Ionic bonding questions reward a small number of habits used consistently: work out the electron arrangement, decide how many electrons move, write the ion charges, then balance the charges to get the formula. Explanations of properties almost always come back to one phrase, strong electrostatic attraction between oppositely charged ions in a giant lattice, so practise saying it precisely.
This set has 14 questions worth 49 marks in total. They move from recall to explanation to a couple of extended-answer questions at the end. Mark allocations match what you would typically see on a school exam paper, so use them to judge how much detail each answer needs. The electron arrangements you need are for the first 20 elements; if you want to check any of them, the electron configuration tool will show them.
Before you start
- Metals in groups 1, 2 and 3 lose 1, 2 and 3 electrons to form positive ions.
- Non-metals in groups 15, 16 and 17 gain 3, 2 and 1 electrons to form negative ions.
- The total positive charge in a formula must equal the total negative charge.
- In a dot-and-cross diagram, use dots for one atom’s electrons and crosses for the other’s, put square brackets around each ion and write its charge at the top right.
Questions
1. State what is meant by an ionic bond. (2 marks)
2. Give the charge on the ion formed by each of these elements: (a) potassium, (b) calcium, (c) aluminium, (d) sulfur, (e) nitrogen. (5 marks)
3. Write the electron arrangement of (a) a sodium atom and (b) a sodium ion. Explain why the ion is more stable. (3 marks)
4. Describe, in terms of electron transfer, what happens when sodium reacts with chlorine to form sodium chloride. (3 marks)
5. Describe the dot-and-cross diagram for magnesium oxide, showing outer shells only. (3 marks)
6. Work out the formula of each ionic compound: (a) sodium oxide, (b) calcium fluoride, (c) aluminium oxide, (d) magnesium nitride. (4 marks)
7. Describe the dot-and-cross diagram for calcium fluoride and explain why the formula is CaF₂ rather than CaF. (4 marks)
8. A student says: “Sodium chloride is made of NaCl molecules.” Explain why this is wrong. (3 marks)
9. Name two ions from question 6 that have the same electron arrangement as a neon atom, and give that arrangement. (3 marks)
10. Explain why sodium chloride has a high melting point (801 °C). (3 marks)
11. Magnesium oxide melts at about 2,852 °C. Explain why this is much higher than the melting point of sodium chloride. (3 marks)
12. Explain why solid sodium chloride does not conduct electricity, but molten sodium chloride and sodium chloride solution both do. (4 marks)
13. Explain why an ionic crystal shatters when hit with a hammer. (3 marks)
14. The Pauling electronegativity values of some elements are: sodium 0.93, magnesium 1.31, carbon 2.55, chlorine 3.16, oxygen 3.44. (a) Calculate the electronegativity difference for Na–Cl and for Mg–O. (2 marks) (b) Use the values to explain why carbon and chlorine do not form an ionic compound. (2 marks) (c) Suggest which of NaCl and MgO would be classed as more ionic using electronegativity difference alone, and state one reason why melting point does not follow the same order. (2 marks)
Total: 49 marks
Answer key
1. Definition (2 marks)
An ionic bond is the strong electrostatic attraction (1) between oppositely charged ions (1). Examiner note: “a bond between a metal and a non-metal” describes where ionic bonds form, not what they are, and does not earn the marks on its own.
2. Ion charges (5 marks, 1 each)
(a) K⁺ (group 1, loses one electron) (b) Ca²⁺ (group 2, loses two) (c) Al³⁺ (group 13, loses three) (d) S²⁻ (group 16, gains two) (e) N³⁻ (group 15, gains three)
3. Sodium atom and ion (3 marks)
(a) Na atom: 2,8,1 (1) (b) Na⁺ ion: 2,8 (1) The ion has a full outer shell (the same arrangement as the noble gas neon), which is a stable arrangement (1).
4. Forming sodium chloride (3 marks)
- Each sodium atom loses its one outer electron to become Na⁺ (1).
- Each chlorine atom gains one electron to become Cl⁻ (1).
- Both ions now have full outer shells, and the oppositely charged ions attract each other (1).
Working: Na 2,8,1 → Na⁺ 2,8; Cl 2,8,7 → Cl⁻ 2,8,8. One electron moves, so the ratio is 1 : 1.
5. Magnesium oxide diagram (3 marks)
- The magnesium ion is drawn in square brackets with an empty outer shell (or with its full second shell of eight shown) and the charge 2+ (1).
- The oxide ion is drawn in square brackets with eight outer electrons: six dots and two crosses (the crosses being the two electrons gained from magnesium), and the charge 2− (1).
- Two electrons are shown transferring from magnesium to oxygen (1).
Working: Mg 2,8,2 loses two → Mg²⁺ 2,8. O 2,6 gains two → O²⁻ 2,8. Charges +2 and −2 balance, so the formula is MgO. For more on drawing these, see dot-and-cross diagrams.
6. Formulas (4 marks, 1 each)
Balance the charges using the lowest whole numbers.
(a) Na⁺ and O²⁻: two Na⁺ give +2 to balance −2 → Na₂O (b) Ca²⁺ and F⁻: two F⁻ give −2 to balance +2 → CaF₂ (c) Al³⁺ and O²⁻: lowest common multiple of 3 and 2 is 6. Two Al³⁺ = +6, three O²⁻ = −6 → Al₂O₃ (d) Mg²⁺ and N³⁻: lowest common multiple of 2 and 3 is 6. Three Mg²⁺ = +6, two N³⁻ = −6 → Mg₃N₂
Shortcut: the “swap and drop” method (the size of each charge becomes the other ion’s subscript) works here, but always simplify: Mg²⁺ and O²⁻ give Mg₂O₂, which reduces to MgO. The guide on naming ionic compounds explains the naming side.
7. Calcium fluoride (4 marks)
- Calcium (2,8,8,2) must lose two electrons to get a full outer shell (1).
- Each fluorine atom (2,7) can accept only one electron (1).
- So two fluorine atoms are needed for each calcium atom (1).
- Diagram: one [Ca]²⁺ in brackets, and two separate [F]⁻ ions in brackets, each with seven dots and one cross in the outer shell (1).
CaF would leave calcium with one outer electron still to lose, and the charges (+2 and −1) would not balance.
8. “NaCl molecules” (3 marks)
- Sodium chloride has a giant ionic lattice (1).
- Each Na⁺ ion is surrounded by, and attracted to, several Cl⁻ ions (six in the sodium chloride structure), and each Cl⁻ is surrounded by six Na⁺ (1).
- There are no separate pairs; the formula NaCl only gives the ratio of ions, 1 : 1 (1).
9. Neon arrangement (3 marks)
Any two of: Na⁺, O²⁻, F⁻, Al³⁺, Mg²⁺, N³⁻ (1 each, max 2). Arrangement: 2,8 (1). (Ca²⁺ is 2,8,8, the same as argon, so it does not count.)
10. NaCl melting point (3 marks)
- Giant ionic lattice (1).
- Strong electrostatic forces of attraction between oppositely charged ions, acting in all directions (1).
- A large amount of energy is needed to overcome these forces so that the ions can move (1).
Examiner note: do not mention intermolecular forces or molecules. That is an automatic loss of marks in an ionic answer.
11. MgO versus NaCl (3 marks)
- MgO contains Mg²⁺ and O²⁻, with charges of 2 rather than 1 (1).
- Higher charges mean stronger electrostatic attraction between the ions (1).
- So more energy is needed to separate the ions (1).
(Better answers also mention that Mg²⁺ and O²⁻ are smaller than Na⁺ and Cl⁻, so the ions are closer together and attract more strongly. Credit this as an alternative for the second mark.)
12. Conductivity (4 marks)
- Electrical conduction requires charged particles that are free to move (1).
- In the solid, the ions are held in fixed positions in the lattice and can only vibrate (1).
- When molten, the lattice breaks down and the ions can move (1).
- In solution, the ions are separated by water molecules and are free to move (1).
Examiner note: “electrons move” is wrong for an ionic compound. It is the ions that carry the charge.
13. Brittleness (3 marks)
- A force makes one layer of ions shift relative to the next (1).
- Ions with the same charge are brought next to each other (1).
- These repel, pushing the layers apart, so the crystal splits (1).
14. Electronegativity (6 marks)
(a) Na–Cl: 3.16 − 0.93 = 2.23 (1). Mg–O: 3.44 − 1.31 = 2.13 (1).
(b) C–Cl: 3.16 − 2.55 = 0.61 (1). The difference is small, so neither atom pulls the electrons away completely; they share electrons and form covalent bonds instead (1).
(c) NaCl has the larger difference, so it would be classed as slightly more ionic (1). Melting point depends mainly on the size of the ion charges (and ion sizes), not on the difference in electronegativity; MgO has 2+ and 2− ions, so its lattice attraction is much stronger (1).
This is a useful reminder that “more ionic” and “stronger bonding” are not the same thing. The post on electronegativity and bond type explores where the cut-off between ionic and covalent really lies.
How did you do?
- 41–49: secure. Try the harder lattice and energy questions in the ionic bonding explainer and then move on to covalent bonding practice.
- 29–40: good understanding, but check your wording on questions 10–13. Examiners look for “electrostatic attraction”, “oppositely charged ions”, “giant lattice” and “free to move”.
- Below 29: revisit ion charges and formula balancing first. Most later marks depend on them.
Summary
Every ionic answer starts from the same chain: electron arrangement → electrons lost or gained → ion charges → balanced formula. Properties then follow from the giant lattice of oppositely charged ions. High melting points come from strong attraction in all directions; conduction needs ions that are free to move, which happens only when molten or dissolved; brittleness comes from like charges repelling when layers slide. Keep the key phrases exact, never mention molecules or intermolecular forces for an ionic solid, and always double-check that the charges in your formula add up to zero.
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