Topic 2: Molecular and Ionic Compound Structure and Properties

Unit 2 moves from individual atoms to the forces that hold them together and the three-dimensional shapes that result. You will learn to classify bonding along a continuum, predict structures with Lewis diagrams and VSEPR, and connect those structures to measurable properties like melting point, conductivity, and bond angle.

Bond Types and the Bonding Continuum

Chemical bonds are not strictly one of three discrete kinds; they fall along a continuum governed by how electrons are shared between atoms. In a pure covalent (nonpolar) bond, two atoms of identical electronegativity share electrons equally. As the electronegativity difference grows, the shared pair shifts toward the more electronegative atom, creating a polar covalent bond with partial charges. When the difference becomes large enough, electrons transfer almost completely and an ionic bond forms between cations and anions. Metallic bonding is a separate model in which metal atoms release valence electrons into a shared 'sea' that moves freely. A useful guideline: metal plus nonmetal tends toward ionic, nonmetal plus nonmetal toward covalent, and metal plus metal toward metallic, but always think in terms of electronegativity difference rather than memorized rules.

Coulomb's Law and Lattice Energy

Coulomb's law is the quantitative heart of this unit. The potential energy between two charged particles is proportional to the product of their charges divided by the distance between their centers (E proportional to q1 times q2 over r). Two consequences matter most for AP Chemistry: larger ion charges produce stronger attractions, and smaller distances (smaller ions) produce stronger attractions. Lattice energy, the energy released when gaseous ions combine into a solid ionic lattice, follows directly. MgO has a much higher lattice energy than NaCl because Mg2+ and O2- each carry double the charge of Na+ and Cl-, and because the ions are smaller. Higher lattice energy translates into higher melting points and greater hardness. When comparing two ionic compounds, evaluate charge first (it has the bigger effect) and then ionic size.

Structure of Ionic Solids

Ionic compounds are not made of discrete molecules; they form an extended three-dimensional lattice in which each cation is surrounded by anions and each anion by cations, maximizing attractions and minimizing repulsions. This arrangement explains characteristic properties. Ionic solids are hard and brittle: applying force can shift one layer so that like charges align, and the resulting repulsion shatters the crystal. They have high melting and boiling points because many strong electrostatic attractions must be overcome. They do not conduct electricity when solid because the ions are locked in place, but they conduct well when molten or dissolved in water because the ions become mobile and can carry charge.

Metals, Alloys, and the Electron Sea

In the electron-sea model, metal cations sit in a fixed lattice while delocalized valence electrons flow throughout the structure. This explains why metals conduct electricity and heat, are malleable and ductile (layers of cations can slide past one another without breaking the bonding), and are often lustrous. Alloys are mixtures of a metal with one or more other elements and come in two structural types. In a substitutional alloy, atoms of similar size replace some host metal atoms in the lattice; brass (copper and zinc) is an example. In an interstitial alloy, much smaller atoms fit into the spaces between the larger host atoms; steel, with small carbon atoms in an iron lattice, is the classic case. Interstitial atoms restrict the sliding of layers, which makes interstitial alloys typically harder and less malleable than the pure metal.

Drawing Lewis Diagrams

Lewis diagrams track valence electrons and predict how atoms connect. A reliable procedure: count the total valence electrons (add electrons for an anion's charge, subtract for a cation's), arrange the atoms with the least electronegative atom usually in the center (hydrogen is never central), connect atoms with single bonds, then distribute remaining electrons as lone pairs to complete octets on outer atoms first. If the central atom lacks an octet, convert lone pairs into double or triple bonds. Several exceptions appear on the exam: hydrogen wants only two electrons, boron and beryllium are often stable with fewer than eight, and elements in period 3 and beyond (such as S and P) can hold expanded octets of ten or twelve electrons.

Resonance and Formal Charge

When more than one valid Lewis diagram can be drawn by moving only electrons (not atoms), the true structure is a resonance hybrid, an average of all contributing structures. Ozone and the carbonate, nitrate, and sulfate ions are common examples; their bonds are identical in length and intermediate between single and double bonds. Formal charge helps select the most reasonable structure. Calculate it as the number of valence electrons in the free atom minus nonbonding electrons minus half the bonding electrons. The best Lewis structure usually keeps formal charges as close to zero as possible and places any negative formal charge on the most electronegative atom. Note that formal charge is a bookkeeping tool and does not represent real, full charges on atoms.

VSEPR: Shapes and Bond Angles

Valence Shell Electron Pair Repulsion theory predicts molecular geometry by assuming that electron domains (bonding pairs and lone pairs) arrange themselves as far apart as possible around a central atom. Count the domains to find the electron geometry: two domains give linear (180 degrees), three give trigonal planar (120 degrees), four give tetrahedral (109.5 degrees), five give trigonal bipyramidal, and six give octahedral. Lone pairs occupy more space than bonding pairs, so they reduce bond angles slightly and change the molecular shape: four domains with one lone pair give a trigonal pyramidal molecule (ammonia, about 107 degrees) and with two lone pairs give a bent molecule (water, about 104.5 degrees). Molecular polarity depends on both bond polarity and shape; symmetric arrangements can cancel dipoles, making a molecule with polar bonds (such as CO2) nonpolar overall.

Hybridization, Sigma Bonds, and Pi Bonds

Hybridization explains how atomic orbitals mix to form bonds matching observed geometry. Count the electron domains on the central atom: two domains correspond to sp hybridization (linear), three to sp2 (trigonal planar), and four to sp3 (tetrahedral or its lone-pair variants). Every single bond is one sigma bond, formed by direct, head-on orbital overlap along the internuclear axis. Double and triple bonds contain pi bonds, formed by side-to-side overlap of unhybridized p orbitals. A double bond is one sigma plus one pi; a triple bond is one sigma plus two pi. Sigma bonds allow free rotation, while pi bonds lock the geometry and prevent rotation, which is why double bonds give rise to cis-trans isomers.