CHEMCA
EXAM MASTER FORMULA SHEET
Chemical Bonding & Molecular Structure
1. Ionic Bonding, Lattice Energy & Fajan's Rules
Energy released when 1 mole of solid ionic crystal is formed from isolated gaseous ions.
Charge dominates over size!
Hydration Energy > Lattice Energy
No ionic bond is 100% ionic. Covalent character arises due to polarization of the anion by the cation.
- High charge on cation (e.g., \(Al^{3+} > Mg^{2+} > Na^+\))
- Small size of cation (e.g., \(Li^+ > Na^+ > K^+\))
- Pseudo-inert gas configuration (\(ns^2np^6nd^{10}\)) polarizes more than inert gas config (\(ns^2np^6\)). E.g., \(CuCl\) is more covalent than \(NaCl\).
- High negative charge on anion
- Large size of anion (e.g., \(I^- > Br^- > Cl^- > F^-\))
2. Hybridization & VSEPR Theory (Shapes)
>
Lone Pair - Bond Pair
>
Bond Pair - Bond Pair
Note: Multiple bonds act as a single super-pair for geometry but exert stronger repulsion than single bonds.
| \(Z\) | Hybridization | Lone Pairs (lp) | Molecular Shape (Geometry) | Examples |
|---|---|---|---|---|
| 2 | \(sp\) | 0 | Linear (180°) | \(BeCl_2, CO_2, HCN\) |
| 3 | \(sp^2\) | 0 | Trigonal Planar (120°) | \(BF_3, SO_3, NO_3^-\) |
| 1 | V-shape / Bent (< 120°) | \(SO_2, O_3, SnCl_2\) | ||
| 4 | \(sp^3\) | 0 | Tetrahedral (109°28') | \(CH_4, NH_4^+, SO_4^{2-}\) |
| 1 | Trigonal Pyramidal (< 109.5°) | \(NH_3, PCl_3, H_3O^+\) | ||
| 2 | V-shape / Bent (<< 109.5°) | \(H_2O, H_2S, OF_2\) | ||
| 5 | \(sp^3d\) | 0 | Trigonal Bipyramidal (TBP) | \(PCl_5, PF_5\) |
| 1 | See-saw | \(SF_4\) | ||
| 2 | T-shape | \(ClF_3, BrF_3\) | ||
| 3 | Linear | \(XeF_2, I_3^-, ICl_2^-\) | ||
| 6 | \(sp^3d^2\) | 0 | Octahedral / Sq. Bipyramidal | \(SF_6\) |
| 1 | Square Pyramidal | \(IF_5, BrF_5\) | ||
| 2 | Square Planar | \(XeF_4, [Ni(CN)_4]^{2-}\) |
3. Dipole Moment (\(\mu\)) & Formal Charge
Units: Debye (D). \(1 \text{ D} = 3.3356 \times 10^{-30} \text{ C}\cdot\text{m}\) = \(10^{-18} \text{ esu}\cdot\text{cm}\)
Apparent charge on an atom in a molecule/ion.
- V: Total valence electrons in free atom
- L: Total non-bonding (lone pair) electrons
- S: Total shared (bonding) electrons
The most stable Lewis structure has formal charges closest to zero.
4. Molecular Orbital Theory (MOT)
\(N_b\) = bonding \(e^-\), \(N_a\) = antibonding \(e^-\)
- B.O. \(\propto\) Bond Dissociation Energy (Strength)
- B.O. \(\propto \frac{1}{\text{Bond Length}}\)
- If B.O. = 0 or negative, the molecule does not exist (e.g., \(He_2, Ne_2\)).
- Resonance Shortcut: \( \text{B.O.} = 1 + \frac{\text{Number of } \pi \text{ bonds}}{\text{Number of } \sigma \text{ bonds}} \) (e.g., \(CO_3^{2-}\) has B.O. = 1.33)
\(\sigma 1s < \sigma^* 1s < \sigma 2s < \sigma^* 2s < \mathbf{(\pi 2p_x = \pi 2p_y) < \sigma 2p_z} < (\pi^* 2p_x = \pi^* 2p_y) < \sigma^* 2p_z\)
\(\sigma 1s < \sigma^* 1s < \sigma 2s < \sigma^* 2s < \mathbf{\sigma 2p_z < (\pi 2p_x = \pi 2p_y)} < (\pi^* 2p_x = \pi^* 2p_y) < \sigma^* 2p_z\)
5. Hydrogen Bonding & Weak Van der Waals Forces
Intermolecular H-bond
Formed between two different molecules (same or different compounds).
Effects: Increases Boiling Point, Viscosity, Surface Tension, and Solubility in water.
Examples: \(H_2O, NH_3, HF\), Alcohols, Carboxylic acids (dimer formation).
Intramolecular H-bond
Formed within the same molecule (results in ring formation / Chelation).
Effects: Decreases Boiling Point (more volatile), prevents association, lowers solubility.
Examples: o-Nitrophenol, Salicylaldehyde.
- Ion-Dipole Interaction: \(\propto 1/r^2\) (Strongest weak force, causes hydration).
- Dipole-Dipole (Keesom): \(\propto 1/r^3\) (Stationary polar molecules) or \(\propto 1/r^6\) (Rotating polar molecules).
- Ion-Induced Dipole: \(\propto 1/r^4\).
- Dipole-Induced Dipole (Debye): \(\propto 1/r^6\) (Polar + Non-polar).
- London Dispersion Forces: \(\propto 1/r^6\) (Non-polar + Non-polar). Depends on polarizability and molecular mass/surface area.
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