Secondary 4 Combined Science Chemistry Periodic Table Quiz

Secondary 4 Combined Science Chemistry Quiz - Periodic Table

Answer Key

Section A: Multiple Choice Questions (1–5)

1. B

• Elements in the same period have the same number of electron shells. [1]
• Common trap: Students may select A (historical arrangement by atomic mass) or D (noble gases are in Group 0, not Group I).

2. C — Group VI, Period 3

• Electronic configuration 2.8.6 → 6 electrons in the outer shell → Group VI; 3 electron shells → Period 3. [1]

3. C — Chlorine

• Halogens are Group VII elements. Chlorine (proton number 17) is in Group VII. [1]

4. C — Atomic radius

• Across a period, atomic radius decreases due to increasing nuclear charge pulling electrons closer to the nucleus. [1]

5. B — They have a full outer shell of electrons.

• Noble gases have a stable electronic configuration (full outer shell), making them chemically unreactive. [1]

Section B: Short Answer Questions (6–13)

6. (a) None of the elements A–D is a noble gas. None has a full outer shell of electrons. [1]

• Accept: "There is no noble gas in the table" with correct reasoning.

(b) C and D are in the same period. [1] Both have 3 electron shells (configurations 2.8.2 and 2.8.7).

(c) Element C forms a cation with a charge of +2. [1] Element C has 2 electrons in its outer shell and loses both to achieve a stable configuration, forming C²⁺.

[3]

7. The atomic radius decreases across Period 3 from sodium to argon. [1]

• This is because the nuclear charge (number of protons) increases across the period, [1] pulling the electrons in the same shell closer to the nucleus, resulting in a smaller atomic radius. [1]

[3]

8. (a) Sodium oxide is a basic oxide. [1]

(b) Na₂O + H₂O → 2NaOH [1]

• Marking: 1 mark for correct formulae, 1 mark for balancing (awarded here as single mark for correct equation).

(c) Add universal indicator (or litmus paper) to the solution. [1] The solution will turn indicator blue/purple (or blue litmus remains blue / red litmus turns blue), confirming the solution is alkaline (pH > 7).

• Accept: Use of pH paper showing pH > 7.

[3]

9. (a) Cl₂ [1]

(b) Chlorine is a greenish-yellow gas at room temperature. [1]

(c) Chlorine has 7 electrons in its outer shell (electronic configuration 2.8.7). [1] Elements in Group VII have 7 valence electrons.

[3]

10. (a) 2.8.8.2 [1]

(b) Both react with water to produce hydrogen gas / both form ions with a 2+ charge / both form white alkaline oxides / both are metals. [1]

(c) The gas is hydrogen. [1] Test: Use a lighted splint — the gas burns with a pop sound. [1]

[3]

11. (a) Ar (argon) — smallest atomic radius in Period 3 due to greatest nuclear charge pulling electrons closest. [1]

(b) Si (silicon) — it is a metalloid, showing properties intermediate between metals and non-metals. [1]

(c) Mg and Cl — Mg (Group II) forms Mg²⁺ and Cl (Group VII) forms Cl⁻, giving MgCl₂ which has the formula XY₂. [1]

[3]

12. From sodium to silicon, the elements have giant metallic (Na, Mg, Al) and giant covalent (Si) structures. [1] These structures have strong bonds (metallic bonds or covalent bonds) throughout the lattice, requiring large amounts of energy to break, hence high melting points. [1]

From phosphorus to argon, the elements exist as simple covalent molecules (P₄, S₈, Cl₂, Ar) with weak intermolecular forces (van der Waals' forces) between molecules. [1] Little energy is needed to overcome these weak forces, so melting points are low.

[3]

13. (a) The gas is hydrogen (H₂). [1]

(b) Element X belongs to Group I. [1] Group I metals react vigorously with cold water to produce an alkaline solution and hydrogen gas. [1]

(c) 2Na + 2H₂O → 2NaOH + H₂ [1]

• Accept: 2X + 2H₂O → 2XOH + H₂ if student uses X.

[3]

Section C: Structured / Data-Based Questions (14–20)

14. (a) Sodium, magnesium, and aluminium have delocalised (mobile) electrons throughout their metallic lattice. [1] These free electrons can carry electrical charge through the structure, making them good conductors.

(b) A semiconductor is a material that conducts electricity poorly under normal conditions but can conduct better when energy (e.g., heat) is supplied. [1] Silicon has a giant covalent structure with no free electrons at low temperatures, but when heated, some electrons gain enough energy to move and conduct.

(c) Phosphorus, sulphur, chlorine, and argon exist as simple molecules with weak intermolecular forces (van der Waals' forces) between them. [1] Very little energy is needed to overcome these weak forces, so they have low melting points. [1]

(d) Silicon has the highest melting point (1410 °C). [1] Silicon has a giant covalent structure where each silicon atom is covalently bonded to four other silicon atoms in a tetrahedral arrangement. [1] A large amount of energy is needed to break the strong covalent bonds throughout the structure.

[5]

15. (a) The reactivity of Group I metals increases going down the group. [1] Lithium reacts slowly, sodium reacts more vigorously, and potassium reacts most vigorously (igniting with a lilac flame).

(b) Going down Group I, the atomic radius increases and the outer electron is further from the nucleus. [1] The attraction between the nucleus and the outer electron is weaker (more electron shielding), so the outer electron is more easily lost, making the metal more reactive. [1]

(c) Rubidium would react even more vigorously than potassium, possibly exploding on contact with water. [1] This is because rubidium is further down Group I, has a larger atomic radius, and its outer electron is even more easily lost. [1]

(d) The gas produced is hydrogen. [1] Test: Hold a lighted splint near the gas — it burns with a pop sound. [1]

[5]

16. (a) Electronic configuration of Y: 2.8.8.2 [1]

(b) Group II, Period 4. [1] (2 electrons in outer shell → Group II; 4 electron shells → Period 4)

(c) Y + Cl₂ → YCl₂ [1]

• Accept: Ca + Cl₂ → CaCl₂

(d) The compound has ionic bonding. [1] Element Y (a Group II metal) loses 2 electrons to form Y²⁺, and each chlorine atom (Group VII non-metal) gains 1 electron to form Cl⁻. [1] The transfer of electrons results in oppositely charged ions that are held together by strong electrostatic forces of attraction.

(e) The compound has a high melting point. [1] This is because ionic compounds have a giant ionic lattice structure with strong electrostatic forces between oppositely charged ions, requiring a large amount of energy to break.

• Accept: Conducts electricity when molten/dissolved (ions are free to move) / soluble in water / hard and brittle.

[5]

17. (a) First ionisation energy is the energy required to remove one mole of electrons from one mole of gaseous atoms to form one mole of gaseous 1+ ions. [1]

• Accept: "Energy required to remove the outermost electron from a gaseous atom."

(b) Across Period 3, the nuclear charge increases (more protons) while the electrons are added to the same shell. [1] The increasing positive charge pulls the outer electrons closer, making them harder to remove, so the first ionisation energy generally increases. [1]

(c) Aluminium has a lower first ionisation energy than magnesium because aluminium's outer electron is in a 3p orbital, which is at a higher energy level than magnesium's 3s orbital. [1] The 3p electron is also slightly further from the nucleus and experiences more shielding from the 3s electrons, making it easier to remove.

(d) Sulphur has a lower first ionisation energy than phosphorus because in phosphorus, the 3p orbitals are singly occupied (3p³, half-filled, stable), whereas in sulphur, one 3p orbital contains a pair of electrons. [1] The electron-electron repulsion in the paired orbital makes it easier to remove one electron from sulphur.

[5]

18. (a) Group II [1]

(b) Element Q is strontium (Sr) or barium (Ba). [1] Since Q is below P (magnesium, proton number 12) in Group II and reacts more vigorously with water, Q must be a Group II metal further down the group (e.g., Ca, Sr, or Ba). [1]

• Accept: Calcium (Ca) — reacts more vigorously than magnesium.

(c) Mg²⁺ [1]

(d) Reactivity increases going down Group II. [1] This is because the atomic radius increases and there is more electron shielding, so the outer electrons are further from the nucleus and are more easily lost. [1]

[4]

19. (a)

• Na₂O: basic [½]
• MgO: basic [½]
• Al₂O₃: amphoteric [½]
• SiO₂: acidic [½]
• P₄O₁₀: acidic [½]
• SO₂: acidic [½]
• Marking: ½ each, round to 3 marks total

(b) Na₂O, MgO, and Al₂O₃ have giant ionic structures with strong electrostatic forces of attraction between oppositely charged ions. [1] A large amount of energy is required to break these strong ionic bonds, resulting in high melting points. [1]

(c) SO₂ exists as simple covalent molecules with weak van der Waals' forces between molecules. [1] Very little energy is needed to overcome these intermolecular forces, so SO₂ has a very low melting point.

(d) Na₂O + 2HCl → 2NaCl + H₂O [1]

[5]

20. (a) Compound Z has ionic bonding. [1] The fact that molten Z conducts electricity indicates the presence of mobile ions, which is characteristic of ionic compounds. [1]

(b) The observations do not fully support the conclusion that Z is a metal oxide. [1] While metal oxides are often ionic and conduct when molten, the observation that the aqueous solution has pH 7 (neutral) is unusual for a metal oxide, which typically forms alkaline solutions. [1] The effervescence with HCl producing CO₂ (limewater turns milky) suggests Z may contain carbonate ions, so Z could be a metal carbonate rather than a metal oxide.

(c) Add aqueous sodium hydroxide and warm — test for ammonia gas (if Z contains ammonium ions); or add silver nitrate solution to test for chloride ions; or perform a flame test to identify the metal ion present. [1]

• Accept: Any reasonable further test.

(d) In solid MgO, the ions (Mg²⁺ and O²⁻) are fixed in position in the ionic lattice and cannot move, so they cannot carry charge. [1] When molten, the ions are free to move and can carry electrical charge through the liquid, allowing it to conduct electricity. [1]

[5]

Total: 40 marks

Marking Notes

• Award marks for correct chemistry even if wording differs from the mark scheme.
• For explanation questions, look for correct scientific reasoning, not just keywords.
• In calculation-based or formula-based answers, award credit for correct formulae even if the final answer is not explicitly required.
• Common mistakes to watch for:
  • Confusing Group number with period number
  • Stating "atomic mass" instead of "proton number" for modern Periodic Table arrangement
  • Forgetting that noble gases are in Group 0 (or Group 18), not Group VIII
  • Confusing intermolecular forces with covalent bonds when explaining melting points
  • Not specifying "gaseous" in ionisation energy definitions