Learning objectives
- Distinguish ionic lattices, covalent molecules or networks, and metallic solids.
- Write neutral ionic formulas by balancing total positive and negative charge.
- Draw introductory Lewis structures and use electron regions to estimate molecular geometry.
- Separate bond polarity from whole-molecule polarity and connect structure to properties with stated limits.
Prerequisites: Complete Atomic structure. You should be able to count valence electrons and work with positive and negative integers.
1. Why atoms bond
Chemical bonding describes attractive interactions that hold atoms or ions in a stable arrangement. Positively charged nuclei attract electrons, and like charges repel. Introductory models ask how an arrangement can lower the system's energy compared with separated atoms.
Electronegativity measures how strongly an atom attracts shared electron density. A large difference often gives a bond with substantial ionic character; similar values favor sharing. This is a continuum rather than a perfect two-box classification. The metal/nonmetal pattern helps predict common compounds, but exceptions need more information.
Valence electrons are the outer electrons involved most directly in main-group bonding. The octet rule is a counting guide, not a universal law. Hydrogen commonly has two electrons, and boron or period 3 elements can have exceptions.
2. Ionic bonding makes a lattice
When a metal atom transfers one or more electrons to a nonmetal atom, the products are cations and anions. Their electrostatic attraction extends in every direction, so a solid ionic compound is a three-dimensional lattice of alternating charges. Sodium chloride is therefore described by formula unit NaCl, the simplest whole-number ratio of Na+ and Cl−, not by separate NaCl molecules.
Ionic formulas must have net charge zero. Use the smallest whole-number ratio and write the cation first. Calcium forms Ca2+ and chloride forms Cl−, so two chlorides balance one calcium. Aluminum oxide uses two Al3+ and three O2−. Strong attractions help explain high melting points and brittleness. In a solid, ions are held in place; when molten or dissolved, mobile ions can conduct.
Worked example 1: an ionic formula
Write the formula for magnesium nitride. Magnesium is Mg2+; nitride is N3−. The least common multiple of 2 and 3 is 6. Three Mg2+ supply +6 and two N3− supply −6, so the formula is Mg3N2. The subscripts describe a lattice ratio, not a small molecule.
Dissolving requires more than “ionic means soluble.” Water must interact favorably enough to separate ions from the lattice, and some ionic solids are sparingly soluble. Dissociation in water is not automatically a new chemical reaction.
3. Covalent bonds share electrons
In a covalent bond, atoms share electron pairs. Two nonmetals can form a discrete molecule such as H2O, CO2, or CH4. A covalent network solid, such as diamond or quartz, extends covalent connections through a large structure rather than making independent molecules. The formula tells which atoms and ratios are present, but the structure determines many properties.
Single, double, and triple bonds share one, two, and three pairs. Bond order generally correlates with shorter, stronger bonds within a related series. Do not call every compound a molecule: ionic lattices and network solids are exceptions.
Many molecular substances have lower melting and boiling temperatures than ionic lattices because heating often separates molecules rather than breaking every covalent bond. Intermolecular attractions, network structures, and shape create exceptions.
4. Metallic bonding
In a metal, many valence electrons are spread through a solid containing positive atomic cores. This simplified “mobile electron” model explains why metals conduct electrical charge and thermal energy, reflect light, and can be hammered or drawn without the lattice shattering along one set of directional bonds. Metallic bonding varies among elements.
Alloys mix elements in a metallic solid and can change strength or flexibility. The model does not mean electrons move without attraction; solid-state models describe their collective behavior more precisely.
5. Lewis structures and molecular shape
A Lewis structure places valence electrons as dots and shared pairs as lines. Count valence electrons, choose a central atom, connect atoms, complete terminal octets, place remaining electrons on the center, and add multiple bonds when needed. Hydrogen is terminal and needs two electrons. Halogens usually form one bond in simple neutral molecules.
Worked example 2: carbon dioxide
CO2 has 4 + 2(6) = 16 valence electrons. Put carbon in the center and connect O–C–O. Completing both oxygen octets and giving carbon an octet produces O=C=O with two lone pairs on each oxygen. Around carbon there are two electron regions, so VSEPR predicts a linear shape near 180°. A double bond counts as one electron region, even though it contains two shared pairs. The polar C=O bond dipoles point in opposite directions.
VSEPR means valence-shell electron-pair repulsion. Two regions give linear electron geometry; three give trigonal planar; four give tetrahedral. Lone pairs occupy more space than bonding pairs, so H2O is bent and NH3 is trigonal pyramidal. Ideal angles are starting points, not guaranteed measurements.
6. Polarity, shape, and properties
A polar covalent bond has unequal electron sharing and a partial negative end near the more electronegative atom. A molecule's polarity depends on bond dipoles and their directions in three dimensions. Symmetric CO2 is nonpolar overall, while bent H2O is polar because its bond dipoles do not cancel. Symmetric CH4 is treated as nonpolar in the introductory model.
Worked example 3: water and carbon dioxide
Water has two O–H bonds and two lone pairs on oxygen. Four electron regions give a tetrahedral electron arrangement, but the two O–H bonds give a bent molecular shape. The bond dipoles add to a net dipole, so water is polar. Carbon dioxide has two electron regions and a linear arrangement. Its equal C=O dipoles cancel, so the molecule is nonpolar. This difference helps explain why water interacts strongly with ions and other polar substances, though polarity alone does not decide every solubility question.
“Like dissolves like” is a starting pattern: polar solvents often dissolve polar or ionic solutes, and nonpolar solvents often dissolve nonpolar solutes. Solubility also depends on lattice energy, intermolecular forces, temperature, pressure, and specific structures. Ionic character, covalent character, hydrogen bonding, and molecular size affect boiling point and conductivity.
Practice with explained answers
Why is CaCl2 neutral? Explain the particle arrangement.
Show reasoning
One Ca2+ gives +2 and two Cl− give −2, so total charge is zero. The solid is a lattice with a 1:2 ion ratio, not a collection of CaCl2 molecules.
Predict the shape and polarity of BF3 in the introductory VSEPR model.
Show reasoning
Boron has three bonding regions and no lone pair in the usual Lewis structure, giving trigonal planar geometry. The three identical bond dipoles are symmetrically arranged, so they cancel and the molecule is treated as nonpolar.
How many valence electrons should be counted for NH3? What is its molecular shape?
Show reasoning
N contributes 5 and three H atoms contribute 3, for 8 total. Nitrogen has three bonds and one lone pair, four electron regions, and a trigonal pyramidal molecular shape. The lone pair prevents cancellation, so NH3 is polar.
Why can solid NaCl not conduct electricity as easily as aqueous NaCl?
Show reasoning
In the solid lattice, ions are locked into place. In water, separated hydrated ions can move and carry charge. The solution conducts only if enough mobile ions are present; water and concentration also matter.
Sources and next steps
The lattice and formula-unit distinction follows OpenStax Chemistry 2e, ionic bonding. Descriptions of electron sharing and bond models are supported by OpenStax, covalent bonding. Geometry, VSEPR limits, and polarity are developed in OpenStax, molecular structure and polarity. These are introductory models: real bonding includes a continuum of electron distribution, resonance, molecular orbitals, and solid-state effects beyond this lesson.