Ionic bonding: what is an ionic bond?

What is an ionic bond?

Atoms are most stable with a full outer shell of electrons, like the noble gases in group 0. Metal atoms have only a few outer electrons, so it is easiest for them to lose them. Non-metal atoms have nearly full outer shells, so it is easiest for them to gain a few more. When a metal reacts with a non-metal, electrons move from one to the other, and both end up with full outer shells.

An atom that loses or gains electrons is no longer neutral. It becomes an ion:

• Metals lose electrons and become positive ions, called cations.
• Non-metals gain electrons and become negative ions, called anions.

The ionic bond is not the electron transfer itself. It is the attraction between the oppositely charged ions that the transfer creates. The AQA GCSE specification describes ionic compounds as held together by "strong electrostatic forces of attraction between oppositely charged ions" that act in all directions.

How an ionic bond forms: sodium chloride

Sodium, in group 1, has the electron arrangement 2,8,1, and chlorine, in group 7, has 2,8,7. Sodium's single outer electron moves to chlorine's outer shell.

Sodium now has one more proton than it has electrons, so its charge is 1+. Chloride has one more electron than protons, so its charge is 1−. Notice the name change: chlorine the element becomes chloride the ion. Non-metal ions end in "ide", so oxygen makes oxide and sulfur makes sulfide. You can add and remove electrons yourself in our Atom Builder to see the charge change.

Dot and cross diagrams for ionic bonding

In a dot and cross diagram, one atom's electrons are dots and the other's are crosses, so you can see where each electron came from. For ionic compounds you draw the ions after the transfer, each in square brackets with its charge outside. Usually only the outer shell is drawn.

Sodium chloride (outer shells shown)

  [ Na ]+      [  x x   ]-
               [ x Cl • ]
               [  x x   ]
               [   x x  ]

Na+ has lost its outer electron, so its outer shell is now the full shell below.
Cl- has 7 crosses (its own) and 1 dot (from sodium): a full shell of 8.

For magnesium oxide, magnesium (2,8,2) gives two electrons to oxygen (2,6), making Mg²⁺ and O²⁻. For magnesium chloride, one magnesium atom gives one electron to each of two chlorine atoms, making one Mg²⁺ and two Cl⁻.

Ion charges from the periodic table

For the main groups, the charge on an ion follows from the group number, because that tells you how many outer electrons the atom has.

Group 4 elements such as carbon rarely form simple ions, and group 0 elements already have full shells. Transition metals such as iron can form more than one ion, so their names include the charge in Roman numerals: iron(II) is Fe²⁺ and iron(III) is Fe³⁺.

Working out the formula of an ionic compound

The charges in an ionic compound always cancel out to zero. Find each ion's charge, then use the smallest number of each ion that adds up to zero.

• Calcium chloride: Ca²⁺ and Cl⁻. One 2+ needs two 1−, so the formula is CaCl₂.
• Potassium oxide: K⁺ and O²⁻. Two 1+ balance one 2−, so K₂O.
• Aluminium oxide: Al³⁺ and O²⁻. The smallest number both 3 and 2 go into is 6: two Al³⁺ make 6+ and three O²⁻ make 6−, so Al₂O₃.
• Magnesium nitride: Mg²⁺ and N³⁻. Three Mg²⁺ make 6+ and two N³⁻ make 6−, so Mg₃N₂.

A quick shortcut is to swap the charge numbers and write them as the small numbers: Al³⁺ and O²⁻ give Al₂O₃. Always simplify afterwards, though: Mg²⁺ and O²⁻ would give Mg₂O₂, which simplifies to MgO.

Ions made of several atoms

Some ions are groups of atoms with an overall charge. Put brackets round them when you need more than one: calcium hydroxide is Ca(OH)₂, and ammonium sulfate is (NH₄)₂SO₄.

Try it: this question gives you two ions and asks how many of each make a neutral compound.

Giant ionic lattice structure

Ions don't pair off one to one like little molecules. Every positive ion attracts every negative ion near it, in all directions. So huge numbers of ions pack together in a regular repeating pattern called a giant ionic lattice. A single grain of salt is one giant lattice.

In sodium chloride, the ions sit in a cubic pattern. Each Na⁺ is surrounded by six Cl⁻ ions, and each Cl⁻ by six Na⁺ ions, as OpenStax Chemistry shows. Because the lattice contains the two ions in equal numbers, the formula NaCl gives the ratio of ions, not a molecule. Our Chemical Bonding course has a 3D model you can turn around.

Every diagram of a lattice has limits. Ball and stick models show gaps between ions that aren't really there, and the sticks suggest bonds between particular pairs of ions, when really each ion attracts all its neighbours. A 2D drawing shows only one layer.

Properties of ionic compounds

Every property comes from the same structure: a giant lattice of charged ions held by strong attractions in all directions.

Bigger charges mean stronger attraction. Sodium chloride (1+ and 1− ions) melts at about 801 °C, while magnesium oxide (2+ and 2− ions) melts at about 2,825 °C, according to figures from the CRC Handbook listed on PubChem. That is why magnesium oxide is used to line furnaces.

Ionic vs covalent bonds

For the sharing side of the story, see our guide to covalent bonding.

Common mistakes

• Calling the electron transfer the bond. The bond is the attraction between the ions that the transfer makes.
• Saying electrons are shared. In ionic bonding they are transferred. Sharing is covalent bonding.
• Talking about molecules of sodium chloride. Ionic compounds form giant lattices, not molecules.
• Saying solid ionic compounds conduct. Only melted or dissolved ones do, because the ions must be free to move.
• Saying electrons carry the current in molten salt. It is the ions that move.
• Leaving out the brackets and charges in dot and cross diagrams of ions.
• Getting the charge sign backwards. Losing negative electrons makes an ion positive.

Read next

Chemistry · Reactions and electrolysisIonic equations: how to write them, with examplesWhat an ionic equation is, how to write one in four steps by removing spectator ions, and examples for neutralisation, precipitation and displacement.