Unit 2 of AP Chemistry: Molecular and Ionic Bonding.
This unit covers ionic bonds, covalent bonds, Lewis structures and VSEPR — essential concepts for AP Chemistry. Use our interactive study games to test your understanding, or review questions in traditional format below.
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This unit covers ionic bonds, covalent bonds, Lewis structures and VSEPR — essential concepts for AP Chemistry. Use our interactive study games to test your understanding, or review questions in traditional format below.
Key Concepts Breakdown
1 Ionic Bonds
Ionic bonds form when electrons are transferred from a metal to a nonmetal, creating oppositely charged ions that attract electrostatically. Students must understand how electronegativity difference drives ion formation and how lattice energy relates to ionic compound stability. Predicting formulas, naming compounds, and comparing properties based on ion charge and radius are all tested.
Key Points
- Form between elements with large electronegativity differences (typically >1.7); metal loses electrons (cation), nonmetal gains electrons (anion)
- Lattice energy increases with higher ion charge and smaller ionic radius — directly affects melting point and solubility
- Ionic compounds are brittle, have high melting points, and conduct electricity only when dissolved or melted
- Charge balance determines the formula: the total positive charge must equal total negative charge
Which has a higher melting point: NaF or MgO? Explain.
MgO has a higher melting point because Mg²⁺ and O²⁻ carry charges of 2+ and 2−, compared to Na⁺ and F⁻ with charges of only 1+ and 1−. Greater ion charges produce stronger electrostatic attraction and higher lattice energy, requiring more thermal energy to overcome. Additionally, Mg²⁺ and O²⁻ have smaller ionic radii than Na⁺ and F⁻, further increasing lattice energy.
2 Covalent Bonds
Covalent bonds form when two nonmetals share electrons to achieve stable electron configurations, and bond properties (length, strength, polarity) are directly tested. Students must distinguish between nonpolar covalent, polar covalent, and ionic bonds using electronegativity differences. Bond order (single, double, triple) inversely correlates with bond length and directly correlates with bond energy.
Key Points
- Electronegativity difference 0–0.4: nonpolar covalent; 0.4–1.7: polar covalent; >1.7: ionic
- Triple bonds are shorter and stronger than double bonds, which are shorter and stronger than single bonds
- Bond polarity creates partial charges (δ+ and δ−); molecular polarity depends on both bond polarity and geometry
- Coordinate (dative) covalent bonds occur when one atom donates both electrons (e.g., NH₃ → BF₃ adducts)
Rank the following bonds from longest to shortest: C–C, C=C, C≡C.
Bond length decreases as bond order increases because more shared electron pairs pull the nuclei closer together. Therefore, the order from longest to shortest is C–C > C=C > C≡C, with approximate lengths of 154 pm, 134 pm, and 120 pm. Correspondingly, bond energy increases in the reverse order: C–C (347 kJ/mol) < C=C (614 kJ/mol) < C≡C (839 kJ/mol).
3 Lewis Structures
Lewis structures show valence electrons as dots and bonds, and must satisfy the octet rule for most atoms (with defined exceptions). Students are tested on drawing correct structures, calculating formal charge to identify the best resonance structure, and recognizing when expanded octets or electron-deficient atoms apply. Resonance structures and delocalization are also high-frequency exam topics.
Key Points
- Formal charge = (valence electrons) − (lone pair electrons) − ½(bonding electrons); best structure minimizes formal charges and places negative formal charge on the more electronegative atom
- Octet rule exceptions: H and He (duet); B and Be (electron deficient, 6 or 4 electrons); Period 3+ elements (expanded octet, e.g., SF₆, PCl₅)
- Resonance structures occur when electrons can be delocalized across equivalent positions; actual bond is an average (e.g., O₃, NO₃⁻)
- Count total valence electrons carefully, adjusting for charge: add 1e⁻ per negative charge, subtract 1e⁻ per positive charge
Draw the Lewis structure of SO₃ and determine the formal charges on each atom in the most stable structure.
SO₃ has 24 total valence electrons (6 from S + 6×3 from O). Placing S in the center with three double bonds to each O gives each oxygen 4 lone-pair electrons and each S–O bond 4 bonding electrons. Formal charge on S = 6 − 0 − ½(12) = 0; formal charge on each O = 6 − 4 − ½(4) = 0. This structure with all formal charges at zero is most stable, and three equivalent resonance structures describe the delocalized bonding.
4 VSEPR Theory
VSEPR (Valence Shell Electron Pair Repulsion) predicts molecular geometry based on minimizing repulsion between all electron groups (bonding pairs and lone pairs) around a central atom. Students must distinguish between electron geometry (all groups) and molecular geometry (only atoms), and predict bond angles including deviations caused by lone pairs. Polarity of the molecule follows from geometry and bond polarity.
Key Points
- Lone pairs repel more strongly than bonding pairs, compressing bond angles below ideal values (e.g., H₂O: 104.5° instead of 109.5°)
- Key geometries to memorize by electron groups: 2=linear, 3=trigonal planar, 4=tetrahedral, 5=trigonal bipyramidal, 6=octahedral
- Molecular polarity requires both polar bonds AND an asymmetric arrangement; symmetric molecules (CO₂, BF₃, CCl₄) are nonpolar despite polar bonds
- For trigonal bipyramidal electron geometry, lone pairs occupy equatorial positions to minimize 90° repulsions (e.g., ClF₃ is T-shaped)
Predict the molecular geometry and bond angle of NH₃, and explain why the bond angle is not exactly 109.5°.
NH₃ has 4 electron groups around nitrogen (3 bonding pairs + 1 lone pair), giving a tetrahedral electron geometry but a trigonal pyramidal molecular geometry. The ideal tetrahedral bond angle is 109.5°, but the lone pair exerts greater repulsion than the N–H bonding pairs, compressing the H–N–H angle to approximately 107°. This deviation is a direct consequence of lone pair–bonding pair repulsion being stronger than bonding pair–bonding pair repulsion.
Questions, answered.
What is Molecular and Ionic Bonding?
Molecular and Ionic Bonding is Unit 2 of AP Chemistry, covering ionic bonds, covalent bonds, Lewis structures and VSEPR.
How to study for AP Chemistry Unit 2?
Start with the Quick Summary above, review the Key Concepts, then test yourself with our interactive study games. Aim for 80%+ accuracy before moving on.
How many questions are in this unit?
This unit has 30+ review questions across 5 different game modes.