Ionic and Covalent Substances

Objective: Predict charges and neutral formulas for simple ions.

An ion has a net electric charge because its electron count differs from its proton count. A cation is positive after electron loss, and an anion is negative after electron gain in common simple models. A neutral ionic formula balances total charges: Mg²⁺ and Cl⁻ combine as MgCl₂ in a 1:2 ratio, not MgCl. The formula states the simplest whole-number ratio of ions in an extended crystal, not an individual molecule. Some ions are polyatomic groups whose atoms remain bonded internally. Charge balance is necessary for formula writing but does not prove that every imagined ion combination forms a stable substance.

Objective: Count atoms and interpret simple shared-electron diagrams.

Covalent molecules contain atoms connected by shared electron density; a molecular formula counts atoms of each element but does not uniquely reveal bond arrangement. H₂O has two hydrogen atoms and one oxygen atom, while CO₂ has one carbon and two oxygens. Lewis structures represent bonding pairs and lone pairs; a single line commonly represents one shared electron pair. Some molecules have double or triple bonds, and resonance or more advanced electron descriptions may be needed. Molecular shape affects properties: water is bent, not a straight H–O–H arrangement. Avoid treating a flat Lewis sketch as a complete three-dimensional structure.

Objective: Compare conductivity, melting, and solubility with qualifications.

A solid ionic crystal usually does not conduct well because its ions are fixed in place, while a molten ionic substance or suitable aqueous solution can conduct as ions move. Many molecular covalent substances do not conduct in their ordinary pure forms, but broad labels alone cannot predict every material property; graphite and network covalent solids are important exceptions. Melting temperatures depend on the strength and organization of interactions throughout a substance, not merely on whether one covalent bond exists. Solubility depends on solute–solvent interactions and conditions. Compare actual substances and states rather than memorizing “ionic equals high” and “covalent equals low” as universal laws.