Chemical Bonding Exam Questions with Mark-Scheme Answers

Bonding questions tend to lose marks in the same places year after year: a missing charge on a dot-and-cross diagram, “the covalent bonds break” when a molecular substance boils, or a shape named without an angle. This paper is built around those traps. It has 14 questions worth 41 marks, which is roughly a 45-minute sitting.

Work through it under exam conditions first. Then mark your answers against the scheme below, where each [1] is one marking point, just as an examiner would award it. Notes marked Accept and Do not accept show where the line falls.

Electronegativity values (Pauling scale), where you need them: H 2.20, C 2.55, O 3.44, Cl 3.16, K 0.82.

Section A: Ionic, covalent and metallic bonding

1. (a) State what is meant by a covalent bond. [1] (b) Define electronegativity. [1]

2. Magnesium reacts with oxygen to form magnesium oxide. Draw a dot-and-cross diagram for magnesium oxide, showing outer electrons only. Include the charges on the ions. [3]

3. Sodium chloride melts at 801 °C. Solid sodium chloride does not conduct electricity, but molten sodium chloride does. (a) Explain why sodium chloride has a high melting point. [2] (b) Explain the difference in conductivity between solid and molten sodium chloride. [2]

4. Describe the bonding in a metal and use it to explain why metals are malleable. [3]

5. Diamond and graphite are both forms of carbon. (a) Explain why graphite conducts electricity but diamond does not. [2] (b) Explain why graphite is soft and slippery. [2]

Section B: Shapes and polarity

6. Ammonia, NH₃, has a bond angle of about 107°. Name the shape of the molecule and explain why the angle is smaller than 109.5°. [3]

7. Both carbon dioxide and water contain polar bonds. Explain why water is a polar molecule but carbon dioxide is not. [3]

8. Xenon forms the compound XeF₄. State the number of bonding pairs and lone pairs around the xenon atom, name the shape, and give the F–Xe–F bond angle. [3]

9. Ammonia reacts with a hydrogen ion to form the ammonium ion, NH₄⁺. (a) Name the type of bond that forms and explain how it forms. [2] (b) State the H–N–H bond angle in NH₄⁺. [1]

Section C: Forces and properties

10. Water boils at 100 °C. Hydrogen sulfide, H₂S, boils at about −60 °C, even though it has more electrons than water. Explain this difference. [3]

11. Iodine is a shiny grey solid that turns to a purple vapour on gentle heating. A student writes: “Iodine vaporises easily because the I–I bond is weak.” Explain why the student is wrong and give the correct explanation. [2]

12. Magnesium oxide melts at a far higher temperature than sodium chloride, although both have the same type of lattice. Explain why. [2]

13. Use the electronegativity values given above to classify the bonding in Cl₂, HCl and KCl as non-polar covalent, polar covalent or ionic. Show the values you use. [3]

14. Ethene has the structure H₂C=CH₂. (a) State the number of sigma (σ) bonds and pi (π) bonds in one molecule of ethene. [2] (b) Describe how the π bond forms. [1]

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Mark scheme

Question 1 (2 marks)

(a) A shared pair of electrons between two atoms (attracted to both nuclei). [1] Accept “electrostatic attraction between two nuclei and a shared pair of electrons”. Do not accept “atoms sharing electrons” without the word “pair”.

(b) The ability (or power) of an atom to attract the bonding pair of electrons in a covalent bond. [1] Do not accept answers that describe attracting a single electron or “gaining electrons”; that describes electron affinity.

Question 2 (3 marks)

• Magnesium ion drawn with no outer electrons (or with the outer shell shown empty) and the charge 2+. [1]
• Oxide ion drawn with eight outer electrons, six of one symbol (oxygen’s own) and two of the other (transferred from magnesium). [1]
• Charge of 2− on the oxide ion, with brackets around each ion. [1]

The most common error is transferring only one electron, which gives Mg⁺ and O⁻. Magnesium is in group 2 and loses two electrons; oxygen is in group 16 and gains two.

Question 3 (4 marks)

(a)

• Giant ionic lattice with strong electrostatic attraction between oppositely charged ions. [1]
• A large amount of energy is needed to overcome these attractions (acting in all directions). [1] Do not accept any mention of molecules or intermolecular forces: award 0 for part (a) if “intermolecular forces” is the explanation.

(b)

• In the solid, the ions are held in fixed positions and cannot move. [1]
• In the molten state, the ions are free to move and carry charge. [1] Do not accept “electrons move” in molten NaCl; the charge carriers are ions.

Question 4 (3 marks)

• A lattice (regular arrangement) of positive metal ions. [1]
• Surrounded by a sea of delocalised electrons; the bonding is the attraction between the ions and the delocalised electrons. [1]
• When a force is applied, layers of ions slide over each other, and the delocalised electrons move with them so the bonding is maintained (no strong repulsion develops). [1]

Question 5 (4 marks)

(a)

• In graphite each carbon atom forms three covalent bonds, leaving one delocalised electron per atom that can move through the layers. [1]
• In diamond each carbon forms four covalent bonds, so all outer electrons are used in bonding and none are free to move. [1]

(b)

• Graphite has layers (sheets) of carbon atoms. [1]
• The forces between the layers are weak, so the layers slide over each other easily. [1] Do not accept “weak covalent bonds between the layers”.

Question 6 (3 marks)

• Trigonal pyramidal. [1]
• Nitrogen has three bonding pairs and one lone pair (four electron pairs in total). [1]
• The lone pair repels more strongly than the bonding pairs, pushing the N–H bonds closer together. [1]

Accept “pyramidal”. Do not accept “triangular” or “trigonal planar”.

Question 7 (3 marks)

• Both molecules contain polar bonds because oxygen is more electronegative than carbon and hydrogen. [1]
• CO₂ is linear (180°), so the two bond dipoles are equal and opposite and cancel. [1]
• H₂O is bent, so the bond dipoles do not cancel and the molecule has an overall dipole. [1]

A labelled diagram showing δ+ and δ− and the direction of the dipoles can earn the second and third marks.

Question 8 (3 marks)

• 4 bonding pairs and 2 lone pairs (6 electron pairs in total). [1] Working: Xe has 8 valence electrons; four are used in bonds to F, leaving four electrons as two lone pairs.
• Square planar. [1]
• 90°. [1]

Do not accept “octahedral” for the shape: that is the electron-pair geometry, not the molecular shape. See the list of all 13 VSEPR shapes for the full set.

Question 9 (3 marks)

(a)

• A dative covalent (coordinate) bond. [1]
• The lone pair on nitrogen is donated to the H⁺ ion, so both electrons in the shared pair come from nitrogen. [1]

(b) 109.5°. [1] (Four bonding pairs and no lone pairs, so the ion is tetrahedral. Once formed, the dative bond is identical to the other three N–H bonds.)

More on this type of bond: coordinate covalent bonds.

Question 10 (3 marks)

• Water forms hydrogen bonds between molecules (H bonded to highly electronegative O, with lone pairs on O). [1]
• H₂S cannot form hydrogen bonds; its molecules are held by weaker permanent dipole and London (dispersion) forces. [1]
• More energy is needed to overcome the hydrogen bonds between water molecules, so water’s boiling point is higher. [1]

Do not accept any answer stating that O–H covalent bonds break on boiling; award a maximum of 1 mark if this appears.

Question 11 (2 marks)

• Vaporising iodine does not break the I–I covalent bonds; I₂ molecules are still present in the vapour. [1]
• Only the weak intermolecular (London dispersion) forces between I₂ molecules are overcome, which needs little energy. [1]

Question 12 (2 marks)

• Mg²⁺ and O²⁻ carry double the charge of Na⁺ and Cl⁻ (and Mg²⁺ is also smaller than Na⁺). [1]
• So the electrostatic attraction between the ions is much stronger, and more energy is needed to overcome it. [1]

Accept “larger lattice energy” for the second mark, provided the charges are mentioned. See lattice energy for the underlying idea.

Question 13 (3 marks)

• Cl₂: Δχ = 3.16 − 3.16 = 0 → non-polar covalent. [1]
• HCl: Δχ = 3.16 − 2.20 = 0.96 → polar covalent. [1]
• KCl: Δχ = 3.16 − 0.82 = 2.34 → ionic. [1]

Each mark needs both the difference and the correct classification. Examiners accept any sensible cut-off (for example, “above about 1.7 is ionic”), because the boundaries are guidelines rather than sharp lines. The worked examples in using electronegativity to predict bond type show where those guidelines fail.

Question 14 (3 marks)

(a) 5 σ bonds (four C–H and one in the C=C) [1] and 1 π bond [1]. Do not accept “6 σ” (double bonds contain only one σ bond).

(b) Sideways (side-on) overlap of p orbitals on the two carbon atoms, above and below the line joining the nuclei. [1]

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How the marks were lost: an examiner’s summary

Looking back across the scheme, most lost marks come from a small set of habits:

1. Mixing up models. Using “intermolecular forces” for an ionic compound (Q3) or “covalent bonds break” for a molecular substance (Q10, Q11) is the single biggest source of zero-mark answers.
2. Forgetting what carries the charge. Ions conduct in molten salts; electrons conduct in metals and graphite. Swapping them costs the mark.
3. Incomplete shape answers. A question asking for shape and angle needs both. The electron-pair geometry is not the molecular shape (Q8).
4. Missing charges and brackets on dot-and-cross diagrams (Q2).
5. Treating electronegativity cut-offs as laws. Quote the difference, apply a stated guide, and move on (Q13).

If you scored below 25 out of 41, revisit the topic articles linked above, then try the paper again in a week. A score above 35 means you are in good shape for this part of the course.

Key takeaways

• Ionic lattices: strong attractions between oppositely charged ions in all directions; conduction needs mobile ions.
• Metals: positive ions in a sea of delocalised electrons; conduction and malleability both come from those electrons.
• Molecular substances: boiling and melting overcome intermolecular forces, never the covalent bonds inside molecules.
• Shapes: count electron pairs, place lone pairs, then name the shape from the atoms and give the angle.
• Polarity needs both polar bonds and an asymmetric shape.