DSE Chemistry

Course Overview

The Hong Kong Diploma of Secondary Education (DSE) Chemistry examination assesses candidates' knowledge and understanding of chemical principles, their ability to apply chemistry to real-world situations, and their skills in chemical investigation and analysis.

The course is divided into compulsory topics (Papers 1A and 1B) and elective topics (Paper 2). All candidates must study the compulsory topics and choose ONE elective.

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Examination Structure

Paper	Component	Duration	Weighting
1A	Compulsory Part -- Multiple-Choice Section	1 h 30 m	30%
1B	Compulsory Part -- Structured Questions	2 h	50%
2	Elective -- Structured Questions	1 h	20%

Total: 100%

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Compulsory Topics

The compulsory syllabus covers the following major areas:

I. Atomic Structure and Periodic Table

• Subatomic particles (protons, neutrons, electrons)
• Atomic number, mass number, isotopes
• Electron arrangements and electronic configuration
• Periodic trends: atomic radius, ionisation energy, electronegativity
• Properties of Groups 1, 2, 17, and 18

See ./dse-chemistry-atomic-structure-and-periodic-table for detailed notes.

II. Chemical Bonding

• Ionic bonding: formation, properties, electron transfer
• Covalent bonding: formation, properties, electron sharing, dot diagrams
• Metallic bonding
• Intermolecular forces: van der Waals forces, dipole-dipole interactions, hydrogen bonding
• Structures: giant ionic, simple molecular, giant covalent, metallic

See ./dse-chemistry-chemical-bonding for detailed notes.

III. Stoichiometry and Mole Concept

• Relative atomic and molecular mass
• The mole concept and Avogadro's number
• Molar volume of gas at RTP and STP
• Empirical and molecular formulae
• Percentage composition
• Reacting mass and gas volume calculations
• Concentration (molarity) and titration

See ./dse-chemistry-stoichiometry-and-mole-concept for detailed notes.

IV. Acids, Bases, and Salts

• Properties of acids and bases
• Strong and weak acids
• pH scale and indicators
• Neutralisation reactions
• Salt preparation methods
• Ionic equations
• Common acids and bases

See ./dse-chemistry-acids-bases-and-salts for detailed notes.

V. Rate of Reaction and Energetics

• Factors affecting the rate of reaction
• Collision theory and activation energy
• Exothermic and endothermic reactions
• Enthalpy changes, Hess's law, and bond energies
• Calorimetry

See ./dse-chemistry-rate-of-reaction-and-energetics for detailed notes.

VI. Redox and Electrochemistry

• Oxidation and reduction (OIL RIG)
• Oxidation numbers and redox equations
• Electrochemical cells and reactivity series
• Extraction of metals
• Rusting and corrosion prevention
• Electrolysis (aqueous and molten)

See ./dse-chemistry-redox-and-electrochemistry for detailed notes.

VII. Carbon Chemistry

• Alkanes, alkenes, alkynes and homologous series
• IUPAC nomenclature and isomerism
• Functional groups and reactions
• Polymers and plastics
• Alcohols, carboxylic acids, and esters
• Macromolecules: proteins, starch, cellulose, DNA

See ./dse-chemistry-carbon-chemistry for detailed notes.

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Elective Topics (Choose One)

Elective	Topic
E1	Industrial Chemistry
E2	Analytical Chemistry
E3	Materials Chemistry
E4	Chemical Energy and the Environment

Paper 2 consists of structured questions based on the chosen elective. Each elective carries equal weighting (20% of the total mark).

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Assessment Objectives

Candidates are expected to demonstrate the ability to:

1. Recall and understand chemical facts, terminology, principles, and relationships
2. Apply and analyse chemical knowledge to familiar and unfamiliar situations
3. Evaluate and synthesise information from chemical contexts, form reasoned arguments
4. Plan and carry out chemical investigations, interpret experimental data, and evaluate methods

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Key Quantities and Constants

Quantity / Constant	Symbol	Value / Unit
Avogadro's number	$N_A$	$6.02 \times 10^{23} \mathrm{ mol^{-1}}$
Molar volume (STP)	$V_m$	$22.4 \mathrm{ dm^3/mol}$
Molar volume (RTP)	$V_m$	$24.0 \mathrm{ dm^3/mol}$
Molar gas constant	$R$	$8.31 \mathrm{ J/(mol \cdot K)}$
Planck's constant	$h$	$6.63 \times 10^{-34} \mathrm{ J \cdot s}$
Faraday constant	$F$	$96500 \mathrm{ C/mol}$

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Examination Tips

• Show all working in stoichiometry calculations. Marks are awarded for method even if the final answer is wrong.
• Write balanced chemical equations wherever possible; state symbols are required when specified.
• In ionic equation questions, cancel spectator ions and include state symbols.
• Memorise the colours of common ions and precipitates (e.g., $\mathrm{Cu^{2+}}$ = blue, $\mathrm{Fe^{3+}}$ = yellow-brown, $\mathrm{AgCl}$ = white).
• For organic chemistry questions, clearly draw structural formulae and use correct IUPAC names.
• Pay attention to significant figures in numerical answers.
• In titration calculations, always read the burette to two decimal places.

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Worked Examples

Worked Example 1

A student needs to convert a volume of gas from STP conditions to RTP conditions. If a gas occupies $11.2 \mathrm{ dm^3}$ at STP, how many moles of gas are present?

Solution

At STP, one mole of gas occupies $22.4 \mathrm{ dm^3}$:

$n = \frac{V}{V_m} = \frac{11.2}{22.4} = 0.500 \mathrm{ mol}$

Worked Example 2

In a DSE Paper 1B question, a student is asked to write the ionic equation for the reaction between aqueous calcium chloride and aqueous sodium carbonate. Write the full balanced equation and the net ionic equation.

Solution

Full equation:

$\mathrm{CaCl_2}(aq) + \mathrm{Na_2CO_3}(aq) \to \mathrm{CaCO_3}(s) + 2\mathrm{NaCl}(aq)$

Ionic equation (spectator ions $\mathrm{Na^+}$ and $\mathrm{Cl^-}$ cancel):

$\mathrm{Ca^{2+}}(aq) + \mathrm{CO_3^{2-}}(aq) \to \mathrm{CaCO_3}(s)$

Worked Example 3

A student sets up an electrochemical cell with a $\mathrm{Zn}$ electrode in $\mathrm{ZnSO_4}(aq)$ and a $\mathrm{Cu}$ electrode in $\mathrm{CuSO_4}(aq)$. Identify the anode, cathode, and the direction of electron flow.

Solution

Zinc is higher in the reactivity series than copper, so zinc is more readily oxidised.

• Anode (oxidation): $\mathrm{Zn} \to \mathrm{Zn^{2+}} + 2e^-$
• Cathode (reduction): $\mathrm{Cu^{2+}} + 2e^- \to \mathrm{Cu}$
• Electron flow: from the zinc electrode (anode) through the external wire to the copper electrode (cathode).

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Problem Set

Problem 1: How many significant figures should the answer to the following calculation be given to? $0.150 \times 24.0 / 0.100$

If you get this wrong, revise: Examination Tips — Significant Figures

Solution

The result is $36.0$. All three values have 3 significant figures, so the answer should be given to 3 significant figures: $36.0$ (not $36$ or $36.00$).

Problem 2: State the colour of the precipitate formed when aqueous sodium hydroxide is added to a solution containing $\mathrm{Fe^{3+}}$ ions.

If you get this wrong, revise: Examination Tips — Colours of Common Ions

Solution

A brown precipitate of $\mathrm{Fe(OH)_3}$ is formed.

For reference: $\mathrm{Cu^{2+}}$ gives blue, $\mathrm{Fe^{2+}}$ gives green (turning brown on standing), $\mathrm{Zn^{2+}}$ gives white.

Problem 3: Identify the compulsory topic that each of the following exam questions belongs to:

(a) Explain why graphite conducts electricity. (b) Calculate the percentage yield of a reaction. (c) Write the ionic equation for the neutralisation of a strong acid by a strong base. (d) Explain the trend in first ionisation energy across Period 3.

If you get this wrong, revise: Compulsory Topics Overview

Solution

(a) II. Chemical Bonding — giant covalent structures (b) III. Stoichiometry and Mole Concept — percentage yield (c) IV. Acids, Bases, and Salts — neutralisation and ionic equations (d) I. Atomic Structure and Periodic Table — ionisation energy trends

Problem 4: Using the Faraday constant ($F = 96500 \mathrm{ C/mol}$), calculate the mass of copper deposited when a current of $2.00 \mathrm{ A}$ is passed through $\mathrm{CuSO_4}(aq)$ for 30.0 minutes.

If you get this wrong, revise: VI. Redox and Electrochemistry — Electrolysis

Solution

$Q = It = 2.00 \times 30.0 \times 60 = 3600 \mathrm{ C}$

$n(e^-) = \frac{Q}{F} = \frac{3600}{96500} = 0.0373 \mathrm{ mol}$

$\mathrm{Cu^{2+}} + 2e^- \to \mathrm{Cu}$

$n(\mathrm{Cu}) = \frac{0.0373}{2} = 0.0187 \mathrm{ mol}$

$m(\mathrm{Cu}) = 0.0187 \times 63.5 = 1.18 \mathrm{ g}$

Problem 5: Explain why the first ionisation energy of magnesium is higher than that of sodium, but lower than that of aluminium.

If you get this wrong, revise: I. Atomic Structure and Periodic Table — Ionisation Energy

Solution

Na ($[\mathrm{Ne}]\, 3s^1$): the single $3s$ electron is relatively easy to remove.

Mg ($[\mathrm{Ne}]\, 3s^2$): the $3s$ electrons are closer to the nucleus (higher effective nuclear charge) and experience a full $3s$ subshell effect, making them harder to remove than Na's single $3s$ electron.

Al ($[\mathrm{Ne}]\, 3s^2\, 3p^1$): the $3p$ electron is in a higher energy subshell than $3s$ and is shielded by the $3s^2$ electrons, making it easier to remove than Mg's $3s$ electrons.

So: IE(Na) $\lt$ IE(Mg) $\gt$ IE(Al).

Problem 6: State two differences between a strong acid and a weak acid of the same concentration, and explain each difference.

If you get this wrong, revise: IV. Acids, Bases, and Salts — Strong and Weak Acids

Solution

1. pH: A strong acid has a lower pH than a weak acid at the same concentration. Strong acids dissociate completely in water ($\mathrm{H^+}$ concentration equals the acid concentration), while weak acids only partially dissociate.

2. Electrical conductivity: A strong acid solution is a better conductor than a weak acid solution of the same concentration. Complete dissociation in a strong acid produces more mobile ions, carrying charge more effectively.

Problem 7: A hydrocarbon contains $85.7\%$ carbon and $14.3\%$ hydrogen by mass. Its molar mass is $42.0 \mathrm{ g/mol}$. Determine its molecular formula and identify the homologous series it belongs to.

If you get this wrong, revise: VII. Carbon Chemistry — Homologous Series

Solution

Element	Mass (g)	Moles	Ratio
C	85.7	$85.7/12 = 7.14$	$7.14/7.14 = 1$
H	14.3	$14.3/1 = 14.3$	$14.3/7.14 \approx 2$

Empirical formula: $\mathrm{CH_2}$

$M_r(\mathrm{empirical}) = 12 + 2 = 14$

$n = \frac{42.0}{14} = 3$

Molecular formula: $\mathrm{C_3H_6}$

This is propene, an alkene (homologous series with general formula $\mathrm{C_nH_{2n}}$).

Problem 8: Explain the difference between an exothermic reaction and an endothermic reaction in terms of enthalpy change ($\Delta H$), and give one example of each.

If you get this wrong, revise: V. Rate of Reaction and Energetics — Enthalpy Changes

Solution

Exothermic: $\Delta H \lt 0$ (enthalpy of products is lower than enthalpy of reactants; heat is released to the surroundings). Example: combustion of methane:

$\mathrm{CH_4} + 2\mathrm{O_2} \to \mathrm{CO_2} + 2\mathrm{H_2O} \quad \Delta H = -890 \mathrm{ kJ/mol}$

Endothermic: $\Delta H \gt 0$ (enthalpy of products is higher than enthalpy of reactants; heat is absorbed from the surroundings). Example: thermal decomposition of calcium carbonate:

$\mathrm{CaCO_3} \to \mathrm{CaO} + \mathrm{CO_2} \quad \Delta H = +178 \mathrm{ kJ/mol}$

Problem 9: In the electrolysis of concentrated aqueous sodium chloride, explain why hydrogen is produced at the cathode rather than sodium, and why chlorine is produced at the anode rather than oxygen.

If you get this wrong, revise: VI. Redox and Electrochemistry — Electrolysis (Aqueous)

Solution

Cathode: $\mathrm{Na^+}$ and $\mathrm{H^+}$ (from water) are both present. Although sodium is more reactive, $\mathrm{H^+}$ is preferentially discharged because it has a much less negative discharge potential. The reaction is: $2\mathrm{H_2O} + 2e^- \to \mathrm{H_2} + 2\mathrm{OH^-}$.

Anode: $\mathrm{Cl^-}$ and $\mathrm{OH^-}$ (from water) are both present. In concentrated solution, $\mathrm{Cl^-}$ is preferentially discharged over $\mathrm{OH^-}$ because the overpotential of chlorine is lower at high chloride concentrations. The reaction is: $2\mathrm{Cl^-} \to \mathrm{Cl_2} + 2e^-$.

Problem 10: A student titrates $25.0 \mathrm{ cm^3}$ of ethanoic acid ($\mathrm{CH_3COOH}$) with $0.100 \mathrm{ mol/dm^3}$ sodium hydroxide. The average titre is $16.7 \mathrm{ cm^3}$. Calculate the concentration of the ethanoic acid and state whether it is a strong or weak acid.

If you get this wrong, revise: IV. Acids, Bases, and Salts and III. Stoichiometry — Titration

Solution

$\mathrm{CH_3COOH} + \mathrm{NaOH} \to \mathrm{CH_3COONa} + \mathrm{H_2O}$

Molar ratio: $1 : 1$

$n(\mathrm{NaOH}) = 0.100 \times \frac{16.7}{1000} = 1.67 \times 10^{-3} \mathrm{ mol}$

$n(\mathrm{CH_3COOH}) = 1.67 \times 10^{-3} \mathrm{ mol}$

$c(\mathrm{CH_3COOH}) = \frac{1.67 \times 10^{-3}}{0.0250} = 0.0668 \mathrm{ mol/dm^3}$

Ethanoic acid is a weak acid — it only partially dissociates in aqueous solution, producing a lower concentration of $\mathrm{H^+}$ ions than a strong acid of the same concentration would.