neco model questions vol1 2017 chemistry | Essay

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Question 1 View Details
The diagram shows a Bohr model of an element X. The nucleus is labelled with an atomic number of 15 and a mass number of 31. The electrons are arranged in three shells: the innermost shell contains 2 electrons, the second shell contains 8 electrons and the outermost shell contains 5 electrons.
Question Parts
(a)
State the name of the element and its chemical symbol.
(b)
Calculate the number of neutrons present in the nucleus of the element.
(c)
Write the ground‑state electron configuration of the element (i) using Bohr notation and (ii) using orbital (Aufbau) notation.
(d)
The element is excited so that one electron from the 3p subshell is promoted to the 4s subshell. (i) Write the new electron configuration. (ii) State qualitatively what happens to the energy of the atom during this transition and explain why such a transition is possible.
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Question 2 View Details
The illustration presents a portion of the periodic table covering periods 1 to 3 and groups 1, 2 and 13–18. For each element shown, the atomic number, symbol and, where indicated, the atomic radius (in picometres) are given.
Question Parts
(a)
Identify the period and group of chlorine and describe two general chemical properties of elements in this group.
(b)
The diagram gives the atomic radii of sodium (186 pm) and chlorine (99 pm). Explain, using periodic trends, why the radius of chlorine is smaller than that of sodium even though both belong to the same period.
(c)
Predict the type of bond formed when magnesium reacts with chlorine and write the chemical formula of the compound produced. Justify your answer using the oxidation states inferred from the periodic table.
(d)
Both magnesium and aluminium are in period 3. Using periodic trends, explain why magnesium metal reacts more vigorously with dilute hydrochloric acid than aluminium does, even though both can produce hydrogen gas.
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Question 3 View Details
Magnesium chloride (MgCl₂) is a typical inorganic compound formed between a metal and a non‑metal. The electronegativity values are: Mg = 1.31, Cl = 3.16. Answer the following questions.
Question Parts
(a)
State the type of bond present in MgCl₂ and justify your answer using the given electronegativity values. Explain why MgCl₂ forms a crystal lattice rather than discrete molecules.
(b)
Draw the Lewis dot diagram for the carbonate ion (CO₃²⁻). Indicate the resonance structures and explain the concept of resonance as it applies to this ion.
(c)
Using Pauling’s empirical formula, estimate the percent ionic character of the Mg–Cl bond. Pauling’s formula: % ionic character = \(\left(1- e^{-0.25(Δχ)^2}\right)\times100\) where \(Δχ\) is the electronegativity difference.
(d)
Compare qualitatively the lattice energy of MgCl₂ with that of NaCl and CaO. Discuss how the ionic charges and ionic radii influence the magnitude of lattice energy.
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Question 4 View Details
A crystalline solid has been analysed and its composition is represented by the structure shown in the illustration. The solid is a hydrate of a transition‑metal complex. Using the diagram, answer the following questions.
Question Parts
(a)
Write the correct molecular formula of the compound, including the water of crystallisation.
(b)
Calculate the molar mass of the compound (to three significant figures).
(c)
A 5.00 g sample of the compound is dissolved in water. Calculate:
() The number of moles of Cr present.
() The mass of chloride ions (both coordinated and counter‑ions) present in the sample.
(d)
The solution is treated with excess silver nitrate, precipitating all chloride ions as AgCl. Calculate the mass of AgCl that would be formed from the 5.00 g sample.
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Question 5 View Details
A 15.0 g sample of impure calcium carbonate (CaCO₃) is heated to decompose it according to CaCO₃(s) → CaO(s) + CO₂(g). The sample is 80 % pure CaCO₃ by mass. After heating, 5.0 g of CaO is collected.
Question Parts
(a)
Determine the theoretical mass of CaO that could be obtained from the pure CaCO₃ present in the sample.
(b)
Calculate the percent yield of CaO based on the actual mass obtained.
(c)
If the same 15.0 g impure sample were instead reacted with excess hydrochloric acid according to CaCO₃ + 2 HCl → CaCl₂ + H₂O + CO₂, determine the theoretical mass of CaCl₂ that could be formed and discuss which method (heating vs acid) gives a higher efficiency, assuming 100 % yield for the acid reaction.
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Question 6 View Details
A 2.00 mol sample of a pure substance X has a melting point of 120 °C and a boiling point of 250 °C at 1 atm. Its specific heat capacities are: solid = 30 J mol⁻¹ K⁻¹, liquid = 45 J mol⁻¹ K⁻¹, gas = 35 J mol⁻¹ K⁻¹. The latent heat of fusion is 6.0 kJ mol⁻¹ and the latent heat of vaporisation is 30 kJ mol⁻¹. The substance is initially at 25 °C.
Question Parts
(a)
Calculate the total amount of heat required to convert the entire sample from solid at 25 °C to gas at 260 °C (i.e., heating to just above the boiling point).
(b)
Sketch a qualitative heating curve for this process, labeling the key temperatures and phases and indicating the relative amounts of heat absorbed in each segment.
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Question 7 View Details
A 0.250 mol sample of the weak monoprotic acid HA (Ka = 1.8 × 10⁻⁵) is dissolved in 500.0 mL of water. To this solution 0.150 mol of its sodium salt NaA is added and the mixture is allowed to equilibrate at 25 °C. Answer the following:
Question Parts
(a)
Calculate the pH of the resulting solution. Show all steps and state any assumptions made.
(b)
The solution is to be titrated with 0.100 M HCl until the pH becomes 4.00. Calculate the exact number of moles of HCl required and the volume of the HCl solution that must be added. Show the derivation of the equation used.
(c)
Suggest a suitable indicator for monitoring this titration and justify your choice based on the pH range of colour change.
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Question 8 View Details
A student carries out a solubility experiment to determine the solubility of calcium fluoride (CaF₂) in water at two different temperatures using the apparatus shown. The apparatus consists of a 250 mL beaker placed on a thermostatically controlled water‑bath, a calibrated thermometer, a magnetic stirrer with a stir bar, a filter funnel fitted with filter paper, and a digital analytical balance. The student adds excess solid CaF₂ to 250 mL of distilled water, heats the water‑bath to the required temperature, stirs until equilibrium is reached, then quickly filters the saturated solution and weighs the amount of solid that remained undissolved. The data obtained are: - At 25 °C, 0.50 g of CaF₂ remained undissolved. - At 40 °C, 0.85 g of CaF₂ remained undissolved. Using this information, answer the following questions:
Question Parts
(a)
Write the balanced net ionic equation for the dissolution of CaF₂ in water.
(b)
Calculate the solubility of CaF₂ in mol L⁻¹ at 25 °C and 40 °C. (Assume the volume of the solution remains 250 mL after dissolution.)
(c)
Using the solubilities obtained in part (b), calculate the solubility product constant (Ksp) of CaF₂ at 25 °C and at 40 °C.
(d)
Explain qualitatively why the solubility of CaF₂ increases with temperature. Then, using the Ksp expression, discuss how the addition of a soluble fluoride salt such as NaF would affect the solubility of CaF₂ at 25 °C.
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