C5: Ionic Bonding
Ions, cations, anions, dot and cross diagrams, giant ionic structures and their properties
📋 Key Definitions
Ion: An atom or group of atoms that has lost or gained one or more electrons, giving it an overall electric charge.
Cation: A positively charged ion formed when an atom loses one or more electrons. Metals typically form cations.
Anion: A negatively charged ion formed when an atom gains one or more electrons. Non-metals typically form anions.
Ionic bond: The strong electrostatic attraction between oppositely charged ions, formed by the transfer of electrons from a metal atom to a non-metal atom.
Giant ionic lattice: A regular, repeating three-dimensional arrangement of positive and negative ions held together by strong ionic bonds.
⚛️ Forming Ions
When metals react with non-metals, electrons are transferred from the metal to the non-metal. The metal becomes a positive ion (cation) and the non-metal becomes a negative ion (anion). This happens because metals have low electronegativity and tend to lose electrons, while non-metals have high electronegativity and tend to gain electrons.
Ionic charge rule: The charge on an ion is equal to the number of electrons lost or gained. Group 1 elements form 1+ ions, Group 2 form 2+ ions, Group 6 form 2− ions, Group 7 form 1− ions.
Common Ions
| Element | Group | Electrons lost/gained | Ion formed |
|---|---|---|---|
| Lithium (Li) | 1 | Loses 1 | Li⁺ |
| Sodium (Na) | 1 | Loses 1 | Na⁺ |
| Potassium (K) | 1 | Loses 1 | K⁺ |
| Magnesium (Mg) | 2 | Loses 2 | Mg²⁺ |
| Calcium (Ca) | 2 | Loses 2 | Ca²⁺ |
| Aluminium (Al) | 3 | Loses 3 | Al³⁺ |
| Oxygen (O) | 6 | Gains 2 | O²⁻ |
| Chlorine (Cl) | 7 | Gains 1 | Cl⁻ |
| Fluorine (F) | 7 | Gains 1 | F⁻ |
Polyatomic ions: Some ions contain more than one atom. Examples include hydroxide (OH⁻), nitrate (NO₃⁻), sulfate (SO₄²⁻) and carbonate (CO₃²⁻).
✏️ Dot and Cross Diagrams
Dot and cross diagrams show how electrons are transferred in ionic bonding. One element's outer electrons are shown as dots, the other as crosses. Only the outer shell electrons need to be shown.
Worked Example - Sodium Chloride (NaCl)
Sodium (2,8,1) transfers its one outer electron to chlorine (2,8,7).
Before transfer:
Na: [2,8,1] · Cl: [2,8,7] with 7 outer electrons shown as ×
After transfer:
Na⁺: [2,8]⁺ (lost its 1 outer dot electron, now has full outer shell)
Cl⁻: [2,8,8]⁻ (gained 1 dot electron from Na, now has full outer shell)
Na · → Na⁺ + e⁻ Cl × + e⁻ → Cl⁻
The dot and cross diagram shows Na⁺ with empty outer shell brackets [2,8]⁺ and Cl⁻ with 8 electrons in outer shell (7 × and 1 ·) [2,8,8]⁻
Na⁺ [2,8]⁺ Cl⁻ [2,8,8]⁻
The electrostatic attraction between Na⁺ and Cl⁻ forms the ionic bond.
Worked Example - Magnesium Oxide (MgO)
Magnesium (2,8,2) transfers two outer electrons to oxygen (2,6).
Before transfer:
Mg: [2,8,2] · · O: [2,6] with 6 outer electrons shown as ×
After transfer:
Mg²⁺: [2,8]²⁺ (lost 2 dot electrons)
O²⁻: [2,8]²⁻ (gained 2 dot electrons from Mg, now has 8 outer electrons: 6 × + 2 ·)
Mg²⁺ [2,8]²⁺ O²⁻ [2,8]²⁻
Mg transfers 2 electrons because it needs to lose 2 to achieve a full outer shell. O needs to gain 2 to achieve a full outer shell.
Worked Example - Sodium Oxide (Na₂O)
Two sodium atoms each transfer one electron to one oxygen atom.
Before transfer:
2 × Na: [2,8,1] · O: [2,6] with 6 outer electrons shown as ×
After transfer:
2 × Na⁺: [2,8]⁺ (each Na lost 1 dot electron)
O²⁻: [2,8]²⁻ (gained 2 dot electrons from the two Na atoms, now has 8 outer electrons: 6 × + 2 ·)
2Na⁺ [2,8]⁺ O²⁻ [2,8]²⁻
The formula Na₂O comes from balancing the charges: 2 × (+1) + (−2) = 0
🏗️ Giant Ionic Structures
Ionic compounds form giant ionic lattices. The ions are arranged in a regular repeating pattern where each positive ion is surrounded by negative ions and vice versa. The strong electrostatic forces of attraction between the oppositely charged ions act in all directions throughout the lattice.
Key feature: In a giant ionic lattice, each ion is attracted to all the oppositely charged ions around it. This is why ionic compounds have high melting points - a large amount of energy is needed to overcome all the strong ionic bonds.
Examiners often ask why ionic compounds have high melting points. You must state that strong electrostatic forces of attraction between oppositely charged ions require large amounts of energy to overcome. Both parts are needed for full marks.
📊 Properties of Ionic Compounds
| Property | Explanation |
|---|---|
| High melting and boiling points | Strong electrostatic forces between oppositely charged ions in the lattice require large amounts of energy to overcome |
| Do not conduct electricity when solid | Ions are held in fixed positions in the lattice and cannot move, so they cannot carry charge |
| Conduct electricity when molten or dissolved | Ions are free to move and can carry charge through the liquid or solution |
| Brittle | If the lattice is shifted, ions of the same charge are brought next to each other and the lattice repels and splits |
| Soluble in water (often) | Water molecules surround and separate the ions from the lattice, allowing them to dissolve |
A very common exam question: "Explain why sodium chloride conducts electricity when molten but not when solid." The key is mentioning that ions are free to move when molten but held in fixed positions in the solid lattice.
Conductivity requirement: For any substance to conduct electricity, it must contain charged particles that are free to move. In ionic compounds, the charged particles are the ions. When solid, the ions cannot move. When molten or in solution, the ions are mobile and can carry charge.
🔬 Working Out Ionic Formulae
The overall charge on an ionic compound must be zero. This means the total positive charge must equal the total negative charge. You can use the charges on ions to work out the formula of an ionic compound.
Rule: Positive charges must balance negative charges
Worked Example - Finding Ionic Formulae
Calcium chloride: Ca²⁺ and Cl⁻
Need two Cl⁻ to balance one Ca²⁺: (+2) + 2(−1) = 0
Formula: CaCl₂
Aluminium oxide: Al³⁺ and O²⁻
Need two Al³⁺ and three O²⁻: 2(+3) + 3(−2) = +6 − 6 = 0
Formula: Al₂O₃
Magnesium nitrate: Mg²⁺ and NO₃⁻
Need two NO₃⁻ to balance one Mg²⁺: (+2) + 2(−1) = 0
Formula: Mg(NO₃)₂ (brackets needed around the polyatomic ion)
❓ Practice Questions
Q1: Foundation Explain why ionic compounds have high melting points.
Ionic compounds have high melting points because there are strong electrostatic forces of attraction between the oppositely charged ions in the giant ionic lattice. A large amount of energy is needed to overcome these forces and separate the ions.
Q2: Foundation Explain why solid sodium chloride does not conduct electricity, but molten sodium chloride does.
In solid sodium chloride, the ions are held in fixed positions in the giant ionic lattice and cannot move, so they cannot carry charge. When molten, the ionic lattice breaks down and the ions become free to move. These mobile charged particles can carry charge through the liquid, allowing it to conduct electricity.
Q3: Higher Draw a dot and cross diagram for the ionic bonding in magnesium chloride, MgCl₂. Explain the formation of the ions.
Magnesium (2,8,2) loses its two outer shell electrons to become Mg²⁺ [2,8]²⁺. Each of two chlorine atoms (2,8,7) gains one of these electrons to become two Cl⁻ [2,8,8]⁻ ions. In the dot and cross diagram: Mg shows 2 dot electrons in its outer shell; each Cl shows 7 cross electrons in its outer shell. After transfer: Mg²⁺ has no outer electrons shown (empty outer shell bracket); each Cl⁻ has 7 crosses and 1 dot in its outer shell. The charges balance: (+2) + 2(−1) = 0.
Q4: Higher Write the formula for aluminium sulfate, given that Al forms Al³⁺ ions and the sulfate ion is SO₄²⁻. Show your working.
Al³⁺ and SO₄²⁻. To balance the charges: 2(+3) + 3(−2) = +6 − 6 = 0. So we need 2 Al³⁺ ions and 3 SO₄²⁻ ions. Formula: Al₂(SO₄)₃. Brackets are needed around the sulfate ion because more than one is required and it is a polyatomic ion.
🎯 Exam Tips
When drawing dot and cross diagrams, only show outer shell electrons. Use dots for electrons from one atom and crosses for electrons from the other. Always include the charge on each ion.
Never say "ions share electrons" in ionic bonding. Electrons are transferred, not shared. Sharing electrons is covalent bonding.
When explaining why ionic compounds conduct when dissolved, mention that water molecules separate the ions from the lattice so they become free to move. The key phrase is "ions are free to move and can carry charge."
For questions about brittleness: when a force is applied, the layers of ions can shift so that ions of the same charge align. The like charges repel, causing the crystal to split.
🧮 Maths Skills
Working Out Ionic Charges from Group Numbers
Charge from group number: Group 1 elements lose 1 electron → form 1+ ions. Group 2 elements lose 2 electrons → form 2+ ions. Group 6 elements gain 2 electrons → form 2− ions. Group 7 elements gain 1 electron → form 1− ions. Transition metals have variable charges (given in exam data).
Charge balance rule: The total positive charge must equal the total negative charge in an ionic compound. E.g. Al³⁺ and O²⁻: need 2 × (+3) = +6 and 3 × (−2) = −6. Formula: Al₂O₃.
Formulae of Ionic Compounds by Charge Balance
Worked examples: Ca²⁺ and Cl⁻ → CaCl₂ (need 2 Cl⁻ to balance one Ca²⁺). Mg²⁺ and NO₃⁻ → Mg(NO₃)₂ (brackets around polyatomic ion when more than one is needed). Al³⁺ and SO₄²⁻ → Al₂(SO₄)₃ (cross-multiply charges: 2 × 3 = 6 and 3 × 2 = 6).
Cross-multiply method: Swap the magnitudes of the charges to get the subscripts. Al³⁺ + O²⁻ → Al₂O₃ (swap 3 and 2). Always simplify if possible: Mg²⁺ + O²⁻ → Mg₂O₂ → simplify to MgO.
Calculating relative formula mass: Mᵣ of NaCl = 23 + 35.5 = 58.5. Mᵣ of Al₂O₃ = (2 × 27) + (3 × 16) = 54 + 48 = 102. Show all working in the exam.
Verifying charge balance: For Mg(NO₃)₂: Mg²⁺ gives +2, two NO₃⁻ gives 2 × (−1) = −2. Total = 0 ✓ Always check the total charge equals zero.
❌ Common Misconceptions
Misconceptions About Ionic Bonding and Structure
Wrong: Ionic compounds conduct electricity when they are solids
Correct: Ionic compounds only conduct electricity when molten or dissolved in water. In the solid state, ions are held in fixed positions in the lattice and cannot move. When molten or in solution, ions are free to move and carry charge.
Wrong: A single NaCl molecule exists — sodium chloride is made of NaCl molecules
Correct: Ionic compounds form giant ionic lattices, not individual molecules. Each Na⁺ ion is surrounded by 6 Cl⁻ ions and each Cl⁻ ion is surrounded by 6 Na⁺ ions. The formula NaCl represents the simplest ratio of ions in the lattice, not a molecule.
Wrong: Ionic bonds are weak because ionic compounds dissolve in water
Correct: Ionic bonds are strong — that is why ionic compounds have high melting points. They dissolve in water because water molecules surround and separate the ions (hydration), not because the ionic bonds are weak.
✍️ 6-Mark Extended Question
Question
Explain why ionic compounds have high melting points and conduct electricity when molten but not when solid.
High melting points: In an ionic compound, the ions are arranged in a giant ionic lattice held together by strong electrostatic forces of attraction between the oppositely charged ions. These forces act in all directions throughout the lattice. A very large amount of energy is needed to overcome these strong forces and separate the ions, so ionic compounds have high melting points [2 marks].
Do not conduct when solid: In the solid state, the ions are held in fixed positions within the ionic lattice and cannot move. Since electrical conductivity requires charged particles that are free to move, solid ionic compounds cannot conduct electricity [2 marks].
Conduct when molten: When the ionic compound is heated above its melting point, the ionic lattice breaks down and the ions become free to move. These mobile, charged ions can carry charge through the liquid, allowing the molten ionic compound to conduct electricity [2 marks].
Mark scheme: 2 marks for high melting point (strong electrostatic forces + large energy needed); 2 marks for not conducting when solid (ions in fixed positions, cannot move); 2 marks for conducting when molten (ions free to move, can carry charge). Must mention ions specifically, not just "particles".🔍 AO3: Analyse and Evaluate
Explaining Melting Point Trends in Ionic Compounds
Scenario: A student looks up the melting points of three ionic compounds: NaF (993°C), NaCl (801°C), NaBr (747°C). All three contain Na⁺ ions but different halide ions. Explain the trend in melting points.
Analysis: The melting point decreases from NaF to NaCl to NaBr. All three compounds contain Na⁺ ions and singly-charged halide ions (F⁻, Cl⁻, Br⁻), so the charges are the same (+1 and −1). The difference is the size of the halide ion: F⁻ is the smallest, Cl⁻ is larger, Br⁻ is the largest. Smaller ions can pack closer together in the lattice, so the distance between the centre of the Na⁺ and the halide ion is smaller in NaF. The electrostatic attraction is stronger when the ions are closer together, so more energy is needed to overcome the attraction, giving a higher melting point.
Evaluation: The trend is consistent with the ionic bonding model — stronger electrostatic attraction (smaller ions, same charge) leads to higher melting points. The student could also compare NaCl (801°C) with MgO (2852°C) — here both ions are smaller AND the charges are higher (2+ and 2−), giving much stronger attraction and a much higher melting point. Charge has a greater effect than ionic size on melting point.
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