Overview
This lecture reviews core topics from AP Chemistry Unit 2: Compound Structure and Properties, focusing on chemical bonding, molecular structure, polarity, Lewis diagrams, lattice energy, and related trends for the summative assessment.
Types of Chemical Bonds and Polarity
- Electronegativity increases left to right and decreases top to bottom in the periodic table.
- Polar covalent bonds form when atoms with unequal electronegativity share electrons; the more electronegative atom gets a partial negative charge.
- The bond with the largest electronegativity difference (e.g., Ge–O) has the greatest dipole moment.
- Bonds between identical atoms (e.g., C–C) have no dipole moment (zero polarity).
Potential Energy Curves and Bond Strength
- Bond length corresponds to the minimum potential energy between two atoms.
- Bond energy is the energy required to break the bond, always reported as a positive value.
- Shorter bonds (smaller atoms or higher bond order) have higher bond energies.
- If two atoms are closer (smaller atomic radius), more energy is needed to break their bond.
Ionic vs. Covalent Compounds
- Ionic compounds form between metals and nonmetals, have high melting points, are solid at room temp, and conduct electricity when molten or dissolved.
- Covalent (molecular) compounds form between nonmetals, often have lower melting points, and do not conduct electricity.
Lewis Structures, Bond Order, and Resonance
- Lewis diagrams require correct valence electron counts, bonds, and lone pairs to satisfy octets when possible.
- Multiple bonds (double/triple) may be needed when there are not enough electrons for octets.
- Higher bond order correlates with higher bond energy (triple > double > single bond).
- Resonance structures share electron pairs among multiple positions; bond order may be fractional (e.g., 1.5).
Lattice Energy and Coulomb’s Law
- Lattice energy increases with higher ionic charge and decreases with increased ionic radius.
- Stronger attractions (smaller ions, higher charge) mean higher lattice energies and more difficult separation.
Molecular Geometry and Polarity
- Electron domains determine geometry: 4 domains = tetrahedral (109.5°), 3 = trigonal planar (120°), 2 = linear (180°).
- Lone pairs increase repulsion, reducing bond angles (e.g., NH3 bond angle < CH4).
- Polar molecules have asymmetric electron distributions or different terminal atoms; nonpolar molecules are symmetric and have identical outer atoms.
Alloys and Steel Structure
- Interstitial alloys form when small atoms fit between larger atoms; steel (Fe and C) is interstitial.
- Higher carbon content in steel increases rigidity by obstructing atom movement.
Formal Charge and Preferred Lewis Structures
- Assign formal charge: valence electrons – (dots + bonds).
- Most preferred structure has the smallest charges and places negative charge on the most electronegative atom.
Counting Sigma and Pi Bonds
- Single bond: 1 sigma.
- Double bond: 1 sigma, 1 pi.
- Triple bond: 1 sigma, 2 pi.
Key Terms & Definitions
- Electronegativity — measure of an atom’s ability to attract electrons.
- Dipole Moment — separation of charges in a bond or molecule.
- Bond Order — number of shared electron pairs between atoms.
- Lattice Energy — energy required to separate one mole of an ionic solid into gaseous ions.
- Resonance — delocalization of electrons across multiple structures.
- Sigma Bond — the first covalent bond between two atoms.
- Pi Bond — additional bonds in double/triple bond situations.
- Interstitial Alloy — alloy with smaller atoms filling spaces between larger atoms.
- Formal Charge — calculated charge on atoms in a Lewis structure.
Action Items / Next Steps
- Review periodic trends and how they affect bonding.
- Practice drawing Lewis structures and identifying molecular shapes.
- Study resonance and formal charge to determine preferred Lewis structures.
- Complete and check assigned packet problems for Unit 2.