inorganic

General and Inorganic Chemistry (PHCM101) Lecture 8

  • Presented by Dr. Nesrine El Gohary.

Competencies

  • 1-1-1 Identify different types of intermolecular forces.

  • 1-1-2 Demonstrate understanding of the valence bond theory.

  • 1-1-3 Demonstrate understanding of hybridization and hybrid orbitals.

  • 2-2-1 Apply different types of orbital hybridization.

  • 2-2-2 Determine the effect of multiple bonds.

Introduction to Intermolecular Forces

  • Intermolecular forces are attraction forces between molecules.

  • Responsible for various physical properties of substances.

  • Weaker than intramolecular forces, which are chemical bonds within a molecule.

Types of Intermolecular Forces

  1. Van der Waals Forces

  2. Ion-Dipole Forces

  3. Hydrogen Bonding

  4. Ionic Bonding

Van der Waals Forces

Types of Van der Waals Forces

  • Dipole-Dipole Forces: Present in polar molecules with permanent dipoles.

    • Molecules attract when +ve and -ve ends align.

    • Weaker in gas phase due to increasing distance between dipoles.

  • London Dispersion Forces (Induced Fluctuating Dipole):

    • Occur in non-polar molecules creating temporary dipoles.

    • Instantaneous dipoles induce neighboring molecule dipoles, leading to attraction.

  • Ion-Induced Dipole Forces: Attraction between an ion and induced dipoles.

  • Dipole-Induced Dipole Forces: Attraction between dipoles and temporary dipoles.

Hydrogen Bonding

  • Strong dipole-dipole attraction involving Hydrogen and highly electronegative atoms (e.g., N, O, F).

  • Strength comes from:

    1. High polarity due to electronegativity difference.

    2. Proximity due to small size of H atom.

Ionic Bonding

  • Found in compounds of metallic and nonmetallic elements (e.g., NaCl).

  • Electrostatic attraction between oppositely charged ions.

  • Increased attraction with higher charges and smaller ionic radii.

Conductivity of Ionic Compounds

  • Ionic compounds do not conduct electricity in solid form (fixed in lattice).

  • When dissolved in water, ions are free to move, making solutions good conductors.

Comparison of Strengths of Intermolecular Forces

  • Ion-Dipole: +40 to -600 kJ/mol.

  • Dipole-Dipole: +5 to +25 kJ/mol.

  • Hydrogen Bonding: 10-40 kJ/mol.

  • Dipole-Induced and Ion-Induced: 2-10 kJ/mol.

  • London Dispersion: 0.05-40 kJ/mol.

Theoretical Frameworks for Bonding

  1. Valence Bond Theory

  2. Molecular Orbital Theory

  • Valence Bond Theory Conditions:

    1. Overlapping orbitals exist between two atoms.

    2. No more than two electrons in overlapping region.

    3. Bond formation releases energy (exothermic reaction).

Hybridization and Bond Formation

  • Carbon Example:

    • Ground state: C (1s2, 2s2, 2p2) expects to bind only 2 H atoms.

    • Actual: C binds with 4 H (CH4) through hybridization (sp3).

    • Four new orbitals formed by combining 2s and 2p.

    • Tetrahedral shape with 109.5° angles.

Different Types of Hybridization

  • LD - sp: Linear arrangement (2 effective pairs).

  • LD - sp2: Triangular planar (3 effective pairs).

  • LD - sp3: Tetrahedral (4 effective pairs).

  • LD - dsp3: Triangular bipyramidal (5 effective pairs).

  • LD - d2sp3: Octahedral (6 effective pairs).

Practical Applications of Hybridization

  • Predicting hybridization based on effective electron pairs:

    • Lone pairs and bonds count as one effective pair.

    • Examples include BF4- (sp3) and XeF2 (dsp3).

Important Terminology

  • Bond Order: Number of shared electron pairs.

    • Single bond: bond order 1, double bond: bond order 2, triple bond: bond order 3.

  • Bond Energy: Energy needed to break a bond, higher for stronger bonds.

Styles of Hybridization Involving d-orbitals

  • dsp3 Hybridization: SF4 and PCl5 with five effective pairs (trigonal bipyramidal).

  • d2sp3 Hybridization: SF6 and XeF4 (octahedral with lone pairs).