Ionic compounds are formed when a metal transfers electrons to a non-metal, producing oppositely charged ions held together by strong electrostatic forces in a giant lattice. This structure gives them high melting and boiling points, makes them brittle solids, and allows them to conduct electricity only when the ions are free to move.
What is the giant ionic lattice structure?
In an ionic compound, positive ions (cations) and negative ions (anions) arrange themselves into a regular, repeating three-dimensional structure called a giant ionic lattice. Each ion is surrounded by several ions of the opposite charge, maximising the attractive forces between them. Sodium chloride (NaCl), for example, arranges its Na⁺ and Cl⁻ ions so that each Na⁺ is surrounded by 6 Cl⁻ ions, and each Cl⁻ is surrounded by 6 Na⁺ ions.
The lattice structure is not a molecule — there is no defined unit of "one NaCl pair" in a solid crystal. The formula NaCl simply indicates the ratio of ions (1:1). A single grain of table salt contains many billions of ions arranged in this repeating lattice.
Why do ionic compounds have high melting and boiling points?
To melt an ionic compound, you must supply enough energy to overcome the strong electrostatic attractions between oppositely charged ions throughout the entire lattice. These attractions are strong because:
- The forces act in all directions (each ion attracts all neighbouring opposite-charge ions).
- The charges involved (e.g. 2+ and 2− for calcium oxide) can be large.
- The ions are small and closely packed, keeping the charges near to each other.
| Ionic compound | Ions | Melting point |
|---|---|---|
| NaCl (sodium chloride) | Na⁺ and Cl⁻ (+1, −1) | 801 °C |
| MgO (magnesium oxide) | Mg²⁺ and O²⁻ (+2, −2) | 2,852 °C |
| CaCl₂ (calcium chloride) | Ca²⁺ and Cl⁻ (+2, −1) | 772 °C |
Magnesium oxide has a much higher melting point than sodium chloride because its ions carry larger charges (+2 and −2 vs +1 and −1), making the electrostatic attractions between ions roughly four times stronger (Coulomb's law: force ∝ q₁q₂).
Why are ionic compounds brittle?
In the giant ionic lattice, each ion is surrounded by ions of opposite charge in fixed positions. When an external force tries to distort the lattice by shifting layers of ions, ions of the same charge are brought next to each other. These like charges repel strongly, causing the crystal to cleave (split) along a plane rather than deform gradually.
This brittleness is the opposite of the behaviour of metals, whose atoms can slide past one another (because the delocalised electrons are not attached to fixed positions) — which is why metals are malleable and ionic solids are not. Ionic crystals break with a sharp, clean fracture rather than bending.
When do ionic compounds conduct electricity?
Ionic compounds conduct electricity only when their ions are free to move. The rule is:
| State | Ions free to move? | Conducts electricity? |
|---|---|---|
| Solid ionic compound | No — ions are fixed in lattice positions | No |
| Molten ionic compound | Yes — ions can move through the liquid | Yes |
| Aqueous solution (dissolved in water) | Yes — ions are free to move through solution | Yes |
In a solid, the ions are locked in place by the lattice forces — they vibrate but cannot migrate. In a melt or solution, the lattice has been disrupted, ions are mobile, and a current can be carried through the substance by the movement of these charged particles towards oppositely charged electrodes.
This is why electrolysis requires either a molten ionic compound or an aqueous ionic solution — the solid form cannot be electrolysed.
Why do many ionic compounds dissolve in water?
Water molecules are polar: the oxygen end carries a partial negative charge (δ−) and each hydrogen end carries a partial positive charge (δ+). When an ionic lattice is placed in water, water molecules cluster around the surface ions:
- The δ− oxygen atoms surround and attract positive ions (cations).
- The δ+ hydrogen atoms surround and attract negative ions (anions).
This process — called hydration — releases energy as the ions are stabilised by their water shells. If the hydration energy released is greater than the lattice energy (the energy needed to pull the lattice apart), the compound dissolves. For many salts this condition is met; for others (such as calcium carbonate, CaCO₃), the lattice energy is too large and the compound is insoluble.
How do the properties of ionic compounds compare with covalent substances?
| Property | Ionic compound | Simple molecular covalent substance |
|---|---|---|
| Melting point | High (hundreds to thousands of °C) | Low (often below room temperature) |
| Electrical conductivity | Only when molten or dissolved | None (unless reacts with water to produce ions) |
| Solubility in water | Often soluble | Often insoluble (exceptions: small polar molecules like HCl, NH₃) |
| Physical state at room temperature | Solid | Often gas or liquid |
| Structure | Giant ionic lattice | Discrete small molecules |
The contrast arises from the fundamental difference in bonding: ionic compounds are held together by strong electrostatic forces throughout the entire lattice, while simple molecular covalent substances are held together by much weaker intermolecular forces (London dispersion forces, dipole–dipole interactions) between discrete molecules.
Frequently asked questions
Why can ionic compounds conduct electricity when dissolved but not when solid?
Electrical conduction requires charge carriers that can move through the material. In a solid ionic compound, the ions are locked in fixed lattice positions and cannot migrate even though they carry charge. When the compound dissolves in water or melts, the lattice breaks up and ions become free to move. These mobile ions carry current through the liquid — positive ions move towards the negative electrode, negative ions towards the positive electrode. This is ionic conduction, as distinct from electronic conduction in metals.
Why does magnesium oxide have a much higher melting point than sodium chloride?
Magnesium oxide (MgO) contains Mg²⁺ and O²⁻ ions, each carrying a charge of magnitude 2. Sodium chloride (NaCl) contains Na⁺ and Cl⁻ ions, each carrying a charge of magnitude 1. Electrostatic force is proportional to the product of the charges (Coulomb's law), so the force in MgO is approximately 2 × 2 = 4 times greater than in NaCl, and the ions are also smaller, keeping them closer together. More energy is therefore needed to pull the MgO lattice apart, giving it a melting point of 2,852 °C compared with 801 °C for NaCl.
What is the difference between ionic and covalent bonding in terms of properties?
Ionic bonding involves the transfer of electrons between a metal and a non-metal, forming oppositely charged ions that attract each other in a giant lattice. This gives high melting points, brittleness, and electrical conductivity in the molten or dissolved state. Covalent bonding involves the sharing of electrons between non-metals, forming discrete molecules held together by much weaker intermolecular forces. This typically gives low melting points, often gas or liquid states at room temperature, and no electrical conductivity.
Can all ionic compounds be dissolved in water?
No — solubility in water depends on whether the energy released by hydration (water molecules surrounding and stabilising the ions) is sufficient to overcome the lattice energy (the energy holding the ions together). Most sodium, potassium, and ammonium salts are soluble. Many carbonates, sulfates, and hydroxides are insoluble or only slightly soluble. For example, barium sulfate (BaSO₄) is so insoluble (Ksp ≈ 1 × 10⁻¹⁰ at 25 °C) that it is used as an X-ray contrast agent — it can be swallowed safely because it does not dissolve in digestive fluids.
For Socratic GCSE chemistry with Professor Curie — predicting from the particle picture whether an ionic solid will melt at a low or high temperature before comparing any two compounds — visit aitutors.me.