Ionic Bonding

Every chemical bond falls into one of three types, and each depends on how the outer-shell electrons behave. Knowing which type applies comes down to which kinds of element are involved.

• Ionic bonding occurs between a metal and a non-metal. Electrons are transferred, forming charged ions that attract each other.
• Covalent bonding occurs between non-metals. Electrons are shared in pairs between atoms.
• Metallic bonding occurs in metallic elements and alloys. Atoms share a pool of delocalised electrons.

Bond type	Between	What happens to electrons
Ionic	Metal + non-metal	Transferred from metal to non-metal
Covalent	Non-metal + non-metal	Shared in pairs
Metallic	Metal(s) / alloy	Delocalised and shared across the structure

All three bonds are electrostatic: they involve the attraction between positive and negative charges. In ionic bonding those charges are whole ions.

This lesson covers ionic bonding in full: the transfer of electrons between a metal and a non-metal.

When a metal reacts with a non-metal, the metal atom loses its outer-shell electrons and the non-metal atom gains them. Both atoms end up with a full outer shell, matching the stable electronic structure of a noble gas.

A metal atom that loses electrons has more protons than electrons, so it becomes a positive ion (a cation). A non-metal atom that gains electrons has more electrons than protons, so it becomes a negative ion (an anion).

Take sodium (electronic structure 2,8,1) reacting with chlorine (2,8,7):

• Sodium loses its single outer electron: it becomes ${\text{Na}}^{+}$ with the structure 2,8 (like neon).
• Chlorine gains that one electron: it becomes ${\text{Cl}}^{-}$ with the structure 2,8,8 (like argon).

$\text{Na}\to {\text{Na}}^{+}+{\text{e}}^{-}$ $\text{Cl}+{\text{e}}^{-}\to {\text{Cl}}^{-}$

The electron leaves sodium and joins chlorine. Nothing is destroyed: the total number of electrons is conserved, which is why the charges balance.

You are expected to work out the charge on an ion directly from its group in the periodic table, for Groups 1, 2, 6 and 7. The rule follows from how many electrons an atom must lose or gain to reach a full outer shell.

Group	Outer electrons	Electrons lost or gained	Ion charge	Examples
1	1	lose 1	1+	Na⁺, K⁺, Li⁺
2	2	lose 2	2+	Mg²⁺, Ca²⁺
6	6	gain 2	2−	O²⁻, S²⁻
7	7	gain 1	1−	Cl⁻, Br⁻

Metals in Groups 1 and 2 form positive ions by losing electrons. Non-metals in Groups 6 and 7 form negative ions by gaining electrons. Each resulting ion has the electronic structure of a noble gas.

Worked example. What is the charge on a calcium ion and an oxide ion? Calcium is in Group 2, so it loses 2 electrons to form ${\text{Ca}}^{2+}$. Oxygen is in Group 6, so it gains 2 electrons to form ${\text{O}}^{2-}$. The 2+ and 2− charges balance one-to-one, so calcium oxide is $\text{CaO}$.

A dot-and-cross diagram shows where each electron came from. One atom's electrons are drawn as dots, the other's as crosses. For ions you draw the transferred electron ending up on the non-metal, and you write the charge on each ion outside square brackets.

The diagram below shows sodium transferring its single outer electron to chlorine. The lone outer electron (drawn as a cross) leaves sodium; chlorine's outer shell (dots) gains it, completing an octet.

For magnesium oxide, magnesium (Group 2) transfers two electrons to oxygen (Group 6):

$\text{Mg}\to {\text{Mg}}^{2+}+2{\text{e}}^{-}\text{O}+2{\text{e}}^{-}\to {\text{O}}^{2-}$

Written as a diagram, both of magnesium's outer electrons move into oxygen's outer shell, giving $[\text{Mg}{]}^{2+}$ and $[\text{O}{]}^{2-}$. The formula is MgO.

In the exam, show the outer shells only, write the charge on each ion, and put square brackets around each ion. Empty inner shells are not needed.

An ionic compound is not made of separate molecules. It is a giant ionic lattice: a regular, repeating 3D arrangement of huge numbers of positive and negative ions. The ions are held together by strong electrostatic forces of attraction that act in all directions between oppositely charged ions.

The named example you must know is sodium chloride. In the NaCl lattice, the Na⁺ and Cl⁻ ions alternate in a cubic pattern that repeats in every direction, so each ion is surrounded on all sides by ions of the opposite charge.

There are no individual "NaCl molecules". The formula NaCl gives the ratio of ions (1:1), not the number in one particle.

Because the electrostatic forces act throughout the whole structure and are strong, a great deal of energy is needed to break them apart. That is why ionic compounds have high melting and boiling points, a property covered in the properties lesson. The strength and all-directional nature of the bonding is the point to remember here.

The spec expects you to deduce the empirical formula of an ionic compound from a diagram or model of its lattice. The empirical formula is the simplest whole-number ratio of the ions present.

There are two reliable methods.

Method 1 — from the model. Count the ions shown, then simplify the ratio. If a model shows 8 sodium ions and 8 chloride ions, the ratio is 8:8, which simplifies to 1:1, giving NaCl.

Method 2 — balancing charges. The compound must be electrically neutral, so the positive and negative charges must cancel.

Worked example — magnesium chloride. Magnesium forms ${\text{Mg}}^{2+}$ and chloride is ${\text{Cl}}^{-}$. To balance one 2+ charge you need two 1− charges, so you need one ${\text{Mg}}^{2+}$ and two ${\text{Cl}}^{-}$:

$1\times (2+)+2\times (1-)=0$

The ratio is 1:2, so the formula is ${\text{MgCl}}_2$.

Worked example — calcium fluoride. Calcium is ${\text{Ca}}^{2+}$ (Group 2) and fluoride is ${\text{F}}^{-}$ (Group 7). One 2+ needs two 1−, giving ${\text{CaF}}_2$.

No single diagram shows the whole truth about ionic bonding. The spec asks you to describe the limitations of each representation, so you must know what each one leaves out.

Representation	What it shows well	What it fails to show
Dot-and-cross	Which atom each electron came from; the transfer	The real 3D lattice; ion sizes; the forces
2D diagram	The ratio of ions in one layer	The third dimension; that the pattern repeats
Ball-and-stick	The 3D geometry and repeating arrangement	The sticks imply covalent bonds; ions look too far apart; ignores their real size
3D space-filling	Relative ion sizes and close packing	Hard to see ions behind the front layer; still finite, not truly giant

A common exam phrase: ball-and-stick models show gaps and sticks that do not really exist between ions, and no model can show that the lattice extends indefinitely in all directions.

1. Getting the direction of electron transfer wrong

The metal loses electrons and the non-metal gains them. Metals become positive, non-metals become negative. Reversing this is a frequent slip in dot-and-cross questions.

2. Forgetting the charges on the ions

A dot-and-cross diagram for an ionic compound must show the charge on each ion and square brackets around it. $\text{Na}$ and $\text{Cl}$ without charges describe atoms, not the ions in the compound.

3. Saying ionic compounds are made of molecules

Ionic compounds are giant lattices, not molecules. Do not describe "NaCl molecules". The formula shows the ratio of ions only.

4. Confusing which group gains and which loses

Groups 1 and 2 lose electrons (they have few outer electrons to shed); Groups 6 and 7 gain electrons (they need only one or two to fill the shell). Check the group number before deciding.

5. Not simplifying to the empirical formula

${\text{Mg}}_2{\text{O}}_2$ is not accepted; it simplifies to MgO. Always give the simplest whole-number ratio of ions.