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.

i) Write a balanced equation with state symbols for this reaction.

ii) Deduce the oxidation state of chlorine in all of the chlorine-containing reactants and products.

iii) State what specific type of redox reaction this reaction is. Explain your answer.

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)

i) What are the oxidation states of the nitrogen atom in NH₃ and in N₂H₄?

ii) What is the oxidising agent in this reaction? Explain your answer.

iii) Explain why the risks of producing chlorine are greater than the risks of producing hydrazine.

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. 

i) Identify the oxidising and reducing agents in this reaction.

ii) Write the half-equation for the oxidation of hydrogen peroxide to oxygen gas.

iii) The manganate (VII) ions themselves get reduced to manganese (II) ions. Write down the half-equation for the reduction of manganate (VII) ions.

iv) Deduce the overall redox reaction for this reaction. 

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)

i) Deduce the oxidation state of silver in AgNO₃ and AgX, and deduce the oxidation state of the halide in X^- and in AgX.

ii) Explain whether the precipitation of silver halides is a redox reaction.

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)

i) State what type of redox reaction this is.

ii) Using the overall redox reaction above, deduce the half-equation for chlorine. State whether chlorine is oxidised or reduced.

iii) Using the overall redox reaction above, deduce the half-equation for bromine. State whether bromine is oxidised or reduced.

iv) Use the reaction above and your knowledge of the halogens, to explain whether chlorine or bromine is a stronger oxidising agent. 

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

i) Deduce the half-equation for the reduction of chlorine in aqueous solution.

ii) Deduce the half-equation for the oxidation of sulfur dioxide in aqueous solution.

iii) Use the two half-equations in part (i) and (ii) to write the overall redox equation.

iv) Bromine reacts with sulfur dioxide in a different way to chlorine. Explain why.

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. 

i) Write the overall redox reaction of chlorine with water.

ii) Deduce the oxidation states of chlorine in all of the chlorine-containing compounds.

iii) The reaction between chlorine and water is a disproportionation reaction. Define the term disproportionation and use oxidation states to explain why the above reaction is an example of this.

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.

i) Write the overall equation for the thermite reaction, including state symbols.

ii) Determine the oxidation states of each of the metals in the reactants and products.

iii) Deduce the reduction and oxidation half-equations for this reaction.

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

i) Explain what type of reaction the thermite reaction is.

ii) What can be deduced about the chemical properties of the aluminium and iron metals involved in this 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)

i) Define the term reduction.

ii) What is the oxidising agent in this reaction? Explain your answer.

iii) A student investigates the rate of rusting of a piece of iron under different conditions.

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.

i) Sulfur can react in different ratios with oxygen to form sulfur dioxide or sulfur trioxide. Write equations to show the oxidation of sulfur to form sulfur dioxide and to form sulfur trioxide. Include state symbols in your equations. 

ii) Explain which reaction has the greatest increase in oxidation state.

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.

i) Write an ionic equation for the conversion of nitrogen to ammonia and hydrogen gas by nitrogen fixation.

ii) Explain whether the reaction in part (i) is reduction or oxidation.

iii) The ammonia produced by nitrogen fixation can be further converted to nitrite ions (NO₂^-). Nitrite ions in turn are oxidised to nitrate ions, which are mainly used by plants for protein synthesis.

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.

i) Explain the role of the oxidising agent in a redox reaction in terms of electron transfer.

ii) State the most common oxidation state of an oxygen atom when in a compound.

iii) Explain which compounds are an exception to your answer in part (ii).

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. 

i) Write down the overall equation for the extraction of aluminium from aluminium (III) oxide by electrolysis.

ii) State whether the aluminium (III) oxide is oxidised or reduced in the electrolysis reaction. Explain your answer.

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.

i) Write the overall equation for the reaction of iron (III) oxide with carbon.

ii) What type of redox reaction is the reaction of iron (III) oxide with carbon? Explain your answer.

iii) Use your answer to part (ii) to explain why aluminium cannot be extracted from aluminium (III) oxide by reacting it 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.

i) Write the reduction half-equation for this reaction. Include state symbols.

ii) Write the oxidation half-equation for this reaction. Include state symbols.

iii) Deduce the overall redox equation using your answers to parts (i) and (ii). Include state symbols.

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.

i) Write a balanced symbol equation for this reaction.

ii) Identify the oxidant in this reaction.

iii) Describe the expected observation to show that this reaction had gone to completion.

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

2 NaHSO₃ + S₂Cl₂ → Na₂S₄O₆ + 2 HCl

i) Deduce the oxidation state of sulfur in all of the sulfur-containing chemicals.

ii) Evaluate if this reaction is a disproportionation reaction.

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.

i) Deduce which student has identified the method with the highest atom economy.

Table 1 shows the relative molecular mass of the chemicals required for both methods.

Table 1

ii) Explain why neither student has given the correct justification for their choice of the highest atom economy.

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.

i) Use your knowledge of oxidation states, explain why a student might mistakenly think that the formula for a cyanide ion is CN⁺.

ii) Use your knowledge of oxidation states to identify the correct oxidation state of carbon in the ionic compound, KCN.

iii) Use your knowledge of chemical bonding to explain how the cyanide ion has the correct formula of CN^-.

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

i) Write the half equation involving copper (I) thiocyanate forming copper sulfate and hydrogen cyanide.

ii) The carbon in copper (I) thiocyanate, CuSCN, has the same oxidation state as the carbon in dichlorodifluoromethane, CCl₂F₂. Using your knowledge of oxidation states, state whether each component element in copper (I) thiocyanate, CuSCN, is oxidised or reduced.

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

i) Write the half equation involving potassium iodate forming potassium chloride and iodine.

ii) Using your knowledge of oxidation states, state whether each component element in potassium iodate is oxidised or reduced.

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.

i) Write a balanced symbol equation, including state symbols, for the reaction described.

ii) Explain, using oxidation states, if this is a redox, disproportionation reaction.

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)

i) Identify the reducing agent in this reaction.

ii) Identify the specific oxidising agent in this reaction.

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.

i) Write the balanced, ionic half equations to show the reduction of the manganate ion and the oxidation of the Fe²⁺ ion.

ii) Use your answers to part (i) to write an overall redox equation for the titration of iron (II) sulfate 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.