Amounts of substance in equations

This question is about ammonia.

Nitrogen gas reacts with hydrogen gas to produce ammonia gas in a reversible reaction.

The equation for the reaction is:

N₂(g) + 3 H₂(g) 2 NH₃(g)

What is the effect of increasing the pressure on the equilibrium position in this reaction?

Tick (✓) one box.

The equilibrium position shifts towards the side with the smaller number of molecules.

The equilibrium position is unchanged.

The equilibrium position shifts towards the side with the larger number of molecules.

The reaction to produce ammonia gas is exothermic.

What is the effect of increasing the temperature on the relative amount of ammonia at equilibrium?

Tick (✓) one box.

The relative amount of ammonia decreases.

The relative amount of ammonia stays the same.

The relative amount of ammonia increases.

How must the apparatus for the reaction be designed so that dynamic equilibrium can be reached?

Explain what happens to the relative amounts of nitrogen, hydrogen and ammonia after dynamic equilibrium has been reached.

Ammonia reacts with hydrogen chloride to produce ammonium chloride (NH₄Cl).

The equation for the reaction is:

6.8 g of ammonia reacts with excess hydrogen chloride.

Explain why ammonia is described as the limiting reactant in this reaction.

Calculate the maximum mass of ammonium chloride that could be produced from 6.8 g of ammonia.

Relative atomic masses (A_r): N = 14 H = 1 Cl = 35.5

Magnesium and hydrochloric acid react to produce magnesium chloride and hydrogen.

Give the test for hydrogen gas.

Give the result of the test if hydrogen gas is present.

During the reaction between magnesium and hydrochloric acid, hydrogen ions form hydrogen molecules.

The half equation for this process is:

What type of reaction does this half equation show?

Give the reason for your answer.

0.72 g of magnesium reacted completely with dilute hydrochloric acid.

The equation for the reaction is:

Calculate the mass of magnesium chloride produced.

Relative atomic masses (A_r): Mg = 24 Cl = 35.5

This question is about the industrial process used to produce ammonia.

Ammonia is produced from nitrogen and hydrogen.

Complete the dot and cross diagram for the ammonia molecule shown in Figure 11.

Show only the electrons in the outer shell of each atom.

Hydrogen is obtained from natural gas.

One stage in the process is to react carbon monoxide with steam.

The equation for the reaction is:

The forward reaction is exothermic.

In the industrial process, the reaction occurs at approximately 400 °C

The reaction between carbon monoxide and steam reaches dynamic equilibrium.

Which statement describes dynamic equilibrium?

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The reaction between carbon monoxide and steam is faster at 600 °C than at 400 °C

Explain why a temperature of 600 °C is not used in the industrial process.

Predict the effect of increasing the pressure on the yield of hydrogen.

Give one reason for your answer.

Give two reasons why a catalyst is used in the reaction between carbon monoxide and steam.

The carbon dioxide produced in this reaction can be used in fizzy drinks.

In fizzy drinks carbon dioxide gas dissolves in water to form an aqueous solution of carbonic acid.

Carbonic acid is a weak acid.

Explain the difference between a weak acid and a strong acid.

The equation for the production of ammonia from nitrogen and hydrogen is:

1.275 × 10¹¹ kg of ammonia is used each year to manufacture fertilisers.

This is 85% of the ammonia produced each year.

Calculate the mass of nitrogen needed to manufacture the total mass of ammonia produced each year.

Relative atomic masses (A_r): H = 1 N = 14

A student planned to make copper sulfate crystals from excess copper oxide and dilute sulfuric acid.

The equation for the reaction is:

Why is it necessary to add excess copper oxide?

This is the method used.

1. Add 25 cm³ of dilute sulfuric acid to a conical flask.

2. Gently warm the dilute sulfuric acid.

3. Add excess copper oxide to the dilute sulfuric acid.

4. Stir the mixture.

5. Heat to evaporate all the water from the mixture.

Suggest two improvements to the method.

Explain why each improvement is needed.

The student used:

• 2 g of copper oxide (in excess)

• 25 cm³ of a solution of dilute sulfuric acid with a concentration of 49 g/dm³

Determine by how many moles the copper oxide (CuO) was in excess.

Relative atomic masses (A_r): Cu = 63.5 O = 16

Relative formula mass (M_r) of sulfuric acid = 98

A teacher demonstrated the extraction of copper from copper oxide.

This is the method used.

1. Mix 1.30 g of zinc and 1.59 g of copper oxide.

2. Heat the mixture strongly.

3. When the mixture starts to glow, stop heating.

4. Let the glow spread through the mixture.

5. Leave the mixture to cool.

6. Add hydrochloric acid to the cooled mixture.

7. Filter the mixture obtained in step 6.

A student concluded that an exothermic reaction had taken place.

Explain how an observation made during the demonstration shows this.

The equation for the reaction between zinc and copper oxide is:

1.59 g of copper oxide reacted.

Calculate the mass of copper produced.

Relative atomic masses (A_r): Cu = 63.5 O = 16 Zn = 65

Explain why steps 6 and 7 result in only copper being obtained as the residue.

The ionic equation for the reaction is:

Which statement about the reaction between zinc and copper ions is correct?

Tick one box.