Calculating Mass of Substances (AQA GCSE Chemistry: Combined Science)

The amount of substances can be described in different ways.

Complete the sentences. 

One mole of a substance is the relative formula mass in ____________________

The relative atomic mass of an element compares the mass of an atom of an element with the mass of an atom of ____________________ 

Explain the meaning of the term “mole”.

Describe the link between the mole and the relative atomic mass of an element.

Calcium sulfate is used as a coagulant in the production of tofu and to make Plaster of Paris. 

The molecular formula of calcium sulfate is CaSO₄. 

Calculate the relative formula mass (M_r) of CaSO₄. 

Relative atomic masses: O = 16; S = 32; Ca = 40. 

Relative formula mass (M_r) = _______________

What is the mass of one mole of calcium sulfate?

What mass of calcium sulfate would be needed to provide 18 grams of calcium?

Calculate the number of moles in 320 g of Fe₂O₃. 

You may need to use the periodic table to answer this question.

Calculate the number of moles in 14 g of C₂H₄.

Magnesium and steam react to produce magnesium oxide and hydrogen gas.

A student carried out the reaction using the apparatus shown in Figure 1 below.

 Figure 1

The balanced equation for the reaction is:

Mg (s) + H₂O (g) → MgO (s) + H₂ (g)

The student used 2.00 g of Mg.

Using the equation, calculate the maximum mass of magnesium oxide produced. 

Give your answer to three significant figures. 

Relative atomic masses (A_r): O = 16; Mg = 24

Maximum mass MgO = ____________________ g

A student uses 75 cm ³ of a copper (II) sulfate solution of 90 g / dm³ in an experiment. 

What mass of copper (II) sulfate is dissolved in 75 cm ³ of this solution? 

Mass = ____________________ g

Another student uses a 0.20 mol / dm³ solution of sodium hydroxide in an acid - base reaction. 

Calculate the mass of sodium hydroxide in 250 cm³ of a 0.20 mol / dm³ solution. 

Relative formula mass (M_r): NaOH = 40

Mass of sodium hydroxide = ___________________ g

How many moles are present in 9 kg of glucose, C₆H₁₂O₆?

Relative formula mass (M_r):       C₆H₁₂O₆ = 180

What is the mass of 0.25 moles of hydrated copper(II) sulfate, CuSO₄.5H₂O?

Relative formula mass (M_r):  H₂O = 18  CuSO₄ = 160

This question is about Avogadro's constant. 

How many grams are there in one mole of carbon?

How many atoms are present in one mole of chlorine atoms?

Tick (✔) one box.

How many atoms are present in one mole of oxygen molecules?

Tick (✔) one box.

Methanol is produced on an industrial scale by reacting carbon monoxide with hydrogen. 

The equation for the reaction is: 

CO (g) + 2H₂ (g) ⇌ CH₃OH (g)

How many moles of CO react completely with 6.0 × 10³ moles of H₂?

Tick (✔) one box

A chemist prepares some zinc iodide (ZnI₂) to use as a catalyst in the production of some organic compounds. 

Zn + I₂ ⟶ ZnI₂ 

Calculate the minimum mass of zinc that should be added to 0.600 g of iodine so that the iodine fully reacts. 

Relative atomic masses (A_r): Zn = 65; I = 127

Minimum mass of Zn = ____________________ g

Titanium is extracted from titanium dioxide in a two step industrial process:

      Step 1     TiO₂ + 2C + 2Cl₂ ⟶ TiCl₄ + 2CO

      Step 2     TiCl₄ + 4Na ⟶ Ti + 4NaCl

In step 2, 40 kg of TiCl₄ is added to 20 kg of Na.

The equation for the reaction is:

TiCl₄ + 4Na ⟶ Ti + 4NaCl

Relative atomic masses (A_r): Na = 23; Cl = 35.5; Ti = 48

Explain why titanium chloride is the limiting reactant.

You must show your working.

A solution of CuSO₄ has a concentration of 35 g/dm³.

What mass of CuSO₄ would there be in 250 cm³ of the solution?

Copper oxide (CuO) was reacted with methane (CH₄) in the apparatus shown below in Figure 1.

Figure 1

Balance the equation for the reaction: 

Use the equation to calculate the mass of copper that could be made from 6.0 g of copper oxide.

Mass of copper = _______________________ g

Iron (III) chloride is used in industrial waste treatment, in water purification and in the production of electronic circuit boards.

One method of producing iron (III) chloride is reacting iron with chlorine:

2Fe + 3 Cl₂ → 2FeCl₃

Calculate the maximum mass of FeCl₃ that can be produced from 14.0 g of iron. 

Relative atomic masses (A_r): Cl = 35.5; Fe = 56

Maximum mass of FeCl₃ = _______________ g

A student wanted to make 12.0 g of copper chloride by reacting copper carbonate with hydrochloric acid. The equation for the reaction is as follows: 

Relative atomic masses, A_r: H = 1; C = 12; O = 16; Cl = 35.5; Cu = 63.5

Calculate the mass of CuCO₃ the student needs to react with dilute hydrochloric acid to make 12.0 g of copper chloride.

Mass of copper carbonate = _________________________ g

This question is about the reaction of iron and chlorine.

A teacher uses the apparatus shown in Figure 1 to react iron with chlorine.

Figure 1

This is the teacher's method.

• Record the mass of the weighing boat
• Add iron and record the mass again
• Heat the weighing boat and iron strongly for two minutes and then allow to cool
• Record the mass of the weighing boat and its contents

Give an addition to the method to check that the iron has fully reacted with the chlorine.

The iron is the limiting reagent in the reaction.

Explain one resulting safety precaution for this reaction.

Table 1 shows the teacher's results.

Table 1

Use the teacher's results to show that the empirical formula of iron chloride is FeCl₃.

Relative atomic masses (A_r):     Fe = 56,     Cl = 35.5

Write the balanced symbol equation, including state symbols, for this reaction of iron and chlorine.

This question is about the redox reaction of tungsten oxide.

The word equation for the reaction is: 

tungsten oxide + hydrogen → tungsten + water

Explain, in terms of redox, the type of reaction that the tungsten oxide is undergoing.

A teacher uses this apparatus in Figure 1 to for the reaction of tungsten oxide with hydrogen.

Figure 1

The teacher wears appropriate eye protection and a lab coat during the demonstration.

Explain one other safety precaution the teacher should take.

The teacher weighed:

• the glass tube
• the glass tube and tungsten oxide before the reaction
• the glass tube and tungsten after the reaction.

Table 1 shows the teacher’s results.

Table 1

Calculate the simplest whole number ratio of:

moles of tungsten atoms : moles of oxygen atoms

Determine the balanced equation for the reaction.

Relative atomic masses (A_r):     W = 184,     O = 16 

This question is about the reaction of zinc with sulfuric acid.

3.75 g of zinc oxide, ZnO (s), was added to 150 cm³ of 98 g/dm³ of sulfuric acid (aq) producing a salt.

Write a balanced symbol equation for this reaction. 

Using the equation in part (a), deduce which is the limiting reactant in the reaction.

Show your full workings. 

Relative atomic masses (A_r):     Zn = 65,     O = 16,     H = 1,     S = 32

Use your answer to part (b), to calculate the amount, in grams, of the salt produced.

Give your answer to 3 significant figures.

Calculate the amount, in moles, of sulfuric acid at the end of the reaction.

Give your answer to 3 significant figures.

This question is about aluminium.

Aluminium can be extracted from aluminium oxide by electrolysis. 

The overall equation for the electrolysis of aluminium oxide is:

2Al₂O₃ → 4Al + 3O₂ 

Write the balanced half-equation, including state symbols, for the process at the negative electrode.

Calculate the mass of oxygen, in kg, produced when 5000 kg of aluminium oxide is completely electrolysed.

Relative atomic masses (A_r):     O = 16,     Al = 27

The thermite reaction of aluminium and iron oxide is used to weld train tracks together.

The equation for this reaction is:

2Al + Fe₂O₃ → 2Fe + Al₂O₃ 

A 5.00 kg thermite mixture contains 20% aluminium and 80% iron oxide.

Show that aluminium is the limiting reactant.

Relative atomic masses (A_r):     O = 16,     Al = 27,     Fe = 56

Explain why having aluminium as the limiting reactant in the thermite reaction is better for industrial applications.

Zinc-bromine batteries are rechargeable batteries that are being developed as an alternative to lithium-ion batteries.

They are less likely to overheat and catch fire compared to lithium-ion batteries.

The zinc-bromine battery has a zinc electrode and a carbon electrode with a solution of zinc bromide as the electrolyte.

Write the balanced symbol equation for the formation of zinc bromide, ZnBr₂, using hydrobromic acid.

Calculate the minimum mass of zinc that needs to be added to 5.00 cm³ of hydrobromic acid so that the hydrobromic acid fully reacts.

The density of hydrobromic acid is 1.49 g / cm³.

Relative atomic masses (A_r):     Zn = 65,     H = 1,     Br = 80

In a second experiment, 5.00 g of zinc powder was reacted with hydrobromic acid.

Calculate the number of atoms of hydrogen produced.