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.
What is the molarity of a solution containing 4.0 g of sodium hydroxide in 250 cm³ of solution? [Molar mass of NaOH = 40 g/mol]
A test-taker could mistakenly select 0.10 mol/dm³ by calculating the moles of solute without dividing by the solution volume in cubic decimeters. To find the correct molarity, one must first determine that 4.0 g of sodium hydroxide equals 0.1 moles, and then divide that value by the volume of 0.25 dm³ to arrive at 0.40 mol/dm³. Common mistake: Forgetting to divide the calculated number of moles by the volume in liters or cubic decimeters.
During the electrolysis of a salt of metal M, a current of 0.05 A flows for 32 minutes 10 seconds, depositing 0.107 g of M. What is the charge of the metal ion? [Molar mass of M = 27 g/mol, F = 96500 C/mol]
Options like 1+ or 3+ might be selected through hasty division errors during the Faraday constant calculation steps. By applying Faraday's laws of electrolysis, the total charge is found to be 96.5 C based on the current and time, and combining this with the deposited mass and molar mass yields a calculated valence of approximately 2. Common mistake: Incorrectly converting time from minutes and seconds into total seconds before computing the charge.
Which of the following reactions occurs at the anode during the electrolysis of a dilute aqueous solution of sodium chloride?
A student might choose chloride ion oxidation because the solution involves sodium chloride, overlooking the fact that the solution is dilute. Because the concentration of chloride ions is extremely low in a dilute setup, water undergoes oxidation at the anode instead, producing oxygen gas alongside hydrogen ions and electrons. Common mistake: Assuming the solute always undergoes discharge at the electrode regardless of its concentration in solution.
Given the standard electrode potentials: Cu²⁺/Cu = +0.34 V, Fe²⁺/Fe = -0.44 V, Br₂/2Br⁻ = +1.07 V, Zn²⁺/Zn = -0.76 V, which pair will react spontaneously?
Options like Cu²⁺ + Fe might look plausible to a student who misreads standard reduction potentials or confuses cathode and anode assignments. A chemical reaction proceeds spontaneously only when the calculated standard cell potential is positive, which occurs for the pairing of bromine and zinc where E°cell equals 1.07 minus negative 0.76, resulting in a favorable 1.83 V. Common mistake: Subtracting standard electrode potential values in the wrong order when calculating cell potential.
The oxidation states of chlorine in HOCl, HClO₂, and HClO₄ are respectively
Options featuring negative oxidation numbers for chlorine might be picked by students who confuse binary chlorides with complex oxyacids. In these specific oxygen-containing compounds, assigning standard oxidation numbers of -2 to oxygen and +1 to hydrogen allows us to solve for chlorine, yielding +1 in HOCl, +3 in HClO₂, and +7 in HClO₄. Common mistake: Inverting the algebraic signs of oxidation states for halogens bonded to oxygen.
A reaction occurs spontaneously if
A student could mistakenly select options involving positive Gibbs free energy values because they confuse thermodynamic favorability with unfavorable endothermic conditions. For any process to occur spontaneously, the change in Gibbs free energy must be less than zero, which mathematically requires the enthalpy change to be less than the temperature multiplied by the entropy change according to the governing equation. Common mistake: Believing that positive Gibbs free energy values indicate spontaneous processes.
Given the standard enthalpies of formation: CO(g) = -110 kJ/mol, H₂O(g) = -242 kJ/mol, CO₂(g) = -394 kJ/mol, what is the enthalpy change for CO(g) + H₂O(g) → CO₂(g) + H₂(g)?
Options like +42 kJ/mol might be chosen by a student who reverses the subtraction order between reactant and product enthalpy values. The enthalpy change is calculated by taking the sum of the standard enthalpies of formation of the products and subtracting the sum of the standard enthalpies of the reactants, yielding -394 minus the sum of -110 and -242, which results in -42 kJ/mol. Common mistake: Swapping the reactant and product terms when applying the summation formula for enthalpy of formation.
For the reaction 2A(g) ⇌ B(g), which condition would shift the equilibrium to the right?
A student might incorrectly select an increase in pressure because pressure changes frequently shift gaseous equilibria, though in this case it would shift the system left toward fewer moles. According to Le Chatelier's principle, adding more of reactant A forces the system to consume the excess by shifting the equilibrium position directly to the right toward product B. Common mistake: Applying pressure shifts to reactions without verifying the stoichiometric mole ratios on both sides.
10 g of a solid is in equilibrium with its own vapor at 25°C. If 1 g of the solid is added, the vapor pressure will
Options indicating an increase or decrease in vapor pressure might be selected under the false assumption that adding more solid reactant alters the physical equilibrium state. As long as some solid remains present in the sealed container, the vapor pressure of a solid at a constant temperature remains entirely independent of the total mass of the solid. Common mistake: Assuming that adding more of a pure solid or liquid changes its equilibrium vapor pressure.
In the reaction NaCl(s) + H₂SO₄(l) → HCl(g) + NaHSO₄(s), H₂SO₄ behaves as
A student might guess that sulfuric acid acts as an oxidizing agent or dehydrating agent due to its well-known strong chemical properties in other contexts. In this particular displacement reaction, concentrated sulfuric acid functions strictly as a strong acid by supplying the hydrogen ions necessary to release hydrogen chloride gas from sodium chloride. Common mistake: Overlooking the specific acid-base role of sulfuric acid in halide salt reactions in favor of its oxidizing properties.
Which of the following oxides will NOT produce oxygen gas when heated?
Options like silver nitrate or calcium nitrate might appear plausible if a student thinks all metal compounds behave identically upon thermal decomposition. Unlike nitrates that thermally decompose to release oxygen gas, sodium carbonate breaks down strictly into sodium oxide and carbon dioxide without liberating any elemental oxygen. Common mistake: Assuming all metal salts release oxygen gas when subjected to high temperatures.
To obtain pure carbon monoxide (CO) from a gaseous mixture containing CO and CO₂, the mixture should be
A student might suggest passing the mixture over heated copper(II) oxide to remove carbon dioxide, but that setup is typically used for reducing copper or oxidizing carbon monoxide. Bubbling the gas mixture through a sodium hydroxide solution selectively traps and removes carbon dioxide by converting it into sodium carbonate, leaving pure carbon monoxide gas untouched. Common mistake: Selecting reducing agents instead of acid-base scrubbers for acidic gas removal.
Which of the following oxides will NOT produce its metal, oxygen, and nitrogen gas when heated?
Options like silver nitrate or lithium nitrate might be picked by test-takers who fail to recall the unique thermal stability and decomposition pathways of group 1 alkali metal nitrates. Sodium nitrate decomposes upon heating to form sodium nitrite and oxygen gas rather than producing the free metal along with nitrogen gas. Common mistake: Generalizing the thermal decomposition products of heavy metal nitrates to all alkali metal nitrates.
Which of the following is a property of ionic chlorides?
Options involving explosions with dry ammonia or decomposition upon heating might seem exotic enough to be correct for ionic halides. When ionic chlorides encounter aqueous silver nitrate, they immediately produce a characteristic white precipitate of silver chloride that redissolves upon the addition of excess ammonia due to complex ion formation. Common mistake: Forgetting the characteristic solubility behavior of silver chloride in excess ammonia solution.
When dilute aqueous solutions of lead(II) nitrate and potassium bromide are mixed, the products are
A student might select options containing lead oxide or nitrogen dioxide gas through confusion over double displacement and gas evolution reactions. Mixing lead(II) nitrate and potassium bromide results in the precipitation of insoluble lead(II) bromide while potassium and nitrate ions remain spectator ions dissolved in the aqueous mixture. Common mistake: Misidentifying insoluble precipitate products in double decomposition reactions.
Bronze is an alloy of
Options like copper and zinc might look tempting because that specific combination forms brass, leading to common alloy confusion. Bronze is a traditional metallic alloy composed specifically of copper combined with tin, prized for its strength and resilience against environmental degradation. Common mistake: Confusing the metal constituents of bronze with those of brass.
Copper metal reacts with concentrated nitric acid to produce
1) Distractor check: A student might mistakenly select option C, thinking that a lower oxide of copper is formed rather than the copper salt, or pick options like A or D regarding the specific nitrogen gas produced. 2) Reasoning to the answer: Concentrated nitric acid acts upon copper to produce copper(II) nitrate alongside water and nitrogen dioxide gas, governed by the chemical reaction represented as Cu + 4HNO₃ → Cu(NO₃)₂ + 2NO₂ + 2H₂O. The metal is oxidized specifically to the Cu²⁺ state by this concentrated acid. Common mistake: Confusing the gaseous product of concentrated nitric acid (which yields NO₂) with that of dilute nitric acid (which typically yields NO).
The active reducing agent in the blast furnace for the extraction of iron is
1) Distractor check: Students might choose option A (carbon) since carbon is a primary element in the process, or option C (limestone) because it is a standard raw material charged into the blast furnace. 2) Reasoning to the answer: Iron ore is chemically reduced within the blast furnace through the action of carbon monoxide, detailed by the equation Fe₂O₃ + 3CO → 2Fe + 3CO₂. This specific gaseous compound serves as the active reducing agent that strips oxygen from the iron oxide. Common mistake: Conflating the raw fuel/reagent materials like limestone or carbon with the actual gaseous species functioning as the active reducing agent.
Al₂O₃ reacts with both acids and bases. This property classifies Al₂O₃ as
1) Distractor check: A test-taker might mistakenly select option A or C by focusing exclusively on how the substance reacts with only one class of compounds, ignoring its dual reactivity. 2) Reasoning to the answer: Because aluminum oxide reacts with both acids, as demonstrated in Al₂O₃ + 6HCl → 2AlCl₃ + 3H₂O, and bases, shown through Al₂O₃ + 2NaOH + 3H₂O → 2NaAl(OH)₄, it is categorized chemically as amphoteric. Common mistake: Assuming an oxide that reacts with acids must exclusively be a base, rather than recognizing dual-reactivity characteristics.
The functional groups in the compound HO-CH₂-CH₂-NH₂ are
1) Distractor check: One might mistakenly choose option B, C, or D by misinterpreting the terminal heteroatom groups attached to the carbon chain, particularly confusing the amine or hydroxyl groups with carbonyl-containing functionalities. 2) Reasoning to the answer: Looking closely at the given formula HO-CH₂-CH₂-NH₂ (known as ethanolamine), the structure explicitly incorporates both a hydroxyl group (-OH), which classifies it as an alcohol, and an amino group (-NH₂), which classifies it as an amine. Common mistake: Overlooking the specific structural classification of the -OH group when paired with a nitrogen-containing moiety.
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