Difficulty: Introductory to Intermediate | Prerequisites: Atomic Structure and the Periodic Table study notes (valence electrons, electron configuration, periodic trends).
This topic explains why and how atoms join together. Everything from the water you drink to the plastic in your phone exists because atoms form bonds. You will learn two main bond types (ionic and covalent), how to name the resulting compounds, how to draw Lewis structures, how to predict molecular shapes, and how polarity determines a molecule's physical behaviour. This is the bridge between atomic theory and chemical reactions.
Atoms bond by transferring electrons (ionic bonds, metals + nonmetals) or sharing electrons (covalent bonds, nonmetals + nonmetals). Lewis structures map out bonding and lone pairs. VSEPR theory uses electron-group count to predict molecular geometry. Electronegativity differences determine whether a bond is nonpolar covalent, polar covalent or ionic, and molecular shape determines whether the molecule as a whole is polar.
Ionic bond
A bond formed by the transfer of one or more electrons from a metal to a nonmetal, producing oppositely charged ions that attract each other. Think of it as one atom giving electrons away and the other accepting them.
Covalent bond
A bond formed when two nonmetal atoms share one or more pairs of electrons. Neither atom fully loses or gains an electron.
Cation
A positively charged ion, formed when an atom loses one or more electrons. Metals form cations. In simple terms, fewer electrons than protons means a positive charge.
Anion
A negatively charged ion, formed when an atom gains one or more electrons. Nonmetals form anions.
Polyatomic ion
An ion made of two or more atoms covalently bonded together that carries a net charge. Sulfate (SO4 2-) and nitrate (NO3-) are common examples.
Lewis structure (electron dot structure)
A diagram showing all valence electrons in a molecule as dots or lines. Lines represent bonding pairs; dots represent lone pairs.
Octet rule
Atoms tend to gain, lose or share electrons until they have eight electrons in their valence shell, matching the electron configuration of a noble gas. Hydrogen is the exception: it aims for two.
VSEPR theory
Valence Shell Electron Pair Repulsion theory. It predicts molecular geometry by assuming that electron groups around a central atom arrange themselves to be as far apart as possible.
Electronegativity
A measure of an atom's ability to attract shared electrons in a covalent bond. Fluorine is the most electronegative element.
Polar bond
A covalent bond in which electrons are shared unequally because the two atoms differ in electronegativity. One end of the bond is slightly negative, the other slightly positive.
Nonpolar bond
A covalent bond in which electrons are shared equally, either because the atoms are identical or very similar in electronegativity.
Molecular polarity
Whether the molecule as a whole has a net dipole moment. Depends on both bond polarity and molecular shape. A molecule can have polar bonds and still be nonpolar if the bond dipoles cancel symmetrically.
Atoms bond to achieve a stable, noble-gas-like electron configuration. There are two main routes.
Ionic bonding occurs when a metal transfers electrons to a nonmetal. The metal becomes a cation (+), the nonmetal becomes an anion (-), and the electrostatic attraction between the two holds the compound together. Example: Na+ and Cl- combine to form NaCl.
Covalent bonding occurs when two nonmetals share electron pairs. Neither atom fully surrenders electrons. Example: two hydrogen atoms share a pair of electrons to form H2.
A useful shortcut: metal + nonmetal = ionic. Nonmetal + nonmetal = covalent. This rule covers the vast majority of introductory-level compounds.
Cations are positive ions. They form when atoms lose electrons. Since protons now outnumber electrons, the ion carries a positive charge.
Anions are negative ions. They form when atoms gain electrons.
Group trends make ion charges predictable for main-group elements:
Group 1A metals lose 1 electron to form +1 cations (Na+, K+).
Group 2A metals lose 2 electrons to form +2 cations (Ca2+, Mg2+).
Group 6A nonmetals gain 2 electrons to form -2 anions (O2-, S2-).
Group 7A nonmetals gain 1 electron to form -1 anions (Cl-, Br-).
Some transition metals form cations with different charges (variable charge metals). Iron, for instance, can be Fe2+ or Fe3+. These are named with Roman numerals: iron(II) and iron(III). An older naming system uses the suffix "-ous" for the lower charge and "-ic" for the higher charge (ferrous and ferric).
Formation and formulas
Ionic compounds are composed of cations and anions arranged in a crystal lattice. The formula is written so that total positive charge equals total negative charge (net charge = 0). Example: calcium chloride is CaCl2 because Ca2+ needs two Cl- ions to balance.
Naming ionic compounds
Name the cation first (the metal), then the anion.
The cation keeps the element name. If the metal has a variable charge, add a Roman numeral in parentheses.
The anion name changes its ending to "-ide" (chlorine becomes chloride, oxygen becomes oxide).
Examples: NaCl = sodium chloride. Fe2O3 = iron(III) oxide.
Polyatomic ions are groups of covalently bonded atoms that carry a net charge and act as a single unit in ionic compounds. Common ones to memorise: sulfate (SO4 2-), nitrate (NO3-), carbonate (CO3 2-), hydroxide (OH-), ammonium (NH4+), phosphate (PO4 3-).
Physical properties of ionic compounds
Crystalline solids at room temperature.
High melting and boiling points (strong electrostatic attractions).
Conduct electricity when dissolved in water or melted, because the ions are free to move.
Covalent bonds form when nonmetal atoms share electron pairs. The resulting particle is called a molecule. Covalent compounds are named with prefixes indicating the number of each atom: mono-, di-, tri-, tetra-, penta-, and so on. The second element's name ends in "-ide." Example: CO2 is carbon dioxide; N2O is dinitrogen monoxide.
Lewis structures show every valence electron in a molecule. Bonding pairs (shared electrons) are drawn as lines between atoms. Lone pairs (unshared electrons) are drawn as dots on individual atoms.
Drawing a Lewis structure:
Count the total valence electrons for all atoms.
Place the least electronegative atom in the centre (hydrogen is always on the outside).
Connect atoms with single bonds (each bond uses 2 electrons).
Distribute remaining electrons as lone pairs, starting with outer atoms, to satisfy the octet rule.
If the central atom still lacks an octet, convert lone pairs on adjacent atoms into double or triple bonds.
The octet rule: atoms tend to bond until they have 8 electrons in their valence shell. Hydrogen is the exception, aiming for 2. When single bonds alone do not satisfy the octet, double bonds (4 electrons shared) or triple bonds (6 electrons shared) form.
VSEPR (Valence Shell Electron Pair Repulsion) theory predicts the three-dimensional shape of a molecule by counting the electron groups around the central atom. An "electron group" is any bonding pair, lone pair, double bond or triple bond attached to the central atom. Each double or triple bond counts as one group.
Electron groups repel each other and arrange themselves as far apart as possible, producing predictable geometries:
2 groups, 180°: linear. Example: CO2.
3 groups, 120°: trigonal planar. Example: BF3.
4 groups, 109.5°: tetrahedral. Example: CH4.
4 groups with 1 lone pair, ~107°: trigonal pyramidal. Example: NH3.
4 groups with 2 lone pairs, ~105°: bent. Example: H2O.
Lone pairs take up more space than bonding pairs, which is why the bond angles in trigonal pyramidal and bent shapes are slightly less than the ideal tetrahedral angle.
Electronegativity measures how strongly an atom attracts shared electrons. It increases across a period (left to right) and decreases down a group, following the same pattern as ionisation energy.
The difference in electronegativity between two bonded atoms determines the bond type:
Nonpolar covalent: difference close to 0. Electrons shared equally. Example: H2, O2.
Polar covalent: moderate difference (roughly 0.4 to 1.7). Electrons shared unequally, creating partial charges. Example: H-Cl.
Ionic: large difference (roughly > 1.7). Electrons effectively transferred. Example: NaCl.
Molecular polarity depends on both the polarity of individual bonds and the shape of the molecule. If polar bonds are arranged symmetrically so that their dipoles cancel, the molecule is nonpolar overall. CO2 is linear with two equal polar bonds pointing in opposite directions, so it is nonpolar. Water is bent, so its two polar O-H bonds do not cancel, making water a polar molecule.
Molecular polarity matters because it determines solubility ("like dissolves like"), boiling points and many other physical properties.
Water's bent shape and polarity make it an exceptional solvent for ionic and polar substances, which is why biological chemistry runs in water.
Table salt (NaCl) dissolves in water because the polar water molecules pull the Na+ and Cl- ions apart, overcoming the ionic lattice energy.
Diamond (pure carbon, covalent network) and graphite (carbon, layered covalent sheets) show how bonding arrangement, not just composition, determines material properties.
Students often think that if a molecule contains polar bonds, the molecule must be polar. CO2 has polar C=O bonds but is nonpolar because its linear shape causes the dipoles to cancel.
Students confuse ionic compounds with molecules. Ionic compounds form extended lattices, not discrete molecules. NaCl is a formula unit, not a molecule.
Students sometimes think double bonds are exactly twice as strong as single bonds. They are stronger, but not exactly double (the second pair of shared electrons, in a pi bond, is weaker than the first sigma bond).
Students forget to count lone pairs as electron groups when using VSEPR. A lone pair on the central atom changes the molecular shape even though it is invisible in the molecular formula.
⚠️ Drawing Lewis structures is one of the most frequently tested skills. Practise counting valence electrons and satisfying octets.
⚠️ Know how to name ionic compounds (including those with polyatomic ions and variable-charge metals) and covalent compounds (using prefixes).
⚠️ Expect VSEPR questions asking you to predict the shape and bond angle from a Lewis structure.
⚠️ Be able to determine whether a molecule is polar or nonpolar given its shape and bond polarities.
⚠️ Writing correct formulas for ionic compounds (balancing charges to zero) is a foundational skill tested throughout the course.
True or false: A bond between sodium and chlorine is covalent.
False. Sodium is a metal and chlorine is a nonmetal, so the bond is ionic.
What is the molecular geometry of a molecule with 4 electron groups and 1 lone pair on the central atom?
Trigonal pyramidal.
Fill in the blank: CO2 has polar bonds but is a ______ molecule because its shape is linear.
Nonpolar.
True or false: The formula for aluminium oxide is Al2O3.
True. Al3+ and O2- balance as 2(+3) + 3(-2) = 0.
How many valence electrons does nitrogen contribute to a Lewis structure?
Five.
Q: Draw the Lewis structure for water (H2O) and predict its molecular geometry.
A: Oxygen has 6 valence electrons; each hydrogen has 1. Total = 8 electrons. Oxygen is central, bonded to two hydrogens (2 bonds = 4 electrons used), with 2 lone pairs on oxygen (4 electrons). Four electron groups, 2 lone pairs: the shape is bent, with a bond angle of roughly 105°.
Q: Name the compound Fe2O3.
A: Iron(III) oxide. Each oxygen is -2 (total -6), so each iron is +3 to balance.
Q: Is CCl4 polar or nonpolar? Explain.
A: Nonpolar. Carbon is central with four identical C-Cl polar bonds arranged in a tetrahedral geometry. The four dipoles cancel symmetrically.
Q: Write the formula for the ionic compound formed by calcium and the nitrate ion.
A: Ca(NO3)2. Calcium is Ca2+ and nitrate is NO3-. Two nitrate ions balance one calcium ion.
Q: What is the difference between a double bond and a triple bond in terms of electron count?
A: A double bond shares 4 electrons (2 pairs). A triple bond shares 6 electrons (3 pairs).
Bonding leads directly into intermolecular forces, where the polarity you determine here governs boiling points, solubility and viscosity. Lewis structures are the starting point for understanding reaction mechanisms in organic chemistry. Ionic compound formulas and naming conventions carry through to stoichiometry and solution chemistry. VSEPR shapes reappear in biochemistry when you study enzyme active sites and drug-receptor interactions.
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