---
title: "AP Chem Unit 2 Review: Compound Structure and Properties"
description: "AP Chemistry Unit 2 covers Types of Chemical Bonds and Intramolecular Force and Potential Energy. Study guides, practice questions, and key terms."
canonical: "https://fiveable.me/ap-chem/unit-2"
type: "unit"
subject: "AP Chemistry"
unit: "Unit 2 – Compound Structure and Properties"
---

# AP Chem Unit 2 Review: Compound Structure and Properties

## Overview

Unit 2 covers how atoms bond, how those bonds store energy, and how the resulting structures determine properties. You will move from electronegativity trends through ionic and metallic solid models, then build Lewis structures, apply resonance and formal charge, and finally use VSEPR theory to predict three-dimensional molecular geometry and polarity.

## AP CED Alignment

This unit hub is organized around AP Course and Exam Description topics, skills, and exam task types when they are available in the source data.
- Topic 2.1: Types of Chemical Bonds
- Topic 2.2: Intramolecular Force and Potential Energy
- Topic 2.3: Structure of Ionic Solids
- Topic 2.4: Structure of Metals and Alloys
- Topic 2.5: Lewis Diagrams
- Topic 2.6: Resonance and Formal Charge
- Topic 2.7: VSEPR and Hybridization
- Practice 6 - Argumentation
- Practice 5 - Mathematical Routines
- FRQ 6 – Short Answer
- FRQ 4 – Short Answer
- FRQ 7 – Short Answer

## Topics

- [Topic 2.1: Types of Chemical Bonds](/ap-chem/unit-2/types-chemical-bonds/study-guide/syV9vHl1BKGwM5tnaOt0): 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.
- [Topic 2.2: Intramolecular Force and Potential Energy](/ap-chem/unit-2/intramolecular-force-potential-energy/study-guide/2ybBItwMrCkB9MAOO9lC): 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.
- [Topic 2.3: Structure of Ionic Solids](/ap-chem/unit-2/structure-ionic-solids/study-guide/3khaTI6A3tdnaMTMJy2u): 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.
- [Topic 2.4: Structure of Metals and Alloys](/ap-chem/unit-2/structure-metals-alloys/study-guide/yTmM9j22ZoTxtr6TCC3p): 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).
- [Topic 2.5: Lewis Diagrams](/ap-chem/unit-2/lewis-diagrams/study-guide/KjqTRYr5TVr2C3Be3u0J): 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.
- [Topic 2.6: Resonance and Formal Charge](/ap-chem/unit-2/resonance-formal-charge/study-guide/rxtiCdBiDF6V9V7S2UEi): 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.
- [Topic 2.7: VSEPR and Hybridization](/ap-chem/unit-2/vsepr-bond-hybridization/study-guide/OslsAmh8LcVoqbpnjPAu): 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.

## Hardest Topics And Analytics

Snapshot: practice snapshot
This snapshot uses Fiveable practice activity to show where students tend to miss questions and which review moves are worth prioritizing first.
- **66% average MCQ accuracy** (Across 19k multiple-choice practice attempts for this unit.)
- **19k MCQ attempts** (Practice activity included in this snapshot.)
- **57% average FRQ score** (Across 24 scored free-response attempts for this unit.)
- **Topic 2.4: Structure of Metals and Alloys**: 34% MCQ miss rate across 2016 attempts. Review Structure of Metals and Alloys with attention to how the concept appears in AP-style source and evidence questions.
- **Topic 2.2: Intramolecular Force and Potential Energy**: 33% MCQ miss rate across 4669 attempts. Review Intramolecular Force and Potential Energy with attention to how the concept appears in AP-style source and evidence questions.
- **Topic 2.6: Resonance and Formal Charge**: 29% MCQ miss rate across 1879 attempts. Review Resonance and Formal Charge with attention to how the concept appears in AP-style source and evidence questions.

## Review Notes

### Topic 2.1: Types of Chemical Bonds

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.

- **Nonpolar covalent**: Shared electrons between atoms of similar electronegativity; C-H bonds are treated as effectively nonpolar on the AP exam.
- **Polar covalent**: Unequal sharing creates partial charges; the more electronegative atom carries delta-.
- **Ionic bonding**: Electron transfer between atoms with large electronegativity differences; results in cation-anion attraction.
- **Metallic bonding**: Delocalized valence electrons move freely among positive metal ions, explaining conductivity and malleability.
- **Electronegativity trend**: Increases across a period (more protons, smaller radius) and decreases down a group (more shielding, larger radius).

**Checkpoint:** Given two elements, can you predict bond type from their electronegativity values and explain the resulting property differences?

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

### Topic 2.2: Intramolecular Force and Potential Energy

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.

- **Equilibrium bond length**: The internuclear distance at the minimum of the potential energy curve; where attractive and repulsive forces balance.
- **Bond energy**: The depth of the potential energy well; energy required to separate the two atoms to infinite distance.
- **Bond order effect**: Double and triple bonds are shorter and have larger bond energies than single bonds between the same atoms.
- **Coulomb's law**: Force is proportional to the product of ion charges and inversely proportional to the square of the distance; larger charges and smaller radii mean stronger ionic interactions.
- **Repulsive region**: At distances shorter than the equilibrium bond length, electron cloud overlap causes energy to rise sharply.

**Checkpoint:** Can you read a potential energy curve to identify equilibrium bond length and bond energy, and predict how those values change with bond order or ion charge?

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

### Topic 2.3: Structure of Ionic Solids

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.

- **Crystal lattice**: The regular, repeating 3-D arrangement of cations and anions that minimizes electrostatic potential energy.
- **Lattice energy**: Energy released when gaseous ions form one mole of ionic solid; increases with higher charges and smaller ionic radii.
- **Coulomb's law in ionic solids**: Stronger attractions come from larger ion charges (q1 and q2) and shorter interionic distances (r).
- **Conductivity rule**: Ionic solids conduct only when ions are free to move: in the liquid state or in aqueous solution.
- **Particulate model**: A diagram showing alternating cations and anions in a lattice; must reflect charge ratio and relative sizes.

**Checkpoint:** Can you draw a particulate model of an ionic solid and use Coulomb's law to compare lattice energies of two compounds, such as MgO versus NaCl?

Compound | Ion Charges | Relative Interionic Distance | Relative Lattice Energy
--- | --- | --- | ---
NaCl | +1 / -1 | Larger | Lower
MgO | +2 / -2 | Smaller | Higher

### Topic 2.4: Structure of Metals and Alloys

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).

- **Sea of electrons model**: Delocalized valence electrons move throughout the metallic lattice, holding positive metal ions together.
- **Malleability and ductility**: Layers of metal ions can shift without disrupting the electron sea, so metals bend and draw into wires.
- **Interstitial alloy**: Small atoms occupy gaps in the host metal lattice; steel (carbon in iron) is the standard example.
- **Substitutional alloy**: Atoms of similar radius replace host atoms in the lattice; brass (zinc replacing copper) is the standard example.

**Checkpoint:** Can you classify an alloy as interstitial or substitutional based on atomic radius information, and explain how the sea of electrons model accounts for metallic conductivity?

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)

### Topic 2.5: Lewis Diagrams

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.

- **Valence electron count**: Sum of valence electrons from all atoms, adjusted for ion charge, is the total electron budget for the Lewis structure.
- **Octet rule**: Most atoms in a Lewis structure are surrounded by eight electrons; hydrogen follows the duet rule with two.
- **Lone pair**: A pair of valence electrons not involved in bonding; affects geometry and reactivity.
- **Expanded octet**: Period 3 and heavier elements can hold more than eight electrons; SF6 has twelve electrons on sulfur.
- **Electron-deficient molecules**: Molecules like BF3 where the central atom has fewer than eight electrons in the most reasonable Lewis structure.

**Checkpoint:** Can you draw a correct Lewis structure for a molecule like SO2 or a polyatomic ion like NO3-, including proper electron count and placement of lone pairs?

### Topic 2.6: Resonance and Formal Charge

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.

- **Resonance hybrid**: The actual electron distribution in a molecule described by resonance; electrons are delocalized across equivalent bonds.
- **Formal charge formula**: Formal charge = valence electrons - lone pair electrons - (1/2)(bonding electrons); best structure minimizes these values.
- **Equivalent resonance structures**: Structures with the same energy that differ only in placement of multiple bonds; CO3^2- and NO3- are classic examples.
- **Nonequivalent resonance structures**: Structures with different formal charges; use formal charge to select the dominant contributor.
- **Odd-electron species**: Molecules like NO with an odd number of valence electrons cannot satisfy the octet rule for every atom.

**Checkpoint:** Can you calculate formal charges for each atom in SO2 and use those values to identify the best Lewis structure among two or more options?

### Topic 2.7: VSEPR and Hybridization

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 domain**: Any region of electron density around a central atom: a single bond, double bond, triple bond, or lone pair each count as one domain.
- **Lone pair compression**: Lone pairs occupy more space than bonding pairs, reducing bond angles below the ideal geometry value.
- **Hybridization shortcut**: Count electron domains on the central atom: 2 = sp, 3 = sp2, 4 = sp3.
- **Sigma and pi bonds**: Every bond contains one sigma bond; double bonds add one pi bond, triple bonds add two pi bonds. Pi bonds restrict rotation.
- **Molecular polarity**: A molecule is polar only when bond dipoles do not cancel; geometry must be asymmetric. CO2 is nonpolar (linear); H2O is polar (bent).

**Checkpoint:** Given a molecular formula, can you draw the Lewis structure, assign VSEPR geometry and bond angles, identify hybridization, and determine whether the molecule has a net dipole moment?

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°

## Study Guides

- [2.3 Structure of Ionic Solids](/ap-chem/unit-2/structure-ionic-solids/study-guide/3khaTI6A3tdnaMTMJy2u)
- [2.5 Lewis Diagrams](/ap-chem/unit-2/lewis-diagrams/study-guide/KjqTRYr5TVr2C3Be3u0J)
- [2.7 VSEPR and Bond Hybridization](/ap-chem/unit-2/vsepr-bond-hybridization/study-guide/OslsAmh8LcVoqbpnjPAu)
- [2.6 Resonance and Formal Charge](/ap-chem/unit-2/resonance-formal-charge/study-guide/rxtiCdBiDF6V9V7S2UEi)
- [2.4 Structure of Metals and Alloys](/ap-chem/unit-2/structure-metals-alloys/study-guide/yTmM9j22ZoTxtr6TCC3p)
- [2.2 Intramolecular Force and Potential Energy ](/ap-chem/unit-2/intramolecular-force-potential-energy/study-guide/2ybBItwMrCkB9MAOO9lC)
- [2.1 Types of Chemical Bonds](/ap-chem/unit-2/types-chemical-bonds/study-guide/syV9vHl1BKGwM5tnaOt0)

## Practice Preview

### Multiple-choice practice

- **AP-style practice question**: Practice 6 - Argumentation | A student determines the copper content of a brass alloy (Cu and Zn) by measuring its density. Copper ($8.96\ g/cm^3$) is denser than zinc ($7.14\ g/cm^3$). The student measures the mass correctly but measures the volume by water displacement using a graduated cylinder where air bubbles adhere to the submerged alloy sample. Which of the following best explains the effect on the result?
- **AP-style practice question**: Practice 5 - Mathematical Routines | In a molecule of phosphorus pentafluoride ($$PF_5$$), calculate the difference in degrees between the bond angle of two equatorial fluorine atoms and the bond angle formed by an axial fluorine and an equatorial fluorine.
- **AP-style practice question**: Practice 5 - Mathematical Routines | Xenon tetrafluoride ($$XeF_4$$) is a nonpolar molecule with a square planar geometry. Calculate the number of lone pairs of electrons on the central xenon atom based on its Lewis electron-dot diagram.
- **AP-style practice question**: Practice 5 - Mathematical Routines | Calculate the formal charge on the central oxygen atom in the ozone molecule, $$O_3$$, which has a resonance structure containing one double bond and one single bond.
- **AP-style practice question**: Practice 5 - Mathematical Routines | Calculate the enthalpy change ($$\Delta H$$) for the reaction $$2H_2(g) + O_2(g) \rightarrow 2H_2O(g)$$ using the following bond energies: H-H ($$436$$ kJ/mol), O=O ($$498$$ kJ/mol), and O-H ($$463$$ kJ/mol).
- **AP-style practice question**: Practice 5 - Mathematical Routines | Acetonitrile has the chemical formula $$CH_3CN$$. Based on its complete Lewis electron-dot diagram, determine the total number of sigma ($$\sigma$$) and pi ($$\pi$$) bonds in one molecule.

### FRQ practice

- **Carbon tetrahalide molecular structure and properties**: FRQ 6 – Short Answer | Carbon tetrahalide molecular structure and properties
- **Ionic bonding, lattice energy, bond strength comparison**: FRQ 4 – Short Answer | Ionic bonding, lattice energy, bond strength comparison
- **Urea bonding, resonance, bond length**: FRQ 7 – Short Answer | Urea bonding, resonance, bond length

## Key Terms

- **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.

## Common Mistakes

- **Confusing electron geometry with molecular geometry**: 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.
- **Forgetting to adjust valence electron count for ion charge**: 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.
- **Treating resonance structures as real alternating structures**: 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.
- **Assuming a molecule with polar bonds is always polar**: 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.
- **Misreading the potential energy curve minimum**: 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.

## Exam Connections

- **Particulate model and justification tasks**: 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.
- **Lewis structure to geometry to polarity chain**: 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.
- **Comparing bond or lattice strength with evidence**: 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.

## Final Review Checklist

- **Final Unit 2 review checklist**: Use this list to confirm you can handle every major skill in Unit 2 before exam day.
- **Predict bond type from electronegativity**: Given two elements, classify the bond as nonpolar covalent, polar covalent, ionic, or metallic and connect that classification to at least one observable property.
- **Interpret a potential energy curve**: Identify equilibrium bond length and bond energy from a graph, and predict how those values shift with bond order or ion charge using Coulomb's law.
- **Compare ionic solids using lattice energy**: Rank compounds like NaCl and MgO by lattice energy, melting point, and hardness using ion charge and ionic radius as evidence.
- **Classify alloys and explain metallic properties**: Distinguish interstitial from substitutional alloys by atomic radius, and use the sea of electrons model to explain conductivity, malleability, and ductility.
- **Draw and refine Lewis structures**: Construct Lewis diagrams for molecules and polyatomic ions, apply resonance where needed, calculate formal charges, and select the best structure.
- **Apply VSEPR to predict geometry and polarity**: From a Lewis structure, assign electron domain geometry, molecular geometry, bond angles, hybridization, and determine whether the molecule has a net dipole moment.

## Study Plan

- **Step 1: Electronegativity, bond types, and potential energy (Topics 2.1-2.2)**: Review electronegativity trends on the periodic table and practice classifying bonds as nonpolar covalent, polar covalent, ionic, or metallic. Then sketch potential energy versus internuclear distance curves for pairs like Cl2 and Br2, labeling equilibrium bond length and bond energy. Use the Topic 2.1 and 2.2 guides on Fiveable to check your reasoning.
- **Step 2: Ionic solids and metallic structures (Topics 2.3-2.4)**: Draw particulate models of ionic solids like NaCl and MgO, then rank them by lattice energy using ion charge and radius. Next, practice explaining metallic properties with the sea of electrons model and classify alloy examples as interstitial or substitutional. The Topic 2.3 and 2.4 guides include worked particulate model examples.
- **Step 3: Lewis structures (Topic 2.5)**: Work through at least ten Lewis structure problems covering molecules, polyatomic ions, and exceptions like BF3 and SF6. Practice the full procedure: count valence electrons, connect atoms, complete octets, and add multiple bonds as needed. Check your work against the Topic 2.5 guide on Fiveable.
- **Step 4: Resonance and formal charge (Topic 2.6)**: Take your Lewis structures from Step 3 and identify any that require resonance, such as CO3^2- and NO3-. Calculate formal charges for each atom in at least five molecules and use those values to select the best structure among nonequivalent options. Review the Topic 2.6 guide for formal charge calculation practice.
- **Step 5: VSEPR, hybridization, and molecular polarity (Topic 2.7)**: For each molecule from Steps 3 and 4, assign electron domain geometry, molecular geometry, bond angles, and hybridization. Then determine whether the molecule has a net dipole moment by checking geometry for symmetry. Use the Topic 2.7 guide and available practice questions on Fiveable to test your full Lewis-to-geometry workflow.

## More Ways To Review

- [Topic study guides](/ap-chem/unit-2#topics)
- [FRQ practice](/ap-chem/frq-practice)
- [Cram archive videos](/cram-archives?subject=ap-chemistry&unit=unit-2)
- [Cheatsheets](/ap-chem/cheatsheets/unit-2)
- [Key terms](/ap-chem/key-terms)

## FAQs

### What topics are covered in AP Chem Unit 2?

AP 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).

### How much of the AP Chem exam is 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.

### What's on the AP Chem Unit 2 progress check (MCQ and FRQ)?

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).

### How do I practice AP Chem Unit 2 FRQs?

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](/ap-chem/unit-2).

### Where can I find AP Chem Unit 2 practice questions?

The best place to find AP Chem Unit 2 practice questions, including multiple-choice and FRQ-style prompts, is [/ap-chem/unit-2](/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.

### How should I study AP Chem Unit 2?

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.

### What graphs do I need to know for AP Chem Unit 2?

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.

## Structured Data

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{"@context":"https://schema.org","@type":"FAQPage","inLanguage":"en","mainEntity":[{"@type":"Question","@id":"https://fiveable.me/ap-chem/unit-2#what-topics-are-covered-in-ap-chem-unit-2","name":"What topics are covered in AP Chem Unit 2?","acceptedAnswer":{"@type":"Answer","text":"AP 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 <a href=\"/ap-chem/unit-2\">/ap-chem/unit-2</a>."}},{"@type":"Question","@id":"https://fiveable.me/ap-chem/unit-2#how-much-of-the-ap-chem-exam-is-unit-2","name":"How much of the AP Chem exam is Unit 2?","acceptedAnswer":{"@type":"Answer","text":"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."}},{"@type":"Question","@id":"https://fiveable.me/ap-chem/unit-2#whats-on-the-ap-chem-unit-2-progress-check-mcq-and-frq","name":"What's on the AP Chem Unit 2 progress check (MCQ and FRQ)?","acceptedAnswer":{"@type":"Answer","text":"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 <a href=\"/ap-chem/unit-2\">/ap-chem/unit-2</a>."}},{"@type":"Question","@id":"https://fiveable.me/ap-chem/unit-2#how-do-i-practice-ap-chem-unit-2-frqs","name":"How do I practice AP Chem Unit 2 FRQs?","acceptedAnswer":{"@type":"Answer","text":"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 <a href=\"/ap-chem/unit-2\">/ap-chem/unit-2</a>."}},{"@type":"Question","@id":"https://fiveable.me/ap-chem/unit-2#where-can-i-find-ap-chem-unit-2-practice-questions","name":"Where can I find AP Chem Unit 2 practice questions?","acceptedAnswer":{"@type":"Answer","text":"The best place to find AP Chem Unit 2 practice questions, including multiple-choice and FRQ-style prompts, is <a href=\"/ap-chem/unit-2\">/ap-chem/unit-2</a>. 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."}},{"@type":"Question","@id":"https://fiveable.me/ap-chem/unit-2#how-should-i-study-ap-chem-unit-2","name":"How should I study AP Chem Unit 2?","acceptedAnswer":{"@type":"Answer","text":"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)\n- Practice drawing Lewis diagrams and assigning formal charges (Topics 2.5-2.6)\n- Apply VSEPR to predict geometry and link geometry to hybridization (Topic 2.7)\n- 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."}},{"@type":"Question","@id":"https://fiveable.me/ap-chem/unit-2#what-graphs-do-i-need-to-know-for-ap-chem-unit-2","name":"What graphs do I need to know for AP Chem Unit 2?","acceptedAnswer":{"@type":"Answer","text":"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."}}]}
```