Reaction Calculations (Edexcel AS Chemistry): Exam Questions

This question is about sodium carbonate.

Sodium carbonate forms a number of hydrates with the general formula Na₂CO₃.xH₂O.

A 250 cm³ standard solution of one of these hydrates contained 10.0 g of the compound.

Describe, including the names of any relevant apparatus, how to make this standard solution when provided with 10.0 g of the hydrate in a beaker.

25.0 cm³ portions of the standard solution described in (a) are titrated with hydrochloric acid solution of concentration 0.300 mol dm^–3, using methyl orange as an indicator.

The table shows the results for this titration.

i) What is the colour change at the end-point of the reaction?

(1)

ii) State why the value for the total titre in Titration 1 should not be used to calculate the mean titre.

(1)

iii) Calculate the mean titre.

(1)

iv) Calculate the relative formula mass, M_r, of the hydrated sodium carbonate, Na₂CO₃.xH₂O.

The equation for the reaction in the titration is

Na₂CO₃ + 2HCl → 2NaCl + H₂O + CO₂

(4)

In an experiment, the M_r of a different hydrated sodium carbonate was found to be 286 g mol^–1.

i) Calculate the relative formula mass of anhydrous sodium carbonate, Na₂CO₃.

(1)

ii) Calculate the number of molecules of water of crystallisation, x, for this hydrated sodium carbonate, Na₂CO₃.xH₂O.

(1)

Sodium carbonate is manufactured from sodium chloride in a two-stage process.

NaCl + NH₃ + CO₂ + H₂O → NaHCO₃ + NH₄Cl                                   2NaHCO₃ → Na₂CO₃ + H₂O + CO₂

Calculate the maximum mass of sodium carbonate, Na₂CO₃, which could be obtained from 500 kg of sodium chloride.

The equation for the complete combustion of butane is

2C₄H₁₀ (g) + 13O₂ (g) → 8CO₂ (g) + 10H₂O (l)

What is the minimum volume of oxygen, at room temperature and pressure (r.t.p.), needed for the complete combustion of 0.200 mol of butane? [Molar volume of a gas at r.t.p. = 24.0 dm³ mol⁻¹]

This question is about the molar volume of gases.

i) Calculate the volume of one mole of an ideal gas, A, at 60 °C and 500 kPa pressure.

Give your answer to two significant figures and include units.

[The ideal gas equation is pV = nRT. Gas constant (R) = 8.31 J K⁻¹ mol⁻¹] 

(3)

ii) At room temperature and pressure (r.t.p) another gas B, with formula XH₃ , has a density of 1.42 g dm⁻³. 

Calculate the molar mass of the gas XH₃ and deduce the identity of the element X. 

[The molar volume of gas B = 24000 cm³ mol⁻¹ at r.t.p.]

(2)

The apparatus shown was used to measure the volume of gas evolved when a weighed mass of sodium carbonate reacted with dilute hydrochloric acid.

 The following procedure was used.

1. Solid sodium carbonate was placed in a container and weighed accurately.

2. The delivery tube and rubber bung were removed and the sodium carbonate was transferred to the test tube.

3. The container was then reweighed.

4. The syringe plunger was pushed in, to zero the syringe.

5. 10.0 cm³ of 0.400 mol dm⁻³ hydrochloric acid was then added to the sodium carbonate and the rubber bung and delivery tube rapidly replaced.

6. The mixture was shaken and, when the reaction had finished, the reading of the syringe was noted.

Results 

Mass of container and sodium carbonate before transfer = 20.135 g

 Mass of container after transfer of the sodium carbonate = 19.893 g

 Mass of sodium carbonate used = 0.242 g

 The equation for the reaction is

Na₂CO₃ (s) + 2HCl (aq) → 2NaCl (aq) + CO₂ (g) + H₂O (l)

i) Calculate the moles of hydrochloric acid and the moles of sodium carbonate used in this experiment.

 Use your answers to decide which reactant is in excess.

Calculate the maximum volume of carbon dioxide which could be produced.

 Molar mass of Na₂CO₃ = 106.0 g mol⁻¹

 Molar volume of gas = 24000 cm³ mol⁻¹ at r.t.p.

(5)

 ii) The actual volume of carbon dioxide collected was less than calculated.

Give two reasons for this. 

(2)

The reaction of sulfuric acid with potassium hydroxide is a neutralisation.

The equation for this reaction is

H₂SO₄ (aq) + 2KOH (aq) → K₂SO₄ (aq) + 2H₂O (l)

A titration was carried out using the following method.

1. Potassium hydroxide solution of unknown concentration was placed in a burette and the initial reading was recorded.

2. 25.0 cm³ of sulfuric acid solution, concentration 0.0800 mol dm⁻³, was transferred to a conical flask.

3. Three drops of phenolphthalein indicator were added to the sulfuric acid.

4. Potassium hydroxide was added from the burette until the solution just changed colour and then the burette reading was recorded.

5. Repeat titrations were carried out until concordant titres were obtained.

Select the most appropriate piece of apparatus to measure the 25.0 cm³ of sulfuric acid.

What is the colour of the solution when neutralisation has just occurred? 

i) Complete the table of results for titration number 1, using the diagrams to find the initial and final burette readings.

 (2)

Table of results 

ii) The best value for the mean titre of this reaction is 

(1)

iii) Calculate the concentration, in mol dm⁻³, of the potassium hydroxide solution, giving your answer to an appropriate number of significant figures. 

(3)

Many vehicles are fitted with airbags which provide a gas-filled safety cushion to protect the occupant of the vehicle if there is a crash.

The first reaction in airbags is the thermal decomposition of sodium azide, NaN₃, to form sodium and nitrogen gas.

i) Write the equation for this decomposition of sodium azide.

State symbols are not required.

(1)

ii) In the reaction in (a)(i), a typical airbag is inflated by about 67 dm³ of gas. Calculate the minimum mass of sodium azide, in grams, needed to produce this volume of gas. Use the Ideal Gas Equation and give your answer to an appropriate number of significant figures.

For the purpose of this calculation, assume that the temperature is 300 °C and the pressure is 140000 Pa.

(4)

The second reaction in the airbag is between the sodium produced in the reaction (a)(i) and potassium nitrate.

Balance the above equation, justifying your answer in terms of the changes in oxidation numbers.

The third reaction in the airbag is between the metal oxides and silicon dioxide.

State the type of reaction taking place and justify why this reaction is necessary.

The Maxwell-Boltzmann distribution diagram shows the molecular energies for the gaseous system immediately after the airbag has been deployed.

What is the change in shape of the curve when the airbag cools?

This question is about the reaction of magnesium with dilute hydrochloric acid.

Write an equation for the reaction of magnesium with hydrochloric acid.

Include state symbols.

The apparatus shown in the diagram can be used to collect the gas produced during the reaction of magnesium with dilute hydrochloric acid.

The following procedure was used.

i) A measuring cylinder was used to measure the 10.0 cm³ of dilute hydrochloric acid in Step 1. The uncertainty for a volume measurement is ± 0.5 cm³. Calculate the percentage uncertainty in the volume of hydrochloric acid.

(1)

ii) Determine which reactant is in excess by calculating the number of moles of magnesium and of hydrochloric acid used in the experiment.

(3)

iii) Calculate the maximum number of moles of gas that could be produced, using your answers to (a) and (b)(ii).

(1)

iv) Under the conditions of the experiment, the temperature was 23°C and the pressure 98 000 Pa.

Calculate the maximum volume of gas, in cm³, that could be produced using your answer in (b)(iii). Give your answer to an appropriate number of significant figures.

[The ideal gas equation is pV = nRT.  Gas constant (R) = 8.31 J mol^–1 K^–1]   

(4)

i) Deduce two possible reasons why the volume of gas collected in the experiment was smaller than that calculated in (b)(iv).

(2)

ii) Describe two changes to the procedure that would enable the volume of gas collected to be closer to that calculated in (b)(iv).

(2)

This question is about the titration of a weak acid with a strong base.

A standard solution of ethanedioic acid, which is a weak, diprotic acid, can be used to determine the concentration of a sodium hydroxide solution. 25.0 cm³ of the ethanedioic acid solution, with concentration 3.80 g dm⁻³, was pipetted into a conical flask. A few drops of indicator solution were added. The ethanedioic acid was titrated with the sodium hydroxide solution which was in the burette. The titration was repeated and the following results were obtained. 

[Molar mass of ethanedioic acid = 90.0 g mol⁻¹]

i) In the appropriate row, tick (✔) those titre values that should be used to find the mean, and use these titres to calculate it.

Write the value of the mean titre in the box provided in the table of results.

(2)

ii) Ethanedioic acid is a weak acid. Name a suitable indicator for this titration and state the colour change at the end-point.

(2)

Name of indicator ..................................................

 Colour change at the end-point from
.............................. to ..............................

iii) The equation for the reaction of ethanedioic acid with sodium hydroxide is 

C₂H₂O₄ + 2NaOH → C₂O₄Na₂ + 2H₂O

 Calculate the concentration of the sodium hydroxide solution, in mol dm⁻³.

Give your answer to three significant figures.

(4)

The uncertainty in each burette reading is ± 0.05 cm³. The uncertainty in the pipette volume is ± 0.06 cm³. 

i) Calculate the percentage uncertainties for titre 4, and the pipette volume.

(2) 

ii) Which of the following changes would halve the percentage uncertainty in the volume of liquid measured by the burette?

(1)

A group of students analysed a hydrated salt with the formula KH₃(C₂O₄)_y.zH₂O where y and z are whole numbers.

The students carried out experiments to determine the values of y and z.

Experiment 1 – to determine the value of y

One student was provided with a 0.0235 mol dm^-3 solution of the salt.

25.0 cm³ portions of the salt solution were acidified with excess dilute sulfuric acid and heated to about 60 °C.

Each portion was titrated with 0.0203 mol dm^-3 potassium manganate(VII).

The results of four titrations are shown in the table.

i) Complete the diagram to show the final burette reading in Titration 1.

(2)

ii) Explain why this student should use a mean titre of 23.15 cm³ and not 23.43 cm³ in the calculation.

(2)

iii) The uncertainty in each burette reading is ±0.05 cm³.

Calculate the percentage uncertainty in the titre volume of potassium manganate(VII) solution used in Titration 2.

(1)

iv) The equation for the reaction is

2MnO₄^– + 5C₂O₄^2– + 16H⁺ → 2Mn²⁺ + 10CO₂ + 8H₂O

Deduce, by calculation, the value of y, to the nearest whole number, in the formula KH₃(C₂O₄)y.zH₂O.

Use the mean titre of 23.15 cm³ and other data from Experiment 1.

You must show your working.

(4)

This question is about a titration experiment carried out by a group of students to determine the concentration of a solution of ethanoic acid using sodium hydroxide.

A student weighed about 4.00 g of sodium hydroxide pellets and added them to a beaker containing 50 cm³ of deionised water.

The mixture was stirred with a glass rod to dissolve the pellets and to give a homogenous solution.

The solution was poured through a funnel into a 250.0 cm³ volumetric flask and deionised water was added up to the mark and then the flask was shaken.

i) Describe how you would ensure that all the sodium hydroxide was transferred to the volumetric flask.

(2)

ii) A student adds deionised water above the mark and shakes the flask.

State why the procedure has to be restarted rather than using a teat pipette to remove the excess water.

(1)

Two students each cleaned a burette, then poured sodium hydroxide solution into their burettes.

i) Student 1 used a funnel to pour sodium hydroxide solution into the burette.

Give two steps needed before the student takes the initial burette reading.

(2)

ii) Student 2 cleaned the burette by rinsing it with deionised water immediately before filling it with the sodium hydroxide solution.

Give the effect, if any, on the value of the first titre. Justify your answer.

(1)

The sketch shows the pH changes during a titration of 25.0 cm³ of ethanoic acid with sodium hydroxide of the same concentration.

The ideal indicator for this titration will change colour on the addition of a very small volume of sodium hydroxide solution at a titre value very close to the equivalence point of the reaction.

i) Assess the suitability of methyl red as an indicator for this titration. Make use of the Data Booklet in answering this question.

(4)

ii) Complete the table, with a tick (✓) or a cross (✗), to show whether or not the indicator would be suitable for use in this titration.

(1)

Each student used a pipette to measure 25.0 cm³ of the ethanoic acid solution into four separate conical flasks and added an indicator.

The results of one student’s titrations are shown in the table.

i) Complete the table.

(1)

ii) The low titre for titration 4 was queried by the teacher. The student had wanted to refill the burette and continue the titration but had been told the measurement uncertainty would increase if this was done.

Calculate the total percentage measurement uncertainty if the burette had been refilled to 0.00, and then a further 0.30 cm³ had been added from the burette, to the conical flask.

The measurement uncertainty for each burette reading is ±0.05 cm³.

(1)

The teacher carried out the experiment and obtained the following results.

Mass of sodium hydroxide used to make 250.0 cm³  solution = 3.80 g

Volume of ethanoic acid solution = 25.00 cm³

Mean titre of sodium hydroxide = 11.90 cm³

CH₃COOH (aq) + NaOH (aq) → CH₃COONa (aq) + H₂O (l)

Calculate the concentration of the ethanoic acid solution in g dm⁻³.

Give your answer to an appropriate number of significant figures.

A student made crystals of a metal chloride, JCl₂.6H₂O, by reacting the metal carbonate, JCO₃ , with hydrochloric acid, HCl(aq). The product was purified.

Procedure

1. 150 cm³ of hydrochloric acid, concentration 0.80 mol dm⁻³, was transferred to a 400 cm³ conical flask. The flask was warmed gently using a Bunsen burner. A spatula measure (about 1.0 g) of metal carbonate was added to the acid.

2. When the reaction in Step 1 was finished, more metal carbonate was added until the metal carbonate was in excess.

3. The resulting mixture was filtered into an evaporating basin.

4. The evaporating basin was heated using a Bunsen burner to concentrate the solution. The concentrated solution was allowed to cool and crystallise.

5. Once crystal formation was complete, the resulting mixture was filtered for a second time.

6. The resulting white crystals were rinsed with a small volume of ice-cold water.

The equation for the reaction between the metal carbonate and hydrochloric acid is

JCO₃ (s) + 2HCl (aq) → JCl₂ (aq) + H₂O (l) + CO₂ (g)

i) Describe two observations that the student might make which show that the reaction in Step 1 has finished. 

(2)

 ii) State the purpose of the filtration in Step 3. 

(1)

 iii) Explain the use of a small volume of ice-cold water in Step 6. 

(2)

The student obtained a mass of 14.26 g of hydrated crystals.

Assuming that the percentage yield is 100%, use the information in the procedure to give a possible identity of J. 

The student was surprised by the white colour of the crystals of JCl₂.6H₂O in Step 6.

This did not agree with the possible identity for J from the calculation in (b).

The student decided to perform a flame test on the crystals. 

i) Explain why the student was surprised and decided to carry out a flame test. 

(2)

ii) The flame test colour was crimson red. Identify J. 

(1)

iii) Calculate the actual percentage yield of the reaction, which produced 14.26 g of crystals.

 Give your answer to two significant figures. 

(2)

Malachite is a green mineral with the formula Cu₂CO₃(OH)₂ . It has a molar mass of 221 g mol^–1.

What is the percentage by mass of copper in pure malachite?

Describe what you would expect to see when an excess of dilute hydrochloric acid is added to a sample of pure solid malachite.

i) Describe how you would carry out a flame test on a sample of powdered malachite.

(3)

ii) When the atoms of some elements are heated, they produce a characteristic flame colour. For example, the copper in malachite gives a blue-green colour. Explain how atoms of different elements can produce different characteristic flame colours when heated.

(4)

i) When malachite is heated to approximately 300 °C, water, carbon dioxide and copper(II) oxide are formed.

The equation for this decomposition is

Cu₂CO₃(OH)₂ → 2CuO + CO₂ + H₂O

Calculate the maximum volume of carbon dioxide that could be produced when 0.810 g of malachite is thermally decomposed.

Assume that the gas is collected at a temperature of 25 °C and 101 kPa pressure.

Give your answer to an appropriate number of significant figures and state the units.[The ideal gas equation is pV = nRT. Gas constant (R) = 8.31 J mol^–1K^–1]

(5)

ii) The gas was collected in a gas syringe with a stated accuracy of ±0.5 cm³.

Calculate the percentage uncertainty in the volume of gas collected.

(1)

iii) Malachite ore is a mixture of malachite and rock. A 0.810g sample of malachite ore was thermally decomposed, producing 0.571g of copper(II) oxide.

Calculate the percentage purity of this malachite ore sample.

Give your answer to an appropriate number of significant figures.

(3)

Wine and gin are aqueous solutions of ethanol with traces of other organic compounds which give these drinks their characteristic flavours and aromas.

When a bottle of wine is opened, oxidation of the ethanol in the wine produces ethanoic acid.

An experiment was carried out to determine the percentage of the ethanol that had been oxidised.

• A bottle of white wine, with an ethanol concentration of 2.50 mol dm^–3,was opened and left to stand at room temperature for three weeks.

• A 25.0 cm³ sample of the wine was transferred to a conical flask and phenolphthalein indicator added.

• Aqueous sodium hydroxide of concentration 0.235 mol dm^–3 was added from a burette until the colour of the indicator permanently changed.

• The titration was repeated and the titre values, in cm³, were 27.90, 26.75 and 26.85.

The equation for the neutralisation reaction is

CH₃COOH + NaOH → CH₃COO^–Na+ + H₂O

i) Name the piece of apparatus used to measure 25.0 cm³ of wine.

(1)

ii) To improve the accuracy, the burette should be rinsed.

State what should be used.

(1)

iii) What is the colour change at the end-point?

(1)

iv) State what is meant by the term ‘concordant results’ as applied to a titration experiment.

(1)

v) Calculate the percentage of ethanol that has oxidised, given that one mole of ethanol forms one mole of ethanoic acid.

Give your answer to an appropriate number of significant figures.

(5)

vi) Deduce why this method would not be effective for the analysis of the acid content of a red wine.

(1)

The ethanol content of alcoholic drinks is usually measured as the percentage of alcohol by volume (ABV).

ABV = × 100

The ABV values of four different brands of gin are shown.

A sample of one of these gins was found to contain an ethanol concentration of 7.50 mol dm^–3.

By calculating the percentage of ethanol by volume (ABV) of this sample, deduce the brand of this gin.

[Assume the density of ethanol, C₂H₅OH = 0.79 g cm^–3]

A student determines the percentage purity of a sample of calcium carbonate, CaCO₃, using a back-titration.

The student adds an excess of 2.00 mol dm^-3 hydrochloric acid to 2.50 g of the impure calcium carbonate sample. The excess acid is then titrated against 0.100 mol dm^-3 sodium hydroxide solution.

The equation for the reaction between calcium carbonate and hydrochloric acid is:

CaCO₃ (s) + 2HCl (aq) → CaCl₂ (aq) + H₂O (l) + CO₂ (g)

The student uses 25.0 cm³ of 2.00 mol dm^-3 hydrochloric acid and obtains the following titration results.

Identify the concordant titres and calculate the mean titre.

Calculate the percentage purity of the calcium carbonate sample.

The uncertainty in each burette reading is ±0.05 cm³.

Calculate the percentage uncertainty in the mean titre.

A second student carries out the same procedure but begins the titration before all of the calcium carbonate solid has fully dissolved in the hydrochloric acid.

Explain how this procedural error would affect the calculated percentage purity of the calcium carbonate.

Before carrying out the back-titration, the student washes the conical flask with distilled water but does not dry it.

Explain how, if at all, using a wet conical flask affects the value of the titre.