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Table salt, chalk, baking soda, limestone, the minerals in your bones — all are ionic compounds. They share a set of characteristic properties: they’re solid at room temperature, they melt only at high temperatures, they’re hard but shatter when struck, and they conduct electricity only when melted or dissolved. Every one of these properties follows from how ionic compounds are built. This article explains each property in turn, with real data and the important exceptions.
The structure behind the properties
Ionic compounds form when metals transfer electrons to non-metals, producing positive and negative ions (see ionic bonding explained). These ions attract each other in all directions, building a giant ionic lattice: a regular 3D arrangement in which each ion is surrounded by ions of opposite charge (see ionic lattices).
Keep two facts in mind as you read:
- The attractions are strong and act in every direction.
- The ions are fixed in place in the solid, but can move when the solid melts or dissolves.
Property 1: high melting and boiling points
To melt an ionic solid, you must supply enough energy to overcome the strong electrostatic attractions holding the ions in their lattice positions. Because there are so many attractions, this takes a lot of energy.
| Compound | Ions | Melting point |
|---|---|---|
| Sodium chloride, NaCl | Na⁺, Cl⁻ | 801 °C |
| Potassium chloride, KCl | K⁺, Cl⁻ | 770 °C |
| Sodium fluoride, NaF | Na⁺, F⁻ | 993 °C |
| Calcium oxide, CaO | Ca²⁺, O²⁻ | about 2,600 °C |
| Magnesium oxide, MgO | Mg²⁺, O²⁻ | about 2,850 °C |
Two factors control the melting point
- Charge on the ions: higher charges mean stronger attraction. MgO and NaF have similar ion sizes, but MgO’s ions carry twice the charge (2+ and 2− versus 1+ and 1−), and it melts nearly 1,900 °C higher.
- Size of the ions: smaller ions can get closer together, so they attract more strongly. NaF (small F⁻) melts higher than NaCl (larger Cl⁻), and NaCl higher than KCl (larger K⁺).
Both factors are captured by the lattice energy — the energy released when gaseous ions form the lattice (see lattice energy).
Magnesium oxide’s very high melting point makes it useful as a refractory material, lining furnaces and kilns.
Property 2: hard but brittle
Ionic solids are hard: the ions are held firmly in place, so the surface resists scratching and denting.
But they’re also brittle: hit a salt crystal with a hammer and it shatters rather than bending. Here’s why:
- In the lattice, each ion is next to ions of opposite charge.
- A sharp blow can push one layer of ions sideways by the distance of one ion.
- Now ions of the same charge are lined up next to each other: Na⁺ beside Na⁺, Cl⁻ beside Cl⁻.
- These like charges repel strongly, and the crystal splits apart along that plane.
Metals behave very differently: their layers can slide without creating repulsion, because the delocalised electrons keep holding the positive ions together. That’s why metals bend and ionic compounds shatter (see how metallic bonding explains conductivity, lustre and malleability).
Property 3: electrical conductivity — only when ions can move
Electricity is a flow of charged particles. Ionic compounds contain charged particles (ions), but whether they conduct depends on whether the ions can move:
| State | Conducts? | Why |
|---|---|---|
| Solid | No | Ions are locked in the lattice |
| Molten (liquid) | Yes | Ions are free to move |
| Dissolved in water | Yes | Ions separate and move through the solution |
When an electric current passes through a molten or dissolved ionic compound, the ions move to the electrodes: positive ions (cations) to the negative electrode, negative ions (anions) to the positive electrode. There, they gain or lose electrons and form new substances. This is electrolysis. Molten sodium chloride can be electrolysed to produce sodium metal and chlorine gas, and molten aluminium oxide (dissolved in cryolite) is electrolysed to extract aluminium.
It’s important to note that ionic compounds don’t conduct by electrons moving through them, as metals do — the ions themselves carry the charge.
Property 4: solubility in water
Many ionic compounds dissolve in water. Water molecules are polar — oxygen is δ− and hydrogen is δ+ (see the shape of a water molecule). When an ionic crystal is placed in water:
- the δ− oxygen ends of water molecules are attracted to positive ions;
- the δ+ hydrogen ends are attracted to negative ions;
- water molecules pull ions out of the lattice and surround them (hydration).
Whether a compound dissolves depends on a balance: the energy needed to break up the lattice versus the energy released when the ions are hydrated. That’s why not all ionic compounds are soluble. Sodium chloride dissolves readily (about 36 g per 100 g of water at 20 °C), but calcium carbonate (chalk, limestone), barium sulfate and silver chloride are almost insoluble (see solubility rules).
Ionic compounds are generally insoluble in non-polar solvents such as hexane, which can’t stabilise the ions.
Property 5: crystalline form
Because the ions are arranged in a regular repeating pattern, ionic compounds tend to form crystals with flat faces and definite angles. Sodium chloride crystals are cubic; slowly grown copper(II) sulfate crystals form blue parallelograms. The shape of the crystal reflects the shape of the unit cell.
Summary table
| Property | Typical ionic compound | Explanation |
|---|---|---|
| Melting point | High | Strong attractions between many ions in all directions |
| State at 20 °C | Solid | As above |
| Hardness | Hard | Ions held firmly in the lattice |
| Brittleness | Brittle | Shifting layers brings like charges together → repulsion |
| Conducts as solid | No | Ions fixed |
| Conducts molten/dissolved | Yes | Ions free to move |
| Solubility in water | Often soluble (not always) | Polar water hydrates the ions |
| Solubility in non-polar solvents | Insoluble | No attraction to stabilise ions |
Exceptions and grey areas
- Some ionic compounds are insoluble in water, as noted above.
- Some compounds of metals and non-metals are largely covalent: aluminium chloride sublimes at about 180 °C and exists as Al₂Cl₆ molecules in the vapour — a sign of significant covalent character (see ionic vs covalent bonds).
- Ionic liquids — salts made of large, awkwardly shaped ions — are liquid at room temperature, because their ions can’t pack efficiently. They’re used as specialist solvents.
Common misconceptions
- “Solid salt conducts because it has ions.” The ions can’t move in the solid.
- “Molten salt conducts because of free electrons.” The current is carried by moving ions.
- “All ionic compounds dissolve in water.” Many don’t.
- “Brittle means weak.” Ionic solids are strong and hard, but can’t bend.
Key takeaways
- Ionic compounds form giant lattices of ions held by strong attractions in all directions.
- They have high melting points, which rise with higher ion charges and smaller ions.
- They’re hard but brittle: shifting layers brings like charges together.
- They conduct only when molten or dissolved, because the ions must be free to move.
- Many dissolve in water because polar water molecules hydrate the ions.
Compare them with properties of covalent compounds.
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