Ionic Bonding & Ionic Compounds
What Is Ionic Bonding?
When metals react with non-metals, atoms transfer electrons from one to the other. The metal atom loses electrons and becomes a positively charged ion called a cation. The non-metal atom gains those electrons and becomes a negatively charged ion called an anion. The strong electrostatic attraction between these oppositely charged ions is an ionic bond.
An ion is simply an atom (or group of atoms) that has gained or lost electrons, giving it a net charge. In a neutral atom the number of protons equals the number of electrons. Once electrons are transferred, this balance is broken and the particle carries a charge.
Ionic bonds form by the transfer of electrons from a metal atom to a non-metal atom. The metal loses electrons to form a positive ion (cation); the non-metal gains electrons to form a negative ion (anion). Both ions end up with a full outer shell of electrons, giving them a stable electronic configuration.
Ion Formation — Groups 1, 2, 6 and 7
You need to predict which ions elements form based on their position in the periodic table. The Edexcel specification limits ionic compound formation to elements in Groups 1, 2, 6 and 7:
Group 1 elements (e.g. sodium, potassium) lose 1 electron to form ions with a +1 charge — for example, \(\text{Na}^+\) and \(\text{K}^+\). Group 2 elements (e.g. magnesium, calcium) lose 2 electrons to form ions with a +2 charge — for example, \(\text{Mg}^{2+}\) and \(\text{Ca}^{2+}\).
Group 7 elements (e.g. chlorine, bromine) gain 1 electron to form ions with a −1 charge — for example, \(\text{Cl}^-\) and \(\text{Br}^-\). Group 6 elements (e.g. oxygen, sulfur) gain 2 electrons to form ions with a −2 charge — for example, \(\text{O}^{2-}\) and \(\text{S}^{2-}\).
Dot-and-cross diagrams show how electrons are transferred during ionic bonding. The electrons from the metal atom are drawn as crosses and the non-metal’s electrons as dots (or vice versa). After transfer, you draw square brackets around each ion and write its charge. For sodium chloride: sodium (2, 8, 1) loses its outer electron (shown as a cross) to chlorine (2, 8, 7), which gains it. Both ions achieve a full outer shell — \(\text{Na}^+\) (2, 8) and \(\text{Cl}^-\) (2, 8, 8).
For magnesium oxide: magnesium (2, 8, 2) loses two outer electrons, both transferring to oxygen (2, 6), giving \(\text{Mg}^{2+}\) (2, 8) and \(\text{O}^{2-}\) (2, 8).
Calculating Protons, Neutrons and Electrons in Ions
When an atom forms an ion, the number of protons stays the same but the number of electrons changes. A sodium atom has 11 protons and 11 electrons. When it forms \(\text{Na}^+\), it still has 11 protons but only 10 electrons. A chlorine atom has 17 protons and 17 electrons; as \(\text{Cl}^-\) it has 17 protons and 18 electrons.
To find the electrons in any simple ion: start with the atomic number (number of electrons in a neutral atom), then subtract the charge for positive ions or add the charge for negative ions.
Naming Ionic Compounds
Ionic compounds are named using specific rules. When a compound contains just two elements, its name ends in –ide. For example, sodium and chlorine form sodium chloride; magnesium and oxygen form magnesium oxide. When a compound contains a metal, a non-metal and oxygen, its name typically ends in –ate. For example, copper sulfate contains copper, sulfur and oxygen; calcium carbonate contains calcium, carbon and oxygen.
Names ending in –ide tell you the compound contains only two elements (e.g. sodium chloride = sodium + chlorine). Names ending in –ate tell you the compound also contains oxygen (e.g. sodium sulfate = sodium + sulfur + oxygen).
Deducing Formulae of Ionic Compounds
To work out the formula of an ionic compound, you need to balance the charges so they cancel out to zero. For sodium chloride: \(\text{Na}^+\) has a +1 charge and \(\text{Cl}^-\) has a −1 charge, so one of each gives NaCl. For magnesium chloride: \(\text{Mg}^{2+}\) needs two \(\text{Cl}^-\) ions to balance, giving \(\text{MgCl}_2\).
Positive ions: \(\text{Na}^+\), \(\text{K}^+\), \(\text{Mg}^{2+}\), \(\text{Ca}^{2+}\). Negative ions: chloride \(\text{Cl}^-\), bromide \(\text{Br}^-\), oxide \(\text{O}^{2-}\), sulfide \(\text{S}^{2-}\), hydroxide \(\text{OH}^-\), nitrate \(\text{NO}_3^-\), carbonate \(\text{CO}_3^{2-}\), sulfate \(\text{SO}_4^{2-}\). To deduce a formula, balance the total positive charge against the total negative charge so they sum to zero.
The Ionic Lattice
Ionic compounds do not exist as individual pairs of ions. Instead, billions of ions arrange themselves into a regular, repeating three-dimensional structure called a lattice. Each positive ion is surrounded by negative ions, and each negative ion is surrounded by positive ions. The strong electrostatic forces of attraction between these oppositely charged ions act in all directions throughout the lattice, holding the structure firmly together.
Ionic compounds form giant lattice structures. The strong electrostatic forces between oppositely charged ions require a large amount of energy to overcome, which is why ionic compounds have high melting and boiling points.
Properties of Ionic Compounds
Ionic compounds share several distinctive properties, all explained by their lattice structure:
High melting and boiling points — the strong ionic bonds throughout the lattice need a lot of energy to break. Conductivity when solid — ionic compounds do not conduct electricity as solids because the ions are locked in fixed positions and cannot move. Conductivity when molten or dissolved — when melted or dissolved in water, the ions become free to move and can carry an electrical current, so the compound conducts.
When an ionic compound dissolves in water, the lattice breaks apart and the individual ions become surrounded by water molecules. These free ions can now move through the solution towards electrodes. Positive ions (cations) move to the negative electrode and negative ions (anions) move to the positive electrode, carrying the charge and completing the circuit. This is why ionic compounds only conduct when molten or in solution — never as a solid.