MS

Intermolecular Forces and Bond Polarities Questions

Bond Polarity and Electronegativity

  • Bond Polarity: Refers to the distribution of electron density in a chemical bond. It affects the behaviour of molecules in various chemical contexts.

  • Electronegativity (EN):

    • Definition: The relative strength of attraction of an atom to a shared pair of electrons in a covalent bond.

    • Measured using the Pauling Electronegativity Scale.

    • Description:

    • Has no units, established as a relative value with a ranking system.

    • For any pair of atoms participating in a chemical bond, electronegativity (EN) can be calculated as the average of the two atoms' electronegativities.

  • Van Arkel-Ketelaar Triangle: A triangular representation used to indicate the type of bond an element forms based on its electronegativity.

    • Shows categories of bonding, ranging from ionic to covalent bonds.

  • Types of Bonds:

    • Pure Covalent Bond:

    • Definition: Equal sharing of electrons between two atoms.

    • Example: Hydrogen (H-H) bond.

    • Characteristic: 0 difference in electronegativity, thus considered zero percent ionic and 100 percent covalent.

    • Electrostatic: If the ionic character is below 5%, still categorized as a pure covalent bond.

    • Polar Covalent Bond:

    • Definition: Electrons are shared unequally between two atoms, where one atom is more electronegative than the other.

    • Characteristic: Represented by a “dipole moment” indicating electron density.

    • Explanation: The dipole moment points towards the more electronegative atom (higher EN), resulting in slight negative (δ−) and slight positive (δ+) charges.

    • Implication: Electrons are retained more by the more electronegative atom, generating an uneven distribution of charge.

Molecular Polarity

  • Definition: Molecular polarity indicates whether a molecule has an overall dipole moment, which leads it to be classified as either polar or non-polar.

  • Molecular Categories:

    • Non-Polar Molecules:

    • Example: Hydrogen (H-H).

    • Characteristic: Even charge distribution across the molecule, resulting in no net dipole.

    • Polar Molecules:

    • Characteristic: Uneven charge distribution, resulting in an overall dipole moment.

    • Example: H-Cl (polar covalent bond) leading to an overall dipole.

    • Important Consideration: The net dipoles can cancel out; thus, the overall polarity depends on the angle and number of bonds.

Intermolecular Forces (Van der Waals Forces)

  • Overview: Intermolecular forces are attractions between neutral molecules within a sample, affecting physical states (liquids/solids).

    • Distinction: Not the same as intramolecular forces, which are covalent bonds within a single molecule.

    • Importance: Influences melting and boiling points; overcome when substances transition phases.

Types of Intermolecular Forces:

  1. London Dispersion Forces:

    • Description: The weakest intermolecular forces found in non-polar molecules.

    • Cause: Result from momentary electron distributions, causing temporary dipoles.

    • Characteristics:

      • Weaker than dipole-dipole attractions;

      • Larger molecules exhibit stronger London forces due to increased electron mobility.

  2. Dipole-Dipole Attractions:

    • Definition: Attractive forces between polar molecules, characterized by the partial positive (δ+) end of one molecule and the partial negative (δ−) end of another.

    • Characteristics: Configuration leads to orderly crystal formation in solid states; only observed among polar molecules.

  3. Hydrogen Bonds:

    • Definition: A stronger type of dipole-dipole attraction occurring between polar molecules featuring Hydroxyl (H) atoms bonded to highly electronegative atoms (N, O, F).

    • Nature: Considered an extreme state of dipole-dipole attraction due to significant electronegativity differences.

Detailed Exploration of Intermolecular Forces:

Hydrogen Bonding

  • Characteristic: Extremely strong dipole-dipole interaction due to high difference in electronegativities.

  • Formation Mechanism:

    1. Hydrogen (H) attached to highly electronegative atoms (e.g., N, O, or F) becomes relatively positively charged (δ+).

    2. The positive hydrogen attempts to interact with lone pairs of another electronegative atom in a different molecule.

  • Comparative Strength: Hydrogen bonds are about 1/10 the strength of covalent bonds.

Specific Examples of Hydrogen Bonding:

  1. Hydrogen Fluoride (HF):

    • Hydrogen bonding occurs when H interacts with the lone pair of fluorine (F) from a neighboring HF molecule, illustrated through dashed lines representing hydrogen bonds.

  2. Ammonia (NH3):

    • Similar to HF, H from one NH3 can bond to the lone pair of N in another molecule.

  3. Water (H2O):

    • Each water molecule can form up to two hydrogen bonds due to the presence of two hydrogen atoms extensively bonded to one electronegative oxygen atom (O).

    • Effect: This leads to a much higher boiling point compared to related compounds without hydrogen bonding.

    • Unique Feature: Water’s structure becomes less dense in solid form due to its hydrogen bonding arrangement, causing ice to float.

Summary Questions:

  1. Define and explore the differences between intramolecular and intermolecular forces.

  2. Discuss van der Waals Forces, dipole-dipole forces, and London Forces in detail.

  3. Describe dipoles and their role in molecular interactions, including characteristic diagrams.

  4. Analyze what constitutes an instantaneous dipole and the subsequent establishment of induced dipoles explaining London Forces with diagrams.

  5. Solve problems based on determining intermolecular forces among given compounds and provide thorough methodical explanations for each case.

  6. Explain the phenomena leading to the boiling point discrepancies among polar molecules under varied conditions, citing water and hydrogen sulfide (H2S) as examples.