1 : Lewis Structures

• Lewis structures are diagrams that represent the bonding and valence electron distribution in a molecule or ion.
• Octet Rule: Atoms tend to gain, lose, or share electrons to achieve a stable configuration with eight valence electrons.

2 : Formal Charge (FC)

• Formula: FC = Valence Electrons - (Number of Bonds + Number of Lone Electrons)
• A more balanced distribution of formal charges is preferred in Lewis structures.

3 : VSEPR Theory (Valence Shell Electron Pair Repulsion)

• VSEPR predicts the molecular geometry based on the repulsion between electron pairs around the central atom.
• Common molecular geometries: Tetrahedral, Trigonal Planar, Linear, Bent, Trigonal Pyramidal.

4 : Bond Polarity

• Electronegativity (EN) is the measure of an atom’s ability to attract electrons in a chemical bond.
• Nonpolar Covalent Bond: EN difference < 0.5
• Polar Covalent Bond: EN difference between 0.5 and 2.0
• Ionic Bond: EN difference > 2.0

5 : Hybridization

• Hybridization theory explains the mixing of atomic orbitals to form hybrid orbitals for bonding.
• Common hybridization states: sp³, sp², sp.

6 : Molecular Orbital Theory

• Molecular orbitals are formed by the overlap of atomic orbitals. The theory predicts the electronic structure of molecules.
• Bond Order (BO) = (Number of Bonding Electrons - Number of Antibonding Electrons) / 2
• BO > 0: Stable molecule, BO = 0: Non-existent molecule, BO < 0: Unstable molecule

7 : Molecular Orbital Theory

• Dipole moment (μ) measures the polarity of a molecule.
• Formula: μ = charge (Q) × distance between charges (r)
• Direction: From the positive to the negative charge.

8 : Resonance

• Resonance occurs when multiple Lewis structures can be drawn for a molecule.
• Resonance hybrid: A combination of all resonance structures.
• Delocalized electrons: Electrons that are not localized to one bond.

9 : Resonance

• Occurs when a hydrogen atom is bonded to a highly electronegative atom (O, N, F) and is attracted to another electronegative atom’s lone pair.
• Important in water (H₂O), ammonia (NH₃), and biological molecules.

10 : Ionic Compounds

• Ionic compounds form between metals and non-metals through the transfer of electrons.
• Formula: Cation (metal) followed by Anion (non-metal).

11 : Lattice Energy

• Lattice energy is the energy required to separate ions in an ionic compound.
• It increases with smaller ion size and higher charge magnitude.

12 : Allotropy

• Allotropy refers to different structural forms of the same element in the same physical state.
• Examples: Carbon (diamond, graphite), oxygen (O₂, O₃).

13 : Metallic Bonding

• In metallic bonding, a “sea of electrons” holds metal ions together.
• Delocalized electrons contribute to properties like electrical conductivity and malleability.

14 : Solubility Rules

Memorize common solubility rules for ionic compounds to predict their solubility in water.

15 : Chemical Reactions

Understand how chemical bonds are broken and formed in chemical reactions, leading to the formation of new compounds.