, H2O, PCl5, SF6, or XeF4. Marks are lost by using molecular geometry without counting lone pairs, or by failing to account for the greater repulsion of lone pairs.
• Assigning hybridisation: Questions may ask for sp, sp2, sp3, sp3d, or sp3d2 hybridisation and the corresponding geometry. The steric number is the total number of sigma bonds and lone pairs around the central atom; confusing electron geometry with molecular geometry leads to incorrect answers.
• Relating bonding to bond length and strength: Questions may compare single, double, and triple bonds or ask about bond dissociation enthalpy. The required relationship is that higher bond order generally gives shorter, stronger bonds and larger bond dissociation enthalpy.
• Using polarity, resonance, or molecular-orbital theory: Questions may require use of μ = q × r, percentage ionic character = observed dipole moment / calculated dipole moment × 100%, resonance principles, or BO = 1/2(Nb − Na). Common errors include assuming every polar bond gives a polar molecule, changing atom positions between resonance structures, or reversing bonding and antibonding electron counts.