Electron Orbitals and Molecular Shapes

An orbital is a region with a high probability of finding an electron, with each orbital holding up to 2 electrons. Electrons fill orbitals in increasing energy order (1s, 2s, 2p, etc.), and they prefer to occupy orbitals singly before pairing up. Remember these key exceptions: 4s fills before 3d (but empties first when writing electron configurations), and elements like chromium (3d⁵4s¹) and copper (3d¹⁰4s¹) have unusual configurations so each electron can have its own orbital.

Molecular shapes depend on the arrangement of electron pairs around the central atom. Linear molecules $like Cl-B-Cl$ have a 180° bond angle with two bonding pairs and no lone pairs. Tetrahedral molecules (like CH₄) have four bonding pairs arranged at 109.5° angles. Other important shapes include non-linear (like H₂O, 104.5°), trigonal planar (like BF₃, 120°), pyramidal (like NH₃, 107.5°), and octahedral (six bonding pairs at 90° angles).

A molecule's polarity depends on both bond polarity and molecular shape. Even with polar bonds, a molecule can be non-polar if its dipole moments cancel out. For example, CO₂ has polar C-O bonds, but the linear arrangement causes the dipoles to cancel, making the molecule non-polar. NH₃, however, has dipoles that don't cancel, making it polar.

Remember this! Lone pairs repel more strongly than bonding pairs, which is why water's H-O-H angle (104.5°) is smaller than the perfect tetrahedral angle (109.5°).

Electronegativity and Bond Polarity

Electronegativity refers to an atom's ability to attract bonding electrons in a covalent bond. This property follows predictable trends across the periodic table. Across a period (left to right), electronegativity increases as nuclear charge increases while the number of electron shells stays the same, creating stronger attraction for bonding electrons.

Down a group, electronegativity decreases despite increasing proton number. This happens because the additional electron shells create more shielding, which negates the increased nuclear charge and results in less attraction between the nucleus and bonding electrons.

Fluorine is the most electronegative element, with electronegativity generally decreasing as you move left and down the periodic table. When two atoms of the same element share electrons (like in Cl₂), they have identical electronegativity, resulting in equal electron sharing and a non-polar bond.

When atoms with different electronegativities bond $like H-Cl$, the more electronegative atom (Cl) pulls the shared electrons closer, creating a partial negative charge (δ-) on that atom and a partial positive charge (δ+) on the less electronegative atom (H). This separation of charge creates a dipole moment, resulting in a polar covalent bond.

Chemistry hack: To quickly determine if a molecule is polar, check both the electronegativity differences (are the bonds polar?) and the molecular shape (do the dipoles cancel?). Both factors matter!