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Chemistry Unit 2: Chemical bonding and compound structure

Atoms bond to reach stable electron arrangements. The type of bond and the molecule's shape decide its properties.

Covers: Ionic, covalent and metallic bonds, Lewis structures, VSEPR shapes, polarity, naming compounds

Study guide

Atoms bond to reach stable electron arrangements, often a full octet. In ionic bonds, a metal transfers electrons to a nonmetal, forming oppositely charged ions held together in a crystal lattice; ionic compounds have high melting points and conduct electricity when melted or dissolved. In covalent bonds, nonmetals share electron pairs; single, double and triple bonds share one, two and three pairs. In metallic bonds, metal atoms share a "sea" of delocalized electrons, which makes metals conductive and malleable.

Lewis structures show valence electrons as dots and bonds as lines. To draw one, count total valence electrons, connect atoms with single bonds, complete octets on outer atoms, then form double or triple bonds if the central atom lacks an octet. Formal charges help choose the best structure, and some molecules have resonance structures.

VSEPR theory predicts shape: electron groups around a central atom spread out as far as possible. Two groups give linear (180°, like CO₂), three give trigonal planar (120°), and four give tetrahedral (109.5°, like CH₄). Lone pairs take up space, so NH₃ is trigonal pyramidal and H₂O is bent.

Polarity depends on bonds and shape. A bond is polar when the atoms have different electronegativities. A molecule is polar if its polar bonds don't cancel, so CO₂ is nonpolar (symmetric) while H₂O is polar (bent).

Naming: ionic compounds name the cation then the anion (sodium chloride, iron(III) oxide), with polyatomic ions like sulfate and nitrate. Covalent compounds use prefixes (carbon dioxide, dinitrogen tetroxide). Acids follow their own rules, like hydrochloric acid and sulfuric acid.

Hybridization follows from the number of electron groups around an atom: two groups give sp (linear), three give sp² (trigonal planar) and four give sp³ (tetrahedral). A double bond is one sigma bond plus one pi bond, and a triple bond is one sigma plus two pi bonds. Formal charge = valence electrons − nonbonding electrons − ½(bonding electrons). The best Lewis structure has formal charges as close to zero as possible, with any negative charge on the more electronegative atom. The formal charges in a structure always add up to the ion's overall charge.

As bond order increases from single to double to triple, bonds get shorter and stronger. In a resonance hybrid, the bond order is the average over the resonance structures, so all the equivalent bonds have the same length. Lattice energy measures how strongly the ions in an ionic solid attract each other. By Coulomb's law, it grows with higher ion charges and smaller ions. MgO has a much stronger lattice than NaCl, and RbCl has a weaker one than NaCl because Rb⁺ is larger.

Lone pairs push bonded atoms closer together, so NH₃ has bond angles near 107° and H₂O near 104.5°. Atoms in period 3 and below can hold more than eight electrons. Five electron groups give a trigonal bipyramidal shape (PCl₅), and six give an octahedral one (SF₆).

Key ideas

• Ionic bonds transfer electrons (metal + nonmetal); covalent bonds share electrons (nonmetals).
• Lewis structures show valence electrons; most atoms want an octet.
• VSEPR: electron groups spread out to minimize repulsion, setting the shape.
• A molecule is polar if it has polar bonds and an asymmetric shape.
• Metallic bonding is a sea of shared electrons, which makes metals conductive.

Formulas and rules

Common mistakes

• Calling CO₂ polar. Its polar bonds cancel because it's linear.
• Forgetting lone pairs when picking a VSEPR shape.

Words to know

Ionic bond

Electrons transferred from a metal to a nonmetal.

Tetrahedral

Molecular shape with four bonds at 109.5°, like CH₄.

Polar molecule

Uneven charge from polar bonds and an asymmetric shape (like bent H₂O).

Valence electrons

Outer-shell electrons used in bonding (nitrogen has 5).

Linear molecular geometry

Two bonding groups at 180°, like CO₂.

Covalent bond

Nonmetals sharing electron pairs.

Metallic bond

Metal atoms sharing a "sea" of delocalized electrons.

Octet rule

Atoms tend to bond until they have 8 valence electrons.

Lewis structure

A diagram of valence electrons as dots and bonds as lines.

Resonance structures

Two or more valid Lewis structures for the same molecule.

Formal charge

Valence electrons − (lone pair electrons + bonds); helps pick the best structure.

VSEPR theory

Electron groups spread as far apart as possible, which sets molecular shape.

Trigonal planar

Three electron groups at 120°.

Bent

Shape of H₂O: two bonds and two lone pairs.

Polyatomic ion

A charged group of atoms, like sulfate (SO₄²⁻).

Naming ionic compounds

Cation then anion, with a Roman numeral for transition metals: iron(III) oxide.

Naming covalent compounds

Use prefixes: carbon dioxide, dinitrogen tetroxide.

Sample practice questions

5 more questions like these, from level 1 to 5, are in the practice set.