Forces of Attraction

Forces of Attraction

Comparison Between Ionic Bonding, Covalent Bonding & Metallic Bonding

Overview

  • Types of Bonding:

    • Ionic Bonding (Giant crystal lattice structure)

    • Covalent Bonding (Simple discrete molecules)

    • Metallic Bonding

Properties Comparison

  • Bonds With:

    • Ionic Bonding: Between metallic elements and non-metallic elements only (cations and anions only)

    • Covalent Bonding: Between non-metallic elements only

    • Metallic Bonding: Within pure metallic elements only

  • Forces of Attraction:

    • Ionic Bonding: Strong electrostatic forces of attraction between oppositely charged ions

    • Covalent Bonding: Strong covalent bonds within the molecule, weak intermolecular forces of attraction between the molecules

    • Metallic Bonding: Strong electrostatic forces between closely packed metallic cations and 'sea of delocalized electrons'

  • Physical State at r.t.p.:

    • Ionic Bonding: Crystalline solid

    • Covalent Bonding: Usually volatile liquid or gas

    • Metallic Bonding: Solid

  • Solubility in Water:

    • Ionic Bonding: Soluble

    • Covalent Bonding: Insoluble

    • Metallic Bonding: Insoluble

  • Solubility in Organic Solvents:

    • Ionic Bonding: Insoluble

    • Covalent Bonding: Low except for giant molecular structures (diamond, graphite, silicon (IV) oxide)

    • Metallic Bonding: Insoluble

  • Boiling and Melting Points:

    • Ionic Bonding: Very high; due to strong electrostatic forces between oppositely charged ions throughout the entire structure.

    • Covalent Bonding: Low; weak intermolecular forces of attraction between molecules and absence of mobile ions and delocalized electrons.

    • Metallic Bonding: High; due to strong electrostatic forces between cations and the delocalized electrons.

  • Electrical Conductivity:

    • Ionic Bonding: Conducts electricity in aqueous and molten state; oppositely charged ions are mobile.

    • Covalent Bonding: Does not conduct electricity in any states due to absence of mobile ions and delocalized electrons.

    • Metallic Bonding: Conducts electricity in solid and molten states; delocalized electrons are mobile.

  • Hardness, Ductility & Malleability:

    • Ionic Bonding: Hard but brittle

    • Covalent Bonding: Gas: soft; Liquid: soft; Solid: brittle

    • Metallic Bonding: Variable hardness; ductile and malleable

Electrostatic Forces

  • Strength of Forces:

    • Electromagnetic static forces are proportional to charge and inversely proportional to the radius of ions

Van-der-Waals Forces

  • Present in all covalent compounds, regarded as weak forces of attraction, including instantaneous dipoles, London dispersion forces.

  • Larger molecules have larger electron clouds leading to more VDW forces.

  • Molecules with larger surface areas have more VDW forces due to bigger exposed electron clouds.

Effects on Boiling and Melting Points

  • Overcoming intermolecular forces is necessary to boil a liquid, whereby stronger intermolecular forces require more energy; hence liquids with stronger VDW forces have higher boiling points.

  • Similarly, solids with more VDW forces exhibit higher melting points.

Dipole-Dipole Forces of Attraction

  • Definition: Polar molecules have permanent dipoles, which are stronger than VDW forces.

  • Dipole results from separation of charge due to differences in electronegativity (E/N).

Hydrogen Bonding

  1. Occurs only when hydrogen is covalently bonded to fluorine, nitrogen, or oxygen, which have available lone pairs.

  2. Exhibit strong polarizations resulting in significant positive charge on hydrogen atoms.

  3. Examples of molecules with hydrogen bonding include water, ammonia, and hydrogen fluoride (HF).

Characteristics of Hydrogen Bonding

  • Hydrogen bonding contributes to higher boiling points due to the significant energy required to overcome the strong forces.

  • Water exhibits anomalous properties; despite being a small molecule, it has a higher boiling point due to hydrogen bonding compared to related compounds.

Properties of Key Molecules

  • Water (H₂O):

    • Capable of forming four hydrogen bonds per molecule but effectively forms fewer bonds in the liquid state compared to ice, contributing to its density characteristics.

    • Most dense at 4 degrees Celsius.

  • Hydrogen Fluoride (HF):

    • Forms strong hydrogen bonds, leading to a high boiling point, but potentials for hydrogen bonds are limited by insufficient hydrogen atoms relative to lone pairs.

Group Hydrides

  • Trends observed in boiling points among Group 1 and Group 2 hydrides relate to VDW forces and dipole-dipole interactions, with boiling point patterns influenced by molecular size and structure.

Conclusion

  • The interplay of various types of bonding and intermolecular forces dictates molecular behavior, physical properties, and the differences in boiling and melting points across diverse compounds.