Ionic Bonding and Ionic Compounds

We have seen that compounds have new properties because their atoms are joined by chemical bonds. But what is a chemical bond? There are three main types — ionic, covalent and metallic — and this lesson covers the first. Ionic bonding happens between metals and non-metals, and it explains why table salt is a hard, high-melting crystal that conducts electricity only when molten or dissolved. The crucial skill in Topic C2 of OCR Gateway Science A is to link the structure of an ionic compound to its properties — and that chain of reasoning, structure → property, runs through the rest of the topic.

By the end of this lesson you should be able to explain how ionic bonds form by electron transfer, work out ion charges from the periodic table, draw dot-and-cross diagrams, describe the giant ionic lattice, and explain the properties of ionic compounds from their structure.

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How Ionic Bonds Form

An ionic bond forms between a metal and a non-metal. The key event is the transfer of electrons from the metal atom to the non-metal atom:

• a metal atom loses its outer-shell electron(s), becoming a positive ion (a cation);
• a non-metal atom gains those electron(s), becoming a negative ion (an anion).

Both atoms do this in order to achieve a full outer shell — the stable electron arrangement of a noble gas. Once the electrons have transferred, the oppositely charged ions are held together by a strong electrostatic attraction: that attraction is the ionic bond.

Take sodium chloride. Sodium (Group 1) has 1 outer electron; chlorine (Group 7) has 7. Sodium transfers its single outer electron to chlorine. Now sodium has a full outer shell and a charge of +1 ($\text{Na}^+$Na+); chlorine has a full outer shell and a charge of −1 ($\text{Cl}^-$Cl−). The two ions attract strongly, forming the ionic compound $\text{NaCl}$NaCl.

Exam Tip: The three words examiners want for ionic bonding are transfer, electrostatic and attraction: electrons are transferred from metal to non-metal, forming oppositely charged ions held by strong electrostatic attraction. Electrons are transferred, never shared (that is covalent bonding).

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Working Out Ion Charges

You can predict the charge of an ion from its group number in the periodic table, because that tells you how many electrons it must lose or gain to reach a full outer shell.

Group	Outer electrons	Loses/gains	Ion charge	Example
1	1	loses 1	+1	$\text{Na}^+$Na+
2	2	loses 2	+2	$\text{Mg}^{2+}$Mg2+
3	3	loses 3	+3	$\text{Al}^{3+}$Al3+
5	5	gains 3	−3	$\text{N}^{3-}$N3−
6	6	gains 2	−2	$\text{O}^{2-}$O2−
7	7	gains 1	−1	$\text{Cl}^-$Cl−

Metals (Groups 1–3) lose electrons to form positive ions; non-metals (Groups 5–7) gain electrons to form negative ions. Notice that the charge equals the number of electrons lost or gained.

Exam Tip: For Groups 1, 2, 3 the ion charge is +1, +2, +3; for Groups 7, 6, 5 it is −1, −2, −3. Learn this and you can predict ion charges and balance ionic formulae instantly.

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Dot-and-Cross Diagrams

A dot-and-cross diagram shows where the electrons go when an ionic bond forms. The electrons from one atom are drawn as dots and from the other as crosses, so you can see clearly that electrons have been transferred. The ions are usually drawn in square brackets with the charge outside.

For sodium chloride, sodium gives up its one outer electron (shown as a dot) to chlorine, which now has eight electrons in its outer shell (seven of its own crosses plus the one transferred dot). Both ions now have full outer shells.

Some further examples, all formed by electron transfer:

• Magnesium oxide ($\text{MgO}$MgO): $\text{Mg}$Mg (Group 2) transfers 2 electrons to $\text{O}$O (Group 6), giving $\text{Mg}^{2+}$Mg2+ and $\text{O}^{2-}$O2− — a $1:1$1:1 ratio.
• Magnesium chloride ($\text{MgCl}_2$MgCl2​): $\text{Mg}^{2+}$Mg2+ needs to lose 2 electrons, but each $\text{Cl}$Cl can only gain 1, so two chlorine atoms are needed — giving $\text{Mg}^{2+}$Mg2+ and two $\text{Cl}^-$Cl−.
• Sodium oxide ($\text{Na}_2\text{O}$Na2​O): each $\text{Na}$Na loses only 1 electron, but $\text{O}$O needs to gain 2, so two sodium atoms are needed — giving two $\text{Na}^+$Na+ and one $\text{O}^{2-}$O2−.

Exam Tip: In a dot-and-cross diagram, use dots for one atom's electrons and crosses for the other's, so the transfer is visible, and draw the ions in square brackets with the charge outside. Show only the outer shell unless asked for all shells.

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The Giant Ionic Lattice

An ionic compound is not made of separate molecules. The electrostatic attraction between the ions acts in all directions, so the ions pack together into a giant ionic lattice — a regular, repeating three-dimensional arrangement of alternating positive and negative ions that extends throughout the whole crystal. Each positive ion is surrounded by negative ions and vice versa.

The formula of an ionic compound (such as $\text{NaCl}$NaCl) gives the simplest ratio of ions in the lattice, not the number in a molecule — because there are no molecules. A single grain of salt contains countless ions in this $1:1$1:1 ratio.

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Properties Explained by the Structure

The giant ionic lattice explains every characteristic property of ionic compounds. This structure → property reasoning is exactly what the exam rewards.

Property	Explanation from structure
High melting and boiling points	There are many strong electrostatic forces between the oppositely charged ions throughout the lattice; a large amount of energy is needed to overcome them
Conduct when molten or dissolved	When melted or dissolved in water, the ions are free to move and can carry charge
Do NOT conduct when solid	In the solid the ions are fixed in the lattice and cannot move, so no charge flows
Often soluble in water	Water can surround and separate the ions
Brittle	A blow can shift the layers so like charges line up and repel, splitting the crystal