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General Chemistry - Bonding and Molecular Theory

Periodic Table of Elements

  • Elements: H, Li, Be, Na, Mg, C, N, O, F, Ne, Ar, K, Ca, etc.

  • Atomic numbers and mass presented.

  • Color code for solid, liquid, gas, metal, metalloid, nonmetal.

Valence Bond Theory

  • Covalent bonds formed by overlapping atomic orbitals.

  • Conditions:

    • Two orbitals overlap.

    • Total electrons ≤ 2.

Hybridization

  • Mixing of atomic orbitals into hybrid orbitals for covalent bonds.

  • Types and geometric arrangements:

    • sp: Linear (BeF2)

    • sp²: Trigonal planar (BF3)

    • sp³: Tetrahedral (CH4)

    • sp³d: Trigonal bipyramidal (PCl5)

    • sp³d²: Octahedral (SF6)

Molecular Geometry

  • VSEPR model to determine geometry based on electron pairs.

  • Example: H2O with 4 pairs around O.

Types of Bonds

  • σ bonds: Direct overlap along bond axis.

  • π bonds: Side-to-side overlap.

Multiple Bonding

  • Each bond + lone pair counted per atomic orbital overlap.

  • Example: C=C in ethene has 1 σ bond and 1 π bond.

Molecular Orbital Theory

  • MOs formed by the combination of atomic orbitals.

  • Bonding and antibonding orbitals exist with stability comparison.

  • Example: H2 forms bonding (σ) and antibonding (σ*) orbitals.

Bond Order

  • Defined as: ext{Bond Order} = \frac{nb - na}{2}

    • $n_b$: electrons in bonding orbitals

    • $n_a$: electrons in antibonding orbitals.

Magnetic Properties

  • Paramagnetism: Unpaired electrons attract to magnets (e.g., O2).

  • Diamagnetism: Paired electrons repel from magnets (e.g., N2).

Summary of MO Theory

  • Pros: Detailed explanation of covalent bonding and magnetic properties.

  • Cons: Complex for large molecules and does not address molecular shape.