igcse model questions vol1 2018 chemistry | Essay

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Question 1 View Details
The diagram shows a neutral atom of chlorine (Cl).
Question Parts
(a)
State the atomic number, mass number and the number of electrons in the atom.
(b)
Write the ground‑state electron configuration for this atom using the Aufbau principle.
(c)
The atom gains one electron to form an anion. Explain how the electron arrangement changes and state the charge on the ion.
(d)
Two isotopes of chlorine are shown: one with mass number 35 and another with mass number 37. If the natural abundances are 75 % for mass 35 and 25 % for mass 37, calculate the average atomic mass of chlorine.
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Question 2 View Details
The diagram shows a section of the periodic table covering periods 3 and 4 and groups 1, 2, 13, 14, 15 and 16.
Question Parts
(a)
Identify the element located in period 4, group 13 and give its chemical symbol and atomic number.
(b)
Write the ground‑state electron configuration for the element identified in part (a).
(c)
Compare the first ionisation energy of the element in part (a) with that of the element directly above it in the same group. Explain the difference using atomic‑structure concepts.
(d)
Predict the type of oxide formed by the element in part (a) and write the chemical formula of the oxide. Justify your answer using oxidation‑state concepts.
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Question 3 View Details
The diagram shows the Lewis structure of sulfur dioxide (SO₂).
Question Parts
(a)
State the type of bond formed between the sulphur atom and each oxygen atom.
(b)
Determine the hybridisation of the central sulphur atom.
(c)
Predict the molecular geometry of SO₂ and give the approximate bond angle between the two S–O bonds.
(d)
Explain whether SO₂ is a polar molecule and justify your answer.
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Question 4 View Details
A colourless liquid organic compound was isolated and its structural diagram is shown in the illustration. The labelled atoms are C1, C2, C3, C4 (carbons) and O1, O2 (oxygens).
Question Parts
(a)
State the type of functional group present in the compound and describe the nature of the bonding between the atoms.
(b)
A 5.00 g sample of the compound is completely combusted. The masses of carbon dioxide and water produced are 10.0 g and 4.1 g respectively. Calculate the mass of carbon, hydrogen and oxygen originally present in the sample.
(c)
Using the masses obtained in part (b), determine the empirical formula of the compound.
(d)
The molar mass of the compound is known to be 88.1 g mol⁻¹. From the empirical formula obtained in (c), calculate the molecular formula and confirm that it is consistent with the structural diagram.
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Question 5 View Details
A sample of aluminium metal reacts with aqueous hydrochloric acid according to the equation: 2 Al + 6 HCl → 2 AlCl₃ + 3 H₂. A student places 15.0 g of Al into a beaker containing 80.0 cm³ of 1.5 M HCl and allows the reaction to proceed to completion.
Question Parts
(a)
Identify the limiting reactant.
(b)
Calculate the theoretical mass of aluminium chloride (AlCl₃) that can be formed.
(c)
The experiment actually yields 4.20 g of AlCl₃. Calculate the percent yield.
(d)
Determine the mass of excess aluminium that remains unreacted after the reaction is complete.
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Question 6 View Details
A rigid, sealed container of volume 2.00 L initially contains 0.500 mol of nitrogen gas at 298 K. In the same container there is 5.00 g of liquid water. The container is then heated to 373 K. The saturated vapour pressure of water at 373 K is 101.3 kPa. Assume ideal‑gas behaviour for the gases and that the water vapour also behaves ideally.
Question Parts
(a)
Calculate the partial pressure of nitrogen gas after the temperature is raised to 373 K.
(b)
Determine the maximum amount (in moles) of water that can exist as vapour in the container at 373 K.
(c)
Will all of the 5.00 g of water convert to vapour on heating? If not, calculate the mass of water that remains liquid.
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Question 7 View Details
A student prepares a buffer solution using acetic acid (CH₃COOH) and its conjugate base sodium acetate (CH₃COONa). She dissolves 0.20 mol of acetic acid in 500 mL of water and then adds solid sodium acetate until the pH of the solution is 5.00 at 25 °C. The acid dissociation constant of acetic acid is Ka = 1.8 × 10⁻⁵.
Question Parts
(a)
Write the acid‑dissociation equilibrium equation for acetic acid and define the meaning of Ka.
(b)
Using the Henderson–Hasselbalch equation, calculate the number of moles of sodium acetate that must be added to obtain a pH of 5.00. Show all steps and state any assumptions made.
(c)
Explain why the resulting solution functions effectively as a buffer. In your answer, describe the role of the conjugate base and the effect of adding a small amount of a strong acid or a strong base.
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Question 8 View Details
A laboratory investigation studies the solubility of lead(II) iodide, PbI₂, at 25 °C. The solubility‑product constant is Ksp = 8.5 × 10⁻⁹. In a first experiment the solid is placed in pure water until equilibrium is reached. In a second experiment 0.010 mol of NaI is dissolved in 250 mL of water and excess solid PbI₂ is added, also reaching equilibrium. In a third experiment solid NaCl (0.20 mol) is added to the saturated solution from the second experiment.
Question Parts
(a)
Write the dissolution equilibrium equation for PbI₂ and express the Ksp expression in terms of ion concentrations.
(b)
Calculate the molar solubility of PbI₂ in pure water at 25 °C.
(c)
In the solution that already contains 0.010 mol of I⁻ from NaI (volume = 250 mL), determine the concentration of Pb²⁺ ions at equilibrium. Show all reasoning.
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Question 9 View Details
A piece of clean zinc metal is immersed in 250 cm³ of a 0.010 mol dm⁻³ copper(II) sulphate solution at 25 °C. The reaction proceeds until the solution becomes colourless. The standard reduction potentials are: \[\text{Cu}^{2+}+2e^-\rightarrow\text{Cu}\quad E^{\circ}=+0.34\text{ V}\] \[\text{Zn}^{2+}+2e^-\rightarrow\text{Zn}\quad E^{\circ}=-0.76\text{ V}\] During the experiment a constant current of 0.50 A is passed for 30 minutes. Answer the following:
Question Parts
(a)
Write the balanced overall redox equation for the reaction between zinc metal and copper(II) ions.
(b)
Write the two half‑reactions (oxidation and reduction) and state which species acts as the oxidising agent and which as the reducing agent.
(c)
Calculate the mass of copper metal deposited on the zinc surface after the 30‑minute period. Assume 100 % current efficiency.
(d)
Determine the standard electromotive force (E°cell) of the Zn|Zn^{2+} || Cu^{2+}|Cu cell.
(e)
Explain qualitatively how the cell potential would change if the concentration of Cu^{2+} ions were halved, using the Nernst equation (no numerical calculation required).
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Question 10 View Details
An aqueous solution of sodium chloride (0.5 mol dm⁻³) is electrolysed in an electrolytic cell containing inert graphite electrodes. A constant current of 2.0 A is passed for 30 minutes. The temperature is 25 °C and the cell operates with 100 % current efficiency for the reactions shown. Answer the following:
Question Parts
(a)
Write the overall chemical equation for the electrolysis of the solution.
(b)
State the half‑reaction occurring at the cathode and at the anode, indicating which species is oxidised and which is reduced.
(c)
Calculate the volume of hydrogen gas (at STP) produced at the cathode during the 30‑minute period.
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