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A Level H2 Chemistry Periodic Table Quiz

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A-Level Chemistry H2 Quiz - Periodic Table

Name: __________________________
Class: __________________________
Date: __________________________
Score: ________ / 50

Duration: 45 minutes
Total Marks: 50

Instructions:

Answer all questions in the spaces provided.
You may use the Data Booklet.
The number of marks is given in brackets [ ] at the end of each question or part question.
Show all working for calculations.

Section A: Periodicity and Trends (Questions 1–5)

1. The first ionisation energies of the elements in Period 3 are shown below.

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

(a) Explain why the first ionisation energy of aluminium is lower than that of magnesium, despite aluminium having a higher nuclear charge. [2]

(b) Explain why the first ionisation energy of sulfur is lower than that of phosphorus. [2]

2. Describe and explain the trend in melting points of the Period 3 elements from sodium to aluminium. [3]

3. Silicon(IV) oxide (SiO₂) and phosphorus(V) oxide (P₄O₁₀) are both oxides of Period 3 elements.

(a) State the type of structure and bonding in SiO₂. [1]

(b) State the type of structure and bonding in P₄O₁₀. [1]

4. Sodium oxide (Na₂O) and sulfur dioxide (SO₂) react with water.

(a) Write an equation for the reaction of Na₂O with water. [1]

(b) Write an equation for the reaction of SO₂ with water. [1]

Na₂O solution: _______________
SO₂ solution: _______________

5. Explain why the atomic radius decreases across Period 3 from sodium to chlorine. [2]

Section B: Group 2 and Group 17 Trends (Questions 6–10)

6. The thermal stability of Group 2 nitrates decreases down the group.

(a) Write an equation for the thermal decomposition of magnesium nitrate, Mg(NO₃)₂. Include state symbols. [2]

(b) Explain, in terms of charge density and polarisation, why barium nitrate is more thermally stable than magnesium nitrate. [3]

7. A student heats a sample of calcium carbonate strongly.

(a) Write the equation for the decomposition. [1]

(b) The student repeats the experiment with strontium carbonate. Would you expect strontium carbonate to decompose at a higher or lower temperature than calcium carbonate? Explain your answer. [2]

8. Chlorine reacts with cold dilute aqueous sodium hydroxide.

(a) Write the ionic equation for this reaction. [2]

(b) State the oxidation states of chlorine in the products formed. [2]

9. Explain the trend in the oxidising power of the halogens from chlorine to iodine. Refer to atomic structure in your answer. [3]

10. Aqueous silver nitrate is added to separate solutions of sodium chloride and sodium iodide, followed by dilute aqueous ammonia.

(a) Describe the observations for the chloride ion test. [2]

With AgNO₃: __________________________________________________
With dilute NH₃: _______________________________________________

(b) Describe the observations for the iodide ion test. [2]

With AgNO₃: __________________________________________________
With dilute NH₃: _______________________________________________

Section C: Transition Elements (Questions 11–15)

11. Define the term transition element. [1]

12. Explain why transition elements exhibit variable oxidation states, whereas s-block elements generally do not. [2]

13. Iron(II) ions, Fe²⁺(aq), are pale green. When aqueous sodium hydroxide is added, a precipitate forms.

(a) State the colour of the precipitate. [1]

(b) Write the ionic equation for the formation of this precipitate. [1]

(c) The precipitate is left standing in air. Describe the change in appearance and explain the chemical change that occurs. [2]

14. Copper(II) ions form a complex ion with ammonia.

(a) Describe the observations when aqueous ammonia is added dropwise, then in excess, to aqueous copper(II) sulfate. [2]

(b) Write the formula of the complex ion formed in excess ammonia. [1]

15. Why are many transition metal complexes coloured? Explain in terms of d-orbitals and light absorption. [3]

Section D: Application and Synthesis (Questions 16–20)

16. An unknown white solid, X, is a Group 2 metal carbonate.

0.500 g of X is heated strongly until constant mass.
The residue weighs 0.282 g.
The gas evolved turns limewater milky.

(a) Identify the gas evolved. [1]

(b) Calculate the relative atomic mass ( $A_r$ ) of the Group 2 metal in X. Show your working. [3] (Assume Ar of C = 12.0, O = 16.0)

17. Vanadium is a transition element. The standard electrode potentials for some vanadium species are given below:

\begin{align*} \text{VO}_2^+ + 2\text{H}^+ + e^- &\rightleftharpoons \text{VO}^{2+} + \text{H}_2\text{O} & E^\ominus &= +1.00 \text{ V} \\ \text{VO}^{2+} + 2\text{H}^+ + e^- &\rightleftharpoons \text{V}^{3+} + \text{H}_2\text{O} & E^\ominus &= +0.34 \text{ V} \\ \text{V}^{3+} + e^- &\rightleftharpoons \text{V}^{2+} & E^\ominus &= -0.26 \text{ V} \end{align*}

Zinc has $E^\ominus (\text{Zn}^{2+}/\text{Zn}) = -0.76 \text{ V}$ .

(a) Determine the final oxidation state of vanadium when excess zinc is added to an acidic solution containing $\text{VO}_2^+$ ions. Explain your answer using $E^\ominus$ values. [3]

(b) State the colour of the vanadium species in the final oxidation state. [1]

18. Aluminium chloride (AlCl₃) behaves differently in the solid state and in the vapour phase at low pressure.

(a) In the vapour phase at low pressure, AlCl₃ exists as discrete molecules. Draw the dot-and-cross diagram for an AlCl₃ molecule, showing only outer shell electrons. [2]

(b) Explain why AlCl₃ has a relatively low sublimation point compared to NaCl. [2]

(c) In the solid state near its melting point, AlCl₃ exists as a dimer, Al₂Cl₆. Draw the structure of Al₂Cl₆, clearly showing any coordinate (dative) bonds. [2]

19. The table below shows the electrical conductivity of three Period 3 chlorides in the liquid state.

Chloride	NaCl	AlCl₃	SiCl₄
Conductivity	High	Low/None	None

(a) Explain why liquid NaCl conducts electricity. [1]

(b) Explain why liquid SiCl₄ does not conduct electricity. [1]

(c) AlCl₃ does not conduct electricity well in the liquid state. What does this suggest about the bonding in liquid AlCl₃? [1]

20. A solution contains both Fe²⁺(aq) and Fe³⁺(aq).

(a) Describe a chemical test to confirm the presence of Fe³⁺(aq). Include reagents and observations. [2]

(b) Describe a chemical test to confirm the presence of Fe²⁺(aq). Include reagents and observations. [2]

Answers

A-Level Chemistry H2 Quiz - Periodic Table (Answer Key)

1. (a) 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/shielded by 3s electrons) than the 3s orbital, so it requires less energy to remove. [1] (b) In sulfur, the electron is removed from a paired 3p orbital. [1] Electron-electron repulsion between the paired electrons makes it easier to remove one electron compared to phosphorus, where the 3p electron is unpaired. [1]

2. Na, Mg, and Al have metallic bonding. [1] The strength of metallic bonding increases from Na to Al because the number of delocalised electrons increases (1, 2, 3) and the ionic radius decreases, leading to higher charge density. [1] This requires more energy to overcome, so melting point increases. [1]

3. (a) Giant covalent / Macromolecular. [1] (b) Simple molecular. [1] (c) SiO₂ has strong covalent bonds throughout the giant lattice which require much energy to break. [1] P₄O₁₀ has weak intermolecular forces (van der Waals) between molecules which require little energy to overcome. [1]

4. (a) $\text{Na}_2\text{O}(s) + \text{H}_2\text{O}(l) \rightarrow 2\text{NaOH}(aq)$ [1] (b) $\text{SO}_2(g) + \text{H}_2\text{O}(l) \rightleftharpoons \text{H}_2\text{SO}_3(aq)$ [1] (c) Na₂O solution: pH 13–14 (Strongly alkaline). [1] SO₂ solution: pH 2–4 (Weakly/Acidic). [1]

5. Across the period, the number of protons (nuclear charge) increases. [1] Electrons are added to the same principal quantum shell, so shielding remains similar. This results in a greater effective nuclear charge, pulling the outer electrons closer to the nucleus. [1]

6. (a) $2\text{Mg(NO}_3)_2(s) \rightarrow 2\text{MgO}(s) + 4\text{NO}_2(g) + \text{O}_2(g)$ [1 for formulae, 1 for balancing/state symbols] (b) The Mg²⁺ ion is smaller than the Ba²⁺ ion, so it has a higher charge density. [1] Mg²⁺ polarises the nitrate ion more strongly than Ba²⁺. [1] This weakens the N-O bonds in the nitrate ion, making it easier to decompose (less thermally stable). [1]

7. (a) $\text{CaCO}_3(s) \rightarrow \text{CaO}(s) + \text{CO}_2(g)$ [1] (b) Higher temperature. [1] Sr²⁺ is larger than Ca²⁺, so it has lower charge density and polarises the carbonate ion less. The C-O bonds are less weakened, requiring more heat to break. [1]

8. (a) $\text{Cl}_2 + 2\text{OH}^- \rightarrow \text{Cl}^- + \text{ClO}^- + \text{H}_2\text{O}$ [1 for species, 1 for balancing] (b) Cl in Cl⁻: -1 [1]; Cl in ClO⁻: +1 [1]

9. Oxidising power decreases down the group. [1] Atomic radius increases and shielding increases, so the attraction between the nucleus and an incoming electron decreases. [1] It becomes harder to gain an electron to form the halide ion. [1]

10. (a) AgNO₃: White precipitate. [1] Dilute NH₃: Precipitate dissolves (forming a colourless solution). [1] (b) AgNO₃: Yellow (or cream) precipitate. [1] Dilute NH₃: Precipitate does not dissolve (or is insoluble). [1]

11. An element that forms at least one stable ion with a partially filled d-subshell. [1]

12. The energy difference between the 4s and 3d orbitals is small. [1] Therefore, different numbers of d-electrons can be involved in bonding/ionisation, allowing variable oxidation states. [1]

13. (a) Green precipitate. [1] (b) $\text{Fe}^{2+}(aq) + 2\text{OH}^-(aq) \rightarrow \text{Fe(OH)}_2(s)$ [1] (c) The precipitate turns brown/rust-coloured. [1] Fe(II) is oxidised to Fe(III) by oxygen in the air. [1]

14. (a) Dropwise: Pale blue precipitate forms. [1] Excess: Precipitate dissolves to form a deep blue solution. [1] (b) $[\text{Cu(NH}_3)_4(\text{H}_2\text{O})_2]^{2+}$ or $[\text{Cu(NH}_3)_4]^{2+}$ [1]

15. Ligands cause the d-orbitals to split into different energy levels. [1] Electrons absorb visible light photons to jump from lower to higher d-orbitals ( $d-d$ transition). [1] The colour observed is the complementary colour of the light absorbed. [1]

16. (a) Carbon dioxide ( $\text{CO}_2$ ). [1] (b) Mass of $\text{CO}_2$ lost = $0.500 - 0.282 = 0.218 \text{ g}$ . [1] Moles of $\text{CO}_2$ = $0.218 / 44.0 = 0.004955 \text{ mol}$ . Moles of $\text{MCO}_3$ = Moles of $\text{CO}_2$ = $0.004955 \text{ mol}$ . [1] $M_r$ of $\text{MCO}_3$ = $0.500 / 0.004955 = 100.9$ . $A_r(\text{M}) + 12.0 + 3(16.0) = 100.9$ . $A_r(\text{M}) = 100.9 - 60.0 = 40.9$ . [1] (c) Calcium (Ca). [1] (Accept close calculation leading to Ca)

17. (a) Zn is a strong reducing agent ( $E^\ominus = -0.76 \text{ V}$ ). $\text{VO}_2^+ \rightarrow \text{VO}^{2+}$ : $E^\ominus_{cell} = 1.00 - (-0.76) = +1.76 \text{ V}$ (Feasible). $\text{VO}^{2+} \rightarrow \text{V}^{3+}$ : $E^\ominus_{cell} = 0.34 - (-0.76) = +1.10 \text{ V}$ (Feasible). $\text{V}^{3+} \rightarrow \text{V}^{2+}$ : $E^\ominus_{cell} = -0.26 - (-0.76) = +0.50 \text{ V}$ (Feasible). Final oxidation state is +2. [1 for each correct potential comparison or logical step, max 3] (b) Violet / Purple. [1]

18. (a) Diagram showing Al sharing 3 electrons with 3 Cl atoms. Al has 6 valence electrons (electron deficient). Cl has 8 (octet). [2] (b) AlCl₃ is simple molecular (covalent) with weak intermolecular forces. [1] NaCl is giant ionic with strong electrostatic forces. [1] (c) Diagram showing two AlCl₃ units linked by two dative bonds from Cl lone pairs to Al atoms. Each Al has octet. [2]

19. (a) Contains mobile ions ( $\text{Na}^+$ and $\text{Cl}^-$ ) that can carry charge. [1] (b) Consists of simple covalent molecules with no free ions or electrons. [1] (c) It is covalent / molecular in nature (not ionic). [1]

20. (a) Add potassium thiocyanate (KSCN) or sodium hydroxide. [1] With KSCN: Blood red solution. With NaOH: Red-brown precipitate. [1] (b) Add potassium manganate(VII) ( $\text{KMnO}_4$ ) or sodium hydroxide. [1] With $\text{KMnO}_4$ : Purple solution decolourises. With NaOH: Green precipitate (turning brown). [1]