Chapter 14: Atomic Structure, Periodicity, and Chemical Bonding

The atom

An atom has a nucleus of protons (charge +1) and neutrons (neutral), surrounded by electrons (charge −1). The atomic number is the number of protons; the mass number is protons plus neutrons. Isotopes have the same but different numbers of neutrons. The relative atomic mass is the weighted mean of the isotopic masses relative to of a carbon-12 atom; a mass spectrometer measures isotopic abundances.

Electronic structure

Electrons occupy orbitals characterised by quantum numbers: (energy level), (shape: ), (orientation), and ( spin). Three rules fill the orbitals: the Aufbau principle (lowest energy first), the Pauli exclusion principle (no two electrons share all four quantum numbers, so an orbital holds at most two with opposite spins), and Hund’s rule (degenerate orbitals fill singly first). The order is . Examples: Na ; Fe . Chromium () and copper () are exceptions because half-filled and filled subshells are especially stable. When forming cations, the electrons are lost before the .

Periodic trends

There are two well-known exceptions in ionization energy within period 2: beryllium is higher than boron (the boron electron is in a higher-energy orbital), and nitrogen is higher than oxygen (the oxygen pair has electron–electron repulsion).

Ionic and covalent bonding

Ionic bonding results from the electrostatic attraction between ions formed by electron transfer, typically between a metal and a non-metal. The strength of an ionic lattice is measured by its lattice enthalpy, which increases with higher ionic charge and smaller ionic radius (so MgO has a higher melting point than NaCl). A Born–Haber cycle applies Hess’s law to find lattice enthalpy.

Covalent bonding is the sharing of electron pairs. Draw Lewis structures by counting valence electrons, placing bonds, and completing octets; calculate formal charge ( valence electrons, non-bonding electrons, bonding electrons). Exceptions to the octet include electron-deficient molecules (BF₃), odd-electron molecules (NO), and expanded octets (PCl₅, SF₆).

Shapes of molecules: VSEPR

The valence shell electron-pair repulsion theory says that electron pairs around a central atom arrange themselves to minimise repulsion; lone pairs repel more than bonding pairs.

A sigma () bond is formed by head-on overlap; a pi () bond by sideways overlap of orbitals. A double bond is one plus one ; a triple bond is one plus two . A molecule is polar if its bond dipoles do not cancel (H₂O and NH₃ are polar; CO₂ and CCl₄ are not). Molecular orbital theory explains why O₂ is paramagnetic (two unpaired electrons) and gives bond order .

Intermolecular forces and properties

London (dispersion) forces: exist between all molecules and increase with the number of electrons.

Dipole–dipole forces: between polar molecules.

Hydrogen bonding: between H bonded to N, O, or F and a lone pair on another N, O, or F; it explains the high boiling point of water and the lower density of ice.

Metallic bonding: a lattice of positive ions in a sea of delocalized electrons, which gives conductivity and malleability.

Boiling points, solubility, and volatility follow from the type and strength of these forces. Giant covalent structures (diamond, graphite, silica) have high melting points; graphite conducts because of delocalized electrons in layers.

Common mistakes

Writing configurations that violate Hund’s rule or the Pauli principle.

Forgetting that lone pairs affect shape but are not counted in the name of the shape.

Calling hydrogen bonds “bonds within the molecule”. They are intermolecular attractions.

Practice questions

Write the electron configuration of Cu and explain why it is an exception. []

Predict the shape and hybridization of SF₄. [See-saw; ]

Explain why NH₃ has a higher boiling point than PH₃.