Ionic vs Covalent Bonds

• Chemical bonds are attractive forces that hold atoms together.

Ionic Bonding

• Ionic bonding is the electrostatic attraction between oppositely charged ions.
• It occurs after the transfer of electrons from one atom to another.
• Ionic bonds typically form between metals and non-metals.
• Metals lose electrons to form positive ions (cations).
• Non-metals gain electrons to form negative ions (anions).
• The resulting ions have stable electron configurations, usually with a full outer shell (octet rule).
• Example: Sodium chloride (NaCl) is formed by the transfer of an electron from sodium (Na) to chlorine (Cl).
• Ionic compounds form giant ionic lattices.
• Physical properties of ionic compounds:
  • High melting and boiling points due to strong electrostatic forces.
  • Conduct electricity when molten or dissolved, as ions are free to move.
  • Often soluble in polar solvents.
  • Tend to be brittle.

Covalent Bonding

• Covalent bonding is the sharing of electron pairs between atoms.
• It occurs when atoms need to gain electrons to achieve a full outer shell but neither atom is able to fully remove electrons from the other.
• Covalent bonds typically form between non-metals.
• Shared electrons are attracted to the nuclei of both atoms, holding them together.
• A single covalent bond involves one shared pair of electrons.
• A double covalent bond involves two shared pairs of electrons.
• A triple covalent bond involves three shared pairs of electrons.
• Example: Methane (CH4) has four single covalent bonds between carbon and hydrogen.
• Covalent compounds can form simple molecular structures or giant covalent structures.

Simple Molecular Structures

• Simple molecular structures consist of a fixed number of atoms covalently bonded together.
• There are weak intermolecular forces between molecules.
• Physical Properties of simple molecular structures:
  • Low melting and boiling points due to weak intermolecular forces
  • Don't conduct electricity because they have no mobile ions or delocalised electrons
  • Often insoluble in water

Giant Covalent Structures

• Giant covalent structures consist of a large network of covalently bonded atoms.
• Examples: Diamond, graphite, silicon dioxide (SiO2).
• Physical properties of giant covalent structures:
  • High melting and boiling points due to strong covalent bonds throughout the structure.
  • Diamond is very hard and doesn't conduct electricity.
  • Graphite conducts electricity due to delocalized electrons between layers.
  • Graphite has layers that can slide over each other, making it slippery.

Metallic Bonding

• Metallic bonding is the electrostatic attraction between positive metal ions and delocalized electrons.
• It occurs in metals.
• Metal atoms lose their outer electrons, forming positive ions.
• The released electrons become delocalized, forming a "sea" of electrons around the metal ions.
• The delocalized electrons are free to move throughout the structure.
• Metals form giant metallic lattices.
• Physical properties of metals:
  • High melting and boiling points due to strong metallic bonds.
  • Conduct electricity and heat due to delocalized electrons.
  • Malleable and ductile because the layers of ions can slide over each other without breaking bonds.
  • Lustrous (shiny) due to the reflection of light by delocalized electrons.

Intermolecular Forces

• Intermolecular forces are weak attractive forces between molecules.
• They are much weaker than covalent or ionic bonds.
• Types of Intermolecular Forces:
  • Van der Waals forces (London Dispersion Forces)
    • Present in all molecules.
    • Temporary dipoles are caused by momentary uneven distribution of electrons.
    • Larger molecules have stronger London dispersion forces.
    • Induced dipole-dipole interactions
  • Dipole-Dipole interactions
    • Occurs between polar molecules.
    • Stronger than London dispersion forces.
    • Positive end of one polar molecule attracts the negative end of another.
  • Hydrogen Bonding
    • Strongest type of intermolecular force.
    • Occurs between molecules containing hydrogen bonded to a highly electronegative atom (oxygen, nitrogen, or fluorine).
    • Hydrogen atom is attracted to a lone pair of electrons on the electronegative atom of another molecule.

Bond Polarity

• Bond polarity is the unequal sharing of electrons in a covalent bond.
• It arises when the atoms in a covalent bond have different electronegativities.
• Electronegativity is the ability of an atom to attract electrons in a covalent bond.
• A polar bond has a dipole, with a partial positive charge (δ+) on one atom and a partial negative charge (δ-) on the other.
• If the electronegativity difference is large enough, the bond becomes ionic.
• Nonpolar bond: electrons shared equally.
• Polar bond: electrons shared unequally, resulting in a dipole moment.
• Example: In water (H2O), oxygen is more electronegative than hydrogen, so the O-H bonds are polar.

Shapes of Molecules

• Valence Shell Electron Pair Repulsion (VSEPR) theory predicts the shapes of molecules based on the repulsion of electron pairs around a central atom.
• Electron pairs repel each other and position themselves as far apart as possible to minimize repulsion.
• Electron pairs can be bonding pairs (shared in covalent bonds) or lone pairs (non-bonding).
• Lone pairs repel more strongly than bonding pairs, affecting the bond angles.
• Common molecular shapes:
  • Linear: Two electron pairs around the central atom (180° bond angle). Example: Carbon dioxide (CO2).
  • Trigonal planar: Three electron pairs around the central atom (120° bond angle). Example: Boron trifluoride (BF3).
  • Tetrahedral: Four electron pairs around the central atom (109.5° bond angle). Example: Methane (CH4).
  • Pyramidal: Three bonding pairs and one lone pair around the central atom (approximately 107° bond angle). Example: Ammonia (NH3).
  • Bent or V-shaped: Two bonding pairs and two lone pairs around the central atom (approximately 104.5° bond angle). Example: Water (H2O).
  • Octahedral: Six electron pairs around the central atom (90° bond angle). Example: Sulfur hexafluoride (SF6).