Properties of Ionic Compounds
Ionic compounds are formed when electrons are transferred from a metal atom to a nonmetal atom, producing oppositely charged ions that attract each other in a three-dimensional lattice structure. This arrangement — known as the ionic or crystal lattice — is not merely a pair of ions but an enormous, repeating network of alternating cations and anions held together by strong electrostatic forces. The physical and chemical properties of ionic compounds are direct consequences of this lattice architecture. The most striking properties of ionic compounds include their characteristically high melting and boiling points, their ability to dissolve in polar solvents like water, and their capacity to conduct electricity when molten or dissolved in solution — but not as a solid. Each of these properties can be explained at the atomic and molecular level by considering how the ions are arranged and how they interact with each other and with their environment. Understanding these properties is essential not only for passing examinations but for real-world applications: ionic compounds form the basis of table salt (NaCl), many fertilizers (KNO₃), antacids (CaCO₃), and industrial salts used in metallurgy. Predicting how an ionic compound will behave under different conditions requires a solid grasp of lattice energy, electrostatic attraction, and polarity.
- High Melting and Boiling Points: The Role of Lattice Energy
- Solubility in Polar Solvents: Ion-Dipole Interactions
- Electrical Conductivity: Free Ions as Charge Carriers
- Brittleness: A Consequence of Lattice Geometry
High Melting and Boiling Points: The Role of Lattice Energy
Ionic compounds have high melting points because a enormous amount of energy is needed to overcome the strong electrostatic attractions holding the lattice together. This energy is quantified as the lattice energy (U), defined as the energy released when gaseous ions come together to form one mole of ionic solid. For NaCl, this value is approximately −787 kJ/mol — a very large negative number indicating highly stable attractions. Coulomb's law governs this attraction: F = kq₁q₂/r², meaning that higher ionic charges and smaller ionic radii both increase the lattice energy. This is why MgO (charges 2+ and 2−, small radii) melts at 2852 °C, while NaCl (charges 1+ and 1−) melts at only 801 °C. In both cases, however, these melting points are far higher than those of most covalent molecular substances.
Solubility in Polar Solvents: Ion-Dipole Interactions
When an ionic compound like NaCl is placed in water, the polar water molecules are attracted to the surface ions. The partially negative oxygen atoms of water (δ−) are attracted to Na⁺ cations, while the partially positive hydrogen atoms (δ+) are attracted to Cl⁻ anions. This process — called hydration — releases energy (the hydration enthalpy), which can compensate for the energy required to break apart the lattice (the lattice energy). If the hydration enthalpy is large enough to overcome the lattice energy, the compound dissolves. Compounds with very high lattice energies (like MgO or Al₂O₃) are essentially insoluble in water because the hydration energy cannot compensate for the energy required to separate the ions. The general rule: ionic compounds dissolve in polar solvents because polar solvent molecules can stabilize the separated ions through ion-dipole interactions. They do not dissolve in non-polar solvents like hexane because non-polar molecules cannot provide the necessary electrostatic stabilization.
Electrical Conductivity: Free Ions as Charge Carriers
In the solid state, ions in an ionic lattice are held rigidly in fixed positions and cannot move — therefore, solid ionic compounds do not conduct electricity. When the compound is melted (molten state) or dissolved in water (aqueous solution), the ionic bonds are overcome and the ions become free to move. These mobile ions can carry electrical charge through the liquid: cations migrate toward the negative electrode (cathode) and anions migrate toward the positive electrode (anode). This is electrolytic conduction. For example, molten NaCl undergoes electrolysis: Na⁺ ions are reduced at the cathode to form sodium metal, and Cl⁻ ions are oxidised at the anode to form chlorine gas. This is how sodium metal is industrially produced (the Downs process). Crucially, neither electrons nor lattice vibrations carry the charge — only the physical movement of ions is responsible.
Brittleness: A Consequence of Lattice Geometry
Although not in the short description, brittleness is a classic ionic property worth noting for completeness. When a mechanical force is applied to an ionic crystal, layers of ions can shift relative to each other. When they shift by even one ion spacing, identical charges (cation next to cation, or anion next to anion) are suddenly aligned. The resulting repulsive forces cause the crystal to fracture sharply along cleavage planes. This contrasts with metals, where layers of atoms can slide over each other without this charge-repulsion problem, making metals malleable and ductile.
The Key Equation
- U — Lattice Energy (J mol⁻¹): Energy released when 1 mole of ionic solid forms from its gaseous ions; always negative for a stable lattice
- N_A — Avogadro's Number (mol⁻¹): 6.022 × 10²³ — scales the expression to a molar quantity
- M — Madelung Constant (dimensionless): Depends on crystal geometry; for NaCl structure, M ≈ 1.748
- z^+, z^- — Ionic Charges (dimensionless): Magnitude of the charges on the cation and anion respectively
- e — Elementary Charge (C): 1.602 × 10⁻¹⁹ C, the charge on one proton
- r_0 — Nearest-Neighbour Distance (m): Distance between the centres of adjacent cation and anion in the lattice
Worked Example
Problem. Explain why magnesium oxide (MgO) has a much higher melting point (2852 °C) than sodium chloride (NaCl, 801 °C), even though both have the same rock-salt crystal structure. Then predict whether MgO or NaCl is more soluble in water, and explain why.
Answer. MgO has a higher melting point than NaCl because its doubly charged ions (Mg²⁺ and O²⁻) create much stronger electrostatic attractions and a larger lattice energy. MgO is essentially insoluble in water while NaCl is soluble, because the hydration energy of water molecules cannot compensate for MgO's very large lattice energy.
Origins · 1 minMax Born and how this was discovered
Max Born, alongside Alfred Landé, developed the Born–Landé equation (1918) for calculating lattice energies of ionic crystals from fundamental electrostatic principles. He also co-developed the Born–Haber cycle with Fritz Haber, which allows lattice energies to be determined experimentally using Hess's law. The Born–Haber cycle remains a cornerstone of ionic bonding theory taught in high school and university chemistry. Born won the 1954 Nobel Prize in Physics not for his ionic bonding work, but for his statistical interpretation of the quantum mechanical wave function — showing that |ψ|² gives the probability of finding a particle at a given location.
Related Concepts
- Formation of Ionic Bonds — Understanding how ionic bonds form through electron transfer is the essential prerequisite for understanding why the resulting lattice has the properties described here.
- Lattice Energy and the Born–Haber Cycle — The Born–Haber cycle allows quantitative calculation of lattice energies — the key quantity determining melting point and solubility of ionic compounds.
- Properties of Covalent Compounds — Contrasting the low melting points and poor electrical conductivity of molecular covalent compounds with ionic compounds reinforces understanding of both bond types.
- Electrolysis of Ionic Compounds — Electrolysis directly exploits the electrical conductivity of molten and aqueous ionic compounds, and is a major industrial application of this property.
- Solubility and Intermolecular Forces — The 'like dissolves like' principle and ion-dipole interactions explain why ionic compounds dissolve in polar solvents but not in non-polar ones.