Chemistry - Structure of Matter and Chemical Bonding
This section tests the bridge from electron arrangement to observable properties: why sodium chloride conducts when molten but not as a solid, why water is bent, and why boiling points differ. Expect electron configurations, periodic trends, bond types, Lewis structures, VSEPR shapes, polarity, and intermolecular forces.
Core rules
- Atomic number Z = protons; mass number A = protons + neutrons. Isotopes differ in neutrons only.
- Ionic bonding: metal loses electrons, nonmetal gains; lattice energy and high melting point follow from electrostatic attraction.
- Covalent bonding: shared electron pairs. Bond polarity depends on electronegativity difference; molecular polarity also depends on geometry.
- VSEPR: electron pairs repel. Two bonding pairs = linear; three = trigonal planar; four = tetrahedral; lone pairs compress angles.
- Intermolecular forces: hydrogen bonding > dipole–dipole > London dispersion, for similar molar mass.
Worked example
Explain the shape and polarity of H2O. Oxygen has six valence electrons; two form single bonds with hydrogen, leaving two lone pairs. Four electron domains give a tetrahedral electron geometry, but the two lone pairs make the molecular shape bent, with an H–O–H angle near 104.5°. Oxygen is more electronegative than hydrogen, so the O–H bonds are polar, and the bent geometry means the bond dipoles do not cancel. Water is therefore polar and forms hydrogen bonds.
Traps
- Confusing mass number with atomic number when counting neutrons.
- Assuming any molecule with polar bonds is polar overall; CO2 is linear and nonpolar.
- Ignoring lone pairs when predicting shape.
- Treating all covalent bonds as nonpolar.
- Using melting point alone to identify bond type without considering the lattice.