Review AP Chem Unit 2 to understand how bonding type, Lewis structures, and molecular geometry connect atomic-level structure to observable properties. This unit builds the foundation for predicting polarity, intermolecular forces, and material behavior across the rest of the course.
Use the topic guides, practice questions, and FRQ practice available on Fiveable to work through every concept from electronegativity to VSEPR geometry.
Unit 2 asks you to explain why compounds look and behave the way they do at the particle level. Every topic connects atomic identity to structure, and structure to properties.
Topics 2.1 and 2.2 establish that electronegativity differences determine whether a bond is nonpolar covalent, polar covalent, ionic, or metallic, and that a potential energy versus internuclear distance graph reveals equilibrium bond length and bond energy for any pair of atoms.
Topics 2.3 and 2.4 extend bonding to bulk solids. Ionic solids arrange cations and anions in a 3-D lattice governed by Coulomb's law, while metallic solids use a sea of delocalized electrons to explain conductivity and malleability. Alloys are either interstitial or substitutional depending on atomic radius differences.
Topics 2.5, 2.6, and 2.7 form a three-step sequence: draw a Lewis structure, refine it with resonance and formal charge, then apply VSEPR to assign geometry, bond angles, hybridization, and molecular polarity.
Every macroscopic property in Unit 2, from melting point to electrical conductivity to molecular polarity, traces back to the arrangement of particles and the forces between them. Coulomb's law is the quantitative thread connecting ionic lattice energy, bond strength on a potential energy curve, and the repulsions that VSEPR theory uses to predict geometry.
Electronegativity trends across the periodic table determine whether atoms form nonpolar covalent, polar covalent, ionic, or metallic bonds. Bond type predicts properties like melting point, conductivity, and polarity.
A potential energy versus internuclear distance graph shows equilibrium bond length and bond energy. Higher bond order means shorter length and greater energy. Coulomb's law explains how ion charge and radius affect interaction strength.
Ionic solids form 3-D lattices that maximize cation-anion attractions and minimize like-charge repulsions. Lattice energy increases with higher ion charges and smaller ionic radii, raising melting point and hardness.
Metallic bonding is modeled as positive ions in a sea of delocalized electrons, explaining conductivity and malleability. Alloys are interstitial (small atoms in gaps, like steel) or substitutional (similar-sized atoms swapped in, like brass).
Lewis diagrams distribute valence electrons as bonding pairs and lone pairs following the octet rule. Exceptions include electron-deficient molecules like BF3 and expanded octets for period 3 elements like SF6.
When multiple valid Lewis structures exist, the molecule is a resonance hybrid with delocalized electrons. Formal charge (valence minus lone pairs minus half bonding electrons) identifies the best structure among nonequivalent options.
VSEPR arranges electron domains to minimize repulsion, predicting geometry, bond angles, and hybridization. Lone pairs compress angles below ideal values. Molecular polarity depends on both bond polarity and geometry.
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Review Structure of Metals and Alloys with attention to how the concept appears in AP-style source and evidence questions.
Review Intramolecular Force and Potential Energy with attention to how the concept appears in AP-style source and evidence questions.
Review Resonance and Formal Charge with attention to how the concept appears in AP-style source and evidence questions.
Electronegativity increases left to right across a period and decreases down a group, following trends in effective nuclear charge and atomic radius. The difference in electronegativity between two bonded atoms predicts bond type: similar values give nonpolar covalent bonds (C-H is the classic example), larger differences give polar covalent bonds with partial charges (delta+ and delta-), and very large differences produce ionic bonds. Metallic bonding occurs between metal atoms and is modeled as a sea of delocalized electrons. Bond type predicts properties such as melting point, conductivity, and solubility.
| Bond Type | Electronegativity Difference | Electron Behavior | Example | Key Property |
|---|---|---|---|---|
| Nonpolar covalent | ~0 | Equally shared | Cl2, C-H | Low conductivity, low melting point |
| Polar covalent | Moderate | Unequally shared | H-Cl, H2O | Dipole moment, partial charges |
| Ionic | Large | Transferred | NaCl | High melting point, conducts when dissolved |
| Metallic | N/A (metals) | Delocalized sea | Fe, Cu | Conducts electricity, malleable, ductile |
A potential energy versus internuclear distance graph shows how the energy of two atoms changes as they approach each other. At large distances, energy is near zero. As atoms approach, attraction lowers energy to a minimum, which marks the equilibrium bond length. Closer approach causes repulsion and energy rises steeply. The depth of the energy well equals the bond energy. Higher bond order (double or triple) means shorter bond length and deeper energy well. Coulomb's law explains ionic interactions: larger charges and smaller interionic distances produce stronger attractions and higher lattice energy.
| Bond Type | Relative Bond Length | Relative Bond Energy |
|---|---|---|
| C-C single bond | Longest | Lowest |
| C=C double bond | Intermediate | Intermediate |
| C triple bond C | Shortest | Highest |
Ionic solids consist of cations and anions arranged in a repeating three-dimensional lattice. The arrangement maximizes attractions between opposite charges and minimizes repulsions between like charges, consistent with Coulomb's law. Lattice energy measures the strength of the ionic solid: higher ion charges and smaller ionic radii both increase lattice energy, which raises melting point and hardness. Ionic solids do not conduct electricity in the solid state because ions are locked in place, but they conduct when melted or dissolved because ions become mobile. Particulate models of ionic solids must show alternating cations and anions in a regular array.
| Compound | Ion Charges | Relative Interionic Distance | Relative Lattice Energy |
|---|---|---|---|
| NaCl | +1 / -1 | Larger | Lower |
| MgO | +2 / -2 | Smaller | Higher |
Metallic solids are modeled as an array of positive metal ions surrounded by a sea of delocalized valence electrons. This model explains electrical and thermal conductivity (electrons move freely), malleability and ductility (layers of ions slide without breaking bonds), and the generally high melting points of metals. Alloys are mixtures of metals with other elements. Interstitial alloys form when small atoms (like carbon in steel) fit into the gaps between larger metal atoms. Substitutional alloys form when atoms of similar radius replace each other in the lattice (like zinc replacing copper in brass).
| Alloy Type | Atom Size Relationship | Mechanism | Example |
|---|---|---|---|
| Interstitial | Solute much smaller than host | Small atoms fill gaps in lattice | Steel (C in Fe) |
| Substitutional | Solute similar size to host | Solute atoms replace host atoms | Brass (Zn for Cu) |
Lewis diagrams show how valence electrons are distributed in a molecule or polyatomic ion as bonding pairs (lines) and lone pairs (dots). The standard procedure is: count total valence electrons (add one per negative charge, subtract one per positive charge), connect atoms with single bonds, complete octets on terminal atoms, place remaining electrons on the central atom, and convert lone pairs to multiple bonds if the central atom lacks an octet. Hydrogen always gets two electrons (duet rule). Exceptions include electron-deficient molecules like BF3 (six electrons on boron) and expanded octets for period 3 and beyond elements like SF6.
When more than one valid Lewis structure can be drawn with the same connectivity, the molecule is described by resonance. The actual molecule is a resonance hybrid with bond orders between whole numbers; for example, the three bonds in NO3- are each 1.33, not alternating single and double. Formal charge is calculated as: valence electrons minus lone pair electrons minus half the bonding electrons. The best Lewis structure minimizes formal charges overall and places any negative formal charge on the more electronegative atom. Formal charge also reveals limitations of the Lewis model: odd-electron species like NO cannot satisfy the octet rule for all atoms.
VSEPR theory predicts molecular geometry by arranging electron domains around a central atom to minimize repulsion. Lone pairs repel more strongly than bonding pairs, which compresses bond angles below ideal values. The electron domain geometry determines hybridization: two domains give sp (linear, 180 degrees), three give sp2 (trigonal planar, 120 degrees), four give sp3 (tetrahedral, 109.5 degrees). Molecular geometry is named by atom positions only, not lone pairs, so four domains with one lone pair gives trigonal pyramidal (NH3), and four domains with two lone pairs gives bent (H2O). Sigma bonds form from end-to-end orbital overlap; pi bonds form from side-to-side overlap and are present in all double and triple bonds. A molecule has a net dipole moment only when polar bonds do not cancel due to molecular geometry.
| Electron Domains | Electron Geometry | Example Molecular Geometry | Hybridization | Ideal Bond Angle |
|---|---|---|---|---|
| 2 | Linear | Linear (BeCl2) | sp | 180° |
| 3 | Trigonal planar | Bent with 1 lone pair (SO2) | sp2 | 120° |
| 4 | Tetrahedral | Trigonal pyramidal with 1 lone pair (NH3) | sp3 | 109.5° |
| 4 | Tetrahedral | Bent with 2 lone pairs (H2O) | sp3 | <109.5° |
| 5 | Trigonal bipyramidal | Seesaw with 1 lone pair (SF4) | sp3d | 90°/120° |
Try AP-style multiple-choice questions and written prompts after you review the notes.
| Term | Definition |
|---|---|
| Electronegativity | A measure of an atom's ability to attract shared electrons in a bond; increases left to right across a period and decreases down a group, driving bond type classification. |
| Coulomb's Law | States that electrostatic force is proportional to the product of ion charges and inversely proportional to the square of the distance between them; used to explain bond strength, lattice energy, and potential energy curves. |
| Bond Energy | The energy required to break a chemical bond and separate the atoms; represented by the depth of the well on a potential energy versus internuclear distance graph. |
| Bond Length | The average distance between two bonded nuclei at the energy minimum; shorter for higher bond orders and smaller atomic radii. |
| Bond Order | The number of shared electron pairs between two atoms; single = 1, double = 2, triple = 3. Higher bond order means shorter length and greater bond energy. |
| Lattice Energy | The energy released when gaseous ions form one mole of ionic solid; increases with larger ion charges and smaller ionic radii. |
| Crystal Lattice | The regular, repeating 3-D arrangement of cations and anions in an ionic solid that maximizes attractions and minimizes repulsions between ions. |
| metallic bonding | A bonding model in which delocalized valence electrons move freely among positive metal ions, explaining conductivity, malleability, and ductility. |
| Substitutional Alloy | An alloy in which atoms of similar radius replace host metal atoms in the lattice; brass (zinc replacing copper) is the standard example. |
| Lewis Structures | Diagrams showing the arrangement of valence electrons in a molecule as bonding pairs and lone pairs, following the octet rule with defined exceptions. |
| Resonance Structures | Two or more valid Lewis structures for the same molecule that differ only in electron placement; the actual molecule is a hybrid with delocalized electrons and fractional bond orders. |
| lone pair | A pair of valence electrons not involved in bonding; occupies more space than a bonding pair and compresses bond angles in VSEPR geometry. |
| Molecular Geometry | The three-dimensional arrangement of atoms in a molecule, named by bonded atom positions only; determined from the Lewis structure using VSEPR theory. |
| dipole moment | A measure of charge separation in a molecule; a net dipole exists only when polar bonds are arranged asymmetrically so their vectors do not cancel. |
| sp3 hybridization | Hybridization of one s and three p orbitals forming four equivalent orbitals; associated with four electron domains, tetrahedral electron geometry, and 109.5 degree bond angles. |
VSEPR names molecular geometry based on atom positions only, not lone pairs. A molecule with four electron domains and one lone pair is trigonal pyramidal, not tetrahedral. Always count lone pairs for hybridization but name the shape from bonded atoms only.
For polyatomic anions like NO3-, add electrons equal to the negative charge before drawing the Lewis structure. For cations like NH4+, subtract electrons. Skipping this step produces structures with the wrong number of electrons.
The molecule does not switch between resonance forms. The actual bond order in NO3- is 1.33 for all three N-O bonds simultaneously. Resonance structures are a modeling tool, not a description of the molecule flipping between states.
Molecular polarity requires both polar bonds and an asymmetric geometry. CO2 has two polar C=O bonds but is linear, so the dipoles cancel and the molecule is nonpolar. Always check geometry before assigning a net dipole.
The equilibrium bond length is the internuclear distance at the lowest point on the curve, not where the curve crosses zero energy. The bond energy is the depth of the well from the minimum to the zero-energy baseline, not the height of the repulsive wall.
AP Chemistry frequently asks you to draw or interpret a particulate-level diagram of an ionic solid, metallic solid, or molecule and then justify a property such as melting point, conductivity, or solubility using the model. For Unit 2, practice connecting your diagram to Coulomb's law or the sea of electrons model in one or two sentences, because explanation tasks require explicit reasoning, not just a correct picture.
A common multi-part task starts with constructing a Lewis structure, then asks for molecular geometry, bond angles, hybridization, and whether the molecule is polar. Each step depends on the previous one, so an error in the Lewis structure propagates. Practice the full chain for molecules with lone pairs on the central atom, such as NH3, H2O, SO2, and SF4, where geometry and polarity are not immediately obvious.
Unit 2 skills appear in questions that ask you to rank compounds by melting point, bond energy, or lattice energy and support your ranking with evidence. The expected reasoning pattern is to identify the relevant variable (ion charge, ionic radius, bond order, or electronegativity difference) and apply Coulomb's law or bond order logic explicitly. Avoid simply stating that one compound has a higher melting point without connecting it to a structural cause.
Open the individual guides for Unit 2 when you want a closer review of one topic.
browse guidesPractice free-response reasoning and compare your answer with scoring guidance.
practice FRQsWatch past review streams filtered to Unit 2 when you want a video walkthrough.
open videosUse unit cheatsheets for a quick visual review after you work through the notes.
open cheatsheetsEstimate your broader AP score goal after you review the course and exam format.
open calculatorAP Chem Unit 2 covers 7 topics: Types of Chemical Bonds, Intramolecular Force and Potential Energy, Structure of Ionic Solids, Structure of Metals and Alloys, Lewis Diagrams, Resonance and Formal Charge, and VSEPR and Hybridization. Together these topics explain how atomic arrangement and bonding forces determine the properties of compounds. See the full topic breakdown at /ap-chem/unit-2.
AP Chem Unit 2 makes up 7-9% of the AP exam. That weight covers everything from types of chemical bonds and Lewis diagrams to resonance, formal charge, VSEPR, and hybridization. It's a focused unit, but the concepts show up again in later units, so a strong grasp here pays off across the whole exam.
The AP Chem Unit 2 progress check includes both MCQ and FRQ parts drawn from all 7 topics in the unit. MCQ questions test your ability to identify bond types, interpret potential energy diagrams, and apply VSEPR theory to predict molecular geometry. FRQ prompts typically ask you to draw Lewis diagrams, assign formal charges, explain resonance structures, or justify hybridization for a given molecule. Practicing these question types before the progress check is the best way to spot gaps. Find matched practice at /ap-chem/unit-2.
AP Chem Unit 2 FRQs most often come from Lewis Diagrams (Topic 2.5), Resonance and Formal Charge (Topic 2.6), and VSEPR and Hybridization (Topic 2.7). A typical question gives you a molecule or ion and asks you to draw the Lewis structure, identify resonance structures, assign formal charges, predict geometry using VSEPR, and state the hybridization of the central atom. To practice, work through each step in writing rather than just thinking through it, since partial credit depends on showing your reasoning clearly. You can find FRQ-style practice questions at /ap-chem/unit-2.
The best place to find AP Chem Unit 2 practice questions, including multiple-choice and FRQ-style prompts, is /ap-chem/unit-2. That page has resources covering all 7 topics, from chemical bonds and ionic solids to Lewis diagrams, resonance, VSEPR, and hybridization. For a practice test feel, work through MCQ sets timed and then review any question involving molecular geometry or bond polarity, since those show up most often.
Start AP Chem Unit 2 by building a solid foundation in chemical bonds and electronegativity, since those ideas run through every other topic. Then work through Lewis diagrams by hand until drawing them feels automatic, because resonance and formal charge both depend on that skill. Once Lewis structures click, VSEPR geometry and hybridization become much more manageable. A practical study plan looks like this: - Review bond types and potential energy diagrams (Topics 2.1-2.2) - Practice drawing Lewis diagrams and assigning formal charges (Topics 2.5-2.6) - Apply VSEPR to predict geometry and link geometry to hybridization (Topic 2.7) - Connect ionic solids and metallic structures to the broader bonding picture (Topics 2.3-2.4) Spacing your practice over several sessions and writing out every step, rather than just reading, will lock in the reasoning College Board expects on FRQs.
The key graph in AP Chem Unit 2 is the intramolecular potential energy curve (Topic 2.2). It shows how potential energy changes as two atoms move closer together, with a minimum at the equilibrium bond length. You need to read off bond length and bond energy from the curve, and explain how those values shift for different bond types (single vs. double vs. triple) or atoms with different electronegativity. Questions may give you two curves and ask you to compare bond strength or bond length, so practice interpreting the shape rather than just memorizing what it looks like.