A Level H1 Chemistry Periodic Table Quiz

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

1. Definition: The energy required to remove one mole of electrons from one mole of gaseous atoms to form one mole of gaseous 1+ ions. [2]

2. Mg vs Al: Magnesium has a stable $3s^2$ configuration. Aluminium has a $3s^2 3p^1$ configuration. The electron in the $3p$ orbital is higher in energy and further from the nucleus (more shielded), making it easier to remove. [2]

3. Na vs K: Potassium has a larger atomic radius. [1] Potassium has more occupied principal energy levels (n=4) compared to sodium (n=3), increasing the distance between the nucleus and valence electrons. [1]

4. F vs Cl: Fluorine has a smaller atomic radius. [1] The bonding pair of electrons is closer to the nucleus and less shielded, resulting in a stronger attraction for the shared pair of electrons. [1]

5. Trend: Generally increases across the period. [1] This is due to the increase in nuclear charge with a constant shielding effect, increasing the attraction between the nucleus and valence electrons. [1]

6. Atomic Radius: Nuclear charge increases across the period [1] while electrons are added to the same principal energy level (shielding remains relatively constant), pulling electrons closer to the nucleus. [1]

7. Na 2nd IE: The second electron is removed from a stable, full inner shell (2p⁶) [1] which is much closer to the nucleus and experiences significantly less shielding, requiring much more energy. [1]

8. Element: Aluminium (Al). [1]

9. Equation: $\text{SiO}_2(s) + 2\text{NaOH}(aq) \rightarrow \text{Na}_2\text{SiO}_3(aq) + \text{H}_2\text{O}(l)$ [2] (1 mark for correct formula/states, 1 mark for balancing).

10. NaCl vs MgCl₂: $\text{MgCl}_2$ has a higher melting point. [1] $\text{Mg}^{2+}$ has a higher charge density (smaller radius, higher charge) than $\text{Na}^+$. [1] This results in stronger electrostatic attractions between the cation and the chloride anions in the giant ionic lattice. [1]

11. $\text{P}_4\text{O}_{10}$: It is a strong Lewis acid. [1] The P-O bonds are highly polar, and the phosphorus atoms are electron-deficient, leading to a very exothermic reaction with water to form phosphoric acid. [1]

12. Sulfur Oxide: $\text{SO}_2$ (or $\text{SO}_3$). [1] Acidic. [1]

13. Nature Trend: Oxides change from basic ($\text{Na}_2\text{O}, \text{MgO}$) to amphoteric ($\text{Al}_2\text{O}_3$) to acidic ($\text{SiO}_2, \text{P}_4\text{O}_{10}, \text{SO}_2, \text{Cl}_2\text{O}_7$). [2]

14. Amphoteric: It reacts with both acids (e.g., $\text{HCl}$) [1] and strong bases (e.g., $\text{NaOH}$) to form salts and water. [1]

15. Group 1 Reactivity: Increases down the group. [1] Atomic radius increases and shielding increases. [1] The attraction between the nucleus and the valence electron decreases, making it easier to lose the electron. [1]

16. He vs Ne: Neon has a larger molecular size/more electrons. [1] This leads to stronger London dispersion forces (Van der Waals forces) between neon atoms. [1]

17. (a) Formula: $\text{HX}$ (where X is halogen). [1] (b) HCl vs NaCl: $\text{HCl}$ is a simple covalent molecule with weak intermolecular forces. [1] $\text{NaCl}$ is a giant ionic lattice with strong electrostatic attractions between ions. [1]

18. F vs I: Fluorine is a stronger oxidizing agent. [1] Fluorine has a smaller atomic radius and higher electronegativity. [1] It has a greater tendency to attract and gain an electron to achieve a stable noble gas configuration. [1]

19. Al Structure: Giant metallic structure. [1] Consists of $\text{Al}^{3+}$ cations surrounded by a sea of delocalized valence electrons held by strong electrostatic forces. [1]

20. (a) Element: Aluminium (Al). [1] (b) Shape/Angle: Trigonal planar. [1] $120^\circ$. [1]