A Lewis structure is a 2D representation of a molecule showing how valence electrons are distributed as shared pairs (bonds) and lone pairs. The goal is usually to satisfy the octet rule (8 electrons) for each atom (or duet rule for H).
Formal Charge (FC): A "bookkeeping" charge to determine the most plausible Lewis structure.
Formula: FC = (Valence e-) - (Non-bonding e-) - (1) / (2)(Bonding e-)
The best structure has FCs closest to zero and any negative FC on the most electronegative atom.
VBT describes a covalent bond as the overlap of half-filled atomic orbitals. The electrons in the overlapping orbitals must have opposite spins.
Multiple Bonding: A double bond is (1 σ + 1 π). A triple bond is (1 σ + 2 π).
Hybridization (proposed by Pauling) is the concept of mixing atomic orbitals of slightly different energies to form a new set of degenerate (equal energy) hybrid orbitals. These new orbitals are better suited for bonding and explain observed molecular geometries.
Hybridization is an "energy-neutral" process in theory. Energy is "spent" to promote electrons (e.g., 2s → 2p in Carbon), but this energy is more than "repaid" by the formation of more, stronger, and more stable bonds (e.g., 4 C-H bonds in CH4 vs. the 2 expected for C 2p2).
Resonance is used when a single Lewis structure cannot adequately describe the bonding in a molecule. The actual structure is an average or "hybrid" of two or more resonance structures (or canonical forms), which differ only in the placement of π-electrons and lone pairs.
Example: Ozone (O3). We can draw two valid Lewis structures. The real O3 molecule is a hybrid of these, with both O-O bonds being identical (length of 1.5) and a charge of -0.5 on each outer oxygen.
Definition: The difference in energy between the actual resonance hybrid and the most stable of its contributing resonance structures.
Resonance leads to delocalization of electrons, which stabilizes the molecule. The larger the resonance energy, the more stable the molecule.
MOT is a more advanced model where all atomic orbitals (AOs) from all atoms combine to form an equal number of molecular orbitals (MOs) that are delocalized over the *entire* molecule.
Formula: BO = (1) / (2) × (No. of Bonding e- - No. of Antibonding e-)
For homonuclear diatomics B2, C2, and N2 (and their ions), the 2s and 2pz orbitals are close enough in energy to interact (mix). This mixing raises the energy of the σ2pz MO above that of the π2p_{x,y} MOs.
For O2, F2, and Ne2, the 2s-2p energy gap is too large for mixing, so the "normal" order is followed.
| For B2, C2, N2 (s-p mixing) | For O2, F2 (no s-p mixing) |
|---|---|
| σ1s < σ*1s < σ2s < σ*2s < π2p_{x,y} < σ2pz < π*2p_{x,y} < σ*2pz | σ1s < σ*1s < σ2s < σ*2s < σ2pz < π2p_{x,y} < π*2p_{x,y} < σ*2pz |
VSEPR is a model used to predict the 3D geometry of molecules based on the idea that electron pairs in the valence shell of a central atom repel each other and will arrange themselves to be as far apart as possible, minimizing repulsion.
LP-LP > LP-BP > BP-BPReason: A lone pair is only held by one nucleus, so it is "fatter" and "wider" than a bonding pair, which is held by two nuclei.
A = Central Atom, X = Bonded Atom (BP), E = Lone Pair (LP)
| Total Pairs (X+E) | Type | BP (X) | LP (E) | Electron Geometry | Molecular Shape | Angle(s) | Example |
|---|---|---|---|---|---|---|---|
| 2 | AX2 | 2 | 0 | Linear | Linear | 180° | BeF2, CO2 |
| 3 | AX3 | 3 | 0 | Trigonal Planar | Trigonal Planar | 120° | BF3 |
| 3 | AX2E | 2 | 1 | Trigonal Planar | Bent (V-shape) | < 120° | SO2, O3 |
| 4 | AX4 | 4 | 0 | Tetrahedral | Tetrahedral | 109.5° | CH4, NH4+ |
| 4 | AX3E | 3 | 1 | Tetrahedral | Trigonal Pyramidal | < 109.5° (e.g., 107°) | NH3, PCl3 |
| 4 | AX2E2 | 2 | 2 | Tetrahedral | Bent (V-shape) | < 109.5° (e.g., 104.5°) | H2O, SCl2 |
| 5 | AX5 | 5 | 0 | Trigonal Bipyramidal | Trigonal Bipyramidal | 90°, 120° | PCl5 |
| 5 | AX4E | 4 | 1 | Trigonal Bipyramidal | See-Saw | < 90°, < 120° | SF4 |
| 5 | AX3E2 | 3 | 2 | Trigonal Bipyramidal | T-shape | < 90° | ClF3 |
| 5 | AX2E3 | 2 | 3 | Trigonal Bipyramidal | Linear | 180° | XeF2 |
| 6 | AX6 | 6 | 0 | Octahedral | Octahedral | 90° | SF6 |
| 6 | AX5E | 5 | 1 | Octahedral | Square Pyramidal | < 90° | BrF5 |
| 6 | AX4E2 | 4 | 2 | Octahedral | Square Planar | 90° | XeF4 |
When two different atoms form a covalent bond (e.g., H-Cl), the shared electrons are not shared equally. The more electronegative atom (Cl) pulls the electrons closer, creating a polar covalent bond.
Formula: μ = q × d
Where: q = magnitude of charge, d = distance of separation. Unit = Debye (D).
Symmetrical molecules (like CO2, CCl4, BF3) can have polar bonds, but their bond moments cancel out, resulting in a zero dipole moment (μ = 0). They are non-polar.
Asymmetrical molecules (like H2O, NH3) have bond moments that do not cancel, resulting in a net dipole moment. They are polar.
No bond is 100% ionic. An "ionic bond" (e.g., NaCl) always has some covalent character because the cation (Na+) pulls on, or polarizes, the electron cloud of the anion (Cl-).
Fajan's rules predict the degree of covalent character in an ionic bond. Covalent character is favored by:
Increased covalent character leads to: