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Ions and ionic bonds

7 learning objectives • 4 core • 3 extended

Topic 2.4: Ions and Ionic Bonds Revision Notes

1. Overview

Ionic bonding is the process by which atoms achieve a stable, full outer shell of electrons by transferring electrons from metals to non-metals. This topic explores how these transfers create charged particles called ions and how the resulting electrostatic forces create strong, high-melting-point structures that are essential to chemical life and industry.

Key Definitions

  • Ion: An atom or group of atoms that has become electrically charged by losing or gaining electrons.
  • Cation: A positively charged ion, formed when a metal atom loses electrons.
  • Anion: A negatively charged ion, formed when a non-metal atom gains electrons.
  • Ionic Bond: The strong electrostatic attraction between oppositely charged ions.
  • Giant Lattice: A regular, repeating three-dimensional arrangement of alternating positive and negative ions.
  • Valency: The combining power of an element, dictated by the number of electrons an atom needs to lose or gain to achieve a full outer shell.

Core Content

The Formation of Ions

Atoms react to achieve a full outer shell of electrons (a stable noble gas configuration).

  • Metals: Located in Groups I, II, and III. They lose their outer shell electrons to form positive cations.
    • Example: A Sodium atom ($Na$) loses 1 electron to become a Sodium ion ($Na^{+}$).
  • Non-metals: Located in Groups V, VI, and VII. They gain electrons to form negative anions.
    • Example: A Chlorine atom ($Cl$) gains 1 electron to become a Chloride ion ($Cl^{-}$).

Ionic Bonding in Group I and Group VII

When a Group I metal (e.g., Lithium) reacts with a Group VII non-metal (e.g., Fluorine), one electron is transferred from the metal to the non-metal.

Word Equation: Lithium(s) + Fluorine(g) → Lithium fluoride(s)

Symbol Equation: $2Li(s) + F_2(g) \rightarrow 2LiF(s)$

📊Dot-and-cross diagram showing a Lithium atom with one outer electron and an arrow pointing to a Fluorine atom with seven outer electrons. The resulting diagram shows $[Li
^+$ and $[F]^-$ with square brackets and charges, with Fluorine's eighth electron represented by a different symbol (e.g., a cross among dots).]

Properties of Ionic Compounds

  1. High Melting and Boiling Points: They are solids at room temperature.
  2. Electrical Conductivity:
    • Solid: Poor conductor (insulator).
    • Molten (l) or Aqueous (aq): Good conductor.

Extended Content (Extended Only)

The Giant Ionic Lattice

Ionic compounds do not exist as simple molecules. Instead, they form a giant lattice structure. This is a regular arrangement of alternating positive and negative ions extending in three dimensions. The structure is held together by strong electrostatic attractions acting in all directions.

📊A 3D cube-like grid showing alternating spheres labeled + and - representing ions in a Sodium Chloride lattice.

Formation of Bonds between Metallic and Non-Metallic Elements

For elements not in Group I or VII, the number of electrons transferred must ensure the final compound is neutral.

  • Example: Magnesium (Group II) and Oxygen (Group VI)
    • $Mg$ loses 2 electrons to become $Mg^{2+}$.
    • $O$ gains 2 electrons to become $O^{2-}$.
    • $Mg(s) + \frac{1}{2}O_2(g) \rightarrow MgO(s)$
  • Example: Magnesium (Group II) and Chlorine (Group VII)
    • One $Mg$ atom transfers 1 electron to each of two $Cl$ atoms.
    • Result: $Mg^{2+}$ and $2 \times Cl^{-}$ ions, forming $MgCl_2$.

Explaining Properties via Structure and Bonding

  • High Melting/Boiling Points: The giant lattice contains millions of strong electrostatic attractions (ionic bonds). A large amount of heat energy is required to overcome these strong forces to break the lattice.
  • Electrical Conductivity:
    • Solid: In the solid state, ions are held in fixed positions within the lattice and cannot move. Therefore, they cannot carry an electric current.
    • Molten/Aqueous: When melted or dissolved in water, the lattice breaks down and the ions become free to move. These mobile ions act as charge carriers, allowing electricity to flow.

Key Equations

Process Equation Notes
Formation of Sodium ion $Na \rightarrow Na^+ + e^-$ Loss of $e^-$ (Oxidation)
Formation of Oxide ion $O + 2e^- \rightarrow O^{2-}$ Gain of $e^-$ (Reduction)
Formation of Magnesium Chloride $Mg(s) + Cl_2(g) \rightarrow MgCl_2(s)$ Balanced neutral compound
Dissolving Sodium Chloride $NaCl(s) \rightarrow Na^+(aq) + Cl^-(aq)$ Showing free ions in solution

Common Mistakes to Avoid

  • ❌ Wrong: "Ionic compounds conduct electricity because electrons move through them."
  • ✅ Right: Ionic compounds conduct because ions are free to move when molten or aqueous. (Electrons only move in metals/graphite).
  • ❌ Wrong: Describing $NaCl$ as a "molecule."
  • ✅ Right: Describe it as a giant lattice or a formula unit.
  • ❌ Wrong: Forgetting to put square brackets or charges on dot-and-cross diagrams.
  • ✅ Right: Always use $[ ]$ and indicate the charge (e.g., $2+$) at the top right for ions.

Exam Tips

  • Command Words: If asked to "Describe" the bonding, talk about electron transfer from metal to non-metal. If asked to "Explain" the melting point, you must mention the "strong electrostatic attractions" and "large amount of energy needed to break them."
  • Diagrams: Read the question carefully—does it ask for all electron shells or only the outer shell? Most IGCSE questions only require the outer shell.
  • Formulas: To find the formula of an ionic compound, swap and drop the numerical values of the charges.
    • Example: $Al^{3+}$ and $O^{2-}$ becomes $Al_2O_3$.
  • State Symbols: Always include them in equations unless the question says otherwise. $(s)$ for solid, $(l)$ for liquid, $(g)$ for gas, and $(aq)$ for aqueous (dissolved in water).

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