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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: 45 minutes
Total Marks: 45

Instructions:

Answer all questions.
Write your answers in the spaces provided.
The number of marks is given in brackets [ ] at the end of each question or part question.
You may use a scientific calculator.
A Data Booklet is provided for reference (standard constants and periodic table).

Section A: Atomic Structure and Subatomic Particles (Questions 1–5)

1. Complete the following table for the species given. [3]

Species	Proton Number	Nucleon Number	Number of Neutrons	Number of Electrons
$^{27}\text{Al}^{3+}$
$^{18}\text{O}^{2-}$
$^{35}\text{Cl}$

2. Define the term first ionisation energy. [2]

3. The first seven ionisation energies of an element X are shown below.

Ionisation Number	1st	2nd	3rd	4th	5th	6th	7th
Energy / kJ mol⁻¹	1012	1903	2912	4957	6274	21267	25397

(a) Identify the group in the Periodic Table to which element X belongs. [1]

(b) Explain your answer to (a) by referring to the data above. [2]

4. Write the electronic configuration, using s, p, and d notation, for: [2]

(a) A neutral nitrogen atom (N).

(b) A copper(II) ion ( $\text{Cu}^{2+}$ ).

5. Sketch the shape of a $p$ -orbital and a $d_{z^2}$ -orbital. [2]

Answer space

Section B: Chemical Bonding and Molecular Shape (Questions 6–12)

6. Draw a 'dot-and-cross' diagram to show the bonding in magnesium chloride ( $\text{MgCl}_2$ ). Show outer electrons only. [2]

Answer space

7. Boron trifluoride ( $\text{BF}_3$ ) reacts with ammonia ( $\text{NH}_3$ ) to form an adduct, $\text{F}_3\text{B-NH}_3$ .

(a) Name the type of bond formed between the boron atom and the nitrogen atom in the adduct. [1]

(b) Explain how this bond is formed, identifying the electron donor and acceptor. [2]

8. Use the Valence Shell Electron Pair Repulsion (VSEPR) theory to deduce the shape and bond angle of the following species. [4]

Species	Shape	Bond Angle
$\text{BeCl}_2$
$\text{NH}_4^+$

9. Explain why the bond angle in ammonia ( $\text{NH}_3$ ) is $107^\circ$ , whereas the bond angle in the ammonium ion ( $\text{NH}_4^+$ ) is $109.5^\circ$ . [3]

10. Sulfur hexafluoride ( $\text{SF}_6$ ) is a non-polar molecule, whereas sulfur tetrafluoride ( $\text{SF}_4$ ) is polar.

(a) Draw the shape of $\text{SF}_6$ . [1]

Answer space

(b) Explain why $\text{SF}_6$ is non-polar despite having polar S–F bonds. [2]

11. Describe the structure and bonding in graphite. Explain why graphite conducts electricity but diamond does not. [4]

12. Silicon(IV) oxide ( $\text{SiO}_2$ ) and carbon dioxide ( $\text{CO}_2$ ) are both Group 14 oxides.

(a) State the type of structure and bonding in $\text{SiO}_2$ . [1]

(b) Explain why $\text{SiO}_2$ has a much higher melting point than $\text{CO}_2$ . [2]

Section C: Intermolecular Forces and Physical Properties (Questions 13–20)

13. Name the strongest intermolecular force present between molecules of: [2]

(a) Hydrogen chloride ( $\text{HCl}$ ).

(b) Methanol ( $\text{CH}_3\text{OH}$ ).

14. Explain why water ( $\text{H}_2\text{O}$ ) has a higher boiling point than hydrogen sulfide ( $\text{H}_2\text{S}$ ), even though $\text{H}_2\text{S}$ has a higher relative molecular mass. [3]

15. The table below shows the boiling points of three alkanes.

Alkane	Formula	Boiling Point / $^\circ\text{C}$
Propane	$\text{C}_3\text{H}_8$	-42
Butane	$\text{C}_4\text{H}_{10}$	-0.5
2-methylpropane	$\text{CH}(\text{CH}_3)_3$	-11.7

(a) Explain the trend in boiling points between propane and butane. [2]

(b) Explain why butane has a higher boiling point than its isomer, 2-methylpropane. [2]

16. Ethanol ( $\text{C}_2\text{H}_5\text{OH}$ ) is miscible with water, but hexanol ( $\text{C}_6\text{H}_{13}\text{OH}$ ) is only slightly soluble. Explain this difference in solubility. [3]

17. Iodine ( $\text{I}_2$ ) is a solid at room temperature, while chlorine ( $\text{Cl}_2$ ) is a gas. Explain this difference in physical state in terms of intermolecular forces. [2]

18. Which of the following compounds can form hydrogen bonds with water? Circle all that apply. [1]

$\text{CH}_4$ $\quad$ $\text{CH}_3\text{NH}_2$ $\quad$ $\text{CH}_3\text{OCH}_3$ $\quad$ $\text{CH}_3\text{Cl}$

19. Explain why ice is less dense than liquid water. [2]

20. A student attempts to dissolve sodium chloride in hexane ( $\text{C}_6\text{H}_{14}$ ). It does not dissolve. Explain why, referring to the energy changes involved in the solution process. [3]

End of Quiz

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Answers

A-Level Chemistry H1 Quiz - Atomic Structure Bonding (Answer Key)

1. [3 marks]

$^{27}\text{Al}^{3+}$ : Protons: 13, Nucleon: 27, Neutrons: 14, Electrons: 10
$^{18}\text{O}^{2-}$ : Protons: 8, Nucleon: 18, Neutrons: 10, Electrons: 10
$^{35}\text{Cl}$ : Protons: 17, Nucleon: 35, Neutrons: 18, Electrons: 17 (1 mark for each correct row)

2. [2 marks] The energy required [1] to remove one mole of electrons from one mole of gaseous atoms to form one mole of gaseous 1+ ions [1]. (Must specify gaseous state for atoms and ions)

3. (a) Group 15 (or V) [1] (b) There is a large jump in ionisation energy between the 5th and 6th electrons [1]. This indicates that the 6th electron is being removed from an inner shell closer to the nucleus / with less shielding [1]. Therefore, there are 5 valence electrons.

4. (a) $1s^2 2s^2 2p^3$ [1] (b) $[Ar] 3d^9$ or $1s^2 2s^2 2p^6 3s^2 3p^6 3d^9$ [1] (Note: Cu is $[Ar] 3d^{10} 4s^1$ , so $\text{Cu}^{2+}$ loses the 4s and one 3d electron)

5. [2 marks]

$p$ -orbital: Dumbbell shape / two lobes [1].
$d_{z^2}$ -orbital: Distinctive shape with two lobes along z-axis and a ring/doughnut in xy-plane [1].

6. [2 marks]

Correct ions: $\text{Mg}^{2+}$ and two $\text{Cl}^-$ [1].
Correct electron transfer shown (Mg empty outer shell, Cl full outer shell with 8 electrons including cross/dot distinction) [1].

7. (a) Dative covalent bond / Coordinate bond [1]. (b) The nitrogen atom in ammonia has a lone pair of electrons [1]. The boron atom in $\text{BF}_3$ is electron deficient (has an empty orbital) [1]. Nitrogen donates both electrons to form the bond.

8. [4 marks]

$\text{BeCl}_2$ : Linear [1], $180^\circ$ [1].
$\text{NH}_4^+$ : Tetrahedral [1], $109.5^\circ$ [1].

9. [3 marks]

$\text{NH}_3$ has 3 bond pairs and 1 lone pair [1].
$\text{NH}_4^+$ has 4 bond pairs and 0 lone pairs [1].
Lone pair-bond pair repulsion is greater than bond pair-bond pair repulsion, compressing the angle in $\text{NH}_3$ [1].

10. (a) Octahedral [1]. (b) The S–F bonds are polar due to electronegativity difference [1]. However, the symmetrical octahedral shape causes the individual bond dipoles to cancel out vectorially, resulting in no net dipole moment [1].

11. [4 marks]

Graphite has a giant covalent structure consisting of layers of hexagonal rings [1].
Each carbon is bonded to 3 others; the fourth electron is delocalised between layers [1].
These delocalised electrons can move and carry charge, allowing conductivity [1].
In diamond, each carbon is bonded to 4 others in a tetrahedral structure with no delocalised electrons [1].

12. (a) Giant covalent / Giant molecular structure [1]. (b) $\text{SiO}_2$ has strong covalent bonds throughout the lattice that require large amounts of energy to break [1]. $\text{CO}_2$ consists of simple molecules held by weak intermolecular forces (van der Waals) which require little energy to overcome [1].

13. (a) Permanent dipole-dipole forces [1]. (b) Hydrogen bonding [1].

14. [3 marks]

Water molecules form hydrogen bonds [1].
$\text{H}_2\text{S}$ molecules only have permanent dipole-dipole / van der Waals forces [1].
Hydrogen bonds are significantly stronger than van der Waals forces, requiring more energy to break [1].

15. (a) Butane has a larger molecular size / more electrons than propane [1]. This leads to stronger instantaneous dipole-induced dipole (van der Waals) forces [1]. (b) Butane has a linear shape allowing for greater surface area contact between molecules compared to the spherical/branched 2-methylpropane [1]. Greater surface area leads to stronger van der Waals forces [1].

16. [3 marks]

Both alcohols have an -OH group that can form hydrogen bonds with water [1].
Ethanol has a short hydrocarbon chain, so the polar -OH group dominates, making it miscible [1].
Hexanol has a long non-polar hydrocarbon chain which disrupts the hydrogen bonding network of water / the hydrophobic effect dominates, reducing solubility [1].

17. [2 marks]

Iodine has more electrons than chlorine [1].
This results in stronger van der Waals forces between $\text{I}_2$ molecules, sufficient to hold them in a solid lattice at room temperature [1].

18. [1 mark] $\text{CH}_3\text{NH}_2$ (Primary amine can H-bond). (Note: $\text{CH}_3\text{OCH}_3$ can accept H-bonds from water but cannot donate; usually considered soluble but $\text{CH}_3\text{NH}_2$ is the stronger H-bonder. If multiple choice allows multiple, ether is also acceptable as soluble, but amine is the classic H-bond example. Given "Circle all", $\text{CH}_3\text{NH}_2$ is the definite yes. $\text{CH}_3\text{OCH}_3$ is also capable of H-bonding with water as an acceptor. Award mark for $\text{CH}_3\text{NH}_2$ . If student circles ether too, accept if they justify acceptor capability, but strictly "form hydrogen bonds" usually implies the full interaction. For H1, Amine is the key answer.) Correction for H1 Level: Both $\text{CH}_3\text{NH}_2$ and $\text{CH}_3\text{OCH}_3$ can form hydrogen bonds with water (Amine: donor/acceptor; Ether: acceptor). $\text{CH}_4$ and $\text{CH}_3\text{Cl}$ cannot. Award mark if $\text{CH}_3\text{NH}_2$ is circled.

19. [2 marks]

In ice, water molecules form an open tetrahedral lattice held by hydrogen bonds [1].
This structure holds molecules further apart than in liquid water, resulting in lower density [1].

20. [3 marks]

Energy is required to break the strong ionic lattice of NaCl (Lattice Energy) [1].
Energy is required to overcome van der Waals forces in hexane [1].
The energy released when ions interact with hexane (solvation) is very small because hexane is non-polar. The energy required is not compensated, so dissolution does not occur [1].