Percentage yield

When iron rusts, iron(III) oxide forms.

State two conditions needed for iron to rust.

What is the name of the process used to coat iron with zinc?

What is the correct order of reactivity of these four metals?

Describe a test to show that a solution contains Fe³⁺ ions.

test

result

A sample of iron(III) oxide reacts with excess carbon monoxide.

This is the equation for the reaction.

Explain why this is a redox reaction.

Calculate the minimum mass of iron(III) oxide needed to produce a theoretical yield of 28 g of iron.

[for Fe₂O₃ M_r = 160]

The actual yield of iron from this sample is 21 g.

Calculate the percentage yield of iron.

A student does a titration to find the volume of phosphoric acid solution (H₃PO₄) needed to neutralise a sample of potassium hydroxide solution (KOH).

One of the products of the reaction is potassium phosphate.

This is the student’s method.

• use a measuring cylinder to transfer 25 cm³ of potassium hydroxide solution to a conical flask

• add two drops of an indicator

• fill a burette with phosphoric acid solution and record the initial reading on the burette

• add the acid to the conical flask

• record the final burette reading when the indicator changes colour

This is the equation for the reaction.

Explain two improvements to the student’s method that would give more accurate results.

The student repeats the titration correctly.

The potassium hydroxide solution has a concentration of 0.150 mol / dm³.

The phosphoric acid solution has a concentration of 0.0400 mol / dm³.

Calculate the volume, in cm³, of phosphoric acid solution needed to neutralise 25.0 cm³ of 0.150 mol / dm³ potassium hydroxide solution.

Give your answer to three significant figures.

Potassium phosphate solution reacts with calcium chloride solution.

One of the products is the insoluble solid, calcium phosphate.

Complete the equation for the reaction.

The student reacts a solution containing 0.0400 mol of potassium phosphate with an excess of calcium chloride solution.

The student obtains 5.70 g of calcium phosphate.

Show that the percentage yield for the reaction is about 90%

[for Ca₃(PO₄)₂ M_r = 310]

Describe how the student can obtain a pure, dry sample of calcium phosphate from the mixture at the end of the reaction.

This question is about the manufacture of ethanol by the fermentation of glucose.

Fermentation needs to be done in the absence of air.

Give two other conditions needed for fermentation.

Explain why fermentation needs to be done in the absence of air.

Complete the equation for fermentation.

A mixture containing 500 mol of glucose is fermented.

A mass of 5750 g of ethanol is obtained.

Calculate the percentage yield.

[for ethanol, M_r = 46]

This method is used in an experiment to produce hydrated zinc nitrate crystals.

• pour 50 cm³ of dilute nitric acid into a beaker

• add a spatula of zinc powder to the acid

• add more zinc until it is in excess

• filter the mixture

• obtain crystals of zinc nitrate from the filtrate

State why the mixture is filtered.

Describe how a pure, dry sample of hydrated zinc nitrate crystals could be obtained from the filtrate.

This equation represents the formation of hydrated zinc nitrate in the experiment.

In another experiment, 9.75 g of zinc is completely reacted with nitric acid.

Show that the maximum possible mass of hydrated zinc nitrate crystals that could be formed is approximately 45 g.

[for Zn(NO₃)₂.6H₂O, M_r = 297]

The actual yield of hydrated zinc nitrate crystals is 36.4 g.

Calculate the percentage yield of hydrated zinc nitrate crystals.

This question is about the insoluble salt, lead(II) bromide.

Lead(II) bromide can be made by a precipitation reaction.

This is the equation for the reaction.

Pb(NO₃)₂(aq) + 2KBr(aq) PbBr₂(s) + 2KNO₃(aq)

Describe how solutions of lead nitrate and potassium bromide can be used to make a pure, dry sample of lead(II) bromide.

A solution containing 0.150 mol of lead(II) nitrate is reacted with an excess of potassium bromide solution.

A mass of 49.6 g of pure, dry lead(II) bromide is produced.

Show, by calculation, that the percentage yield of lead(II) bromide is 90.1%.

[for PbBr₂, M_r = 367]

A student investigates the change in electrical conductivity as dilute lead(II) nitrate solution is added to dilute potassium bromide solution.

This is the student’s method.

Step 1add 50 cm³ of potassium bromide solution to a beaker

Step 2 measure the electrical conductivity of the solution

Step 3add 10 cm³ of lead(II) nitrate solution to the beaker

Step 4 stir the mixture

Step 5 measure the electrical conductivity of the mixture

Repeat steps 3, 4 and 5 until a total of 50 cm³ of lead(II) nitrate solution has been added.

The table shows the student’s results.

Plot the student’s results on the grid.

Draw a line of best fit, ignoring the anomalous result.

Total volume of lead(II) nitrate in cm³

Explain the shape of the graph.

Suggest a mistake the student could have made to cause the anomalous result.

Further 10 cm³ volumes of lead(II) nitrate are added to the beaker so the lead(II) nitrate is in excess.

Predict what will happen to the conductivity of the mixture when the lead(II) nitrate is in excess.

The diagram shows the electrolysis of molten lead(II) bromide, PbBr₂

This is the ionic half‑equation for the formation of bromine at electrode A.

Give a reason why this half‑equation shows oxidation.

A teacher uses this apparatus to reduce tungsten oxide.

This is the teacher’s method.

• record the mass of a weighing boat

• add tungsten oxide and record the mass again

• heat the weighing boat and tungsten oxide strongly for two minutes and then allow to cool

• record the mass of the weighing boat and its contents

Complete the equation by adding the state symbols.

Give an addition to the method to check that the tungsten oxide has been completely reduced.

The table shows the teacher’s results.

Use the teacher’s results to show that the empirical formula of tungsten oxide is WO₃

[for tungsten, A_r = 184 for oxygen, A_r = 16]

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

Give one other safety precaution the teacher should take.

In industry, tungsten oxide is reduced on a large scale using hydrogen.

The percentage yield of tungsten is 73.5%

This is the equation for the reaction.

Calculate the mass, in tonnes, of tungsten that is produced when 2784 tonnes of tungsten oxide are reacted with an excess of hydrogen.

[1 tonne = 110⁶g]

[for tungsten, A_r = 184 for oxygen, A_r = 16]

mass of tungsten =_........ tonnes

A student uses this apparatus to heat crystals of hydrated zinc sulfate and collect the liquid produced.

Describe a chemical test to show that the colourless liquid contains water.

Describe a physical test to show the colourless liquid is pure water.

The equation for the decomposition of hydrated zinc sulfate is

The student records these masses.

mass of boiling tube = 41.64 g

mass of boiling tube + ZnSO₄·7H₂O = 54.46 g

Calculate the maximum volume, in cm³, of pure water that could be produced.

Give your answer to 1 decimal place.

[1.00 cm³ of pure water has a mass of 1.00 g]

[M_r of ZnSO₄·7H₂O = 287 M_r of H₂O = 18]

In an experiment using a different mass of ZnSO₄·7H₂O the maximum volume of pure water that could be produced is 8.5 cm³.

The student collected the pure water and calculated the percentage yield to be 20.3 %.

Calculate the volume, in cm³, of pure water collected.

Explain an improvement to the apparatus that would increase the percentage yield of pure water.

This question is about salts.

Soluble salts can be prepared by the reaction between a metal oxide and an acid.

The equation for this type of reaction is

metal oxide + acid salt + water

State the name given to this type of reaction.

State, in terms of protons, what happens in this reaction.

A student is given 50 cm³ of dilute sulfuric acid and a bottle of solid copper(II) carbonate.

Describe the method that the student should use to prepare a saturated solution of copper(II) sulfate.

In your answer, refer to the pieces of apparatus that the student should use.

The student produces dry crystals of hydrated copper(II) sulfate from the saturated solution.

He calculates that 6.40 g of dry crystals should be formed.

The mass of dry crystals he actually obtains is 1.80 g less than he calculated.

Calculate the student’s percentage yield.

Give your answer to one decimal place.

Gypsum is hydrated calcium sulfate.

A sample of gypsum contains 79% of calcium sulfate by mass.

Calculate the value of x in CaSO₄.xH₂O

Describe a test for calcium ions in the sample of gypsum.

This question is about the insoluble salt silver chloride (AgC l).

Silver chloride can be made by the reaction between copper(II) chloride and silver nitrate.

Describe how a student could prepare a pure, dry sample of silver chloride starting with copper(II) chloride solution and silver nitrate solution.

A student investigates the quantity of silver chloride produced when different volumes of silver nitrate solution are added to copper(II) chloride solution.

This is the student’s method. pour 5.0 cm³ of copper(II) chloride solution into a test tube add 1.0 cm³ of silver nitrate solution to the test tube

allow the silver chloride precipitate to settle

measure the height of the precipitate

The student repeats the method using different volumes of silver nitrate solution.

The table shows the student’s results.

Plot the student’s results.

Draw two straight lines of best fit, ignoring the anomalous result.

Volume of silver nitrate added in cm³

Suggest a mistake the student made to cause the anomalous result.

Give a reason why the last three heights are the same.

The equation for the reaction between copper(II) chloride and silver nitrate is

CuCl₂(aq) + 2AgNO₃(aq) 2AgCl(s) + Cu(NO₃)₂(aq)

A student measures 25.0 cm³ of 0.500 mol/dm³ copper(II) chloride solution and reacts it with silver nitrate solution.

Name a piece of apparatus suitable for measuring 25.0 cm³ of copper(II) chloride solution.

Calculate the maximum mass, in grams, of silver chloride that could be produced.

[M_r of AgCl = 143.5]

In an experiment using different solutions, the mass of silver chloride produced is 0.744 g.

The maximum mass of silver chloride that could be produced is 0.850 g.

Calculate the percentage yield.

A student uses the reaction between zinc and dilute sulfuric acid to prepare some zinc sulfate crystals.

Complete the equation for this reaction by giving the correct state symbols.

State what would be observed during this reaction.

The student adds excess zinc to a beaker of dilute sulfuric acid.

Explain why it is necessary to add excess zinc.

Draw a diagram of the apparatus the student should use to remove the unreacted zinc and collect the zinc sulfate solution.

The student obtains a pure, dry sample of zinc sulfate crystals.

The formula of zinc sulfate crystals is ZnSO₄.7H₂O

Calculate the relative molecular mass (M_r) of zinc sulfate crystals.

The student uses 0.0200 mol of dilute sulfuric acid in her preparation.

Show that the maximum mass of zinc sulfate crystals that the student could obtain is about 6 g.

The student obtains a mass of 4.28 g of zinc sulfate crystals.

Calculate the percentage yield of the zinc sulfate crystals.

Give your answer to three significant figures.

Sodium hydrogencarbonate (NaHCO₃) is also known as baking soda.

Baking soda can be used to make cakes increase in size in an oven.

This is the equation for the reaction that takes place when baking soda is heated.

What type of reaction is this?

Suggest why the reaction makes the cakes increase in size.

A student uses this apparatus to investigate the reaction that takes place when sodium hydrogencarbonate is heated.

This is the student’s method.

● weigh a crucible and record the mass

● add some sodium hydrogencarbonate to the crucible, reweigh it and record the mass

● heat the crucible and contents for five minutes, then allow to cool before weighing and recording the mass

● heat the crucible and contents again for a further three minutes, then allow to cool before weighing and recording the mass

Give a reason why the crucible and contents are heated for a further three minutes.

The student considered using a lid on the crucible in the experiment.

Suggest an advantage and a disadvantage of using a lid on the crucible.

The table shows some of the student’s results.

Calculate the mass of sodium hydrogencarbonate that the student uses.

Using this equation, calculate the maximum mass of sodium carbonate (Na₂CO₃) that could form in the student’s reaction.

[M_r of NaHCO₃ = 84 M_r of Na₂CO₃ = 106]

In a second experiment, the student uses a larger mass of sodium hydrogencarbonate.

She calculates that she should obtain 4.8 g of sodium carbonate.

She actually obtains 4.2 g of sodium carbonate.

Calculate the percentage yield from the student’s experiment.

Other than spillages, suggest a possible reason why the student’s actual yield is less than expected.

Complete the chemical equation for the reaction in stage 1.

Give the meaning of the symbol

State the purpose of the platinum catalyst.

Give a chemical equation for the reaction of nitrogen monoxide and oxygen in stage 2.

The equation for the reaction in stage 3 is

Calculate the maximum mass, in tonnes, of nitric acid that could be produced in this reaction from 11.5 tonnes of nitrogen dioxide.

mass of nitric acid =_........ tonnes

Suggest what use can be made of the nitrogen monoxide gas formed in stage 3.

When copper(II) oxide reacts with dilute nitric acid, copper(II) nitrate is produced.

The equation for the reaction is

0.200 mol of nitric acid reacts with excess copper(II) oxide.

A mass of 15.3 g of copper(II) nitrate is produced.

Calculate the percentage yield of copper(II) nitrate.

[M_r of copper(II) nitrate = 187.5]

(a) Ethanol, C₂H₅OH, can be produced by the fermentation of glucose, C₆H₁₂O₆

Complete the equation for the fermentation of glucose.

State why it is necessary for fermentation to be done in the absence of air.

Explain why the temperature should not be higher than 40 °C.

When 4 mol of glucose is fermented, a mass of 55.2 g of ethanol is produced.

Show that the percentage yield of ethanol is 15%.

[M_r of C₂H₅OH = 46]

Ethanol can also be produced by the reaction between ethene and steam.

The equation for the reaction is

This reaction is in dynamic equilibrium.

Give two features of a reaction in dynamic equilibrium.

When the equilibrium mixture is heated, the yield of ethanol decreases.

Explain whether the forward reaction is exothermic or endothermic.

Carboxylic acids react with alcohols to form esters.

The displayed formula of an ester is

H H O

H H H H

Carboxylic acid A and alcohol B react to produce this ester.

Give the displayed formula of carboxylic acid A and of alcohol B.

Indicators can be used to test for carboxylic acids.

Describe a different chemical test for a carboxylic acid.