What you'll learn
This topic covers how metals and non-metals form ionic compounds through electron transfer, creating charged particles called ions. You'll learn to predict ion formation using the periodic table, draw dot-and-cross diagrams to represent ionic bonding, and explain why ionic compounds have characteristic properties like high melting points and electrical conductivity when molten or dissolved. Understanding ionic bonding is essential for explaining the behaviour of many common compounds, from sodium chloride (table salt) to calcium carbonate (limestone).
Key terms and definitions
Ion — a charged particle formed when an atom loses or gains electrons to achieve a full outer shell
Ionic bonding — the electrostatic force of attraction between oppositely charged ions
Electrostatic attraction — the force between opposite charges (positive and negative ions)
Giant ionic lattice — a regular three-dimensional arrangement of ions held together by strong ionic bonds in all directions
Cation — a positively charged ion formed when an atom loses electrons (metals form cations)
Anion — a negatively charged ion formed when an atom gains electrons (non-metals form anions)
Electron transfer — the process by which electrons move from metal atoms to non-metal atoms during ionic bond formation
Empirical formula — the simplest whole number ratio of ions in an ionic compound
Core concepts
Formation of ions
Atoms form ions to achieve a stable electronic structure with a full outer shell, similar to the nearest noble gas.
Metal atoms form positive ions (cations):
- Metals in Groups 1, 2 and 3 lose electrons from their outer shell
- Group 1 metals lose 1 electron to form 1+ ions (e.g. Na⁺, K⁺)
- Group 2 metals lose 2 electrons to form 2+ ions (e.g. Mg²⁺, Ca²⁺)
- Group 3 metals lose 3 electrons to form 3+ ions (e.g. Al³⁺)
Non-metal atoms form negative ions (anions):
- Non-metals in Groups 5, 6 and 7 gain electrons into their outer shell
- Group 7 elements gain 1 electron to form 1− ions (e.g. Cl⁻, Br⁻)
- Group 6 elements gain 2 electrons to form 2− ions (e.g. O²⁻, S²⁻)
- Group 5 elements gain 3 electrons to form 3− ions (e.g. N³⁻)
When sodium reacts with chlorine:
- Each sodium atom loses one electron: Na → Na⁺ + e⁻
- Each chlorine atom gains one electron: Cl + e⁻ → Cl⁻
- Both ions now have a full outer shell (8 electrons, or 2 for ions with a single shell)
How ionic bonding works
Ionic bonding occurs between metals and non-metals through electron transfer.
The process:
- Metal atoms transfer electrons to non-metal atoms
- This creates oppositely charged ions
- Strong electrostatic forces of attraction develop between the positive and negative ions
- These forces act in all directions throughout the structure
Key point: Ionic bonding is the electrostatic attraction between oppositely charged ions, not the transfer of electrons itself. The electron transfer creates the ions; the attraction between these ions is the ionic bond.
The strength of ionic bonding depends on:
- The size of the charges on the ions (higher charges = stronger attraction)
- The size of the ions (smaller ions = stronger attraction as charges are closer together)
Representing ionic bonding with dot-and-cross diagrams
Dot-and-cross diagrams show the electronic structure of atoms and ions, using dots and crosses to distinguish electrons from different atoms.
Rules for drawing ionic bonding diagrams:
- Show only outer shell electrons (not all electrons)
- Use dots for electrons from one element, crosses for the other
- Draw square brackets around ions with the charge outside
- Show the transfer of electrons with arrows (for some questions)
For sodium chloride (NaCl):
- Sodium (2,8,1) loses 1 electron to form Na⁺ with structure (2,8)
- Chlorine (2,8,7) gains 1 electron to form Cl⁻ with structure (2,8,8)
- Draw [Na]⁺ and [Cl]⁻ showing 8 electrons around chloride ion
For magnesium oxide (MgO):
- Magnesium (2,8,2) loses 2 electrons to form Mg²⁺ with structure (2,8)
- Oxygen (2,6) gains 2 electrons to form O²⁻ with structure (2,8)
- Two electrons transfer from magnesium to oxygen
- Draw [Mg]²⁺ and [O]²⁻ showing 8 electrons around oxide ion
Important: In compounds like magnesium chloride (MgCl₂), one magnesium atom transfers its 2 outer electrons to two different chlorine atoms. Show one Mg²⁺ ion and two Cl⁻ ions in your diagram.
Structure of ionic compounds
Ionic compounds form a giant ionic lattice structure.
Characteristics of the giant ionic lattice:
- Regular arrangement of alternating positive and negative ions
- Ions held in fixed positions by strong ionic bonds
- Ionic bonds act in all directions throughout the structure
- Very large numbers of ions (millions or billions) in each crystal
- The lattice extends in three dimensions
In sodium chloride:
- Each Na⁺ ion is surrounded by 6 Cl⁻ ions
- Each Cl⁻ ion is surrounded by 6 Na⁺ ions
- The ratio of ions is 1:1, giving the formula NaCl
The empirical formula represents the simplest ratio of ions. For example:
- Sodium chloride: 1 Na⁺ to 1 Cl⁻ = NaCl
- Magnesium chloride: 1 Mg²⁺ to 2 Cl⁻ = MgCl₂
- Calcium oxide: 1 Ca²⁺ to 1 O²⁻ = CaO
- Aluminium oxide: 2 Al³⁺ to 3 O²⁻ = Al₂O₃
To work out formulae, the total positive charge must equal the total negative charge.
Properties of ionic compounds
The giant ionic lattice structure explains the characteristic properties of ionic compounds.
High melting and boiling points:
- Strong electrostatic forces between oppositely charged ions throughout the lattice
- Large amounts of energy needed to overcome these forces
- Breaking ionic bonds requires very high temperatures
- Most ionic compounds are solid at room temperature
Example: Sodium chloride melts at 801°C because energy is needed to overcome the strong ionic bonds holding Na⁺ and Cl⁻ ions together.
Electrical conductivity:
- Solid ionic compounds do NOT conduct electricity because ions are in fixed positions and cannot move
- When melted (molten) or dissolved in water (aqueous solution), ionic compounds DO conduct electricity
- The ions are free to move and carry charge through the liquid
- This movement of charged particles (ions) creates an electric current
Brittleness:
- Ionic crystals are hard but brittle
- A force can cause layers of ions to shift
- Ions of the same charge align and repel each other
- This repulsion causes the crystal to split or shatter
Solubility:
- Many ionic compounds dissolve in water
- Water molecules can attract and surround ions, pulling them from the lattice
- Not all ionic compounds are soluble (this depends on the specific ions involved)
Working out formulae of ionic compounds
The charges on ions must balance in any ionic compound (total positive charge = total negative charge).
Method:
- Write down the ions with their charges
- Work out the ratio needed to balance charges
- Write the formula without charges
Examples:
Sodium fluoride: Na⁺ and F⁻
- 1+ and 1− already balance
- Formula: NaF
Calcium bromide: Ca²⁺ and Br⁻
- Need 2× Br⁻ (2−) to balance one Ca²⁺ (2+)
- Formula: CaBr₂
Aluminium oxide: Al³⁺ and O²⁻
- Need 2× Al³⁺ (6+) and 3× O²⁻ (6−) to balance
- Formula: Al₂O₃
Magnesium nitrate: Mg²⁺ and NO₃⁻
- Nitrate is a compound ion (polyatomic ion)
- Need 2× NO₃⁻ (2−) to balance one Mg²⁺ (2+)
- Formula: Mg(NO₃)₂
- Use brackets when you need more than one polyatomic ion
Common polyatomic ions to know:
- Hydroxide: OH⁻
- Nitrate: NO₃⁻
- Carbonate: CO₃²⁻
- Sulfate: SO₄²⁻
- Ammonium: NH₄⁺
Worked examples
Example 1: Describe how ionic bonding occurs in magnesium oxide (4 marks)
Model answer:
- Magnesium atoms transfer/lose two electrons (1 mark)
- Oxygen atoms gain two electrons (1 mark)
- This forms Mg²⁺ ions and O²⁻ ions (1 mark)
- Strong electrostatic attraction/force between oppositely charged ions (1 mark)
Mark scheme notes:
- Must mention electron transfer (not sharing)
- Must state that ions are oppositely charged
- Must use term "electrostatic attraction" or similar
Example 2: Draw a dot-and-cross diagram to show the bonding in calcium fluoride (3 marks)
Model answer: A diagram showing:
- [Ca]²⁺ with no outer electrons (1 mark)
- Two [F]⁻ ions, each with 8 outer electrons in a complete shell (1 mark)
- Electrons shown as dots and crosses correctly distributed (1 mark)
Common errors to avoid:
- Forgetting square brackets and charges on ions
- Showing wrong number of electrons transferred
- Not showing the 2:1 ratio (need 2 fluoride ions for each calcium ion)
Example 3: Explain why sodium chloride has a high melting point but conducts electricity only when molten or dissolved (4 marks)
Model answer:
- Giant ionic lattice with strong electrostatic forces between oppositely charged ions in all directions (1 mark)
- High melting point because large amounts of energy are needed to overcome these strong ionic bonds (1 mark)
- Solid does not conduct because ions are in fixed positions and cannot move (1 mark)
- When molten or dissolved, ions are free to move and carry charge (1 mark)
Examiner tip: Always link structure to properties. Explain the structure first, then connect it to the property.
Common mistakes and how to avoid them
Confusing ionic and covalent bonding: Remember ionic bonding involves transfer of electrons from metals to non-metals, creating oppositely charged ions. Covalent bonding involves sharing electrons between non-metals. Check whether both elements are non-metals (covalent) or a metal with a non-metal (ionic).
Forgetting that ionic bonding is the attraction, not the transfer: Electron transfer creates the ions, but ionic bonding is the electrostatic force of attraction between these ions. Don't write "ionic bonding is when electrons are transferred" without mentioning the attraction.
Stating ionic compounds conduct electricity without specifying state: Solid ionic compounds do NOT conduct electricity. Always state that they conduct when molten or dissolved in water because the ions are then free to move.
Drawing incorrect dot-and-cross diagrams: Only show outer shell electrons, not all electrons. Make sure to draw square brackets around ions with the charge outside. Remember that the number of electrons around each ion should give them a full outer shell (usually 8, or 2 for small ions).
Getting the ratio wrong in compound formulae: The total positive and negative charges must balance. For magnesium chloride, you need two Cl⁻ ions (total 2−) to balance one Mg²⁺ ion (2+), giving MgCl₂, not MgCl.
Not using brackets for polyatomic ions: When you need more than one polyatomic ion like sulfate or nitrate, use brackets: Al₂(SO₄)₃ not Al₂SO₄₃. The bracket shows that the entire polyatomic ion is multiplied.
Exam technique for ionic bonding and ionic compounds
Command word awareness: "Describe" means state what happens without detailed explanation. "Explain" requires you to give reasons using scientific principles. For ionic bonding, "explain" questions need you to link structure (giant ionic lattice, strong electrostatic forces) to properties (high melting point, electrical conductivity).
Structure your answers for properties: Follow a clear pattern: (1) describe the structure, (2) state the nature of the bonding/forces, (3) connect this to the property, (4) add specific detail. For example, for electrical conductivity: "Giant ionic lattice → ions in fixed positions in solid → cannot move → cannot conduct. When molten, ions free to move → carry charge → conducts electricity."
Answer marks per point: In AQA Chemistry, most mark schemes award 1 mark per valid scientific point. A 4-mark question on ionic bonding typically requires four distinct points. Don't repeat the same idea in different words—make sure each sentence adds new information.
Use correct terminology consistently: Use "electrostatic attraction/force" not just "attraction", specify "ions" not "particles" or "atoms", distinguish between "molten" (melted) and "aqueous" (dissolved in water). Precision in terminology gains marks.
Quick revision summary
Ionic bonding occurs when metals transfer electrons to non-metals, forming oppositely charged ions held together by strong electrostatic forces. These ions arrange in a giant ionic lattice with ionic bonds acting in all directions. Ionic compounds have high melting points due to strong ionic bonds requiring large amounts of energy to break. They conduct electricity only when molten or dissolved because ions must be free to move to carry charge. Use dot-and-cross diagrams to represent electron transfer, showing ions with square brackets and charges. Balance charges when working out formulae: total positive charge equals total negative charge.