Unit 2 · 7–9% of the exam
Molecular and Ionic Compound Structure and Properties
Unit 2 explains how atoms are held together, whether by ionic, covalent or metallic bonding, and how to draw and use Lewis structures. VSEPR shapes, resonance and formal charge let you predict a molecule's geometry and polarity, which sets up the intermolecular forces in Unit 3.
VSEPR shape explorer
Choose bonding pairs and lone pairs, then rotate the 3D geometry and read the bond angles.
Topics in this unit
2.1
Types of Chemical Bonds
Bond type depends on the electronegativity difference: electrons are transferred (ionic), shared unequally (polar covalent) or shared equally (nonpolar covalent); metals share a sea of delocalized electrons.
2.2
Intramolecular Force and Potential Energy
A potential energy curve has its minimum at the bond length, where attraction and repulsion balance; the depth of the well is the bond energy.
2.3
Structure of Ionic Solids
Ionic solids are 3-D lattices of alternating cations and anions; larger charges and smaller ions give stronger Coulombic attraction and higher melting points.
2.4
Structure of Metals and Alloys
Metals are cations in a sea of delocalized electrons; substitutional alloys swap in similar-sized atoms, while interstitial alloys fill holes with small atoms like C, making the metal less malleable.
2.5
Lewis Diagrams
Lewis diagrams show how valence electrons are arranged as bonding pairs and lone pairs, usually giving each atom an octet (H gets two).
2.6
Resonance and Formal Charge
When one Lewis structure cannot describe the bonding, the real structure is a resonance hybrid; formal charges near zero identify the best structure.
2.7
VSEPR and Bond Hybridization
Electron domains around a central atom spread out to minimize repulsion, which sets the molecular geometry, bond angles, polarity and hybridization (sp, sp2, sp3).
Must know
- Formal charge = valence electrons − nonbonding electrons − ½(bonding electrons)
- Electron domains → geometry: 2 linear (180°), 3 trigonal planar (120°), 4 tetrahedral (109.5°), 5 trigonal bipyramidal, 6 octahedral
- Lone pairs compress bond angles: NH3 ≈ 107°, H2O ≈ 104.5°
- Hybridization from electron domains: 2 → sp, 3 → sp2, 4 → sp3
- A single bond is 1 σ; a double bond is 1 σ + 1 π; a triple bond is 1 σ + 2 π
- Higher bond order → shorter, stronger bond
- Lattice energy (Coulombic attraction) ∝ (q1q2)/r: bigger charges, smaller ions → stronger
- A molecule is polar if it has polar bonds whose dipoles do not cancel because of its shape
Common mix-ups
Resonance structures flip back and forth between forms.
The molecule exists only as the hybrid; in NO3− all three N–O bonds are identical, with bond order 1⅓.
Any molecule with polar bonds is polar.
Symmetric shapes cancel bond dipoles: CO2 and CCl4 have polar bonds but are nonpolar molecules.
Breaking a bond releases energy.
Breaking bonds always requires energy; energy is released when bonds form.
Practice questions
AP-style multiple choice. Every answer comes with the reasoning, including why the tempting wrong answer is wrong.
Which molecule has a trigonal pyramidal molecular geometry?
Free-response practice
The nitrate ion, NO3−, and the nitrite ion, NO2−, both contain N–O bonds. (a) Draw a valid Lewis structure for NO3− and state how many resonance structures it has. (b) Give the molecular geometry, O–N–O bond angle and hybridization of N in NO3−. (c) Is the N–O bond in NO3− longer, shorter or the same length as the N–O bond in NO2−? Justify your answer.
Flashcards
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