Redox Processes (DP IB Chemistry: SL)

Common household bleach is a cleaning product which 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(I) 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 oxidize hydrogen peroxide to oxygen gas. The reaction is carried out with both species under acidic conditions. 

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

Metals can often be seen written as a list, from the most reactive metal to the least reactive metal. This list is known as the reactivity series of metals and can be used to predict the feasibility of a reaction. 

Below is a section of the reactivity series of metals, ordered from most to least reactive: 

            Calcium

            Magnesium

            Aluminium

            Zinc

            Iron

            Tin

            Lead 

A piece of zinc was placed into a solution of iron(II) sulfate and a solution of magnesium sulfate.

Predict, giving a reason, whether a reaction would occur in each solution. 

Copper is below lead on the reactivity series shown in part (a). A piece of zinc was placed into a solution of copper(II) sulfate. Write the half equation for the zinc and identify the type of reaction taking place.

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

The following reaction is an example of a common redox reaction:

      5Fe²⁺ (aq) + MnO₄^- (aq) +8H⁺ (aq) → 5Fe³⁺ (aq) + Mn²⁺ (aq) + 4H₂O (l) 

Deduce the oxidation numbers of iron and manganese in the above reaction, both as reactants and as products.           

State which substance is reduced.

The amount of iron in some dietary iron supplements was analyzed by redox titration. Four tablets were crushed and dissolved in 50.0 cm³ of 2.00 mol dm^-3 sulfuric acid. The solution was then transferred to a 250 cm³ volumetric flask and made up to 250 cm³ with distilled water. 

A 25.0 cm³ sample of the iron tablets solution was titrated against 0.00500 mol dm^-3 potassium manganate(VII) and 25.8 cm³ was needed for complete reaction. 

Determine the amount of iron, in mol, in one tablet.

Halide ions, such as chloride, Cl^-, 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 other halogens 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 oxidizes sulfur dioxide (SO₂) in aqueous solutions to sulfate ions (SO₄^2-) under acidic conditions.

Use the two half-equations from part (c) to construct the overall redox equation for this reaction.  

The iron of 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) 

Define the term reduction.

State, with a reason, the oxidizing agent in this reaction in part (a).

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.

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

                      Table 1 

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 oxide, Al₂O₃. Electrolysis is then carried out on the molten Al₂O₃, to extract the aluminium.

Another ionic compound which can undergo electrolysis is molten lead bromide.

State the products formed at each electrode during the electrolysis of molten lead bromide, giving the equations at each electrode with state symbols.  

Draw a labelled diagram of the apparatus suitable to carry out the electrolysis of molten lead bromide. Include the direction of electron flow, the negative electrode (cathode), the positive electrode (anode) and the electrolyte.

The list below shows three metals from the activity series in order of reactivity.

Deduce which of the three metals is the strongest reducing agent.

A voltaic cell can be made by joining two half-cells together, such as  Zn/Zn²⁺ and Ni/Ni²⁺.

Write a balanced equation for the overall reaction taking place when the two half-cells are connected together, and state which species is undergoing oxidation.

Cell diagrams are a way to represent the redox reactions taking place in voltaic cells.

Write a cell diagram for the overall cell reaction taking place in part (b).

Complete the partially labelled diagram in Figure 1, of the apparatus used in the voltaic cell in part (b). Show the direction of the movement of the electrons and ions in the cell.                                                      

Figure 1

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 0.0100 mol dm^-3 potassium manganate(VII) solution. The average titre was found to be 26.65 cm³ of potassium manganate(VII) solution. 

Iron sulfate reacts with chromium to produce chromium(III) sulfate, Cr₂(SO₄)₃ , and iron

Deduce the overall ionic equation for the reaction occurring 

Molten potassium bromide can be electrolysed using graphite electrodes.

State the half equations for the oxidation and reduction processes and deduce the overall cell reaction, including state symbols. 

Explain why solid potassium bromide does not conduct electricity. 

A voltaic cell is made from a half-cell containing a magnesium electrode in a solution of magnesium nitrate and a half-cell containing a silver electrode in a solution of silver(I) nitrate.

State the oxidation state of phosphorus in the following compounds.

The tetrathionate ion is shown below:

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

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

A 150.0 cm³ sample of pond water was analysed using the Winkler method to determine its biological oxygen demand (BOD). Initially it took 29.40 cm³ of 0.010 mol dm^-3 Na₂S₂O₃ to react with iodine.

After five days it required 13.70 cm³ of 0.010 mol dm^-3 Na₂S₂O₃ to react with iodine.

The unbalanced equations for the Winker method are shown below.

......Mn²⁺ (aq) + ......OH^- (aq) + O₂ (aq) → ......MnO₂ (s) + ......H₂O (l) 

MnO₂ (s) + ......I^- (aq) + ......H⁺ (aq) → Mn²⁺ (aq) + I₂ (aq) + ......H₂O (l) 

......S₂O₃^2- (aq) + I₂ (aq) → S₄O₆^2- + ......I^- (aq)  

Balance the equations for the Winkler method. 

Deduce the reducing agent in the reaction between S₂O₃^2- and I₂. Justify your answer. 

Use the information in part a) and section 6 in your data booklet to determine the initial concentration, in ppm, of oxygen. 

Use the information in part a) and section 6 in the data booklet to determine the concentration, in g dm^-3, of oxygen after five days. 

Determine the BOD of the pond water in ppm. 

15.00 cm³ of ethanedioic acid, H₂C₂O₄ (aq), requires 10.30 cm³ of a 0.250 mol dm^-3 solution of sodium hydroxide, NaOH (aq), 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 solution, KMnO₄ (aq), 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. 

H₂C₂O₄ (aq) + 2NaOH (aq) → Na₂C₂O₄ (aq)+ 2H₂O (l) 

[2]

Deduce the following half equations and overall redox equation for the reaction outlined in part a).

MnO₄^- (aq) to Mn²⁺ (aq) .........................

H₂C₂O₄ (aq) to CO₂ (g) ...........................

Overall equation ......................................

Calculate the concentration, in mol dm^-3, of the potassium manganate(VII), KMnO₄, solution.