Covering a broad range of inorganic, physical, and organic concepts, the 2004 JAMB Chemistry paper challenges candidates to apply textbook definitions to practical chemical scenarios. Key recurring themes highlighted across these questions include solution concentration and dilution stoichiometry, chemical equilibrium adjustments under varying stress conditions, and laboratory preparation methods alongside gas properties. Utilizing this page effectively means tracking down why incorrect options are wrong, especially in conceptual multiple-choice items, to ensure genuine mastery of the subject matter.
Topics in this paper
What is the oxidation number of sulfur in H₂SO₄?
A student might incorrectly select option A, B, or D by miscalculating the charges or forgetting the standard oxidation numbers for hydrogen and oxygen. For a neutral molecule of H₂SO₄, the sum of all oxidation numbers must equal zero. Using the known values of +1 for hydrogen and -2 for oxygen, we set up the equation 2(+1) + x + 4(-2) = 0, which simplifies to 2 + x - 8 = 0, yielding x = +6 for sulfur's oxidation number. Common mistake: failing to account for the subscript multiplier on the oxygen atoms.
What is the concentration of a solution containing 2g of NaOH in 100cm³ of solution? [Na = 23, O = 16, H = 1]
A student might mistakenly choose option A, C, or D by failing to correctly convert units or dividing by the wrong molar mass. First, find the molar mass of NaOH, which is 40 g/mol. Next, compute the moles by dividing the given mass by this molar mass (2 / 40 = 0.05 mol). Convert the volume from 100 cm³ to 0.1 dm³, and then divide the moles by this volume (0.05 / 0.1) to arrive at a concentration of 0.50 mol dm⁻³. Common mistake: omitting the volume unit conversion from cm³ to dm³.
Which of the following properties is NOT peculiar to matter?
A student might mistakenly pick option A, B, or C because those choices correctly describe real trends in the kinetic energy, random motion, or orderliness of matter across different states. Option D is incorrect—and therefore the correct answer to the question asking for properties that are NOT peculiar—because random motion of particles actually decreases, rather than increases, when moving from a gas to a solid due to increasing orderliness. Common mistake: confusing the direction of the physical state transition when evaluating particle behavior.
The principle of column chromatography is based on the ability of the constituents to
A student might mistakenly select option B, C, or D by assuming that stationary or reacting phases dominate the chromatographic process. Column chromatography separates components based on their differing rates of movement, meaning the constituents move at different speeds through the column. Common mistake: assuming chemical reactions take place between the stationary and mobile phases during separation.
Which of the following is a characteristic of an ideal gas?
A student might mistakenly choose option A, C, or D because they confuse real gas deviations with ideal behavior. An ideal gas follows the ideal gas law equation, PV = nRT, which assumes that the particles have negligible volume and experience no intermolecular forces. Common mistake: assuming real gas properties like strong intermolecular forces apply to ideal models.
Which of the following statements is correct about the periodic table?
A student might mistakenly pick option A, C, or D by confusing trends across periods with trends down groups in the periodic table. As atomic number increases across a period, the valence electrons of the elements increase progressively. Common mistake: mixing up atomic properties across rows versus columns of the periodic table.
The relative atomic mass of a naturally occurring lithium consisting of 90% ⁷Li and 10% ⁶Li is
A student might mistakenly select option B, C, or D by incorrectly weighting the isotope percentages or miscalculating the arithmetic sum. The relative atomic mass is found by multiplying each isotope's mass by its fractional abundance and adding them together: (0.9 × 7) + (0.1 × 6), which gives 6.3 + 0.6, resulting in a final value of 6.9. Common mistake: using whole-number percentages directly without converting them to decimals.
An isotope has an atomic number of 15 and a mass number of 31. The number of protons it contains is
A student might mistakenly pick option A, C, or D by confusing the mass number with the atomic number or subtracting values incorrectly. By definition, the atomic number of an element equals the number of protons it contains, which is directly given as 15. Common mistake: subtracting the atomic number from the mass number when asked specifically for the proton count.
The molecular lattice of iodine is held together by
A student might mistakenly choose option A, B, or C by assuming iodine forms stronger intramolecular bonds or metallic crystals in the solid state. The molecular lattice of iodine is held together by weak van der Waal's forces. Common mistake: confusing covalent bonding within the diatomic molecule with the intermolecular forces holding the lattice together.
The arrangement of particles in crystal lattices can be studied using
A student might mistakenly select option B, C, or D by guessing other types of radiation without considering wavelength compatibility with crystal spacings. X-rays are used in X-ray diffraction because their wavelength is appropriate for studying crystal lattice structures. Common mistake: selecting high-energy gamma rays instead of X-rays for diffraction analysis.
The solubility of a solute is 0.6 mol dm⁻³ at 55°C and 0.5 mol dm⁻³ at 40°C. How much solute is deposited when 200 cm³ of the solution is cooled from 55°C to 40°C?
A student might mistakenly choose option B, C, or D by forgetting to factor in the volume of the solution or miscalculating the solubility difference. First, find the solubility difference between the two temperatures (0.6 - 0.5 = 0.1 mol dm⁻³). Then, convert the volume to dm³ (200 cm³ = 0.2 dm³) and multiply it by the solubility difference to find the deposited solute: 0.1 × 0.2 = 0.02 mol. Common mistake: failing to convert cubic centimeters to cubic decimeters before multiplying by solubility.
The importance of sodium aluminate (III) in the treatment of water is to
A student might mistakenly pick option B, C, or D by confusing water coagulants with disinfectants or pH adjusters. Sodium aluminate promotes coagulation, aiding in water purification and treatment. Common mistake: thinking coagulants are used primarily to kill germs or prevent dental decay.
What type of bond exists between an element with atomic number 12 and an element with atomic number 17?
A student might mistakenly choose option A, C, or D by misinterpreting the bonding behavior of non-metals and metals. Element 12 is Magnesium (a metal) and element 17 is Chlorine (a non-metal), which combine to form MgCl₂ via ionic bonding. Common mistake: assuming elements with higher atomic numbers automatically form covalent bonds.
Hardness of water is mainly due to the presence of
A student might mistakenly select option A, C, or D by confusing temporary or permanent hardness compounds with other hydroxides or soluble salts. Hardness of water is mainly due to the presence of calcium hydrogen carbonate or calcium sulfate, represented here by calcium trioxocarbonate (IV) or calcium tetraoxosulphate (VI). Common mistake: confusing soluble sodium salts with hardness-causing calcium and magnesium salts.
A suitable solvent for iodine and naphthalene is
A student might mistakenly pick option A, C, or D by assuming non-polar solutes dissolve easily in polar solvents like water. Ethanol effectively dissolves both iodine and naphthalene due to similar intermolecular characteristics. Common mistake: assuming water is a universal solvent for all solid substances.
Which of the following noble gases is commonly found in the atmosphere?
A student might mistakenly choose option A, B, or C by selecting a rarer noble gas or helium. Argon is the most abundant noble gas found in the atmosphere. Common mistake: picking helium because it is well-known, overlooking its low atmospheric concentration.
N₂O₄(g) ⇌ 2NO₂(g) ∆H = +ve. In the reaction above, an increase in temperature will
A student might mistakenly select option B, C, or D by applying Le Chatelier's principle incorrectly to endothermic versus exothermic systems. For an endothermic reaction (where ∆H is positive), a higher temperature increases the equilibrium constant K and shifts the equilibrium to the right. Common mistake: assuming temperature increases always favor the reverse direction in chemical equilibria.
CH₃COOH(aq) + OH⁻(aq) ⇌ CH₃COO⁻(aq) + H₂O(l). In the reaction above, CH₃COO⁻(aq) is
A student might mistakenly pick option B, C, or D by misidentifying acid-base pairs in the reverse direction. In this equilibrium, CH₃COO⁻ is the conjugate base formed after the acid CH₃COOH loses a proton (H⁺). Common mistake: labeling the species left behind after proton loss as an acid instead of a conjugate base.
How many cations will be produced from a solution of potassium aluminium tetraoxosulphate (VI)?
A student might mistakenly choose option A, B, or C by miscounting the total ions or missing the aluminum ion contribution. Potassium aluminium tetraoxosulphate (VI), KAl(SO₄)₂, dissociates into K⁺, Al³⁺, and 2SO₄²⁻, which produces 2 distinct cations (K⁺ and Al³⁺). Common mistake: counting anions alongside cations or forgetting to include the aluminum ion.
Which of the following is NOT an alkali?
A student might mistakenly pick option B, C, or D by assuming all substances containing hydrogen or hydroxyl groups are true alkalis. NH₃ is a weak base, not an alkali, whereas the others are soluble metal hydroxides that qualify as true alkalis. Common mistake: treating all basic substances as soluble alkalis.
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