AP Chemistry: Chemical Bonding

Chemical bonding explains why atoms stick together and, just as importantly, why substances behave so differently in the real world. The bonding model you choose affects everything that follows: how to draw a correct Lewis structure, how to predict molecular geometry with VSEPR theory, whether a molecule is polar, and why a solid melts at 30°C instead of 3000°C. For AP Chemistry, the goal is not memorization of isolated facts but building a connected toolkit for predicting structure and properties.

Why atoms form bonds

Atoms bond because the bonded arrangement is typically lower in potential energy than separated atoms. Two broad ideas help guide prediction:

• Attractive forces (such as attraction between nuclei and shared or transferred electrons) lower energy.
• Repulsive forces (electron-electron and nucleus-nucleus repulsions) raise energy.

Stable bonding reflects a balance between these forces. In practice, you will classify bonding into ionic, covalent, and metallic categories, then refine that picture using electronegativity, molecular geometry, and intermolecular forces.

Distinguishing bond types: ionic, covalent, metallic

Ionic bonding

Ionic bonding is most common between a metal and a nonmetal. Electrons are transferred (conceptually) to form cations and anions, which are held together by electrostatic attraction in a crystal lattice.

Key consequences:

• High melting and boiling points due to strong attractions in the lattice.
• Electrical conductivity when molten or dissolved (ions can move), but not as solids (ions fixed in place).
• Brittleness because shifting layers can bring like charges together, causing repulsion.

In AP-style reasoning, you often connect ionic bonding to Coulomb’s law, which describes how attraction increases with larger charge and shorter distance:

$$F=k\frac{|q_1{q}_2|}{r^2}$$

Stronger ionic attractions (larger charges, smaller ionic radii) generally mean higher lattice energy and higher melting points.

Covalent bonding

Covalent bonding involves sharing electron pairs between atoms, typically nonmetals. Covalent bonds can be:

• Nonpolar covalent when electrons are shared roughly equally.
• Polar covalent when electrons are drawn more toward the more electronegative atom, producing partial charges (${\delta }^{-}$ and ${\delta }^{+}$).

Bond polarity trends with electronegativity difference, but AP Chemistry emphasizes that electronegativity is a model for electron distribution, not a strict switch between “ionic” and “covalent.” Many bonds are best described as polar covalent with partial ionic character.

Covalent substances split into two major property groups:

• Molecular covalent (discrete molecules like CO₂, H₂O): typically lower melting points, often gases or liquids, and poor conductors.
• Network covalent (extended covalent lattices like diamond or SiO₂): very high melting points and hard solids because breaking the solid requires breaking covalent bonds.

Metallic bonding

Metallic bonding occurs in metals, where valence electrons are delocalized across a lattice of metal cations. This “sea of electrons” model explains:

• High electrical and thermal conductivity (mobile electrons).
• Malleability and ductility (layers can slide without shattering, unlike ionic lattices).
• Luster (interaction of light with delocalized electrons).

Metallic bonding is distinct from covalent bonding because the electrons are not localized between specific pairs of atoms.

Lewis structures: building accurate electron models

Lewis structures are the starting point for predicting shape, polarity, and sometimes reactivity. A reliable workflow:

1. Count total valence electrons (adjust for ionic charge).
2. Choose a central atom (often least electronegative, never hydrogen).
3. Connect atoms with single bonds (each bond uses 2 electrons).
4. Complete octets on terminal atoms.
5. Place remaining electrons on the central atom.
6. Form multiple bonds if the central atom lacks an octet.

Formal charge and best structure choice

When multiple valid structures exist, formal charge helps identify the most reasonable one:

$$\text{Formal charge}={\text{valence e}}^{-}-({\text{nonbonding e}}^{-}+\frac{1}{2}{\text{bonding e}}^{-})$$

Strong candidates usually minimize formal charges and place negative formal charge on more electronegative atoms.

Resonance

Some molecules cannot be represented by a single Lewis structure (for example, ozone). Resonance indicates delocalized electrons. The real molecule is a resonance hybrid; resonance structures are bookkeeping tools for electron distribution, not rapidly flipping pictures.

VSEPR theory: predicting molecular geometry

Once a correct Lewis structure is drawn, VSEPR (Valence Shell Electron Pair Repulsion) predicts geometry by arranging electron domains to minimize repulsion.

Electron-domain geometry vs molecular geometry

Count electron domains around the central atom:

• Each single, double, or triple bond counts as one domain.
• Each lone pair counts as one domain.

Electron-domain geometries:

• 2 domains: linear
• 3 domains: trigonal planar
• 4 domains: tetrahedral
• 5 domains: trigonal bipyramidal
• 6 domains: octahedral

Molecular geometry describes the shape formed by atoms only (ignoring lone pairs), so lone pairs change the observed shape. For example, four domains with one lone pair gives a trigonal pyramidal molecular geometry.

Why lone pairs matter

Lone pairs repel more strongly than bonding pairs because their electron density is concentrated on the central atom. This commonly compresses bond angles relative to ideal values. Recognizing these deviations can help you justify observed trends rather than treating angles as isolated facts.

Hybridization: connecting shape to bonding model

Hybridization is a model that matches geometry to atomic orbital mixing on the central atom. It links VSEPR domain count to orbital sets:

• 2 domains: $sp$
• 3 domains: $s{p}^2$
• 4 domains: $s{p}^3$
• 5 domains: $s{p}^{3}d$
• 6 domains: $s{p}^3{d}^2$

Hybridization also supports bonding descriptions:

• Single bonds are typically $\sigma$ bonds.
• Double bonds consist of one $\sigma$ and one $\pi$ bond.
• Triple bonds consist of one $\sigma$ and two $\pi$ bonds.

Because $\pi$ bonds restrict rotation, molecules with double bonds often have more rigid shapes than molecules connected only by single bonds.

Molecular polarity: combining bond polarity and geometry

A molecule’s overall polarity depends on both:

1. Individual bond dipoles (from electronegativity differences)
2. Molecular geometry (whether dipoles cancel)

A molecule can have polar bonds but be nonpolar overall if the geometry is symmetric enough for dipole cancellation. Carbon dioxide is a classic example: each C=O bond is polar, but the linear geometry makes the net dipole zero. Water, in contrast, has polar O-H bonds and a bent geometry, so the dipoles do not cancel, producing a polar molecule.

When AP questions ask for polarity, they are testing whether you can move from Lewis structure to VSEPR geometry to vector-like dipole cancellation reasoning.

How bonding affects physical properties

Bonding type sets the baseline, but many property questions require a second step: intermolecular forces.

Ionic solids

• High melting points and low volatility
• Conduct when molten or aqueous
• Often soluble in polar solvents due to ion-dipole interactions

Molecular covalent substances

Properties vary widely because intermolecular forces vary:

• London dispersion forces exist in all molecules and increase with molar mass and surface area.
• Dipole-dipole forces occur in polar molecules.
• Hydrogen bonding (a strong dipole-dipole case) occurs when H is bonded to N, O, or F and interacts with lone pairs on N, O, or F.

These forces influence boiling point, melting point, viscosity, and solubility. For example, a polar molecule with hydrogen bonding generally has a higher boiling point than a similarly sized nonpolar molecule, because more energy is required to separate molecules.

Metallic solids

• Conduct electricity as solids
• Usually malleable and ductile
• Melting points vary, but bonding is collective rather than discrete molecules

Network covalent solids

• Extremely high melting points
• Hard and often insoluble
• Poor conductors (with notable exceptions such as graphite, where delocalized electrons enable conductivity)

Putting it all together for AP Chemistry

A strong bonding answer in AP Chemistry follows a predictable chain of reasoning:

1. Determine bonding type and write a correct Lewis structure.
2. Use VSEPR to predict geometry and approximate bond angles.
3. Assign hybridization and identify $\sigma$ and $\pi$ bonding where relevant.
4. Decide molecular polarity using geometry-based dipole cancellation.
5. Connect structure to macroscopic properties using bonding category and intermolecular forces.

Chemical bonding is the bridge between microscopic electron models and the tangible behavior of matter. Mastering that bridge is what turns chemistry from a set of rules into a predictive science.