Revision on Bonding (AQA Spec)

Overview

• This video provides an overview of bonding for the AQA specification.
• PowerPoint slides used in the video can be purchased via a link in the description.

Ionic Bonding

• Ionic bonding involves charged ions held together by strong electrostatic attractions.
• Example: Sodium (Na) and Chlorine (Cl) form NaCl.
  • Sodium donates an electron to Chlorine, forming oppositely charged ions.
• Ionic compounds have regular cubic structures.
• Group Trends:
  • Group 1: +1 ions
  • Group 2: +2 ions
  • Group 3: +3 ions
  • Group 5: -3 ions
  • Group 6: -2 ions
  • Group 7: -1 ions
• Molecular Ions: (
  • Hydroxide (OH⁻)
  • Nitrate (NO₃⁻)
  • Ammonium (NH₄⁺)
  • Sulfate (SO₄²⁻)
  • Carbonate (CO₃²⁻))

Properties of Ionic Compounds

• Conduct electricity when molten or in solution due to free-moving ions.
• High melting points due to strong electrostatic attractions.
• Often soluble in water (polar solvent).

Covalent Bonding

• Covalent bonding involves sharing electrons between atoms.
• Single, Double, and Triple Bonds:
  • Single bond: sharing of one pair of electrons.
  • Double bond: sharing two pairs of electrons.
  • Triple bond: sharing three pairs of electrons.
• Dative Covalent Bonds:
  • One atom donates both electrons in the bond.
  • Example: NH₃ bonding with H⁺ to form NH₄⁺.

Giant Covalent Structures

• Graphite:
  • Layers of hexagonally arranged carbon atoms.
  • High melting point, conducts electricity, low density, insoluble.
• Diamond:
  • Each carbon atom bonded to four others in a 3D structure.
  • Very high melting point, hard, doesn’t conduct electricity, insoluble.

Shapes of Molecules

• Determined by the number of bond pairs and lone pairs:
  • Linear (e.g., BeCl₂): 180°
  • Trigonal Planar (e.g., BF₃): 120°
  • Tetrahedral (e.g., CH₄): 109.5°
  • Trigonal Bipyramidal (e.g., PCl₅): 120° and 90°
  • Octahedral (e.g., SF₆): 90°
  • Bent/Non-linear (e.g., H₂O): <109.5° due to lone pairs
  • Pyramidal (e.g., NH₃): ~107°

Electronegativity

• Electronegativity: an atom's ability to attract electrons in a covalent bond.
• Increases across a period and up a group (excluding noble gases).
• Fluorine is the most electronegative element.

Polar Bonds & Molecules

• Polar Bond: Difference in electronegativity between bonded atoms.
  • Example: HCl (Cl more electronegative, creates dipole).
• Non-Polar Bond: Equal sharing of electrons.
  • Example: Cl₂ (same electronegativity).
• Symmetrical molecules can be non-polar even if they have polar bonds.

Intermolecular Forces

• Van der Waals Forces:
  • Weak forces caused by temporary dipoles.
  • Present in all molecules with electrons.
• Dipole-Dipole Forces:
  • Occur between molecules with permanent dipoles (e.g., HCl).
• Hydrogen Bonding:
  • Strongest intermolecular force.
  • Occurs in molecules with N, O, or F bonded to H.
  • Example: H₂O; explains properties like ice being less dense than water.

Metallic Bonding

• Metals have a giant lattice structure with delocalized electrons.
• Properties:
  • Good electrical and thermal conductors.
  • High melting points.
  • Insoluble in water.