What you'll learn
Ionic bonding and ionic compounds is the first of the three bonding types in AQA GCSE Chemistry, and it establishes the pattern that the rest of the unit follows: work out the structure, then deduce the properties from it. Ionic bonding occurs between a metal and a non-metal, and it involves the complete transfer of electrons from one atom to the other, producing charged particles that attract each other strongly in every direction. Almost every property of an ionic compound — the high melting point, the brittleness, the peculiar behaviour of conducting only when molten or dissolved — follows from that one arrangement. By the end of this guide you should be able to explain how ions form, predict the charge on an ion from its group, work out the formula of an ionic compound, describe the giant ionic lattice, and explain each property of ionic compounds in terms of structure and bonding.
Key terms and definitions
Ion — an atom or group of atoms with an overall electrical charge, formed by losing or gaining electrons
Cation — a positively charged ion, formed when a metal atom loses electrons
Anion — a negatively charged ion, formed when a non-metal atom gains electrons
Ionic bond — the strong electrostatic force of attraction between oppositely charged ions
Electrostatic attraction — the attraction between opposite electrical charges
Giant ionic lattice — a regular repeating three-dimensional arrangement of oppositely charged ions
Empirical formula — the simplest whole-number ratio of ions in a compound
Polyatomic ion — an ion made of more than one atom carrying an overall charge, such as sulfate or nitrate
Brittle — liable to shatter rather than bend when struck
Electrolyte — a molten or dissolved ionic compound that conducts electricity
Core concepts
How ions form
Atoms react in order to achieve the stable electronic structure of a noble gas, meaning a full outer shell.
Metals have few outer electrons, so the easiest route to a full outer shell is to lose them. A sodium atom, with electronic structure 2,8,1, loses its single outer electron to leave 2,8 — the structure of neon. Having lost a negative electron while keeping all its protons, it now carries a charge of plus 1.
Non-metals have nearly full outer shells, so the easiest route is to gain electrons. A chlorine atom, with structure 2,8,7, gains one electron to reach 2,8,8 — the structure of argon — and carries a charge of minus 1.
When a metal reacts with a non-metal, the electrons lost by the metal are precisely the electrons gained by the non-metal. Electrons are transferred, not shared, and this is the essential difference from covalent bonding.
Predicting ionic charge from the group
The charge on a simple ion follows directly from its position in the periodic table, which makes prediction reliable.
Group 1 metals lose one electron to form ions of charge plus 1. Group 2 metals lose two, forming plus 2. Group 3 metals lose three, forming plus 3.
Group 5 non-metals gain three electrons to form ions of charge minus 3. Group 6 gains two, forming minus 2. Group 7 gains one, forming minus 1.
Group 0 forms no ions at all, because the outer shell is already full.
A useful pattern: metals form positive ions with a charge equal to the group number, and non-metals form negative ions with a charge equal to eight minus the group number.
Transition metals are the exception, since many form ions of more than one charge — iron, for example, forms both plus 2 and plus 3 ions. Their charges must be given in the question or indicated by a name such as iron(III).
Working out formulae
An ionic compound has no overall charge, so the positive and negative charges must cancel exactly. The formula is the simplest whole-number ratio of ions that achieves this.
Sodium chloride combines ions of plus 1 and minus 1, so one of each is needed and the formula has a one-to-one ratio.
Magnesium chloride combines plus 2 with minus 1, so two chloride ions are needed for each magnesium ion.
Aluminium oxide combines plus 3 with minus 2. The lowest common multiple of 3 and 2 is 6, so two aluminium ions supply plus 6 and three oxide ions supply minus 6, giving a two-to-three ratio.
The polyatomic ions worth memorising are hydroxide with a charge of minus 1, nitrate minus 1, carbonate minus 2, sulfate minus 2 and ammonium plus 1. When more than one polyatomic ion is needed, it is placed in brackets with the number outside, so calcium hydroxide requires two hydroxide ions written in brackets.
The giant ionic lattice
Ionic compounds do not form molecules. Instead the oppositely charged ions arrange themselves into a giant lattice: a regular repeating three-dimensional structure extending throughout the substance.
In sodium chloride, each sodium ion is surrounded by six chloride ions and each chloride ion by six sodium ions. There is no single pair that constitutes a molecule, which is why the formula gives a ratio rather than a molecular composition.
The ionic bond itself is the strong electrostatic force of attraction between the oppositely charged ions, and crucially it acts in all directions throughout the lattice. Describing it in exactly those words — strong electrostatic attraction between oppositely charged ions, acting in all directions — is what secures the mark.
Explaining the properties
High melting and boiling points follow directly. There are very many strong electrostatic forces of attraction throughout the lattice, and melting requires enough energy to overcome them. Sodium chloride melts at 801 degrees Celsius.
A related and frequently examined point is that compounds with more highly charged ions have higher melting points. Magnesium oxide, with ions of plus 2 and minus 2, melts at about 2,850 degrees Celsius, far above sodium chloride, because the attraction between more highly charged ions is stronger.
Electrical conductivity is the property students most often get wrong, and the reasoning has two halves.
When solid, an ionic compound does not conduct. The ions are charged, but they are held in fixed positions in the lattice and cannot move, so no charge can flow.
When molten or dissolved in water, it does conduct. The lattice has been broken apart, so the ions are free to move and can carry charge through the liquid. A compound in this state is called an electrolyte, which is the link to the electrolysis topic.
Brittleness is explained by the lattice arrangement. If a force displaces one layer of ions relative to the next, ions of the same charge come to lie alongside one another. Like charges repel, so the layers push apart and the crystal shatters rather than bending. This contrasts directly with metals, whose layers slide without any such repulsion.
Solubility in water is common, because water molecules are able to separate the ions and surround them, though not all ionic compounds dissolve.
Comparing the three structure types
Because bonding questions frequently ask you to classify an unknown substance, it is worth holding the comparison ready.
Ionic compounds have high melting points and conduct only when molten or dissolved.
Simple molecular substances have low melting points and never conduct.
Giant covalent structures have very high melting points and do not conduct even when molten, with graphite as the exception.
Metals have high melting points and conduct in the solid state.
Conductivity in the solid and molten states together distinguishes all four.
Worked examples
Example 1: Explaining ion formation (4 marks)
Describe what happens to the electrons when magnesium reacts with oxygen, and state the charges on the ions formed.
Magnesium has two electrons in its outer shell. It loses both, achieving the full outer shell of neon, and because it has lost two negative electrons while keeping all its protons it becomes an ion with a charge of plus 2.
Oxygen has six electrons in its outer shell and gains two, achieving the full outer shell of neon and becoming an ion with a charge of minus 2.
The two electrons lost by each magnesium atom are the two gained by each oxygen atom, so the electrons are transferred from the metal to the non-metal. The oppositely charged ions then attract one another strongly.
Example 2: Deducing a formula (3 marks)
Deduce the formula of calcium nitrate.
Calcium is in Group 2, so it forms an ion with a charge of plus 2. The nitrate ion is polyatomic with a charge of minus 1.
For the compound to have no overall charge, the plus 2 must be balanced by two nitrate ions supplying minus 1 each. The formula therefore contains one calcium ion and two nitrate ions, with the nitrate written in brackets and a subscript 2 outside, since the whole nitrate group is duplicated rather than a single atom within it.
Example 3: Explaining conductivity (4 marks)
Explain why magnesium chloride does not conduct electricity as a solid but does conduct when dissolved in water.
In the solid, the magnesium ions and chloride ions are held in fixed positions within the giant ionic lattice by strong electrostatic forces of attraction. Although the ions are charged, they are not free to move, so no charge can flow and the solid does not conduct.
When the compound dissolves, the lattice is broken apart and the ions become free to move through the solution. Because the ions are charged and mobile, they can carry charge from one place to another, so the solution conducts electricity.
Common mistakes and how to avoid them
The most frequent error is stating that ionic compounds conduct because they contain ions. Containing ions is not sufficient; the ions must be free to move, which is why the solid does not conduct.
Students often describe electrons as being shared in ionic bonding. Sharing is covalent bonding; ionic bonding involves complete transfer.
Another common slip is calling the ionic bond a bond between atoms. It is between ions, and describing it as electrostatic attraction between oppositely charged ions is what the mark scheme requires.
Many candidates give the formula of an ionic compound with the charges still shown, or fail to bracket a polyatomic ion that appears more than once. Brackets are needed whenever the whole group is multiplied.
Finally, in melting point explanations, answers frequently say the bonds are strong without saying there are many of them throughout the lattice. Both the strength and the number matter.
Exam technique for "Ionic bonding and ionic compounds"
Identify the elements before anything else. Metal plus non-metal means ionic, and that single observation determines every answer that follows.
For formula questions, write the charge above each ion first, then find the ratio that cancels them. Doing this on paper is faster than attempting it mentally and leaves working the examiner can credit.
For any property question, use the three-step chain: name the structure, describe the forces, then link to the property. The chain works for every substance in the unit.
When conductivity is mentioned, check the state given in the question before answering. Solid, molten and aqueous give different answers for the same compound, and the state is always stated for a reason.
Quick revision summary
Ionic bonding occurs between metals and non-metals by complete transfer of electrons, giving each a full outer shell. Metals lose electrons to form positive ions with a charge equal to the group number; non-metals gain electrons to form negative ions with a charge of eight minus the group number; transition metals may form more than one charge. Formulae are the simplest ratio that makes the charges cancel, with brackets around any polyatomic ion that is multiplied. The ions form a giant lattice held by strong electrostatic attraction acting in all directions. Melting and boiling points are high because many strong forces must be overcome, and are higher still for more highly charged ions. Solids do not conduct because the ions are fixed, but molten or dissolved compounds do because the ions become free to move. Ionic compounds are brittle because displacing a layer brings like charges together, which repel and split the crystal, and many dissolve in water.