Bonding

1. Ionic Bonding

  • Ionic Bonding - the transfer of electrons
  • The transfer of electrons - the electrostatic force of attraction between the oppositely charged particles

Characteristics

  1. Well-defined crystals (solids at room temp.)

  2. High melting point

    1. Above 400ºC

  3. Soluble in water - forms ions in solutions

  4. Conduct electricity in molten states

  5. Between metals and nonmetals

  6. Electron difference in Ionic Bonds is greater than or equal to 1.67

2. Covalent Bonding

Covalent Bonding - sharing of electrons

Characteristics

  1. Exist as three states of matter

    1. Solids - table sugar
    2. Liquids - ethers
    3. Gases - methane

  2. Low Melting Point

  3. Low solubility in water

  4. No or poor conduction

  5. Between nonmetals

  6. Brittle

  7. Electron difference is less than 1.67

    molecules in constant motion - bonds can bend, rotate, or wag

Coordinate Covalent Bonds

  • Coordinate Covalent Bonds - both electrons in the shared pair come from the same atom

Polyatomic Ions

  • Polyatomic Ions - 2 or more atoms covalently bonded together but resulting in a net charge
  • The ion as a whole bonds ionically with something else
    • Ex: SO4

Metallic Bonding

  • Metallic Bonding - metals with free electrons that are delocalized and able to move throughout the metals

Dot Diagrams

Ionic Dot Diagrams

  • Must use enough of each atom to get a total charge of zero'
  • Metals lose electrons
    • Cations
  • Nonmetals gain electrons
    • Anion

Covalent Dot Diagrams

  • IA
    • Always Ionic
    • No shape
  • IIA
    • Linear molecule
    • Bond on right and bond on left
    • 180º bond angle
  • IIIA
    • Trigonal planar molecule
    • Flat, no unshared pare of electrons
  • IVA
    • Tetrahedron
    • 4 polar bonds
  • VA
    • Trigonal Pyramid
    • One pair of unshared electrons
    • 3 polar bonds
    • Larger electron cloud
  • VIA
    • Bent or V-Shaped
    • Not flat
    • Not linear
    • 104.5º bond angle
  • VIIA
    • Can only be linear

Bond Polarity

  • Bond polarity - electrons aren’t shared equally between 2 atoms in a bond
  • Non-polar bond - Electronegativity less than or equal to 0.4
  • Polar bond, dipole - 0.4<Electronegativity<1.67
  • Dipole - electrons moving closer to one atom
  • Electrons are closer to one atom because it has a higher electronegativity
  • Exception: HF
    • very weak bond (both atoms are tiny)
    • the bond is too small
    • Not Ionic

Molecular Polarity

  • The poles aren’t balanced
  • Electrons are not shard equally throughout the molecule
  • Trigonal Pyramid - same atom on all sides: non-polar molecule
    • they all cancel out
  • Tetrahedron - 4 polar bonds, non polar molecule
  • Trigonal Pyramid & Bent or V-Shaped
    • polar molecule
    • unshared pair of electrons
  • VIIA - 1 polar bond, polar molecule

Coordinate Covalent Bonds

Coordinate Covalent Bonds - atoms in bond come from the same atom

Resonance

  • structures where moving the double-bond does nothing

Intermolecular Forces

Intermolecular forces - forces of attraction between molecules

Dispersion Forces < Dipole Forces

  • van der Waals Forces - weak forces of attraction between the nucleus of one atom and another

  1. Dipole-Dipole

  2. Dipole-Induced-Dipole

  3. Dispersion Forces - only attractive free between non-polar molecules

    1. In all covalent bonds

      1. Caused by the movement of electrons
  • As the atomic number increases, so does the bond strength
  • Polar substances have both dipole and dispersion forces

Hydrogen Bonding

  • strongest intermolecular force of attraction
  • H in one molecule is attatched to the very electronegative elements
    • O, N, F: in a neighboring molecule
  • the molecules have very high dipoles
  • a bare proton is left after the bond (no nucleus)
  • gives water of its special properties

Paramagnetic & Diamagnetic

  • diamagnetic - all electrons are paired
  • paramagnetic - some electrons unpaired, shows strong attraction to external magnetic fields