An ionic compound is a substance made of positively and negatively charged atoms (ions) held together by the electrical attraction between them. Table salt, baking soda, and limestone are all ionic compounds you encounter daily. They form when one atom gives up electrons to another, creating oppositely charged particles that lock together in a rigid, repeating structure.
How Ionic Compounds Form
Ionic compounds start with electron transfer. When a metal atom meets a nonmetal atom, the metal gives away one or more of its outermost electrons, and the nonmetal accepts them. Take sodium chloride (table salt) as the classic example: a sodium atom hands off one electron to a chlorine atom. Sodium becomes positively charged (Na⁺) because it lost a negative electron, and chlorine becomes negatively charged (Cl⁻) because it gained one. These opposite charges pull the two ions together, and that attraction is the ionic bond.
This transfer happens because of a large gap in how strongly the two atoms attract electrons. Chemists measure this with a scale called electronegativity. When the difference between two atoms is above about 1.8 on the Pauling scale, the bond is classified as ionic rather than covalent (where atoms share electrons instead of transferring them). Metals sit at the low end of the electronegativity scale, nonmetals at the high end, which is why ionic compounds almost always involve a metal paired with a nonmetal.
The Crystal Lattice Structure
Ionic compounds don’t exist as isolated pairs of ions. Instead, each positive ion surrounds itself with negative ions, and each negative ion surrounds itself with positive ions, forming a three-dimensional grid called a crystal lattice. In sodium chloride, every sodium ion sits next to six chloride ions, and every chloride ion sits next to six sodium ions. This repeating pattern extends in all directions, which is why salt crystals have flat faces that meet at sharp angles.
The lattice is what gives ionic compounds their characteristic rigidity. Because every ion is locked in place by the pull of its neighbors, the structure resists deformation. It also means that when you write the formula for an ionic compound, like NaCl, you’re not describing a single molecule. You’re describing the simplest whole-number ratio of ions in that lattice.
Physical Properties
The strong electrical forces in the crystal lattice make ionic compounds hard, rigid, and brittle. They’re solids at room temperature, and they need a lot of heat energy to melt because you have to overcome the attraction between billions of ions all at once. Sodium chloride, for instance, melts at 801 °C. The energy holding ionic lattices together typically ranges from 600 to 4,000 kJ/mol, far higher than the 150 to 400 kJ/mol found in single covalent bonds. Compounds with higher lattice energies tend to be harder, less soluble, and have even higher melting points.
Brittleness is another hallmark. If you hit a salt crystal with a hammer, it shatters rather than bending. That’s because shifting one layer of ions slightly pushes positive ions next to other positive ions, and the sudden repulsion cracks the crystal apart.
Electrical Conductivity
Solid ionic compounds do not conduct electricity. Even though they’re made entirely of charged particles, those particles are locked in fixed positions within the lattice and can’t move toward an electrode. If you put solid salt in an electrical circuit, the light bulb stays dark.
Dissolve that salt in water or melt it, though, and the story changes. Once the lattice breaks apart, the individual ions are free to move. Positive ions drift toward one electrode while negative ions drift toward the other, completing the circuit and allowing current to flow. This is why saltwater conducts electricity but dry salt does not, and it’s one of the simplest ways to confirm that a substance is ionic.
Solubility in Water
Many ionic compounds dissolve in water because water molecules are polar: they have a slightly positive end and a slightly negative end. When an ionic solid meets water, the positive ends of water molecules cluster around the negative ions, and the negative ends cluster around the positive ions. This “hydration” process pulls ions away from the lattice and into solution.
Not all ionic compounds dissolve easily, though. Whether a particular compound dissolves depends on a tug-of-war between two forces: the lattice energy holding the ions together and the hydration energy released when water molecules surround the freed ions. If the lattice energy is too strong relative to the hydration energy, the compound stays mostly intact. Silver chloride, for example, is nearly insoluble in water because the bond between silver and chloride ions is unusually strong. Charge density (how concentrated an ion’s charge is relative to its size) is a useful general predictor of solubility, but exceptions like silver salts show it’s not the whole story.
How Ionic Compounds Are Named
Naming ionic compounds follows a straightforward pattern. The positive ion (cation) comes first, and the negative ion (anion) comes second. The cation keeps the element’s name: Na⁺ is simply “sodium.” The anion gets an “-ide” ending: Cl⁻ becomes “chloride.” So NaCl is sodium chloride, and MgO is magnesium oxide. Unlike molecular compounds, you don’t use Greek prefixes like “di-” or “tri-” to indicate how many of each atom are present. Na₂O is “sodium oxide,” not “disodium oxide.”
When a metal can form ions with different charges (like iron, which can be Fe²⁺ or Fe³⁺), a Roman numeral in parentheses indicates which one you mean. FeCl₂ is iron(II) chloride; FeCl₃ is iron(III) chloride.
Polyatomic Ions
Not every ionic compound is a simple pairing of two elements. Many contain polyatomic ions, which are clusters of atoms that carry a collective charge and act as a single unit. Some of the most common ones include:
- Ammonium (NH₄⁺), the only common polyatomic cation, found in fertilizers
- Nitrate (NO₃⁻), found in potassium nitrate and sodium nitrate
- Sulfate (SO₄²⁻), found in Epsom salt (magnesium sulfate)
- Carbonate (CO₃²⁻), found in limestone and baking soda
- Hydroxide (OH⁻), found in lye (sodium hydroxide)
- Phosphate (PO₄³⁻), essential in biological molecules and detergents
These polyatomic ions follow the same naming and formula rules as simple ions. Calcium carbonate, for instance, is CaCO₃: one calcium ion (Ca²⁺) balancing one carbonate ion (CO₃²⁻). Compounds with polyatomic ions are still classified as ionic because the overall structure is held together by the attraction between charged particles.
Ionic vs. Covalent Compounds
The key difference comes down to what happens to electrons. In ionic compounds, electrons transfer from one atom to another, creating ions. In covalent compounds, atoms share electrons. This leads to very different physical properties. Ionic compounds are typically hard crystalline solids with high melting points, while covalent compounds are more often gases, liquids, or soft solids at room temperature. Ionic compounds conduct electricity when dissolved or melted; most covalent compounds do not conduct electricity in any state.
In practice, the line between ionic and covalent isn’t perfectly sharp. Many bonds have some character of both types. But the distinction is useful because it predicts how a substance will behave: whether it dissolves in water, whether it conducts electricity, how much heat it takes to melt, and how it interacts with other chemicals.