Ionic Bonding: Transferring Electrons Step by Step

Sprinkle table salt on your food and you’re handling one of the most studied substances in chemistry: sodium chloride, the classic ionic compound. Its atoms aren’t sharing electrons. Instead, each sodium atom has handed an electron to a chlorine atom, and the resulting charged particles — ions — cling together in a perfectly ordered crystal. This guide walks through ionic bonding step by step: how ions form, why the process happens at all, and what it means for the properties of ionic compounds.

What is an ionic bond?

An ionic bond is the strong electrostatic attraction between oppositely charged ions. It forms when one or more electrons are transferred from one atom to another — typically from a metal to a non-metal.

• The metal atom loses electrons and becomes a positive ion (a cation).
• The non-metal atom gains electrons and becomes a negative ion (an anion).

For a refresher on ions themselves, see what is an ion?.

Step 1: the metal loses electrons

Take sodium, element 11 (see sodium). Its electron configuration is 2, 8, 1 — one electron in its outer shell, far from the nucleus and well shielded by the inner electrons. That lone electron is relatively easy to remove. The energy needed is the first ionisation energy: 5.14 eV per atom, or about 496 kJ mol⁻¹.

Na → Na⁺ + e⁻ (2,8,1 → 2,8)

The sodium ion has the same electron arrangement as neon. It has 11 protons but only 10 electrons, so it carries a charge of +1.

Step 2: the non-metal gains electrons

Chlorine, element 17 (see chlorine), has the configuration 2, 8, 7 — one electron short of a full outer shell. When a chlorine atom gains an electron, energy is released: its electron affinity is 3.62 eV per atom, or about 349 kJ mol⁻¹.

Cl + e⁻ → Cl⁻ (2,8,7 → 2,8,8)

The chloride ion has the same arrangement as argon: 17 protons and 18 electrons, charge −1.

Step 3: the ions attract

Na⁺ and Cl⁻ attract each other strongly because of their opposite charges. The force follows Coulomb’s law: it’s proportional to the product of the charges and inversely proportional to the square of the distance between them.

F ∝ (q₁ × q₂) / r²

This means:

• Higher charges give stronger attraction. Mg²⁺ and O²⁻ attract each other about four times more strongly than Na⁺ and Cl⁻ at the same distance.
• Smaller ions can get closer together, so they attract more strongly (see ionic radius).

Why does it happen? The energy balance

Here’s a puzzle. Removing an electron from sodium costs 496 kJ mol⁻¹, but adding it to chlorine releases only 349 kJ mol⁻¹. On their own, those two steps cost a net 147 kJ mol⁻¹. So why does sodium react so vigorously with chlorine?

The answer is the third step. When gaseous Na⁺ and Cl⁻ ions come together to form a solid crystal, a huge amount of energy is released — the lattice energy, about 787 kJ mol⁻¹ for sodium chloride. That’s far more than enough to pay for the electron transfer.

In the full reaction, starting from solid sodium and chlorine gas, there are a few more steps (turning solid sodium into gaseous atoms, and breaking Cl–Cl bonds). Chemists account for all of them using an energy cycle called the Born–Haber cycle. The overall result is that forming sodium chloride from its elements releases about 411 kJ per mole — a strongly exothermic reaction (see exothermic vs endothermic).

The key lesson: ionic compounds form not because electron transfer on its own is favourable, but because the attraction between the ions in the lattice releases so much energy.

Working out ion charges from the periodic table

For main-group elements, the charge on an ion is usually predictable from its group:

Group	Outer electrons	Tendency	Typical ion	Examples
1	1	lose 1	+1	Li⁺, Na⁺, K⁺
2	2	lose 2	+2	Mg²⁺, Ca²⁺
13	3	lose 3	+3	Al³⁺
15	5	gain 3	−3	N³⁻, P³⁻
16	6	gain 2	−2	O²⁻, S²⁻
17	7	gain 1	−1	F⁻, Cl⁻, Br⁻

Transition metals often form more than one ion — iron forms Fe²⁺ and Fe³⁺, for example — which is why their compound names include Roman numerals, like iron(III) oxide (see how to name ionic compounds).

Working out formulas

An ionic compound has no overall charge, so the positive and negative charges must balance.

Example 1: magnesium oxide. Mg²⁺ and O²⁻: one of each balances. Formula MgO.

Example 2: calcium chloride. Ca²⁺ and Cl⁻: two chloride ions are needed to balance one calcium ion. Formula CaCl₂.

Example 3: aluminium oxide. Al³⁺ and O²⁻: the lowest common multiple of 3 and 2 is 6, so two Al³⁺ (+6) and three O²⁻ (−6). Formula Al₂O₃.

A shortcut is the “swap and drop” method: write the charge numbers, swap them to become subscripts, and simplify if possible. It works, but always check that the charges balance.

Showing ionic bonding: dot-and-cross diagrams

Dot-and-cross diagrams show where electrons come from. For sodium chloride, sodium’s electron is drawn as a cross and chlorine’s as dots. After transfer, the sodium ion is drawn with an empty outer shell (or its full second shell), and the chloride ion with eight outer electrons (seven dots and one cross), each in square brackets with its charge. See how to draw dot-and-cross diagrams.

The giant ionic lattice

Ionic compounds don’t form separate molecules. Each ion attracts every oppositely charged ion around it in all directions, so ions pack into a regular 3D pattern called a giant ionic lattice. In sodium chloride, each Na⁺ is surrounded by six Cl⁻, and each Cl⁻ by six Na⁺, forming a cubic arrangement. That’s why salt crystals are often little cubes.

Because there are no molecules, the formula NaCl is a ratio, not a molecular formula: there’s one Na⁺ for every Cl⁻.

Properties of ionic compounds

The structure explains the properties:

Property	Explanation
High melting and boiling points	Many strong electrostatic attractions must be overcome. NaCl melts at 801 °C; MgO at about 2,850 °C (higher charges, stronger attraction).
Conduct electricity when molten or dissolved	Ions become free to move and carry charge.
Don’t conduct when solid	Ions are fixed in place.
Often soluble in water	Polar water molecules surround and separate the ions (see hydrogen bonding). Not all ionic compounds dissolve, though — calcium carbonate is almost insoluble.
Hard but brittle	A strong blow shifts one layer of ions; like charges end up next to each other and repel, so the crystal cracks.

Is any bond purely ionic?

No bond is 100 % ionic. Even in sodium chloride, the electron cloud of the chloride ion is slightly distorted by the sodium ion’s positive charge, giving a small amount of covalent character. The larger the difference in electronegativity between the two elements, the more ionic the bond. Caesium fluoride (electronegativities 0.79 and 3.98) is among the most ionic compounds known. See ionic vs covalent bonds.

Common misconceptions

• “Sodium gives its electron to chlorine because it wants a full shell.” The process happens because of the overall energy change, especially the lattice energy.
• “NaCl is a molecule.” It’s a giant lattice; NaCl is the ratio of ions.
• “An ionic bond is between one pair of ions.” Each ion is attracted to all its neighbours.
• “Solid ionic compounds conduct because they contain ions.” The ions must be free to move.

Key takeaways

• An ionic bond is the electrostatic attraction between oppositely charged ions formed by electron transfer.
• Metals lose electrons to form cations; non-metals gain electrons to form anions. Group number predicts the usual charge.
• Electron transfer alone costs energy; the lattice energy released when ions come together makes the process favourable.
• Ions form giant lattices, giving high melting points, conductivity when molten or dissolved, and brittleness.
• Balance charges to find formulas: MgO, CaCl₂, Al₂O₃.

Next: see how non-metals share electrons in covalent bonding explained.