Atomic Structure, Periodicity, and Chemical Bonding
This domain covers the quantum-mechanical atom, electron configuration, periodic trends, and the models that explain bonding and intermolecular forces. Items often require you to move between a formula, a particle diagram, and a measurable property.
Atomic structure and periodicity
- Orbitals fill by Aufbau order; Hund's rule requires unpaired electrons before pairing; Cr and Cu are exceptions because half-filled or filled d subshells are favored.
- Effective nuclear charge increases across a period, so atomic radius decreases and first ionization energy generally increases; radius increases down a group as n grows.
- Electronegativity follows a similar trend, with noble gases excluded.
Bonding and shape
- Ionic bonds transfer electrons between metals and nonmetals; covalent bonds share electrons, and polarity is set by electronegativity difference.
- Draw Lewis structures, assign formal charge, then apply VSEPR: lone pairs, single bonds, and double bonds each count as one electron domain; lone pairs compress bond angles.
- Molecular polarity depends on geometry as well as bond dipoles. CO2 is nonpolar; SO2 is polar.
- For similar molar mass, intermolecular forces rank roughly hydrogen bonding > dipole-dipole > London dispersion.
Worked example
For SO2: total valence electrons = 6 + 2(6) = 18. Place S central, connect two O atoms, and add lone pairs. A double bond to one O and a single bond to the other gives formal charges that sum to zero. The S center has three electron domains (two bonds and one lone pair), so VSEPR predicts a bent shape with an angle below 120 degrees. The two S-O bond dipoles do not cancel, so SO2 is polar and has dipole-dipole forces.
Traps that cost marks
- Counting a double bond as two VSEPR domains.
- Assuming a molecule with polar bonds is always polar.
- Confusing ionization energy with electron affinity.
- Ignoring d-orbital exceptions when writing electron configurations.