Topic Two: Bonding and Structure - Revision Notes

Introduction

• Video for Edexcel specification on bonding and structure.
• Presented by Chris Harris from AlleryTutors.com.
• Purpose: Overview of topic for revision.
• Slides available for purchase (link in description).

Ionic Bonding

• Definition: Bonding between oppositely charged ions through electrostatic attraction.
• Formation of Ions:
  • An element loses an electron (e.g., Sodium) and another gains an electron (e.g., Chlorine) to achieve full electron shells.
  • Sodium (1 electron) gives to Chlorine (7 electrons) leading to Na⁺ and Cl⁻ ions.

Ion Groups

• Group 1: 1⁺
• Group 2: 2⁺
• Group 3: 3⁺
• Group 5: Gains 3 electrons (3⁻)
• Group 6: 2⁻
• Group 7: 1⁻

Common Molecular Ions

• Hydroxide (OH⁻)
• Nitrate (NO₃⁻)
• Ammonium (NH₄⁺)
• Sulfate (SO₄²⁻)
• Carbonate (CO₃²⁻)

Formula of Ionic Compounds

• Method: Swap and drop charges.
  1. Write ions (e.g., Ca²⁺, NO₃⁻).
  2. Swap the charges.
  3. Drop the charges to lowest terms and simplify.
  • Example: Ca(NO₃)₂ = Calcium Nitrate

Characteristics of Ionic Compounds

• Giant ionic structures (e.g., NaCl).
• High melting and boiling points due to strong electrostatic forces.
• Conduct electricity when molten or dissolved; soluble in water due to polar molecules.
• Brittle due to repulsion when layers shift.

Factors Affecting Ionic Bond Strength

• Charge Size: Larger charges lead to stronger attractions and higher melting points (e.g., CaO has a higher melting point than KCl).
• Ionic Radii: Smaller ions lead to stronger attractions and higher melting points (e.g., NaCl vs KCl).

Trends in Ionic Radii

• Increases down a group (more electron shells).
• Isoelectronic ions: Same number of electrons, differing protons lead to smaller ionic radii with increasing atomic number.

Covalent Bonding

• Definition: Sharing of outer electrons to obtain full shells.
• Types of bonds:
  • Single Bond: 2 shared electrons
  • Double Bond: 4 shared electrons
  • Triple Bond: 6 shared electrons

Dative Covalent Bonds

• One atom donates both electrons (e.g., NH₃ and H⁺ forming NH₄⁺).

Bond Enthalpy

• Strength of a bond related to bond length; shorter bonds = higher enthalpy.

Molecular Shapes

• Shapes determined by bond pairs and lone pairs.
• Lone pairs exert greater repulsion than bond pairs, altering bond angles.
• Common shapes include:
  • Linear (180°)
  • Trigonal Planar (120°)
  • Tetrahedral (109.5°)
  • Octahedral (90°)

Giant Covalent Structures

• Graphite:
  • Layers allow sliding, conducts electricity.
  • High melting point due to strong covalent bonds.
• Diamond:
  • Rigid structure with high melting point; does not conduct electricity.
• Graphene:
  • One layer of graphite, excellent electrical conductor, strong.

Metallic Bonding

• Occurs in metals; positive ions in a sea of delocalized electrons.
• Properties:
  • Good conductors of heat and electricity.
  • High melting points.
  • Malleable and ductile.

Electronegativity

• Ability of an atom to attract electrons in a bond.
• Fluorine is most electronegative (4.0).
• Differences in electronegativity determine bond type (ionic or covalent).

Polar Bonds

• Covalent bonds can be polar if there's a difference in electronegativity.
  • Example: HCl (polar) vs Cl₂ (non-polar).

Intermolecular Forces

• London Forces: Weakest forces, exist between any atoms/molecules.
• Permanent Dipole-Dipole: Occur in polar molecules (e.g., HCl).
• Hydrogen Bonding: Strongest force, occurs between hydrogen and highly electronegative atoms (N, O, F).

Solubility

• Polar substances dissolve in polar solvents (like water).
• Nonpolar substances dissolve in nonpolar solvents (like hydrocarbons).

Conclusion

• Summary of bonding types and their characteristics.
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