Oxidation, Reduction & Redox Equations (AQA A Level Chemistry)

Household bleach smells like chlorine gas and is therefore also called chlorine bleach. It contains a mixture of sodium chlorate (NaOCl), sodium chloride and water and can be made by dissolving chlorine gas in a solution of sodium hydroxide.

The mixing of household bleach with ammonia during cleaning should be avoided, as a redox reaction between the ammonia and the chlorate ions in bleach will generate toxic chlorine gas and hydrazine (N₂H₄).

The overall redox reaction for this reaction is shown below.

2NH₃ (aq) + 2ClO^- (aq) → N₂H₄ (aq) + Cl₂ (g) + 2OH^- (aq)

Due to the risks associated with chlorine-based bleach, alternative bleaches are often used instead. These bleaches are based on peroxides such as hydrogen peroxide. 

Manganate (VII) ions oxidise hydrogen peroxide to oxygen gas. The reaction is carried out with both species under acidic conditions. 

Explain how the oxidation state of the oxygen atom in H₂O₂ is different from its oxidation state in other compounds.

Halide ions can be identified using chemical tests. If an unknown compound is dissolved in dilute nitric acid, and then silver nitrate solution is added, a precipitate will form if the unknown solution contains halide ions. The precipitate formed will be a silver     halide.

The general equation for the precipitation reaction of halide ions with silver nitrate solution is:

AgNO₃ (aq) + X^- (aq) → AgX (s) + NO₃^- (aq)

Halide ions can also react with each other in aqueous solutions. Chlorine reacts in a redox reaction with an aqueous solution of sodium bromide, to form sodium chloride and bromine.

Cl₂ (aq) + NaBr (aq) → NaCl (aq) + Br₂ (aq)

Chlorine also oxidises sulfur dioxide (SO₂) in aqueous solutions to sulfate ions (SO₄^2-) under acidic conditions.

Chlorine is widely used in water purification. Besides killing dangerous germs like bacteria and viruses, chlorine also helps reduce unwanted tastes and odors in water by reacting with organic chemicals in water. 

Thermite is a mixture of finely powdered aluminium and iron oxide. 

In a thermite reaction, the aluminium reacts with iron oxide to produce hot molten iron. This process is often utilised in remote locations for welding railway lines.

Initially, a lot of heat is required to start off the reaction. However, once started, the reaction itself releases a significant amount of heat which is enough to melt the iron.

The thermite reaction in part (a) is a special type of redox reaction. 

Deduce the oxidation state of each of the stated elements in the ions and compounds to complete Table 1 below.

Table 1

The iron of the railway lines rusts when it comes into contact with water and oxygen. The overall redox equation for the rusting of iron is as follows:

4Fe(s) + 3O2(g) + 6H2O(g) → 4Fe(OH)3(s)

Figure 1 shows the set-up of the students’ experiment.

Figure 1

Predict in which test tube(s) the iron metal will not rust. Explain your answer.

Fossil fuels contain relatively large amounts of sulfur. When these fossil fuels are burnt, sulfur gets into the atmosphere and acts as a pollutant. 

The sulfur is oxidised in the air and rises very high into the atmosphere, where it mixes and reacts with water, oxygen and other chemicals to form more acidic pollutants known as acid rain.

Acid rain causes corrosion of buildings, statues and can damage aquatic and plant life.

Acid rain can also be caused by oxides of nitrogen when they dissolve in water.

Deduce the oxidation state of the nitrogen atom in the given ions and compounds to complete Table 1 below.

Table 1

Nitrogen is very unreactive, due to the strong triple bonds within the molecule, and will only react with oxygen in air to form nitrogen oxides under extreme conditions (such as lightning).

Nitrogen fixing bacteria in the soil convert the unreactive molecular nitrogen in the atmosphere into ammonia or other related nitrogen compounds. This process is called nitrogen fixation.

Write an equation for the conversion of nitrite ions to nitrate ions and deduce the oxidation states of nitrogen in the nitrogen-containing compounds.

Many chemical reactions are redox reactions as they involve the transfer of electrons.

Aluminium is present in the Earth’s crust in aluminium ore, called bauxite. A number of processes are done to this ore, to extract the aluminium from it. The bauxite is initially purified to produce aluminium (III) oxide. Electrolysis is then carried out, to extract the aluminium from the aluminium oxide. Oxygen gas is also formed as a byproduct of this part of the process. 

Other metals can be extracted from their metal ores by reacting the ore with carbon. Iron, for example, is extracted from iron (III) oxide (Fe₂O₃) in a huge container called a blast furnace, during a redox reaction with carbon.

Ferric chloride (FeCl₃) is used to treat sewage, industrial waste and to purify water. FeCl₃ can be formed from the reaction of iron with chlorine.

When ferric chloride is dissolved in water, it undergoes hydrolysis and gives off heat. The resulting acidic and corrosive solution is used to purify water and treat sewage.

The hydrolysis reaction of ferric chloride in water is as follows:

FeCl₃ (aq) + 3H₂O (l) → Fe(OH)₃ (aq) + 3HCl (aq)

State whether the hydrolysis of ferric chloride is a redox reaction. Explain your answer.

The tetrathionate ion is shown in Figure 1.

Figure 1

Calculate and explain the oxidation state of sulfur in the tetrathionate ion.

Sodium tetrathionate can be formed by reacting sodium thiosulfate, Na₂S₂O₃, with iodine.

Sodium tetrathionate can also be made by reacting sodium bisulphite, NaHSO₃, with disulfur dichloride, S₂Cl₂.

2 NaHSO₃ + S₂Cl₂ → Na₂S₄O₆ + 2 HCl

Two students are asked to comment on the atom economy for the formation of sodium tetrathionate using two different methods.

Method 1:        2 Na₂S₂O₃ + Cl₂ → Na₂S₄O₆ + 2 NaCl

Method 2:        2 NaHSO₃ + S₂Cl₂ → Na₂S₄O₆ + 2 HCl

Student 1 concludes that method 1 has the higher atom economy because the second reactant is smaller.

Student 2 concludes that method 2 has the higher atom economy because the waste product is smaller.

Cyanide-containing compounds are commonly used in organic chemistry to lengthen the carbon chain. They are also used in gold and silver mining because of their ability to dissolve these metals and their ores.

Copper (I) thiocyanate, CuSCN, can react with potassium iodate, KIO₃, in the presence of hydrochloric acid according to the following unbalanced equation.

CuSCN + KIO₃ + HCl → CuSO₄ + KCl + HCN + ICl + H₂O

Copper (I) thiocyanate can react with potassium iodate in the presence of hydrochloric acid according to the following unbalanced equation.

CuSCN + KIO₃ + HCl → CuSO₄ + KCl + HCN + ICl + H₂O

Use your answers from parts (b) and (c) to balance the symbol equation for the reaction of copper (I) thiocyanate with potassium iodate in the presence of hydrochloric acid.

__ CuSCN + __ KIO₃ + __ HCl → __ CuSO₄ + __ KCl + __ HCN + __ ICl + __ H₂O

Photochromic glass contains an evenly distributed mixture of copper (I) chloride and silver (I) chloride. When sunlight passes through the glass the silver (I) chloride is separated into its ions. The chloride ions are then converted into chlorine atoms and the
silver (I) ions into silver atoms. The silver atoms cluster together causing the lenses of photochromic glasses to darken.

Write equations for the processes involved in the darkening of photochromic glasses and explain if the reaction is reduction or oxidation.

Prescription photochromic sunglasses react to ultraviolet light and are adjusted to the needs of the wearer.

Explain why someone wearing photochromic sunglasses might not receive the full benefit of the glasses when driving a car on a cold, bright day.

The darkening process caused by the formation of silver atoms is reversible.

Write two balanced symbol equations to show how copper (I) chloride can react with the products from part (a) to remove the silver atoms and cause the lens to lighten.

Copper (I) chloride, similar to that used in photochromic glass, can be made in the laboratory by the following method:

1. Warm 0.5 g copper (II) oxide with 5 cm³ concentrated hydrochloric acid for 1 minute.
2. Add 1.0 g of copper turnings.
3. Gently boil for 5 minutes.
4. Filter the solution into 250 cm³ of deionised water.
5. Allow the copper (I) chloride precipitate to settle. 
6. Decant the liquid.
7. Allow the copper (I) chloride to dry.
   
   

The reaction of sodium iodide with concentrated sulfuric acid forms a variety of products as shown in Table 1.

Complete Table 1 by identifying which products are formed by oxidation or reduction.

Table 1      

Sulfur dioxide reacts with a solution of copper (II) chloride according to the following equation.

SO₂ (g) + 2 H₂O (l) + 2 CuCl₂ (aq) → H₂SO₄ (aq) + 2 HCl (aq) + 2 CuCl (s)

Sodium iodate is one of the main sources of iodine in the world. To extract the iodine, sodium iodate is reacted with sodium hydrogen sulfite, NaHSO₃, according to the following ionic equation.

2 IO₃ (aq) + 5 HSO₃^- (aq) → 3 HSO₄^- (aq) + 2 SO₄^2- (aq) + I₂ (aq) + H₂O (l)

Explain why sodium does not appear in the equation.

Explain the role of the hydrogen sulfite ions in the reaction shown in part (c).

A student sets up a titration to determine the amount of iron (II) sulfate in an iron tablet. They titrate the iron (II) sulfate solution with potassium manganate (VII) solution.

The iron (II) sulfate solution is acidified before titration to stop the manganate ion forming unwanted manganese dioxide.

Explain the effect that not acidifying the iron (II) sulfate would have on the final calculation of the estimated mass of iron.

The student dissolved the iron tablet in excess sulfuric acid and made the solution up to 250 cm³ in a volumetric flask.

25.0 cm³ of this solution was titrated with 0.01 mol dm^-3 potassium manganate (VII) solution.

The average titre was found to be 26.65 cm³ of potassium manganate (VII) solution.

Calculate the mass of iron, in mg, in the tablet. Give your answer to the appropriate number of significant figures.

A student wanted to repeat the same titration while reducing the percentage uncertainty of the 50 cm³ burette. 

Use your answer to part (c) to suggest a suitable titre value to reduce uncertainty in then burette. Calculate the concentration of the potassium manganate (VII) solution required. Give your answer in standard form.