Inorganic Chemistry Practicals (A Level only) (AQA A Level Chemistry)

The most common oxidation states of iron in iron compounds are +2 and +3. Iron(II)ions are readily oxidised to iron(III). 

Write full electronic configurations for Fe²⁺ and Fe³⁺ ions and explain their relative stabilities.

Iron(II) ions may be oxidised by bubbling chlorine gas through an aqueous solution of the ions.

Write an ionic equation to represent the reaction and state the colour change which would occur.

Describe, including equations, how it is possible to distinguish between Fe²⁺ and Fe³⁺ ions using aqueous sodium hydroxide solution.

The concentration of iron in iron supplement tablets can be determined using colorimetry of a complex of iron with thiocyanate ions;

Fe³⁺ + SCN^- → [FeSCN]²⁺

Outline a method which could be used.

Water pipes can release lead, copper and zinc ions as a result of containing acidic drinking water.  Describe how you could detect aqueous copper(II) ions in a solution using aqueous sodium hydroxide solution, including an equation in your answer.

The concentration of copper(II) ions in solution can be determined by titration with a standard solution of EDTA.

The equipment used for this experiment is shown in Figure 1 below.

Figure 1

Explain why the complex formed between copper(II) ions and EDTA is more stable than the complex in the aqueous solution.

Aqueous copper(II) ions will react with aqueous sodium carbonate to form a precipitate. 
Give one observation associated with this reaction and write a balanced symbol equation to represent the reaction.

The percentage of iron in a steel wire can be determined from a redox titration using acidified potassium manganate(VII) solution. 

Explain why the iron wire is dissolved in an excess of sulfuric acid and not ethanoic acid.

1.60 g of iron wire was dissolved in an excess of dilute sulfuric acid and the solution made up to 250 cm³. A 25.0 cm³ portion of this solution required 26.0 cm³ of 0.02 mol dm^–3 potassium manganate(VII) solution for complete reaction.

Calculate the percentage of iron in the steel wire.

If an excess of sodium hydroxide is added to the flask after the titration is completed it will react with both manganese(II) and iron(III) ions.   Deduce what would be observed and write an equation, including state symbols for the reaction which would be taking place when the sodium hydroxide reacts with the iron(III) ions.

Malachite has the formula CuCO₃.Cu(OH)₂ and is readily converted to copper (II) oxide by heating.  The oxide can be reduced to copper by heating strongly with carbon.

Malachite dissolves in nitric acid to form a solution of copper(II) nitrate.  When excess aqueous ammonia is added dropwise to the solution colour changes occur.

State and explain the observations of adding the ammonia solution until in excess and write equations for the reactions that occur.

Redox potentials are used to predict chemical reactions.  Use the standard electrode potentials below to write an equation for the reaction which occurs when bromine is added to aqueous iron(II) ions and deduce the EMF.

Fe²⁺(aq) + 2e^- ⇌Fe(s)             E^ϴ = -0.44 V

Fe³⁺(aq) + e^- ⇌ Fe²⁺(aq)         E^ϴ = +0.77 V

Br₂(aq) + 2e^- ⇌ 2Br^-(aq)          E^ϴ = +1.09 V

Ethanedioic acid,H₂C₂O₄, is used commercially as a laundry rinse, wood-bleaching agent and limescale remover.  The ethanedioate ion, C₂O₄^2-, is a bidentate ligand and can participate in ligand substitution reactions.

Fe³⁺(aq) + e^- ⇌ Fe²⁺(aq)         E^ϴ = +0.77 V

2CO₂(g) + 2e^- ⇌  C₂O₄^2-(aq)   E^ϴ = -0.49 V

A sample of phosphorus(V) oxide was prepared in the laboratory using white phosphorus (P₄) which is a hazardous form of the element.

Give one reason why white phosphorus is stored under water.

Give an equation and state one observation for the complete reaction of phosphorus with air.

Phosphorus(V) oxide has a melting point of 300 ^oC.  Outline a suitable method to determine the melting point of the sample produced and its link to the purity of the sample.

Phosphorus(V) oxide is an acidic oxide. 

Write two equations to demonstrate the amphoteric nature of aluminium oxide.

Outline the mechanism to show how an aqueous Al³⁺ ion can act as a Brønsted-Lowry acid.

Outline an experiment to determine the pH of the period three oxides. Include expected results and equations in your answer.

Explain how it is possible to establish the purity of a sample of P₄O₁₀.

Write an equation for the reaction between ethanedioic acid, H₂C₂O₄, and sodium hydroxide, NaOH.

15.00 cm³ of a H₂C₂O₄ requires 10.30 cm³ of a 0.25 mol dm^-3 solution of NaOH for complete neutralisation using a phenolphthalein indicator for the first permanent colour change.

15.00 cm³ of the same H₂C₂O₄ solution required 12.35 cm³ of potassium permanganate, KMnO₄, solution for complete oxidation to carbon dioxide and water in the presence of dilute sulfuric acid to further acidify the H₂C₂O₄ solution for the first permanent colour change.

Calculate the moles of H₂C₂O₄ in the solution and therefore the concentration of the solution. 

Write the full redox equation, including state symbols, for this redox titration.

Calculate the concentration of the potassium permanganate, KMnO₄, solution. 

Vanadium is a transition metal with oxidation states ranging from +5 to +2, which can each be identified by the colour of the solution that they form.

State the formula of ammonium vanadate(V).

A demonstration of the oxidation states of vanadium was carried out in a fume hood against a white background to help observe the colour changes.

Ammonium vanadate(V) is not very soluble in water. A quarter spatula of ammonium vanadate was mixed with 1.0 mol dm^-3 hydrochloric acid forming a yellow solution containing the dioxovanadium(V) ion, VO₂⁺.

An excess of zinc powder was added and the test tube was again shaken. Over a period of fifteen minutes the colours changed from yellow to green to blue to green and finally violet.

Using the data in Table 1, explain the first reaction that takes place after the addition of the excess zinc powder.

Your answer should include oxidation states and relevant equations.

Table 1

During the demonstration described in part (b), a colour change from yellow to green is observed. Explain this colour change. 

The experiment outlined in part (b) was repeated and zinc was replaced with tin.

Use the redox potentials in Table 1, part (b), to explain the extent to which tin can reduce the dioxovanadium(V) ion, VO₂⁺.

Your answer should include any expected colour changes.

A student carried out a series of experiments to identify four different transition metal compounds. Table 1 shows the student’s results which include incorrect and incomplete observations.

Identify the transition metal ions in each compound and correct the student’s observations.

Table 1

Write an equation, including state symbols, for the separate reactions between the aqueous transition metal complex of Compound C with the following reagents:

Write an equation for the reaction between the aqueous transition metal complex of Compound B with sodium carbonate solution to produce the white precipitate observed.

Explain why compound A releases a gas when reacted with sodium carbonate solution. 

A sample of rusted iron was analysed to determine the percentage of iron that had been oxidised to hydrated iron(III) oxide. A small amount of this sample was dissolved in excess sulfuric acid to give 500 cm³ of solution. This solution contains both Fe²⁺ and Fe³⁺ ions.

20.00 cm³ of this solution required 18.05 cm³ of 0.05 mol dm^-3 potassium permanganate for complete oxidation.

An oxidising agent was also added to a second 20.00 cm³ of the rust solution to convert
all of the Fe²⁺ ions to Fe³⁺ ions. The Fe³⁺ ions reacted with 18.40 cm³ of 0.05 mol dm^-3 of EDTA^4- ions.

Write the full equation for the reaction between iron(II) ions and manganate(VII) ions

Calculate the number of moles of Fe²⁺ ions present in the sample titrated against potassium permanganate.

Assuming 1 mole of EDTA reacts with 1 mole of Fe³⁺ ions, calculate the moles of Fe³⁺ ions present in the second solution. 

Using your answers to part (a) and part (b) calculate the percentage of rust in the original sample of rusted iron. 

Using the redox potentials given in Table 1, explain why hydrochloric acid is not a suitable choice for this titration. Your answer should include relevant equations.

Table 1