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Ionic Bonding and Ionic Structures

Ionic Bonding and Ionic Structures

This lesson covers ionic bonding in detail, including electron transfer, dot-cross diagrams, lattice structures, properties of ionic compounds and experimental evidence for the existence of ions. A thorough understanding of ionic bonding is essential for AQA A-Level Chemistry specification section 3.1.3.

Electron Transfer and Ion Formation

Ionic bonding occurs between metals and non-metals. Metal atoms lose electrons to form positive ions (cations), while non-metal atoms gain electrons to form negative ions (anions). The driving force is that both atoms achieve a stable noble gas electron configuration.

Key Definition: An ionic bond is the strong electrostatic attraction between oppositely charged ions.

Examples of Ion Formation

Sodium chloride (NaCl):

  • Na (2,8,1) loses 1 electron → Na⁺ (2,8) — isoelectronic with neon
  • Cl (2,8,7) gains 1 electron → Cl⁻ (2,8,8) — isoelectronic with argon

Magnesium oxide (MgO):

  • Mg (2,8,2) loses 2 electrons → Mg²⁺ (2,8) — isoelectronic with neon
  • O (2,6) gains 2 electrons → O²⁻ (2,8) — isoelectronic with neon

Calcium fluoride (CaF₂):

  • Ca (2,8,8,2) loses 2 electrons → Ca²⁺ (2,8,8) — isoelectronic with argon
  • Each F (2,7) gains 1 electron → F⁻ (2,8) — isoelectronic with neon
  • Two fluorine atoms are needed to accept the two electrons from one calcium atom

Exam Tip: When drawing dot-cross diagrams for ionic compounds, always show the charge on each ion and use square brackets around the ion. The dots and crosses represent electrons from different atoms — they do not represent different types of electron.

Dot-Cross Diagrams for Ionic Compounds

Dot-cross diagrams show the outer-shell electrons of each ion. Electrons from one atom are shown as dots (•) and from the other as crosses (×).

Sodium Chloride (NaCl)

     [Na]⁺    [  :Cl:  ]⁻
               [× ×× × ]
               [×      ×]
               [× •× × ]
               [  ××    ]

Na donates its single 3s electron (shown as •) to Cl. The chloride ion now has 8 outer electrons: 7 original (×) plus 1 from sodium (•).

Magnesium Oxide (MgO)

     [Mg]²⁺    [  :O:  ]²⁻
                [× ×• × ]
                [×      •]
                [× ×× × ]

Mg donates both its 3s electrons to O. The oxide ion has 8 outer electrons: 6 original (×) plus 2 from magnesium (•).

Calcium Chloride (CaCl₂)

Ca donates one electron to each of two Cl atoms. Each chloride ion has one electron (•) from calcium. This shows the 1:2 ratio in the formula.

Factors Affecting the Formation of Ionic Compounds

For ionic bonding to occur, the overall energy change must be favourable. Several factors influence this:

Factor Effect
Ionisation energy of the metal Must be relatively low — metals on the left of the periodic table form cations more easily
Electron affinity of the non-metal Must be sufficiently exothermic — non-metals on the right of the periodic table (especially Group 7) gain electrons readily
Lattice enthalpy The strong electrostatic attractions in the lattice release a large amount of energy, which drives the overall process

Common Misconception: Students often think that ionic bonds form because atoms "want" a full outer shell. In reality, the formation of an ionic compound is energetically favourable because the lattice enthalpy released more than compensates for the energy required to form the ions. The Born-Haber cycle (Lesson 9) makes this clear.

Ionic Lattice Structures

Ionic compounds do not exist as discrete molecules. Instead, they form giant ionic lattices — regular three-dimensional arrangements of alternating positive and negative ions.

Sodium Chloride Lattice

The NaCl lattice has a face-centred cubic structure:

  • Each Na⁺ ion is surrounded by 6 Cl⁻ ions (octahedral coordination)
  • Each Cl⁻ ion is surrounded by 6 Na⁺ ions
  • The coordination number is 6:6
  • The lattice extends in all three dimensions with no discrete molecular units

Caesium Chloride Lattice

The CsCl lattice has a body-centred cubic structure:

  • Each Cs⁺ ion is surrounded by 8 Cl⁻ ions (cubic coordination)
  • Each Cl⁻ ion is surrounded by 8 Cs⁺ ions
  • The coordination number is 8:8
  • Cs⁺ is larger than Na⁺, allowing more anions to pack around it

Properties of Ionic Compounds

Property Explanation
High melting and boiling points Strong electrostatic attractions between ions require a large amount of energy to overcome
Hard but brittle Ions are held in fixed positions; displacing a layer causes like charges to align, and the repulsion shatters the crystal
Conduct electricity when molten or dissolved Ions are free to move and carry charge; solid ionic compounds do not conduct because ions are fixed in the lattice
Often soluble in water Polar water molecules can surround and stabilise the ions (hydration); this depends on whether the hydration enthalpy exceeds the lattice enthalpy

Worked Example: Explaining Brittleness

When a force is applied to an ionic crystal, layers of ions are displaced. This brings ions of the same charge into alignment:

Before force:    + − + − + −
                 − + − + − +
                 + − + − + −

After displacement:
                 + − + − + −
                   − + − + − +
                 + − + − + −

Same-charge ions now adjacent → repulsion → crystal shatters

Exam Tip: When explaining why ionic compounds are brittle, always mention that layers slide so that ions of the same charge become adjacent and repel, causing the crystal to shatter. Simply saying "the bonds break" is insufficient.

Evidence for the Existence of Ions

Electrolysis

When molten ionic compounds are electrolysed, metals are deposited at the cathode (negative electrode) and non-metals are produced at the anode (positive electrode). This demonstrates that the compound contains charged particles that migrate under an electric field.

X-ray Diffraction

X-ray crystallography reveals the positions of ions in the lattice and shows that the electron density is centred on individual ions rather than shared between atoms. The measured internuclear distances are consistent with ionic radii.

Electron Density Maps

Electron density maps from X-ray data show regions of high electron density around each ion with very low electron density between ions. This confirms that electrons have been transferred rather than shared.

Comparison of Atomic and Ionic Radii

  • Cations are smaller than their parent atoms (loss of outer shell and increased effective nuclear charge)
  • Anions are larger than their parent atoms (gain of electrons increases electron-electron repulsion)
Species Radius (pm) Explanation
Na 186 Neutral atom
Na⁺ 95 Lost outer shell; increased Zeff
Cl 99 Neutral atom (covalent radius)
Cl⁻ 181 Gained electron; increased e⁻-e⁻ repulsion
Mg 160 Neutral atom
Mg²⁺ 65 Lost outer shell; high Zeff
O 73 Neutral atom (covalent radius)
O²⁻ 140 Gained 2 electrons; much increased repulsion

Summary

  • Ionic bonds form by electron transfer from metals to non-metals
  • The ionic bond is the electrostatic attraction between oppositely charged ions
  • Dot-cross diagrams show how electrons are transferred
  • Ionic compounds form giant lattices with high coordination numbers
  • Properties include high melting points, brittleness, and electrical conductivity when molten or in solution
  • Evidence for ions comes from electrolysis, X-ray diffraction and electron density maps
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