Ionic and Molecular Compounds

Core Concepts

Ionic compounds: electron transfer and the octet rule

Atoms react to reach stable configurations, and the octet rule is a useful guide: atoms tend to gain, lose, or share electrons so their valence shell resembles a noble gas — eight electrons (two for hydrogen and helium). A metal on the left has few valence electrons and loses them easily; sodium (1 valence electron) becomes Na⁺. A nonmetal on the right has nearly a full shell and gains electrons; chlorine (7 valence electrons) becomes Cl⁻. The electrostatic attraction between the ions is the ionic bond. Ionic compounds do not exist as individual molecules: a formula unit shows only the ratio of ions, packed into an extended three-dimensional crystal lattice. That lattice explains their properties: high melting points (NaCl melts at 801 °C), hardness and brittleness, and conduction of electricity when molten or dissolved in water — but not when solid.

Molecular compounds: shared electrons

When nonmetals combine, neither atom readily gives up electrons; instead they share pairs, forming covalent bonds. The resulting particle — a group of atoms held together by covalent bonds — is a molecule. Water (H₂O), methane (CH₄), ammonia (NH₃), carbon dioxide (CO₂), and oxygen (O₂) are molecular compounds. Molecular substances tend to have low melting and boiling points, are often gases or volatile liquids, and generally do not conduct electricity.

Polyatomic ions

Some ions are groups of atoms carrying a net charge; these are polyatomic ions. The atoms within are held by covalent bonds, but the group behaves as a single charged unit in formulas and reactions. Common examples: ammonium (NH₄⁺), hydroxide (OH⁻), nitrate (NO₃⁻), carbonate (CO₃²⁻), sulfate (SO₄²⁻), phosphate (PO₄³⁻). Polyatomic ions let compounds of only nonmetals be ionic: ammonium chloride (NH₄Cl) contains no metal, yet it is an ionic compound of NH₄⁺ and Cl⁻.

Predicting ion charges from the periodic table

For main-group elements the charge is predictable from the group number: groups 1, 2, and 13 form 1+, 2+, and 3+ cations; groups 15, 16, and 17 form 3−, 2−, and 1− anions. Transition metals are the exception — most have more than one common charge (Fe²⁺/Fe³⁺, Cu⁺/Cu²⁺, Sn²⁺/Sn⁴⁺) — which is why their names carry Roman numerals in the next topic.

Writing formulas by charge balance

The guiding rule: in a neutral ionic compound, total positive charge equals total negative charge. Write the cation first, then the anion, and choose subscripts that make the charges cancel. The criss-cross method swaps charge magnitudes as subscripts (Ca²⁺ with Cl⁻ gives CaCl₂; Al³⁺ with O²⁻ gives Al₂O₃), then reduces to simplest whole numbers (Mg²⁺ with O²⁻ gives MgO, not Mg₂O₂). When a polyatomic ion is used more than once, enclose it in parentheses: Ca(OH)₂, Mg₃(PO₄)₂. Two cautions: the metal–nonmetal rule is a strong guideline rather than an absolute law — compounds like AlCl₃ and BeCl₂ are largely covalent — and a formula gives the ratio of ions, not the mass (the subject of Chapter 3).

How It Works / Step-by-Step Process

1. Identify the elements: metal + nonmetal (or polyatomic ion) → ionic; nonmetal + nonmetal → molecular.
2. Determine charges: main-group charges from group number; polyatomic charges from memory; transition metals from context or Roman numerals.
3. Ionic: balance total positive and negative charge, reduce to simplest ratio, use parentheses for repeated polyatomic ions.
4. Molecular: write symbols and subscripts from the given name or atom counts.
5. Predict properties (melting point, conductivity) from the class, and check the formula's charge balance.

Common Confusions