Describe ionic bonding as the electrostatic attraction between oppositely charged ions in a regular three-dimensional lattice, predict the formula of binary ionic compounds, and relate physical properties (melting point, electrical conductivity, brittleness, solubility) to lattice structure

What this dot point is asking

QCAA wants you to describe ionic bonding as electrostatic attraction between cations and anions arranged in a 3D lattice, predict the formula of a binary ionic compound from the charges of its ions, and explain the characteristic properties of ionic compounds (high melting points, conductivity only when molten or dissolved, brittleness, variable water solubility) in terms of the lattice structure.

The answer

An ionic bond is the electrostatic attraction between oppositely charged ions, typically formed by complete transfer of one or more electrons from a metal atom to a non-metal atom. In the solid state, the ions are arranged in a regular three-dimensional lattice in which each ion is surrounded by ions of opposite charge. The properties of ionic compounds follow from the strength and geometry of this lattice.

Forming an ionic compound

Metals lose electrons to form cations (positive ions); non-metals gain electrons to form anions (negative ions). Both species typically reach a noble-gas electron configuration (closed-shell stability).

Worked examples:

• Na: $...3s^1 \rightarrow Na^+: ...2p^6$ (Ne configuration).
• Mg: $...3s^2 \rightarrow Mg^{2+}: ...2p^6$ (Ne configuration).
• Al: $...3s^2 3p^1 \rightarrow Al^{3+}: ...2p^6$ (Ne configuration).
• Cl: $...3s^2 3p^5 \rightarrow Cl^-: ...3p^6$ (Ar configuration).
• O: $...2s^2 2p^4 \rightarrow O^{2-}: ...2p^6$ (Ne configuration).

The electrostatic attraction between cation and anion is the ionic bond. It is not directional in the way a covalent bond is; an ion attracts every oppositely charged ion in its vicinity.

Predicting formulae

The compound must be electrically neutral overall. Use the smallest whole-number ratio that balances charge.

The cross-over method (swap charge magnitudes as subscripts, then simplify) is reliable for binary ionic compounds. For polyatomic ions (NH_4^+, NO_3^-, SO_4^2-, CO_3^2-, OH-, PO_4^3-), use brackets when there is more than one of the polyatomic unit.

The ionic lattice

In the solid state, ions are arranged in a regular repeating 3D pattern (a crystal lattice). The geometry depends on the relative sizes of the cation and anion (the radius ratio). Common QCE-relevant lattices:

• Rock-salt structure (NaCl). Each Na+ is surrounded by 6 Cl-; each Cl- is surrounded by 6 Na+. Cubic.
• Caesium chloride structure (CsCl). Each ion has 8 nearest neighbours of opposite charge. Suited to similar-sized cations and anions.
• Fluorite structure (CaF_2). Each Ca^2+ has 8 F-; each F- has 4 Ca^2+.

The lattice energy is the energy released when gaseous ions combine to form the solid lattice. It governs melting point and solubility.

where q_1 and q_2 are the charges and r is the inter-ionic distance. Higher charges and smaller ions give a larger lattice energy.

Physical properties from lattice structure

High melting and boiling points

Strong electrostatic attractions throughout the lattice; large lattice energy. Comparing within ionic compounds, higher charges and smaller ions give higher melting points. MgO (charges 2+/2-) melts above 2800 degrees C; NaCl (1+/1-) melts at 801 degrees C.

Hard but brittle

The lattice resists deformation, so the solid is hard. But under stress, one layer of ions can shift over another so that like charges align; the resulting repulsion shatters the crystal along that plane. This is why ionic crystals cleave cleanly under impact.

Electrical conductivity only when molten or dissolved

Conduction requires mobile charge carriers. In the solid, ions are fixed in the lattice. In the molten state, the lattice has broken down and ions can migrate. In aqueous solution, ions are dissociated and surrounded by water (solvation shells) and free to move. Solid ionic conductors are an exception covered in materials chemistry, not in QCE Unit 1.

Variable solubility in water

Water is a polar solvent that solvates ions effectively. Solubility depends on the balance between lattice energy (must be overcome to dissociate the solid) and hydration energy (released when ions are solvated). Compounds with comparable lattice and hydration energies tend to dissolve; compounds with very high lattice energy relative to hydration energy do not. AgCl is famously insoluble (very high lattice energy, modest hydration energy); NaCl is very soluble. QCE Chemistry provides solubility rules to memorise; Unit 2 builds on these.

Naming ionic compounds

• Cation name first, anion name second.
• Monatomic anion: use the element root plus -ide (chloride, oxide, nitride).
• Polyatomic anion: use the standard name (sulfate, nitrate, carbonate, hydroxide, phosphate).
• For transition metals with variable oxidation state, indicate the cation charge with a roman numeral: iron(II) chloride is FeCl_2; iron(III) chloride is FeCl_3.

Worked example: predicting properties

Question. Predict whether sodium oxide or magnesium oxide has the higher melting point, and explain your reasoning.

Na_2O has charges +1 / -2; MgO has charges +2 / -2. The product of the charges (proportional to lattice energy) is 1 x 2 = 2 for Na_2O versus 2 x 2 = 4 for MgO. Also Mg^2+ is smaller than Na+ (smaller r), increasing the lattice energy further. Therefore MgO has the higher melting point. Actual values: Na_2O 1132 degrees C; MgO 2852 degrees C. Consistent with the prediction.

This kind of comparative-strength reasoning is a routine QCAA EA short response.

Examples in context

Example 1. Galena and sphalerite at Mount Isa. Mount Isa Mines processes lead sulfide ($\text{PbS}$, galena) and zinc sulfide ($\text{ZnS}$, sphalerite). Both are ionic lattices: $\text{Pb}^{2+}$ and $\text{Zn}^{2+}$ cations balanced by $\text{S}^{2-}$ anions in $1\!:\!1$ ratio. Lattice energy magnitudes are about $3{,}030 \, \text{kJ mol}^{-1}$ (PbS) and $3{,}625 \, \text{kJ mol}^{-1}$ (ZnS), reflecting the smaller $\text{Zn}^{2+}$ ionic radius ($74 \, \text{pm}$ versus $119 \, \text{pm}$ for $\text{Pb}^{2+}$). Stronger lattice energy means $\text{ZnS}$ has a higher melting point ($1{,}185^{\circ}\text{C}$ versus $1{,}113^{\circ}\text{C}$). Both are brittle: layer slip brings like charges together causing fracture along cleavage planes during crushing.

Example 2. Magnesium chloride brine from Gladstone alumina red mud. Bayer-process residue from Queensland Alumina at Gladstone contains residual $\text{MgCl}_2$ at $\sim 5\%$. Magnesium chloride forms when $\text{Mg}$ transfers two electrons to two $\text{Cl}$ atoms: $\text{Mg} \rightarrow \text{Mg}^{2+} + 2 e^-$ and $\text{Cl} + e^- \rightarrow \text{Cl}^-$, giving formula $\text{MgCl}_2$. Solid $\text{MgCl}_2$ does not conduct because $\text{Mg}^{2+}$ and $\text{Cl}^-$ ions are locked in the lattice. Once dissolved in red-mud leachate the ions become mobile and conductance jumps, which operators use as a real-time process monitor for residue dewatering.

Try this

Q1. Predict and write the formula of the ionic compound formed between (a) aluminium and oxygen, (b) calcium and phosphate. [4 marks]

• Cue. (a) $\text{Al}_2 \text{O}_3$, $2 \times 3+ = 3 \times 2-$. (b) $\text{Ca}_3 (\text{PO}_4)_2$.

Q2. Lattice energies (in kJ mol$^{-1}$): $\text{NaCl}\,786$, $\text{MgO}\,3{,}795$, $\text{CaO}\,3{,}414$. Account for the differences. [4 marks]

• Cue. Charge magnitude ($1+/1-$ vs $2+/2-$) dominates; ionic radius ($\text{Mg}^{2+}$ smaller than $\text{Ca}^{2+}$) explains MgO $>$ CaO.

Q3. A student tests solid and molten $\text{KBr}$ with a conductivity meter. (a) Describe the bonding. (b) Predict observations in both states. (c) Explain in terms of ion mobility. [2+2+3 marks]

• Cue. (a) Ionic, $\text{K}^+$ and $\text{Br}^-$ in 3-D lattice. (b) Solid no current, molten conducts. (c) Lattice locks ions; melting frees them. ISMG application and communication.