Bonding In Chemistry 11

Bonding in Chemistry 11

• Types of Bonds:

  • Ionic Bonds: Formed between a metal and a non-metal through the transfer of electrons.

  • Covalent Bonds: Formed between non-metals by sharing electrons.

  • Metallic Bonds: Found in metals where electrons are delocalized and shared among atoms.

• Lewis Structures:

  • Diagrams that show the bonding between atoms and the lone pairs of electrons.

  • Follow the octet rule for main group elements.

• Polarity:

  • Polar Covalent Bonds: Unequal sharing of electrons leading to partial charges.

  • Nonpolar Covalent Bonds: Equal sharing of electrons.

• Intermolecular Forces:

  • Van der Waals Forces: Weak forces between molecules.

  • Hydrogen Bonding: Strong dipole-dipole attraction between molecules with H-F, H-O, or H-N bonds.

• VSEPR Theory:

  • Predicts the shape of molecules based on the repulsion between electron pairs around the central atom.

  • Shapes include linear, trigonal planar, tetrahedral, trigonal bipyramidal, and octahedral.

• Hybridization:

  • Mixing of atomic orbitals to form new hybrid orbitals.

  • Examples include sp, sp2, sp3 hybridization.

• Resonance:

  • Describes the delocalization of electrons in molecules with multiple Lewis structures.

  • The actual structure is a combination of resonance structures.

• Molecular Orbital Theory:

  • Describes the behavior of electrons in molecules using molecular orbitals.

  • Molecular orbitals are formed by the combination of atomic orbitals.

• Bond Energy:

  • The energy required to break a bond.

  • Stronger bonds have higher bond energies.

• Electronegativity:

  • The ability of an atom to attract electrons in a chemical bond.

  • Increases across a period and decreases down a group in the periodic table