A Level H2 Chemistry Atomic Structure Bonding Quiz

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A-Level Chemistry H2 Quiz - Atomic Structure Bonding

Name: ____________________ Class: __________ Date: __________ Score: ________ / 50

Duration: 60 Minutes
Total Marks: 50 Marks

Instructions:

• Answer all questions in the spaces provided.
• Use of the Data Booklet is permitted and required for specific questions.
• Show all working for calculations.
• Ensure all curly arrows in mechanisms are drawn clearly from lone pairs/bonds to electrophilic centres.

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Section A: Atomic Structure & Periodicity (Questions 1–7)

1. An ion $\text{Y}^{3+}$ has 36 electrons and 48 neutrons. Identify element $\text{Y}$ and write its full electron configuration. [3]
   
   
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2. Explain why the first ionisation energy of Magnesium is higher than that of Aluminium, despite Aluminium having a higher nuclear charge. [2]
   
   
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3. Compare the first ionisation energies of Nitrogen and Oxygen. Explain the observed trend with reference to electron configuration. [3]
   
   
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4. Define the term first ionisation energy. [2]
   
   
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5. An element $\text{Z}$ in Period 3 has a second ionisation energy significantly higher than its first. Suggest the Group of element $\text{Z}$. Justify your answer. [3]
   
   
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6. Explain why the atomic radius of Potassium is larger than that of Calcium. [2]
   
   
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7. Write the electron configuration of the $\text{Cu}^{2+}$ ion in the $\text{[Cu(H}_2\text{O)}_6]^{2+}$ complex. [2]
   
   
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Section B: Chemical Bonding & Molecular Geometry (Questions 8–14)

8. Using VSEPR theory, predict the shape and the bond angle of $\text{SF}_6$. [2]
   
   
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9. Explain why $\text{BF}_3$ is a non-polar molecule despite having polar $\text{B}-\text{F}$ bonds. [2]
   
   
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10. Compare the boiling points of $\text{H}_2\text{O}$ and $\text{H}_2\text{S}$. Explain the difference in terms of intermolecular forces. [3]
    
    
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11. Describe the bonding in a metallic lattice. Why are metals typically good conductors of electricity? [3]
    
    
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12. Draw the Lewis structure of the $\text{CO}_3^{2-}$ ion. Indicate all formal charges. [2]
    
    
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13. $\text{PCl}_5$ exists as a trigonal bipyramidal molecule in the gas phase. Explain why the axial $\text{P}-\text{Cl}$ bonds are slightly longer than the equatorial $\text{P}-\text{Cl}$ bonds. [3]
    
    
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14. Explain why $\text{CCl}_4$ does not exhibit hydrogen bonding, whereas $\text{CH}_4$ does not either, but $\text{CCl}_4$ has a significantly higher boiling point than $\text{CH}_4$. [3]
    
    
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Section C: Advanced Bonding & Applications (Questions 15–20)

15. $\text{BrF}_3$ is a covalent compound that conducts electricity in the liquid state. Suggest an equation for its auto-ionisation and explain how this leads to conductivity. [3]
    
    
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16. Compare the lattice energy of $\text{NaCl}$ and $\text{MgO}$. Explain which compound has a higher melting point and why. [3]
    
    
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17. Using the concept of orbital overlap, explain why the $\text{C}-\text{C}$ bond in ethene is shorter and stronger than the $\text{C}-\text{C}$ bond in ethane. [3]
    
    
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18. Predict the shape of the $\text{I}_3^-$ ion. Justify your answer using the number of bonding pairs and lone pairs on the central atom. [3]
    
    
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19. Explain the difference between a $\sigma$-bond and a $\pi$-bond in terms of the region of electron density. [2]
    
    
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20. An unknown compound $\text{X}_2\text{O}$ is amphoteric. Write an ionic equation to show its reaction with hot aqueous sodium hydroxide. [2]
    
    
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Answer Key - A-Level Chemistry H2 Quiz: Atomic Structure Bonding

1. Identity: $\text{Y}$ is Rubidium ($\text{Rb}$). Calculation: Protons = electrons + charge = $36 + 3 = 39$. Atomic number 39 is $\text{Y}$ (Yttrium). Correction: Atomic number 39 is Yttrium. Configuration: $1\text{s}^2 2\text{s}^2 2\text{p}^6 3\text{s}^2 3\text{p}^6 3\text{d}^{10} 4\text{s}^2 4\text{p}^6 4\text{d}^1$ (for $\text{Y}$). For $\text{Y}^{3+}$, remove $4\text{d}^1$ and $4\text{s}^2 \rightarrow [ \text{Kr} ]$. Marks: 1 for element, 2 for configuration.

2. Mg has a stable $3\text{s}^2$ configuration. Al has a $3\text{p}^1$ electron which is further from the nucleus and more shielded by the $3\text{s}^2$ electrons, making it easier to remove. [2]

3. Nitrogen has a half-filled $2\text{p}^3$ subshell, which is relatively stable. Oxygen has $2\text{p}^4$; the repulsion between the two electrons in the same p-orbital makes it easier to remove the first electron. Thus, $\text{IE}_1(\text{N}) > \text{IE}_1(\text{O})$. [3]

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

5. Group: Group 1. Justification: The first electron is removed from the valence shell ($s^1$). The second electron must be removed from a complete inner shell (noble gas configuration), which is much closer to the nucleus and experiences much less shielding, resulting in a massive jump in energy. [3]

6. Potassium has one more principal energy level (shell) than Calcium's core, but specifically, Calcium has a higher nuclear charge which pulls the valence electrons closer to the nucleus, reducing the atomic radius. [2]

7. $\text{Cu}$ is $[ \text{Ar} ] 3\text{d}^{10} 4\text{s}^1$. $\text{Cu}^{2+}$ is $[ \text{Ar} ] 3\text{d}^9$. [2]

8. Shape: Octahedral. Angle: $90^\circ$. [2]

9. $\text{BF}_3$ has a trigonal planar geometry. The three polar $\text{B}-\text{F}$ bond dipoles cancel each other out due to the symmetrical arrangement, resulting in a net dipole moment of zero. [2]

10. $\text{H}_2\text{O}$ has hydrogen bonding (strongest IMF) due to the high electronegativity difference between $\text{O}$ and $\text{H}$. $\text{H}_2\text{S}$ only has permanent dipole-dipole and London forces. $\text{H}_2\text{O}$ requires more energy to overcome these forces, hence a higher boiling point. [3]

11. Bonding: A lattice of positive metal ions surrounded by a "sea" of delocalised valence electrons. Conductivity: These delocalised electrons are free to move through the lattice when a potential difference is applied. [3]

12. [Structure: Central $\text{C}$ with three $\text{O}$ atoms. One $\text{C}=\text{O}$ double bond, two $\text{C}-\text{O}^-$ single bonds. Resonance arrows indicated]. [2]

13. The axial positions experience more repulsion from the equatorial bonding pairs than the equatorial positions do from each other. To minimize repulsion, the axial bonds lengthen. [3]

14. Neither has $\text{H}-\text{O}$, $\text{H}-\text{N}$, or $\text{H}-\text{F}$ bonds, so no H-bonding. $\text{CCl}_4$ is a larger molecule with more electrons than $\text{CH}_4$, leading to stronger London dispersion forces. [3]

15. Equation: $2\text{BrF}_3 \rightleftharpoons \text{BrF}_2^+ + \text{BrF}_4^-$ Explanation: The auto-ionisation produces mobile ions in the liquid state, which can carry an electric current. [3]

16. $\text{MgO}$ has higher lattice energy. $\text{Mg}^{2+}$ and $\text{O}^{2-}$ have higher charges than $\text{Na}^+$ and $\text{Cl}^-$. Stronger electrostatic attraction requires more energy to break, leading to a higher melting point. [3]

17. Ethane has a $\text{C}-\text{C}$ $\sigma$-bond (head-on overlap). Ethene has a $\sigma$-bond and a $\pi$-bond (side-on overlap of p-orbitals). The $\pi$-bond pulls the nuclei closer together, increasing bond strength and shortening the length. [3]

18. Shape: Linear. Justification: Central $\text{I}$ has 5 valence electrons + 2 from other $\text{I}$ atoms + 1 from charge = 8 electrons (4 pairs). 2 bonding pairs and 2 lone pairs $\rightarrow$ linear geometry (lone pairs occupy equatorial positions). [3]

19. $\sigma$-bond: Electron density is concentrated along the internuclear axis. $\pi$-bond: Electron density is concentrated in two lobes above and below the internuclear axis. [2]

20. $\text{X}_2\text{O}(\text{s}) + 2\text{OH}^-(\text{aq}) + \text{H}_2\text{O}(\text{l}) \rightarrow 2[\text{XO}_2](\text{aq})^- + 2\text{H}_2\text{O}$ (or similar based on $\text{Al}_2\text{O}_3$ pattern). [2]