A Level H2 Chemistry Atomic Structure Bonding Quiz

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A-Level Chemistry H2 Quiz - Atomic Structure Bonding

Name: ____________________
Class: ____________________
Date: ____________________
Score: ________ / 55

Duration: 60 minutes
Total Marks: 55
Instructions: Answer all questions. Use of the Data Booklet is permitted. Show all working for calculations.

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Section A: Atomic Structure & Electron Configuration

Questions 1–5

1. An ion $\text{Y}^{3+}$ contains 36 electrons and 48 neutrons. Identify the element $\text{Y}$ and write its full electron configuration in the ground state. [3]
   
   
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2. Compare the first ionisation energy of Magnesium ($\text{Mg}$) and Aluminium ($\text{Al}$). Explain the trend observed with reference to electron configuration. [3]
   
   
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3. Write the electron configuration of the $\text{Cu}^{2+}$ ion. Explain why the $4s$ electrons are removed before the $3d$ electrons. [3]
   
   
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4. Define the term 'first ionisation energy' and explain why the second ionisation energy of Sodium ($\text{Na}$) is significantly higher than its first. [3]
   
   
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5. An element $\text{Z}$ has the ground state electron configuration $[\text{Ar}] 3\text{d}^{10} 4\text{s}^2 4\text{p}^3$. State the group and period of $\text{Z}$ in the Periodic Table. [2]
   
   
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Section B: Chemical Bonding & Molecular Geometry

Questions 6–12

6. Predict the shape and bond angle of the $\text{BF}_3$ molecule. Explain your answer using VSEPR theory. [3]
   
   
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7. Draw the Lewis structure of the $\text{NO}_3^-$ ion, showing all lone pairs and formal charges. [3]
   
   
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8. Explain why $\text{H}_2\text{O}$ has a higher boiling point than $\text{H}_2\text{S}$, despite $\text{H}_2\text{S}$ having a larger relative molecular mass. [3]
   
   
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9. Describe the bonding in graphite and explain why it is an electrical conductor while diamond is not. [4]
   
   
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10. Predict the shape of the $\text{SF}_4$ molecule. Explain why the bond angles deviate from the ideal tetrahedral angle. [3]
    
    
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11. Compare the strength of the covalent bond in $\text{HF}$ versus $\text{HCl}$. Justify your answer based on atomic radii and orbital overlap. [3]
    
    
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12. Explain the term 'electronegativity' and describe how it varies across Period 3. [3]
    
    
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Section C: Advanced Bonding & Application

Questions 13–20

13. $\text{XeF}_2$ is a known compound. Predict its molecular geometry and explain the presence of lone pairs on the central atom. [3]
    
    
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14. Write an ionic equation to represent the reaction of $\text{Al}_2\text{O}_3(\text{s})$ with hot aqueous sodium hydroxide. [2]
    
    
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15. $\text{ICl}_3$ is a covalent compound that exhibits electrical conductivity in the liquid state. Suggest an explanation for this property with the aid of an equation. [4]
    
    
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16. Using the concept of hybridisation, explain the geometry of the carbon atoms in ethene ($\text{C}_2\text{H}_4$). [3]
    
    
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17. Describe the difference between a $\sigma$-bond and a $\pi$-bond in terms of orbital overlap. [3]
    
    
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18. Explain why $\text{BeCl}_2$ exists as a polymer in the solid state but as discrete molecules in the gas phase. [4]
    
    
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19. Compare the lattice energy of $\text{MgO}$ and $\text{NaCl}$. Explain which compound has a higher melting point and why. [4]
    
    
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20. Predict the shape of the $\text{ClO}_4^-$ ion and state the hybridisation of the central chlorine atom. [2]
    
    
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Answer Key - A-Level Chemistry H2 Quiz: Atomic Structure Bonding

1. Element Y: Rubidium (Rb) [1]

   • Protons = electrons + charge = $36 + 3 = 39$. Atomic number 39 is Rb. [1]
   • Configuration: $1\text{s}^2 2\text{s}^2 2\text{p}^6 3\text{s}^2 3\text{p}^6 4\text{s}^2 3\text{d}^{10} 4\text{p}^6 5\text{s}^1$ (or $[\text{Kr}] 5\text{s}^1$) [1]

2. Al has a lower 1st IE than Mg [1]

   • Mg: $3\text{s}^2$ (full subshell) [1]
   • Al: $3\text{s}^2 3\text{p}^1$. The $3\text{p}$ electron is higher in energy/further from nucleus/more shielded than $3\text{s}$ electrons, making it easier to remove. [1]

3. Configuration: $[\text{Ar}] 3\text{d}^9$ [1]

   • Neutral Cu: $[\text{Ar}] 3\text{d}^{10} 4\text{s}^1$. [1]
   • Electrons are removed from the highest principal quantum number shell first ($n=4$ before $n=3$). [1]

4. Definition: Energy required to remove one mole of electrons from one mole of gaseous atoms to form one mole of gaseous $1+$ ions. [1]

   • Na 1st IE removes $3\text{s}^1$ electron. [1]
   • Na 2nd IE removes electron from $2\text{p}^6$ (inner shell), which is closer to the nucleus and experiences much stronger electrostatic attraction/less shielding. [1]

5. Group 15 (or Group V) [1]

   • Period 4 [1]

6. Trigonal Planar [1]

   • Bond angle: $120^\circ$ [1]
   • 3 bonding pairs, 0 lone pairs; electron pairs repel to maximum distance. [1]

7. Structure: Central N with three O atoms. [1]

   • One $\text{N}=\text{O}$ double bond, two $\text{N}-\text{O}$ single bonds (resonance). [1]
   • Formal charge: $-1$ on the single-bonded oxygens, $0$ on N and double-bonded O. [1]

8. Hydrogen Bonding [1]

   • $\text{H}_2\text{O}$ has strong H-bonds due to high electronegativity difference between O and H. [1]
   • $\text{H}_2\text{S}$ only has weaker dipole-dipole and London forces. [1]

9. Graphite: Hexagonal layers, $\text{C}-\text{C}$ bonds within layers, weak London forces between layers. [2]

   • Each C is bonded to 3 others; one delocalised electron per C atom. [1]
   • Delocalised electrons are free to move and carry charge. Diamond has all electrons localised in $\sigma$-bonds. [1]

10. See-saw [1]

    • 4 bonding pairs, 1 lone pair. [1]
    • Lone pair-bond pair repulsion > bond pair-bond pair repulsion, pushing bonds closer together (angle $< 90^\circ$ and $< 120^\circ$). [1]

11. HF bond is stronger [1]

    • F is smaller than Cl. [1]
    • Better orbital overlap between $1\text{p}$ (H) and $2\text{p}$ (F) compared to $3\text{p}$ (Cl). [1]

12. Definition: The ability of an atom to attract the shared pair of electrons in a covalent bond. [1]

    • Increases across Period 3. [1]
    • Nuclear charge increases while shielding remains constant, increasing attraction for bonding electrons. [1]

13. Linear [1]

    • 2 bonding pairs, 3 lone pairs on Xe. [1]
    • Lone pairs occupy equatorial positions to minimize repulsion. [1]

14. $\text{Al}_2\text{O}_3(\text{s}) + 2\text{OH}^-(\text{aq}) + 3\text{H}_2\text{O}(\text{l}) \rightarrow 2[\text{Al}(\text{OH})_4]^-(\text{aq})$ [2]

15. Autoionization [1]

    • $2\text{ICl}_3(\text{l}) \rightleftharpoons \text{ICl}_2^+(\text{l}) + \text{ICl}_4^-(\text{l})$ [2]
    • Production of ions allows the liquid to conduct electricity. [1]

16. $\text{sp}^2$ hybridisation [1]

    • One $\text{s}$ and two $\text{p}$ orbitals mix to form three $\text{sp}^2$ hybrid orbitals. [1]
    • Results in trigonal planar geometry around each C atom. [1]

17. $\sigma$-bond: Head-on overlap of orbitals. [1.5]

    • $\pi$-bond: Sideways overlap of parallel $\text{p}$-orbitals. [1.5]

18. Solid: Be is electron-deficient; forms coordinate bonds with Cl atoms of adjacent molecules to achieve octet. [2]

    • Gas: Thermal energy overcomes these weak intermolecular coordinate bonds, leaving discrete $\text{BeCl}_2$ molecules. [2]

19. $\text{MgO}$ has higher lattice energy [1]

    • $\text{Mg}^{2+}$ and $\text{O}^{2-}$ have higher charges than $\text{Na}^+$ and $\text{Cl}^-$. [1]
    • Stronger electrostatic attraction requires more energy to break. [1]
    • Therefore, $\text{MgO}$ has a higher melting point. [1]

20. Tetrahedral [1]

    • $\text{sp}^3$ hybridisation [1]