The Mole & the Avogadro Constant (Cambridge (CIE) IGCSE Chemistry)

Dilute hydrochloric acid reacts with sodium carbonate solution.

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

Explain why effervescence is seen during the reaction.

Separate: Chemistry and Extended Only

Dilute hydrochloric acid was titrated with sodium carbonate solution.

• 10.0 cm³ of 0.100 mol / dm³ hydrochloric acid were placed in a conical flask.

• A few drops of methyl orange indicator were added to the dilute hydrochloric acid

• The mixture was titrated with sodium carbonate solution.

• 16.2 cm³ of sodium carbonate solution were required to react completely with the acid.

i) What colour would the methyl orange indicator be in the hydrochloric acid?

[1]

ii) Calculate how many moles of hydrochloric acid were used.

[1]

iii) Use your answer to (b)(ii) and the equation for the reaction to calculate the number of moles of sodium carbonate that reacted.

[1]

iv) Use your answer to (b)(iii) to calculate the concentration of the sodium carbonate solution in mol / dm³.

[2]

Extended Only

In another experiment, 0.020 mol of sodium carbonate were reacted with excess hydrochloric acid.

Calculate the maximum volume (at r.t.p.) of carbon dioxide gas that could be made in this reaction.

Separate: Chemistry and Extended Only

Two salts can be made from potassium hydroxide and sulfuric acid. They are potassium sulfate, K₂SO₄, and the acid salt potassium hydrogen sulfate, KHSO₄. They are both made by titration.

25.0 cm³ of potassium hydroxide, concentration of 2.53 mol / dm³, was neutralised by 28.2 cm³ of dilute sulfuric acid.

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

Calculate the concentration of the sulfuric acid. number of moles of KOH used = ............................ number of moles of H₂SO₄ needed to neutralise the KOH = ............................ concentration of dilute sulfuric acid = ............................ mol / dm³

In the conical flask there is a neutral solution of potassium sulfate which still contains the indicator used in the titration.

i) Describe how you could obtain a solution of potassium sulfate without the indicator.

[2]

ii) Potassium hydrogen sulfate can be made by the following reaction.

KOH (aq) + H₂SO₄ (aq) → KHSO₄ (aq) + H₂O (l)

Suggest how you could make a solution of potassium hydrogen sulfate without using an indicator.

[2]

Describe a test which would distinguish between aqueous solutions of potassium sulfate and sulfuric acid.

   Test:

   Result:

Extended Only

Define the following

i) The mole

[1]

ii) The Avogadro constant

[1]

Extended Only

Which two of the following contain the same number of molecules? Show how you arrived at your answer.    2.0 g of methane, CH₄    8.0 g of oxygen, O₂    2.0 g of ozone, O₃    8.0 g of sulfur dioxide, SO₂

Extended Only

4.8 g of calcium is added to 3.6 g of water. The following reaction occurs.

Ca + 2H₂O → Ca(OH)₂ + H₂

i) The number of moles of Ca = .................... The number of moles of H₂O = ....................

[1]

ii) Which reagent is in excess? Explain your choice.

[2]

iii) Calculate the mass of the reagent named in (ii) which remained at the end of the experiment.

[1]

Separate: Chemistry and Extended Only

Quantities of chemicals, expressed in moles, can be used to find the formula of a compound, to establish an equation and to determine reacting masses.

A compound contains 72% magnesium and 28% nitrogen. What is its empirical formula?

A compound, Al₄C₃, contains only aluminium and carbon. 0.03 moles of this compound reacted with excess water to form 0.12 moles of Al(OH)₃ and 0.09 moles of CH₄.

Write a balanced equation for this reaction.

Extended Only

0.07 moles of silicon reacts with 25 g of bromine.

Si + 2Br₂ → SiBr₄

i) Which one is the limiting reagent? Explain your choice.

[3]

ii) How many moles of SiBr₄ are formed?

[1]

Extended Only

Chemists use the concept of the mole to calculate the amounts of chemicals involved in a reaction.

Define mole.

Extended Only

3.0 g of magnesium was added to 12.0 g of ethanoic acid.

Mg + 2CH₃COOH → (CH₃COO)₂Mg + H₂

The mass of one mole of Mg is 24 g.

The mass of one mole of CH₃COOH is 60 g.

i) Which one, magnesium or ethanoic acid, is in excess? You must show your reasoning.

[3]

ii) How many moles of hydrogen were formed?

[1]

iii) Calculate the volume of hydrogen formed, measured at r.t.p.

[2]

Separate: Chemistry and Extended Only

In an experiment, 25.0 cm³ of aqueous sodium hydroxide, 0.4 mol / dm³, was neutralised by 20.0 cm³ of aqueous oxalic acid, H₂C₂O₄.

2NaOH + H₂C₂O₄ → Na₂C₂O₄ +2H₂O

i) Calculate the number of moles of NaOH in 25.0 cm³ of 0.4 mol / dm³ solution.

[1]

ii) Use your answer to (i) and the mole ratio in the equation to find out the number of moles of H₂C₂O₄ in 20 cm³ of solution.

[1]

iii) Calculate the concentration, mol / dm³, of the aqueous oxalic acid.

[2]

Extended Only

Iron(III) sulphate decomposes when heated. Calculate the mass of iron(III) oxide formed and the volume of sulphur trioxide produced when 10.0 g of iron(III) sulphate was heated.  Mass of one mole of Fe₂(SO₄)₃ is 400 g.

Fe₂(SO₄)₃ (s)→ Fe₂O₃ (s) + 3SO₃ (g)

Number of moles of Fe₂(SO₄)₃ =

Number of moles of Fe₂O₃ formed = 

Mass of iron(III) oxide formed in g =

Number of moles of SO₃ produced = 

Volume of sulphur trioxide at r.t.p. in dm³ =

[5]

Separate: Chemistry and Extended Only

Sulphur trioxide can be made from sulphur dioxide.

i) Why is this reaction important industrially?

[1]

ii) Complete the word equation. sulphur dioxide + .............................................. → sulphur trioxide

[1]

iii) What are the conditions for this reaction?

[2]

Separate: Chemistry and Extended Only

Sulphur dioxide is easily oxidised in the presence of water.

SO₂ + 2H₂O – 2e^– → SO₄^2- + 4H⁺

What colour change would be observed when an excess of aqueous sulphur dioxide is added to an acidic solution of potassium manganate(VII)?

Extended Only

Sulphur dioxide reacts with chlorine in an addition reaction to form sulphuryl chloride.

SO₂ + Cl₂ → SO₂Cl₂

8.0 g of sulphur dioxide was mixed with 14.2 g of chlorine. The mass of one mole of SO₂Cl₂ is 135 g.

Calculate the mass of sulphuryl chloride formed by this mixture.

Calculate the number of moles of SO₂ in the mixture = 

Calculate the number of moles of Cl₂ in the mixture = 

Which reagent was not in excess? 

How many moles of SO₂Cl₂ were formed =

Calculate the mass of sulphuryl chloride formed in g =

Separate: Chemistry and Extended Only

A compound, X, contains 55.85% carbon, 6.97% hydrogen and 37.18% oxygen.

i) How does this prove that compound X contains only carbon, hydrogen and oxygen?

[1]

ii) Use the above percentages to calculate the empirical formula of compound X.

[2]

iii) The M_r of X is 86. What is its molecular formula?

[2]

Separate: Chemistry Only

i) Bromine water changes from brown to colourless when added to X.

What does this tell you about the structure of X?

[1]

ii) Magnesium powder reacts with an aqueous solution of X. Hydrogen is evolved. What does this tell you about the structure of X?

[1]

iii) X contains two different functional groups. Draw a structural formula of X.

[1]

Separate: Chemistry and Extended Only

Titanium is a transition element. It is isolated by the following reactions.

titanium ore → titanium(IV) oxide → titanium(IV) chloride → titanium

                              TiO₂                          TiCl₄                  Ti

Why is it usually necessary to include a number in the name of the compounds of transition elements?

Separate: Chemistry and Extended Only

Titanium(IV) chloride is made by heating the oxide with coke and chlorine.

TiO₂ + 2Cl₂ TiCl₄ + O₂

2C + O₂  2CO

Explain why the presence of coke ensures the maximum yield of the metal chloride.

Separate: Chemistry and Extended Only

Explain why the change, titanium(IV) chloride to titanium, is reduction.

Separate: Chemistry and Extended Only

The titanium ore contains 36.8% iron, 31.6% titanium and the remainder is oxygen.

i) Determine the percentage of oxygen in this titanium compound.

[1]

ii) Calculate the number of moles of atoms for each element. The number of moles of Fe is shown as an example. Number of moles of Fe = 36.8/56 = 0.66

[1]

iii) What is the simplest ratio for the moles of atoms, Fe : Ti : O?

[1]

iv) What is the formula of this titanium compound?

[1]

Separate: Chemistry and Extended Only

The food additive E220 is sulfur dioxide. It is a preservative for a variety of foods and drinks.

Sulfur dioxide is a reductant (reducing agent). Describe what you would see when aqueous sulfur dioxide is added to acidified potassium manganate(VII).

Extended Only

Sulfur dioxide can also be made by the reaction between a sulfite and an acid.

Na₂SO₃ + 2HCl → 2NaCl + SO₂ + H₂O

Excess hydrochloric acid was added to 3.15 g of sodium sulfite. Calculate the maximum volume, measured at r.t.p., of sulfur dioxide which could be formed. The mass of one mole of Na₂SO₃ is 126 g.

Across the world, food safety agencies are investigating the presence of minute traces of the toxic hydrocarbon, benzene, in soft drinks. It is formed by the reduction of sodium benzoate by vitamin C.

Sodium benzoate is a salt, it has the formula C₆H₅COONa. It can be made by the neutralisation of benzoic acid by sodium hydroxide.

i) Deduce the formula of benzoic acid.

[1]

ii) Write a word equation for the reaction between benzoic acid and sodium hydroxide.

[1]

iii) Name two other compounds that would react with benzoic acid to form sodium benzoate.

[2]

Separate: Chemistry and Extended Only

Benzene contains 92.3% of carbon and its relative molecular mass is 78.

i) What is the percentage of hydrogen in benzene?

[1]

ii) Calculate the ratio of moles of C atoms: moles of H atoms in benzene.

[1]

iii) Calculate its empirical formula and then its molecular formula. The empirical formula of benzene is ................... The molecular formula of benzene is ...................

[2]

The structural formula of Vitamin C is drawn below.

i) What is its molecular formula?

[1]

ii) Name the two functional groups which are circled.

[2]

The alkanes are a family of saturated hydrocarbons. Their reactions include combustion, cracking and substitution. 

i) What is meant by the term hydrocarbon?

[1]

ii) What is meant by the term saturated?

[1]

Separate: Chemistry Only

i) What is the general formula for the homologous series of alkanes?

[1]

ii) Calculate the mass of one mole of an alkane with 14 carbon atoms.

[2]

Extended Only

The complete combustion of hydrocarbons produces carbon dioxide and water only.

i) Write the equation for the complete combustion of nonane, C₉H₂₀.

[2]

 ii) 20cm³ of a gaseous hydrocarbon was mixed with an excess of oxygen, 200 cm³. The mixture was ignited. After cooling, 40 cm³ of oxygen and 100 cm³ of carbon dioxide remained.

Deduce the formula of the hydrocarbon and the equation for its combustion. All volumes were measured at r.t.p..

[3]

Cracking is used to obtain short-chain alkanes, alkenes and hydrogen from long-chain alkanes.

i) Give a use for each of the three products listed above.

 short-chain alkanes ..................................................

 alkenes ..................................................

 hydrogen ..................................................

[3]

 ii) Write an equation for the cracking of decane, C₁₀H₂₂, which produces two different alkenes and hydrogen as the only products.

[1]

Separate: Chemistry and Extended Only

Chlorine reacts with propane in a substitution reaction to form 1-chloropropane.

 CH₃–CH₂–CH₃ + Cl₂ → CH₃–CH₂–CH₂–Cl + HCl

i) What is the essential condition for the above reaction?

[1]

 ii) There is more than one possible substitution reaction between chlorine and propane.

Suggest the structural formula of a different product.

[1]

In a titration, a student added 25.0 cm³ of 0.200 mol/dm³ aqueous sodium hydroxide to a conical flask. The student then added a few drops of methyl orange to the solution in the conical flask. Dilute sulfuric acid is then added from a burette to the conical flask. The volume of dilute sulfuric acid needed to neutralise the aqueous sodium hydroxide was 20.0 cm³ The reaction is shown by the equation.

2NaOH + H₂SO₄ → Na₂SO₄ + 2H₂O

State the colour of methyl orange in aqueous sodium hydroxide.

Separate: Chemistry and Extended Only

Determine the concentration of the dilute sulfuric acid in g/dm³ using the following steps.

• Calculate the number of moles of aqueous sodium hydroxide added to the conical flask.

..................... mol

• Calculate the number of moles of dilute sulfuric acid added from the burette.

...................... mol

• Calculate the concentration of the dilute sulfuric acid in mol/dm³.

............ mol/dm³

• Calculate the concentration of the dilute sulfuric acid in g/dm³.

............... g/dm³

Separate: Chemistry and Extended Only

Ethanoic acid is a weak acid and hydrochloric acid is a strong acid. Both ethanoic acid and hydrochloric acid dissociate in aqueous solution.

Hydrochloric acid produces salts called chlorides. Magnesium carbonate reacts with hydrochloric acid to produce magnesium chloride.

MgCO₃ + 2HCl→ MgCl₂ + H₂O + CO₂

A student used 50.00 cm³ of 2.00 mol/dm³ hydrochloric acid in an experiment to produce magnesium chloride.

Calculate the mass, in g, of magnesium carbonate needed to react exactly with 50.00 cm³ of 2.00 mol/dm³ hydrochloric acid using the following steps.

• Calculate the number of moles of HCl present in 50.00 cm³ of 2.00 mol/dm³ HCl.

.............................. mol

• Determine the number of moles of MgCO₃ which would react with 50.00 cm³ of 2.00 mol/dm³ HCl.

.............................. mol

• Calculate the relative formula mass, M_r, of MgCO₃.

M_r of MgCO₃ = ..............................

• Calculate the mass of MgCO₃ needed to react exactly with 50.00 cm³ of 2.00 mol/dm³ HCl.

mass = .............................. g

Extended Only

When copper(II) oxide is heated at 800 °C it undergoes the reaction shown by the equation.

4CuO → 2Cu₂O + O₂

Calculate the volume of oxygen, measured at r.t.p., which is formed when 1.60 g of CuO reacts as shown in the equation.

Separate: Chemistry and Extended Only

Oxygen is produced by the decomposition of hydrogen peroxide. Manganese(IV) oxide is the catalyst for this reaction. The equation for the decomposition of hydrogen peroxide is shown.

2H₂O₂ (aq) → 2H₂O (l) + O₂ (g)

25.0 cm³ of aqueous hydrogen peroxide forms 48.0 cm³ of oxygen at room temperature and pressure (r.t.p.). Calculate the concentration of aqueous hydrogen peroxide at the start of the experiment using the following steps.

• Calculate the number of moles of oxygen formed.

• Deduce the number of moles of hydrogen peroxide that decomposed.

• Calculate the concentration of hydrogen peroxide in mol/dm³.

Extended Only

This question is about elements X, Y and Z.

What is the name of the amount of any substance that contains 6.02 × 10²³ particles?

Extended Only

The constant 6.02 × 10²³ has a name. What is the name of this constant?

Separate: Chemistry and Extended Only

Sulfur dioxide, SO₂, is used in the manufacture of sulfuric acid.

The next stage of the process is a reaction which can reach equilibrium. The equation for this stage is shown.

2SO₂ (g) + O₂ (g)  2SO₃ (g)

Calculate the percentage by mass of sulfur in sulfur trioxide, SO₃.

percentage = ..............................%

Separate: Chemistry and Extended Only

This question is about reactions of bases and acids.

A student wanted to find the concentration of some dilute sulfuric acid by titration. The student found that 25.0 cm³ of 0.0400 mol/dm³ NaOH (aq) reacted exactly with 20.0 cm³ of H₂SO₄ (aq). Calculate the concentration of the H₂SO₄(aq) in mol/dm³ using the following steps.

• Calculate the number of moles of NaOH in 25.0 cm³.

moles = ..............................

• Deduce the number of moles of H₂SO₄ that reacted with the 25.0 cm³ of NaOH (aq).

moles = ..............................

• Calculate the concentration of H₂SO₄ (aq) in mol/dm³.

concentration = .............................. mol/dm³

Separate: Chemistry and Extended Only

A student wanted to find the concentration of some dilute sulfuric acid by titration. The student found that 25.0 cm³ of 0.0400 mol/dm³ NaOH (aq) reacted exactly with 20.0 cm³ of H₂SO₄ (aq).

Calculate the concentration of the 0.0400 mol/dm³ NaOH (aq) in g/dm³.

concentration = .............................. g/dm³

Separate: Chemistry and Extended Only

Ammonia, NH₃, is used to produce nitric acid, HNO₃. This happens in a three-stage process. Stage 1 is a redox reaction.

4NH₃ + 5O₂ → 4NO + 6H₂O

In this reaction the predicted yield of NO is 512 g. The actual yield is 384 g. Calculate the percentage yield of NO in this reaction.

percentage yield of NO = ................................ %

Extended Only

The equation for the reaction in stage 3 is shown.

4NO₂ + 2H₂O + O₂ → 4HNO₃

Calculate the volume of O₂ gas, at room temperature and pressure (r.t.p.), needed to produce 1260 g of HNO₃.

Use the following steps.

• Calculate the number of moles of HNO₃.

moles of HNO₃ = ..............................

• Deduce the number of moles of O₂ that reacted.

moles of O₂ = ..............................

• Calculate the volume of O₂ gas that reacts at room temperature and pressure (r.t.p.).

volume of O₂ gas = .............................. dm³

Separate: Chemistry and Extended Only

The following method is used to make crystals of hydrated nickel sulphate. An excess of nickel carbonate, 12.0 g, was added to 40 cm³ of sulphuric acid, 2.0 mol/dm³. The unreacted nickel carbonate was filtered off and the filtrate evaporated to obtain the crystals.

NiCO₃ + H₂SO₄  → NiSO₄ + CO₂ + H₂O

NiSO₄ + 7H₂O  → NiSO₄.7H₂O

Mass of one mole of NiSO₄.7H₂O = 281 g Mass of one mole of NiCO₃ = 119 g

i) Calculate the mass of unreacted nickel carbonate.

Number of moles of H₂SO₄ in 40 cm³ of 2.0 mol/dm³ acid = 0.08 Number of moles of NiCO₃ reacted = ...................................... 
Mass of nickel carbonate reacted = ...................................... g Mass of unreacted nickel carbonate =  ...................................... g

[3]

ii) The experiment produced 10.4 g of hydrated nickel sulphate. Calculate the percentage yield.

The maximum number of moles of NiSO₄.7H₂Othat could be formed = ......................................
The maximum mass of NiSO₄.7H₂O that could be formed = ...................................... g
The percentage yield = ...................................... %

[3]

Extended Only

In the above method, a soluble salt was prepared by neutralising an acid with an insoluble base. Other salts have to be made by different methods.

i) Give a brief description of how the soluble salt, rubidium sulphate could be made from the soluble base, rubidium hydroxide.

[3]

ii) Suggest a method of making the insoluble salt, calcium fluoride.

[3]

Soluble salts can be made using a base and an acid.

Complete this method of preparing dry crystals of the soluble salt cobalt(II) chloride-6-water from the insoluble base cobalt(II) carbonate.

   step 1    Add an excess of cobalt(II) carbonate to hot dilute hydrochloric acid.   step 2

   step 3

   step 4

i)

5.95 g of cobalt(II) carbonate were added to 40 cm³ of hydrochloric acid, concentration 2.0 mol / dm³. Calculate the maximum yield of cobalt(II) chloride-6-water and show that the cobalt(II) carbonate was in excess.

CoCO₃ + 2HCl → CoCl₂ + CO₂ + H₂O CoCl₂ + 6H₂O → CoCl₂.6H₂O

maximum yield:

number of moles of HCl used = .........................................

number of moles of CoCl₂ formed = ................................

number of moles of CoCl₂.6H₂O formed = .......................................................

mass of one mole of CoCl₂.6H₂O = 238 g

maximum yield of CoCl₂.6H₂O = ...................................................................g

to show that cobalt(II) carbonate is in excess:

number of moles of HCl used = ..................................... (use your value from above)

mass of one mole of CoCO₃ = 119 g

number of moles of CoCO₃ in 5.95 g of cobalt(II) carbonate = ..............................

[5]

ii) Explain how these calculations show that cobalt(II) carbonate is in excess.

[1]

Crystals of sodium sulphate-10-water, Na₂SO₄.10H₂O, are prepared by titration.

25.0 cm³ of aqueous sodium hydroxide is pipetted into a conical flask. A few drops of an indicator are added. Using a burette, dilute sulphuric acid is slowly added until the indicator just changes colour. The volume of acid needed to neutralise the alkali is noted.

Suggest how you would continue the experiment to obtain pure, dry crystals of sodium sulphate-10-water.

Separate: Chemistry and Extended Only

Using 25.0 cm³ of aqueous sodium hydroxide, 2.24 mol / dm³, 3.86 g of crystals were obtained. Calculate the percentage yield.

2NaOH + H₂SO₄ → Na₂SO₄ + 2H₂O

Na₂SO₄ + 10H₂O → Na₂SO₄.10H₂O

Number of moles of NaOH used = Maximum number of moles of Na₂SO₄.10H₂O that could be formed = Mass of one mole of Na₂SO₄.10H₂O = 322g Maximum yield of sodium sulphate-10-water, in g = Percentage yield = 

Calcium and other minerals are essential for healthy teeth and bones. Tablets can be taken to provide these minerals.

Healthy Bones

Each tablet contains calcium magnesium zinc copper boron

Boron is a non-metal with a macromolecular structure.

i) Predict two physical properties of boron.

[2]

ii) Name another element and a compound that have macromolecular structures.

[2]

iii) Sketch the structure of one of the above macromolecular substances.

[2]

Describe the reactions, if any, of zinc and copper(II) ions with an excess of aqueous sodium hydroxide.

i) Zinc ions Addition of aqueous sodium hydroxide: Excess sodium hydroxide:

[2]

ii) Copper(II) ions

Addition of aqueous sodium hydroxide: Excess sodium hydroxide:

[2]

Extended Only

Each tablet contains the same number of moles of CaCO₃ and MgCO₃. One tablet reacted with excess hydrochloric acid to produce 0.24 dm³ of carbon dioxide at r.t.p.

CaCO₃ + 2HCl → CaCl₂ + CO₂ + H₂O MgCO₃ + 2HCl → MgCl₂ + CO₂ + H₂O

i) Calculate how many moles of CaCO₃ there are in one tablet. Number of moles CO₂ = Number of moles of CaCO₃ and MgCO₃ = Number of moles of CaCO₃ =

[3]

ii) Calculate the volume of hydrochloric acid, 1.0 mol / dm³, needed to react with one tablet. Number of moles of CaCO₃ and MgCO₃ in one tablet =  Use your answer to (c)(i). Number of moles of HCl needed to react with one tablet =

Volume of hydrochloric acid, 1.0 mol / dm³, needed to react with one tablet =

[2]

The soluble salt hydrated lithium sulfate is made by titration from the soluble base lithium hydroxide.

The sulfuric acid is added slowly from the burette until the indicator just changes colour. The volume of sulfuric acid  needed to just neutralise the lithium hydroxide is noted.

Describe how you would continue the experiment to obtain pure dry crystals of hydrated lithium sulfate.

Separate: Chemistry and Extended Only

Using 25.0 cm³ of aqueous lithium hydroxide, concentration 2.48 mol / dm³, 2.20 g of hydrated lithium sulfate was obtained.

Calculate the percentage yield, giving your answer to one decimal place.

2LiOH + H₂SO₄ → Li₂SO₄ + 2H₂O Li₂SO₄ + H₂O → Li₂SO₄.H₂O

 Number of moles of LiOH used = .......................    

Number of moles of Li₂SO₄.H₂O which could be formed = .......................    

Mass of one mole of Li₂SO₄.H₂O = 128g    

Maximum yield of Li₂SO₄.H₂O = ....................... g    

Percentage yield = .......................%

Separate: Chemistry and Extended Only

An experiment was carried out to show that the formula of the hydrated salt is Li₂SO₄.H₂O. A sample of the hydrated salt was weighed and its mass recorded. It was then heated and the anhydrous salt was weighed. This procedure was repeated until two consecutive masses were the same. This procedure is called ‘heating to constant mass’.

i) What is the reason for heating to constant mass?

[1]

ii) The mass of the hydrated salt is m₁ and the mass of the anhydrous salt is m₂. Explain how you could show that the hydrated salt has one mole of water of crystallisation per mole of the anhydrous salt.

[3]

Separate: Chemistry and Extended Only

The elements in Period 3 and some of their common oxidation states are shown below.

i) Why do the oxidation states increase from sodium to silicon?

[1]

ii) After Group(IV) the oxidation states are negative and decrease across the period. Explain why.

[2]

Aluminium sulfide contains two elements. Predict its formula.

Separate: Chemistry and Extended Only

Choose a different element from Period 3 that matches each description.

i) It has a similar structure to diamond.

[1]

ii) It reacts violently with cold water to form a solution pH = 14.

[1]

iii) It has a gaseous oxide of the type XO₂, which is acidic.

[1]

Extended Only

Draw a diagram that shows the arrangement of the outer electrons in the ionic compound sodium phosphide. Use o to represent an electron from sodium. Use x to represent an electron from phosphorus.

Extended Only

Sodium reacts with sulphur to form sodium sulfide.

2Na + S → Na₂S

An 11.5 g sample of sodium is reacted with 10 g of sulfur. All of the sodium reacted but there was an excess of sulfur.

Calculate the mass of sulfur left unreacted.

i) Number of moles of sodium atoms reacted =  [2 moles of Na react with 1 mole of S]

[1]

ii) Number of moles of sulfur atoms that reacted =

[1]

iii) Mass of sulfur reacted, in grams =

[1]

iv) Mass of sulfur left unreacted, in grams =

[1]

Separate: Chemistry and Extended Only

Sulfuric acid is made by the Contact process.

2SO₂ + O2 2SO₃

This is carried out in the presence of a catalyst at 450°C and 2 atmospheres pressure.

i) Name the catalyst used.

[1]

ii) If the temperature is decreased to 300°C, the yield of sulfur trioxide increases. Explain why this lower temperature is not used.

[1]

Separate: Chemistry and Extended Only

Sulfuric acid was first made in the Middle East by heating the mineral, green vitriol, FeSO₄.7H₂O. The gases formed were cooled. FeSO₄.7H₂O (s) → FeSO₄ (s) + 7H₂O (g) green crystals      yellow powder 2FeSO₄ (s) → Fe₂O₃ (s) + SO₂ (g) + SO₃ (g) On cooling SO₃ + H₂O → H₂SO₄ sulfuric acid SO₂ + H₂O → H₂SO₃ sulfurous acid

i) How could you show that the first reaction is reversible?

[2]

ii) Sulfurous acid is a reductant (reducing agent). What would you see when acidified potassium manganate(VII) is added to a solution containing this acid?

[2]

iii) Suggest an explanation why sulfurous acid in contact with air changes into sulfuric acid.

[1]

Extended Only

9.12 g of anhydrous iron(II) sulfate was heated. Calculate the mass of iron(III) oxide formed and the volume of sulfur trioxide, at r.t.p., formed. 

2FeSO₄ (s) → Fe₂O₃ (s) + SO₂ (g) + SO₃ (g)

mass of one mole of FeSO₄ = 152 g number of moles of FeSO₄ used = _____________ number of moles of Fe₂O₃ formed = _____________  mass of one mole of Fe₂O₃ = _____________ g mass of iron(III) oxide formed = _____________ g number of moles of SO₃ formed = _____________  volume of sulfur trioxide formed = _____________ dm³

Separate: Chemistry and Extended Only

Sulfur dioxide, SO₂, and sulfur trioxide, SO₃, are the two oxides of sulfur.

Sulfur trioxide can be made from sulfur dioxide.

i) Why is this reaction important industrially?

[1]

ii) Complete the word equation.

sulfur dioxide + .............................................. → sulfur trioxide

[1]

iii) What are the conditions for this reaction?

[2]

Separate: Chemistry and Extended Only

Sulphur dioxide is easily oxidised in the presence of water.

SO₂ + 2H₂O – 2e^– → SO₄^2– + 4H⁺

i) What colour change would be observed when an excess of aqueous sulphur dioxide is added to an acidic solution of potassium manganate(VII)?

[2]

ii) To aqueous sulphur dioxide, acidified barium chloride solution is added. The mixture remains clear. When bromine is  added, a thick white precipitate forms. What is the white precipitate? Explain why it forms.

[3]

Extended Only

Sulphur dioxide reacts with chlorine in an addition reaction to form sulphuryl chloride.

SO₂ + Cl₂ → SO₂Cl₂

8.0 g of sulphur dioxide was mixed with 14.2 g of chlorine. The mass of one mole of SO₂Cl₂ is 135 g. Calculate the mass of sulphuryl chloride formed by this mixture.

Calculate the number of moles of SO₂ in the mixture = ..................

Calculate the number of moles of Cl₂ in the mixture = ..................

Which reagent was not in excess? ...............................

How many moles of SO₂Cl₂ were formed = ...................

Calculate the mass of sulphuryl chloride formed = ............. g

Until recently, arsenic poisoning, either deliberate or accidental, has been a frequent cause of death. The symptoms of arsenic poisoning are identical with those of a common illness, cholera. A reliable test was needed to prove the presence of arsenic in a body.

In 1840, Marsh devised a reliable test for arsenic

Hydrogen is formed in this reaction. Any arsenic compound reacts with this hydrogen to form arsine which is arsenic hydride, AsH₃.

i) The mixture of hydrogen and arsine is burnt at the jet and arsenic forms as a black stain on the glass. Write an equation for the reaction which forms hydrogen.

 [2]

ii) Draw a diagram which shows the arrangement of the outer electrons in one molecule of the covalent compound arsine. The electron distribution of arsenic is 2 + 8 + 18 + 5. Use x to represent an electron from an arsenic atom. Use o to represent an electron from a hydrogen atom.

[2]

Separate: Chemistry and Extended Only

Another hydride of arsenic has the composition below.

arsenic 97.4% hydrogen 2.6%

i) Calculate the empirical formula of this hydride from the above data. Show your working.

[2]

ii) The mass of one mole of this hydride is 154 g. What is its molecular formula?

[1]

iii) Deduce the structural formula of this hydride.

[1]

Separate: Chemistry and Extended Only

In the 19th Century, a bright green pigment, copper(II) arsenate(V) was used to kill rats and insects. In damp conditions, micro-organisms can act on this compound to produce the very poisonous gas, arsine.

i) Suggest a reason why it is necessary to include the oxidation states in the name of the compound.

[1]

ii) The formula for the arsenate(V) ion is AsO₄^3–.

Complete the ionic equation for the formation of copper(II) arsenate(V).

......Cu²⁺ + ......AsO₄^3– → ..................................

[2]