A Level H1 Chemistry Atomic Structure Bonding Quiz

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A-Level Chemistry H1 Quiz - Atomic Structure Bonding

Name: ____________________ Class: ____________________ Date: ____________________ Score: ________ / 45

Duration: 60 Minutes
Total Marks: 45
Instructions: Answer all questions. Use a calculator where necessary. Show all working for calculation questions.

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Section A: Atomic Structure & Basic Bonding (Questions 1-7)

1. Complete the table below for the following species. [3] | Species | Nucleon Number | Atomic Number | Protons | Neutrons | Electrons | | :--- | :---: | :---: | :---: | :---: | :---: | | $^{31}\text{P}^{3-}$ | | | | | |

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2. Define the term isotope and explain why isotopes of the same element exhibit identical chemical properties. [2]

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3. State and describe the structure and bonding of solid magnesium. [2]

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4. Write the full electronic configuration for a $\text{Cu}^+$ ion. [1]

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

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6. Describe the difference between a $\sigma$-bond and a $\pi$-bond in terms of orbital overlap. [2]

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7. State the shape and the bond angle of a molecule of $\text{BF}_3$. [2]

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Section B: Molecular Geometry & Intermolecular Forces (Questions 8-14)

8. For the molecule $\text{PCl}_5$: (a) State the geometry of the molecule. [1] (b) Explain why the axial bonds are slightly longer than the equatorial bonds. [2]

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9. 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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10. Draw the dot-and-cross diagram for the $\text{I}_3^-$ ion, including all lone pairs and the overall charge. [3]

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11. Explain why $\text{CCl}_4$ is a non-polar molecule even though the $\text{C}-\text{Cl}$ bond is polar. [2]

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12. Predict the shape and bond angle of the $\text{NH}_4^+$ ion. [2]

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13. Which of the following substances would have the highest melting point? Justify your answer by comparing the types of bonding/forces involved. $\text{NaCl}$, $\text{Cl}_2$, $\text{H}_2\text{O}$ [3]

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14. Describe the nature of the bonding in a molecule of $\text{CO}_2$. [2]

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Section C: Advanced Bonding & Application (Questions 15-20)

15. Boron trifluoride ($\text{BF}_3$) and ammonia ($\text{NH}_3$) react in a 1:1 ratio to form a white crystalline compound. (a) Draw a diagram to illustrate the bonding in the product. [2] (b) Name and explain the type of bond formed. [2]

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16. Explain why graphite is a good conductor of electricity while diamond is not, referring to their structures. [3]

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17. Define electronegativity and explain its trend across Period 3 from $\text{Na}$ to $\text{Cl}$. [3]

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18. A compound has the formula $\text{XeF}_2$. (a) State the number of lone pairs on the central $\text{Xe}$ atom. [1] (b) Predict the shape of the molecule. [1]

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19. Explain why the melting point of $\text{MgO}$ is significantly higher than that of $\text{NaF}$. [3]

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20. Describe the structure and bonding in a sample of solid sodium. [2]

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Answer Key - A-Level Chemistry H1 Quiz: Atomic Structure Bonding

1. Table Completion [3]

   • Nucleon Number: 31
   • Atomic Number: 15
   • Protons: 15
   • Neutrons: 16 (31 - 15)
   • Electrons: 18 (15 + 3)

2. Isotopes [2]

   • Definition: Atoms of the same element with the same number of protons but different number of neutrons. [1]
   • Reason: Chemical properties are determined by the electronic configuration, which is identical for isotopes of the same element. [1]

3. Solid Magnesium [2]

   • Structure: Giant metallic structure. [1]
   • Bonding: $\text{Mg}^{2+}$ cations held together by a sea of delocalized valence electrons by strong electrostatic forces. [1]

4. Electronic Configuration [1]

   • $\text{Cu}^+$: $1\text{s}^2 2\text{s}^2 2\text{p}^6 3\text{s}^2 3\text{p}^6 3\text{d}^{10}$ (or $[\text{Ar}] 3\text{d}^{10}$) [1]

5. Ionisation Energy [2]

   • In Al, the electron is removed from a $3\text{p}$ orbital, whereas in Mg, it is removed from a $3\text{s}$ orbital. [1]
   • The $3\text{p}$ electron is higher in energy/further from the nucleus/more shielded, making it easier to remove. [1]

6. $\sigma$ vs $\pi$ bonds [2]

   • $\sigma$-bond: Formed by head-on (end-to-end) overlap of orbitals. [1]
   • $\pi$-bond: Formed by sideways overlap of p-orbitals. [1]

7. $\text{BF}_3$ [2]

   • Shape: Trigonal planar [1]
   • Bond angle: $120^\circ$ [1]

8. $\text{PCl}_5$ [3]

   • (a) Trigonal bipyramidal [1]
   • (b) Axial bonds experience greater repulsion from the equatorial bonds (90° vs 120°), pushing them further away and lengthening the bond. [2]

9. $\text{H}_2\text{O}$ vs $\text{H}_2\text{S}$ [3]

   • $\text{H}_2\text{O}$ has a higher boiling point. [1]
   • $\text{H}_2\text{O}$ exhibits strong hydrogen bonding (due to high electronegativity of O). [1]
   • $\text{H}_2\text{S}$ only exhibits permanent dipole-dipole and van der Waals forces, which are weaker. [1]

10. $\text{I}_3^-$ Diagram [3]

    • Linear arrangement of 3 Iodine atoms. [1]
    • Central I has 3 lone pairs and 2 bonding pairs. Terminal I atoms have 3 lone pairs. [1]
    • Correct brackets and $-1$ charge. [1]

11. $\text{CCl}_4$ Polarity [2]

    • The molecule is tetrahedrally symmetrical. [1]
    • The individual $\text{C}-\text{Cl}$ bond dipoles cancel each other out. [1]

12. $\text{NH}_4^+$ [2]

    • Shape: Tetrahedral [1]
    • Bond angle: $109.5^\circ$ [1]

13. Melting Point Comparison [3]

    • $\text{NaCl}$ has the highest melting point. [1]
    • $\text{NaCl}$ has a giant ionic lattice with strong electrostatic attractions between $\text{Na}^+$ and $\text{Cl}^-$. [1]
    • $\text{Cl}_2$ (van der Waals) and $\text{H}_2\text{O}$ (Hydrogen bonding) have much weaker intermolecular forces. [1]

14. $\text{CO}_2$ Bonding [2]

    • Covalent bonding. [1]
    • Each $\text{C}=\text{O}$ bond consists of one $\sigma$-bond and one $\pi$-bond. [1]

15. $\text{BF}_3 + \text{NH}_3$ [4]

    • (a) Diagram showing arrow from $\text{N}$ lone pair to $\text{B}$ center. [2]
    • (b) Coordinate covalent bond (or dative bond). [1] Nitrogen donates both electrons to the empty orbital of Boron. [1]

16. Graphite vs Diamond [3]

    • Graphite: Each C is bonded to 3 others in layers; one delocalized electron per C is free to move and conduct electricity. [2]
    • Diamond: Each C is bonded to 4 others in a rigid 3D lattice; no delocalized electrons. [1]

17. Electronegativity [3]

    • Definition: The ability of an atom to attract the bonding pair of electrons in a covalent bond. [1]
    • Trend: Increases across Period 3. [1]
    • Reason: Nuclear charge increases while shielding remains constant, increasing the attraction for electrons. [1]

18. $\text{XeF}_2$ [2]

    • (a) 3 lone pairs [1]
    • (b) Linear [1]

19. $\text{MgO}$ vs $\text{NaF}$ [3]

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

20. Solid Sodium [2]

    • Structure: Giant metallic structure. [1]
    • Bonding: $\text{Na}^+$ cations in a regular lattice surrounded by a sea of delocalized valence electrons. [1]