This 2019 Chemistry examination paper covers a broad spectrum of inorganic chemistry, general physical chemistry, and electrochemical principles. Recurring themes evident in the questions include stoichiometry and gas laws, electrolysis and redox reactions, and chemical equilibrium alongside enthalpy changes. Make the most of your practice time by working through the math-based questions independently before looking at the provided solutions, ensuring you truly understand the formulas and stoichiometric ratios involved.
Topics in this paper
The dissolution of common salt in water is a physical change because
A student might mistakenly think a new substance is formed or that heat generation indicates a chemical change. Dissolving salt in water is a physical change because the sodium chloride can be fully recovered by evaporating the water, leaving the original substance chemically unaltered. Common mistake: assuming that processes involving solutions or slight thermal changes must be chemical reactions.
Which of the following substances is a mixture?
One might incorrectly select sulphur powder, distilled water, or ethanol, mistaking them for solutions. Brine solution is a mixture because it has a variable composition consisting of sodium chloride dissolved in water, whereas sulphur powder, distilled water, and ethanol are pure substances. Common mistake: confusing solutions, which are homogeneous mixtures, with chemically pure compounds.
How many moles of oxygen gas are produced from the decomposition of 2.5 moles of potassium chlorate (KClO₃)? [2KClO₃ → 2KCl + 3O₂]
An examinee might mistakenly calculate a 1:1 ratio or misread the stoichiometric coefficients, leading to incorrect values like 2.50 or 7.50 moles. Based on the balanced equation 2KClO₃ → 2KCl + 3O₂ where 2 moles of KClO₃ yield 3 moles of O₂, multiplying 2.5 moles of KClO₃ by the ratio of 3/2 yields 3.75 moles of oxygen gas. Common mistake: failing to apply the correct molar ratio from the balanced reaction coefficients.
A balanced chemical equation obeys the law of
A student might mistakenly choose the law of definite proportions, multiple proportions, or conservation of energy due to confusion over fundamental chemical laws. A balanced chemical equation ensures that the number of atoms for each element remains identical on both sides, which directly obeys the law of conservation of mass. Common mistake: confusing mass conservation in balanced reactions with energy conservation principles.
A gas occupies 1.50 dm³ at 25°C and 1 atm. What volume will it occupy at 100°C and 1 atm?
One might mistakenly use Boyle's Law or invert the temperature ratios, resulting in incorrect volumes such as 0.60 dm³. Applying Charles’s Law (V₁/T₁ = V₂/T₂) requires converting temperatures to Kelvin (T₁ = 298 K, T₂ = 373 K), so calculating V₂ as (1.50 × 373) / 298 gives approximately 1.88 dm³. Common mistake: forgetting to convert Celsius temperatures to Kelvin before applying gas laws.
A mixture contains 80.0 g of oxygen and 56.0 g of nitrogen with a total pressure of 1.8 atm. What is the partial pressure of oxygen? [Molar mass: O₂ = 32 g/mol, N₂ = 28 g/mol]
A test-taker might mistakenly compute mole ratios directly from masses without dividing by molar masses, leading to an incorrect pressure value like 1.4 atm. Calculating the moles gives 2.5 moles of O₂ and 2.0 moles of N₂, resulting in a total of 4.5 moles; multiplying the mole fraction of oxygen (5/9) by the total pressure of 1.8 atm yields a partial pressure of 1.0 atm. Common mistake: using mass grams directly instead of converting to moles for partial pressure calculations.
Which of the following properties is characteristic of an ideal gas at constant temperature?
An examinee could mistakenly think pressure increases with volume or is independent of it, confusing different gas law relationships. For an ideal gas at constant temperature, Boyle’s Law dictates that pressure is inversely proportional to volume (P ∝ 1/V). Common mistake: confusing inverse relationships like Boyle's Law with direct proportionality.
For ionic crystals to sublime on heating, the molecules must acquire energy that is
A student might mistakenly believe that lesser energy or just enough energy to melt the solid is sufficient. Sublimation of ionic crystals requires overcoming the strong electrostatic forces holding the lattice together, thus demanding energy greater than the lattice energy to transition directly from solid to gas. Common mistake: equating the energy required for sublimation with standard melting enthalpy.
An element E has the electronic configuration 1s² 2s² 2p⁶ 3s² 3p³. The reaction of E with a halogen X can produce
One might mistakenly select EX only or EX₃ only by ignoring the element's maximum valence capacity. Element E is phosphorus with a 3s² 3p³ electronic configuration providing 5 valence electrons, allowing it to react with halogens X to form both EX₃ and EX₅ by utilizing either 3 or all 5 of its valence electrons, meaning both options are possible under EX and EX₃ parameters. Common mistake: overlooking higher oxidation states when determining possible halide compounds for group 15 elements.
The atoms ²³⁵U and ²³⁸U are
A test-taker could mistakenly choose isobars, isomers, or anomers by confusing nuclear terminology with structural chemistry terms. ²³⁵U and ²³⁸U are classified as isotopes because they share the same atomic number of 92 but possess different mass numbers due to differing neutron counts. Common mistake: confusing isotopes (same atomic number, different mass numbers) with isobars.
As the difference in electronegativity between bonded atoms increases, the polarity of the bond
An examinee might mistakenly believe that polarity decreases or drops to zero as electronegativity differences grow larger. A larger difference in electronegativity between bonded atoms creates a greater dipole moment, which inherently increases the polarity of the bond. Common mistake: assuming electronegativity differences reduce bond polarity instead of enhancing it.
Which group of elements forms hydrides that are pyramidal in structure?
A student might mistakenly choose Group IV or Group VI because they also contain non-metals that form simple hydrides, leading to confusion about molecular geometry. VSEPR theory dictates that elements in Group V, such as nitrogen and phosphorus, possess a distinct lone pair alongside three bonding pairs of electrons. This specific valence electron arrangement forces the resulting hydrides, like NH₃, into a three-dimensional pyramidal structure. Common mistake: Confusing the molecular shape of Group V hydrides with tetrahedral or linear geometries.
Water has a high boiling point despite its low molecular mass because of
Options like covalent bonding or ionic bonding might seem plausible since they describe intramolecular forces found in many high-melting substances. However, the high boiling point of water despite its relatively small molecular mass is specifically governed by intermolecular hydrogen bonding. This strong attractive force between adjacent molecules requires significantly more thermal energy to separate them into the gas phase. Common mistake: Attributing the high boiling point of water to its internal covalent bonds rather than intermolecular forces.
Argon is used in gas-filled electric lamps because it helps to
A student might guess that argon increases the intensity of light or keeps the filament glowing brightly by reacting with the tungsten. In reality, argon is completely unreactive and serves to prevent the oxidation of the tungsten filament by residual oxygen gas. By fulfilling this inert protective role, the gas successfully extends the operational lifespan of the lamp without altering its illumination. Common mistake: Assuming inert fill gases actively boost bulb brightness instead of protecting the filament from degradation.
The air around a petroleum refinery is most likely to contain
Options containing N₂O or NH₃ might appear correct to a test-taker familiar with general industrial pollutants, but those gases are not typical byproducts of this specific industrial sector. Petroleum refining operations specifically emit sulfur dioxide from the processing of sulfur-containing crude components, carbon monoxide from incomplete combustion processes, and unreacted nitrogen drawn directly from the surrounding air. Common mistake: Grouping all common atmospheric pollutants together rather than identifying the specific emissions tied to petroleum processing.
Water can be identified by the use of
A student could mistakenly select calcium hydroxide or sodium hydroxide because they are common laboratory reagents used in various chemical identification tests. To reliably test for the presence of water, chemists utilize anhydrous copper(II) sulfate, which transitions from white to blue as it absorbs moisture to form the pentahydrate. Common mistake: Confusing standard alkali solutions with specific desiccants or hydration indicators.
The phenomenon whereby sodium carbonate removes temporary hardness of water is known as
Options such as effervescence or hygroscopy might look tempting due to their association with gas evolution or moisture interaction during chemical processes. Sodium carbonate eliminates temporary water hardness by reacting with calcium ions to form an insoluble solid, which is a precipitation reaction represented by the equation Na₂CO₃ + Ca(HCO₃)₂ → CaCO₃ + 2NaHCO₃. Common mistake: Mistaking precipitation reactions for physical phenomena like effervescence or efflorescence.
A student prepares 0.5 M solutions of hydrochloric acid and ethanoic acid and measures their pH. The result shows that
A student might mistakenly select equal pH values because both acids share the exact same initial molar concentration of 0.5 M. Hydrochloric acid is a strong acid that fully dissociates to yield a high hydrogen ion concentration and a low pH of approximately 0.3, whereas ethanoic acid is a weak acid that only partially dissociates, resulting in a much higher pH near 2.5. Common mistake: Assuming solutions of equal concentration possess identical pH values regardless of acid strength.
The solubility of which of the following salts increases most rapidly with temperature? [Solubility data at 0°C and 100°C: CaSO₄: 0.21 g to 0.24 g; NaCl: 35.7 g to 39.1 g; KCl: 27.6 g to 56.7 g; KNO₃: 13.3 g to 247 g per 100 g of water]
Options like NaCl or KCl might be chosen because their solubility also rises with heating, making them seem like strong candidates for rapid dissolution changes. Examining the data reveals that potassium nitrate skyrockets from 13.3 g to 247 g per 100 g of water, representing an enormous 1857% increase that vastly surpasses the modest percentage gains of the other listed salts. Common mistake: Comparing absolute mass changes rather than evaluating the percentage increase relative to the starting solubility.
In the reaction NH₃ + H₂O ⇌ NH₄⁺ + OH⁻, what role does H₂O play?
A student might view water as a redox agent or oxidizing agent because water is frequently involved in electron-transfer electrochemical reactions. In this specific proton-transfer equilibrium, water donates a hydrogen ion to ammonia to generate the ammonium ion, fulfilling the exact definition of a Bronsted-Lowry acid. Common mistake: Classifying water as a redox reagent in acid-base proton transfer equilibria.
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