Reacting Masses & Volumes (of Solutions & Gases) (Cambridge (CIE) A Level Chemistry): Exam Questions

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

0.50 g of quartz sand and 0.50 g of powdered silica are mixed with 2.40 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.

i) Construct an equation for this reaction.

Include state symbols.

[2]

ii) Deduce which reagent is the limiting reagent. Show your working.

[2]

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

Mg₂Si (s) + 4HCl (aq) → 2MgCl₂ (aq) + SiH₄ (g)

The gaseous silane formed spontaneously ignites in air.

State the type of reaction and construct an 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.

Si (s) + 3HCl (aq) → HSiCl₃ (l) + H₂ (g)

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

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)

Calculate the volume, in dm³, of hydrogen gas required to produce 10.0 dm³ of silane gas. Explain your answer.

Assume all volumes are measured under the same conditions.

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.

i) Calculate the amount, in mol, of nitrous oxide in the measuring cylinder.

[1]

ii) Calculate the number of carbon disulfide molecules required to react with the nitrous oxide at room conditions. Show your working.

[2]

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

i) Deduce which reagent is in excess. Show your working.

[2]

ii) Calculate the maximum mass of sulfur that will be produced from this demonstration.

[1]

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 conditions. This produces 14.32 g of sulfur.

Calculate the percentage yield for this demonstration. Show your working.

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.

i) Construct an equation for this reaction of iron with chlorine.

State symbols are not required.

[1]

ii) Calculate the mass of iron(III) chloride produced from 5.02 g of iron, assuming a 78% yield.

Show your working.

[3]

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, in dm³, produced at room conditions by this reduction of 9.86 g of iron(III) oxide.

Show your working.

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.

Deduce the equation for this reaction. Show your working.

Verdigris is a green pigment that contains both copper(II) carbonate, CuCO₃, and copper(II) hydroxide, Cu(OH)₂, in varying amounts. Both copper compounds react with dilute hydrochloric acid.

CuCO₃ (s) + 2HCl (aq) → CuCl₂ (aq) + CO₂ (g) + H₂O (l)

Cu(OH)₂ (s) + 2HCl (aq) → CuCl₂ (aq) + 2H₂O (l)

You are to plan an experiment to determine the percentage of copper(II) carbonate in a sample of verdigris. Your method should involve the reaction of verdigris with excess dilute hydrochloric acid.

You are provided with the following:

• 0.494 g of verdigris

• 10.0 mol dm^–3 hydrochloric acid, HCl (aq)

• commonly available laboratory reagents and equipment

You may assume that any other material present in verdigris is unaffected by heating and is not acidic or basic.

i) A student suggests that finding the volume of dilute hydrochloric acid required to react with a known mass of verdigris would be a suitable method to determine the percentage of copper(II) carbonate in a sample of verdigris.

Suggest why this method would not work.

[1]

ii) The 10.0 mol dm^–3 HCl is too concentrated for use in the experiment. Instead, a more dilute solution should be prepared.

Describe how you would accurately prepare 250.0 cm³ of 0.500 mol dm^–3 hydrochloric acid from the 10.0 mol dm^–3 HCl provided.

Your answer should state the name and capacity in cm³ of any apparatus you would use.

[3]

iii) The percentage of copper(II) carbonate in a sample of verdigris can be determined by measuring the volume of gas produced when excess hydrochloric acid is added to the sample of verdigris.

Draw a diagram to show how you would set up the apparatus and chemicals to measure the total volume of gas produced in this reaction.

Label your diagram.

[2]

iv) Sketch a graph on the axes to show how the volume of gas produced would change during your experiment. The independent variable should be on the x-axis.

• Label both axes.

• Extend the graph beyond the point at which the reaction is complete.

[2]

v) A student thinks that their 0.494 g sample of verdigris only contains CuCO₃.

Calculate the minimum volume, in cm³, of 0.500 mol dm^–3 HCl that is needed to completely react with this sample if the student is correct. Show your working. [M_r: CuCO₃ = 123.5]

[2]

Azurite is a blue copper-containing mineral. The copper compound in azurite has the formula Cu₃(CO₃)₂(OH)₂. This copper compound reacts with sulfuric acid according to the equation.

Cu₃(CO₃)₂(OH)₂ (s) + 3H₂SO₄ (aq) → 3CuSO₄ (aq) + 2CO₂ (g) + 4H₂O (l)

A student carries out a series of titrations on 1.50 g samples of solid azurite using 0.400 mol dm^–3 sulfuric acid.

Assume that any other material present in azurite does not react with sulfuric acid. Some titration data is given in Table 1.1.

Table 1.1

The indicator for the titration is bromophenol blue. Bromophenol blue is blue at pH 4.6 and yellow at pH 3.0.

i) Complete Table 1.1.

[1]

ii) Calculate the percentage uncertainty in titre 1.

[1]

iii) The student concludes that 24.15 cm³ of 0.400 mol dm^–3 sulfuric acid completely reacts with 1.50 g of azurite. Calculate the percentage by mass of Cu₃(CO₃)₂(OH)₂ in the sample of azurite using the student's value of 24.15 cm³ of 0.400 mol dm^–3 sulfuric acid.

Write your answer to three significant figures. Show your working. [M_r: Cu₃(CO₃)₂(OH)₂ = 344.5]

[3]

iv) Identify two possible problems with the student's titration experiment and suggest improvements to it.

problem 1

improvement 1

problem 2

improvement 2

[4]

Activated charcoal is a form of carbon with a very high surface area. It can be used to remove impurities from mixtures. It does this by a process called adsorption, where particles of the impurity bond (adsorb) to the activated charcoal surface.

A student wants to determine the ability of activated charcoal to adsorb a blue dye (the impurity) from aqueous solution.

The equation that links the mass of activated charcoal with the amount of blue dye adsorbed is shown.

log = A + b log [X]

D = difference in concentration of dye (in g dm^–3) before and after adsorption
m = mass of activated charcoal (in g)
[X] = final concentration of dye (in g dm^–3) after adsorption
A and b are constants

The student uses the following procedure to investigate the ability of activated charcoal to adsorb a blue dye from an aqueous solution.

• Place a 50.0 cm³ sample of a 25.00 g dm^–3 solution of blue dye in a conical flask.

• Add a weighed mass of activated charcoal to the flask.

• Stir the contents of the flask for three minutes and then leave for one hour.

• Filter the mixture.

• Determine the final concentration of the blue dye, [X].

• Repeat the procedure using different masses of activated charcoal.

The procedure is carried out. The final concentrations of blue dye, [X], are shown in Table 2.1.

i) Process the results to complete Table 2.1.

Record your data to two decimal places.

Table 2.1

[2]

ii) Identify the dependent variable in this experiment.

[1]

iii) State and explain the effect, if any, of increasing the mass of activated charcoal, m, on the amount of adsorption that occurs.

[2]

Plot a graph on the grid to show the relationship between log and log [X].

Use a cross (×) to plot each data point. Draw the straight line of best fit.

Circle the most anomalous point on the graph.

Suggest why this anomaly may have happened during the experimental procedure.

i) Use the graph to determine the gradient of the line of best fit. State the coordinates of both points you used in your calculation. These must be selected from your line of best fit.

Write your answer to three significant figures

coordinates 1 ....................

coordinates 2 ..........................

gradient = ..........................................

[2]

ii) Use the graph to deduce a value for A.

A = ....................

[1]

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.

Show your working.

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

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, at room conditions, of carbon dioxide that will react with 5.00 g of the superoxide. Give your answer to 3 significant figures.

Show your working.

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.

Show your working.

This question is about the reactions of ionic compounds.

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

NaCl + NH₃ + CO₂ + H₂O → NaHCO₃ + NH₄Cl

2NaHCO₃ → Na₂CO₃ + H₂O + CO₂

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

Show your working.

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.

Show your working.

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.

Show your working.

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.

Show your working.