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A Level H1 Chemistry Periodic Table Quiz

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Questions

A-Level Chemistry H1 Quiz - Periodic Table

Name: __________________________
Class: __________________________
Date: __________________________
Score: ______ / 40

Duration: 45 minutes
Total Marks: 40

Instructions:

Answer all questions.
Write your answers in the spaces provided.
The number of marks is given in brackets [ ] at the end of each question or part question.
You may use a scientific calculator.
A Data Booklet is provided for reference.

Section A: Multiple Choice & Short Answer (Questions 1–5)

1. Which property generally increases across Period 3 from sodium to chlorine?
A. Atomic radius
B. Metallic character
C. First ionisation energy
D. Electrical conductivity

[1]

2. Which oxide of Period 3 elements is amphoteric?
A. Na₂O
B. MgO
C. Al₂O₃
D. SiO₂

[1]

3. Explain why the first ionisation energy of aluminium is lower than that of magnesium, despite aluminium having a higher nuclear charge.

[2]

4. State the trend in the melting points of the Group 17 elements (halogens) from chlorine to iodine and explain this trend in terms of structure and bonding.

[3]

5. Write a balanced equation, including state symbols, for the reaction of phosphorus(V) chloride with water.

[2]

Section B: Structured Questions (Questions 6–15)

6. The table below shows the first ionisation energies of the elements in Period 3.

Element	Na	Mg	Al	Si	P	S	Cl	Ar
1st IE / kJ mol⁻¹	496	738	578	789	1012	1000	1251	1521

(a) Explain the general increase in first ionisation energy from sodium to argon.

[2]

(b) Explain the drop in first ionisation energy between phosphorus and sulfur.

[2]

7. Magnesium and sulfur are both Period 3 elements.

(a) Describe the structure and bonding in solid magnesium.

[2]

(b) Sulfur exists as S₈ molecules. Explain why sulfur has a much lower melting point than magnesium.

[2]

8. Sodium reacts vigorously with water, while magnesium reacts very slowly with cold water.

(a) Write the equation for the reaction of sodium with cold water.

[1]

(b) Suggest one observation when sodium reacts with water.

[1]

[2]

9. Chlorine is used to purify water. When chlorine is added to water, it undergoes disproportionation.

(a) Define the term disproportionation.

[1]

(b) Write the equation for the reaction of chlorine with water.

[1]

10. Aluminium oxide is described as an amphoteric oxide.

(a) Define the term amphoteric.

[1]

(b) Write the equation for the reaction of aluminium oxide with hydrochloric acid.

[1]

11. The chlorides of Period 3 elements show a trend in their behaviour with water.

(a) Describe the pH of the solution formed when silicon tetrachloride (SiCl₄) is added to water.

[1]

(b) Write the equation for the reaction of SiCl₄ with water.

[1]

[2]

12. Consider the elements in Group 2 (alkaline earth metals).

(a) State the trend in the solubility of Group 2 sulfates as you go down the group.

[1]

(b) State the trend in the thermal stability of Group 2 nitrates as you go down the group.

[1]

[2]

13. Nitrogen and phosphorus are both in Group 15.

(a) Nitrogen exists as N₂ gas, while phosphorus exists as P₄ solid. Explain this difference in physical state at room temperature.

[2]

(b) Nitrogen is relatively unreactive compared to phosphorus. Explain why.

[2]

14. Transition elements are located in the d-block of the Periodic Table.

(a) State two characteristic properties of transition elements.

[2]

(b) Explain why scandium (Sc) is not always considered a typical transition element.

[2]

15. The atomic radius of potassium is larger than that of sodium.

(a) Explain this trend.

[2]

(b) Predict and explain which element, potassium or sodium, will have a lower first ionisation energy.

[2]

Section C: Data Response & Application (Questions 16–20)

16. The following data refers to the oxides of Period 3 elements.

Oxide	Na₂O	MgO	Al₂O₃	SiO₂	P₄O₁₀	SO₃
Melting Point / °C	1132	2852	2072	1610	340	17
Nature	Basic	Basic	Amphoteric	Acidic	Acidic	Acidic

(a) Explain the high melting point of MgO compared to P₄O₁₀.

[2]

(b) Why is Al₂O₃ classified as amphoteric while MgO is basic?

[2]

17. A student investigates the reaction of Group 2 metals with water.

(a) Write the general equation for the reaction of a Group 2 metal, M, with water.

[1]

(b) The pH of the resulting solution increases down the group. Explain why the hydroxides become more soluble and more alkaline down the group.

[2]

18. Chlorine reacts with cold dilute sodium hydroxide to form bleach.

(a) Write the ionic equation for this reaction.

[1]

(b) Identify the useful product formed in this reaction.

[1]

19. Consider the halogens chlorine, bromine, and iodine.

(a) Describe the colour and physical state of bromine at room temperature.

[1]

(b) Chlorine gas is bubbled through aqueous potassium bromide. Describe the observation and write the ionic equation for the reaction.

Observation: ___________________________________________________________
Equation: ______________________________________________________________

[2]

20. The table shows the electronegativity values of Period 3 elements.

Element	Na	Mg	Al	Si	P	S	Cl
Electronegativity	0.9	1.2	1.5	1.8	2.1	2.5	3.0

(a) Define electronegativity.

[2]

(b) Explain the trend in electronegativity across Period 3.

[2]

*** End of Quiz ***

Answers

A-Level Chemistry H1 Quiz - Periodic Table (Answer Key)

1. C
[1]

2. C
[1]

The electron removed from Al is from a 3p orbital, whereas the electron removed from Mg is from a 3s orbital. [1]
The 3p orbital is higher in energy and further from the nucleus (and experiences more shielding) than the 3s orbital, making it easier to remove. [1]
[2]

Trend: Melting point increases from chlorine to iodine. [1]
Explanation: The molecules become larger with more electrons. [1]
This leads to stronger van der Waals’ forces (London dispersion forces) between molecules, requiring more energy to overcome. [1]
[3]

5.
PCl₅(s) + 4H₂O(l) → H₃PO₄(aq) + 5HCl(aq)
[1] for correct formulae, [1] for balancing and state symbols.
[2]

6.
(a)

Across the period, the number of protons (nuclear charge) increases. [1]
Electrons are added to the same principal quantum shell, so shielding remains similar. The increased nuclear attraction pulls electrons closer, making them harder to remove. [1]
[2]

(b)

In sulfur, the electron is removed from a paired 3p orbital. [1]
Electron-electron repulsion within the paired orbital makes it easier to remove the electron compared to phosphorus, where the 3p electron is unpaired. [1]
[2]

7.
(a)

Giant metallic structure. [1]
Consists of Mg²⁺ cations in a sea of delocalised electrons, held by strong electrostatic forces. [1]
[2]

(b)

Sulfur consists of simple molecular structures (S₈) held together by weak van der Waals’ forces. [1]
Magnesium has strong metallic bonding. Less energy is required to overcome the weak intermolecular forces in sulfur than the strong metallic bonds in magnesium. [1]
[2]

8.
(a)
2Na(s) + 2H₂O(l) → 2NaOH(aq) + H₂(g)
[1]

(b)

Sodium floats / moves rapidly on the surface / melts into a sphere / effervescence. (Any one) [1]
[1]

(c)

Sodium has a larger atomic radius and lower nuclear charge compared to magnesium (within the same period context, but primarily group trend logic applies if comparing Na/Mg reactivity generally, but here it's period).
Correction for Period 3 context: Sodium has only one valence electron to lose, while magnesium has two. The first ionisation energy of Na is much lower than Mg, making it easier for Na to lose its electron and react. [1]
Also, Na⁺ has a lower charge density than Mg²⁺, resulting in weaker attraction for the remaining electrons, facilitating ionisation. [1]
[2]

9.
(a)

A reaction in which the same element is both oxidised and reduced. [1]
[1]

(b)
Cl₂(g) + H₂O(l) ⇌ HCl(aq) + HClO(aq)
[1]

(c)

In HCl: -1 [0.5]
In HClO: +1 [0.5]
[1]

10.
(a)

An oxide that can react with both acids and bases to form a salt and water. [1]
[1]

(b)
Al₂O₃(s) + 6HCl(aq) → 2AlCl₃(aq) + 3H₂O(l)
[1]

11.
(a)

Acidic (pH < 7). [1]
[1]

(b)
SiCl₄(l) + 2H₂O(l) → SiO₂(s) + 4HCl(aq)
[1]

(c)

Carbon has no available d-orbitals in its valence shell to expand its octet and accept lone pairs from water molecules. [1]
Silicon has empty 3d orbitals that can accept lone pairs from water, allowing hydrolysis to occur. [1]
[2]

12.
(a)

Solubility decreases down the group. [1]
[1]

(b)

Thermal stability increases down the group. [1]
[1]

(c)

As the cation size increases down the group, the charge density decreases. [1]
Lower charge density means less polarisation of the nitrate ion, making it harder to decompose (more stable). [1]
[2]

13.
(a)

N₂ consists of small non-polar molecules with weak van der Waals’ forces, so it is a gas. [1]
P₄ consists of larger molecules with stronger van der Waals’ forces, so it is a solid. [1]
[2]

(b)

Nitrogen has a triple bond (N≡N) which is very strong and requires high energy to break. [1]
Phosphorus has single bonds (P-P) which are weaker and more reactive. [1]
[2]

14.
(a)

Form coloured ions/compounds. [1]
Have variable oxidation states. [1]
(Other acceptable answers: act as catalysts, form complex ions)
[2]

(b)

Scandium only forms the Sc³⁺ ion. [1]
Sc³⁺ has an empty d-subshell (3d⁰), so it does not exhibit typical transition element properties like coloured ions or variable oxidation states. [1]
[2]

15.
(a)

Potassium has an additional electron shell compared to sodium. [1]
This increases the atomic radius and shielding, outweighing the increased nuclear charge. [1]
[2]

(b)

Potassium will have a lower first ionisation energy. [1]
The outer electron is further from the nucleus and more shielded, so the attraction is weaker and less energy is required to remove it. [1]
[2]

16.
(a)

MgO has a giant ionic lattice structure with strong electrostatic forces between Mg²⁺ and O²⁻ ions. [1]
P₄O₁₀ has a simple molecular structure with weak van der Waals’ forces between molecules. [1]
[2]

(b)

Al³⁺ has a high charge density, which polarises the oxide ion, giving the bond significant covalent character. This allows it to react with bases. [1]
Mg²⁺ has a lower charge density, so MgO is predominantly ionic and basic. [1]
[2]

17.
(a)
M(s) + 2H₂O(l) → M(OH)₂(aq) + H₂(g)
[1]

(b)

Down the group, the lattice energy of the hydroxide decreases more rapidly than the hydration energy. [1]
This makes the hydroxides more soluble, releasing more OH⁻ ions into solution, resulting in a higher pH. [1]
[2]

18.
(a)
Cl₂ + 2OH⁻ → Cl⁻ + ClO⁻ + H₂O
[1]

(b)

Sodium chlorate(I) (or bleach/disinfectant). [1]
[1]

19.
(a)

Red-brown liquid. [1]
[1]

(b)

Observation: Solution turns orange/brown. [1]
Equation: Cl₂ + 2Br⁻ → 2Cl⁻ + Br₂ [1]
[2]

20.
(a)

The power of an atom to attract the bonding pair of electrons in a covalent bond. [2]
[2]

(b)

Nuclear charge increases across the period. [1]
Shielding remains constant (same shell), so the attraction for bonding electrons increases. [1]
[2]