Atoms, Molecules & Stoichiometry (CIE A Level Chemistry)

This question is about the formation and reaction of silane, SiH₄.

0.5 g of quartz sand and 0.5 g of powdered silica are mixed with 2.4 g of magnesium powder. The resulting mixture is placed in a pot half filled with sand. A magnesium ribbon fuse is added and the mixture is ignited to form magnesium oxide and solid magnesium silicide, Mg₂Si.

The magnesium silicide is reacted with hydrochloric acid according to the following equation:

The gaseous silane formed spontaneously ignites in air. 

State the type of reaction and write the chemical equation for this reaction.

Include state symbols.

An alternative method to produce silane involves the reaction of silicon with hydrochloric acid to form trichlorosilane, HSiCl₃.

Calculate the maximum mass of trichlorosilane produced from the reaction of 35.125 g of silicon.

Another method of silane production involves the reduction of silicon dioxide by aluminium and hydrogen gas.

3SiO₂ (s) + 6H₂ (g) + 4Al (s) → 3SiH₄ (g) + 2Al₂O₃ (s)

What is the volume, in dm³, of hydrogen gas required to produce 10 dm³ of silane gas? Explain your answer.

Cobalt(II) sulfate is used in storage batteries, electroplating baths and as an additive to soils and animal feeds.

A sample of hydrated cobalt sulfate, CoSO₄•xH₂O, has a relative formula mass of 173.0.

Calculate a value for 'x'.

A student completed an experiment to determine the value of ‘x’ in a hydrated salt.

They heated a solid sample of the salt, CoSO₄.xH₂O, in a crucible for 60 seconds and recorded the following set of results.

Use the data obtained by the student in Table 2.1 to calculate a value for ‘x’.

The actual value for x in CoSO₄.xH₂O is 6.

Cobalt(II) sulfate undergoes a displacement reaction with magnesium because magnesium is a more reactive metal than cobalt.

Write an ionic equation for this reaction. State symbols are not required.

Chevreul's salt or copper(I, II) sulfite dihydrate, Cu₃(SO₃)₂•2H₂O, is an unusual hydrated salt as it contains copper in both of its common oxidation states.

The first step in the formation of Chevreul's salt is the reaction of aqueous copper sulfate with a solution of potassium metabisulfite, K₂S₂O₅.

Balance the equation for the reaction to form the anhydrous Chevreul's salt.

The chemical formula of Chevreul's salt is sometimes written as Cu₂SO₃•CuSO₃•2H₂O to show that it contains copper(I) sulfite and copper(II) sulfite.

Calculate the percentage by mass of copper(I) ions in copper(I) sulfite

Show your working.

Chevreul's salt can react with ammonia solution to form a deep blue solution of the tetraamminecopper(II) complex, [Cu(NH₃)₄]²⁺ (aq). It can also react with hydrochloric acid to form a white solid of copper(I) chloride, water and sulfur dioxide.

Write an ionic equation for the formation of copper(I) chloride from Chevreul's salt. State symbols are not required.

When Chevreul's salt is heated in an inert atmosphere, it is stable to around 200 ^oC. Above this temperature, a variety of compounds can be produced although the final product of sustained heating at high temperatures is copper(II) oxide.

During the heating process, a sample of the mixture is taken. After separation, the most common compound was analysed and found to contain 39.8% copper, 20.1% sulfur and the rest oxygen.

Calculate the empirical formula of the most common compound. Show your working.

The reaction of nitrous oxide gas, N₂O (g), with carbon disulfide vapour, CS₂ (g), is an example of a chemical luminescence reaction producing a flame with a bright blue light.

2N₂O (g) + CS₂ (g) → 2S (s) + 2N₂ (g) + CO₂ (g)

The brightness of the light produced was a reason why this reaction was once used in low-light photography.

When nitrous oxide and carbon disulfide are placed together there is no immediate reaction.

Explain how a flame can help initiate the reaction.

A measuring cylinder containing 100 cm³ of nitrous oxide is used in a demonstration of the reaction between nitrous oxide and carbon disulfide.

In a second demonstration, 10.8 dm³ of nitrous oxide is reacted with 5.6 dm³ carbon disulfide vapour.

The demonstration is performed a third time to determine the percentage yield of sulfur of the reaction.

For this demonstration, 16.50 dm³ of nitrous oxide is reacted with an excess of carbon disulfide at room temperature and pressure. This produces 14.32 g of sulfur.

Calculate the percentage yield for this demonstration. 

Iron compounds have a variety of uses. Many iron compounds are found in fertilisers and as food additives and supplements.

Iron sulfate is used to treat and prevent iron deficiency anaemia.

Suggest two possible formulae for iron sulfate.

Iron reacts with chlorine to form iron(III) chloride.

Iron(III) oxide is reduced to Fe with carbon at a temperature of 1200 °C in the blast furnace. 

2Fe₂O₃ + 3C → 4Fe + 3CO₂ 

An example reaction is set up by heating 9.86 g of iron(III) oxide with an excess of carbon powder. After the reaction is complete, it is left to cool to room temperature. 

Calculate the volume of carbon dioxide, at room temperature, that is produced by this reduction of 9.86 g of iron(III) oxide.

A student reacts 4.95 g of iron(III) oxide with 1.12 g of carbon.

The resulting products are 3.46 g of iron and 2.60 g of carbon monoxide.

Determine the balanced equation for the reaction. Show your working.

Magnesium carbonate is a common antacid that is used to react with and neutralise excess stomach acid, HCl. 

MgCO₃ (s) + 2HCl (aq) MgCl₂ (aq) + CO₂ (g) + H₂O (l)

The total mass of the reactants is 9.91 g.

Calculate the volume of carbon dioxide that is produced at room temperature and pressure.

Another magnesium salt, magnesium sulfate, can exist in hydrated or anhydrous forms.

6.31 g of a sample of magnesium sulphate, MgSO₄.xH₂O, was strongly heated until a mass of 3.61 g of solid remained with no further change in mass recorded.

Calculate the number of moles of water of crystallisation, x, in the magnesium sulphate sample.

moles of water = ..................... mol

Another sample of hydrated magnesium sulphate, MgSO₄.xH₂O, is found to contain 51.1% water by mass.

Calculate the number of moles of water of crystallisation, x, in this sample.

moles of water = ..................... mol

Describe the effect on the value of x obtained if the salt was not heated strongly to a constant mass.

Fireworks contain elements to give them their colour, as well as oxygen-containing compounds to facilitate speedy combustion.

In a firework, solid potassium nitrate, KNO₃, decomposes to form solid potassium nitrite, KNO₂, and oxygen, O₂.

Calculate the mass, in g, of potassium nitrate, KNO₃, required to make 1.5 g of oxygen. Give your answer to 2 significant figures.

mass of KNO₃ = .................. g

Calculate the volume, in dm³, of gas that is produced when solid potassium nitrate, KNO₃, decomposes at room temperature and pressure. Give your answer to 2 significant figures. 

volume of gas = ................ dm³

Potassium can form a superoxide, KO₂ (s), which will react with carbon dioxide, CO₂ (g), to produce potassium carbonate, K₂CO₃ (s) and oxygen, O₂ (g).

Calculate the volume, in dm³, of carbon dioxide which will react with 5.00 g of the superoxide. Give your answer to 3 significant figures. 

volume of CO₂ = .................. dm³

2.61 g of potassium carbonate, K₂CO₃, was produced during the reaction in part (c).

Calculate the percentage yield of potassium carbonate. Give your answer to 2 decimal places. 

% yield = ................... % 

The coca-cola company states that a 12 oz can of coke contains 34 mg of caffeine, M_r = 194.19 g mol^-1.

Calculate the number of moles of caffeine in a can of coke.

moles of caffeine = .................... mol

Caffeine is made up of 49.48%  carbon, 5.20% hydrogen, 28.85% nitrogen and the rest is oxygen.

Calculate the empirical formula of caffeine.

Calculate the molecular formula of caffeine using your answer from part (b).

This question is about the reactions of ionic compounds.

Sodium carbonate is manufactured in a two-stage process as shown.

Calculate the maximum mass of sodium carbonate that can be obtained from 0.75 kg of sodium chloride.

Norgessaltpeter, Ca(NO₃)₂, was the first nitrogen fertiliser to be manufactured in Norway. It is produced by the reaction of calcium carbonate with nitric acid.

In an experiment, an excess of powdered calcium carbonate was added to 37.4 cm³ of 0.531 mol dm^–3 nitric acid.

Calculate the minimum mass of CaCO₃ that should be added to react with all of the nitric acid. Give your answer to an appropriate number of significant figures.

mass of CaCO₃ = ....................

A 0.0830 mol sample of pure zinc oxide was added to 75.0 cm³ of 0.90 mol dm⁻³ hydrochloric acid.

Assuming the reaction has a percentage yield of 60.6%, calculate the mass of anhydrous zinc chloride that could be obtained from the products of this reaction.

mass of zinc chloride = ................. g

Magnesium nitride, Mg₃N₂, reacts with water to form magnesium hydroxide and ammonia.

Calculate the number of molecules of ammonia gas produced by the reaction of 12.62 g of magnesium nitride.

number of molecules = ............................

Apatite is a general term used for a family of various calcium phosphate compounds commonly found in rocks and minerals, bones and teeth.

Fluorapatite, Ca₅(PO₄)₃F, reacts with sulfuric acid to form hydrogen phosphate, calcium sulfate and hydrogen fluoride.

Write the balanced chemical equation for this reaction.

Fluorapatite, Ca₅(PO₄)₃F, can also react with sulfuric acid to form calcium dihydrogenphosphate, Ca(H₂PO₄)₂, calcium sulfate and hydrogen fluoride.

2Ca₅(PO₄)₃F + 7H₂SO₄ → 3Ca(H₂PO₄)₂ + 7CaSO₄ + 2HF

Calculate the mass of calcium dihydrogenphosphate formed when 50.2 g of Ca₅(PO₄)₃F reacts with 87.9 g H₂SO₄.

mass of Ca(H₂PO₄)₂ = ............................ g